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Prafessor of Analytical Chemistry in Purdue University 

Second Edition 

Revised and Enlarged 

Second Impression 




6 & 8 BOUVERIE ST., E. C. 


Copyright, 1914, 1919, by the 
McGraw-Hill Book Company, Inc. 



Since the publication of the first edition of this book numerous 
changes have been made in the standardized methods of analysis 
of certain industrial materials, as adopted by official committees 
of various scientific societies. In the present edition the dis- 
cussions and detailed procedure have been modified to conform to 
the revised methods, wherever possible. This statement applies 
to the analysis of coal, water, fertilizers, dairy products and 

A number of other analytical methods are now' described, 
of which may be mentioned the gravimetric determination 
of the chloride, sulphate and phosphate radicals and the in- 
direct determination of the halogens; the per chlorate method for 
potassium; the determination of chromium and vanadium in 
steel and the glyoxime method for nickel in steel; the analysis 
of brass and of soft bearing metals; and the volumetric determina- 
tion of zinc. 

A part of the discussion of metallography and treatment of 
steel has been rewritten and several new sections have been 
added to this portion of the book. New photomicrographs have 
been substituted for the old ones and about fifteen new figures 
have been used at various points throughout the book, the latter 
being the work of the author's students in Chemical Engineering, 
to whom grateful acknowledgment is due. Finally, a new system 
of chapter division has been introduced in order to give greater 
emphasis to the sections dealing with industrial analysis. 

In the process of making these changes a considerable poition 
of the book has been rewritten and the discussions have been 
amplified, in the interest of added clearness. It is hoped that 
these changes will serve to enhance the usefulness of the book as a 
college text and that the favor with which the first edition was so 
generously received will not be undeserved by the present 

December, 1918 


This book cannot be classed as a complete reference work on 
quantitative analysis, neither is it a bare outline of laboratory 
exercises. The author has felt a desire that has probably been 
felt by every teacher of quantitative analysis, to produce a book 
that would cover the ground that he wishes to cover in the college 
courses, providing a reasonable degree of latitude in the selection 
of exercises for other possible users of the book, and at the same 
time to present a theoretical and practical discussion of the sub- 
ject, sufficiently simple to be comprehended by the average stu- 
dent but not so elementary as to destroy his self-respect. 

One of the most difficult tasks connected with the teaching of 
quantitative analysis is to produce in the mind of the student a 
clear comprehension of the scientific development of quantitative 
methods. There seems to be a more or less unconscious tendency 
toward; the acceptance of the present laboratory method of pro- 
cedure as a gift of Providence. How well this situation has been 
met in the present volume must be shown by the test of experience. 
The general discussions have been given a large share of attention 
although elaborate or involved theoretical discussions have been, 
so far as possible, avoided. References to original papers have 
been carefully selected with a view to actual reading by the stu- 
dent and such references are, in nearly all cases, to discussions 
that will serve either to impress more clearly upon the reader's 
mind the principles mentioned in the text or to bring to his mind a 
realization of the labor involved in the development of the finished 
method. It is believed that careful and systematic reading and 
discussion of such original papers by the student with his in- 
structor is a most valuable aid in the understanding of quanti- 
tative analysis as a truly scientific study. 

The mathematical development of quantitative calculations in 
this book is somewhat unusual, in that the "rule of three" has 
been carefully excluded. It is the firm belief of the author, after 
several years of experience in teaching quantitative analysis, that 
the use of this rule of proportion has produced much harm and 
has been the greatest of all obstacles to the student in his attempt 
to grasp the principles of quantitative calculations, and par- 



ticularly those of volumetric analysis. The ideas involved in the 
solution of proportions are so labored and so unnecessary and 
require such cumbersome solutions of problems involving them 
that it is difficult to see why emphasis has so generally been placed 
upon this rule in chemical calculations. Other teachers will, no 
doubt, differ with the author upon this point. It is desired only 
that the method of presentation involved in these pages be tested, 
not in part but in whole, before final judgment is given. Most 
of the calculations involved in the laboratory exercises have been 
left to the student. Principles of such calculations are first 
fully explained but ready-made calculations that leave nothing 
to the ingenuity of the student furnish poor preparation for 
scientific analysts. 

Acknowledgment is here gladly expressed to Mr. H. C. Mahin 
for all of the original drawings in this book, also to Wm. Ainsworth 
and Sons, Bausch and Lomb Optical Company, the Bureau of 
Standards, Eimer and Amend and The Scientific Materials Com- 
pany for several cuts which they have loaned. 

Edward G. Mahin. 
Lafayette, Ind., 
December, 1913. 



Preface to the Second Edition v 

Preface to the First Edition vii 

Introduction xiii 


General Principles . \ 

Cleanliness and care — Limit of accuracy — Classes of methods. 

Gravimetric Analysis 5 

Factors — Factor weights. 

General operations: Preparation of samples — Solution — Precipita- 
tion — Colloids — Enlargement of particles — Filtration — Washing 
— Drying of precipitates — Ignition — Fusion — Weighing — Cali- 
bration of weights. 

Reagents — Glassware — Records. 


Experimental Gravimetric Analysis 76 

Calcium — Silver — Chlorides, bromides and iodides — Aluminium — 
Barium — Sulphates — Strontium — Potassium and sodium — Re- 
covery of platinum from waste — Magnesium — Phosphates — 
Manganese — Halogen compounds — Carbonates and carbon 

Electro-analysis 138 

Nature of electrolyte — Solvent — Temperature — Electrolytic pres- 
sure — Current density — Nature of electrodes — Other apparatus. 

Copper — Silver — Iron — Lead — Nickel — Separations — Moving elec- 
trodes — Mercury cathode. 


Volumetric Analysis 168 

Apparatus — Units of volume — Calibration by weighing — Calibra- 
tion by standardized bulbs — Calculation of results — Weight of 


one substance equivalent to a stated weight of another — Stand- 
ard solution for titration of but one substance — Burette reading a 
percentage reading — No system — Normal system — Decimal 
system — Choice of system — Temperature correction for standard 
solutions — Adjustment to exact concentration. 


Color Change of Indicators 207 

Ionization theory — Theory of chromophors — Classification of 
indicators — Description of indicators. 

Standardization 216 

Direct weighing — Weighing a substance produced by a measured 
volume of solution — Measuring the volume of solution required 
to react with a known weight of a substance of known purity — 
Titration against another standard solution — Primary standards. 


Experimental Volumetric Analysis 221 

Standard acids: Materials for standardization — Standardization 
by direct weighing — Preparation of pure sodium carbonate. 

Standard hydrochloric acid: Soda ash — Mixtures of carbonates 
and bases — Mixtures of carbonates and bicarbonates — Hardness 
and alkalinity of water. 

Standard bases: Selection of base for standard solutions — Stand- 

Standard sodium hydroxide: Concentration of the laboratory 
acids — Citric acid — Vinegar — Boric acid. 

Use of two standards : Limestone for agricultural purposes. 

Oxidation and Reduction 240 

Apparent valence. 

Potassium permanganate: Iron — Reduction of permanganate by 
chlorides — Primary standards — Reduction of iron solutions — 
Calcium — Manganese — Available oxygen. 

Potassium dichromate: Iron — Chromium. 

Iodine and sodium thiosulphatej Oxidizing power of peroxides — 
Copper — Bleachmgpowder — Standard iodine solution — Arsen- 
ical insecticides — Total arsenic and copper in Paris green. 


Titrations Involving Formation of Precipitates 276 

Silver — Halides and cyanides — Zinc. 






Rock Analysis 284 

Carbonate minerals : Carbon dioxides-Silica — Iron and aluminium 

— Calcium — Magnesium — Sodium and potassium. 
Silicate minerals: Moisture — Silica — Iron and aluminium — Man- 
ganese — Calcium — Magnesium — Sodium and potassium. 


Fuels 297 

Coal: Proximate analysis — Fusing point of ash — Ultimate analysis 

— Calorimetry. 
Gas mixtures: Apparatus — Absorbents — Illuminating gas — Chim- 
ney gases. 


Oils, Fats and Waxes 345 

"Burning oils: Specific gravity — Flash point — Burning point — 

Fractional distillation. 
Lubricating oils: Viscosity — Specific gravity — Separation of 

saponifiable from mineral oils — Chill test — Cold test. 
Edible fats and oils : Composition — Specific gravity — Melting point 
of fats — Iodine absorption number — Acid value — Saponification 
number — Insoluble and soluble acids — Reichert-Meissl number 
— Polenske value — Acetyl value — Maumen6 number — Specific 
temperature reaction — Qualitative reactions — Fish and marine 
animal oils in mixtures with vegetable oils — Examination of an 
unknown oil — Hardened oils. 

Water 391 

Industrial analysis: Corrosives — Incrustants — Foam producers — 
Hypothetical compounds — Treatment. 

Sanitary examination : Potability — Collection of samples — Physical 
examination — Chlorine — Organic nitrogen — Nitrogen in am- 
monia — Nitrites — Nitrates — Required oxygen. 

Steel and Alloys 432 

Steel and cast iron: Sampling — Standard methods — Solution and 
evaporation — Standard samples — Carbon — Silicon — Sulphur — 


Phosphorus — Titanium — Manganese — Tungsten — Chromium — 
Nickel — Vanadium — Oxygen. 

Treatment of steel: Thermal changes — Allotropism — Proximate 
constituents of slowly cooled steel — Ferrite — Cementite — 
Pearlite — Relation between structure and carbon percent — 
Austenite — Relation of proximate constituents to critical points 
of steel — Steel as a solid solution — Martensite — Critical points 
related to rate of cooling or heating — Hardening and annealing — 
Troostite — Sorbite — Quenching media — Tempering — Granula- 
tion — Overheating — Uneven carbon distribution — Streaks — 
Sulphur prints — Case hardening — Effect of working — Fatigue — 
Slag — Apparatus for metallographic work — Experiments. 

Brass and bronze: Tin — Lead — Copper — Zinc. 

Anti-friction metals: Tin — Antimony — Lead — Copper. 

Agricultural Materials 510 

Fertilizers: Moisture— Nitrogen — Availability of nitrogen — Phos- 
phorus — Potassium. 

Milk : Adulteration of milk — Specific gravity — Total solids — Ash — 
Total nitrogen — Casein — Albumin — Lactose — Fat. ' 

Cream: Ash — Fat. 

Condensed milk: Total solids — Ash — Nitrogen — Lactose — Fat — 

Butter and substitutes: Moisture — Fat — Casein — Salt — Examina- 
tion of fat — Coloring matter. 


The Fire Assay ." 564 

Gold and silver ores: Sampling — Weighing — Crucible process — 
Fluxes — Reducing agents — Oxidizing agents — Crucible charge — 
Reducing ores — Ores containing copper, arsenic or antimony — 
Cupellation — Inquartation — Parting— Annealing — Scorification. 

Table of Logarithms and Antilogarithms 586 

Index 593 


In the study of such divisions of chemistry as naturally pre- 
cede quantitative analysis the work has been mainly descriptive. 
The chemical and physical properties of elements and compounds 
have been determined. Laws of chemical action have been devel- 
oped and theories have been evolved for the explanation of such 
action and as generalizations upon which to plan further studies 
and investigations. In the courses in qualitative analysis an 
effort was made to detect and recognize elements and their 
compounds by certain characteristic reactions of these sub- 
stances and to separate more or less complicated mixtures into 
their constituents. 

Quantitative analysis is the next. logical step in the study of 
the composition of matter. The qualitative analysis should pre- 
cede the quantitative, unless the nature of the substance is 
already known, because in nearly all cases the presence of sub- 
stances other than those whose percent is being determined will 
make necessary certain modifications in the. method to be em- 
ployed. Whether or not a qualitative analysis is made it is 
a fact to be constantly kept in mind that an intelligent understand- 
ing of quantitative processes can be obtained only by a continued 
application of the facts and laws earlier learned to the newer 
processes which are being studied.- The industrial development 
of the world, as well as the evolution of chemistry as a science, 
would be a more rapid and substantial change were it not for 
the numbers of inadequately trained chemists who have found a 
place in industrial work and who have been content to allow 
their study to go no farther than the routine of following direc- 
tions without understanding. This is the inevitable end of the 
student who does not begin his scientific study with the definite 
determination patiently and persistently to think out each prob- 
lem to its logical conclusion as it presents itself, and who does not 
continue this process in his work in quantitative analysis, re- 
viewing his earlier work until the principles that were imperfectly 
understood expand and illuminate the newer problems. 






Cleanliness and Care. — In no other line of scientific work are 
neatness and care more essential than in quantitative analysis. 
It is desirable that the student should acquire speed so that his 
efficiency may be increased by his ability to accomplish much 
work in the time at his disposal but speed attained through care- 
less manipulation or through a sacrifice of a close study of the 
analytical process is of little avail because the results, when ob- 
tained, are not dependable. Speed should rather be acquired by 
an intelligent application of methods that are thoroughly under- 
stood. Several experiments may be performed at the same time 
without confusion if the analyst has cultivated the habit of delib- 
erate and clear thinking. It is highly important that as far as 
possible the operations should be systematized so that the mind 
may be left free for the more important problems that are not 
of a routine nature. For example a complete system of marking 
vessels and materials makes impossible errors due to confusion 
and lack of identification. Apparatus should be kept scrupu- 
lously clean, note books and other records with perfect system 
and the desk always cleared of apparatus that is not in use. Re- 
agents should be added carefully and, whenever practicable, 
should be measured approximately so that in case of error the 
defect may be corrected. A disregard of this precaution often 
makes it necessary to begin over again a determination that 
might have been saved. 



All of these points will be amplified as the work proceeds and 
as occasion offers. They are here mentioned in order to give 
some idea of the requirements of the work that is to follow. 

Limit of Accuracy. — A question frequently asked by beginners 
in quantitative analysis concerns the degree of agreement that 
is regarded as reasonable and necessary for the results of deter- 
minations made in duplicate. This question can never be 
answered without qualification. The ideal of all scientific 
work is absolute accuracy. How closely this ideal may be ap- 
proached will depend partly upon the individual analyst, but also 
upon the possibilities of the method with which he happens to be 
working. The latter is a variable quantity. Certain inorganic 
methods are capable of giving results that are quite reliable to 
within one hundredth of one percent while other (particularly 
organic) methods give only an approximation of whole percents. 
Thus no definite limit to the accuracy of methods in general can 
be set. Each case must be considered by itself and, while the 
student will use the utmost possible care in all cases, he must, 
for a time, be content to allow his instructor to be the judge of 
the accuracy that may be reasonably required. 

A very common fallacy is to the effect that no part of the work 
need be performed more carefully than that part which is necessarily 
least accurate. For example, it is said that if a certain part of an 
analytical process involves an unavoidable error of 0.10 percent, 
it is a waste of time to attempt to avoid errors in other parts of 
the work amounting to 0.05 percent or even 0.09 percent. This 
is a most unfortunate attitude for the analyst. This logic woul< 
lead one to the conclusion that if a method cannot give results 
nearer than 0.10 percent to the truth it should be given an excel- 
lent chance to depart 0.10 percent +0.09 percent from the trutl 
or 0.10 percent+nX0.09 percent, if there are n other places 
where errors may occur. Of course it is not necessarily true 
that all errors will have the same algebraic sign and the] 
may, to a certain extent, counteract each other in effect. 
is important to note, however, that there is no assurance tht 
they do counteract each other instead of accumulating. This 
misconception, no doubt, arises from the unquestioned fact that 
when a certain minimum error is unavoidable it is not wise to 
expend an undue amount of time in rectifying other possible. 


errors where the ratio of these errors to the larger error is very- 
small, because in such a method the analytical results have little 
significance as expressed in small fractions. For example, in 
the determination of volatile combustible matter in coal it is 
difficult to obtain results from analyses performed in duplicate, 
agreeing more closely than 0.2 to 0.5 percent, because of the 
impossibility of exactly duplicating heating conditions. It is 
then not profitable to weigh the sample so accurately as to justify 
an agreement within 0.005 percent unless there are several other 
places where rectification of similar errors is possible, simply 
because the results will have no significance in this remote decimal 
place and will not be expressed to the third or even the second 
place. If, however, there are five other points in the method 
where errors may be made as small as 0.005 percent but by undue 
haste or neglect might be as large as 0.05 percent, then the argu- 
ment already referred to would lead one to commit an error as 
large as 0.5 percent +5X0. 05 percent = 0.75 percent where it 
could have been kept as low as 0.5 percent+5X 0.005 percent 
= 0.525 percent. 

There still remains to be answered a legitimate question re- 
garding the number of decimal places to be reported. The opera- 
tions of multiplication and division that are involved in calcula- 
tions of analyses often give five or more figures in decimal places 
and yet even the novice understands that these figures have no 
significance since the probability of their correctness is very slight. 
A rule that is very generally followed is to report one decimal 
place farther than the one that is considered to be certainly correct. 
Usually the student has little evidence concerning the accuracy of 
his work other than the agreement of his duplicate determinations. 
If two analyses of an iron ore give 52.75 percent and 52.71 per- 
cent, respectively, it would be considered, from this standpoint 
alone, that 52.7 percent is certainly correct and that the average 
of 52.73 percent is a close approximation for the second decimal 
place. This average would therefore be reported. If the re-' 
suits were 52.75 percent and 52.67 percent the disagreement 
would be 0.08 percent and there is again agreement in the first 
decimal place although, at first sight, this would not appear to be 
true. The average 52.71 percent would then be reported. If 
the results were 52.75 percent and 52.45 percent the disagree- 


ment would be 0.30 percent. Evidently the first decimal place is 
but an approximation and the average 52.6 percent would be 

The rule for reports will then be as follows: Determine the 
disagreement of the two determinations by subtracting the smaller 
percent from the larger. Calculate the average and report as far 
as the first significant figure in the difference. 

Classes of Methods.— Quantitative analysis is, for the sake of 
convenience, divided into two general classes which are desig- 
nated as Gravimetric and Volumetric. The general nature of the 
work is indicated by these terms, since the first class consists of 
analyses made by means of measurements of weight while the 
second class consists of those made by means of measurements of 
volume. The discussion of volumetric analysis will be left for a 
later section of the book because the complete volumetric proc- 
ess cannot be carried on without a knowledge of gravimetric 
methods. Gravimetric analysis will therefore comprise the first 
part of the course. 

The problem, as it presents itself to the quantitative analyst, 
is to take a substance whose qualitative composition is at least 
partially known and so to treat it that a part or all of the constitu- 
ents may be expressed in terms of percents. Thus we speak of 
the " analysis" of a substance and of the "determination" or 
"estimation" of the constituents, the last two terms being used 
with practically the same meaning. 


The gravimetric determination of any constituent of a com- 
pound or mixture usually depends upon the precipitation of 
this constituent in such a form that it may be separated and 
weighed. This makes it necessary that the following conditions 
shall be fulfilled: (a) The precipitate must have an extremely 
small solubility in order that the amount lost in the filtrate may 
be negligible, (b) The precipitate must be of such a nature that 
it may be readily retained upon the filter and washed free from 
all impurities, (c) The precipitate must be susceptible to drying 
or strong heating, changing its composition either not at all, 
or in a perfectly definite and well-known manner, (d) The pre- 
cipitate should, if an electrolyte, be a strongly ionized one. The 
last condition will be understood when the nature of precipita- 
tion is considered. 

The determination of calcium in a calcium compound may be 
taken as an illustration of the gravimetric process. A definite 
weight of sample is taken, dissolved and the calcium precipitated 
as oxalate. This is separated by filtration, washed free from 
soluble salts and then strongly heated and weighed as pure cal- 
cium oxide. It is known from the formula for calcium oxide 
(expressing as it does the results of many careful analyses) that 


rp n7 of its weight is calcium. That is, 

r fin7 = weight of calcium, (1) 

where P = weight of calcium oxide found. 

Since this experimentally determined weight of calcium was ob- 
tained from a known weight of the calcium compound it follows 

weight of calcium X 100 . . . . ,. , 

— ^ — - = percent of calcium in the sample, 

S representing the weight of sample used. 


Combining these two expressions: 

40.07 X 100P ' , , . . ,, . ' 

— — , fi 07 ^ — = percent of calcium in the sample. (2) 

Factors. — It is seen from an inspection of the above formula 
that so long as determinations of calcium are being made in this 

manner the constant fraction — ' gg „„ will enter into all cal- 


culations. The calculation will be somewhat shortened if these 

factors are collected into the decimal fraction 71.464. 

Expression (2) then becomes 


— -~ = percent of calcium. (3) 

This is a typical form of expression for the experimentally deter- 
mined percent of any element, radical or compound. The deci- 
mal fraction 71.464 is known as a gravimetric factor, which may 
be defined as the percent of a given element or group of elements 
in or equivalent to a pure compound weighed. The pure substance 
weighed usually contains the element or radical which is being 
determined but does not always do so. 

If the factor in general is represented by F, equation (3) may 
be written 


-5- = required percent. (4) 

Atomic weights have been calculated from the results of the 
most careful experimental work of modern investigators. For 
this reason the factor is the most reliable of all of the experimental 
data used by the analyst and the calculation may be carried far- 
ther than that of analytical percents. Five significant figures 
should usually be recorded. 

Factor Weights. — In equation (4), F is a constant for all deter- 
minations of the particular element or group of elements for whic] 
it has been calculated. It is possible so to choose the weight oi 
sample taken as to simplify the calculation represented by equa- 
tion (4). For instance, by taking a sample weight equal in grai 


to the value of the factor, « = 1 and equation (4) becomes 

P = required percent. 


In such a case the weight of the precipitate, expressed in terms 
of grams or fractions, would indicate, without calculation, the 
required percent of the constituent being determined. 

A weight of sample equal in grams to the value of the factor is 
usually too large a quantity to be handled readily and a definite 
fraction of this weight (0.5, 0.2, 0.1, 0.01, etc.) may be used in- 
stead. Any such weight is called a factor weight, which may be 
defined as a quantity equal in weight units to the value of the gravi- 
metric factor, or to some simple fraction of this value. 

The use of such weights involves adjusting the quantity of 
the sample on the balance to a certain specified value. If this 
device is to result in the saving of time it is obvious that the 
weighing must not require any extra time. Therefore the use of 
factor weights is generally confined to the determination of such 
constituents as occur in small amounts in a sample that is to be 
analyzed, so that relatively large amounts of the sample may be 
used and the accuracy of the weighing may be correspondingly 
less without impairing the accuracy of the determination. 

Use of Logarithms. — Most of the results of quantitative analy- 
sis involve chiefly multiplications and divisions. On this account 
a four or five place table of logarithms will be found extremely 
serviceable. Many students hesitate to use such tables because 
they find that mistakes are more readily made and that the cal- 
culations require more time than is the case when ordinary 
arithmetical calculations are employed. If this is true it is be- 
cause of a lack of facility in handling such tables and this comes 
through lack of practice. The beginner in this work is strongly 
urged to make use of his tables even though 'this should, at first, 
involve more work and care than direct arithmetical solutions. 
Wherever a constant occurs in a solution the logarithm of this 
constant may be recorded and used in all similar calculations. 
This is particularly helpful in the case of factors. 

Chemists , Slide Rules. — A slide rule may be substituted for 
the logarithmic table but it is not accurate beyond the third 
significant figure and should not be used for the calculation of 
careful analyses. Chemists' slide rules have been devised, having 
marked at the proper points the constants that represent the 
analytical factors. By the use of such a rule a percent may be 
read from one adjustment of the slide. 


Fig. 1. — Chemist's slide rule. 

Following is the international table of atomic weights for 1918. 
International Atomic Weights, 1918 




Aluminium. . 
Antimony. . . 




Bismuth. . . . 


Bromine. . . . 
Cadmium . . . 
Caesium. . . . 




Chlorine. . . . 
Chromium . . 


Columbium . 


Dysprosium . 
Erbium ...... 

Europium. . . 




Germanium . 
Glucinum. . . 



Holmium. . . 
Hydrogen. . . 





Krypton. . . . 



Lutecium. . . 
Manganese. . 
Mercury. . . . 






















207 . 20 



Molybdenum. . 
Neodymium. . . 








Phosphorus. . . . 













Sodium , 












Uranium , 











1. Calculate the factors and their logarithms indicated in the follow- 
ing table, recording in the proper spaces for future use. The atomic 
weights as recorded in the preceding table should be used and the factors 
should be calculated in five significant figures. 

Table of Gravimetric Factors and Logarithms 




Log Factor 









A1 2 3 


BaS0 4 


S0 4 
S0 3 


SrS0 4 


K 2 PtCl 6 



K 2 S0 4 



K 2 S0 4 

Na 2 S0 4 




Mg 2 P 2 7 

P • 

Mn 2 P 2 7 


Mg 2 As 2 7 


General Operations 

Certain operations enter into the majority of gravimetric 
analyses and the general conduct of these will be discussed briefly, 
principles and precautions being indicated. It will be readily 
understood that variations in standard procedure will be de- 
pendent upon the nature of the substance being analyzed and 
such variations will be discussed as the necessity arises. 

Preparation of Samples. — The object of all preliminary work 
with samples is to make it possible to obtain, for the actual analy- 


sis, a portion that shall truly represent the average composition 
of the entire material at hand. This matter is likely to be treated 
lightly by the beginner, but proper sampling is often one of the 
most difficult problems of quantitative analysis. It is generally 
necessary to use a quantity of 1 gm or less and if the substance 
is not homogeneous this small quantity may have an average com- 
position that is very different from the average composition of 
the entire material being investigated. No matter how carefully 
an analysis may be performed or how accurate the results ob- 
tained, if the substance used does not represent the average of 
the substance originally at hand the results become nearly or en- 
tirely valueless. If the substance is practically homogeneous the 
operation of sampling involves nothing more difficult than grind- 
ing down to a degree of fineness required for the work. This is 
the case when the substance is an approximately pure chemical 
compound, such as will be used for the earlier exercises. 

The gross sample, as the analyst receives it, may be in the 
form of lumps, as is frequently the case with minerals, or it may 
be in the form of small pieces, crystals, powder, or solution. In 
any case except that of liquid samples, the object is to reduce the 
size of pieces to that required for the analysis (usually a rather 
fine powder) and at the same time to select from the total mass 
such a quantity as is required for the experimental work. The 
original sample is often quite large. It may vary from a few 
grams, or less, to many pounds. It is obviously unnecessary 
and practically impossible to grind the entire amount into a fine 
powder. The operation then resolves itself into a thorough mix- 
ing and progressive grinding and dividing. Many forms of 
both hand and power grinders are in common use. For the first 
exercises nothing more complicated than a porcelain mortar an 
pestle will be required. 

Mixing and Dividing.- — The mixing and dividing is best carried 
out by using a sheet of oilcloth or paper and a spatula. In many 
laboratories it is customary to use oilcloth, particularly for mixing 
minerals. This is convenient but offers the possibility of con 
tamination (" salting") of one sample by the remnant of one that 
has preceded it. It is better to use a large sheet of tough, flexible 
paper, which can be discarded after using. The sample, after 
haying been broken down to the proper maximum size of pieces, 



is placed on the paper and thoroughly mixed by rolling diagonally 
across the paper and alternating the direction of rolling as illus- 
trated in Fig. 2. The proper rapid manipulation of the paper is 
attained only after considerable practice. One precaution is 
essential: the corner of the paper that is lifted must be drawn 
across, low down, in such a manner that the pile of material is not 
caused to slide along the paper but is turned over upon itself so 
that the bottom is brought entirely to the top. In this way only 
can a segregation of larger and smaller particles be prevented. 
Since the larger and smaller particles usually have different com- 
position it is essential that they should be thoroughly mixed. 

Fig. 2. — Manipulation of paper when mixing samples. 

The number of times that the sample is rolled before dividing will 
depend upon the degree of homogeneity and the accuracy re- 
quired in the analysis. In the assaying of gold and silver ores 
it is not unusual to require one hundred times. 

Quartering. — When the first mixing is finished the pile is 
made approximately circular and is then divided, by means of a 
spatula, into quarters. Opposite quarters are carefully scraped 
to another sheet of paper, ground finer if necessary, remixed and 
quartered as before. This process of grinding, rolling, and quar- 
tering is continued until a sample is finally obtained, small 
enough in quantity and fine enough in texture to serve the pur- 
pose of the final weighing and analysis. The maximum size of 


particles to be allowed in any particular mixing and quartering 
will depend upon the total quantity of material being handled 
in this operation. No particle should be so large that its inclusion 
in any quarter would cause the average composition of this quar- 
ter to be appreciably different from the average composition of 
the entire pile. This means that the ratio of the size of the largest 
particle to the size of the quarter should not be greater than a 
certain maximum value. What this maximum value shall be 
must be arbitrarily determined by the nature of the sample and 
the degree of accuracy required in the analysis. It is obvious 
that the part can perfectly represent, in composition, the whole 
only when the largest particle is infinitesimal. It is equally 
obvious that this limit is impossible and unnecessary in practice 
and we may say that, in general, the ratio of the largest particle 
to the portion that includes it 'should not be greater than 0.01 
percent. If this condition is met, then, after thorough mixing of 
the sample, the chance inclusion or exclusion of any given particle 
cannot modify the results of the analysis to any appreciable 

The maximum size of the particles to be obtained in the final 
portion that is to be weighed and used in the analysis must be 
determined, not only from the above considerations, but 
also by the nature of the operation to follow the weighing. 
This is usually solution or fusion. If the substance is considered 
to be almost absolutely homogeneous and if it is easily soluble (as, 
for example, a crystal of cupric sulphate) then the grinding need 
be carried no farther than is necessary to permit the easy adjust- 
ment, between fairly narrow limits, of the weight taken for analy- 
sis. In such a case, if a sample of 0.3 to 0.5 gm is required, then 
no particle should weigh more than about 0.1 gm. If, however 
the process of solution or fusion is a difficult one to accomplish 
or if the material is far from being homogeneous, the grinding 
is carried much farther, in order to provide a very large surface 
of contact between the particle and the solution or flux, or in 
order to conform to the rule of maximum size of particles, stated 
above. In many cases, as with minerals, the maximum size o 
particles is fixed by causing the sample to pass through a siev 
having meshes of stated dimensions. A gold ore may be groun 
to pass a sieve having 100 or 200 meshes to the linear inch. I 





such a case one should not make the mistake of grinding and 
sifting a portion until a sufficient quantity is passed, discarding 
the remainder. This would cause an error because the particles 

Fig. 4. — Division by quartering. 

that resist grinding longest are less brittle and have a composition 
different from that of the particles which pulverize easily. 

The reason for dividing into quarters after each mixing and for 
selecting opposite quarters will be understood from the following: 



Close examination of the pile of unmixed material will reveal the 
fact that, even after the most thorough and careful mixing, it is 
not entirely homogeneous. Around the circumference of the 
base the particles are coarser and they may be gathered toward 
one side. Around the apex of the conical pile there is a collection 
of coarser particles. If we simply dig in at random for the por- 
tion to be removed the lack of homogeneity will alter the char- 
acter of the portion removed. 


Fig. 5. — Riffle; for automatic division of dry samples. 

Fig. 3 shows how quartering will properly divide a pile with 
coarse material over the top. Fig. 4 shows how the opposite 
quarters, no matter in what direction the cuts be made, will ob- 
tain the average of a non-homogeneous pile, while a cut into 
halves will do so only in case the cut is made in the direction ab. 
In these diagrams the conditions are purposely exaggerated. 

An automatic sampler, called a riffle, is illustrated in Fig. 
5. The nature of its action is easily seen. 

In case the substance to be analyzed is a liquid, the operation 


of sampling is usually a simple one, consisting of thorough mixing 
before the removal of the proper quantity for the analysis. 

Solution. — After the sample of solid substance has been prop- 
erly selected and weighed, the next operation is usually one of 
solution. What the solvent shall be is determined by the nature 
of the case. It may be water, concentrated or dilute acids, bases 
or salts, organic solvents, or a solid that is to act as a dissolving 
flux when heated with the substance. In any event it is desirable 
to use a relatively small amount of the solvent, not only because it 
must finally be entirely removed, but also because all precipitates 
dissolve to some extent and it is only by keeping the amount of 
solvent down to the least quantity that is workable that the loss 
of precipitate is reduced to the practicable minimum. 

Precipitation. — The process of producing precipitates for use in 
quantitative analysis is one that has received much study from 
the quantitative standpoint. Every substance has a more or less 
definite solubility, under definite conditions, and it is necessary 
to reduce this (already small) solubility to the least possible 
quantity. According to the ionic hypothesis reactions between 
most inorganic and many organic compounds are reactions of the 
ions of the compounds that change. 

Mass Law. — The reaction between potassium chloride and 
silver nitrate is, under proper conditions, nearly complete — 
nearly enough, in fact, that we regard it, for practical analytical 
purposes, as complete. Both of these salts are highly ionized 
in moderately concentrated solutions. When the solutions are 
mixed, silver chloride is formed and precipitated. The usual 
expression for the reaction is 

KCl+AgN0 3 <=>AgCl+KN0 3 . 

This really expresses merely the substances as originally taken for 
making the solution and the products as they are when removed 
from solution. The equation that expresses the real reaction is 

Ag+CteAgCl, (1) 

also, to some extent, 

K+N0 3 ^KN0 3 . (2) 

Reaction (1) is the one that is important in this instance. 


The principle known as the Mass Law, as formulated in 1867 
by Guldberg and Waage, 1 states that the velocity of any chemica 
reaction is directly proportional to the active masses of the re- 
acting substances. In a reversible reaction this applies to the 
reverse as well as to the direct reaction, and in the case already 
cited this means that the velocity is proportional to the concen- 
trations of the reacting bodies, whether these are ions or molecules. 
Stated symbolically for reaction (1), we have 

vel = kC^Cci and 

vel=k'C Agcl 

where vel and vel are the velocities of the reactions taking place 
in the direction of the arrows, k and k' are constants and C^, C c -i 
and C Agcl the concentrations, stated in weight equivalents per 
unit volume, of silver ions, chlorine ions and silver chloride mole- 
cules at any given instant. At the moment of adding the silver 
nitrate to the potassium chloride C A + g and C c -ihave their maximum 
values while C AgcI =0. As the reaction proceeds C Agcl increases 
while C A + G and Cci decrease, owing to the formation of silver 

chloride molecules. For this reason vel decreases while vel in- 
creases. Equilibrium occurs when the direct and reverse re- 
actions proceed with equal velocity. Therefore at equilibrium 

vel = vel and 

kC A + g C c -j = k'C AgcL 

Therefore C^Ce^^C.^^KC^, where K = ~ (3) 

Solubility Product. — Equation (3) is an expression of the fact 
that at equilibrium there is a definite and constant relation 
between the concentrations of the reacting bodies and of the 
bodies formed. The laws of solubility tell us that there is simi- 
larly a constant relation between the concentration of dissolved 
substance and that of the same substance undissolved after pre- 

l Ostwald's Klassikor, No. 104. 


cipitation has begun and the stable condition known as saturation 
is reached. In the case of silver chloride, 

C A gciK , C Agc i_ soI ici 
If concentration is taken to mean the ratio of the weight of 
solute (expressed in gram equivalents) to the weight of solution, 
then C^ci-soHd must equal 1, no matter how much or how little 
solid may be present. Therefore K'C Agcl _ solid has a constant 
value : 

C Ag ci = K / C A g Cl _ solid =K and 

C; g C c ^ = KC Agcl = KK' = K'' (4) 

Equation (4) is an expression of the fact that in a solution satu- 
rated with a precipitated substance, the product of the concen- 
tration of the ions of this substance is a constant quantity. This 
product is known as the solubility product and it is of great impor- 
tance in quantitative analysis. This product is a maximum 
value for the condition known as saturation. It is exceeded when 
the solution is supersaturated as every solution is before precipi- 
tation begins. Such precipitates as calcium oxalate and silver 
chloride attain the solubility product comparatively quickly, 
magnesium ammonium phosphate slowly. 

According to condition (a) on page 5 it becomes necessary 
that the solubility product as well as the concentration of dis- 
solved molecules shall be extremely small, if the substance is to 
be useful for quantitative purposes, as otherwise the results of 
the analysis will be vitiated by an appreciable loss of the substance 
being precipitated. It should be noticed also that it is the solu- 
bility product that is constant and not the concentration of the 
individual ions. It is usually true that a single gravimetric 
determination concerns a single radical and not the entire com- 
pound; thus, in the case already considered, if the chlorine radical 
were being quantitatively determined the chloride would bfc 
weighed and silver nitrate used as the reagent. It is possible 
so to increase the factor representing the concentration of the ion 
which is not being determined that the other factor must assume a 
very low value. As an illustration, suppose that a sample of a 
soluble chloride is weighed, dissolved in water and a solution of 
silver nitrate added until all of the chlorine is precipitated except 


that represented by an ordinary saturated solution of silver 
chloride. At 18° 100 gm of water will dissolve 1.7 X10~ 4 gm 
or 1.19X10 -6 gram equivalents of silver chloride. 1 In such a 
highly dilute solution ionization is nearly complete (i.e., K is 

very large) and since there is present practically the same number 

of gram equivalents of Ag and CI as represents the number of 
gram equivalents of dissolved silver chloride, 0^0^= (practi- 
cally) (1.19X10- 6 ) 2 = 1.42X10- 12 . This is the solubility pro- 
duct of silver chloride. If to this solution 0.001 gram equivalent 
of silver ions is added (a condition approximately attained, with- 
out appreciable dilution, by the addition of 3.5 cc of 5 per- 
cent silver nitrate solution) 

CAgCcI = (1-19X10~ 6 ) 2 becomes upon substitution of 
C A + g =lX10" 3 , 
10- 3 C c i=1.42X10- 12 , whence 
C c -j=1.42X10- 9 

Thus by adding the precipitating reagent but 3.5 cc in excess of 
the amount required to form silver chloride completely, the 
amount of chlorine remaining in solution has been reduced from 
1.19 X 10 -6 gram equivalent to 1.42 X 10 -9 gram equivalent; 
stated in terms of the actual weight of chlorine, 0.00004 gm 
remaining in solution has been reduced to 0.00000005 gm, 
a quantity so small that it has absolutely no significance in any 
but the most refined methods of analysis, such, for instance, as 
are involved in atomic weight determinations. From this reason- 
ing follows the general rule that the precipitating reagent should 
always be added in a certain slight excess over the amount equiva- 
lent to the substance being precipitated. Of course this is done 
also as a matter of convenience, since it is obviously impossible 
to add exactly the equivalent amount of the reagent and a defi- 
ciency would be fatal to the success of the experiment. 

Attention has already been called (page 5) to the desirability 
of having highly ionized electrolytes for precipitates. This can 
now be understood since the molecular form of the dissolved sub- 
stance constitutes the irrecoverable portion. The effect of an 
excess of reagent upon the solubility of the precipitate becomes 

iRohlrausch and Rose, Z. physik. Chem., 12, 234 (1893). 


less as we pass to precipitates whose ionization is small and 
becomes negligible or zero with non-electrolytes. 

Colloids. — There is a certain class of precipitates with which we 
cannot deal in so definite a manner. This is the class known as 
colloids. Certain important distinctions are to be made between 
the colloids and the crystalloids. While the latter have certain 
very definite effects upon the freezing-point, vapor pressure and 
osmotic pressure of solvents, the former have little that is appre- 
ciable. The solubility of the crystalloids in various solvents is a 
tolerably definite and constant quantity, temperature being speci- 
fied, while there is no definite solubility for the colloids. The 
solubility of those crystalloids which ionize is affected in a definite 
and calculable manner by other electrolytes. The solubility of 
the colloids is affected by the presence of electrolytes but not in 
the same definite way. Examples of well known colloids are 
ferric, chromium and aluminium hydroxides, arsenic sulphide, 
and many organic compounds such as the albumens. It is now 
tolerably well established that the colloidal " solutions" contain 
molecular aggregates that are large enough to exhibit more of the 
properties of invisible mechanical suspensions than of true 
solutions, although there is little doubt that the difference 
between the pseudo solution of a colloid and the true solution of 
a crystalloid is one of degree of molecular association rather than 
of kind, since most true solutions contain associated molecules. 
In order to distinguish the true solution from the pseudo solution 
of a colloid the latter is called a "sol" and if water is the solvent 
the sol is a "hydrosol." 

The importance of colloidal sols to the analyst lies in the fact 
that they do not respond to the effect of excess of reagent when 
attempting a precipitation and that the colloid will remain in 
invisible suspension to an extent that would normally represent 
a greatly supersaturated solution. The absence of effect of 
excess of reagent is, no doubt, due to the absence of any but a 
very slight concentration of ions in the solution of the precipi- 
tating colloid. Colloidal sols are broken down with precipitation 
(flocculation) of the colloid by the addition of certain electrolytes. 
The flocculated colloid may, under certain conditions, return to 
the sol. Certain colloids will not so return after drying. Colloids 
of the former class are known as "reversible" and those of the 


latter class are called " irreversible." Some examples of reversi- 
ble colloids are dextrine, gums, albumens and many organic 
colloids, also silver chloride and molybdic acid. Examples of 
irreversible colloids are metal hydroxides, stannic acid, arsenic 
sulphide and colloidal metals. Reversible colloids may often, by 
certain treatment, be changed into irreversible colloids. Thus 
strong heating causes the change of reversible silica into irre- 
versible silica. 

Enlargement of Particles of the Precipitate. — Complete separa- 
tion of a precipitate by nitration requires that the smallest indi- 
vidual particles of the precipitate shall be larger than the largest 
pores of the filter, although any further considerable increase in 
size is unnecessary and undesirable, because thorough washing 
is then more difficult. Amorphous precipitates offer some diffi- 
culty, on account of the fact that sols are likely to form and these 
cannot be separated by the finest filters. Many crystalline pre- 
cipitates, notably barium sulphate and calcium oxalate, have a 
tendency to precipitate in very fine crystals so that, even when 
the finest grades of filter paper are used, appreciable quantities 
of the precipitate are lost. This tendency can be partly over- 
come by observing certain precautions during the process o: 
precipitation, such as adding the reagent slowly, stirring the 
solution vigorously, heating the solution while adding the pre- 
cipitant, etc. 

It has already been noticed that a state of supersaturation 
always precedes precipitation. The extent to which super- 
saturation occurs depends, among other things, upon the rate at 
which the precipitatng substance is formed in the solution and 
this, in turn, depends upon the rate of addition of the precipitant/ 
The formation of crystal nuclei requires an appreciable amount of 
time. If the reagent is added very rapidly a relatively *large 
degree of supersaturation is produced before the substance which 
is being formed begins to crystallize. If, on the other hand, the 
reagent is added slowly and thoroughly mixed by stirring, crystals 
begin to form before any considerable supersaturation occurs. 
This condition is very important to the formation of large 
crystals. If there is but slight supersaturation at the beginning 
of precipitation and the solution is vigorously stirred as more 
reagent is added, a relatively small number of crystal nuclei 


may serve to maintain a condition of approximate equilibrium 
with the solution. If large supersaturation occurs before pre- 
cipitation begins, and if the rapid addition of reagent is continued 
many crystal nuclei will form at all parts of the solution. If 
we regard the quantity of precipitate ultimately formed as the 
same in either case it is evident that the solutions from which 
large numbers of crystals form will produce the finer precipitates. 
From this follows the rule that the precipitating reagent should 
always be added slowly and the solution should be stirred during the 
addition of the reagent. 

Even after the most careful attention to the method of precipi- 
tating it occasionally happens that a precipitate is produced in 
such a fine state of division that a portion passes through the 
paper. When this occurs it is usually possible to enlarge the 
crystals to a size that will admit a ready separation by filtration, 
by allowing to stand for some time in contact with the mother 
liquor, with or without the application of heat. This enlargement 
is not merely a coherence of small particles to form larger ones, 
but is a real enlargement of individual crystals, the number of 
particles decreasing as the average mass increases. Ostwald 1 
demonstrated that this change in the form of the crystalline aggre- 
gate is due to the fact that small crystals have a greater solution 
tension than the large ones. The result of digestion in contact 
with the solvent is that the solution containing a definite con- 
centration of the dissolved precipitate cannot remain in equilib- 
rium with both large and small crystals. If in equilibrium with 
the larger ones, and thus saturated with respect to these, it is 
under-saturated with respect to the smaller ones so that these 
dissolve to some extent. The solution then becomes supersatu- 
rated with respect to the larger crystals, and some of the sub- 
stance precipitates (crystallizes) on these. The end of such a 
process is the total disappearance of the small particles and the 
appearance of a smaller number of particles having a larger 
average size. Hulett 2 showed that the solubility of barium 
sulphate could be increased from 0.00229 gm per liter (the 
solubility of the precipitated salt) to 0.00618 gm per liter by 
pulverizing the crystals. 

1 Z. physik. Chem., 34, 495 (1900). 

2 Ibid., 34, 69 (1900). 


Filtration. — Materials generally used for filtering are combus 
tible (organic) or non-combustible (inorganic). The precipitate 
is to be weighed later and it is necessary that some means be founc 
for either entirely destroying the filtering medium or discounting 
its weight. Paper filters are, at present, used for the majority 
of gravimetric analyses although several inorganic substitutes 
are being used to a greater extent than formerly because they pos- 
sess the advantage of having no action on most precipitates ai 
high temperatures. 

Filter Paper. — Paper for qualitative work is necessarily made 
with great care since it must combine the quality of considerable 
strength with that of a uniform porosity. For quantitative 
purposes the paper must have even greater uniformity of texture; 
it is nearly always removed by burning, after the precipitate 
has been thoroughly washed, and it is essential that the inorganic 
matter which is always found in organic fibers and left as ash 
shall be either in sufficiently small quantity to be negligible or 
that its quantity shall be uniform so that a definite weight may be 
subtracted from the weight as found for each precipitate. The 
better grades of quantitative papers have been washed with 
acids so that much of the inorganic matter has been removed. 
This results also in softening the fiber, making it a better filtering 
medium although more frail and subject to rupture if suction is 
applied. For the highest grade of paper hydrochloric acid and 
hydrofluoric acid are used and the weight of ash for a paper 
9 cm in diameter is reduced to as low as 0.00011 gm. On account 
of the fact that this weight is negligible in ordinary determina- 
tions such paper is often erroneously called "ashless" paper. 

Much trouble will be obviated in the use of paper filters if some 
care is exercised in the selection of funnels. A circular paper is 
usually folded into quadrants, then opened out to form a cone, the 
sides of which include an angle of 60°. Funnels as purchased for 
chemical work are supposed to be made with a 60° angle but com- 
paratively few are exactly of this form. Others have the correct 
angle between the sides but near the apex of the cone the shape is 
irregular so that the corresponding part of the paper is not prop- 
erly supported. 

Reduction by Burning Paper. — Certain precipitates are 
changed by ignition in contact with organic matter, in such a 


manner that the composition of the resultant substance is un- 
certain. Examples are compounds of easily reducible metals, 
such as silver, platinum, lead, etc. If silver chloride is heated 
in contact with burning paper it is partially reduced to metallic 
silver but the amount of reduction is not known in any given 
case, so that no factor can be used to obtain the weight of either 
silver or chlorine. In such a case the precipitate may be merely 
dried without strong heating, or it may be treated by a process 
such that the altered portion will be reconverted into the original 
form, or, finally, it may be filtered on a filter of some inorganic 
material that is unaltered by high temperatures. The first 
method of treatment is objectionable because it involves drying 
and weighing the paper both before and after filtration and wash- 
ing. The cellulose fiber is somewhat soluble in almost any liquid, 
also it is practically impossible to dry it to the same degree of 
hydration. This leaves the second method of treatment (re- 
conversion of the changed precipitate) as the only desirable 
one, if paper is to be used as the filter. In the case of silver chloride 
the paper and precipitate are dried and most of the precipitate 
removed and preserved. The paper is then burned in air, this 
reducing a compartively small part of the precipitate. This 
small amount of reduced silver is treated in the crucible 
with nitric acid, which redissolves it, and with hydrochloric 
acid, which reconverts the silver nitrate to silver chloride. 
The acids are evaporated, the remainder of the (unchanged) 
silver chloride is added and the whole is heated and weighed. 

Reduction of Pressure for Filtration. — In most cases the only 
pressure needed or desired for causing the liquid to filter is that 
due to gravity. If the funnel is properly made and the paper 
fits well the stem will fill with liquid and this increases the speed 
of filtration on account of the length of the 'column. Some 
precipitates, particularly those of a colloidal nature, clog the pores 
of the filter and render filtration a slow and tedious process. 
In such cases it is necessary to use some means for diminishing 
the pressure beneath the liquid . The funnel is inserted in a rubber 
stopper, placed in the top of a pressure flask or bell jar, to which 
is attached a suction pump. Since .there is no support for the 
apex of the filter paper, a supporting cone of platinum or strong 
paper is placed in the funnel under the paper. The cone need 



not be used if the paper is one that has been partially parch- 
mentized or " hardened" by treatment with sulphuric acid. 

Inorganic Filters. — Many chemists prefer to use inorganic 
materials for such precipitates as are reduced by contact with 
burning paper, and for this purpose the crucible devised by 
Gooch 1 is widely used. This piece of apparatus takes the place 
of both filter paper and crucible, the precipitate being either 
simply dried or strongly heated directly in the filter, which has 

Fig. 6. — Filtering crucible and bell jar. 

previously been weighed. It consists of a tall crucible of 
platinum, with a bottom having fine perforations. This is 
placed in a holder which can be fitted to a flask or bell jar for 
use with diminished pressure. Into the crucible is poured a 
small amount of finely shredded asbestos fiber suspended in 
distilled water, the water is drawn out and the asbestos sucked 
down over the perforated bottom making a thin felt which is an 
admirable substitute for trie usual paper. The crucible may 

1 Proc. Am. Acad., Feb. 13, 1878. 


then be rinsed with alcohol to promote rapid drying; it is dried 
and weighed and is then ready for use as a filter. In using the 
Gooch crucible it is essential that the suction pump be running 
when any liquid is poured into the crucible as otherwise the 
felt will be stirred up and disintegrated and some of the fiber 
will pass through the perforations. The loss of asbestos is the 
most frequent source of error and even with the greatest care it 
occasionally happens that small amounts are lost by washing 
through. The asbestos for this purpose should be as nearly pure 
•silicate as possible and free from oxides of iron or other metals. 
It is first thoroughly shredded, then is digested with concentrated 
hydrochloric acid to remove all soluble material and is finally 
washed free from hydrochloric acid and soluble chlorides. The 
purified material is kept in bottles covered with distilled water, 
ready for use. Asbestos cannot be used in any case where the 
solution to be filtered is basic because it is somewhat soluble in 
bases. In case strong ignition is required the crucible of platinum 
is fitted with a cap which covers the bottom portion, thus pre- 
venting any loss of asbestos during heating. 

The porcelain modification of the Gooch crucible was devised 
by Caldwell. 1 It is not well adapted to strong heating because 
of its liability to crack and also because the asbestos felt curls up 
and partly disintegrates and will inevitably cause a loss of the 

A method has been devised 2 for making a platinum sponge 
of such texture that it is suitable for filtering. This has, as yet, 
not found a very extended use. Experiments have also been 
conducted with a view to adapting porous material, such as 
unglazed porcelain, "alundum" (vitreous aluminium oxide), 
etc., to the purpose of quantitative filters. Such materials 
may be used extensively in the future. They possess the very 
decided advantage that there is no possible loss of loose material 
such as is liable to occur during the use of the ordinary Gooch 

Any filtrate obtained in the processes of quantitative analysis 
should be received in a beaker or other vessel which has been 

1 J. Am. Chem. Soc. 13, 105 (1891). 

2 Munroe: Chem. News, 58, 101 (1888) and Snelling: J. Am. Chem. Soc, 
31, 456 (1909). 


thoroughly cleaned. It is often thought that the nature of the 
receiver is unimportant because the nitrate is to be finally dis- 
carded. It frequently happens, however, that a filter paper 
breaks, allowing the precipitate to escape, or the precipitate is so 
fine as to run through the pores, or it is discovered that precipita- 
tion from the filtrate is not complete. In any of these cases th< 
filtrate must be returned to the original precipitating vesse 
and if the receiving vessel were not clean the determination is 

Washing. — The soluble products of reactions of precipitation, 
as well as soluble impurities originally present, must be washec 
away from the precipitate on the filter. If the precipitate itself 
exerted no action upon the dissolved substances washing would be 
comparatively easy as will be evident if the process is considerec 
in detail. 

If we assume that the filtrate is allowed to drain away fron 
the precipitate until a definite small quantity a remains, that th< 
wash water is then added to make a volume b, stirred up with the 
precipitate and allowed to drain until volume a again remains- 
and that this process of dilution and draining is repeated with 
each additional washing, then each addition of wash water reduces 

the concentration of dissolved matter by the fraction r of the 

previous concentration. If the concentration in the original 
mother liquor is represented by c the first addition of water 

reduces this to-vXc, the second to tXtXc= (-t) c, the nth 
~r) c. After draining the last wash water the quantity of 

-t) ca. If, 

for example, c = 5 percent (which is greater than the usual con- 
centration of dissolved impurities) a = l cc and 6 = 10 cc, 
then after one washing the amount of impurity remaining is 

r^X 0.05X1 =0.005 gm; after two washings the amount is 


X 0.05X1 =0.0005 gm; after three washings there remains 


0.00005 gm. The last is a quantity that would not be appreci- 
able to the ordinary analytical balance. 



Interference by Adsorption.— The above method of reasoning 
is not strictly valid because the precipitate itself exercises an 
influence upon the solution, resulting in diminished efficiency 
of the washing. DeVille 1 first demonstrated the fact that 
wherever a solution is in contact with a solid, the former is 
slightly more concentrated in the region adjoining the surface 
of contact. This difference in concentration is due to a mutual 
attraction between the molecules of solid and those of solute. 
Where this surface is merely the wall of the containing vessel 
the difference in concentration is not made evident by any 
ordinary means of measurement because the thickness of the more 
concentrated layer is very slight and the interior surface of the 
containing vessel is not large. A precipitate, on the other hand, 
has a very much larger surface, owing to the usual fine state of 
division, the same statement being true with regard to the filter 
paper. The portion a of the solution is therefore always one of 

greater concentration than c or ( H n c, and the amount of impurity 

rj ca. This action, known as " ad- 
sorption/ ' does not greatly obstruct the washing of precipitates 
that are decidedly crystalline in character but causes much 
trouble in the case of amorphous and colloidal precipitates, pos- 
sibly because of the very great surface possessed by these bodies. 
The great surface exposed by hydrosols is illustrated by the 
figures of the following table adapted from the work of Wo. 
Ostwald. 2 A cube having a length of side equal to 1 cm is 
subdivided into smaller cubes, with this result : 

Length of side 

1.0 cm 

0.1 cm 

0.01 cm 

0.001 cm (Diam. of blood corpuscles = about 

0.007 cm). 

1.0 /i(0. 0001 cm; Diam. of small coccus) 

0.1 a 

0.01// (limit of ultramicroscopic visibility) 

l.O/i/* ( = 0.000001 mm; Diam. of smallest colloid 

. 1 fift (Diam. of elementary molecules) 

1 Ann. chim. phys., 38, 5 (1853). 
1 Grundriss der Kolloidchemie, 85. 

Number of 

Total surface 


6 sq cm 


60 sq cm 


600 sq cm 


6,000 sq cm 


6 sq meters 

10 1 * 

60 sq meters 


600 sq meters 


6,000 sq meters 


60,000 sq meters 



While colloidal particles are not to be regarded as cubes, their 
surface would vary with continued subdivision in the same way 
At the moment of precipitation a substance having a surface 
relatively so enormous, may show the effect of adsorption to 
marked degree, much of soluble salts being carried out of the 
solution. Mocculation no doubt diminishes the surface consid- 
erably but the flocculated colloid still possesses a much greater 
surface than the same weight of a crystalline solid may have 
The hydroxides of iron, aluminium and chromium retain dis- 
solved salts or bases with great tenacity and are extremely diffi- 
cult of purification by washing 
The number of washings neces- 
sary for the satisfactory purifica- 
tion of a given precipitate wil 
depend upon the nature of both 
precipitate and dissolved sub- 
stance and must be learned by 
experience. A safe plan to fol- 
low is that of testing the wash- 
ings until they are found to b< 
practically free from the sub- 
stance in question. In decidinj 
for what substance the test shal 
be made in the washings one 
must be guided by the reactions 
that are known to take plac< 
.during precipitation and by 
knowledge of what other sub- 
stances may have been present with the element or radica 
being precipitated. This will be dealt with in each specific case 
in the exercises that follow. 

Wash Bottles. — The simplest apparatus for use in washing pre- 
cipitates is shown in Fig. 7. It consists of an ordinary flask 
of convenient size fitted with tubes and rubber stopper as shown 
The tube (a) may be continuous, but the flexibility produced by a 
rubber connection is advantageous in directing the stream oi 
water. For use with hot water the neck of the flask should b< 
wrapped with cork, paper, or heavy twine for the protection oi 
the hands. The usual equipment includes at least two wash 

Fig. 7. 



bottles, one being for hot water and one for cold water. A mis- 
take that is often made is that of allowing the hot water bottle 
to remain over a flame or hot plate when not in use, keeping the 
water boiling meanwhile. Boiling promotes solution of the glass 
of the flask so that the water may become unfit for use. The 
nozzle (b) of the wash bottle may be made in either of two ways. 
A glass tube may be drawn out until a capillary tube is produced 
and then cut off where the bore is such as to give a fine stream 
of proper dimensions. The edges are then rounded. A better 
method is to cut off a piece of tubing and fuse one end until 
it has' contracted to the proper diameter. This 
tube possesses the advantage that it is not easily 
broken by contact with the funnel. 

For use with organic solvents that dissolve 
rubber a wash bottle is used, having a ground 
glass stopper instead of a rubber one and the 
delivery tube is of one piece, omitting the 
rubber connection. 

A wash bottle of any design should be so con- 
structed as to furnish a very fine stream of the 
wash liquid. The stream is directed against the 
upper part of the paper in such a way as to 
wash thoroughly both paper and precipitate. 
It is never directed against the funnel above the 
paper as the precipitate will almost invariably creep up the glass. 

Drying of Precipitates. — Unless the precipitate is to be removed 
from the paper it is generally unnecessary to dry it completely 
before placing in the crucible for ignition. It should be allowed 
to drain thoroughly before removing from the funnel, after which 
the paper may be folded and placed directly in the crucible. 
It sometimes happens that the precipitate is reduced or otherwise 
affected by contact with carbon or reducing gases from the burn- 
ing paper; such precipitates must be largely removed before 
burning the paper and this involves previous drying in order to 
prevent the sticking of the precipitate to other materials with 
which it may come in contact. 

Ovens. — A glance at the pages of the catalogues of dealers 
in chemical apparatus will impress one with the fact that there 
are available many types of ovens for such purposes. These 

Fig. 8. — Incor- 
rect (a) and cor- 
rect (b) forms for 



types need not be described here. It is sufficient to notice tha' 
the oven must possess at least two features: circulation of air 
through the drying chamber in order to remove the water vapor 
as it is formed, and a fairly accurate means of controlling th< 
temperature. Electrically heated ovens are more convenient 
than those heated by gas and, considering the length of life 
are probably not more expensive. There are cases where the 
precipitate is affected by oxygen or carbon dioxide. Such 
a precipitate must be dried in an atmosphere of some gas, such 
as hydrogen or nitrogen, toward which it is chemically inert, anc 
the oven must be provided with means for passing a current of gag 
through it. 

Fig. 9. — Wash bottle for organic solvents. 

Desiccators. — In addition to devices for drying at elevatec 
temperatures we have also those for drying at the ordinary 
temperature of the room. Substances that are not definitely 
and decidedly hygroscopic will lose most of their moisture by 
simple exposure to the air but this is obviously an inconvenient 
procedure and involves much loss of time. Evaporation o 
moisture can be hastened in one or both of two ways withoul 



raising the temperature: (a) by enclosing the moist substance 
in a vessel which also contains a strongly hygroscopic material 
and (b) by keeping the atmosphere which surrounds the material 
practically free from water vapor by mechanical means, such 
as exhausting by means of an air pump or by passing a dry gas 
through the vessel. The vessel known as a " desiccator" is 
of such form that the drying agent, such a sulphuric acid or 
calcium chloride, may be contained in the lower portion, while 
the substance to be dried is placed above. A desiccator for 
drying under reduced pressure is shown in Fig. 10. 

Fig. 10. — Desiccator for drying under reduced pressure. 

In order to understand the action of the various forms of 
desiccators it is necessary to recall the physical law that any 
moist substance, if confined in a vessel at a given temperature 
will continue to lose moisture until a definite pressure of water 
vapor is established, when equilibrium between liquid and vapor 
phases is accomplished. If the pressure of the vapor is reduced 
by extraneous means evaporation begins in an attempt to re- 
establish equilibrium and so long as the vapor pressure is kept 



reduced evaporation continues. It is important to note, however 
that the vapor pressure to be considered is not the total pressure 
(such as that of the atmosphere) but is the partial pressure of the 
vapor of water. The same result is therefore finally accomplishec 
by pumping out the mixture of air and water vapor, by simply dis- 
placing this mixture by means of any gas that has been freed from 
water vapor or by having present and in contact with the confinec 
atmosphere any substance that readily absorbs moisture. Th( 
simplest desiccators involve no principle other than that of con- 
finement with a drying agent. A small desiccator of this descrip- 
tion like Fig. 11 is used by the analyst, not for drying precipitates 

Fig. 11. — Ordinary desiccator. 

but for keeping crucibles, precipitates and small amounts of mate 
rials in a dry atmosphere, previous to weighing. In such cases 
the materials have already been dried and the only function oi 
the desiccator is to prevent the absorption of moisture. 

The desiccator is prepared for use by partly filling the lowei 
chamber with the proper drying agent, a triangle or perforatec 
porcelain plate for supporting small objects being then placec 
upon the shoulder above. The ground-glass joint of the covei 
is lightly smeared with vaseline to make it impervious to air. 
If calcium chloride is used as the drying agent a small piece oi 


sodium hydroxide may be added to keep the atmosphere free 
from carbon dioxide. 

Drying Agents. — The drying agents commonly used in desic- 
cators and for drying gases are sulphuric acid, calcium chloride 
and phosphorus pentoxide. The efficiency of any drying agent 
will depend upon the pressure of water vapor that is maintained 
when equilibrium between the agent and the surrounding atmos- 
phere is established. Every definite chemical substance which 
combines with water maintains a definite tension of water vapor 
at a definite temperature. If this tension is large the compound 
is a poor drying agent. If as large or larger than that of the 
substance to be dried it does not act at all or even adds moisture 
to the latter. If the aqueous tension is exceedingly small the 
substance is a good drying agent. The rapidity of action depends 
also upon the relative surface exposed to the atmosphere. Thus 
a granular or porous solid will absorb moisture more rapidly 
than a liquid, its aqueous tension being the same. 

Phosphorus pentoxide is the most hygroscopic of the three 
substances mentioned above and is therefore the most efficient 
drying agent. In addition to its greater cost its use is also limited 
by the fact that it becomes viscous when moist and that a large 
amount of heat is evolved when it combines with moisture. 
Sulphuric acid ranks next to phosphorus pentoxide in efficiency 
but it is not much used in desiccators that are to be carried about 
the laboratory because of its tendency to splash against crucibles. 
or other articles carried in the desiccator. Calcium chloride, 
although the least efficient of all, is the most convenient for many 
purposes and is generally used for desiccators and for the drying of 
gases in many analytical processes. 

Ignition of Precipitates. — The term " ignition" is used in this 
connection in a sense somewhat beyond its ordinarily accepted 
meaning, since it is applied to the heating to high temperatures of 
substances that are entirely incombustible. The purposes of 
ignition are to destroy the filter, if paper has been used, to expel 
the last traces of moisture and volatile impurities that have not 
been removed by washing and to cause the precipitate to change 
in a definite manner, if a change is to be made. If a paper filter 
has been used it is carefully removed from the funnel by slipping 
up the side. It is then folded as indicated in Fig. 12, the object 



being so to enclose the precipitate that loss is impossible. If 
it is to be dried and removed it is then placed in the oven on a 
cover glass. 

Reducible Precipitates. — The method of treating a precipi- 
tate that is affected by burning paper is as follows : After drying 
completely the paper and cover glass are placed on a sheet of 


Fig. 12. — Folding a filter paper for ignition. 

glazed paper. This paper is to prevent possible loss of traces 
of the precipitate during removal from the filter and should be 
black if the precipitate is white, or white if the precipitate is 
dark in color. The paper is^carefully opened and, by use of a 
spatula of horn, steel or platinum, the precipitate is carefully 
removed and placed in the cover glass. In doing this the pre- 

Paper held by platinum wire for ignition. 

cipitate must be removed as completely as possible but the paper 
must not be scraped, as otherwise fiber will be removed with the 
precipitate. The paper is now re-folded and placed in a crucible 
which has been ignited and weighed, where it is carefully burned. 
The objectionable action of the paper, whereby some of the pre- 
cipitate is changed, is usually that of reduction. It is obvious 



that, if the small amount of precipitate remaining on the paper is 
to escape this action to some extent, the burning of the paper 
must be performed under conditions favorable to vigorous oxida- 
tion. Burning in an atmosphere of pure oxygen would seem to be 
the best remedy but this is usually impracticable. Slow combus- 
tion, at as low a temperature as will support combustion and with 
an excess of air, is easily carried out. One method is that of 
burning on a platinum wire. The paper is tightly rolled and held 
by a heavy platinum wire over the crucible which is placed on the 

Fig. 14. — Crucible inclined for accelerating combustion. 

glazed paper. The paper is fired by touching with the outer edge 
of the burner flame and is allowed to burn slowly, the ash drop- 
ping into the crucible. After treating the partially changed pre- 
cipitate to convert it into the original condition the main portion 
is brushed in, using a small pencil brush of camePs hair, and the 
whole is heated to the desired temperature, then cooled in the 
desiccator and weighed. If burning on the wire is considered 
unnecessary the paper is placed directly in the crucible and ig- 
nited by placing the burner under it. The proper method of 


heating crucibles in order to oxidize the contents is described 
below. The brush already mentioned should be free from loose 
hairs and should be rounded on the end. In using it is drawn 
sidewise in such a manner that very little of the precipitate 
enters the brush itself. 

Oxidation in the Crucible. — The crucible is almost invariably 
heated by means of a naked flame, being supported on a tri- 
angle by means of some kind of stand. When the object is to 
oxidize the paper or precipitate the crucible is placed on its side 
and the cover leaned against it as shown in Fig. 14. The burner 

Fig. 15. — Correct position Fig. 16. — Incorrect position 

of inclined crucible. of inclined crucible. 

is placed under the bottom of the crucible in such a position that, 
the gaseous products of the burner cannot enter the crucible. 
The uprising current of warm air strikes the cover and is deflected 
into the crucible, thus providing an oxidizing atmosphere about 
the paper. If the flame from the burner is applied only enough 
to keep the paper burning the desired condition is attained. 
No harm results if the volatilized combustible material from the 
paper burns with a flame above the crucible. After the paper 
is thoroughly charred the temperature is gradually raised to 
complete the combustion. 

The proper position of the crucible on the triangle is shown 
in Fig. 15. If placed as in Fig. 16 the crucible is liable to fall 
back and it may even sometimes fall through and cause a loss 
of the determination. 

Even in cases where the burning paper has no reducing action 
upon the precipitate it is still desirable to complete the com- 
bustion of the paper at a comparatively low temperature. Crys- 
talline precipitates that are ordinarily regarded as infusible will 
often undergo softening at the sharp corners of the crystals. 


This causes a certain sticking together which results in the enclo- 
sure of a small amount of carbon in such a way as to make its 
oxidation extremely difficult. If the paper containing the 
precipitate is heated to a high temperature at the very beginning 
it is often almost impossible to make it white. One of the best 
examples of this action is in the ignition of magnesium ammo- 
nium phosphate to convert it into magnesium pyrophosphate. 
Premature heating of this substance to very high temperatures 
will frequently result in a black or gray material that cannot be 
whitened by long ignition. 

Decomposition in the Crucible. — After oxidation of the paper 
is completed the temperature is raised in order to volatilize 
completely any volatile impurities that may remain and to cause 
whatever decomposition is desired. Since oxidation is no longer 
an object the crucible is placed in an upright position and the 
cover is placed over the top. This gives an opportunity for the 
flame to bear directly on the bottom of the crucible where the 
precipitate lies. The cover also largely prevents loss of heat 
due to convection currents of air within the crucible. 

Porcelain Crucibles. — The most commonly used crucibles are 
made of a high grade of porcelain. Such a crucible will withstand 
temperatures as high as can be attained by use of the ordinary air 
and gas blast lamp without more than a trifling loss of weight and 
they possess the decided advantage of low cost. They cannot 
be used for fusions because most fluxes, particularly those of a ba- 
sic nature, attack the glaze as well as the porcelain itself. It is 
because of this susceptibility to the action of fusible salts which 
act as fluxes that the life of a porcelain crucible is limited. In 
spite of the most careful washing there will always remain with 
the precipitate traces of fusible materials and these will, in time, 
cause destruction of the glaze lining of the crucible. After the 
crucible is thus roughened it is difficult to clean it and it becomes 
unsuitable for further use. 

Porcelain for Chemical Uses. — The cutting off of imports 
by the war left America in a situation with regard to chemical 
porcelain, similar to that which existed in the case of glassware. 
This country had depended very largely upon foreign porcelain 
for chemical uses and German and Austrian ware had practically 
monopolized the field. American industry has since developed 


an excellent porcelain and the Bureau of Standards has made 
comparative tests 1 of two American makes (Coors and Guernsey) 
one Japanese ware, marked "S. C. P." and two German wares 
(Royal Berlin and Bavarian). The tests included resistance 
to sudden heating and cooling, also treatment with a number of 
acids and bases, in solution and fused. It was concluded that 
the Japanese porcelain was fully equal to the Royal Berlin in 
every respect. American ware suffered in the heat tests, 
although it was stated that later samples had withstood the 
test satisfactorily. 

Platinum Crucibles. — Crucibles of platinum are very desirable 
for ignition and are almost essential for fusions. Platinum fuses 
at 1770° and does not soften enough to preclude its use much 
below this temperature. It resists the action of all single acids 
if these are pure. It is readily dissolved by solutions of chlorine 
and, on this account, by aqua regia. Even " chemically pure" 
hydrochloric acid often contains traces of chlorine and will 
slightly attack platinum. Platinum easily alloys with most 
metals and so should not be heated in contact with compounds of 
easily reducible metals, particularly if carbon be present. When 
heated for a long time in contact with carbon it slowly dissolves 
this and becomes brittle because of the presence of the carbide 
of platinum. This is noticed when the crucible is heated in a 
reducing flame and on this account it is necessary carefully to 
adjust the flame so that the tip of the inner cone is below (not 
against) the bottom of the crucible. A flame showing yellow 
must never be used. Platinum crucibles cannot be heated in 
contact with alkali hydroxides although they are not attacked 
by alkali carbonates. Compounds of phosphorus are reduced 
to some extent by heating with carbon and the phosphorus readily 
combines with platinum, causing destruction of the crucible. 
Ignition of phosphates thus requires especial care if this actioi 
is to be prevented. 

Formerly platinum crucibles and dishes contained a smal 
percentage of iridium, added for the purpose of hardening the 
metal and giving greater resistance to mechanical wear. Such 
an alloy, however, is more readily attacked by reagents and i 
is more volatile and practically pure platinum is now used almos 

1 Bur. Stand. Tech. Paper, 105. 


exclusively. It was also formerly the custom to form crucibles 
and dishes by spinning the metal. This gives an article with a 
fine surface as apparent to the naked eye but it results in the 
formation of minute surface cracks and scales and disintegration 
is aided by this process. The crucible of the present day is 
given its surface by hammering. This results in a more uneven 
surface but the particles of metal are firmly welded together 
and the hammered crucible has a life appreciably longer than 
that of the spun crucible. 

Very great differences are observed in the platinum ware as 
furnished by different refiners and manufacturers and many 
troubles are experienced in the use of such ware for accurate 
work. The continued advance in the cost of platinum has made 
it correspondingly difficult to obtain satisfactory ware and a com- 
mittee was appointed by the American Chemical Society to 
investigate the subject and to make recommendations to the 
Society. Two reports of this committee have been made. 1 
An investigation of the quality of platinum ware has been made 
also by the Bureau of Standards. 2 

In these reports the objections to inferior platinum ware are 
summarized as follows: 

1. Undue loss of weight on ignition. 

2. Undue loss of weight on acid treatment, especially after strong ignition. 

3. Unsightly appearance of the surface of the ware after strong ignition, 
especially in the early stages of heating. 

4. Adherence of dishes and crucibles to triangles. 

5. Basicity of the surface of the ware after strong ignition. 

6. Blistering. 

7. Development of cracks after continued heating. 

So far as these matters have been studied the causes are stated 
as follows : 

Loss on Ignition is chiefly due to the presence of iridium which 
is added to stiffen and harden the ware. The loss at 1200° 
is more than twice as great as at 1100° for nearly all ware used 
in the experiments. 

Appearance of the Surface after Ignition. — Even with good ware 
the surface is frosted after strong ignition. This is due to sur- 

1 J. Ind. Eng. Chem., 3, 686 (1911); Ibid., 6, 512 (1914). 

2 Bur. Stand. Tech. Paper, 254. 


face crystallization, but the crystallization should be fine and 
evenly distributed.. Poor ware becomes, when ignited, unevenly 
coated with a whitish layer and with brown stains, the latter 
being due to small amounts of iron. Sometimes the entire inner 
surface becomes stained brown from the latter cause. 

Adherence of Crucibles and Dishes to Triangles is caused by 
welding of the platinum of the ware with the triangle. This 
cannot well be overcome but is not serious at ordinary blast-lamp 
temperatures unless platinum triangles are used. 

Basicity of the Surface after Ignition is due to the presence of 
traces of calcium alloyed with the platinum. Calcium oxide is 
produced when the crucible or dish is heated and this is made 
evident by testing with moist red litmus paper. 

Blistering is found to occur appreciably only with ware of 
earlier manufacture and is not now an important objection. 

Cracking. — No cause has been determined for this form of 

Specifications. — It is recommended that purchasers specify 
that platinum ware must show no marked uneven discoloration 
on heating, must give no test for iron after heating for two 
hours and that the rate of loss per hour at 1100° over a period 
of four hours shall not exceed 0.2 mg and that 5 percent of 
rhodium be substituted for iridium as a hardening agent. 

While platinum possesses properties that make it an extremely 
valuable metal for the chemist, its use is greatly limited by its 
present high cost. Because of this fact, available substitute 
are always in demand. 

Care of Platinum. — Platinum ware will deteriorate rapidl; 
unless precautions are taken in its use and care. 

1. Handle carefully to avoid bending. Use clean crucib] 
tongs and do not allow the tongs to come into contact with fused 
materials within the crucibles or dishes. 

2. For cleaning apply the appropriate solvent, according tc 
the nature of the material to be removed. Chromic acid is 
suitable for removing organic matter, hydrochloric or nitri< 
acids for insoluble carbonates or metallic oxides and fusing wi1 
sodium carbonate for silica or silicates, or with sodium pyr< 
sulphate for such metals or metallic oxides as resist the actic 
of acids. 


3. Do not heat platinum in contact with the inner cone of the 
laboratory burner, as brittleness results from such exposure. 

4. Do not heat compounds of lead, tin, bismuth, arsenic, 
antimony or zinc in contact with platinum. Reduction may 
occur, the reduced metal alloying with the platinum. 

5. Do not attempt to remove fusions from platinum crucibles 
or dishes by means of files, glass rods or other hard tools. Use a 
rubber-tipped rod or solvents. 

6. Dull surfaces should be polished lightly with wet emery 
slime or fine carborundum. 

Platinum Substitutes. — The increasing scarcity of platinum 
has made the introduction of substitutes a practical necessity. 
While it is true that pure platinum possesses certain properties 
that cannot be duplicated by any other metal or alloy, yet certain 
alloys have been found to be suitable for making into crucibles 
and dishes that will serve for many of the operations of 
the analytical laboratory, in place of the platinum that has been 
in use. Two of these will be mentioned. 

Palau. — This is a trade name for an alloy containing about 
80 percent gold and 20 percent palladium. The alloy is some- 
what darker in color than platinum but resembles it otherwise. 
The melting point is 1370°. This is higher than the melting 
point of gold but it is 400° lower than that of platinum. 

The Bureau of Standards found that a crucible of this material 
was comparatively free from iron and that its loss on heating 
to 1200° was less than that of a platinum crucible containing 
2.4 percent of iridium. The resistance to acids, sodium hydrox- 
ide and ferric chloride solutions and to fused sodium carbonate 
is comparable with that of platinum. The ware is therefore 
suitable for fusions with sodium carbonate but it is decidedly 
attacked by fused pyrosulphates. Crucibles of palau could 
not be used for work at high temperatures on account of its 
relatively low melting point. 

Rhotanium. — Fahrenwald 1 described a series of six gold- 
palladium alloys, some of which compare very favorably with 
I platinum in many of the essential properties. Unfortunately 
the percentage composition of these alloys is not definitely 
stated but it is to be inferred that gold ranges from 60 to 90 

1 J. Ind. Eng. Chem., 9, 590 (1917). 


percent and that rhodium is contained in some members of the 
series. The melting points range from 1150° to 1450° for the 
series. Presumably the alloy having the highest melting point 
contains the highest proportion of palladium but increasing 
palladium also increases the rate of attack by some reagents. 
It would appear that rhotanium of properly chosen composition 
might well replace platinum for most analytical purposes, 
excluding processes where hot concentrated nitric acid is used. 
The alloys cannot be used as anodes in electroanalysis. 

Gold shows about the same resistance to the action of reagents 
as does platinum, but its relatively low melting-point (1035°) 
makes it unsuitable for crucibles or other articles that must be 
strongly heated. 

Silica. — Crucibles have recently been made of fused silica. 
This resists the action of chemicals better than does glass and 
the melting-point is such that no ordinary air-gas flame will fuse 
the crucible. Pure quartz fuses at 1600° and amorphous silica 
at 1750°-1780°. Moreover the coefficient of expansion is very 
small (5.4 X10~ 7 at temperatures between 0° and 1000°) so that 
sudden heating or cooling does not cause cracking. Silica cru- 
cibles have not yet taken the place of those of porcelain, largely 
because of their higher cost. 

Alundum. — Recently a highly refractory form of aluminium 
oxide has been utilized for crucibles. This is commercially 
known as " Alundum." Alundum does not fuse below 2050° 
does not much soften at 1775° and its coefficient of expansion is 
7.8 X10- 6 . 1 

Triangles. — The discussion concerning the relative merits o 
various materials used for crucibles will apply equally well to 
those used for supporting triangles. Porcelain is the cheapes 
serviceable material and will do for supporting crucibles of any 
other material as well as those of porcelain itself. The familial 
"pipe stem" triangle is constructed of three tubes of refractory 
unglazed porcelain, held together by a frame of iron wire (Fig 
17). The chief objection to this form lies in the fact that the 
relatively large tubes, lying on three sides of the crucible, obstruct 
the flame and cause a very noticeable decrease in efficiency. An 
improved form is shown in Fig. 18. The projections serve to sup- 

1 Saunders: Trans. Am. Electrochem. Soc, 19, 333 (1911). 


port the crucible and allow much better contact with the flame. 
Any porcelain triangle becomes practically useless if cracked 
because the supporting iron frame is thus exposed to the flame 
and soon oxidizes and breaks. 

Fig. 17. — Common "pipe stem" triangle. 

Platinum triangles offer the great advantage of long life and 
they also obstruct the flame very little on account of the small 
•size of the framework. Here again the high cost of platinum pre- 
cludes its extensive use. A common form of platinum triangle 

Fig. 18. — Triangle with projections. 

is shown in Fig. 19. A less expensive form can be constructed 
by stretching heavy platinum wire on an iron framework, like 
Fig. 20 or Fig. 21. This is not very satisfactory^because the 
weight'ofthe crucible causes lengthening and sagging of the wire 


at high temperatures. The screws shown in Fig. 21 can be used in 
taking up this slack but the wire eventually weakens and breaks. 
Alloys of nickel and chromium having high fusing points anc 
considerable resistance to oxidation at high temperatures have 
been adapted to the construction of chemists' triangles. The 

Fig. 19. — Heavy platinum triangle. 

proportion of the two elements may be varied somewhat but 
the presence of iron, even in small amounts, makes .the alloy 
oxidizable in the flame and the resulting oxide combines with the 
silicious glaze of porcelain crucibles, thus changing the weights 
of the latter. As it is becoming increasingly difficult to obtaii 

Fia. 20. — Triangle of platinum wire. 

nickel-chromium triangles that are free from iron it is likely 
that their use must soon be restricted to qualitative work or other 
work at not very high temperatures. 

Probably the best triangle that is now obtainable at a moderate 
cost is one of the pipe-stem form, made from silica tubes on i 



frame of nickel wire. These may be used with either porcelain 
or platinum crucibles and they are not easily cracked by sudden 
heating or cooling. 

Fig. 21. — Another form of platinum wire triangle. 

Burners. — The burner that is to be used by the analyst may be 
anything from the cheapest and simplest burner of the Bunsen 
type to the most expensive and complicated burner obtainable. 
The purchaser has his choice and probably certain advantages 

Alloy triangle. 

are possessed by each burner. The only feature that is really 
essential is independent regulation of air and gas supply. The 
requirements are quite different in different cases and the analyst 
must have at his disposal all kinds of flame, from the yellow illumi- 



nating flame to the most intensely hot and oxidizing flame, and 
he requires very small and very large flames of each class. In order 
to obtain this variety of flame there must be some method of 
regulating the gas supply without changing the pressure at the 
gas valve, since this also changes the amount of air drawn in at 
the mixer. The simplest form of Bunsen burner does not permit 
this gas regulation without unscrewing the upper tube and chang- 
ing the gas jet by the use of pliers. Such regulation is not pos- 
sible in practice. 


Two forms of adjustable burners are shown in Figs. 23 and 24. 
In the E. and A. form the gas flow is regulated by the large 
milled disc near the bottom of the burner and the air supply 
by the ring which screws up and down at the base of the burner 
tube. In the Teclu burner (Fig. 24) the gas is controlled by the 
screw on the side of the base while the disc at the bottom of the 
cone controls the air supply. 

It is important to note that in both of these burners the regu- 
lation of gas flow is not accomplished by altering the pressure 


under which it is delivered but by changing the size of the orifice 
in the burner. The maximum pressure is thus used at all times 
and the result is a better mixture of gas with air than is obtainable 
by regulating the gas cock of the supply line. ] 

A very common error on the part of students lies in careless- 
ness with regard to the regulation of flames. If a relatively 
cool flame is required and if a deposit of carbon is not objection- 
able the air should be excluded from the mixer. If, on the other 
hand, the highest efficiency of the burner is desired, careful 
regulation of the air and gas is necessary. The inner blue cone 
should be well defined and it should not show a yellow tip. If 
air is admitted more than that required to completely burn the 
gas with production of a blue flame, the result is a roaring and 
fluttering flame. This means that more air is being admitted 
than can be used and this air, in being heated by the flame, 
lowers the temperature of the latter. 

Blast Lamps. — The blast lamp is used for obtaining tempera- 
tures higher than are attainable by use of the ordinary burner. In 
this burner the gas is used in large quantities and air is delivered 
under pressure so that a flame is produced more intensely hot 
than the slower burning flame of the common burner. Where 
extremely high temperatures are necessary the oxyhydrogen blow- 
pipe is used but this is rarely the case for analytical operations. 

Meker Burner. — A somewhat radical departure from the older 
types is found in the Meker burner. This is shown in section 
in Fig. 25. The air is drawn in through several holes in 
the base of the tube. The delivery of the gas under pressure 
into the inverted cone which forms the burner tube causes a 
greater reduction of pressure within the tube than is the case 
with burners having cylindrical tubes. The result is a greater 
inflow of air, making possible the combustion of a greater amount 
of gas in a given space, and also more complete mixing of gas 
and air. 

The nickel grid through which the mixture flows at the top of 
the burner causes the gas to burn exactly as though each mesh 
were a small individual burner. The tip of the inner reducing 
cone of each small flame is usually about one millimeter above 
the top of the burner and, as all of the small flames unite to form 
one large one, the result is a highly concentrated flame, every 



part of which is oxidizing in character except a zone of about one 
millimeter in depth, immediately above the top of the burner. 
This is a distinct advantage, especially in heating platinum 
articles, since platinum is easily damaged by heating in a re- 
ducing flame. 

Fig. 25. — Section of Meker burner. 

The flame of the M6ker burner is nearly as hot as that of the 
ordinary blast lamp using the same gas and it may be substi- 
tuted for the blast lamp in many cases. There is also a Meker 
blast lamp, similar in construction to the one already described 
but using air under pressure. 

Fusion. — For the purposes of quantitative analysis the fusion 
of materials is almost always accomplished with the end in view of 
producing more soluble substances through the interaction of 


an added agent, called a flux, and the refractory material. 
For instance, most of the natural silicates are practically insoluble 
in water and all ordinary reagents and therefore they cannot 
be analyzed by ordinary methods. % By a preliminary heating 
to a high temperature in contact with a basic substance like sodium 
carbonate, a fusible mixture of new compounds is formed and 
these will, for the most part, be soluble in water and hydro- 
chloric acid so that the solution may be submitted to precipita- 
tion and filtration processes for the separation and determination 
of the elements. Similarly, refractory and insoluble metallic 
oxides may be heated with sodium pyrosulphate with the for- 
mation of a fused mass consisting of soluble sulphates of 
the metals. 

The necessary qualities of any useful flux are (1) that it must 
be of such a nature as to be capable of reacting with the re- 
fractory body when heated with it and (2) that the resulting 
compounds shall fuse at the prevailing temperature. To these 
the analyst adds a third requisite: (3) that the resulting com- 
pounds shall be soluble in water or in the laboratory reagents. 
The first condition is met by choosing as the flux a substance 
of opposite nature to that of the refractory sample. That 
is, if the latter is of an acid nature (as silica and polysilicates) 
the flux should be basic, and conversely. 

No general statement can be made with regard to the relative 
fusibility of various compounds, as based upon the chemical 
composition of these compounds. It may be noted that re- 
fractory silicates are usually made more readily fusible by reduc- 
ing the ratio of silica to metal oxide by introducing more metals, 
and particularly by the introduction of the alkali metals. Both 
of these points are made by using alkali metal carbonates as 
fluxes, since the net result of the reaction at high temperatures 
is to expel carbon dioxide and to combine the alkali metal 
oxide with the refractory silicate. This will explain why these 
carbonates are almost always chosen as fluxes for silicates. 
A reaction such as the following may occur when orthoclase 
is fused with sodium carbonate: 

2KAlSi 3 8 +5Na 2 C03-^K 2 Si03-l-5Na 2 Si03 
+2NaA10 2 +6C0 2 , 


a more or less complicated mixture of aluminates and silicate 
of the alkali metals being formed. 

Basic Fluxes. — Sodium carbonate, potassium carbonate anc 
the double sodium-potassium carbonate are the most importan 
of the basic fluxes that are used for analytical purposes. Thes 
are used chiefly for fusion with silica and the refractory silicates 
Such fluxes as calcium oxide, used for fluxing silicates in th 
blast furnace for iron, are of little use for analytical purposes 
partly because the resulting compounds are not soluble an 
partly because metals that are to be determined in the sampl 
are introduced by the use of such materials. 

Acid Fluxes. — Fluxes of an acid nature are valuable chief! 
for forming fusible, soluble compounds when heated with metalli 
oxides or salts that are over-saturated with metallic oxides 
The most useful of such fluxes are the pyrosulphates and the 
biborates of sodium and potassium. 

Acid sulphates are often used instead of pyrosulphates. Whe 
the former are heated they give off water and they are completel 
converted into pyrosulphates by heating to higher temperatures 

2NaHS0 4 ^Na 2 S 2 7 +H 2 0. 

Because of the excess of sulphur trioxide in the pyrosulphate, 
this readily reacts with metallic oxides when heated with them: 

Fe 2 03+3Na2S207-+3Na 2 S04+Fe 2 (S04)3. 

The biborates likewise combine with metallic oxides because 
of their excess of boric anhydride. 

Fe 2 3 +3Na 2 B407->2Fe(B0 2 )3+6NaB02. 

Weighing. — The balance is the only kind of weighing apparatus 
that is independent of the numerical value of the force of gravity 
so that it actually measures mass and not weight. In its simplesi 
form it consists of a beam, supported at its middle point on i 
fulcrum, having suspended at its ends, at equal distances fron 
the middle fulcrum, two platforms for holding, respectively 
the objects to be weighed and the weights. It is thus seen tc 
be an apparatus for comparing masses and if it is mechanically 
perfect and if the standard weights are correct, then the mas! 
of the object is given by the mass of the weights that counter 
balance it when the balance is at equilibrium. 



Description of Balance. — The balance as used by the analyst is 
a much more sensitive and carefully built piece of apparatus than 
ithe ordinary balance and must be capable of giving weights that 
I are correct to within 0.0001 gm. In order to reduce the friction 
of the bearings to the lowest possible value these are made in the 
form of a knife edge supported by a polished block, the material 
of the bearings being some substance that is quite hard, this 
usually being agate. The method of weighing is to place the 

Fig. 26. — Analytical balance. 

object to be weighed in one pan and to add weights to the other 
until equilibrium is attained, judging this condition by observing 
when a pointer, attached to the beam, swings equal distances to 
the right and left of an experimentally determined "zero point" 
on a scale. The whole balance is enclosed in a glass case in order 
to exclude dust and to prevent the interference of air currents. 
The case is provided with levelling screws and a spirit level or 
plumb bob. 

To prevent needless wear on the knife edges the beam and pans 
should be provided with suitable rests so that the knife edges may 



not be in contact with their bearings when the balance is not bein 
used. These rests are necessary also for the proper control of th 
action of the balance, as will be noticed when exercises with th 
balance are described. The rests for the beam and for the pa 
supports are usually operated by one piece of mechanism. Th 
exact construction of this varies in different balances but th 
action of the beam rests belongs to one of two classes, the vertica 
and the circular. In the first (Fig. 27) the rests move vertically 
upward. The chief defect of this action is the fact that if th 
beam is caught at any position except a horizontal one the knif 

Fig. 27. — Beam rests with vertical action. 

edges are caused to slide upon their bearings, causing unnecessary 
wear. In the circular action (Fig. 28) the arms of the arrests hav 
the same length as the arms of the beam and they move about th 
same axis. No matter in what position the beam is caught o 
how sudden the motion of arrest, no damage results to the knif 
edges. All designs of mechanism for this purpose include device 
for automatically placing beam and pan bearings in their prope 
position in case they have been accidentally twisted out of posi 
tion. These are shown at (a), Fig. 28. The rests under the pan 
are not for the purpose of lifting knife edges from their bearings 
but merely for steadying the pans and for controlling the move 
ments of the balance when it is in use. In some balances the* 



are operated by the same mechanism that operates the other 
rests, while in others a separate knob is provided. 

Sensibility. — The sensibility of the balance is stated in terms of 
the displacement of the beam by a given excess of load in one pan. 
More specifically, it is the number of scale divisions of displace- 
ment of zero point by an excess of 1 mg in one pan. 

The sensibility is affected by a number of factors. The bal- 
ance, in order to possess stability, must have the center of gravity 
of the moving parts slightly below the point of the middle knife 
edge. This distance is one of the determining factors of the sen- 

Fig. 28. — Beam rests with circular action. 

sibility. If large, the sensibility must be relatively small, since a 
given displacement of the zero point will involve a relatively 
large displacement of the center of gravity and will, in conse- 
quence, require a greater difference in load. Every balance has 
some provision for arbitrarily altering the sensibility by altering 
the location of the center of gravity. The most common device 
is a weight that can be moved up and down the pointer. 

The three knife-edge bearings must lie in the same plane, as 
well as parallel to each other. Since the pans swing freely upon 
their own bearings, the whole load of the pans is applied at these 
points. If the plane of these bearings were below that of the 
middle bearing it could easily be that the center of gravity 
would lie between these planes and then an increase in load 



would lower the center of gravity with reference to the centra 
bearing and thus decrease the sensibility. If their plane wer 
above that of the middle bearing an increase in load woul 
raise the center of gravity to a point above the middle poin 
of support and give instability to the balance so that if dis 
placed from its normal horizontal position the beam would no 
return. When the three bearings lie in the same plane an in 
crease in load will raise the center of gravity but can never rais 
it to the level of the middle knife edge. The above methoc 
of reasoning supposes that the balance beam is perfectly rigi< 
a property that is never attained in practice. Increase in loa( 

Fig. 29. — Illustrating the principle of moments. 

therefore, does actually cause decrease in sensibility because 
the beam is somewhat distorted, causing the center of gravity 
to be lowered. In order to combine great strength with lightness 
of weight and so minimize the distortion of the beam, makers 
have tried many designs and many alloys in the manufacture 
of balance beams. 

Another property that has a large influence upon the sensibilit 
is the length of the arms of the beam. It is a well known princip 
of physics that when a balance is in equilibrium the produ< 
of the weight of one side into the length of the correspondin 
arm must equal the product of the other weight into. the lengt 
of its arm. This is the " principle of moments." In Fig. 2 C 
aw = a'w' '. The greater the inequality of the statical moment 
(aw and a'w') when an excess of weight lies on one side the greate 


will be the displacement of zero point and this inequality will 
be greater when a and a' .(the lengths of the arms) are large. 
Lengthening the arms also causes slower swinging so that sensi- 
bility gained by this means results in loss of time in weighing. 
This fact sets a practical limit to the length of the beam. 

The following properties have been discussed as having an 
influence upon the sensibility of the balance: 1. Distance between 
the center of gravity and point of support. 2. Coincidence 
of the planes of the three bearings. 3. Length of the arms 
of the beam. 4. Reduction of friction to a minimum by finely 
ground knife edges. In addition to these should be mentioned 
the weight of the beam. A heavy beam makes a balance that 
is less sensitive than a balance having a light one. 

In addition to those features which affect the sensibility, 
certain others are essential if the balance is to weigh accurately. 
It is extremely difficult to construct a beam having absolutely 
the same distance between the central knife edge and the two end 
ones. Obviously any difference involves a slight difference in 
the two weights required to bring the balance to equilibrium. 
From the equation aw = a'w', if a 9^ a' then w^w' and the weight 
of the substance which is being weighed is not the same as that 
of the weights which counterbalance it. The relative lengths of 
the arms must be determined for each balance and the observed 
weights corrected if these lengths are appreciably different. 
Even if the discrepancy in lengths is sufficiently small to be 
negligible it may be magnified by a change in temperature or by 
a change in load, if the beam is not absolutely uniform in material 
and structure. This last condition is impossible of attainment 
in practice. It will be made clear in the course of the work in 
gravimetric analysis that it is not the absolute weight that is 
important in most cases but only relative weights because the 
object of quantitative analysis is to determine the proportionate 
parts of the constituents of a compound or mixture. From the 
equation aw — a'w^ if a and a! are not equal the error in w' 
bears in all cases a definite ratio to w' '. 

While the balance gives the true mass of the object and is in- 
dependent of the magnitude of the force of gravity, this expres- 
sion is true only if the buoyancy of the air acts upon weights and 
objects alike. This can be the case only when the density of 



weights is the same as the density of objects, a condition that i 
not fulfilled in the majority of cases. A correction must there 
fore be introduced in such cases in order to find the true weigh 
of the object. The amount of correction is negligible in gravi 
metric analysis and becomes serious only when the total weigh 
is considerable. The method for applying this correction wi 
be explained in the section dealing with volumetric analysis 
(See p. 180.) 

Weights. — Sets of analytical weights as purchased frequently 
include weights as small as 1 mg. These are rarely usee 
because the balance provides a more convenient method fo 
making the final adjustment, in the form of a " rider" or sma 
weight of fine platinum or aluminium wire which may be shifte< 
to various positions on the beam. The manner in which th 

Fig. 30. — Balance beam. 

beam is divided varies with the balances of different manufac 
turers. The lowest subdivision should be at most 0.1 mg. Th 
weight of the rider will depend upon the manner of numbering th 
milligram divisions and the weight will be represented 'by th 
number which is directly over the terminal knife edge. This 
because when the rider is placed on this division it is essentially th 
same as though it were in the pan below. To correctly indicat 
weights it must then weigh the number of milligrams indicatec 
by this division, which may be 5, 6, 10, 12, or any other number 
A mechanism is provided for lifting and adjusting the rider with 
out opening the balance case. This may be a simple sliding hook 
or an elaborate carrier, such as are found in more expensive 



balances. It is essential that the carrier be capable of quickly 
and easily shifting the rider without danger of throwing it from 
the beam. 

The Chain Rider.- — The "Chainomatic" balance entirely dis- 
penses with a separate rider. One end of a small gold chain is 
permanently attached to the balance beam. The other end of 
this chain is fastened to a hook which may be moved up and 
down a scale, this action being controlled by a knob outside the 
balance case. Movement of the hook on the scale varies in a 
definite manner the length of side of the loop which is supported 



2(n fioi loX 


Fig. 31. — Usual form of a set of analytical weights. 

by the beam and this may be adjusted while the beam is in 
\ motion. This is a distinct advance in balance design, although 
this improvement adds considerably to the cost of the balance. 
Every balance is rated for a certain maximum load, it being 
understood that this is the load for each pan and not the total 
load. The normal load is fixed by the strength of the knife edges 
and by the capacity of the beam to resits deformation under stress. 
If the knife edges are short and ground to exceeding fineness they 
are injured more readily by a load than if they are slightly more 
blunt. If the beam is overloaded it is temporarily deformed to 


such an extent that there is an unusual loss of sensibility, due 
the excessive lowering of the center of gravity. It is thus evi 
dent that the weighing of heavy objects requires correspondingl 
more sturdy balances and these will, of course, be less sensitive 
The usual form of a set of metric weights is shown in Fig. 3 
The largest weight should not be heavier than the maximum 
load for which the balance is rated and the least weight shoul 
be such that, used in conjunction with the rider, 10 mg ma 
be made up. The larger weights are constructed of brass c 
bronze, plated with platinum or gold to prevent corrosion. Th 
fractional pieces are of platinum in the better sets or of aluminiur 
in the cheaper ones. Even with weights plated with platinui 
or gold it is comparatively easy to damage the surface by carele 
handling or by allowing chemicals to touch the weights. 

Printed directions for setting up always accompany a ne 
balance. The following rules deal only with the balance set u 
and ready for use. 

Cleanliness. — The pans, beam, bearings and all other parts 
inside the glass case must at all times be kept free from dust and 
chemicals. A camel's hair brush 1 inch wide should be provided 
for this purpose. It is not permissible to weigh any soluble mate 
rial in direct contact with the pans because some of this will 
variably stick to the pan and eventually it may cause corrosio 
Volatile acids must never be brought inside the balance case u 
less securely stoppered in an air-tight container. 

Adjustment. — The balance is levelled by means of the screws 
provided for that purpose. Examination is made to determine 
whether the knife edges are in the proper position with respect to 
their bearings. The pan rests are released to determine whether 
the pans hang vertically from the stirrups, or whether they swing 
horizontally when released. If so this swinging is stopped by 
momentarily touching the pans by the pan rests, repeating th 
operation until the pans hang quietly upon release. 

To Set the Beam in Motion. — Various methods are used f< 
starting the oscillation of the balance about the central bearin 
One pan may be lightly touched with a small camel's hair brush 
This is not an easy process to carry out properly because it i 
difficult to control the impulse given to the pan. Anoth 
method is to raise the door and fan one of the pans slightly wit 



I the hand. This is open to the same objection as is the first 

I method and, in addition, it defeats the primary aim of the glass 

case which is to prevent the interference of air currents. Even 

the slight current started by the hand does not at once die and it 

must be a disturbing influence for some time after the balance case 

is closed. A better method than either of the above mentioned is 

| to lower the rider to the beam just before releasing the latter, then 

I to catch up the rider with the carrier after releasing and allowing 

I the proper start. A short practice will enable the operator to 

| give just the desired impulse to the beam to make the pointer 

swing over from five to ten divisions on either side of the zero 

! of the scale. 

The rests should be so adjusted that the three knife edges are 
lifted from their bearings when the rests are raised, but the distance 
between the edge and bearing should be barely, perceptible. If 
this distance is unduly large the shock to the delicate knife edge is 
so great that this edge is soon dulled or chipped with a consequent 
loss of sensibility. 

To Determine the Zero Point. — The zero point may be de- 
fined as that point on the scale at which the pointer would eventually 
\come to rest from swinging over the scale. It is never observed by 
allowing the pointer to actually come to rest because such in- 
fluences as minute air currents would either prevent this con- 
summation entirely or would cause the observation of a fictitious 
zero point. The effect of these influences is counteracted by 
allowing the pointer to swing a number of times to the right and to 
| the left, taking the average of the indications. Proceed as 

If the balance operates all of its rests by one mechanism care- 
fully lower these rests and set the beam in motion as directed 
above. If the pan supports are controlled by a separate button 
[lower the beam and stirrup rests first, then the pan rests. Allow 
the pointer to swing three or four times in each direction and re- 
cord the number of scale divisions over which it swings, taking 
[the last reading on the same side as the first. Record in two 
(columns and take the average of each column. Subtract the less 
iaverage from the greater and divide the remainder by 2. This 
feives the zero point if the proper direction is assigned to it. 

















Average 7 . 25 

Average 5.15 


1.05. Therefore the zero point is 1 division to 

the left of the zero of the scale. Two methods of procedure 
are now open to the operator. He may either make his weighings 
with reference to this observed zero point or he may adjust 
the balance so that the observed zero point is the actual zero of 
the scale, using for this purpose the small screws provided on the 
ends of the beam. The first method is preferable for the beginner 
because any attempt to change the adjustment will probably 
result in more serious derangement. After skill has been gained 
by practice time will be gained and much calculation will be 
saved if the observed zero point is adjusted to coincide with the 
ideal zero of the scale. A method for making a close approxi- 
mation of the zero point without resorting to calculations on 
paper will be explained in a later paragraph. The zero point 
changes and it must be determined each day, or more often if 

To Determine the Sensibility. — The sensibility in the case 
of the analytical balance has already been defined as the number 
of scale divisions that the zero point is displaced by an excess in 
weight of 1 mg on one side. That the sensibility varies with 
the total load has already been explained. To determine the 
sensibility with zero load, first determine the zero point of 
the balance. Place the rider on the beam at the division marked 
1 and redetermine the zero point. The difference is the sensi- 
bility. Determine the sensibility when both pans are loaded 
with 5, 10, 25 and 50 gm, respectively. Record the results on' 
a card and place this in the balance case for future reference. 


To Determine the Relative Lengths of the Arms. — Weigh 
a small object, such as a crucible, placing it on the left pan and 
the weights on the right, then place the object on the right pan 
and the weights on the left. If the arms of the balance are of 
unequal length these weights will not be the same. Let W = 
the true weight of the object, IF' = the sum of the weights when 
the object is on the left pan, a = the weight added to W when the 
object is placed on the right pan, r = the length of the right arm 
and I = that of the left arm. From the principle of moments 

W'r = Wl 
Tf r =(Tf , +a)Z 
WW'r^WW+ajl 2 
W'r 2 =(W'+a)l 2 
r*_W' + a 
P~ W 

r jW ; + a _ L,_a_. 

This is the proper method for testing the equality of arms. 

■ T 

If the arms are equal then a = and 7 = 1. If a is negative, 


|y < 1 and the right arm is therefore shorter than the left, while 

if a is positive ~i>l and the right arm is the longer. As has 

already been explained, the question of equality has little or no 
Importance for purely analytical work. Inasmuch as the chemist 
Lias uses for his balance in other lines of work it should be tested. 
To Weigh. — With all of the rests raised the object to be 
(weighed is placed on the left pan by means of the crucible tongs 
pr by some other method that avoids contact with the fingers, 
jrhe pan rests are lowered and raised momentarily until the 
bans stop swinging on their bearings, then these rests are lowered 
[md fastened down unless all of the rests are governed by one 
mechanism. By means of the weight forceps place one of the 
^eights, judged to be somewhat heavier than the object, on the 
i'ight pan. Lower the beam rest slowly until the pointer just 
starts to move to the right or left. If to the right the weight is 
ioo light, if to the left it is too heavy. If it is too light raise the 



beam rest and exchange the weight for the next heavier one and 
repeat the trial until the first weight is found that is too heavy. 
Remove this and replace by the next lighter one. Continue the 
addition of weights, trying after each addition and adding always 
the next consecutive weight lighter than the one used in the 
preceding trial. When the weights below 1 gm (milligram 
pieces) are reached the difference between the loads on the two 
pans is so small that the pan rests will now readily control the 
movements of the balance. If these rests are operated by a 
separate knob they are then raised and the beam and stirrup 
rests lowered, and the process of trial is continued until the range 
covered by the rider is reached. The balance case is now closed, 
and the pans again steadied if they have shown a tendency 
toward lateral swinging. Make the preliminary rider trials 
by placing the rider on the division estimated to be nearest th 
proper one, slightly releasing the pan rests until the pointer start 
to move. Arrest the motion and move the rider by whole milli- 
gram divisions to the right or left, as may be' required, until the 
milligram nearest to the correct weight is reached. Estimate 
the proper fraction of a milligram, set the beam in motion as 
already directed and drop the rider on the estimated division, ji 
By observing the swinging determine the zero point and calcu- : 
late from this and the sensibility with this load what change 
should be made in the position of the rider to bring the balance 
to equilibrium on the zero as determined with no load. Repeat ' 
the trial with the rider on this calculated position and shift if 
necessary to obtain the exact weight. Raise all of the rests and 
read the weights as follows: Observe the empty places in the 
box. If the weights have been systematically placed in the box, 
none being removed except those on the pan, the empty places 
will give the correct weight. Record this weight in the data 
book. Confirm by counting the weights as they lie on the pan. 
Reconfirm by counting as they are removed and replaced in the 
box. This gives three readings and if these are carefully made 
a mistake is practically impossible. In recording the weights 
mental addition may be made if they are taken in order, pro- 
ceeding from the larger ones to the smaller ones. This is because, 
as the sets of weights are made, no one order of digits can total 
more than 9. Each order can be mentally added and recorded 


I with the certainty that no other order will change the one read. 
| Thus, if there are on the pan the following pieces : 10 gm, 5 gm, 

I I gm, 200 mg, 100 mg, 50 mg, 20 mg, 10 mg, and 5 mg, and on 
I the rider 3.2 mg, we should read and record thus: Of whole 
, grams 16, of tenths (100 mg) 3, of hundredths 8, of thousandths 
(8, of ten-thousandths 2. Writing in the same order we should 
[have 16.3882 gm. 

Use of Rests. — The beam and stirrup rests must be used when 
j changing weights heavier than 1 gm for two reasons : First, the 
shock of the weight against the pan must not be allowed to 
I communicate itself to the knife edges when on their bearings 
jand second, the pan rests are held up by a spring that will not 
support an excess load of more than 1 gm in one pan. In other 
I words, these rests would not control the balance at this point 
in the experiment. When the milligram pieces are being ex- 
changed the shock of impact with the pan is so small that the 
jknife edges are not damaged and the pan rests offer an easier 
| method of control. 

Trial of Weights. — In making a trial of a weight the pointer 
should be allowed to move only far enough to indicate the direc- 
tion of motion. This indicates the proper change to be made in 
weights as well as if it were allowed to move half way across 
Ithe scale and it does not derange the balance. 

Estimation of Zero Point. — It has been stated that in later 
Iwork extended calculations of zero point would not be made. 
On account of the resistance due to friction with the air and in 
phe bearings, any balance decreases the amplitude of vibration 
(with each successive j ourney . The amount of such decrease varies 
[with different balances but a close approximation can be made 
py simple observation. If the zero point of the unloaded balance 
pas been adjusted to coincide with that of the scale, in the final 
[adjustment of weights the loaded balance can also be brought to 
phis ideal zero point without the necessity of extended cal- 
culations, by simply noting that the distance to which the pointer 
swings in one direction is a certain (approximate) fraction of a 
division less than the distance in the opposite direction on the next 
preceding journey. In many cases of quantitative analysis, 
if the longer (even though slightly more exact) method involving 
papulations of zero point were followed, the time consumed 


in weighing would be so great that the weight of the object would 
change appreciably while on the balance, owing to absorption 
or evolution of water, carbon dioxide, etc. 

To Obtain a Specified Weight of Sample. — It is often desirable, 
in order to simplify calculations, to use a certain specified weight 
of sample, as 1 gm, 0.5 gm, 5 gm, etc. This is a difficult operation 
if the ordinary method of weighing by difference is used, because 
the sample that is to be used is poured from a weighing bottle 
and if too much is inadvertently poured out it is not easy to 
return the excess without loss. If the sample is dry and is 
unaffected by free contact with air one can dispense with the 
weighing bottle and weigh directly on a watch glass or scoop. 
For this purpose one may obtain "counterpoised watch glasses," 
which are pairs of glasses the members of which possess so nearly 
the same weight that they can easily be balanced by means of the 
rider so that their weight does not enter into the calculation. 
The method of obtaining a predetermined weight of a sample 
is as follows: The glasses are placed on the pans and exactly 
balanced, if necessary, by using the rider. Weights totaling 
the desired quantity are placed on the right glass and the pan 
rests lowered, steadying the pans at the time if necessary. 
The beam rests are now lowered just enough to allow the central 
knife edge to come to its bearing and the pointer to perceptibly 
move to the left. The balance door is lowered about half way 
and then the sample, in a fine state of division, is carefully poured 
on the left glass until but a slight excess is obtained, as evidenced 
by the swing of the pointer to the right. (The total length ol 
wing as allowed by the beam rests should not exceed two scak 
divisions as otherwise the adjustment of the balance may b( 
disturbed.) Using the spatula, sufficient sample is now removec 
from the glass to make this side slightly too light. This if 
held over the glass and, by gently tapping the spatula, the sampl< 
thus held is gradually dropped to the pan until equilibriun 
of the balance is nearly or quite attained. The excess is dis 
carded. By repeating this process once or twice apparen 
equilibrium is quickly attained and this is confirmed by closin; 
the case and determining the zero point by the usual methoc 
In this way any desired weight of sample can be obtained in 
comparatively short time. The degree of accuracy with whic 



It is finally weighed will depend, as in all other cases, upon 
the nature of the sample, the total weight being taken and 
j:he degree of accuracy possible at other points in the analysis. 

To Correct the Observed Weight for Inequality of Arms. — 
Certain investigations in physics and physical chemistry require 
phat the weight found shall be the absolute weight, correcting 
tor the inequality of arms and for buoyancy of the air. The 
former correction need not be made for analytical work but where 
necessary either of the following methods may be used. 
; Method of Gauss. — Weigh the object first on the left pan and 
then on the right. Let W be the true weight, a the .weights 
required to counterbalance when the object is on the left pan 
ind b the weights when the object is on the right. By the 
Drinciple of moments: 

Wl = ar 
bl = Wr 
WHr = ablr 
W 2 = ab_ 

Therefore the true weight is the square root of the product of 
the two observed weights. Where the inequality of arms is very 
[slight (the usual case) the arithmetical mean of the two weights 
Is a sufficiently close approximation to the square root of the 

! Method of Borda. — This is also known as the method of 
(substitution. Place the object to be weighed on one pan and 
counterbalance with any other material, such as similar weights 
pr dry sand. Remove the object and substitute accurate 
weights until the balance is again in equilibrium. These 
weights are necessarily the same in value as the object for 
vvhich they substitute, irrespective of the relative length of 
he arms. 

Use of Arm Ratio. — Weigh the object, as usual, on the left 
oan with the weights on the right. Multiply the observed 


weight by the ratio .- This gives the true weight. If the arms 

)f the balance are equal in length, j = 1 and the observed 

weight is the true weight. 



Calibration of Weights. — For analytical purposes it is not 
necessary that the various pieces in a set of weights shall have 
the exact values • indicated by the stamp. This is because an 
analysis is always reported as a percent or as some similar 
ratio. The only requirement is, therefore, that the pieces shall 
have the correct relation to each other. That is, the piece 
marked "1 gm" need not weigh exactly one gram. Indeed it 
might conceivably have any other value that is reasonably 
near to one gram; but it is necessary that its weight shall be 
one-tenth of that of the piece marked " 10 gm, " ten times that of 
the piece marked "100 mg, " etc., or that the deviation from 
these ratios be known and corrected in calculations of the weights 
of objects. 

Commercial weights are seldom accurately adjusted unless 
the cost is high. Therefore a calibration should always be made 
and correction applied in such cases as are made necessary by 
errors in manufacture. Also, if standard pieces are available 
it is desirable to correct each set to true gram values, rather 
than to merely relative values, since the chemist frequently 
uses his weights for other than analytical purposes. 

Before beginning the calibration see that in all cases where there 
is more than one piece of a given denomination the different 
pieces bear some distinctive mark. A good plan is to make 
small dots by means of a prick punch. This marks without 
damaging the plate or changing the weight of the pieces. One 
of the 10-gram pieces may be marked (•) and the other (••); the 
three 1-gram pieces (•), ("), ("•); an d similarly with other dupli- 
cate or triplicate pieces. 

Calibration of weights is essentially a comparison of the dif- 
ferent pieces of a set with each other and with a standard, 
followed by a calculation of either relative or absolute values. 
The method to be followed in making this comparison will depend 
upon whether the arms of the balance have been found to be of: 
equal length, within reasonable experimental limits. The 
analytical problem sometimes requires an accuracy represented 
by a maximum error of 0.001 percent, although the maximum 
permissible error is usually larger than this. In order to meet 
these requirements the weights should be calibrated with the 

T . 

same degree of accuracy and if the ratio of arm lengths, ,' is 


jnot greater than 1.00001 nor less than 0.99999, calibration 
[by direct comparison of weights on the two sides of the balance 
may be employed. If the ratio of arms is not within these 
limits the method of direct comparison may still be used, pro- 
vided that a correction is applied. The apparent weight of the 
piece on the right side, in terms of the piece on the left, ismul- 

i . T . 

jtiplied by the ratio j* this giving the true comparative values. 

The principle of Gauss may also be used in making the com- 
parison on a balance having unequal arms but this involves so 
many calculations that it is generally better to use the method 
already described, or else the method of substitution, the latter 
being based upon the principle of Borda. Since these methods 
jare independent of the arm ratio the latter need not be determined. 

Finally it may be stated that the most accurate values are ob- 
tained by ' comparing each -piece with a standard piece of the same 
[denomination. But as this involves the use of an accurately 
standardized complete set and as such a set is necessarily quite 
expensive this method is generally impracticable. ' Instead, one 
standard piece (usually a 1-gram piece) may be used and all other 
lvalues calculated from this. Of course the unavoidable experi- 
mental error is thus cumulative in the larger pieces and the 
work must be done with extreme care if the corrections are to have 
any real significance. 

Calibration. Method of Direct Comparison. — Determine the zero 
[point of the unloaded balance. Place the 1-gram piece marked (•) on 
[the left pan and the one marked (••) on the right. Adjust the rider, if 
necessary, to maintain the zero point at its first position. If the rider 
is not required to restore equilibrium and if the balance arms have been 
found to be of equal length, then the two pieces are of equal value. If 
[the rider is necessary on the right arm the weight on that side is less 
|in value than the one on the left. If the rider was used on the left arm 
(the reverse is true. Record the result as 

1* = 1" + or — n milligram, 

n being the difference between the values of the two pieces. If the 
balance arms are unequal this relative value for 1- is to be multiplied 


Compare similarly the other pieces of the set, as follows: 


Gram Pieces 

!• with standard piece, if one is available 

1— with 1- or I- 

2 with 1- + I- 

5 with 2 + 1- + 1- + 1- 

10- with 5 + 2 + 1- + 1- + 1- 

10- with 10- 

20 with 10- + 10- 

50 with 20 + 10- + 10- + 5 + 2 + 1- + 1- + H 

Milligram Pieces 

500 with 200 + 100- + 100- + 50 + 20 + 10- + 10" + 5 + 

rider at 5 

200 with 100- + 100- 

100- with 50 + 20 + 10- + 10- + 5 + rider at 5 

50 with 20 + 10- + 10- + 5 +rider at 5 

20 with 10- +10- 

10- with 10- 

10- with 5 + rider at 5 

5 with rider at 5 

Also compare all milligram pieces plus the rider at 5 with 1* or with the 
standard piece if the latter has been used. 

If a standardized piece has been used calculate all of the values for the 
other gram pieces from this; otherwise select the one of the gram pieces 
of the set which appears to be most nearly normal with respect to the 
other pieces and call this 1.0000 gram, calculating values for the rest of 
the gram pieces upon this arbitrary assumption. 

In the case of the milligram pieces calculate provisional values for 
each piece, starting with the arbitrary assumption that the smallest 
piece of the set is correct. (The smallest piece is usually a 5- or 10- 
milligram piece. Smaller pieces are not needed because the adjustment 
of the rider provides smaller subdivisions.) Add these provisional 
values. If the total is not that indicated by the comparison of the 
collective milligram pieces with the chosen standard then the provi- 
sional value for each piece is to be multiplied by the factor: — • 

y F J prov. sum 

This will give the new value for the piece, based upon the standard 
finally adopted. The new sum will equal the true sum unless the 
dropping of decimals beyond the fourth place has impaired the accuracy 
of the calculations. 


The ratio, — — » should be calculated as far as six decimal places, 

prov. sum 

This means that considerable work will be involved in multiplying all of 

the provisional values by this ratio. It is much simpler (and sufficiently 

accurate in most cases) to use a somewhat different method for changing 

the provisional values to the true values, thus : 

Subtract the provisional sum from the true sum, applying the correct 
sign to the difference. This gives the total correction to be made. 
Apply 0.5 of this correction to the 500-milligram piece, 0.2 to the 200- 
milligram piece, 0.1 to each of the 100-milligram pieces, and so on. 

Method of Substitution. — In calibrating by this method a second set 
of weights should be provided to serve as counterpoise pieces. This may 
be a low priced set of ordinary weights or a worn out and discarded set 
of analytical weights, since the accuracy of the calibration does not 
depend in any manner upon the accuracy of adjustment of the counter- 
poise. Also it is unnecessary to determine the zero point of the unloaded 
balance, the only requirement being that the adjustment of the rider 
shall be made so as to bring the balance into equilibrium about any 
given scale division as zero point and to maintain this zero point through- 
out the experiment. This point may conveniently be the zero of the 

Place the 1-gram piece marked (•) on the right pan and a 1-gram 
piece of the counterpoise set on the left. Adjust the rider so as to make 
the pointer swing about the true zero of the scale and note the rider 
position. Remove the piece from the right pan and put in its place the 
I-gram piece marked (••). Readjust the rider, if necessary, to restore 
the equilibrium of the balance. If the rider position is the same as 
before the pieces (•) and (••) have the same value. If the rider was 
moved to the right in the second experiment the piece (••) is lighter 
than the piece (•); if the rider was moved to the left the reverse is true. 

In either of the latter cases the amount of rider shift is a measure of 
the numerical difference between the values of the two pieces. 

Continue this process of comparison of pieces, using the combinations 
listed in the directions for calibration by direct comparison. In each 
case the pieces to be compared are placed successively on the right pan, a 
counterpoise piece of the same denomination being placed on the left 

Calculate the true or relative values for all of the pieces of the set 
according to the method already described. 

Reagents. — One of the most vexatious problems with which 
the analyst has to deal is that of obtaining reagents that are 
sufficiently pure to suit his purpose. Methods of manufacture 


are constantly being improved and better chemicals are now 
available than in the past, but even at this time the reagent 
that is assumed to be pure often contains small quantities of 
impurities which interfere with the accuracy of analytical proc- 
esses. Attempts have been made by manufacturers to indicate 
on the label the degree of purity. Thus "c. p." for " chemically 
pure," signifies a reagent containing no impurity in a quantity 
that could be detected by chemical tests. "Com." for "com- 
mercial" means a crude unpurified chemical, "medicinal" 
sufficiently pure for medicinal purposes, "U. S. P." purity as 
specified by the United States Pharmacopoeia, etc. Zinc might 
be labeled "arsenic free" to indicate that it could be used without 
a blank test for a determination of arsenic by Marsh's method, 
or "iron free" so that it could be used for reducing solutions 
in iron analysis without a blank test. If these labels ever did 
have any real value they early lost it. "c. p." has been made a 
cover for a multitude of shortcomings in packages of grossly 
impure reagents. "Medicinal" has meant little more than that 
the manufacturer hoped that the substance so marked might 
be sold to the unsuspecting for medicinal purposes. "Silver 
free" lead (for assaying silver ores) is often lead from which the 
manufacturer has removed a certain fraction (or none at all) 
of the silver originally contained in it. 

Analyzed Chemicals. — On account of the carelessness evident 
in preparing and labeling reagents chemists have come to practi- 
cally disregard all such indications of purported purity and to 
rely upon one or both of two sources of information regarding, 
the purity of reagents. These sources are the reputation that 
the manufacturer bears for producing reliable chemicals and the 
chemist's own personal test of the chemicals themselves. Many 
manufacturers have now entirely discarded the abbreviation 
"c. p." and publish on the label a supposed analysis of the sub- 
stance contained in the package. "Analyzed chemicals" have 
thus become popular, but the inexperienced chemist will make a 
great mistake if he forms the too hasty conclusion that the 
analysis is always correct. It is often very far from being, 
correct. The passage of pure food and drugs acts in this and 
many foreign countries has resulted in great improvement in 
the matter of labeling chemicals that are to be used for medicinal 


purposes, since the label constitutes a legal guarantee as to the 
contents of the package. When these acts are extended to 
include the reagents used for scientific purposes the chemist will 
have a better opportunity for purchasing chemicals of the degree 
of purity of which he can feel reasonably assured. At present 
the only safe plan is to make blank tests for such impurities as 
will interfere in the analysis to be performed. 

Action upon Glass. — While it will be readily conceded that no 
substance can be made absolutely pure yet certain reagents 
can only with difficulty be made even approximately pure. 
Examples are the strong bases, such as sodium hydroxide, potas- 
sium hydroxide, ammonium hydroxide, barium hydroxide, etc., 
which readily attack and dissolve glass so that they are always 
contaminated with silica. On this account their solutions are 
seldom kept as stock reagents in the laboratory but are made from 
the solids as required excepting, of course, ammonium hydrox- 
ide which is a solution of a gas. In such cases the chemist 
will simply require that interfering substances shall be absent. 
Basic solutions often contain precipitated matter. The glass 
bottle is first attacked, the solution accumulating alkali silicates. 
These are later hydrolyzed, causing the precipitation of hydrated 
silica. The rule must never be forgotten that solutions are to 
be filtered just before using for analytical purposes unless they 
are already quite clear and free from sediment. 

Chemical Glassware. — Glassware that is to be used for 
analytical work must possess certain properties that are not 
found in ordinary glass, (a) It must have a very slight solubility 
in water and in solutions of acids, bases and salts, (b) lb must 
have a low coefficient of expansion and be well annealed, in 
order to withstand sudden changes in temperature as well as 
mechanical shock, (c) Its chemical composition must be adapted 
to the work at hand so that traces dissolving shall not affect 
analytical results. For example, a lead glass should not be used 
for solutions in which small amounts of lead are to be determined. 

Prior to the beginning of the world war manufacturing of 
chemical "resistance" glassware had not been developed to any 
great extent in America and most of such material was imported. 
Jena glass, which is probably the best known of imported ware, 
s essentially a borosilicate of sodium, zinc and aluminium. 


The United States Bureau of Standards has made an extensive 
investigation 1 of the qualities of two kinds of imported resistance 
glassware (Jena and Kavalier) and of five American glasses, 
most of the latter having been developed within the past four* 
years, although one or two American glasses have had an excellent 
reputation for a long time. The qualities tested were (a) 
chemical composition, (b) coefficient of expansion, (c) internal 
stresses, (d) resistance to sudden changes of temperature, (e) 
resistance to mechanical shock and (/) solubility in water, 
ammonium hydroxide, mixed solutions of ammonium sulphide 
and ammonium chloride, solutions of sodium phosphate, sodium 
and potassium carbonates, and sodium and potassium hydroxides. 

As a result of this investigation it may be stated that " — — 
all of the American made wares tested are superior to Kavalier 
and equal or superior to Jena ware for general chemical laboratory 

In the following pages Pyrex glass will often be specified but 
it should be understood that any available resistance glass may 
usually be substituted. 

Distilled Water. — Natural waters always contain dissolved 
matter which unfits them for use in analytical work. Besides 
such natural mineral matter and. dissolved gases, water will 
always dissolve certain quantities of the container when allowe 
to stand. In order to remove dissolved solids the water 
distilled and recondensed. Various forms of stills are in use. 
In any form of such apparatus the vessel in which the water is 
boiled should be so far separated from the condensing worm that 
it is impossible for any spray to enter the latter. The boiler 
itself may be of any material, but the condensing worm should 
be of pure tin, silver, or platinum because hot water dissolves 
most other metals and also glass. The cost of platinum 
course precludes its use in any but small stills that are to be use 
for preparing water for exact investigations, such as are carrie 
out in physical chemical work. Pure tin is the metal generally 
used for the purpose. 

Distillation does not free water from dissolved gases and 
work in which carbon dioxide, oxygen, nitrogen, or ammonia v 
interfere it is necessary to boil the water immediately befo 

x Bur. Stand. Tech. Paper 107 (1908). 






'using. Boiling should not be unduly prolonged, since the water 
thus becomes recontaminated with the material of the contain- 
ing vessel. 

Transfer of Liquids.- — The operations involving pouring rea- 
gents from bottles, pouring liquids into a filter or pouring from 
one vessel to another are often so clumsily performed as to cause 
a loss of part of the liquid through splashing or running down 
the outside of the pouring vessel, thus vitiating the results of 
the analysis or at least producing a very disagreeable sort of 
uncleanness of the apparatus. When pouring from a bottle 
the stopper should never be laid on the desk but is held between 
the fingers of the right hand. The bottle is then grasped in 
such a way as to bring the label under the hand and then a glass 
rod is held in a vertical position against the mouth of the bottle. 
The latter is tilted to pour out the required amount of liquid, the 

J. S. Streeter 

Sulphuric Acid, 10% 


Fig. 32. — Form of label for laboratory reagent bottle. 

glass rod being kept in position until the bottle is returned to an 
upright position. No liquid should run down the outside of the 
bottle in this case. If a drop should escape the label will not be 
marred because the method of holding the bottle brings the label 
on the upper side while pouring. When pouring from a beaker 
the stirring rod is use'd in the same way for preventing splashing. 
Records. — In no other part of the analysis is system more 
important than in that of records. Results are many times 
rendered worthless by uncertainty regarding the meaning of the 
experimental figures or regarding the pieces to which recorded 
weights belong. All analyses are run at least in duplicate, some 
in triplicate. If at any stage in the work the beakers, crucibles, 
burettes, or other pieces become interchanged or if the recorded 
weights, volumes, temperatures, or other data are not properly 



labeled and are applied to the wrong pieces, then the calculated 
results of the analysis are, of course, incorrect. At the very- 
outset the duplicate pieces should be numbered (I and II unless 
other marks are preferred) in every case where the mark is not 
objectionable. It is not advisable to mark by labels or pencil 

































Fig. 33. — A convenient blank for reporting the results of a gravimetric analysis. 

any article that is to be weighed because the mark itself often 
causes changes in weight through rubbing off or absorption of 
moisture. For articles that are not to be exactly weighed a 
small label or pencil such as is used for marking glass and porce- 
lain may be used. The latter has a soft core composed of pig 



merit and grease or paraffin so that it will stick to glass. The 
better grades of glassware are now furnished with a spot rough- 
ened by sand blasting, this making it possible to use a common 
graphite pencil or a pen for marking without a label. Even in 
the case of articles that are not to be marked they can be easily 
identified if the analyst follows the plan of always keeping No. I 
on the left and No. II on the right when precipitating, filtering, 
igniting or allowing crucibles to stand in desiccators. 

Reagent bottles should be marked by a label bearing the name 
of the student, that of the reagent, the concentration of the solu- 
tion and the desk number, as shown in figure 32. 

Many systems of note-book records have been used with greater 
or less success. If an ordinary blank note book is used it is 
necessary to so indicate each measured weight or volume that 
there is no possibility of uncertainty regarding the meaning of 
the figures. Loose sheets of paper must not, under any circum- 
stances, be used for records of a permanent character. Loss of such a 
sheet has frequently caused the loss of days or even weeks of 
laboratory work. A better device than that of the blank paged 
note book is found in a systematic record book, with spaces so 
provided and lettered that mere recording figures is all that is 
necessary. Such a page for gravimetric analysis is shown in Fig. 
33. No indication of the identity of weights or percents is 
necessary beyond the filling of the blanks as shown. 



One of the most apparent differences between the methods of 
study as applied to qualitative and to quantitative analysis lies 
in the fact that while in qualitative analysis the metals and acids 
are grouped according to their susceptibility to the action of cer- 
tain " group reagents/' in quantitative analysis widely different 
reagents and methods are often used for elements or acids that 
are closely allied in most respects. A general review of the 
conditions that must be fulfilled for accurate gravimetric analysis 
will serve to show that no such systematic classification as 
is found in qualitative work is desirable in quantitative pro- 
cedure, since the reagent and method must always be selected 
which will give the precipitate which is the most insoluble and 
the most easily separated and purified and which assumes the 
most definite form upon the application of heat. For example, 
calcium, barium and strontium fall in the same periodic as well 
as the same qualitative group and yet there is no logical reason 
for studying these metals together in quantitative analysis 
because barium is most conveniently precipitated as sulphate 
and calcium as oxalate from water solutions, while strontium is 
precipitated as sulphate from solutions containing alcohol. 
Barium sulphate is stable when ignited, and is weighed as such. 
Calcium oxalate decomposes and is weighed as carbonate or oxide 
while strontium sulphate is stable and is weighed in this form. 
Any one of these metals could be precipitated by ammonium car- 
bonate and the carbonate changed to the oxide by ignition but 
this would not, in any case, prove to be the most convenient 
or the most accurate method. 

In view of these facts it becomes desirable to select, for each 
element or radical, that method which is, under the circum- 
stances, the most easily executed and the most accurate and to 
study these, not in order of qualitative or periodic groups, but 



in that order which will best develop manipulative skill and 
accuracy in experimental work. Many methods that were 
formerly used have become almost obsolete because of the 
development of better apparatus or more rapid methods. In 
such cases the older methods will generally receive mention at 
the proper place and the more modern method will be described 
more fully. 

Apparatus. — Most of the ordinary apparatus with which the 
quantitative laboratory is usually stocked is already more or 
less familiar to the student. Special forms of apparatus will be 
described in connection with the determinations for which 
they are to be used. At the beginning of the course all glassware 
should be thoroughly cleaned and other apparatus should be put 
in first class order. A principle that should never be forgotten 
is that both accuracy in results and speed in working are pro- 
moted by following the practice of cleaning apparatus as soon 
as its use in a given experiment is finished, so that it will be ready 
(with perhaps a single rinsing with distilled water) for the next 
operation that may demand it. -The student who works with 
a desk full of soiled, broken or disorderly apparatus or with 
spilled chemicals scattered over the tables may as well make up 
his mind at the outset that he will never be an analyst of any 
useful sort. 

After desk apparatus has been invoiced and cleaned, two desiccators 
like that shown in Fig. 11 are prepared for use. They are first made 
clean and dry and then a layer, one-half inch thick, of fused, granular 
calcium chloride is placed in the bottom and a short piece of sodium 
hydroxide is placed upon this. The former keeps the atmosphere 
within the desiccator free from moisture while the sodium hydroxide 
absorbs carbon dioxide and traces of acid vapors that sometimes come 
from the calcium chloride. The upper part of the desiccator is again 
wiped with a dry towel to remove calcium chloride dust and a clay or 
alloy triangle is bent so as to lie freely on the shoulder above the 
calcium chloride. (Perforated porcelain plates are sometimes used 
| instead of triangles for this purpose.) A thin film of vaseline is rubbed 
on the ground joint of the cover and the latter is then worked down 
until it fits well, the surplus vaseline being removed from the edge. 
The desiccator is now ready for use. 

Prepare also two wash bottles (one each for hot and cold water) 
like the one illustrated in Fig. 7, using liter flasks for the purpose. The 


neck of the one that is to be used for hot water should be wrapped with 
cotton cord or sheet cork. 

A number of stirring rods, 4 to 6 inches long, should be made by 
fusing and rounding the ends of glass rod or tubing. One or two of 
these, tipped with pieces of rubber tubing having an end cemented 
together, are to be used for loosening precipitates from beakers and 
dishes. These are known as the chemist's "policemen." 

' Calcium 

Calcium may be precipitated from ammoniacal solution as 
carbonate by alkali carbonates or as oxalate by alkali oxalates. 
Since ammonium carbonate or oxalate yields volatile ammonium 
salts as byproducts in the reaction and since traces of these will 
be expelled upon ignition if not completely washed out of the 
precipitate, the ammonium salts are always used in preference 
to those of sodium or potassium. 

Solubility. — The solubility of calcium carbonate in water was 
determined by Kohlrausch and Rose 1 by conductivity experi- 
ments and was found to be 0.013 gm (corresponding to 0.0052 
gm of calcium) per liter at 18°, but the solubility is considerably 
increased by ammonium salts, such as must be present when 
ammonium carbonate reacts with a calcium salt in solution. The 
solubility of calcium oxalate in water was found by Kohlrausch 
and Rose to be 0.0056 gm of the crystallized salt, CaC 2 4 H 2 
corresponding to 0.0015 gm of calcium) per liter. The solubility 
is considerably reduced by an excess of ammonium oxalate. On 
account of this difference in solubility the oxalate method is 
preferable. Ammonium chloride should be present in either 
case to prevent precipitation of traces of magnesium. The 
reaction involved in the precipitation may be expressed as follows: 

CaCl 2 +(NH 4 ) 2 C 2 4 ^CaC 2 4 +2NH 4 Cl. 

Ammonium oxalate does not readily dissolve in water, the 
saturated solution at 0° containing 2.2 percent of the salt. As 
the temperature is raised the solubility is increased so that a 
saturated solution at 20° is easily made by heating in contact 
with an excess of the salt and allowing to cool. It is not best 
to keep a stock solution of ammonium oxalate because it readily 
iZ. physik. Chem., 12, 234 (1893); 44, 197 (1903). 


undergoes hydrolysis, yielding ammonium hydroxide which 
attacks the glass, and because there occurs a decomposition in 
solution as follows : 

(NH 4 ) 2 C 2 04->(NH4)2C0 3 +CO. 

It is preferable to make a small amount of the solution as needed 
for the determination. 

A difficulty that is often encountered by the inexperienced 
analyst is the formation of a precipitate of calcium oxalate which 
is so finely crystalline that it passes through the pores of the 
filter, making its complete separation impossible. Refiltration 
of the portion that has passed through will partially remedy this 
trouble but when a precipitate is found to be too finely divided 
for filtration the only satisfactory cure is found in digestion. 
Certain grades of filter paper are made for filtering fine precipi- 
tates, their structure being very dense. This renders filtration 
less rapid than would be the case with papers of ordinary density. 
The difficulty nearly or quite disappears if the proper conditions 
are observed during the precipitation. It has been explained 
(page 20) that too rapid precipitation causes the formation of a 
large number of small particles rather than a small number of 
large particles. Two conditions are found to be suitable for the 
formation of large crystals of calcium oxalate: (1) boiling tem- 
perature for solutions of calcium salt and ammonium oxalate 
and (2) slow addition of reagent. These conditions will be 
elaborated in the directions for the determination. 

Calcium may also be determined by precipitating as sulphate 
from a solution containing alcohol 1 or volumetrically ~by pre- 
cipitating as oxalate and titrating with a standard solution of 
potassium permanganate. The latter method will be described 
in the section dealing with volumetric analysis. 

Calcium precipitated as either carbonate or oxalate may A be 
weighed as carbonate, oxide or sulphate. Calcium oxalate 
decomposes as follows, on the application of heat: 

CaC 2 4 -*CaC0 3 +CO, (1) 

CaC0 3 ->CaO+C0 2 . (2) 

1 Stolberg: Z. angew. Chem., 17, 269 (1904). 



Reaction (1) begins at quite low temperatures. Reaction (2) 
begins as low as 500° but requires long heating at this tempera- 
ture to complete the decomposition. In practice the highest 
temperature attainable by the blast lamp is applied and heating 
is continued until no further loss in weight occurs. Calcium 
oxide is not reduced by hot carbon and the precipitate may be 
heated without removing from the paper. Many chemists 
prefer to ignite the oxalate at a low temperature and to weigh as 
carbonate but this procedure is of doubtful utility, even for an 
experienced analyst, on account of the difficulty in stopping t 
decomposition at a point where all the oxalate has disappeare 
and no oxide has been formed. It is also possible to add a few 
drops of sulphuric acid to the crucible containing the calcium 
oxalate and to weigh the resulting calcium sulphate, after gentle 
ignition to expel the oxalic acid and the excess of sulphuric acid 
The most important source of error in this procedure come 
from the loss by spattering, a loss which even the greatest ca 
can scarcely prevent. There is also danger of decomposing ca 
cium sulphate by strong heating: 

CaS0 4 ->CaO+S0 3 . 

It is much preferable to heat strongly the precipitate convertin 
it quantitatively into calcium oxide. The completion of th 
conversion can be judged by the constancy in the weight of th 
substance upon further heating. Calcium oxide readily absorb 
moisture and carbon dioxide when exposed to the air and tb 
resulting change of weight will become appreciable if the proces 
of weighing is unduly prolonged. If the weight of the crucible 
and Oxide is approximately known most of the weights may be 
placed on the balance pan before the crucible is removed from 
the desiccator and the remainder of the process completed in a 
short time. It is a good plan to keep a small piece of potassiu 
hydroxide in the desiccator in which calcium oxide is to be pr 
served. This lessens the absorption of carbon dioxide b 
keeping the atmosphere free from that gas. 

The converse of this method may be used for the determin 
tion of the oxalate radical, precipitation being made by a soluble 
calcium salt in basic solution. Volumetric methods are, how- 
ever, preferable. 


Determination. — Fill a clean dry weighing bottle with the calcium 
compound to be analyzed. Provide two clean beakers of Pyrex or other 
resistance glass, having a capacity of 250 cc, and mark them I and II. 
If the substance is of such nature that it is altered in any way by free 
exposure to air the sample to be used must be weighed by difference as 
follows: Place the bottle on the balance pan. using for this purpose a 
pair of crucible tongs, having short pieces of rubber tubing drawn over 
the tips, and carefully weigh. Record this weight in the data book at 
the top of the space marked for sample I, reading the weights as directed 
on page 62. Carefully remove the stopper, holding over beaker marked 
I, and pour what is judged to be between 0.2 gm and 0.5 gm into the 
beaker. Replace the stopper, taking great care that any falling par- 
ticles drop into the beaker and are not lost, then reweigh the bottle and 
contents. Record this weight under the first, also at the top of the 
space for sample II. Remove a second portion to beaker II and 
reweigh the bottle, recording under the preceding weight. Subtracting 
the less weights from the greater gives the weights of sample used. 

If the substance is known to be unaffected by contact with air it may 
be poured into one of the counterpoised glasses and weighed directly, 
being then brushed into the beaker by the small pencil brush of camel's 

After having weighed the two samples for analysis determine, by 
qualitative tests on another portion of the substance, whether it is solu- 
ble in water and, if not, whether in dilute hydrochloric acid. If soluble 
in water dissolve in about 100 cc of distilled water, add 5 cc of 10 percent 
ammonium chloride solution and treat each sample as directed below. 
If insoluble in water but soluble in hydrochloric acid first moisten 
each sample with water then cover the beakers with glasses and carefully 
add about 20 cc of dilute acid* After effervescence has ceased rinse 
down the cover glass and the sides of the beaker with water from the 
wash bottle, dilute to about 100 cc and gently boil for one minute to 
expel dissolved carbon dioxide. From this point the procedure is the 
same as for water soluble salts. 

Prepare ammonium oxalate solution by heating to boiling 5 gm of 
powdered ammonium oxalate and 100 cc of water in a Pyrex beaker. 
Part of the salt will crystallize when cooled, leaving a saturated solution. 
Add dilute ammonium hydroxide (filtered unless already quite clear) to 
the solution of calcium salt until the liquid smells very distinctly of 
ammonia. In determining this point the ammonia that is already in the 
air above the liquid must be blown away before testing the odor. Heat 
to boiling and add, drop by drop from a pipette, 20 cc of the recently 
prepared ammonium oxalate solution, stirring vigorously during the 
addition. If this is carefully done the precipitate should settle readily, 


leaving a clear liquid above. When this is the case add a few drops 
more of ammonium oxalate solution, observing whether any precipitate 
forms. If so, more reagent must be added in the same manner as at 
first. When precipitation has been shown to be complete the liquid is 
digested at a temperature somewhat below the boiling-point over a 
low flame or on the steam bath for one-half hour or until the super- 
natant liquid is quite clear, when it is ready for filtration. The odor 
of ammonia should still be easily perceptible at this stage. 

Prepare two filters of extracted paper, marking the funnels I and II. 
Carefully decant each solution, while hot, upon the proper filter, allowing 
to run through into clean beakers. Observe the directions given on 
page 73 for proper method of pouring from beakers. Before completing 
the filtration a few drops of ammonium oxalate solution should be added 
to the clean filtrate that has already run through. If a precipitate 
forms, the filter must be well washed; the filtrate and washings returned 
to the beaker in which precipitation was made and more reagent added 
in the same manner as before, until precipitation is complete. When no 
precipitate is produced in the filtrate, complete the filtration, washing 
the precipitate into the filter by a stream from the hot water bottle, 
rubbing the beakers with a glass rod tipped with rubber tubing (a 
''policeman")- ' Wash the precipitate on the filter until a small amount 
of the last washings fails to give more than a faint precipitate with 
silver nitrate and a drop of nitric acid showing that chlorides have been 

Allow the paper to drain then remove from the funnel, fold as directed 
on page 34 and place in a previously ignited and weighed porcelain or 
platinum crucible. The cover should have been weighed with the cruci- 
ble because the closed crucible will hinder absorption of moisture and 
carbon dioxide while weighing. The crucible is carefully heated by 
the burner until the moisture has been expelled and smoking begins. 
The cover may now be removed and placed on a clean tile while the 
crucible is heated more strongly until the paper is completely charred. 
The crucible is now placed on its side, the cover adjusted and complete 
oxidation of the carbon of the paper is accomplished as described on 
page 36. The crucible is then placed in an upright position, is covered 
and subjected to the hottest flame available from the blast lamp. 
This ignition is continued for 30 minutes, when the crucible is placed 
in the proper desiccator, allowed to cool to the temperature of the 
room and quickly weighed. It is ignited for 10 minutes longer, cooled 
and reweighed. If there is a decrease in weight of more than 0.0003 
gm the crucible is reheated for 10 minutes and weighed, the process 
of heating and weighing being continued until the weight remains 
constant within the limit given above. The preliminary weighings 



should be recorded upon the back of the sheet preceding the one used 
for the final record, the final weight being recorded in the proper 
blank space. 

Calculate the percent of calcium in the sample, using the factor 
already calculated and recorded on page 9, and using a table of loga- 
rithms for the arithmetical work. Do not discard the ignited product 
until after the report has been accepted. Errors may have been made 
which can be corrected if this has been preserved. This is a principle 
that should be observed, when possible, in all analytical work. 


Silver might be gravimetrically determined as chloride, 
bromide or iodide. The solubilities of these salts in water, 
shown in the following table, were determined by Kohlrausch 
and Rose. 1 

Milligrams per liter, soluble at 18° 


Silver equivalent to salt 

Silver chloride 



Silver bromide 


Silver iodide. 


From the comparative solubilities one might conclude that 
silver chloride is the least desirable form in which to weigh silver. 
The bromide and iodide, however, are much more sensitive to 
the action of light, being more readily decomposed into sub- 
halides with liberation of free halogen. The stabilities of 
these salts are in the same relation to each other as are those 
of halides of other metals and of hydrogen. On this account 
the gravimetric determination of silver is invariably made by 
weighing silver chloride, using hydrochloric acid as the precipi- 
tating reagent. Conversely the determination of chloranion 
is made by using a soluble silver salt as the reagent. A small 
excess of either a soluble silver salt or a soluble chloride greatly 
diminishes the solubility of silver chloride as explained in the 
section dealing with the principles of precipitation. If more 
than a very slight excess of a metal chloride is present the solu- 

1 Z. physik. Chem., 12, 234 (1893). 


bility of silver chloride is increased, because of the formation 
of soluble double salts. On this account hydrochloric acid 
is used as the precipitant for silver. 

Silver chloride shows a well defined tendency toward the for- 
mation of a hydrosol in cold water and when this is formed the 
solubility follows no definite rule. The sol can be flocculated 
by boiling with dilute acids or other electrolytes. 

The precipitate of silver chloride is affected by light as are 
the other silver halides, it being reduced to a subchloride, Ag 2 Cl, 
with liberation of free chlorine. The darkening of silver chloride 
under the influence of strong light is due to the appearance of 
the subchloride which is bluish black in color. While some 
decomposition undoubtedly occurs in daylight of any intensity, 
if the precipitation is performed in the darker parts of the room 
ordinary diffused light will not appreciably affect the weight of 
the precipitate in a short time. 

The precipitate cannot be ignited in contact with filter paper 
on account of the ease with which it is reduced to metallic silver. 
Use can be made of any of the devices for dealing with such pre- 
cipitates, as mentioned in the general discussion of the ignition 
of precipitates; we may here follow either the method of removing 
the precipitate from the filter paper or the method involving 
the use of the Gooch crucible, both methods being described. 
Whenever silver chloride is heated, care must be taken to prevent 
a rise of temperature above the point of fusion, which is 451°, 
since it is sensibly volatile at high temperatures. 

By making the proper changes in procedure the halogens may 
be determined, silver nitrate being used as the precipitating 
reagent. These determinations are discussed later (page 87). 

Determination. — Fill a stoppered weighing bottle with the powdered 
silver salt and weigh out two samples of about 0.3 gm each for analy- 
sis, placing in 250-cc beakers, which may be of ordinary hard glass. 
The weighing may be done upon counterpoised glasses or from the weigh- 
ing bottle, by following the directions given for the determination of 
calcium. In this, as in all other cases, care must be exercised to avoid 
spilling any of the substance upon the balance pan, as this invariably 
results in corrosion of the pan. 

Dissolve the weighed sample in about 100 cc of water, and heat 
nearly to boiling. Add, with stirring, 5 cc of dilute hydrochloric acid 


and digest on the steam bath until the precipitate is completely floccu- 
lated and the supernatant liquid is quite clear. Test the clear liquid 
with more hydrochloric acid as soon as is practicable, to determine 
whether precipitation is complete. When the precipitate settles com- 
pletely proceed by one of the following methods. 

(a) Filtration on Paper. — Filter on extracted paper, with or without 
slight suction, and wash with hot water containing 1 cc of dilute nitric 
acid in each 100 cc of water, the acid being used in order to prevent the 
silver chloride from returning to the condition of a hydrosol and thus 
passing through the filter. Wash until free from chlorides, testing the 
washings with silver nitrate. Allow the precipitate to drain, then 
remove the paper, fold over the top and sides (see page 34), place on a 
watch glass and dry in an oven at 100°. 

When the precipitate and paper are completely dried, place a piece 
of black glazed paper on the desk, unfold the filter paper and carefully 
detach as much as possible of the precipitate, using a spatula for this 
purpose and allowing the precipitate to fall upon the central portion of 
the glazed paper. While it is desirable to leave as little as possible of 
the precipitate upon the filter paper it is also essential that no paper 
fiber be removed with the main portion of the precipitate since this 
portion is not to be treated to reconvert reduced silver into silver chloride. 
With the small camel's hair brush the precipitate is now brushed into a 
pile, loosening any particles that may have been caught by the brush, 
and is covered with a watch glass. An ignited and weighed crucible is 
placed upon one corner of the glazed paper. The filter paper is refolded 
in the same manner as before, is rolled into a tight roll and a stiff plati- 
num wire is coiled around it in a manner such that the roll can be held 
over the crucible by means of the wire. Being held in this position it is 
touched with the oxidizing flame of the gas burner until the paper ignites. 
The gas flame is to be used only often enough to keep the paper ignited 
and the outer oxidizing portion of the flame is always to be used for 
this purpose. The paper is thus burned, the ash falling into the crucible, 
where the combustion is completed at a low temperature. Some silver 
has been reduced even with these precautions. To change this again 
into silver chloride the ash is moistened with a drop or two of concentrated 
nitric acid and, after a few minutes, a drop of concentrated hydrochloric 
acid is added. The reduced silver is first dissolved by the nitric acid, 
forming silver nitrate, and this is changed to silver chloride by the 
hydrochloric acid. 

Evaporate the acids by placing the crucible on a water bath, then 
carefully brush into the crucible the main portion of the precipitate. 
Heat gently over the burner until the precipitate shows the first appear- 
ance of fusion where it is in contact with the sides of the crucible. In 


case the precipitate has been unduly reduced by light or if it becomes 
reduced when heated, on account of cellulose derived from the paper, 
it should be moistened by nitric acid and hydrochloric acid, as directed 
above, and warmed, when it will usually become white, after which 
the stronger heating to the fusion point may be performed. Place the 
crucible in the desiccator and weigh as soon as cool. Calculate the per- 
cent of silver in the salt, using the factor for silver in silver chloride as 
already calculated in the table of factors on page 9. 

(b) Filtration on a Gooch Crucible. — A device which is similar to that 
shown on page 24 is used for applying the suction. Place a porcelain 
Gooch crucible in the holder, apply the suction and pour into 
the crucible a suspension of purified and shredded asbestos until 
a mat about 1 mm thick is obtained on the bottom of the crucible. 
Asbestos to be used for this purpose should have been prepared by digest- 
ing for one hour with concentrated hydrochloric acid to dissolve 
any acid-soluble material, then washing with distilled water until free 
from chlorides. The desirable thickness of the mat in the crucible 
will depend somewhat upon the character of the asbestos fiber; if the 
latter has a fine texture a closer felt will result, with a consequent 
increase in the efficiency of the filter. 

Draw all surplus water from the filter by means of the pump, then 
rinse once with redistilled alcohol. This is for the purpose of promoting 
rapid drying. Remove the crucible from the holder, wipe the outside 
then dry in an oven at 105° to 110° for 30 minutes. Cool in a desiccator 
and weigh. Heat again in the oven for 30 minutes, cool and weigh. 
This weight will usually be the same as the first but if there is a decrease 
of more than 0.3 to 0.5 mg the heating and weighing must be continued. 

When the weight of the crucible has become constant replace the 
crucible in the holder and again apply suction. Carefully filter the 
solution from which the silver has been precipitated, finally transferring 
the entire precipitate to the filter. Wash and test the washings as 
directed for the method of filtering on paper. Finally rinse once with 
alcohol, dfy to constant weight at 105° to 110° and calculate the per- 
cent of silver in the sample. 

This method is usually to be preferred to the one first described. The 
chief source of error is in the loss of asbestos during filtration and 
washing. This may be prevented by proper preparation of the asbestos 
suspension before using, the fine material being removed by sedimenta- 
tion and decantation. Even when this has been done it is necessary to 
keep the suction applied at all times during filtration and washing, the 
asbestos thus being held down in the felt. 


Chlorides (Bromides and' Iodides) 

The method for the determination of the halogen of halides is 
the converse of the one just described, silver nitrate being used 
as the precipitating reagent. Nitric acid must be present to 
prevent the precipitation of salts of silver, other than the halide. 

Determination. — From a stoppered bottle weigh into 250-cc beakers 
duplicate samples of about 0.2 gm of the chloride, bromide or iodide. 
Dissolve in about 100 cc of distilled water and add 1 cc of dilute nitric 
acid. Both distilled water and nitric acid must be tested and found 
free from chlorides. Heat the solution to near boiling and add, slowly 
from a pipette, a 5 percent solution of silver nitrate until no further 
precipitate is produced. Digest on the steam bath until the precipitate 
flocculates readily, leaving a clear solution. Test the clear liquid with 
another drop of silver nitrate solution to insure complete precipitation. 

For the filtration either of the methods described for the determination 
of silver may be used but the Gooch crucible is recommended, especially 
for the determination of halogens other than chlorine. Filter the solu- 
tion and wash free from silver with chloride-free distilled water contain- 
ing nitric acid, testing the washings at the last with a drop of dilute 
hydrochloric acid. The preparation of the filter and the treatment of the 
precipitate on the filter are to be exactly as described above for the 
determination of silver. Calculate the percent of halogen in the sample. 

The precipitate must be protected from the action of light, especially 
in the case of silver bromide or iodide, these substances being much 
more sensitive to light than the chloride. 


The gravimetric determination of aluminium, as well as of 
iron, chromium, nickel, cobalt and copper, may be accomplished 
by precipitating them as hydroxides, igniting and weighing these 
as oxides. For reasons that will presently be discussed all of 
these metals excepting aluminium are now usually determined 
by volumetric or electrolytic processes. 

Aluminium is precipitated as hydroxide by any basic solution, 
whether it be that of a pure base or of a hydrolyzed alkali salt 
of a weak acid, such as sodium carbonate or ammonium sul- 
phide. Aluminium hydroxide readily forms hydrosols which 
are flocculated by the addition of electrolytes, Which must for 
this purpose be inorganic salts. Since the flocculated hydroxide 
is also of a colloidal nature (hydrogel) it manifests the phenome- 


non of adsorption to a marked degree and the inorganic salts 
are consequently washed out with considerable difficulty. For 
this reason, as well as for other and more important ones, am- 
monium hydroxide is chosen as the precipitant because the by- 
products of the reaction will thereby be ammonium salts and re- 
maining traces will be volatilized when the precipitate is heated 

Solubility in Bases. — Aluminium hydroxide, besides being 
soluble as a hydrosol, also dissolves in solutions of bases. It 
thus happens that if an excess of the basic precipitant is in 
advertently added, part or all of the precipitate returns to the 
solution, the amount dissolved depending upon the excess anc 
ionization of precipitant. The strong bases, as sodium or 
potassium hydroxide, dissolve aluminium rrydroxide more readily 
than the weaker ones and this furnishes a second reason for the 
use of the weaker base, ammonium hydroxide, as the precipitating 
reagent. In case an excess of this has been added it is possible 
to remove it by boiling. 

The solvent action of bases has been explained upon the sup- 
posed ability of aluminium hydroxide to exist in both acid and 
basic form. When precipitation is taking place, the solution, 
besides holding more or less of the hydrosol, is saturated also 
with the substance in the condition of molecular aluminiu 
hydroxide in equilibrium with two sets of ions. This equilib 
rium within the solution might be expressed simply thus: 

+ + + 

^=>A1 + 3 OH 
A1(0H) ^H + H 2 aTo 3 
(If the substance is an acid it must ionize in three stages: 

Al(OH) 3 -+H+H 2 A10 3 , 

H 2 A10 3 — H+HAIO3, 


Also Al(OH) 3 ->H+Ai02+H 2 0. 

Since it must necessarily be very weakly ionized the ion H 2 A10a 
must predominate, but the equilibrium represented above can 
be but an approximate representation of the real conditions.) 


The addition of either a strong base or a strong acid to such a 
system would cause aluminium hydroxide to dissolve if the 
resulting salt were soluble. The effect of the strong acid upon 
the acid form would be that of suppression of the (already small) 
ionization. Its effect upon the basic form would be interaction 
to form a salt: 

A1(0H) 3 +3HC1-+A1C1 3 +3H 2 0. 

The disturbance of equilibrium resulting from the disappearance 
of hydroxyl ions (due to the formation of water) would cause 
more aluminium hydroxide to ionize in this manner and, since 
nonionized aluminium hydroxide is also in equilibrium with the 
undissolved portion, more would go into solution. 

The effect of a base would be quite similar to that of an acid, 
although by a different process and through the formation of a 
different salt. The basic ionization of the aluminium hydroxide 
would be suppressed while the added base would react with the 
acid form: 

H 3 A10 3 +NaOH^NaH 2 A10 3 +H 2 0, • 

HAIO2 + NaOH^Na AIO2+ H 2 0. 

Here again a salt is for*med, although the aluminium appears in 
the anion. The disturbance of equilibrium has the same ultimate 
result as before, namely, that the solid substance passes into 

Adsorption. — It has already been stated that the washing of the 
gelatinous precipitate is more or less difficult on account of the 
adsorption of dissolved salts. It is impossible to avoid- the 
presence of such salts when making separations or when pre- 
cipitating aluminium from such compounds as the alums. 
(Ammonium salts are always present.) There is also danger 
that, in the case of prolonged washing by water to remove 
alkali salts or other salts, some of the hydrogel may return to 
the condition of the hydrosol. In order to prevent this the 
customary device of having present an ammonium salt in the 
wash water is used. Some of this necessarily remains with the 
precipitate at the last. If this salt is ammonium chloride some 
of the aluminium will be lost by volatilization, the chloride being, 
as with most other metals, more volatile than the salts of other 


acids. The chloride will be formed during ignition by interaction 
of the aluminium oxide or hydroxide and the ammonium chloride : 

Al(OH)s+3NH 4 Cl->AlCl3+3NHs+3H 2 0. 

Ammonium nitrate should therefore be used in the wash water. 
Aluminium nitrate, if formed, is decomposed at high temperatures 
into aluminium oxide and oxides of nitrogen. 

If the precipitate of aluminium hydroxide is filtered and washed 
under diminished pressure, care should be exercised that the 
liquid is not, at any time before the completion of the washing 
process, drawn out so nearly completely as that the precipitate 
should harden and crack. In such a case the wash water that 
is subsequently used will run through the cracks instead of 
through the body of the precipitate and complete washing will 
therefore be accomplished only after the use of much water. 
If it becomes necessary to allow the precipitate to remain in 
the funnel from one day to the next and before the washing is 
completed, the precipitate may be kept moist by plugging the 
stem of the funnel, covering the precipitate with water, and 
placing a watch glass over the top. 

By making suitable changes in the procedure, aluminium 
chloride might be made a reagent for the quantitative determina- 
tion of hydroxyl. Volumetric methods are always used instead. 

Determination. — Fill a weighing bottle with the powdered sample 
of an aluminium salt. Choose the method to be used in weighing 
according to the nature of the substance and weigh two samples of about 
1 gm each into Pyrex beakers. Dissolve in 100 cc of water and add 
dilute, recently filtered ammonium hydroxide, stirring until the liquid 
is distinctly basic, as shown by a bit of litmus paper thrown into the 
beaker. Boil until the precipitate is coagulated and until the odor of 
ammonia above the solution is but faint. Boiling after the odor has 
disappeared may cause some of the precipitate to return to the solution: 

A1(0H)3+3NH 4 C1^A1C1 3 +3NH 3 +3H 2 0. 

Allow the precipitate to settle and then filter through paper, using a 
filter pump attached to a bell jar or filter flask and placing a supporting 
cone of hardened paper or platinum in the funnel (page 23). Wash 
with hot distilled water containing 1 percent of ammonium nitrate until 
the washings are free from chlorides, shown by adding a drop of nitric 
acid and a few drops of silver nitrate solution to a cubic centimeter of the 
washings caught in a test tube; also from sulphates, as shown by adding a 


drop of dilute hydrochloric acid and a few drops of barium chloride solu- 
tion to another portion of the washings. Suck the precipitate as nearly 
dry as possible and transfer the paper and precipitate to a porcelain or 
platinum crucible which has been ignited and weighed, folding the paper 
in the manner already learned. 

Heat very gently in the covered crucible until the moisture is volatil- 
ized, then raise the temperature and burn the paper, inclining the cru- 
cible and placing the cover as in the case of the ignition of the paper 
containing calcium oxalate. When all of the carbon has been burned, 
cover the crucible and heat over the blast lamp for 30 minutes. Cool 
in the desiccator and weigh. Heat again for 10 minutes, cool and weigh. 
If necessary repeat this process until the weight is constant. 

Calculate the percent of aluminium in the salt. 

Aluminium oxide absorbs water from the air, reforming the hydroxide 
with a corresponding gain in weight. On this account the crucible and 
oxide should be weighed rapidly. 

Copper, cobalt and nickel cannot be quantitatively precipi- 
tated by ammonium hydroxide because of the formation of 
soluble complex ammonium salts. Sodium hydroxide or potas- 
sium hydroxide is used as the reagent. Adsorption of the 
reagent by the precipitate causes a large error and volumetric 
or electrolytic methods are preferable. 


Barium may be precipitated as sulphate, carbonate or chro- 
mate. The sulphate and chromate are weighed as such, while 
the carbonate is ignited to form the oxide, in which form it is 
weighed. The solubilities are as follows (determined by 
Kohlrausch and Rose 1 ). 

Milligrams per liter soluble at 18° 


Barium equivalent to salt 

BaC0 3 





BaS0 4 

1 53 

BaCr0 4 


The sulphate is seen to be the most suitable compound for 
the separation of barium from solution. The precipitating re- 
agent may be sulphuric acid or a soluble alkali sulphate. Since 
the latter produces by the reaction alkali salts that must be 

1 Z. physik. Chem., 12, 234 (1893). 



washed from the precipitate, while the former produces volatile 
acids, sulphuric acid is generally used for the purpose: 

Ba(N0 3 )2+Na 2 S04->BaS04+2NaN03. 
Ba(N0 3 )2+H 2 S04->BaS04+2HN0 3> 

Barium sulphate easily precipitates in the form of fine crystals. 
If precipitation takes place rapidly and from a somewhat con- 
centrated solution the crystals may be so small as to pass through 
the filter. Remedies similar to those applied to calcium oxalate 
may be used also with barium sulphate. These are use of a 
dense paper for the filter, precipitating slowly from a hot solu- 
tion and digestion of the precipitate in the mother liquor at a 
temperature near the boiling point. Sometimes the filter paper 
is treated before using with a hot, concentrated solution of 
ammonium chloride which softens and swells the cellulose fibers, 
making a less permeable filter. Such treatment is of doubtful 
utility since the ammonium chloride must later be washed out 
of the paper or cause some volatilization of barium chloride when 
the precipitate is ignited. No trouble will be experienced if 
the precipitation is accomplished under proper conditions. 

The converse of this method may be used for the determina- 
tion of the sulphate radical and for sulphur of any compoun< 
that may be changed to a sulphate. Barium chloride is thei 
the precipitating reagent and the solution is made slightly aci< 
by adding hydrochloric acid. The latter is necessary to prevenl 
the precipitation of barium salts of certain other acids whos< 
salts might be present. Examples of such salts are carbonates, 
oxalates, phosphates and borates, the barium salts of all of thes< 
being insoluble in water but soluble in hydrochloric acid. 

However, barium sulphate itself is appreciably soluble in 
hydrochloric acid as is shown in the following table: 1 

Solubility of barium sulphate, milligrams per liter 

Hydrochloric acid 

Earium sulphate 

Barium equivalent to 
barium sulphate 









Banthisch: J. pr. Chem., 29, 54 (1884). 


Therefore, if hydrochloric acid is present in any considerable 
quantity in the solution it must be nearly neutralized before 
precipitating the barium sulphate, not only because of its solvent 
action above shown but also because of its tendency to increase 
the occlusion of other salts by barium sulphate. 1 

Such occlusion readily takes place if iron salts are present. 
Barium chloride itself is also readily occluded by barium sulphate. 
In the latter case the nature of the resulting error depends 
upon whether barium or sulphuric acid is being determined. 
If the former, the result is a negative error because part of the 
barium is weighed in combination with chlorine, w T hose equivalent 
weight is less than that of the sulphate radical. If the sulphate 
radical itself is being determined the error of occlusion is positive 
because any barium chloride that may be carried down is simply 
a part of the precipitating reagent, contaminating the pre- 
cipitate of barium sulphatej In order to avoid either error the 
concentration of hydrochloric acid should be made as small 
as will serve to hold such salts as the carbonate, oxalate, etc., 
in solution and the precipitation of the sulphate must be ac- 
complished by slow addition of the reagent to the hot solution, 
stirring vigorously meanwhile. 

Change of Weight of Barium Sulphate. — Considerable care 
must be exercised in burning the paper upon which barium 
sulphate has been filtered and in subsequent ignition of the 
precipitate to expel traces of moisture. If the temperature 
is allowed to rise to too high a point barium sulphate will gradu- 
ally decompose, yielding sulphur trioxide and losing weight 

BaS0 4 ->BaO+S0 3 . 

On this account the blast lamp should never be used for heating 
the precipitate and the temperature is not allowed to rise above 
that of dull redness. 

On the other hand, errors may occur through partial reduction 
of barium sulphate by carbon monoxide or organic gases resulting 
from heating of the filter paper. Barium sulphide is thus 
produced and again the material loses weight: 

BaS0 4 +4CO^BaS+4C0 2 . 
1 Richards: Z. anorg. Chem., 8, 413 (1895). 


In order to avoid this reduction the temperature should be 
held at as low a point as will serve to accomplish the combustion 
of the paper and a plentiful supply of air must be maintained 
by inclining the crucible and cover, as directed on page 36. Even 
with these precautions some reduction may occur but if heating 
is continued for a few minutes after the carbon has disappeared, 
reoxidation will take place: 

BaS+20 8 -»BaS0 4 . 

If it should be suspected that either or both of the errors just 
discussed has occurred in any given analysis a correction may be 
made by adding a drop of dilute sulphuric acid to the pre- 
cipitate after the first weighing, then gently reheating to expel 
the excess of acid and water, and reweighing. A gain in weight 
is taken as evidence that sulphide or oxide of barium was present 
in the first case. The second weight is then the correct one. 

This addition of acid, with subsequent heating, also serves 
to correct any error that may have occurred in the determination 
of barium, through the occlusion of barium * chloride by the 
precipitating barium sulphate. It will be recalled that such 
occlusion occasions a negative error in the determination of 
barium, but a positive one in the determination of the sulphuric 
acid radical. Then in the first case sulphuric acid converts 
occluded barium chloride into barium sulphate and gives a 
precipitate of correct composition. In the second case barium 
chloride is an occluded impurity in the precipitate and its con- 
version to sulphate merely serves to increase the error. There- 
fore, when barium chloride is used as the precipitating reagent 
for sulphuric acid it is highly important that the precipitation 
should be carried out very slowly by adding the reagent drop-wise 
and stirring vigorously. This method serves not only to minimize 
occlusion of the reagent but also to prevent the formation of a 
very finely divided precipitate. 

Determination. — Weigh about 0.2 gm of a barium salt into each of 
two beakers, dissolve in water or the least possible quantity of hydro- 
chloric acid, dilute to 100 cc and heat to boiling. Add, drop by drop, 
with vigorous stirring, 2 cc of 25 percent sulphuric acid. Allow the 
precipitate to settle somewhat and test the supernatant liquid, as usual, 
to determine whether precipitation is complete. Digest on the steam 


bath for 15 minutes or longer, until the precipitate settles readily. 
Filter without the use of a pump, on an extracted paper and wash several 
times with hot, distilled water, testing the washings for sulphates. 

Remove the paper from the funnel, fold and place in a weighed cru- 
cible. Incline the crucible as usual for burning the paper and heat at 
moderate temperature until white. Cover the crucible and heat barely 
to dull redness for 15 minutes, cool and weigh. Since no decomposition 
of the precipitate takes place when it is heated at this temperature 
there should be no change in weight after the first few minutes of 
heating unless washing has not been thorough, leaving salts that 
slowly volatilize. 

Calculate the percent of barium in the barium salt, using the factor 
for barium in barium sulphate. 


If a solution of barium chloride is used as the precipitating 
reagent the sulphate radical may be determined by essentially 
the same process as that just described for barium. A small 
concentration of hydrochloric acid must be maintained as other 
insoluble salts of barium might be formed in a neutral or basic 
solution. Phosphate, carbonate and oxalate may be mentioned 
as common examples of such salts. 

Determination. — Weigh duplicate samples of about 0.25 gm of the 
sulphate into beakers and dissolve in 75 cc of distilled water. Add 1 cc 
of dilute hydrochloric acid, heat to boiling and add, drop-wise and with 
constant stirring, a 5 percent solution of barium chloride until the 
sulphate is completely precipitated. Digest on the steam bath until 
the precipitate settles and the solution clears, then filter and wash with 
hot distilled water, testing finally with dilute sulphuric acid to insure 
removal of barium chloride. 

Heat in a weighed crucible as directed above for the determination of 
barium. From the weight of sample and of barium sulphate calculate 
the percent of the sulphate radical in the sample. In some cases it is 
desirable to calculate the percent of sulphur or of sulphur trioxide. 
This may be done by use of the proper factor. 

Free sulphuric acid may be determined by the same process. 
However, since the reaction of this with barium chloride produces 
free hydrochloric acid it is unnecessary to add any of the latter. 



Strontium is best determined as sulphate, precipitating from a 
solution containing alcohol and an excess of dilute sulphuric 
acid. Its solubility in water at 18° is 114 mg per liter. 1 The 
solubility is considerably diminished by a small excess of sul- 
phuric acid and in 50 percent aqueous alcohol the solubility 
is very slight, although no definite figures are now available. 

Determination. — Weigh the proper quantity of strontium salt to 
produce 0.2 to 0.3 gm of strontium sulphate. If soluble in water dis- 
solve in 50 cc of water and add to the solution 60 cc of alcohol. If 
insoluble in water dissolve in hydrochloric acid, evaporate in a cas- 
serole to expel excess of acid and dilute the solution to 50 cc, then 
add 60 cc of alcohol. Add, slowly and with stirring, dilute sulphuric 
acid until precipitation is complete. An excess of about 5 cc is desirable. 
Stir for some time then allow to stand for 12 hours. Filter and wash 
twice with 50 percent alcohol containing a few drops of dilute sulphuric 
acid, then with 50 percent alcohol until the washings fail to give a test 
for sulphates. Ignite in a weighed crucible at as low a temperature 
as possible until white. Weigh the strontium sulphate and calculate 
the percent of strontium in the sample. 

The separation of barium, strontium and calcium is accom- 
plished by converting all of the metals into nitrates, evaporating 
to dryness and taking up with a mixture of alcohol and ether. 
Calcium nitrate dissolves and the calcium is precipitated as 
oxalate after evaporating the alcohol and ether and dissolving 
the residue in water. The barium and strontium nitrates are 
dissolved in water, barium is precipitated as chromate and stron- 
tium as sulphate as described above. 

Potassium and Sodium 

Potassium may be separated from sodium and determined as 
perchlorate or as chlorplatinate. It may also be precipitated 
as potassium sodium cobaltinitrite and this determined by a 
volumetric process or dissolved and the potassium later deter- 
mined by the other gravimetric methods. It is often stated 
that it can also be determined by weighing as sulphate or as 
chloride. Inasmuch as the latter two methods involve no separa 

1 Kohlrausch: Z. physik. Chem., 50, 355 (1905). 


jtion by precipitation and filtration but simply conversion of the 
[potassium into a form other than the one in which it formerly 
existed and since any other metals that might be present would 
also be converted into sulphates or chlorides and weighed with 
the potassium, the value of these methods is not apparent. Of 
jthe first two methods the perchlorate method has the advantage 
jof cheapness, while the chlorplatinate method is more convenient 
and probably more accurate. 

So far as is indicated by present experience the chlorplatinate 
method is the most reliable and, at the same time, the most 
convenient of all of the known methods for the potassium de- 
termination. The cobaltinitrite method has not yet been im- 
proved to the point of a satisfactory quantitative method, 
although useful in qualitative analysis. The perchlorate method 
has also suffered the drawback of questionable reliability, 
as well as that of requiring a reagent which is more or less 
dangerous to prepare and handle. The one serious obstacle 
to the continued use of the chlorplatinate method is the very 
high cost of the reagent itself. The great scarcity and the ex- 
traordinary rise in the price of platinum since the beginning of 
the world war has made it increasingly desirable, if not absolutely 
essential, that some method that does not involve this metal 
shall be used. In the industrial laboratory a considerable 
portion of the used platinum is systematically recovered but the 
recovery involves considerable work and expense, while a certain 
fraction of the metal is lost in each recovery operation. 

The chlorplatinate method has been varied in matters of detail 
but essentially it consists in precipitating potassium chlorplati- 
nate from an alcoholic solution by adding chlorplatinic acid. 

K 2 S0 4 +H 2 PtCl 6 ^K 2 PtCl 6 +H 2 S0 4 . 

Compounds of sodium and magnesium are not so precipitated 
and potassium is separated from these by filtration but other 
metals must be absent because of the small solubility of most 
chlorplatinates. Potassium chlorplatinate may be either weighed 
as such or ignited in a current of hydrogen : 

K 2 PtCl 6 +2H 2 ->2KCl+Pt+4HCl. 

The potassium chloride is washed out of the reduced platinum 




which is then weighed. In practice potassium chlorplatinate 
usually weighed without ignition. 

Washing. — The separation from alcoholic solution can b 
accomplished only under certain conditions, owing to the sma 
solubilities of certain other compounds that may be presenl 
Sodium chlorplatinate is easily soluble in alcohol but ammonium 
chlorplatinate is soluble to a very slight extent. Ammoniu 
compounds must therefore be volatilized by heating, before th 
addition of chlorplatinic acid. Sodium chloride and sulpha 
are nearly insoluble in alcohol, and must be changed into mo 
soluble substances before a separation of sodium and potassium 
can be accomplished. The following table will serve to show t 
nature of the questions that must here be met. There is con 
siderable disagreement between the results as obtained by dif- 
ferent investigators but these figures may be regarded as at leasl 
approximately correct. 

Percent alcohol 
in water 

Solubility, grams 

salt per liter of 


Sodium sulphate 1 






Sodium chloride 


about 5.0 

Ammonium chlorplatinate 2 




Potassium chlorplatinate 3 





















1 de Bruyn: Z. physik. Chem., 32, 63 (1900). 

2 Fresenius: Z. anal. Chem., 36, 322 (1897). 

3 Archibald, Wilcox and Buckley: J. Am. Chem. Soc, 30, 747 (1908). 


The Lindo Method.- — The method of Lindo 1 consisted in obtain- 
hg a solution containing chlorides of no other metals than sodium 
Imd potassium, adding sufficient chlorplatinic acid to convert 
Itll of the chlorides into chlorplatinates, evaporating nearly to 
Iryness and adding strong alcohol. Sodium chlorplatinate 
lissolves and is separated from the potassium chlorplatinate 
>y filtration. The potassium salt is washed with alcohol and 
veighed or treated as already described. Lindo also showed 
|,hat the solubility of very fine crystals of potassium chlor- 
;)latinate is greater than that of larger crystals (c/. page 21). 
! Using this method, it was necessary that no sulphate should 
pe present because of the very slight. solubility of sodium sul- 
bhate in alcohol. The discussion of the laws of precipitation 
[page 15 et seq.) will make it clear that even if enough chlor- 
platinic acid were present to combine with all of the sodium 
present to form sodium chlorplatinate, the less soluble sodium 
lulphate would still be precipitated. Even when all of the metals 
tre present as chlorides there is without doubt some contamina- 
tion of the precipitated potassium chlorplatinate by sodium 
:hloride. It was early observed 2 that high results would be 
)btained by this method for determining potassium if the factor 
or potassium in potassium chlorplatinate were calculated by the 
lse of the atomic weight of platinum as determined by Seubert 3 
194.8) or even 195.2, which is that given in the table of atomic 
veights for 1918, while a factor calculated from the atomic 
veight 197.2, which had previously been accepted, gave correct 
fesults. This may be partly due 4 to the fact that the pre- 
cipitated potassium chlorplatinate contains also some compounds 
ivith composition as represented by such a formula as H 2 PtCl 5 OH 
bi H 2 PtCl 4 0. It is undoubtedly also partly due to the presence 
pf some sodium chloride in the precipitate. On this account 
j;he method has been modified by using 80 percent alcohol 
nstead of absolute alcohol. Reference to the solubility table 
libove will show that alcohol of this concentration will dissolve 
)otassium chlorplatinate to a greater extent than does absolute 

'Chem. News, 44, 77, 86, 97 and 129 (1881). 
? Dittmar and McArthur: J. Soc. Chem. Ind., 6, 799 (1887). 
3 Ann., 207, 1 (1881). 
Dittmar and McArthur; loc. cit. 


alcohol and this negative error seems practically to balance 
the positive error discussed. 

Gladding's Modification. — The fact that volatile acids, organic 
compounds, ammonium salts, etc., can be easily volatilized 
by heating makes it desirable to obtain the sodium and potassium 
in a form in which they can be heated to redness without danger 
of loss. Both chlorides are sensibly volatile at- such temperatures 
while the sulphates are not but, as already stated, sodium sulphate 
must not be allowed to form because of its small solubility inv 
alcohol. To meet this difficulty Gladding suggested 1 a further] 
modification of the Lindo method in which the sodium sulphate 
was to be washed out, after the removal of the excess of chlor-J 
platinic acid, by a water solution of ammonium chloride. The 
solubility of sodium sulphate in water is considerably increased 
by the presence of ammonium chloride. This follows the general ; 
law that the addition of an electrolyte which does not contain 
an ion in common with the first electrolyte increases its solubility. 
To avoid loss of potassium chlorplatinate the washing liquid is^j 
previously saturated with the pure salt. In using such a solution i 
it is important that no great change in temperature shall occur in< 
the solution after it is withdrawn from the bottle and before it is 
used for washing the precipitate. This is because the solution is 
kept saturated by an excess of potassium chlorplatinate in th$ 
bottle. If the temperature should rise the solution which was 
saturated in the bottle now becomes undersaturated and wi 
dissolve some of the precipitate on the filter. On the other hanc 
if the funnel is at a much lower temperature than the reagen 
bottle (due to working in a colder part of the room) it will coo 
the solution and cause a deposition of potassium chlorplatinat 
upon the precipitate already present. This modified methoc 
known as the Lindo-Gladding method, is now used quite generally 
particularly for the determination of potassium (and indirectl 
of sodium) in industrial products, minerals, etc. 

Decomposition. — During the evaporation of sodium, potassium 
and ammonium salts with sulphuric acid the first product is, 
of course, a mixture of the acid sulphates. After the evaporatic 

1 U. S. Dept. of Agr., Chem. Bull., 7, 38 


of the excess of acid and upon further heating, the pyrosulphates 
of sodium and potassium are formed: 

2KHS0 4 ->K 2 S 2 7 +H 2 0; 
2NaHS0 4 -+Na 2 S 2 7 +H 2 0.' 

These reactions begin at about 350°. At a temperature of 
dull redness the normal sulphates begin to form: 

K 2 S 2 7 -^K 2 S0 4 +S0 3 ; 
Na 2 S 2 7 ^Na 2 S0 4 +S0 3 . 

This decomposition requires heating for some time. It is best 
to test the completion of decomposition by repeating the heating 
until the weight becomes constant. 

Ammonium acid sulphate is decomposed and volatilized as 
ammonia, water and sulphur trioxide : 

NH 4 HS0 4 ->NH3+H 2 0+S0 3 . 

If it is desired to accomplish the removal of ammonium salts 
and organic matter by ignition but to avoid the use of the am- 
monium chloride washing solution the sulphates, first obtained 
by evaporating with sulphuric acid and igniting, are dissolved 
and precipitated by barium chloride which precipitates barium 
sulphate and leaves sodium chloride and potassium chloride in 
solution. The excess of barium is then precipitated by sulphuric 

There is no method known for the direct determination of 
sodium if we exclude the weighing as sulphate or chloride, 
methods of very limited usefulness. This is because no sodium 
compound has sufficiently small solubility to make possible its 
separation from the corresponding salt of potassium. Sodium is 
usually determined by weighing it with potassium in the form of 
sulphate or chloride, determining potassium and calculating 
sodium by difference. It should be noted that such a method 
throws all of the errors of the potassium determination upon that 
of sodium, in addition to any errors that may have occurred in 
the weighing of the combined chlorides or sulphates. 

Ammonium. — Chlorplatinic acid is also used as a reagent 
for the quantitative determination of the ammonium radical but 
potassium must be absent. On account of the difficulty ex- 


perienced in the removal of potassium from ammonium the 
latter is more conveniently determined by volumetric methods. 

Platinum. — The converse of the Lindo method for potassium is 
used for the determination of platinum. Either ammonium 
chloride or potassium chloride may be used as the reagent but 
the former is generally used because of its greater solubility in 
alcohol, which makes the removal of the excess of reagent more 

Determination by the Lindo -Gladding Method. — (To be performed 
in an atmosphere which is free from ammonia). Use portions of about 
0.3 gm of a sample containing salts of potassium and sodium and weigh 
into small weighed evaporating dishes. Dissolve in a small amount of 
hot water, add 0.5 cc of concentrated sulphuric acid, evaporate to dry- 
ness under the hood, using care to avoid spattering, and ignite at bright 
redness until no more white fumes are evolved and the residue is white. 
The steam bath should not be used for the evaporation on account of 
appreciable solubility of porcelain in steam, and consequent loss in 
weight. Cool, weigh and ignite again to constant weight. Record the 
weight of sulphates of sodium and potassium. Dissolve the resulting 
potassium sulphate (mixed with sodium sulphate) in 50 cc of hot water 
and then add chlorplatinic acid, using about 1 cc more than the theoret- 
ical amount, calculated upon th 3 assumption that the original salt was 
potassium chloride. Evaporate on the steam bath to a thick paste 
but not to dryness, cool and add 50 cc of 80 percent alcohol, stir up the 
solid matter and allow to stand, covered, for 30 minutes. 

If the liquid is not visibly colored too little reagent has been used. 
In this case new samples should be taken and the quantity of chlor- 
platinic acid increased. Filter and wash the precipitate thoroughly 
with 80 percent alcohol, washing several times after the washings pass 
through colorless. The wash bottle should be provided with ground- 
glass joints so that no rubber will come into contact with the alcohol. 
Remove the nitrate and washings, pouring these into the bottle provided 
for platinum waste residues, and wash the precipitate with five or six 
portions of 10 cc each of 10 percent ammonium chloride solution which 
is saturated with potassium chlorplatinate. Wash again, thoroughly, 
with 80 percent alcohol, using particular care in washing the upper 
part of the paper free from ammonium chloride. Wash until only 
a faint turbidity is produced by the addition of a drop of silver nitrate 
to the last washings. 

Drain most of the alcohol from the paper, slip the latter out of the 
funnel and dry in the oven at 100°. Place a weighed porcelain crucible 


upon a piece of glazed paper, remove most of the precipitate to the 
crucible, brushing up any particles that may have fallen upon the glazed 
paper, and then replace the paper in the funnel. Place the crucible 
under the funnel and dissolve the remainder of the precipitate in the 
smallest amount of nearly boiling water, allowing the solution to run into 
the crucible. Evaporate to dryness on the steam bath, carefully wipe 
the outside of the crucible with a clean towel and dry for 30 minutes at 
100°. Weigh and calculate the percent of potassium in the salt ana- 
lyzed. Calculate also the weight of potassium sulphate, subtract this 
weight from that of the mixed sulphates, and from the remainder 
calculate the percent of sodium. 

Optional Method, Using a Gooch Crucible. — Proceed as above until 
ready to filter out the potassium chlorplatinate. Prepare two Gooch 
filters as directed on page 86, paying attention to the precautions sug- 
gested, and using strong suction in forming the asbestos felt. Rinse 
the crucible with alcohol, remove, wipe the outside and dry at 100° to 
105° for 30 minutes or until the weight is constant. Weigh and replace 
in the holder. Before the suction pump is again turned on moisten the 
asbestos with one or two drops of water. Start the pump and filter and 
wash the precipitate exactly as above directed. Remove the crucible, 
dry in the oven and weigh. Calculate potassium and sodium as before. 

In the foregoing exercise the procedure is based upon the as- 
sumption that sodium or ammonium salts or both may be pres- 
ent. The latter are volatilized by heating with sulphuric acid. 
The former are removed by the ammonium chloride solution. 

Recovery of Platinum from Waste and Preparation of Chlor- 
platinic Acid. — The recovery and purification of platinum from 
miscellaneous filtrates and other waste solutions is a matter 
of increasing importance on account of the extraordinary increase 
in the price of platinum within recent years. The following 
method, described by Delong, 1 has been found to serve well 
for this purpose. 

Place the solutions in an evaporating dish having a capacity of 2 liters 
for each 100 gm of platinum and evaporate until most of the water 
j has been expelled. Make basic with sodium hydroxide solution and add, 
[ stirring, sodium formate, either solid or in concentrated solution. A 
i quantity of sodium formate equal to about half the weight of platinum 
I will be required. If foaming occurs add more sodium hydroxide. 
! Heat on the steam bath for 1 hour, stirring occasionally, then acidify 

1 Chem. Weekblad, 10, 833 (1914). 



with hydrochloric acid, 25 percent solution, stirring during the addition 
of acid. 

Filter off the precipitated platinum on a soft paper, using suction. 
Wash twice with hot 2 percent hydrochloric acid, then with hot water 
until free from acid. Separate the platinum from the paper, dry, ignite 
and weigh. Pour over the platinum in a porcelain dish five times its 
weight of 25 percent hydrochloric acid, heat on the steam bath and add 
slowly 50 percent nitric acid until no more gas is evolved. About 1 cc 
of nitric acid will be required for each gram of platinum. 

After the platinum is in solution add 10 cc of 25 percent hydrochloric 
acid and evaporate to small volume and repeat this process twice. 
This reduces and eliminates nitric acid. Dilute with water and evapo- 
rate two or three times to expel hydrochloric acid. Finally dilute, 
cool and filter on a soft filter whose approximate weight is known. If 
the filtrate is not perfectly clear refilter. Wash the paper free from 
platinum stain and if any appreciable residue remains, dry and weigh 
it on the filter. Correct the weight of platinum for this weight of car- 
bon, etc., then make the solution to the desired concentration. For 
potassium determinations the solution should contain 0.1 gm of 
platinum in 1 cc. 

A method for the recovery of platinum from scrap by elec- 
trolysis is described by Weber. 1 Chlorplatinic acid prepared by 
this method is quite free from traces of nitric acid. 

The Perchlorate Method. — The perchlorate method is based 
upon the fact that potassium perchlorate is almost insoluble 
in '97 percent ethyl alcohol, while sodium perchlorate dissolves 
with greater ease. It involves the use of an aqueous solution 
of perchloric acid, the preparation of which is somewhat trouble- 
some and dangerous. 

The solubility of potassium perchlorate in alcohol of various 
concentrations is as follows: 2 

. Concentration of 
alcohol, percent by 

Grams, potassium 

perchlorate in one 

liter of solvent 

Grams, potassium 

equivalent to potassium 









1 J. Am. Chem. Soc, 

2 Wenze: Z. angew. ( 

30, 29 (1908). 
}hem., 5, 691 (1891). 


The solubility is considerably diminished by an excess of per- 
chloric acid. Sodium perchlorate dissolves easily in alcohol 
although no definite data are on record. 

The perchlorate method has been improved by the sub- 
stitution of a 20 percent solution of perchloric acid for the pure 
acid, as formerly used. This solution keeps well and involves 
little or no danger of accident through handling. As in the 
chlorplatinate method it is necessary to remove ammonium 
salts. This may be done by gently heating the chlorides or, 
more safely, by evaporating with sulphuric acid and heating 
the sulphates rather more strongly. If the latter method is 
followed it becomes necessary to reconvert the salts to chlorides 
before precipitating potassium perchlorate because of the limited 
solubility of sodium sulphate in 97 percent alcohol, which is used 
for washing the precipitate. 

The analytical method for materials containing salts of only 
sodium, potassium and ammonium is described below. 

Determination. — Weigh 1 gm of the sample in which sodium and 
potassium are to be determined, brushing into a porcelain or platinum 
dish. Treat with sulphuric acid and evaporate and heat to expel 
ammonium salts and excess of acid, using the procedure as described 
for the chlorplatinate method. Dissolve the weighed sulphates of 
sodium and potassium in 50 cc of hot water, add one drop of concen- 
trated hydrochloric acid, heat nearly to boiling and then add, dropwise 
and with constant stirring, a 5 percent solution of barium chloride until 
all sulphate is precipitated. This operation should be performed very 
carefully in order to have the least possible excess of barium chloride 
at the end. Digest for a short time over a small flame or on the steam 
bath, then filter and wash the precipitate and paper well with hot water. 

Evaporate the filtrate and washings to about 25 cc and add 10 cc 
of 20 percent perchloric acid solution. Evaporate over a steam bath 
in a hood until the solution becomes viscous, cool and dissolve the 
residue in a small amount of hot water. Again add 5 cc of perchloric 
acid solution and evaporate over the steam bath until the solution 
evolves dense white fumes of perchloric acid. Cool to room tem- 
perature and add 25 cc of a solution made by mixing 1 cc of 20 percent 
perchloric acid with 100 cc of 98 percent alcohol (making practically 97 
percent alcohol). If the insoluble potassium perchlorate is caked it 
should be broken with a stirring rod so that no soluble salts will escape 
the action of the alcohol. 



During the process of evaporation of the various solutions a Gooch 
filter should be prepared, the asbestos felt being washed with the 
perchloric acid-alcohol mixture. The filter is dried for 1 hour at 120° 
to 130°, cooled and weighed. Filter the solution on this prepared filter, 
removing every trace of the precipitate from the beaker by means of a 
policeman and the prepared washing solution, and wash four or five 
times with this solution. Dry for 1 hour at 120° to 130°, cool and weigh. 

From the weight of potassium perchlorate thus obtained calculate 
the percent of potassium in the sample. Also from this weight calcu- 
late the weight of potassium sulphate which is equivalent to it, subtract 
this weight from the combined weights of sodium sulphate and 
potassium sulphate and from the remaining sodium sulphate calculate 
the percent of sodium in the sample. 


The determination of magnesium is usually made by pre- 
cipitating from a basic solution as dimagnesium ammonium 
orthophosphate. This is ignited and weighed as magnesium 
pyrophosphate. The reactions may be expressed thus : 

MgCl 2 +NH 4 OH+Na 2 HP0 4 -^MgNH 4 P04+2NaCl+H 2 0, 
2MgNH 4 P0 4 ^Mg 2 P 2 7 +2NH 3 +H 2 0. 
No other metals than those of the alkali group may be present, 
as the phosphates of practically all others are insoluble in am- 
monium hydroxide. Any soluble phosphate may be used as 
the precipitating reagent but the ones most used in practice are 
disodium orthophosphate and sodium ammonium acid ortho- 
phosphate (microcosmic salt). 

Solubility. — The. following tabular statement from the work of 
Ebermayer 1 shows the solubility of crystallized dimagnesium 
ammonium orthophosphate in mixtures of ammonium hydroxide 
and water. 

Percent by volume of 

ammonium hydroxide 

of sp. gr. 0.96 

Grams of 

MgNH 4 P0 4 .6H 2 

per liter of solvent at 15° 

Equivalent grams of 
Mg per liter of solvent 




» J. prakt. Chem., 60, 41 (1853). 


This statement of the solubility of the precipitate in solutions 
containing various concentrations of ammonium hydroxide 
would lead to the conclusion that precipitation from the more 
concentrated solutions of ammonium hydroxide would result in 
greater accuracy because of the small solubility of magnesium 
ammonium orthophosphate. From this standpoint alone the 
conclusion would be correct. It happens, however, that the 
basicity of the solution, as well as the presence of other salts, has 
an important influence upon the composition of the precipitate. 
The decrease of solubility with increasing concentrations of 
ammonium hydroxide and also of ammonium salts is to be 
expected as a consequence of the mass law since these substances 

increase the concentration of the ion NH 4 , a constituent of 
the precipitate. Substances other than dimagnesium am- 
monium orthophosphate are precipitated to some extent under 
the following conditions: 

Effect of High Basicity. — If the solution is strongly basic 
when the reagent is added there is formed some trimagnesium 
orthophosphate, Mg3(P04)2, and the quantity of this substance 
is increased by slow addition of the reagent. This is not decom- 
posed upon heating and the ignited precipitate is therefore not 
all magnesium pyrophosphate. This fact makes it undesirable 
that too much ammonium hydroxide should be present, even 
though the solubility of the precipitate is lessened thereby. 

The solubility of dimagnesium ammonium orthophosphate 
is much less than that given by Ebermayer, according to the 
work of Bube 1 who states that the saturated solution in pure 
water contains about 0.00014 gm in 1000 cc. It is also stated 
that in such a solution the solubility product of trimagnesium 
orthophosphate is far exceeded and that the solubility of di- 
magnesium ammonium orthophosphate is increased by large 
concentrations of ammonium ions. This would probably 
account for the increased precipitation of trimagnesium ortho- 
phosphate in solutions made strongly basic by ammonium hydrox- 
ide, the magnesium ammonium salt changing into magnesium 
phosphate and ammonium phosphate, the magnesium salt 
precipitating : 

3MgNH 4 P04^Mg3(P04)2+(NH 4 ) 3 P04. 

1 Z. anal. Chem., 49, 525 (1910). 


Effect of Ammonium Salts. — If the solution contains excessive 
quantities of ammonium salts, whether the precipitation takes 
place from a strongly basic or weakly basic solution the dimag- 
nesium ammonium orthophosphate will contain certain quantities 
of monomagnesium ammonium orthophosphate, Mg(NH 4 )4 
(P0 4 )2. This substance, when strongly heated, passes into 
magnesium metaphosphate, Mg(P0 3 )2, a substance which can 
be converted into magnesium pyrophosphate only after prolonged 
heating at high temperatures (2Mg(P03)2->Mg 2 P207+P20 5 ). 
Ammonium salts should thus be nearly or entirely absent, 
with the exception of a certain amount of ammonium chloride, 
which must be present to prevent the precipitation of magnesium 
hyd oxide. If they have accumulated in the solution as a result 
of the use of ammonium hydroxide in the separation of other 
metals, they should be removed before precipitation, by (a) 
evaporating to dryness and heating strongly or (b) evaporating 
to small volume and heating with concentrated nitric acid or (c) 
performing a double precipitation, dissolving the first -impure 
precipitate in hydrochloric acid and reprecipitating. Method 
(a) or (c) is to be preferred. 

Temperature and Rate of Precipitation. — Most chemists 
prefer to precipitate magnesium from a cold solution although 
Gibbs 1 recommends a boiling solution. Whether the cold or 
hot solution is used one of two procedures may be followed in 
order to conform to the principles outlined above. The entire 
amount of disodium phosphate solution may be added at once 
to an acid solution and then dilute ammonium hydroxide slowly 
added until the solution is basic. After standing a short time 
most of the precipitate will form and the remaining magnesium 
can be precipitated by the addition of concentrated ammonium 
hydroxide. Instead of following this method the solution may 
be made neutral or faintly basic and disodium phosphate added 
slowly, thus precipitating nearly all of the substance, when 
concentrated ammonium hydroxide may be added as before. 
The second method is recommended. 

Decomposition upon Heating.— The reason for the difference 
in the mode of decomposition of the precipitate containing more 
of the ammonium radical from that of the dimagnesium am- 

1 Am. J. Sci., [3] 5, 114. 


jmonium salt is apparent when the properties of the three phos- 
phoric acids and of the salts are examined. Phosphorus pent- 
oxide, by combining with different proportions of water, gives 
rise to three different acids : 

P2O5+ H 2 0-^2HP03, metaphosphoric acid, 
P205+2H 2 0-^H 4 P 2 7 , pyrophosphoric acid, 
P 2 5 +3H 2 0->2H 3 P0 4 , orthophosphoric acid. 

Either metaphosphoric or pyrophosphoric acid will be trans- 
formed into the one containing more water if allowed to stand in 
solution. Also the acids may be changed in the opposite sense 
by heating. At about 213° orthophosphoric acid loses water 
and yields pyrophosphoric acid. 

2H 3 P0 4 ->H 4 P 2 7 +H 2 0. 

At about 400° pyrophosphoric acid loses one molecule of water 
and yields metaphosphoric acid. 

H 4 P 2 07-^2HP0 3 +H 2 0. 

When heated to higher temperatures the remaining molecule of 
water is lost and phosphorus pentoxide remains. 

2HP0 3 ->P 2 5 +H 2 0. 

Phosphorus pentoxide is thus seen to be the final product of any 
of the three acids when the acid is heated to a high temperature 
and this is because a volatile substance (water) is produced by 
heating. Just as the acids are compounds of phosphorus pent- 
oxide and water, so the salts may be regarded as compounds 
of phosphorus pentoxide and metallic oxide (which is analogous 
to hydrogen oxide). Consequently the extent to which the 
salts may be decomposed by heating will be conditioned by the 
nature of the metallic oxide or, in other words, by its degree of 
volatility. Thus the normal phosphates of sodium, potassium, 
magnesium, calcium, etc., are not decomposable at all, except 
at extremely high temperatures where phosphorus pentoxide 
begins to be volatile, while the acid phosphates of these metals 
are decomposable to whatever extent is denoted by the propor- 
tion of water that may be formed. Ammonium salts are con- 
verted completely into phosphorus pentoxide because, instead 
of the hypothetical metallic oxide, (NH 4 ) 2 0, there are formed 



ammonia and water and both of these substances are volatile 
Orthophosphoric and pyrophosphoric acids are polybasic and a 
considerable variety of salts may be prepared, containing varyin; 
amounts of metals, ammonium and hydrogen, so that they ma 
be regarded as containing varying amounts of metallic oxide, 
ammonia, water and phosphorus pentoxide. 

Examples. — The composition and decomposition of the three 
phosphoric acids and typical examples of their salts are shown in 
the following statement: 


Composition Decomposition by Heat 


2HP0s.=C=H 2 0.P20 6 

H 4 P207O2H20.P20 6 
2H 3 P0 4 ^:3H20.P205 

2HP0 3 ->P206 + H 2 
H 4 P 2 07-»-P20 6 + 2H20 
2HsP04->P206+3H 2 

Normal Potassium Salts 

2KP03 OK2O.P2O6 

K4P2O7 =C= 2K 2 0.P20 8 

Not decomposed except by slow loss of 
P2O8 at high temperatures 

Potassium Acid Salts 

No acid metaphosphate possible 
2K2HPO4 =C= 2K2O.H2O.P2O5 
2KH2PO4 =C= K2O.2H2O.P2O5 

K2H2P2O7-* 2KP0 8 + H2O 
2K 2 HP04-»K4P20 7 + H2O 
2KH2PO4-+2KPO3 + 2H2O 

Potassium Ammonium Salts 

No double metaphosphate possible 
K 2 (NH4)2P20 7 O K 2 0.(NH 4 )20.P20 6 
2K2NH4PO4 O 2K20.(NH 4 )20.P20 8 
2K(NH 4 )2P04 OK 2 0.2(NH4)20.P20 5 

K2(NH 4 )2P207-+2KP03-f-2NH 3 + H20 
2K 2 NH4P04->K4P20 7 + 2NH 3 + H20 
2K(NH4)2P04-*2KP03 + 4NH3 + 2H 2 

Monomagnesium ammonium orthophosphate, Mg(NH 4 )4(P0 4 ): 
is analogous to monopotassium ammonium orthophosphate 
K(NH4) 2 P0 4 , as may be seen from the structural formulae: 

NH 4x 

nhApo 4 

NH 4 -PO, 


NH 4 ^ 


NH 4 ^P0 4 

NH 4 / 

Its decomposition can therefore proceed as far as magnesium 
metaphosphate : 

Mg(NH 4 ) 4 (P0 4 )2-+Mg(P03)2+4NH 3 +2H 2 0.. 


jvhile dimagnesium ammonium orthophosphate can decompose 
pnly as far as the pyrophosphate: 

2MgNH 4 P04-*Mg 2 P 2 7 +2NH 3 +H 2 0. 

irhis makes it necessary that such conditions shall be maintained 
'is will make possible the formation of but one double salt, in 
order that the composition of the ignited precipitate may be 
definite and constant. 

I Rate of Crystallization. — The complete precipitation of 
liimagnesium ammonium orthophosphate takes place only 
tafter standing for some time. Formerly it was considered neces- 
sary to allow 24 hours for the action to proceed. It is now gener- 
ally considered that from 2 to 3 hours is sufficient for the pre- 
cipitation of all but a minute amount, negligible under ordinary 
(circumstances. The usual method of testing with excess of 
|reagent, to determine whether precipitation is complete, is 
jrendered useless because of the slow crystallization of the pre- 
cipitate unless several hours are allowed for the possible pre- 
cipitation of small amounts. The crystalline precipitate may be 
readily filtered and washed by a dilute solution of ammonium 
[hydroxide or ammonium nitrate. The precipitate is appreciably 
soluble in distilled water. An application of the laws of solubility, 
discussed under the head of "Precipitation" (page 15) would 
Slead to the conclusion that any one of the three classes of soluble 
[^compounds: phosphates, ammonium salts or magnesium salts, 

Swill lessen the solubility of ammonium magnesium phosphate 

+ ++ — 

jsince the latter dissociates into the three ions NH 4 , Mg and P0 4 . 

BThe addition of magnesium salts to the washing fluid is clearly 

lout of the question if magnesium is to be determined. 

Phosphates must themselves be removed by washing because 
| only the ammonium phosphates are entirely volatile and these 
ionly with some difficulty. Either ammonium hydroxide or 
I ammonium nitrate is suitable for the purpose, excess of either 

being driven off during drying and ignition of the precipitate. 

Ignition. — Considerable difficulty is often experienced in 
' obtaining pure, white magnesium pyrophosphate by igniting the 
| magnesium ammonium orthophosphate. This is usually due 
l to imperfect washing, sodium phosphate being left in the pre- 


cipitate. Upon heating, traces of the salt cause partial fusion, 
particles of carbon being enclosed and oxidation made difficult. 
Thorough washing followed by long heating at high temperatures 
is the only remedy. 

A similar method may also be used for the determination of 
arsenic acid, the precipitate of magnesium ammonium arsenate 
being heated at a moderate temperature until it forms magnesium 
pyroarsenate : 

2MgNH4As04-+Mg 2 As 2 07+2NH3+H 2 0. 

Determination. — Weigh into Pyrex beakers portions of about 0.3 gra 
of a magnesium salt. If the salt is soluble in water dissolve in about 
100 cc of distilled water and add a drop of concentrated hydrochloric 
acid. If not soluble in water dissolve in hydrochloric acid (1 part of 
concentrated acid to 1 part of water), warming if necessary. Cool and 
drop in a very small piece of litmus paper and then add, slowly and with 
stirring,- dilute ammonium hydroxide until the solution is faintly basic. 
Now add from a pipette, slowly and with stirring, 15 cc of a clear 
10 percent solution of disodium orthophosphate. Allow to stand for 
15 minutes until a considerable part of the precipitate has appeared, 
then add concentrated ammonium hydroxide solution (sp. gr. 0.90), in 
such quantity that the solution shall finally contain ammonium hydrox- 
ide equivalent to one-ninth of its total volume. Cover and allow to 
stand for three hours or stir continuously for 30 minutes. 

Filter the precipitate on a filter of extracted paper, in a weighed 
platinum Gooch crucible or in an ignited and weighed alundum crucible 
and wash until free from chlorides with a solution containing 2 percent 
of ammonia or 5 percent of ammonium nitrate, finally testing the 
washings with silver nitrate after acidifying with nitric acid. If a 
Gooch crucible has been used place the cap on the bottom and heat 
over the burner until dry, then over the blast lamp for 20 minutes. 
An alundum crucible is treated similarly. If a paper filter was used 
remove the paper from the funnel and, if sufficient precipitate is present 
to make its removal from the paper feasible, dry and remove most of 
the precipitate to a sheet' of glazed paper, refold the paper and place 
in a weighed porcelain or platinum crucible. Incline the crucible with 
the cover leaned against it and heat gently over the burner until the 
paper is completely burned and the precipitate is nearly white. After 
the precipitate is white or gray the main portion is added and the cruci- 
ble is heated for 20 minutes over the blast lamp, cooled in the desiccator 
and weighed. From the weight of magnesium pyrophosphate calculate 
that of magnesium and the percent of magnesium in the original sample. 



For the precipitation of the phosphate radical as magnesium 
ammonium phosphate it is necessary that no metal that can 
form an insoluble phosphate shall be present. In case this 
condition is not fulfilled a preliminary separation of the phosphate 
radical is made by precipitating ammonium phosphomolybdate 
from a solution containing free nitric acid. This operation is 
described on page 453 and following, in the discussion of the 
determination of phosphorus in steel. A sample containing 
no metals other than those of the alkali group does not require 
this treatment. This simpler determination will now be 

Determination. — Prepare a solution of "magnesia mixture" as follows: 
Dissolve 55 gm of crystallized magnesium chloride and 140 gm of 
ammonium chloride in water, add 130 cc of ammonium hydroxide 
(specific gravity 0.90) and dilute to 1000 cc. If this solution is kept 
in stock for any considerable time it will acquire a flocculent precipitate 
of hydrated silica, derived from solution of the glass by the base. The 
solution must be clear when used. This condition may be insured by 
filtering the solution or by preparing only enough of the reagent to last 
a short time. 

Weigh duplicate samples of 0.2 to 0.4 gm of the phosphate into beakers 
of resistance glass, dissolve and dilute to 75 cc. Drop in a very small 
bit of litmus paper and if a basic reaction is not shown add dilute 
ammonium hydroxide until the paper becomes blue. Add 10 cc of 
a 10 percent solution of ammonium chloride, mix and then add, 
very slowly, "magnesia mixture" sufficient in quantity to precipitate 
all of the phosphate. As the precipitate does not form rapidly in a 
barely basic solution it is not always easy to determine when enough 
of the reagent has been added. It is then best to use what is thought 
to be a good excess and to rely upon testing the filtrate which is obtained 

The rest of the procedure is exactly the same as in the determination of 
magnesium, with the single exception that " magnesia mixture" instead 
of sodium phosphate solution is used in testing for the completion of 
precipitation. From the weight of magnesium pyrophosphate finally 
obtained calculate the percent of the phosphate radical, of phosphorus 
pentoxide or of phosphorus, the report depending upon the nature of the 
sample examined. 


The method of Gibbs for manganese depends upon the same 
chemical principles as are involved in the determination of 
magnesium. A soluble orthophosphate is added to the solution 
of the manganese salt and the solution is then made basic with 
ammonium hydroxide. Dimanganese ammonium orthophos- 
phate is precipitated and this, when ignited, gives manganese 

MnS04+Na 2 HP04+NH40H-^MnNH4P04+Na 2 S04+H 2 0, 
2MnNH 4 P04-^Mn 2 P 2 07+2NH 3 +H 2 0. 

If the manganese is already in its lowest state of oxidation, 
precipitation is accomplished without further change. If it 
is in the form of a manganate or permanganate or of manganese 
dioxide it is first reduced by sulphurous acid: 

2KMn04+5H 2 S0 3 ->2MnS0 4 +2H 2 S04-|-K 2 S04+3H 2 0; 
MnO 2 +H 2 S0 3 ->MnS04+H 2 0. 

Determination. — Weigh enough of the sample to contain about 0.1 
gm of manganese. If the sample is a permanganate or manganese di- 
oxide dissolve in 50 cc of a saturated solution of sulphurous acid, con- 
taining also 1 percent of hydrochloric acid, filter if necessary and boil 
to expel the excess of sulphur dioxide. 

If the sample is a soluble manganese salt omit the treatment with 
sulphurous acid. In either case proceed as follows : 

Add to the solution in a Pyrex beaker, 3 percent more than the quan- 
tity of 10 percent disodium phosphate solution calculated to be neces- 
sary for complete precipitation of the manganese. Heat to boiling and 
add dilute ammonium hydroxide solution, drop by drop with constant 
stirring, until a precipitate begins to form. Boil and stir until this pre- 
cipitate becomes crystalline, then add another drop of ammonium hy- 
droxide and stir and boil until the additional amorphous precipitate 
becomes crystalline. Continue this process until further addition of 
ammonium hydroxide produces no precipitate. All of the precipitate 
should now be in the crystalline condition. Add 0.5 cc excess of ammo- 
nium hydroxide, then cool the solution by placing the beaker in ice-water. 
Filter and wash with a clear, slightly basic, 10 percent solution of am- 
monium nitrate or a 2 percent solution of ammonium hydroxide until 
free from chlorides, then ignite with the same precautions as were ob- 
served in the ignition of magnesium ammonium phosphate. From the 
weight of manganese pyrophosphate obtained calculate the percent of 
manganese in the sample. 


Chlorine, Bromine and Iodine 
The members of the halogen group may occur in different 
forms, requiring different methods of procedure. This occur- 
rence may be as free halogens, as oxyacids or salts, as hydracids 
or salts, or as organic compounds. The gravimetric determina- 
tion of the negative radical of the halogen hydracids is invariably 
made by precipitating and weighing the silver salt. The 
solubilities of the latter, as well as the principles involved in 
the precipitation, washing and ignition, were discussed under the 
description of the determination of silver. The procedure is also 
similar to that involved in the determination of silver, the silver 
salt (silver nitrate) being, in this case, the precipitant while the 
chloride, bromide or iodide is the substance being investigated. 
Separation of Chlorine and Iodine. — If chlorides and iodides 
occur together the iodine may be precipitated as palladious iodide 
by a solution of palladious chloride. PdCl 2 . In another portion 
the total halogen may be precipitated by silver nitrate and 
weighed as a mixture of silver iodide and chloride. The proper 
weight of silver iodide, as calculated from the weight of pal- 
ladious iodide found, is subtracted and chlorine calculated from 
the remainder. 

Indirect Method. — The chlorine and iodine may be determined 
indirectly by precipitating by excess of silver nitrate, weighing the 
mixed chloride and iodide, then converting the silver iodide into 
chloride by heating in an atmosphere of chlorine and reweighing. 
If x = weight of chlorine, a = weight of silver chloride and 
iodide, b = weight after conversion of silver iodide into chloride 

and y = weight of iodine, then ^ ' x = weight of silver chloride 

and -, 9 ' y = weight of silver iodide, therefore : 

143.34 234.80 _ v,' 

35.46 x +\2§.§2 v ~ a ' W 

143,34 14^34 

35.46 X± 126. 92^ °' W 

91 46 
Subtracting (2) from (1), j^ilM 2/ = a-&« 

Then y=1.3877(a-6), (3) 

x = 0.63506 -0.3877a. (4) 



In a manner similar to that shown above for chlorine and iodine, derive 
the following formula?: 

2. z = 2.7855 6-1.7004 a, 
?/=2.7004a -3.5380 6, 

where x= weight of bromine, y= weight of iodine, a = weight of silver 
bromide and iodide and 6= weight of silver chloride after chlorination. 

3. re = 1.0451 6-0.7976 a, 
y = 1.7974 (a -6), 

where x = weight of chlorine, y= weight of bromine, a = weight of silver 
chloride and bromide and 6= weight of silver chloride after chlorination. 

Determination of Two Halogens in Mixed Halides, Indirect Method. 

— Use about 0.5 gm of sample. Dissolve in 75 cc of water, add 0.5 cc 
of dilute nitric acid and then, drop by drop and with constant stirring, 
a slight excess of 5 percent silver nitrate solution. 10 to 30 cc will 
be sufficient. Digest at near the boiling temperature until the precipi- 
tate settles readily, leaving a clear supernatant solution. Test for 
completion of the precipitation then filter in a prepared and weighed 
Gooch crucible and wash free from excess of silver nitrate. Dry at 105° 
to constant weight. 

Ignite a small porcelain boat and cool in a desiccator. Place this on a 
sheet of black glazed paper and carefully remove the asbestos filter 
and every particle of the silver halides from the filter, placing these in the 
boat. Allow this to remain in the desiccator for 15 minutes, then weigh. 
From this weight subtract that of the silver halides, already found. 
This gives as the remainder the weight of the boat plus asbestos. 

Prepare a chlorinating apparatus according to Fig. 34. 

A is a distilling flask of 100 cc capacity, fitted with a funnel tube which 
reaches to the bottom of the flask. The latter is connected at the side 
with a glass combustion tube, B, of 12 to 15 mm internal diameter 
and 40 cm length. Corks are used in the ends of the combustion tube. 

Place the entire apparatus under a hood. Insert the boat containing 
the silver halides and asbestos into the middle of the tube. Place 
10 gm of potassium permanganate in the flask and then pour into the 
tube 5 cc of concentrated hydrochloric acid. Warm, if necessary, to 
start the reaction and when the tube is filled with chlorine carefully heat 
directly under the boat, using a wing burner. Add more acid to the 
flask as it may become necessary, in order to maintain a slow evolution 
of chlorine, and heat the boat for 15 minutes to a temperature just under 
the fusing point of the silver chloride. Finally raise the temperature 



luntil fusion barely begins, then remove the flame. Disconnect the tube 
from the flask, attach a tube of calcium chloride to one end and slowly 
draw air through until the tube is quite cool. Remove the boat from 
the tube and test the odor to determine whether all free chlorine has 
been removed. Dry for 30 minutes at 105°, cool in the desiccator and 
weigh. From this weight subtract that of the boat plus asbestos. The 
remainder is the weight of silver chloride. From this and the weight of 
the mixed silver halides and of the sample calculate the percent of each 
halogen in the sample, using one of the formulas derived above . 
The entire experiment should be conducted away from bright light. 



Fig. 34. — Apparatus for chlorination of mixed halides of silver. 

Chlorine and iodine may also be separated by the method of 
graded oxidation, to be described in connection with the separa- 
tion of the three halogens. 

Separation of Bromine and Iodine. — Mixtures of bromides 
and iodides may be analyzed by methods similar to those de- 
scribed above. Palladious bromide is sufficiently soluble to 
make possible the precipitation of palladious iodide in the pres- 
ence of the bromide. Also either the indirect analysis or the 
method of graded oxidation may be used. 

Separation of Chlorine and Bromine. — For mixtures of 
chlorides and bromides either the method of graded oxidation 
or that of indirect analysis may be used. For the latter silver 


chloride and bromide are weighed together and the bromide 
then converted into silver chloride and reweighed. 

Separation of Chlorine, Bromine, and Iodine, by Graded Oxi- 
dation. — In the discussion of the decomposition voltages of 
electrolytes (page 141) it is shown that any electrically neutrs 
element that is capable of ion formation will, when placed ii 
contact with a solution containing its ions, generate a definite 
potential difference whose magnitude depends upon the solution 
tension of the element and the concentration of its ions already in 
solution. Two such systems will generate a definite electro- 
motive force if external connection is made between the non- 
ionized elements and if the two solutions are brought into contact. 
This electromotive force is always in the direction that would 
cause a current to flow externally from the element having the 
less solution tension (if its ions are positive, or the greater if 
its ions are negative) to the one having the greater solution 
tension. When a metal passes into solution and forms positive 
ions it is thereby oxidized. When an element capable of negative 
ion formation passes into solution and forms ions it is thereby 
reduced Conversely, when metallic ions are converted into mas- 
sive, uncharged metal they are reduced and when non-metallic or 
negative ions are discharged they are thereby oxidized. Ac- 
cording to this view oxidation consists in the addition of positive 
charges or the removal of negative ones, while reduction is the 
addition of negative charges or the removal of positive ones. 

Thus the change : Fe— >Fe (change of the ferrous ion to the ferric 

ion) is oxidation, while the reverse is reduction and Mn0 4 — > 

Mn(>4 (change of the permanganate ion to the manganate ion) 
is reduction, while the reverse is oxidation. It follows from 
these statements that the metal having the greater solution 
tension is the stronger reducing agent while the non-metallic 
element having the greater solution tension is the better oxidizing 
agent. It thus becomes possible to compare the activities of 
two oxidizing or reducing agents by measuring the magnitude 
and direction of the electromotive force produced by combining 
two systems made up of these agents in contact with solutions 
containing the respective products of reduction or oxidation, 


If the oxidizing or reducing agents are solids the electrodes 

ire composed of these solids. If gases the electrode is of some 

material that will superficially dissolve the gases, while if the 

igents are solutions they are merely brought into contact with 

blectrodes made of indifferent metals, such as platinum. Thus 

jjilver as a reducing agent would itself be made to form the 

alectrode material, in contact with an ionized silver salt. Oxygen 

■ an oxidizing agent would be caused to bubble into the solution 

In contact with an electrode of platinum which is coated with 

jplatinum black, in which oxygen dissolves, and this would be 

immersed in a solution containing hydroxyl ions. Potassium 

permanganate would be used in simple contact with a platinum 

: (electrode and the solution would also contain the positive 

manganese ion, Mn, the product of reduction of the ion Mn04. 

!|In the last case the force, solution tension, is replaced by the 

(tendency of the permanganate anion, already in solution, to 

'become reduced. 

Oxidation Potential. — All of the elementary halogens are 

J oxidizing agents because they exhibit a tendency toward negative 

ion formation: 

C1->C1, Br-*Br, etc. 

Conversely the change of halogen ions into neutral elements is an 
oxidation of these ions. In order to bring about this oxidation 
it is necessary to apply another oxidizing . agent whose " oxida- 
tion potential, " F (potential of the nonionized electrode minus 
that of the electrolyte) is greater. This is analogous to the 
decomposition of a halide by means of the current, which is an 
oxidizing agent for the negative ion and a reducing agent for 
the positive ion. 

Selective Oxidation. — If an oxidizing agent can be found, 
having an oxidation potential greater than that of one of the 
halogens and less than that of another this agent may be used 
for separating the two halogens by oxidation of the salt (or acid) 
of one, removing the liberated element by distillation and leaving 
the other, which was incapable of being oxidized by this agent. 
This is analogous to the electrolytic separation of metals by 
grading the electromotive force which is applied to two electrodes 


in the solution. The measurement of oxidation potentials! 
should therefore furnish valuable information for assisting m 
the selection of oxidizing agents suitable for graded oxidation 
of the anions of the halogen hydracids. Bancroft has shown 1 
that the following differences exist between the oxidation po- 
tentials of chlorine, bromine and iodine in solutions of sate 
of their respective hydracids: 

Chlorine in potassium chloride — bromine in potassium bro- 
mide =0.241 volt. 

Bromine in potassium bromide — iodine in potassium iodide = 
0.535 volt. 

These differences vary somewhat if the concentrations are^ 

As examples of oxidizing agents which will serve for the iodine 
anion without thereby oxidizing bromine or chlorine anions, 
may be mentioned monopotassium arsenate and nitrous acid. 
These were suggested by Gooch. 2 The oxidation potential of 
nitrous acid (or potassium nitrite and sulphuric acid), according 
to the measurements of Bancroft, is 0.249 volt higher than that 
of iodine in potassium iodide, and 0.285 volt lower than that 
of bromine in potassium bromide. 

For the selective oxidation of the bromine anion in presence 
of the chlorine anion the following substances have been used: 
potassium permanganate in acid solution, lead peroxide in acid 
solution, potassium dichromate in acid solution, ammonium 
persulphate in neutral solution and potassium iodate in acid 
solution. With the exception of the last named oxidizing agent 
all of these substances possess oxidation potentials higher than 
that of chlorine in potassium chloride and could therefore be- 
made to serve for a quantitative separation of bromine and 
chlorine only by carefully regulating the concentrations of oxi- 
dizing and reducing agents and by stopping the distillation at 
exactly the correct time. Attempts have been made to regulate 
the speed of oxidation by acidifying with substances that can 
furnish only a small concentration of hydrogen ions, since it is 
only in presence of hydrogen ions that the reactions can proceed. 
Weakly acid substances that have been used for this purpose 

1 Z. physik. Chem., 10, 387 (1892). 

2 Chem. News, 61, 235 (1890). 


are acetic acid, potassium acid sulphate, ferrous sulphate and 
aluminium sulphate. The last two are weakly acid through 

Potassium permanganate is the oxidizing agent used by Jan- 
nasch and Aschoff 1 and the oxidation of hydrochloric acid is 
prevented by acidifying with an acid no stronger than acetic 
acid and by employing a large dilution. The oxidation potential 
of potassium permanganate with sulphuric acid is 0.097 volt 
higher than that of chlorine with potassium chloride. 

The reduction of potassium permanganate is really a reduction 
of manganese itself, being a change of heptavalent into bivalent 
manganese. The complete equation is 

KMn0 4 +5HBr+3HC 2 H30o->KC 2 H 3 2 +Mn(C 2 H,0 2 ) 2 + 

5Br+4H 2 0. 

The ionic change involving manganese is 
Mn0 4 ->Mn+20 2 . 

The electrical change of manganese itself is not oxidation, as 

would appear from the last equation, butjreduction, because the 

univalent anion, Mn(>4, is composed of one atom of heptavalent 

positive manganese, and four atoms of bivalent negative oxygen, 

so that the change of manganese is really 

V+"+ + + + 

If the substance being analyzed is known to be a pure mixture 
of only two of the halides, one of the halogens may be liberated 
and removed by distillation without subsequent absorption, the 
other being determined in the residual solution. If it is not a 
pure mixture or if it contains salts of three halogens it is necessary 
to absorb at least one of these and make a direct determination 
of it in the absorbing solutions. 

Bugarszky 2 used potassium iodate and sulphuric acid for 
separating bromine and chlorine, distilling the bromine without 
absorption. The oxidation potential of acidified potassium 
iodate was found by Bancroft to be 0.064 volt higher than that 

1 Z. anorg. Chem., 1, 144 and 245 (1892); 5, 8 (1894). 

2 Ibid., 10,387 (1895). 


of bromine and 0.178 volt lower than that of chlorine. The 
reaction is as follows : 

KI0 3 +5KBr+3H 2 S0 4 ^3K 2 S0 4 +5Br+I+3H 2 0. 
Both bromine and iodine are liberated and distilled and account 
of this must be taken if the free halogens are absorbed and subse- 
quently determined. Chlorine is determined in the residual 
solution, after reducing the excess of iodic acid to hydriodic acid 
by means of sulphurous acid, then oxidizing by nitrous acid and 

Andrews 1 modified the method of Bugarszky by substituting 
nitric acid for sulphuric acid and by reducing the excess of iodic 
acid by means of phosphorous acid. His method was not tested, 
however, except for the determination of chlorides in crude 
bromides and of chlorine in crude bromine. In both cases the 
chlorine was present in relatively small quantities (less than 10 
percent) and it was not adapted to the determination of both 
bromine and chlorine. For the determination of chlorine, 
bromine and iodine by direct means ; the method of Jannasch 
and Aschoff is probably the best of all methods yet proposed, 
even though permanganic acid is not an ideal oxidizing agent 
for the separation of chlorine and bromine. In this method 
the solution of mixed chlorides, bromides and iodides is first 
acidified with sulphuric acid and potassium nitrite is added: 

KI+KN0 2 +2H £ S0 4 ->2KHS0 4 +NO+H 2 0+I. 
The liberated iodine is distilled, absorbed, and subsequently 
determined. The sulphuric acid is. then neutralized by sodium 
hydroxide, acetic acid and potassium permanganate are added 
and the bromine is distilled, absorbed and determined. In the 
residual solution the excess of permanganate is reduced and the 
chlorine is determined gravimetrically. 

The absorbent which best serves for iodine and bromine is a 
solution containing sodium hydroxide and hydrogen peroxide. 
Bromine and iodine react with sodium hydroxide to form sodium 
bromide and sodium hypobromite in the one case and sodium 
iodide and sodium hypoiodite in the other: 

2NaOH + 2Br~>NaBr + NaBr + H 2 0, 
2NaOH+2I-»NaIO+NaI+H 2 0. 
1 J. Am. Chem. Soc, 29, 275 (1907). 



(f the solutions are allowed to stand for some time bromates and 
odates are formed: 

3NaBrO-+NaBr0 3 +2Na/Br, 
3NaIO->NaI0 3 +2NaI. 

The last change does not take place if hydrogen peroxide is 
>resent and the solution is kept cold, the hypobromite and 
lypoiodite being reduced as fast as formed: 

NaBrO+H 2 2 -->NaBr+H 2 0+0 2 , 
NaIO+H 2 2 ->NaI-f-H 2 0+0 2 . 

f in the resulting solutions bromine and iodine may be determined 
ei is the silver salts in the usual manner after acidifying with 
;lf Sulphuric acid. 



Fig. 35. — Apparatus for the separation of the halogens. 

If the absorbing solution has been allowed to become warm 
^ some iodate or bromate will be formed. In this case acidification 
will cause the liberation of free halogen which will escape precipi- 
tation. If silver nitrate is added before acidifying, any iodate 
a br bromate will remain in solution as the silver salt. This inter- 
ference of oxysalts may also be prevented by the addition of a 
flpulphite before the addition of acid. The resulting sulphurous 
111 iacid then reduces the iodate or bromate to iodide or bromide. 
During the distillation of bromine and iodine it is essential 
that contact with cork or rubber be avoided, since the halogens 
are thereby reduced and absorbed. Ground-glass stoppers are 
necessary in all parts of the apparatus where such contact would 


occur and, where rubber connections are used, the glass tubes 
inside must be pushed together so as to expose as little of the 
rubber tubing as is possible. All reagents must be tested and 
found free from the halogens. 

Determination. — Weigh about 1 gm of the mixture of halides, placing 
the sample in a round bottomed, glass stoppered distilling flask, having 
a capacity of 1000 cc, and having an inlet tube sealed into the side 
of the neck and reaching to the bottom of the flask. Connect the appa- 
ratus as shown in Fig. 35. A is a vessel in which steam may be generated, 
B is the distilling flask, C and D are bubble tubes having a capacity 
of 150 cc. The tube a should reach to the bottom of the steam generator 
and should extend about 18 inches above. This tube provides an inlet 
for air, in case there is any tendency toward drawing liquid back from B. 

Each of the absorption tubes C and D contains 50 cc of 5 percent 
sodium hydroxide and 50 cc of hydrogen peroxide. The union between 
B and C and between C and D should be made by bringing the glass 
tubes quite together inside the rubber connections. 

Dissolve the weighed sample in about 600 cc of water, add 5 cc of 
25 percent sulphuric acid and 2 gm of sodium nitrite. Heat the solution 
nearly to boiling and pass steam through the flask for twenty minutes 
after the solution is colorless. During this time the tubes C and D 
must be kept cool by immersion in ice water. 

When all of the iodine has been distilled the boiling is interrupted, 
the absorption tubes are disconnected and their contents washed into 
a 300 cc beaker. The tubes are then returned to the apparatus and 
are refilled with sodium hydroxide and hydrogen peroxide as before. 
The solution in the flask is barely neutralized with sodium hydroxide 
solution and evaporated to a volume of about 500 cc. 1.5 gm of 
potassium permanganate and 60 cc of 33 percent acetic acid are added 
and the bromine thus liberated is distilled and absorbed in hydrogen 
peroxide and sodium hydroxide in the cooled tubes. When steam has 
been passed through the solution for some time after the latter has 
become colorless the distillation is again stopped and the contents of 
the tubes washed into another beaker. 

The solutions now containing the iodine and bromine are boiled until 
the excess of hydrogen peroxide is completely decomposed. 0.5 gm 
of sodium sulphite is added and then dilute sulphuric acid until the solu- 
tion is slightly acid in character. If any color appears at this point it is 
due to the presence of iodine or bromine produced by iodate or bromate, 
showing that insufficient sodium sulphite has been added. In this case 
0.5 gm more is at once added to reduce the free halogen. 


When the solutions are acid and colorless a 5 percent solution of silver 
nitrate is added, drop by drop from a pipette, stirring vigorously until 
no further precipitation occurs. The liquid is digested at near the boil- 
ing temperature until the precipitate settles readily, after which it is 
filtered on a Gooch crucible, as directed on page 86, and the precipitates 
are washed free from silver nitrate, testing the washings with dilute 
hydrochloric acid. The crucibles are finally washed once with alcohol 
to promote rapid drying and are then dried at 110° for one-half hour or 
until the weight is constant. The percent of iodine and of bromine is 

The solution in the large distilling flask is boiled with alcohol to reduce 
the excess of potassium permanganate and is then poured into a beaker 
or evaporating dish and evaporated to a volume of not more than 150 cc. 
5 cc of dilute nitric acid is added and the chlorine is precipitated and 
weighed exactly as directed in the case of iodine and bromine. 

Halogen Oxyacids. — The oxyacids of the halogens (or their 
salts) may be reduced to the hydracids by warming with hydrogen 
peroxide, after which the separation and determination may be 
accomplished as above directed. 

Free Halogens existing in solution may be converted into 
oxysalts by treatment with alkali bases, after which their separa- 
tion and determination may be carried out by methods already 
discussed. Their determination is more conveniently made by 
volumetric methods which will be discussed later. Chlorine in 
gaseous mixtures is also determined by absorption followed by a 
volumetric process. 

Organic Halogen Compounds. — Compounds of the halogens 
with organic residues cannot be analyzed by the usual methods 
because such compounds do not, as a rule, ionize to form the 
anions of the halogen acids. The compound must be decomposed 
in such a manner as to leave the halogen in the form of an in- 
organic compound of one of the well-defined acids. In such 
cases either the lime method or the method of Carius may be used. 

In the lime method the material is mixed with pulverized lime, 
free from halogens, and is placed in a hard glass tube, closed at 
one end. Lime is placed in the open end of the tube, which is then 
heated in a combustion furnace. The organic compound is 
decomposed and the halogen unites with the calcium oxide to 
form calcium halide, from the acid solution of which the halogen 
may be precipitated by silver nitrate. If more than one halogen 


is present the separation may be made, after the heating is 
finished, by methods already outlined. 

In the Carius 1 method the material is heated in a closed tube 
in contact with fuming nitric acid and silver nitrate. The organic 
compound is oxidized and the free halogen thus produced is 
converted into the hydracid. The halogen hydracid at once 
reacts with silver nitrate and the silver halide is later weighed. 
The method is not well adapted to separation of the halogens, 
since a mixture of silver salts is obtained in the tube. 

Determination. — Carius tubes of hard glass may be obtained with 
one end already closed. The tube should be approximately 50 cm 
long and 2 cm in diameter. If such a tube is not at hand a good grade 
of combustion tubing may be used. One end is closed as follows: The 
tube is carefully heated at a point about 10 cm from one end by rotating 
in the flame of the blast lamp. When the glass has softened the tube is 
quickly drawn out, until half closed. It is allowed to cool and is then 
removed from the flame and cut at the narrow part. The nearly closed 
end is then fused together until a well-rounded end is produced. This 
must be annealed with great care or disastrous breaks will occur later. 

Having prepared a tube that is clean and dry, another small tube 
about 4 cm long is closed at one end to serve as a weighing tube. About 
0.2 gm of the organic material is weighed into the latter. Into the 
Carius tube is carefully placed about 1.5 gm of powdered silver nitrate 
and 2 cc of fuming nitric acid, free from halogens. The acid is intro- 
duced through a funnel with a long stem which reaches at least half 
way to the bottom of the tube, thus keeping the upper half dry. The 
weigl ing tube containing the substance to be analyzed is inserted into 
the end of the Carius tube, the latter being placed in a slanting position. 
Mixing of the contents of the weighing tube with the acid should not 
occur until after the Carius tube is sealed. The latter is now heated 
about 10 cm from the open end, the tube is drawn out while in the flame 
and the walls are sealed together. A more or less blunt point should be 
left here as shown in Fig. 36. 

Since a high pressure will be generated within the tube when heating 
begins it is necessary to place the tube inside an iron tube having caps 
screwed over the ends. The glass tubes frequently break on account of 
high pressure. The iron tube is now placed in a suitable furnace in which 
it may be gradually and uniformly heated. The temperature and 
time necessary for heating will vary with the nature of the substance 
under examination. Most organic compounds will be completely 

l Z. anal. Chem., 1, 240 (1861); 4, 451 (1864); 10, 103 (1871). 


decomposed by heating for three hours at 300°, while many aliphatic 
campounds will require a temperature no higher than 150°. 

After the decomposition is completed, as shown by the disappearance 
of carbon the furnace is allowed to cool, the iron pipe containing the tube 
is carefully removed, the cap unscrewed and the glass tube taken out. 
The latter is wrapped in a towel, to minimize the danger due to possible 
explosions, and the point of the tube, where it was last sealed off, is 
held in a flame until softened. The internal pressure causes the glass to 
blow out and the gas escapes, after which the tube may be handled 
without risk of injury. A scratch is made near the blown-out end, but 

Fig. 36. — Sealed end of Carius tube. 

on the wide part, and this end is broken off by touching the scratch with 
a hot glass rod. The contents of the tube are rinsed into a beaker, 
diluted with water and filtered; the precipitate is washed and weighed 
by the ordinary process, using either a paper filter or a Gooch crucible. 
From the weight of silver halide found the percent of halogen in the 
organic compound is calculated. 

If the fuming nitric acid contains halogens, blank determinations 
must be made and corrections applied. 

Carbonic Acid and Carbon Dioxide 

The following cases are to be considered: Carbon dioxide in 
gaseous mixtures, solutions of carbonic acid and salts of car- 
bonic acid. 

Carbon Dioxide in Gaseous Mixtures (air, chimney gases, 
etc.) — This determination is best made by gasometric methods 
which will be considered in a later section (pages 333 and 341). 

Carbonic Acid in Solution. — The most frequently occurring 
case is that of underground waters. Such waters, coming from 
regions of low temperature and high pressure, often contain con- 
siderable quantities of carbonic acid. When the water reaches 
the surface, diminished pressure and rise in temperature cause 
the release of more or less carbon dioxide, so that a determina- 
tion is always subject to some uncertainty regarding the rela- 
tion of the original concentration of carbonic acid to that in the 
water as the analyst receives it. Determinations are also re- 



quired of carbonic acid in carbonated drinks. In such cases 
provision must be made for transferring the solution from the 
pressure bottle to the apparatus in which the determination is 
to be made without loss of carbon dioxide. 

The procedure for the determination of carbonic acid in water 
is given on page 401. 

Carbon Dioxide in Carbonates. — Determinations of this class 
are by far the most common in general analytical practice. The 

Fig. 37. — Rohrbeck's appara- 
tus for determination of carbon 
dioxide by loss. 

Fig. 3S.-^Mohr's apparatus for 
determination of carbon dioxide by 

carbonate is decomposed by means of a stronger acid than car- 
bonic acid and the carbon dioxide determined in one of three 
ways: (1) by a determination of loss in weight, (2) by measuring 
the gas disengaged, or (3) by weighing this gas after absorbing 
by reagents in a suitable apparatus. 

Determination by Loss. — Many forms of apparatus may be 
obtained for the determination of carbonic acid by loss. Three 



of these are shown in Figs. 37, 38, and 39. Any such apparatus 
must include means for drying incoming air and outgoing gas. 
It must also be compact and not too heavy to be weighed on the 
[analytical balance. In using such apparatus the sample is 
| weighed and brushed into the lower generating vessel. Hydro- 
chloric or sulphuric acid is placed in the upper bulb and the 
1 bubble tubes are partly filled with concentrated sulphuric 
acid. The whole apparatus is accu r 
Irately weighed, after which the cock is 
'carefully opened so that acid drops 
|upon the carbonate, evolving carbon 
| dioxide at a moderate rate. This 
! carbon dioxide passes out through the 
sulphuric acid in the bubble tube, 
(being freed from moisture by so doing. 
[The apparatus is finally heated and 
air is drawn through to displace the 
remaining carbon dioxide. The loss 
in weight is taken to represent carbon 
dioxide. This is not accurately the 
case unless the air that is drawn 
Ithrough the apparatus is first dried. 
The determination by means of such 
apparatus is quickly made but is sub- 
ject to a rather large error on account 
of the large weight of the apparatus, 
because of the large surface and largely 
because of the difficulty encountered 
in the drying and purification of the 
outgoing gases unless unduly large 
quantities of sulphuric acid are used, as well as an absorbent for 
acid vapors. 

Determination by Absorption. — The direct determination by a 
somewhat more elaborate apparatus is to be preferred if accuracy 
is an object. In such a method the purification of the carbon 
dioxide is rendered complete by elaborating that part in which 
the purification is accomplished, providing better contact of 
the gases with drying agents and acid absorbents. Instead of 
weighing the entire apparatus before and after expulsion of 

Fig. 39. — Schrotter's ap- 
paratus for determination of 
carbon dioxide by loss. 



carbon dioxide from the carbonate the carbon dioxide is ab- 
sorbed in a weighed amount of potassium hydroxide, which is 
again weighed after the absorption. Many variations in the 
apparatus have been employed but the apparatus here described 
embodies the essential features of most of these. 

In Fig. 40, A is a generating flask into which the weighed 
sample of carbonate is placed. B is a dropping funnel having a 
capacity of 50 cc, and having the lower end drawn out to a point 
and turned upward. This part should extend to the bottom of 
the flask. At the top of the dropping funnel a drying tube C is 
connected by means of a rubber stopper and a bent glass tube. 

Fig. 40. — Assembled apparatus for determination of carbon dioxide by 


The drying tube is filled with soda lime for the absorption of 
carbon dioxide from the air that is later to be drawn through. 
Following the generating flask is a short condenser D and the 
U-tubes E, F and G. The first U-tube is omitted if sulphuric acic 
is to be used for decomposing the carbonate, or is filled with an al 
sorbent for hydrochloric acid vapors if this acid is used. Tl 
U-tubes F and G are filled with granular calcium chloride whicl 
absorbs moisture from the gas mixture. Following these is tl 
apparatus H in which potassium hydroxide is placed for the al 
sorption of carbon dioxide. This apparatus also carries a sm£ 


tube filled with calcium chloride to prevent the removal of mois- 
ture from the apparatus, which would occur if the dry entering 
gases were allowed to leave the apparatus saturated with mois- 
ture. To provide a means for drawing air through the whole 
apparatus the aspirator J is placed at the end of the series, while 
to prevent moisture from diffusing backward into the absorption 
apparatus the calcium chloride tube / is interposed. 

Choice of Acid. — The choice of acid to be used in decomposing 
the carbonate will depend upon the nature of the latter. Sul- 
phuric acid is to be preferred where it can be used, because it is 
non-volatile and thus needs no absorbent in the purifying 
apparatus. If, however, the carbonate is one of a metal which 
forms a sulphate of small solubility (e.g., calcium carbonate or 
barium carbonate) sulphuric acid soon coats the particles with 
insoluble sulphate which hinders the decomposition of the 
interior of the particles. Decomposition is slow and uncertain 
|and for this reason hydrochloric acid is used instead of sulphuric 
acid. A preliminary test should be made to ascertain whether 
sulphuric acid forms a complete solution of the carbonate to be 

Absorbent for Hydrochloric Acid. — If hydrochloric acid must 
be used a suitable absorbent is placed in the U-tube E, following 
the generating flask. Absorbents which serve best for this 
purpose are silver sulphate and anhydrous copper sulphate. For 
such a purpose the copper sulphate is prepared by first dropping 
red hot pieces of pumice stone into a concentrated solution of 
copper sulphate, removing the pumice stone, allowing to drain 
and then drying at 200°. A supply of the dry pieces is kept in a 
desiccator and fresh pieces placed in the U-tubes for each deter- 
mination. A more satisfactory absorbent is a saturated solution 
|of silver sulphate in concentrated sulphuric acid. This may be 
absorbed by pieces of pumice and used in the same manner as 
popper sulphate. When sulphuric acid is used in this manner it 
must not be allowed to come into contact with corks, cotton or 
any other organic matter. The evidence of such contact is 
blackening. The result is the formation of both carbon dioxide 
and sulphur dioxide. Both oxides are absorbed in the potas- 
sium hydroxide and give rise to errors in the determination. 
A U-tube with glass stoppers should be used. Silver nitrate 



cannot be used because its reaction with hydrochloric acid pro- 
duces nitric acid, which is nearly as volatile as hydrochloric 
acid, and also chlorine. 

Soda Lime. — The soda lime which is used for the removal of 
carbon dioxide from the entering air should be fresh and in the 
form of lumps. A powdered condition is evidence of having 
been air-slaked, in which case it is unfit for use since it is already 
saturated with carbon dioxide. 

Calcium Chloride. — The calcium chloride used for the absorp- 
tion of moisture should be the granular form which has been 
fused. Fusion is necessary in order to produce an anhydrous 
material. This fusion always produces a certain amount of 
calcium oxide which, if allowed to remain as such, will absorb a 
certain quantity of carbon dioxide as well as of water. It is 
best to treat the material directly in the bottle by passing dried 
carbon dioxide through for several hours, then displacing the 
carbon dioxide by drawing through dried air. In filling U-tubes 
only lumps should be used. The tube is filled to just below the 
side branches and then a loose plug of cotton or glass wool is 
placed on top in each side to prevent drawing out of any grains 
of powder that may subsequently be produced. For sealing 
the tubes a cork is pressed in until it begins to fit closely. It is 
then cut off even with the top and the smaller part is pressed 
into the tube about one-eighth inch farther. The shallow cup 
thus formed is poured full of melted paraffin or sealing wax. 
When this solidifies an air-tight seal should result, unless bubbles 
have formed in the sealing material. In the latter case a flame 
may be lightly touched to the surface of the solid paraffin or 
wax, which will cause the bubbles to break. 

Phosphorus Pentoxide. — Passing gases are dried to a greater 
extent by phosphorus pentoxide than by calcium chloride. The 
former absorbs moisture so rapidly as to make the charging of 
tubes difficult unless the humidity of the atmosphere is low. 
The oxide is usually obtained as a fine powder, formed by sub- 
limation. As this combines with water it forms a sticky mass of 
phosphoric acid, which soon clogs the tube unless some device 
is employed to prevent this. The best method for charging 
drying tubes with phosphorus pentoxide is to arrange a ribbon 
of glass wool, over which the oxide is sifted. The glass wool is 


fyhen quickly folded into a narrow strip which is placed in the 

) Choice of Drying Agent. — It might seem at first sight that 
(practically perfect drying of carbon dioxide before absorption 
would be necessary for accurate determinations and that phos- 
phorus pentoxide would therefore be the ideal drying agent. 
;Vn inspection of the conditions will show that this is not the case. 
Of course the increase in the weight of the absorption bulbs must 
(iccurately represent the weight of carbon dioxide absorbed. 
But in order that this may be the case it is only necessary that 
the unabsorbed air shall pass out of the bulbs with the same degree 
bf hydration that it possesses when entering. Therefore any 
[airly good dehydrating agent will serve provided that the same 
agent that is used preceding the bulb in the train is also used 
In the tube that is attached to the bulbs for drying outgoing 
gases. Either calcium chloride or phosphorus pentoxide, but 
pot both, may be used in the same train. Similar reasoning will 
apply to the use of concentrated sulphuric acid with other drying 

| Absorbent for Carbon Dioxide. — Potassium hydroxide is 
generally used for the absorption of carbon dioxide, a solution 
33 percent by weight being commonly employed. In practice 
t is found that absorption becomes so slow as to be uncertain 
before the point of complete saturation is reached. The prac- 
tical limit is reached when 0.10 gm of carbon dioxide has been 
absorbed by each cubic centimeter of potassium hydroxide 
solution. To determine the amount of gas that can be absorbed 
^y the solution in the apparatus the latter is first filled with 
kater to the height at which the liquid is to stand. This is 
femptied out and measured. The number of cubic centimeters 
times 0.1 gm is the weight of carbon dioxide which can be 
absorbed before the solution becomes inefficient. By adding 
together the weights of gas absorbed in successive experiments it 
is easy to determine when the bulbs need refilling. The bulbs 
in which the absorption is to take place furnish the greatest 
source of error to be encountered in this method. Inaccuracies 
are due to the large weight, the large surface and the possibility 
of moisture being carried out by the outgoing air. The effect 
of the comparatively large weight of the bulbs and their contents 


is to decrease the sensibility of the balance. The surface giv 
rise to a possible error because of the variable amount of moistur 
which is always dissolved in the surface of the glass. This 
error may be considerable if the two weights (before and afte 
the absorption) are observed under different atmospheri 
conditions of humidity. For this reason it is necessary tha 
the two readings of weights shall be made on the same day an 
as near to each other in point of time as possible. 

The danger of loss of moisture from the potassium hydroxic 
solution to the dry air which enters is magnified by the necessary 
limit which the already large weight of the bulbs places upon 
the tube which carries the calcium chloride for drying the 
outgoing air. For this reason many analysts prefer to separate 
this drying tube from the bulbs, using a small U-tube or even 
two such tubes, and weighing the apparatus in the two or three 
parts. Sometimes there is placed in the first half of the U-tube 
so used, or in the first U-tube if two are used, solid potassiu 
hydroxide to insure complete absorption of carbon dioxid 
While this procedure may make more certain the complete dryi 
of the air and thus prevent a loss of weight from this cause 
added uncertainty is introduced due to the accumulation of tr 
errors of four weighings. It is possible to insure complc 
detention of the moisture by passing the gas at a regular, specifie 
rate, not exceeding a maximum found by experience. 

Determination. — Procure the following parts for assembling: 

1 dropping funnel, 50 cc, with 1-hole rubber stopper, 

1 short, wide flask, 75 cc, such as is used for fat extractions, wi 
2-hole rubber stopper, • 

1 condenser with body not more than 6 inches long, 

3 U-tubes with corks to fit, 

1 U-tube with glass stoppers, 

1 straight drying tube with 1-hole rubber stopper, 

1 set "potash bulbs" of some approved form, 

1 aspirator bottle, tubulated near bottom, with 1-hole rubber stopper 
to fit, 

1 piece glass tubing, about 2 feet X % inch, for supporting apparatus 

2 clamps, 
2 pinch cocks, 

1 small screw clamp (Hoffman screw), 

2 retort stands, 



Glass and rubber tubing for connections. 

Fill and connect the apparatus in the manner previously described. 
Measure the capacity of the absorption bulbs by drawing in distilled 
water, then blowing out and measuring the water. This volume will be 
used as a basis for the calculation of absorbing power as already directed. 
When filling the absorption bulbs with potassium hydroxide solution 
the latter should not be warmer than the air of the room. The bulbs 
are detached from the apparatus and the solution is drawn in through 
a tube attached at a, suction being applied at b. The solution should 
about half fill the bulb c when air is bubbling through. The ground- 
glass joint between the drying tube b and the bulbs should be lightly 
coated with vaseline and the tube then twisted on until it fits closely 
enough that there will be no danger of loosening during the course of an 
experiment. Any surplus vaseline is removed from the outside of the 

Place the bulbs in position, close the cock of the dropping funnel 
ind open the pinch cock at e to allow water to flow from the aspirator. 
Bubbles of air will at first pass through the bulbs but this action will 
inally cease unless there is a leak in the apparatus, in which case it 
nust be found and closed . It is important that all glass tubes be brought 
entirely together inside the rubber connections since rubber is slightly 
permeable to gases. 

After the apparatus has been shown to be free from leaks the pinch 
pock at/ is closed, the cock of the separatory funnel slowly opened and, 
after equilibrium is established, the clamp k is so adjusted that when 
plamp / is opened air will pass through the bulbs at a rate not greater 
than 3 bubbles per second. Clamp k is not thereafter changed. This 
provides against too rapid flow of gas under any conditions. Clamp / 
s now closed, the bulbs are removed, the inlet and outlet tubes are 
plosed by short rubber tubes containing glass plugs and the bulbs are 
iviped clean and placed in the balance case. A short glass tube is 
nserted to bridge the gap made by removing the bulbs. The bulbs 
t;hould be allowed to stand for 15 minutes before weighing. In the 
|neantime about 1 gm of the carbonate is weighed and brushed into the 
generating flask and a small amount of water is added to moisten the 
(•ample. The stopcock of the funnel B and the clamp e are now opened 
;tnd 500 cc of air is drawn through the apparatus, measured by the 
;>utflowing water from the aspirator. This frees the apparatus from 
:arbon dioxide. After the absorption bulbs have stood for 15 minutes 
I he tubes carrying the plugs are removed and the bulbs are weighed. 
The plugs are then replaced and left so until the bulbs can be connected 
jn the apparatus. 50 cc of dilute sulphuric acid or hydrochloric acid is 
Placed in the dropping funnel, a test having previously been made to 


determine whether sulphuric acid will form a clear solution with tl 
carbonate. If such a solution is not produced, of course hydrochlor 
acid must be used and silver sulphate and pumice must be placed in tub 
E. Reconnect the apparatus and open all cocks except the stop-co 
in the dropping funnel, leaving the clamp k set for the proper rate of ga^ 
flow, as previously determined. Slowly open the cock of the droppin 
funnel, allowing acid to drop just fast enough to evolve carbon dioxide i 
the prescribed rate. The constant attention of the operator is necessai 
at this point, for by causing too rapid evolution of gas some moistu 
may escape absorption in the small tube of the absorption bulbs and th 
experiment be rendered worthless. 

The acid should be allowed to run in until about 1 cc is left abovt 
the stopcock, this acting as a seal during the subsequent boiling. Aftei 
the decomposition of the carbonate is complete the solution in the flask i$ 
slowly heated until it boils, always with due regard to the rate at whicl 
the gas is made to flow through the absorption bulbs. The boiling i* 
continued for one minute, when the flame is withdrawn, the cock of th< 
dropping funnel being opened at the same time to allow air to enter s( 
that no back suction occurs, due to the cooling effect. Air is now drawi 
through the apparatus until 1000 cc of water has flowed from the aspira 
tor. This amount of air should be sufficient to sweep all of the carbo 
dioxide into the absorption bulbs. 

The clamp / is now closed, and the absorption bulbs are removec 
plugged and placed in the balance case. After 15 minutes they ar 
weighed without the plugs, the increase in weight being the weigh 
of carbon dioxide. From this and the weight of sample the percent o 
carbonic anhydride (combined carbon dioxide) is calculated. 

For the duplicate or any subsequent determination the generatin;, 
flask and the dropping funnel are washed absolutely free from acid, s i 
that no decomposition of the next carbonate sample may occur bef or j 
the bulbs are in place. The first U-tube should also be emptied am 
recharged with absorbent, if such is to be used for the next determinatior j 

If a large number of determinations is to be made with the sam 
apparatus much time will be saved by providing two decompositio I 
flasks and two absorption bulbs. While one determination is bein 
made another sample may be weighed into the duplicate flask and th 
second absorption bulb may be weighed. The next determination ma I 
then be started while the first bulbs are standing in the balance case, pr< I 
liminary to the final weighing. It is also necessary to determine whe j 
the various absorbents have become so saturated as to be inefficier 
for further work. Soda lime in the tube C is good until the lumps ha\ J 
fallen into a powder. Silver sulphate in the pumice of tube E ma I 
become inefficient through absorption of hydrochloric acid or through th 1 



iccumulation of water in the tube. The solubility of silver sulphate in 
tfater is much less than in concentrated sulphuric acid. If the acid 
olution becomes diluted the silver salt crystallizes and will not there- 
after readily absorb hydrochloric acid. As the silver sulphate becomes 
aturated with hydrochloric acid it darkens, on account of the action of 
ight. When the darkening effect has proceeded as far as the middle 
)f the tube the material should be replaced. Calcium chloride must be 
eplaced when it becomes visibly moist for the first third of any absorbing 

A method has already been given for the determination of the amount 
>f carbon dioxide which can be absorbed by the solution in the bulbs. 



We have here to deal with a class of work that, while also 
gravimetric in most cases, is sufficiently different from what has 
already been considered to be treated as a separate division. In 
all of the preceding exercises the element or radical to be deter- 
mined was precipitated from a solution by chemical reactions 
produced by other substances which were added for the purpose. 
In the cases now to be considered the precipitation will be brought 
about by electrical action, the passage of a current through the 
solution causing the deposition of a metal upon a cathode in 
such a form that it can be weighed, or the accomplishment of 
some change which makes possible the determination of a sub- 
stance not a metal. The electrolysis of silver sulphate will serve 
as an example. When a solution of this salt is electrolyzed at 
platinum electrodes the metal is plated on the cathode and 
sulphuric acid is produced at the anode, thus: 


2Ag+S0 4 ^2Ag+S0 4 , 
2S0 4 +2H20->2H 2 S04+02. 

The silver can then be weighed and the sulphuric acid deter- 
mined volumetrically. 

While the electrolysis of a simple salt is frequently a tolerably 
simple and well understood process, the practical accomplish- 
ment of such a process for the purpose of a quantitative analysis 
is usually possible only when a certain set of conditions is main- j 
tained. The principal reasons for failure to attain accuracy an 
three: (1) Deposition may -not occur upon passage of a current. 
(2) The deposit may be contaminated by other products of 
electrolysis. (3) The deposit may not have the proper physical 
character, so that it will not adhere to the electrode but crumbles 
off during the electrolysis or during the process of washing. We 
have thus to consider the nature of salt to be used, solvents, tem- 



Iperature, electrolytic pressure (voltage), current density and 
nature and kind of electrode. 

Nature of Electrolyte. — Electrolytic methods are more fre- 
Jquently applied to the determination of metals than of non- 
metals, although methods have lately been perfected for the 
determination of the latter. If the metal alone is to be deter- 
mined it will usually be possible to obtain it in the form of what- 
ever salt gives the best results. Certain anions must be excluded 
iin specific cases, either because they yield substances that at- 
tack the anode or because the acids that are produced by their 
(electrical discharge cause the metal to deposit in an undesirable 
[physical form. As an example of corrosive action upon the anode 
ft is sufficient to mention here the formation of nascent chlorine 
lat a platinum anode when a chloride is electrolyzed. With regard 
to the effect of the acid that accumulates in the solution as 
(electrolysis proceeds it may be stated that there is little known, 
lat present, of the reasons for the effect of acids, bases and other 
pubstances that may be in the solution, upon the nature of the 
[deposit. Experiment shows, however, that such substances 
[often exert a very important influence upon the physical character 
|of a deposited metal and they are often added for this reason, 
[although they may be objectionable for other reasons. A solu- 
tion of copper sulphate, if electrolyzed without the addition of 
|another substance, usually gives a dark red or brown deposit of 
[finely divided copper which is liable to powder and be lost during 
[washing. If a small amount of sulphuric acid is first added 
[the deposit is improved, while nitric acid causes a still better 
[deposit of bright red, firm and adherent metal. For this reason 
Initric acid is usually added although it gives rise to more or 
|less danger of resolution when the cathode copper is being washed. 
jOn the other hand, a silver salt is best electrolyzed in the absence 
of nitric acid. If silver nitrate is electrolyzed from water solution 
l|with or without the addition of nitric acid (the latter is formed 
|!by the electrolysis) the plate of silver on the cathode is so decidedly 
[[Crystalline that it is very easily detached. The addition of 
Ipotassium cyanide in quantity sufficient to redissolve the pre- 
cipitate of silver cyanide first formed gives a solution from which 
(silver will deposit as a white firm plate. The solution in potas- 
ilsium cyanide has a comparatively high electrical resistance 


so that more energy is consumed in the accomplishment of its 
decomposition, nevertheless potassium cyanide is generally 

Other examples of similar effects will appear in the exercises. 
It is desirable to note that little is known of the cause of such 
effects, also to guard against a very common misconception 
regarding the purpose of adding other electrolytes to solutions 
that are to be electrolyzed. It is frequently stated that such 
substances are added in order to increase the conductivity of the 
solution. If such substances could increase the ionization of the 
salt that is to be electrolyzed, or in any manner diminish the fric- 
tional resistance to the passage of the ions, such an effect would be 
desirable. It is evident, however, that the addition of a foreign 
electrolyte can usually increase the conductivity only by itself 
acting as a carrier of current, in which case it has accomplished 
no desirable effect since the prime object is not to use a large 
current but to make the minimum current do the maximum 
work in discharging an ion already in solution. 

Solvent.— Very little work has been done in any solvents other 
than water. The use of organic solvents may, in some cases, 
prove advantageous in producing good deposits where other con- 
ditions fail to do so. 

Temperature. — Conductivity of solutions usually increases 
with rise in temperature. This is not due to increased ionization 
(ionization usually decreases with rise in temperature) but to re- 
duced viscosity and consequent reduction in frictional resistanc 
to ionic migration. If the complete electro-decomposition of ; 
substance requires considerable time it is not convenient to heat 
to any definite elevated temperature. In most cases, therefore, 
the temperature of the solution is not raised above that of the 
laboratory except at the beginning. 

Decomposition Voltage. — For every electrolyte in solution 
there is a definite minimum voltage, below which no decomposi- 
tion will take place. If but one electrolyte is present and the 
voltage lies below this minimum, a continuous current cannot flow. 
The minimum voltage necessary to produce a continuous flow 
of current is called the "decomposition voltage" for the sub- 
stance in question. If salts of more than one metal are present 
in solution, the deposit on the cathode will consist of any metals, 


the decomposition voltage of whose salts has been exceeded. 
If there is sufficient difference in the values of decomposition 
Ivoltage for the different salts, separation may be made. It is 
only necessary to adjust the voltage so that it shall exceed the 
decomposition voltage of the metal that is most easily discharged. 
After this metal has been completely removed from solution and 
weighed the voltage is raised until it exceeds the decomposition 
voltage of the metal next in order and so on. While this consti- 
tutes the general procedure for electrolytic separations it is 
necessary to make certain changes in the nature of the solution 
after each metal is removed in turn as will be understood from 
the discussion of the solutions to be employed. 

In order to understand the origin of the decomposition voltage 
it will be necessary to briefly consider the underlying principles 
of electrolysis. If a metal is placed in contact with a solution of 
one of its salts it will be found that a difference in potential exists 
between the metal and solution. This difference may be either 
positive or negative, i.e., the metal may be at a higher or lower 
potential than that of the solution. The difference may be 
zero in certain cases but such cases are special. The conception 
of Helmholtz 1 regarding the cause of this potential difference 
may be thus stated: Whenever two dissimilar substances are 
in contact a potential difference is established because of the 
passage of one into the other. In the case of metals and their 
salt solutions, one of two things may happen: either some metal 
atoms pass into solution and become charged ions or some ions 
are discharged by the mass of metal and themselves become 
elementary In the first case the solution assumes a higher 
potential than the metal because positive charges have been 
transferred from metal to solution. In the second case the 
solution is at a potential lower than that of the metal because 
positive electricity has passed from solution to metal. The 
direction of the change is determined by the relative magnitude 
of two opposing forces. The metal shows a tendency to pass 
into solution in the ionic condition in obedience to a force called 
by Nernst 2 " electrolytic solution tension." This force varies 
with different elements, but is constant for a given element and 

iWied. Ann., 7,337 (1879). 

2 Z. physik. Chem., 4, 129 (1889). 



may be relatively large or small. When positive ions have bee 
thrown into the solution the potential difference thus establishe 
gives rise to an attraction of an electrostatic nature betwee 
the positively charged solution and the negatively charged meta 
This may be represented diagrammatically as in Fig. 41. 

Double Layer. — Helmholtz considered that a "double layer 
was thus formed, composed of positive and negative charges 
and that the components of this layer were very close together 
The attraction existing between the components of the double 
layer increases as more ions are formed and finally reaches equilib 
rium with the solution pressure. If ions of the metal in question 

+ - -+ 

4 - -4 

4 - - + 

4 - - + 

4- - + 

4 - - + 

4- - 4 

4- - + 

4 - - 4 

4 - - 4 

4- - + 

4- -4 

4- - + 

4- - + 

4 - - 4 

4 - - 4 

4I- 1 ' ' ' M4 

Fig. 41. — Diagram illustrating the "double layer." 

were already in the solution, then osmotic pressure would oppos 
the entrance of more ions into the solution and would thus act ir 
conjunction with the electrostatic attraction, so that equilibria 
would be reached with a smaller potential difference. Evidently 
then, the potential difference between the electrode and elec 
trolyte, in the case of a given metal and its salt solution, wi 
be numerically greatest when the initial ion concentration 
least, and least when the latter is greatest. If the solution 
pressure is small and the ion concentration large, equilibrium 
may be reached only by the actual deposition of ions upon th 
metal. In this case the potential difference will be positive. 

If the different elements are compared, with regard to the 
potential difference established between them and their ion solu 



Jtions, the ion concentration being the same in all cases, a series 

of different values is obtained. It is to be noted that if an 

element takes a negative charge upon becoming ionized the 

potential difference is reversed in sign. In the following table 1 

several of the elements are given with values for the potential 

difference. These differences are measured by a method that 

need not be here discussed, involving the use of an arbitrary 

standard, so that only the relative values are important. These 

values are for solutions which contain 1 gm-ion per liter. 

Potential difference 


Potential difference 



— 3.20 



< + 0.293 

+ 0.329 







— 2.54 

+ 0.750 

— 1.276 

+ 0.771 









< + 1.079 

+ 1.96 



+ 1.417 



+ 0.993 



+ 0.520 



+ 1.119 



In the electrolysis of a salt solution both positive and negative 
ions are discharged at the electrodes, both electrodes become 
coated with the products of decomposition (polarized), and both 
then become essentially electrodes of the respective elements, 
no matter what the original substance might have been. Hence 
such a system as that just discussed may be considered to exist 
at each electrode. Since, after electrolytic decomposition has 
begun, electrolysis consists of electrical discharge of ions it is 
evident that it must act in opposition to solution pressure and in 
conjunction with osmotic pressure. That is, in order to produce 
continuous electrolysis a voltage must be applied to the electrodes 
that is at least as great as the algebraic difference between the 
single potential differences normally established at the cathode 
and anode. This difference constitutes the theoretical "deeom- 

1 Wilsmore: Z. physik. Chem., 35, 291 (1900). 


position voltage" of a given salt solution, and no appreciabl 
electrolysis can take place as a result of the application of 
lower voltage. The attempt to calculate the decomposition 
voltage of a solution from the single potential differences, ex 
perimentally determined, does not always give results that agre 
with those found by experiment. This is because the ion con 
centration is not necessarily equal to 1 gm-ion per liter and i 
changes as electrolysis proceeds. Also where gases are evolvec 
at the anode (a condition generally noticed) the phenomenoi 
of "overvoltage" exerts a very important influence upon tin 
practical decomposition voltage. 

Importance of this Principle. — Agreement is not at all satis 
factory in the case of salts of oxyacids. This is partly becaus 
oxysalts are not normal with respect to the oxygen or hydroxy 
ion, also because oxygen shows overvoltage to a marked degree 
giving a much larger decomposition voltage than would be 
calculated. The matter that is here important is not the possibiliti 
of calculating decomposition voltages from known single potentia 
differences but the recognition of the reasons for the fact that 
minimum decomposition voltage must exist for every compound 
under definite conditions, and that if this value can be determined 
a method for electrolytic separations is available. We shall under 
stand that " decomposition voltage" refers to the fall in potentia 
measured across the electrodes, below which electrolysis cannol 
take place continuously. So far as we now know there is no 
upper limit to the voltage that, should be used, excepting that set 
by the current density that should be employed, unless 
separations are to be made. 

Because the variable factors which influence the practical de 
composition voltage (such as overvoltage of the anion, varying 
temperature and added electrolytes) and the consequent diffi 
culty that is experienced in its calculation, Sand 1 proposed t( 
measure merely the cathode potential difference and to mak( 
metal separations by properly grading this potential. This is 
no doubt, the ideal method. For making such measurements 
however, more elaborate apparatus is required and much cart 
must be exercised. For most purposes such refinement is entireb 
unnecessary because the practical decomposition voltage is 

1 J. Chem. Soc, 91, 374 (1906). 


usually fairly accurately known as the result of experiments with 
a given salt, electrolyzed under specified conditions. 

Current Density. — The relation between the amount of current 
flowing through a solution of an electrolyte and the amount of 
substance decomposed is stated in the law of Faraday: 1 (a) 
For any given electrolyte the amount of decomposition is directly 
proportional to the amount of current; (b) the amounts of different 
substances decomposed by the same current are proportional to the 
combining weights of the substances. According to the first part 
of this law the rate of decomposition and electro-deposition in 
any experiment will depend upon the current strength, except- 
ing the part played by other electrolytes that may be present. 
Being thus able to limit and control the rate of deposition the 
question arises as to whether there is any suitable current strength, 
above or below which good results will not be attained. If it 
were not possible for any current to pass through a solution ex- 
cept that carried by the ions of the salt which we desire to de- 
compose there would probably be no definite limit to the prac- 
tical current strength to be employed. The ions encounter a 
large frictional resistance in their migration toward the elec- 
trodes. In order to overcome this resistance the pressure (volt- 
age) is raised, often considerably above the decomposition voltage 
of the metal salt being analyzed, in order to hasten the action. 
If no other cation is present there is no upper limit to the prac- 
tical voltage. There is, however, another positive ion that is 
present in all aqueous solutions and particularly when acids are 
present. The ion referred to is that of hydrogen. If the pres- 
sure is raised above its discharge potential it can discharge at 
the cathode and will do so unless the current strength, concen- 
tration of hydrogen ion and concentration of metal ion are so 
related that the current can easily be carried from solution to 
cathode by the metal ion without the necessity for discharge 
of the hydrogen ion. This relation will evidently be such that 
I there is a relatively small current, large metal ion concentration 
and small hydrogen ion concentration. The objection to the 
! deposition of hydrogen on the cathode is based upon the fact 
I that minute bubbles of gas prevent the proper coherence of the 
deposited metal. Apparently, then, the upper limit of current 

1 Pogg. Ann., 33, 301 and 481 (1834). 


will be fixed by the point at which noticeable evolution of hy- 
drogen (or other gas) occurs. This limit cannot well be calcu- 
lated but is determined by experiment. It should be noted that 
the value to be measured is not that of total current flowing 
across the solution but is that of the current flowing into unit 
area of the electrode where the desired deposition is taking 
place, usually the cathode. This gives rise to the term "cur 
rent density," abbreviated to CD. In stating the conditions 
to be observed in electro-depositions the current density may b< 
more specifically defined as the current in amperes flowing int< 
each 100 sq cm of cathode surface. This is denoted by CD i o 

total am peres 

100 ~ square decimeters cathode surface 

The proper current density to be employed in a given case is called 
the ''normal density" for that experiment. This is indicatec 
by ND ioo. The normal density is fixed by the conditions alread; 
discussed. There is considerable variation, however, basec 
also upon the form of the electrodes. This will be taken u] 
in the next paragraph. 

Nature of Electrodes. — Electrodes must possess certain proper- 
ties in order to be suitable for use in quantitative analysis. Th< 
electrode material must be insoluble in the solution of electn 
lyte, with or without current action. In the process of metal 
plating as used in the arts it is customary to make the anode of 
the metal being plated, so that the rate of deposition at the cath- 
ode is equal to the rate of solution at the anode, the mean con- 
centration of the metal in the solution remaining constant. 
This is obviously out of the question in quantitative analysis, 
where the total metal in solution, and no more, is to be deposited. 

The material most used for electrodes is platinum. The con- 
tinued advance in the cost of platinum has led to a search for 
less expensive materials. However, the combination of high 
electrical conductivity and low solubility is a rare one. Other 
metals could be used for cathodes because the current action 
prevents their resolution, but when the deposited metal is to be 
removed after the process is finished the solvent used will gener- 
ally dissolve some of the electrode also. So long as the cost 
of platinum is sufficiently low to make it possible to provide an 


adequate supply of platinum electrodes it is doubtful whether 
any other material will supplant it to any great extent. 

Gooch and Burdick 1 have perfected a method for making 
electrodes, by which a very small amount of platinum is spread 
over a relatively large surface of glass. A mixture of glycerine 
and chlorplatinic acid is spread over the glass surface, which 
is then heated. The glycerine is evaporated and the chlor- 
platinic acid is decomposed, elementary platinum fusing into the 
glass surface. 

Mercury is used as a cathode for a certain class of work and 
this will be discussed in a later paragraph. 

The electrodes must also be chemically unaltered by the 
passage of a current or else altered in a definite manner. The 
first condition is more often realized but there are cases where 
one electrode, is altered in a definite manner, as a silver anode, 
used for the determination of chlorine, bromine or iodine anion 
becomes coated with chloride, bromide or iodide of silver. 

The electrode that is to be weighed and is to receive the deposit 
(usually the cathode) should present the maximum surface 
for the minimum weight of electrode material. Since a practical 
limit is placed upon the current density the duration of the process 
of deposition will be inversely proportional to the total electrode 
surface exposed. The weight of the electrode must not be too 
large for accurate work and these considerations naturally lead 
us to consider the form of material where the ratio of surface 
to weight will be as large as possible. In general, any piece 
of platinum, small enough to be weighed, may be used as a 
cathode for metal deposition. Any chemist who has a dish 
or crucible may make at least an occasional analysis by the use 
of such an article as cathode. The dish is simpler because the 
solution may be placed directly in it and a coil of platinum wire 
used as an anode. The ratio of surface to weight is not large in 
this case, especially as only one surface is effective. Moreover 
if there is any sediment in the solution this may be partly caught 
by the depositing metal and weighed along with the latter. The 
cathode dish designed by Classen is quite thin and presents a 
larger surface. A dish of this kind weighing about 40 gm has 
a capacity of 250 cc and presents an inner surface of about 150 

1 Z. anorg. Chem., 78, 213 (1912). 



sq cm to the solution. A crucible may be used as a cathode b 
connecting as in Fig. 42. A rubber stopper is used to hel 
support the crucible, a metal rod passing through and connectin 
with the negative of the current source. A small platinum wir 
serves to complete the connection of crucib 
with rod. This form of cathode also possess 
a small relative surface. 

Other forms of cathodes are open cylinde: 
and cones of foil, and gauze cylinders an 
plates, made from gauze of small mesh an 
fine wire. These forms are shown in t 
illustrations (Figs. 43 and 44). Of all of these 
forms the gauze electrode is most efficient, n 
only because the relative surface is great! 
increased by constructing of fine wire but al 
because practically all parts of the surface are 
equally effective. The latter condition does 
not obtain for foil electrodes of any form, the 
surface farthest from the anode being in a 
relatively weak electrolytic field. Gauze elec- 
trodes also permit better mixing of the solu- 
tion and very much higher current densities 
Special forms of electrodes for rapid rotation will 




42. — Crucible 

may be used. 

be discussed later. 

Other Apparatus. — The necessary apparatus for electro- 
analysis will include, besides the electrodes, a generator of direct 
current, variable resistance, voltmeter and ammeter. The source 
of current may be a dynamo, any of the forms of primary cells, 
secondary or storage cells or thermoelements. The thermoele- 
ment is not a practical source of current, being both inefficient 
and unreliable. The direct current from a dynamo may be 
used and is better than primary cells, the latter being trouble- 
some in the matter of maintenance. The chief objection to 
the dynamo current lies in the fluctuations usually resulting 
from a variable load on the line from the generator. The best 
and most satisfactory current producer for this class of work is 
the secondary or storage element. Any of the various forms of 
accumulators will prove satisfactory, the lead-lead peroxide cell 
being the best known. The great merit of the secondary cell 



I is its constancy and reliability. The E. M. F. of the lead cell 
| is about 2 volts and the necessary voltage for the work may be 
obtained by connecting several cells in series. 

Any rheostat will do for this work, provided that the range 
in resistance is properly related to the other factors entering 
into the determination of current strength. In the absence 

Fig. 43. — Platinum foil cathodes. 

Fig. 44. — Platinum 
gauze cathode. 

of such a rheostat carbon lamps may be used for the current 
control, if the line voltage is high. The resistance of a 16 c.p. 
carbon lamp is about 220 ohms and by arranging several 
in parallel a fairly satisfactory regulation of current may be 

Voltmeters and ammeters should have the scales graduated 
with a range as limited as is consistent with the current conditions 



to be employed, so that each subdivision may represent a smal 
fraction of a unit. A satisfactory plan is to have double scale 
instruments, the range of one scale being ten times that of the 

The necessary connections for the apparatus used for electro 
analysis are shown in figure 45. B represents the source o 
current. In series with this are connected A, the ammeter fo 
measuring the current strength, R, a rheostat for varying the 
resistance in the circuit and thus providing current control, 
c and a, cathode and anode. The voltmeter V, for measuring 
the pressure, is given a shunt connection across the electrodes. 



Fig. 45.— Diagram of connections for electro-deposition of metals. 

The actual current flowing through V is very small because of 
the high resistance possessed by the winding of the voltmeter. 
On this account the indications of the ammeter are practically 
correct for the current flowing through the electrolyte between 
the electrodes, although not absolutely so. 

The following is a description of the apparatus now in use in 
the Purdue laboratory. 

Purdue University Laboratory for Electro-Analysis. — Current 
is furnished by storage cells of the lead type, each having a 
capacity of 48 amp-hours and a maximum charge and discharge 
rate of 6 amp. They are placed in a closed battery case which is 
outside the room for electro-analysis. The cells are provided 
with sand trays and insulators, and also with glass covers which 



almost entirely prevent the annoyance due to acid spray during 
charging, and the case ventilator makes charging absolutely 
inoffensive. The interior of the case is protected against the 
attack of acid by asphaltum paint. 

Wires from the cells run into the special laboratory where 
they are connected with the distributing switchboard. This is 

Fig. 46. — Distributing switchboard of the Purdue University Laboratory 
for electro-analysis. 

a 28 X 72-inch board of black oiled slate, providing switches 
and plug receptacles for the control of all current which is 
used for any purpose within this room. The cells are con- 
nected in series groups of three each, the outside terminals of 
such group having double receptacles. This arrangement en- 


ables the operator to connect his cells with any number o 
groups in multiple, thus giving greater latitude in the selection 
of voltage and current strength than is possible with the usua 
series connections. 

The 110-volt charging current enters the board through a 
switch which can be connected by plug connectors with any 
cell or combination of cells, and a slide-wire rheostat on the 
back of the board makes it possible to charge any number of 
cells at a time. The cells are protected during charging by an 
under-load circuit-breaker, and also during both charge and 
discharge by a fuse panel which is placed on the back of the 
board. An ammeter with a range of 20 amperes and a double 
scale voltmeter, with ranges of 150 volts and 15 volts, are pro- 
vided for proper control of the charging process. The 15-volt 
scale is used for testing the voltage of the cells when they are not 
in use. 

A switch on the distributing board controls the 110-volt 
alternating current which is used for the lights and motors 
Finally, on this board are the terminals for all of the desks, so 
that any operator may connect with his desk any of the cells 
not then in use, and in almost any combination. It will be seen 
that the connections on this board make the different cells and 
desks practically independent of each other; for instance, a part 
of the cells may be charging while the remainder may be dis- 
tributed to the various desks as wanted. 

At each working desk is a 24 X 36-inch slate panel which 
carries all of the apparatus that will be needed by the analyst, 
making each board an independent working unit. The volt 
meter and ammeter on each board are double-scale instruments 
with ranges of 2 volts and 20 volts and amperes, respectively 
The multiplier for the voltmeter is controlled by a small knife 
switch, and the shunts for the ammeter are joined to plug 

The current to each desk panel is controlled by a slide-wire 
taper rheostat. These rheostats sre wound to give a total re- 
sistance of 130 ohms, in 254 steps. The carrying capacity of 
the first step is 1.3 amperes and that of the last step 25 amperes, 
on continuous work. 

For working with rotating electrodes, 1/30 h.p. alternating- 



current, series motors of the commutator type are mounted on 
the board and are controlled by a switch and 5-step rheostat. 
Induction motors are not used, because it is desirable to make 
variation in speed possible. These motors have a maximum 
speed of 2200 r.p.m. on 110 volts and are provided with three 
pulleys of different sizes. 

Fig. 47. — A single desk panel and rotator, Purdue University Laboratory. 

In many laboratories it is the practice to mount the motor 
directly on the electrolyzing stand, thus avoiding all belting and 
making possible direct connection with the rotating electrode. 
On the other hand, the application of a small belt is a simple 



operation and not only very much decreases the vibration o 
the stand, and consequently the danger of dust particles fall 
ing into the bath, but also removes the motor from the regio 
of the bath, which in many cases contains acids and which i 
frequently hot. Corrosion of the motor parts is in this way 
largely prevented. 

The stand for holding the electrodes is of iron, is quite heavy 
to prevent vibration, and is fitted with rubber feet. Every 
portion of the base and vertical rod is heavily enameled and th 

Fig. 48. — Group of five^desk panels, Purdue University Laboratory. 

electrode supports are clamped to the rod by means of heavy 
thumb screws, but in such a manner that the screw does not come 
into contact with the rod, so that the enamel is not injured by 
the grip. Each clamp is insulated from the rod by a fiber 
bushing and each carries a binding post. There is no glass 
about the stand, perfect insulation of the electrode clamp 
being secured by the fiber bushings. The supporting ring for 
a dish cathode has three brass screw contacts which are ad- 
justable for dishes of different sizes. Finally, for stationary 
electrodes, simpler clamps are provided to take the place of th( 


rotator. The rotator, which carries three pulleys of different 
sizes, is a vertical shaft, the'lower end of which carries a universal 
chuck, electrical contact being insured by a brass brush. 











Amount or 
8ample Take* 

Fraction Used 

Quantity op 





Total Amperes 

Cathode 3urfaoe 

0. D. 



Total Time op 

Description of 


and Speed, •if 


VtasnT oB 


Character of 




obtained ( 
purpose is 

ig. 49.— Blank f 

. — Systematic 

luring the ex 

shown in figi 

or reporting rest 

3 records shou 
periment. A 
ire 49. 

dts of electro-ana 

Id be kept of 


all of the data 
blank for this 



If a copper salt is to be used a nitrate or sulphate is best suited 
Other salts of volatile acids may be converted into the sulphat 
by evaporating with sulphuric acid, stopping the evaporatior 
before any decomposition into copper oxide occurs. Coppe 
deposits in a coherent form from solutions containing sulphuric 
acid, nitric acid, oxalic acid and ammonium oxalate, potassiun 
cyanide, phosphoric acid, formic acid or ammonium hydroxide 
Of all of these, nitric acid produces the best results and probably 
sulphuric acid is next best. Chlorides should not be present 
If metallic copper is to be analyzed it may be dissolved in nitric 
acid and the undesired excess of acid removed by evaporation. 

Determination of Copper in Soluble Salts. — Use enough sample to 
yield 0.25 to 0.50 gm of copper. Dissolve in such a manner that 200 
cc of solution will contain about 2 cc concentrated nitric acid. If the 
sample contains chlorides it must be treated with the least possible 
excess of sulphuric acid and heated until fumes of sulphur trioxide 
appear. It is sometimes desirable to make a larger quantity of solution, 
as 250 cc, and to use an aliquot part for each determination. In this 
case the acid may be added to the solution as used. The electrolysis 
may be begun and finished at the temperature of the room, but it will 
be hastened by warming the solution to about 70° at the beginning. 
Connect the weighed electrodes and add enough water to cover the 
cathode, which must extend entirely to the bottom of the beaker, place 
split cover glasses on the beaker and electrolyze with a pressure not 
below 1.7 volts and not above 2.0 volts unless other metals are known to 
be absent. In the latter case there is no upper limit to the voltage that 
may be used, except that fixed by the current density desired to give 
good deposits. The current density that may be used will depend upon 
the kind of electrodes. For foil cones or cylinders or for dishes, NDi 00 = 
about 0.1 amp. For gauze electrodes iVDioo may sometimes be as high 
as 5 amp. In any case the analyst must use his judgment, watching the 
deposited metal to discover its character. The copper should appear as 
a bright red metal with no spots of brown and no tendency to crumbl 
off the cathode. When the disappearance of color indicates that th 
metal is deposited remove a few drops to a white test plate by means of 
pipette and test for traces of copper by adding a drop of concentrate 
ammonium hydroxide or of potassium ferrocyanide solution. The for 
mer is preferable for the first test because if copper is found the solutio 
can be neutralized with nitric acid and returned to the beaker. If n 
copper is indicated make another test, using four or five drops take 
from the bottom* of the beaker, and applying the ferrocyanide test 


When all of the metal is found to be deposited arrange a small siphon 
tube in such a manner that the solution may be drawn from the bottom 
of the beaker without interrupting the current. As the solution is 
removed wash down the exposed portion of the cathode removing every 
trace of acid before the E. M. F. is removed from the system. It is best 
to add water fast enough to keep the bottoms of both electrodes covered 

I until the acid has become so diluted that no further action is to be feared. 

II Siphon out the remaining liquid and perform the next operations as 
quickly as possible. Instead of siphoning, another plan is to have the 

v beaker supported on blocks, removing these and lowering the beaker 
t. jgradually and washing the cathode as it becomes exposed. 

Lower the beaker and remove the cathode, taking care to avoid 
touching cathode to anode and thus making a short circuit, and quickly 
wash with much water. Set aside until the duplicate cathode has been 
treated in the same way, then wash both cathodes with redistilled alcohol 
and dry at 100°. The alcohol washing may be omitted, but drying is 
hastened by this means. Weigh and calculate the percent of copper in 
the sample. 

Remove the copper from the cathode by dipping into warm dilute 
nitric acid. 

Determination of Copper in Brass and Similar Alloys. 1 — Dissolve 
0.5 gm of the drillings in a mixture of 5 cc each of concentrated nitric 
acid and water, in a covered casserole, heating finally to expel all brown 
oxides of nitrogen. Dilute to 75 cc and filter through a paper of close 
texture. Wash the residue of metastannic acid with hot water, pre- 
serving the washings with the filtrate. (The precipitate may be ignited 
win a porcelain crucible to stannic oxide for an approximate determination 
J of tin.) 

To the filtrate and washings in a casserole add 5 cc of concentrated 
; sulphuric acid. Evaporate under a hood until copious fumes of sul- 
J } phuric acid appear. The nitric acid is thus driven off. Cool, add 35 cc 
of water and boil for 1 minute in order to dissolve all soluble sulphates. 
I Filter on a Gooch crucible and wash with dilute sulphuric acid until a 
B few drops of the washings show no - blue color with a slight excess of 
! I ammonium hydroxide. (In case the Gooch crucible has been ignited 
;, and weighed before filtering the solution, the lead sulphate may be used 
for a determination of lead. It is washed with 50 percent alcohol to 
.[; remove sulphuric acid and is then dried and heated, cautiously at first, 
I to dull redness for 15 minutes. After cooling it is weighed as lead sul- 
phate, from which the percent of lead is calculated.) 

1 The complete analysis of brass is described on page 505. The brief 
i outline here given is with reference to the electrolytic determination of 


The solution and washings (not including alcohol washings from tl 
lead sulphate) now contain copper, zinc and sulphuric acid. In pn 
ence of so much sulphuric acid it is not desirable to add more nitric acic 
even though the latter would improve the character of the deposit 
copper to be obtained. Instead, add 1 gm of ammonium nitrate, whic 
results in the substitution of nitric for sulphuric acid. Rinse the soh 
tion into the beaker in which electrolysis is to take place, arrange tl 
electrodes and add water until the cathode is covered, mixing wel 
Conduct the electrolysis and subsequent treatment as directed on pag 
156, keeping the voltage below 2.5 to avoid the deposition of zinc. 

Determination of Copper in Ores. — The ore will always contain moi 
or less insoluble gangue. It may or may not contain metals whos 
compounds are soluble in aqua regia and whose decomposition voltages ; 
are near to that of copper (bismuth, antimony, arsenic, mercury or 
silver). In case such metals are absent, or are present in quantities so 
small that they may be disregarded, proceed as follows : 

Weigh 0.5 gm of the powdered ore and place in a casserole. Addi 
10 cc of concentrated hydrochloric acid and 5 cc of concentrated nitric 
acid, cover and boil slowly to aid in dissolving. Digest on the steam . 
bath until action seems to be complete then add 7 cc of concentrated; 
sulphuric acid and boil under a hood until volatile acids are expelled; 
and fumes of sulphuric acid appear. Cool, add 25 cc of water, boil n 
minute and filter to remove lead sulphate and gangue, receiving the; 
filtrate in the beaker in which the solution is to be electrolyzed. Wash| 
well with hot water, preserving the washings with the filtrate. Add 1 
gm of dry ammonium nitrate, connect the electrodes in place, add water 
until the cathode is covered and mix well. Electrolyze as directed on 
page 156. 

If any of the metals named above are present in appreciable quantities 
the procedure is the same as that just described for other ores, until 
after the filtrate from lead sulphate and gangue is obtained. This 
filtrate and the washings are received in a casserole. Dilute, if neces- 
sary, to 75 cc. Cut a strip of sheet aluminium about 2.5 cm wide and 
14 cm long, bend into a triangle and place in the casserole. Cover the 
solution and boil gently until all of the copper is precipitated, leaving 
the solution colorless or green from ferrous sulphate. (If this condition 
cannot be obtained it is because nitric acid has not been expelled com- 
pletely when evaporating with sulphuric acid.) 

When all copper is precipitated wash down the sides of the casserole 
with a stream of hydrogen sulphide solution and pour the solution and 
copper into a filter paper. Filter rapidly, to minimize oxidation and 
resolution of copper, and wash the aluminium and copper with hydrogen 
sulphide solution. The latter will cause the precipitation of traces of 


rcopper that may have redissolved but this precipitation should take 
place before the solution has passed through the filter. If the filtrate 
appears brown this indicates a loss of copper and the solution must be 
i refiltered and washed until a clear filtrate is obtained. Place under 
phe filter tht beaker which is to be used for the electrolysis, then pour 
JDver the aluminium in the casserole a mixture of 4 cc of concen trated 
aitric acid with the same volume of water. Heat to dissolve all adher- 
ing copper then pour the acid slowly over the copper in the filter. When 
all copper is dissolved wash the paper very thoroughly with hot water. 
Boil the filtrate and washings to expel oxides of nitrogen, then electro- 
llyze without further addition of acid. 


The deposition of silver from solutions containing free acids 
Imay be accomplished, but the deposit is usually either spongy or 
{crystalline so that it cannot be washed without loss. When 
nitric acid is present there is also a deposit of silver peroxide at 
the anode if high voltage is applied. The addition of potassium 
cyanide, although materially lowering the concentration of silver 
cations, is desirable because it entirely prevents the formation 
of silver peroxide and also yields a coherent deposit of silver at 
the cathode. The deposit is without lustre and should be white. 

The deposition of silver peroxide upon the anode, when high 
voltage is applied to solutions containing nitric acid, is probably 
due to the existence of the anion of a silver oxyacid, H 2 Ag0 2 — - 

2H+Ag0 2 . This is a theoretical derivative of the dihydroxide 
of silver, Ag(OH) 2 . The decomposition potential of this anion 
'is higher than that of the univalent silver cation and its concentra- 
tion is always small, yet it will discharge to some extent at the 
same time that silver is depositing upon the cathode, if high vol- 
tage and current density are used: Ag0 2 — »AgO+0. Such a 
deposit having formed upon the anode it will finally redissolve 
as the silver becomes more dilute in the solution because the 


Ag^Ag0 2 

is disturbed by the removal of silver cations, Ag, which are dis- 
charging at a much greater rate on account of their lower decom- 


position potential. The discharge and deposition of silve 
peroxide is entirely prevented by potassium cyanide. Whe 
this is added to a solution of silver nitrate a pricipitate of silvc 
cyanide is first formed: 

AgN0 3 +KCN->AgCN+KN0 3 . 

An excess of potassium cyanide redissolves this precipitat 
forming a salt of a complex anion: 

AgCN+KCN->KAg(CN) 2 . 

In this way the equilibrium, 

Ag<=»Ag0 2 

is also disturbed by the conversion of silver into the new anio 

Ag(CN)2, which, presumably, has a high decomposition potentia 
The double cyanide of potassium and silver cannot all be in th 
form represented by the formula KAg(CN) 2 as in this case no 
silver could be discharged at the cathode, but only hydrogen. 
There must, therefore, be equilibrium between the salt having the 
composition represented above and the single cyanides: 

KAg(CN) 2 ^KCN+AgCN, 

or the ionic equilibrium 

+ - 

Ag(CN) 2 ^Ag+2CN. 

It is to be supposed that the relatively high decompositio 

potential of the anion Ag(CN) 2 prevents its discharge und< 
ordinary conditions. 

Determination. — Use sufficient sample to give about 0.3 gm of silve 
Dissolve this in a small amount of water and add just double the quai 
tity of a solution containing 3 to 5 gm of potassium cyanide in 25 cc < 
water that is necessary to redissolve the precipitated silver cyanid 
Fasten the weighed electrodes in place and dilute with water until th 
cathode is covered. The E.M.F. required will be about 2 volts and NDi 
= 0.04 to 0.10 amp for electrodes not of gauze or as high as 2 amp 
gauze electrodes are used. If the current density is too large the pota 
sium cyanide will be decomposed around the anode, giving a dark soli 



ition of organic matter that will eventually reach the cathode and darken 
the silver. Such darkening will occur also if the potassium cyanide is 
impure. A cyanide of high grade is required for this purpose. The 
deposition will require 30 minutes to 5 hours, depending upon the 
current density used. 

When the deposition is thought to be complete a few drops of the 
solution may be tested for silver by adding dilute nitric acid in sufficient 
quantity to decompose the potassium cyanide. Hydrocyanic acid is 
removed by boiling and then ammonia is added to make basic, this 
; being followed by ammonium sulphide. When the deposition has been 
I shown to be complete the solution is removed and the cathode washed, 
dried and weighed as in the preceding exercise. Remove the silver from 
the electrode by dipping into dilute nitric acid. 

Caution. — Avoid inhaling the vapors that arise during the progress 
of electrolysis. Do not use pipettes and do not fill the siphon by suction. 
Determination of Silver and Copper in an Alloy. — (Other metals, with 
i the exception of gold, are assumed to be absent.) 

Clean the alloy by polishing, followed by wiping with filter paper. 
Cut into pieces or drill so that about 0.5 gm may be used for each 
analysis. Dissolve in a covered casserole in a mixture of 5 cc each of 
concentrated nitric acid and water. Rinse down the cover glass and 
I evaporate the solution nearly to dryness on the steam bath. Do not 
I heat the dry mass as this would cause the decomposition of some of the 
nitrates. Redissolve the salts in about 10 cc of water. 

Silver. — Place 3 gm of potassium cyanide in the electrolyzing beaker, 
dissolve in 50 cc of water and rinse into this the solution of copper and 
silver nitrates, stirring. Insert the electrodes, dilute until the cathode 
is covered with solution, mix and electrolyze as directed on page 160, 
keeping the voltage below 1.5. 

Copper. — To the solution from which silver has been deposited, 
including the washings from the latter, add 5 cc of concentrated nitric 
acid. This should be done under a hood as the vapors of hydrocyanic 
acid are extremely poisonous. Evaporate to 75 cc or less, to expel all 
of the hydrocyanic acid. Determine the copper as directed on page 156. 


Iron does not deposit well from -solutions containing nitrates, 
chlorides or strong inorganic acids. If the iron salt is derived 
from such an acid this acid will be formed as electrolysis proceeds 
and iron will redissolve from the cathode. To prevent its forma- 
tion a salt of a weak inorganic or organic acid may be added, 


Examples of salts that are used for this purpose are ammoniur 
oxalate, ammonium tartrate and sodium citrate. There is alway 
a possibility of depositing some carbon from such solution 
and this is least likely to occur when the oxalate is used. Th 
iron or iron salt should be converted into the sulphate befor 

Determination. — Calculate the weight of sample that will be require 
to give about 0.3 gm of iron. If the sample is an iron salt, solub] 
in water, dissolve in water, using no more than is necessary. If th 
sample is iron or steel, dissolve in the proper quantity of dilute su 
phuric acid, avoiding an excess. In either case, pour the solution 
slowly into a solution of 5 to 10 gm of ammonium oxalate in as little water 
as possible, stirring until the precipitate of iron oxalate is redissolved. 
Dilute, after the electrodes are in place, until the solution covers the 
cathode. The decomposition voltage for iron in such a solution is about 
andiVDioo= 0.1. The deposited iron should be bright. It does n 
easily redissolve in the solution when the current is interrupted and ma 
be readily washed. The end of the process is tested by the use of pot a 
sium ferri cyanide or potassium thiocyanate. If the latter is used t 
solution should be previously warmed with two or three drops of nitri 
acid, since the iron (if any is present) has been reduced to the ferrous co: 
dition by the current action. After weighing the iron it should be re- 
moved from the cathode by dissolving in warm dilute sulphuric acid. 


When salts of lead in solution are subjected to the action of 
current the lead may deposit upon both cathode and anode 
upon the former as elementary lead and upon the latter 
lead peroxide. Upon the cathode the lead is so spongy that 
becomes impossible to wash and dry it properly. This is the 
familiar action of the lead storage cell during charging, where 
lead sulphate is electrolyzed, producing spongy lead at the 
" negative" and lead peroxide at the "positive." If the lead 
salt solution contains a considerable excess of nitric acid the 
entire quantity of lead deposits upon the anode as peroxide and 
this is the only practicable quantitative method for the electroly- 
sis of lead salts except where the mercury cathode is used. 
Lead dioxide cannot be completely dehydrated unless heated to 
a temperature above 200°. 


The deposition of lead peroxide upon the anode when nitric 
acid is present is to be explained exactly as in the case of silver. 
The solution contains a small concentration of the amphoteric 
lead perhydroxide, Pb(OH) 4 , which furnishes anions of an oxy- 
acid of lead, H 4 Pb0 4 . The electrolysis of this acid and dis- 
charge of its anion produces lead peroxide and oxygen: 

Pb04-*Pb0 2 +0 2 . 

As in the case of silver the formation of the peroxide may be 
prevented by the addition of some substance which will diminish 
the concentration of lead cations, such as ammonium oxalate. 
Since the cathodic deposit of lead cannot well be washed and 
dried without loss or oxidation the anodic deposition of peroxide 
is assisted by addition of nitric acid. 

Determination. — Weigh enough lead salt (preferably nitrate) to 

produce about 0.3 gm of lead, dissolve in the proper quantity of water 

to cover the electrodes and add 20 cc of concentrated nitric acid for 

each 100 cc of solution. Connect the electrodes so that the one with 

the largest surface will be the anode, instead of the cathode as is the 

case in the electrolysis of most other metals. The lead peroxide will not 

adhere well unless the anode surface has been roughened, as by sand 

blasting. Warm to about 50° and keep at this temperature until the 

electrolysis is finished. Use 2.4 volts. NDi 00 = 1.5 amp. At the end 

| of the operation test the solution for lead by adding a few drops of 

I hydrogen sulphide solution to a small amount of the electrolyte. Care- 

! fully remove the anode, having previously washed it in the usual way. 

Dry at 200° to 230° until the weight is constant. 


The best solution from which to deposit nickel is that of the 
I sulphate, containing ammonium sulphate and ammonium hy- 
j droxide. If nitric acid is present there is usually some trouble, 
i due to the oxidation of the deposited nickel. Nickel may also be 
precipitated from solutions containing ammonium oxalate, 
tartrate or citrate or from solutions containing an excess of potas- 
sium cyanide. 

Determination. Separation of Copper, Nickel and Iron. — Thoroughly 

j clean and dry a nickel coin, weigh it accurately, place in a casserole and 

dissolve in nitric acid (sp. gr. 1.2) the casserole being covered while the 

coin is dissolving. Carefully add 10 cc of concentrated sulphuric acid 


and evaporate over a flame until the characteristic dense, white fumes 
01 sulphuric acid appear. The evaporation should be accomplished 
while holding the casserole in the hand, giving it a continuous rotary 
motion to hasten evaporation and prevent spattering. Allow the mate- 
rial to cool then wash into a 250 cc graduated flask and dilute to the 

Measure 50 cc of the solution into the vessel in which electrolysis 
is to be accomplished and deposit the copper in the manner already 
described. The voltage should not exceed 2.7, which is nearly the 
decomposition voltage of nickel. Unusual care should be exercised 
in washing and saving the washings because other metals are to be 

Evaporate the solution from which the copper has been removed, 
until the volume is about 100 cc. Neutralize with ammonium hydrox- 
ide, boil to flocculate the colloidal ferric hydroxide which is always 
present and filter off the precipitate, washing the paper and precipitate 
with hot water. Add to the filtrate 7 gm of ammonium sulphate 
and 20 cc of ammonium hydroxide (sp. gr. 0.90), and electrolyze. 
Decomposition voltage is about 2.8 and any voltage above this value 
may be used, the upper limit being fixed by the nature of the deposit 
obtained. Record the current density. 

Dissolve the ferric hydroxide in the filter paper with 1 cc of oxalic 
acid solution, saturated at about 20°. Wash the solution out of the 
paper with hot water and into a solution of 5 gm of ammonium oxalate 
in 100 cc of water. Electrolyze as previously directed. 

Moving Electrodes. — In the discussion of decomposition 
voltage it was noted that if the voltage is unduly increased in 
order to hasten the decomposition, gas evolution prevents the 
formation of a dense deposit of metal. Migration of the ions 
is comparatively slow and current is transferred, not only across 
the solution, but also from the solution to the cathode, by hydro- 
gen ions. If the migration of the metal ions is aided by stirring 
the solution large currents may sometimes be carried without 
the deposition of enough hydrogen on the cathode to injure the 
metal deposit. Stirring may be accomplished by any one of 
five different methods: (1) Heating to produce convection 
currents, (2) use of stirring apparatus not connected with the 
electrodes, (3) rotation of the anode, (4) rotation of the cathode, 
(5) electromagnetic action. Convection currents are of limited 
usefulness because they are not sufficiently rapid. They will, 



lowever, materially shorten the time of electrolysis. Mechanical 
stirring, whether by the second, third or fourth method, has 
practically the same use. Which method of these three is 
to be chosen will be decided chiefly by matters of convenience, 
[f a stirrer of glass or other non-conducting material is to be 
used it will require room for its movement and, since it is as 
sasy to rotate one of the electrodes, the stirrer is generally made 
one of these. Either electrode may be rotated 
with success. The anode is generally the one 
chosen for this purpose because it has not the 
arge surface of the cathode and is therefore more 
easily manipulated. Fig. 50 shows one of the 
forms of anodes that may be rotated rapidly 
without becoming bent or distorted. The 
cathode is most frequently a dish. In order to 
provide a dish of large surface and capacity 
Classen devised a very thin platinum dish. The 
speed of rotation of the anode should be as high 
as may be attained without danger of throwing 
the solu fcion out of the vessel. 500 to 1000 r.p.m. 
may be used. Recently a comparatively slow 
rotation (150 r.p.m.) of the ordinary spiral anode 
inside a gauze cathode has been used, with a con- 
siderable degree of success. Indeed, it is ques- 
tionable whether this is not more practicable than 
the rapid rotation because the gauze electrodes 
may be used without danger and the apparatus 
needs no watching after starting. The time 
necessary for complete deposition of the metal 
may be made about one-fifth of that required 
when stationary electrodes are used. 

In the electromagnetic stirring apparatus de- 
vised by Frary 1 the solution is placed within an 
electromagnetic field, generated by a current 
passing through a solenoid surrounding" the electrolyte. The 
moving ions constitute a conductor in which the current moves 
radially while the electromagnetic field is vertical. In another 
form the apparatus is changed using a vertical field and vertical 

1 J. Am. Chem. Soc", 29, 1592 (1907). 

Fig. 50.— 
Anode suitable 
for rotating. 



current lines. In either case the mutual action of the fieh 
causes rotation of the solution within. 

The Mercury Cathode. — By making the cathode of mercur 
instead of platinum two important gains are made. The lar$ 
expenditure for electrodes is largely eliminated because of th 
relative cheapness of mercury, also there is no longer any ques 
tion as to the satisfactory nature of the deposit of metal, sine 
the latter amalgamates with the mercury instead of forming 
surface deposit. The one obstacle to a nearly universal use < 
this iorm of electrode lies in the small surface that may be use( 

Fig. 51. — (a) Mercury cathode cell with (b) drying apparatus. 

the large specific gravity of mercury prohibiting the use of more 
than a few cubic centimeters. The most satisfactory form of 
apparatus for this purpose is a glass cup having a small platinum 
wire fused into the bottom for the cathode connection (Fig. 51). 
The upper limit to the normal density is approximately fixe( 
by the undesirable heating effect of large current densities. Th< 
anode should be rotated. 

In the following exercises are given the changes necessary t< 
adapt the foregoing exercises to the use of rotating electrodes an< 
the mercury cathode. 

Determination of Copper by Use of the Rotating Anode. — Set u] 
the apparatus and add to the solution 1 cc of dilute sulphuric aci< 



various devices are used. In the Schellbach burette, Fig. 55, a 
background of white glass bears a stripe of blue. The meniscus 
appears against this as a point. This improvement is of doubtful 
value except in rooms where the light is not good, because it 
does not prevent parallax. The use of floats, Fig. 56, sometimes 
renders readings more easily made, especially when the liquid is 
so dark in color as to be nearly opaque. The mark on the float 
is brought so near to the side of the burette that parallax is also 
largely prevented. Trouble due to sticking of the float is suf- 
jficient cause for dispensing with its use whenever possible. The 

Fig. 54.— Effect of Parallax. 

Fig. 55. — Meniscus as seen in 
Schellbach burette. 

best construction of burettes yet devised is that specified by the 
U. S. Bureau of Standards. 1 Upon burettes made according to 
jjhese specifications the marks for whole cubic centimeters extend 
J3ntirely around the burette while those for subdivisions extend 
half way around. This arrangement absolutely obviates the 
troubles of parallax and makes quite sharp readings possible. 

I Certain devices are sometimes used for promoting rapid fill- 
ing of the burette. Fig 58 illustrates some of these. To 
,?et up such a burette and keep it in working order requires 

I I certain amount of attention and such devices are of value 
chiefly in works laboratories where large numbers of routine 
leterminations are to be made by means of the same standard 

Bur. Stand., Circ. No. 9. 



solution. The burette with a plain glass cock is most serviceabl 
for ordinary use. All burettes should be covered by a cap wh( 
in use. This excludes dust and lessens evaporation of tl 

Units of Volume. — For the requirements of volumetric analys 
the same accuracy may be obtained without regard to the pai 
ticular unit of volume adopted, provided that all of the differei 
pieces are calibrated upon the basis of the same unit. The liter 
defined by the International Bureau* of Weighl 
and Measures to be the volume occupied 
4° C. by water having a mass of 1 kg. Thi 
is almost exactly 1000 cc and for all practical 
purposes may be regarded as such. (The 
milliliter is 1.000029 cc.) It is now customary 
to use this true liter as the standard, calibrat- 
ing apparatus upon this basis, the apparatus 
to be used at the average room temperature. 
In America the working temperature is usual] 
taken to be 20°. Other temperatures are use 
as standard working temperatures, partici 
larly abroad where 17.5°, 15.5° or 15° is us 
as a calibration and working temperature. 
When the true liter is made the basis of ct 
bration and higher temperatures than 4° are used for the expei 
mental part of the calibration corrections must be made for tl 
difference in density of water used for calibrating and, if tt 
water is weighed, also for the buoyant effect of air. In on 
to avoid making such corrections Mohr 1 suggested a different 
unit, the "Mohr liter, " which is defined to be the volume of 
1000 gm of water weighed in air at a standard pressure of 760 
mm of mercury and at a temperature of 17.5°. 

Absolute or Relative Capacities. — In the discussion of the 
calibration of weights (page 66) it was stated that true gram 
values were not required; since accurately known relative values 
of the various pieces are sufficient for analytical values, analytical 
results being always stated in some sort of ratio of the con- 
stituent determined to the material analyzed. For a similai 

1 Lehrbuch, Chemisch-analytischen Titriermethoden, 6th ed. (Rev. b) 
A. Classen), 42. 

Fig. 56.— Burette 





reason the various pieces of volumetric apparatus may be 
graduated upon any desired basis, true liter or otherwise, pro- 
vided only that the different pieces shall have correctly indicated 
relative capacities. However, it is very desirable that all pieces 
that are to be used in a given laboratory shall be interchangeable 

Fig. 57. — Form of burette approved 
by the Bureau of Standards. 


58. — Burette with automatic fill- 
ing and overflow devices. 

and the adoption of a common standard for all workers in a 
laboratory is a practical necessity. In such a case calibration 
to a basis of the true liter, using weights that are calibrated in 
true gram values, is the logical procedure. 

Tolerance. — Certain experimental errors in the graduation of 
volumetric apparatus may be regarded as reasonable errors, on 
account of which the apparatus should not be rejected. This 


does not mean that corrections should not be made after~the 
results of calibration are known. Maximum permissible errors 
in graduation are known as " tolerances" and if the tolerance 
is exceeded in a given case the piece should not be used without 

The following quotation from the bulletin of the U. S. Bureau 
of Standards already referred to gives the requirements of that 
bureau for volumetric flasks, burettes and pipettes to be accepted 
for testing. These requirements are recognized as, at once, 
rigid and scientific. Apparatus used, even by the student 
beginning the study of volumetric analysis, should, whenever 
possible, conform to these requirements. 

General Specifications 

" (a) Units of Capacity. — The liter, defined as the volume occupied by 
a quantity of pure water at 4° C. having a mass of 1 kg, and the one- 
thousandth part of the liter, called the milliliter or cubic centimeter, are 
employed as units of capacity. 

"(b) Standard Temperature. — 20° C. is regarded by the bureau as the 
standard temperature for glass volumetric apparatus. 

" (c) Material and Annealing.- — The material should be of best quality 
of glass, transparent and free from striae, which adequately resists chem- 
ical action, and has small thermal hysteresis. All apparatus should be 
thoroughly annealed at 400° C. for 24 hours and allowed to cool slowly 
before being graduated. 

"(d) Design and Workmanship. — The cross section must be circular 
and the shape must permit of complete emptying and drainage. 

"Instruments having a base or foot must stand solidly on a level sur- 
face, and the base must be of such size that the instruments will stand 
on a plane inclined at 15°. Stoppers and stopcocks must be so ground 
as to work easily and prevent leakage. 
• "The parts on which graduations are placed must be cylindrical for at 
least 1 cm on each side of every mark, but elsewhere may be enlarged to 
secure the desired capacities in convenient lengths. 

"The graduations should be of uniform width, continuous and finely 
but distinctly etched, and must be perpendicular to the axis of the appa- 
ratus. All graduations must extend at least halfway around, and on 
subdivided apparatus every tenth mark, and on undivided apparatus 
all marks must extend completely around the circumference. 

"The space between two adjacent marks must not be less than 1 mm. 
The spacing of marks on completely subdivided apparatus must show 




no evident irregularities, and sufficient divisions must be numbered to 
readily indicate the intended capacity of any interval. Apparatus 
which is manifestly fragile or otherwise defective in construction will 
not be accepted. 

" (e) Inscriptions. — Every instrument must bear in permanent legible 
characters the capacity in liters or cubic centimeters, the temperature 
in Centigrade degrees at which it is to be used, the method of use, i.e., 
whether to contain or to deliver, and on instruments which deliver 
through an outflow nozzle the time required to empty the total nominal 
capacity with unrestricted outflow must be likewise indicated. 

" Every instrument should bear the name or trade-mark of the maker. 
Every instrument must bear a permanent identification number, and 
detachable parts, such as stoppers, stopcocks, etc., belonging thereto, 
must bear the same number. 

Special Requirements 

"(a) Flasks. — At the capacity mark or marks on a flask the inside 
diameter should be within the following limits : 

Capacity of flask (in re) up to and 


Maximum diameter (in mm) .... 
Minimum diameter (in mm) 














"The neck of a flask must not be contracted above the graduation 

"The capacity mark on any flask must not be nearer the end of the 
cylindrical portion of the neck than specified below: 


Distance from upper 
end, cm 

Distance from lower 
end, cm 

1 00 cc or less 




More than 100 cc 


"Flasks of 1 liter or more but not less may be graduated both to 
contain and to deliver, provided the intention of the different marks 
is clearly indicated. 

" (b) Transfer Pipettes. — Pipettes for delivering a single volume are 
designated "transfer" pipettes. 

"The suction tube of each transfer pipette must be at least 16 cm 
long, and the delivery tube must not be less than 3 cm nor more than 
25 cm long. 

"The inside diameter of any transfer pipette at the capacity mark 
must not be less than 2 mm and must not exceed the following limits: 



Capacity of pipettes (in cc) up to and including 25 

Diameter (in mm) 4 



"The outside diameter of the suction and delivery tubes of transfer 
pipettes exclusive of the tip must not be less than 5 mm. 

"The capacity mark on transfer pipettes must not be more than 6 
cm from the bulb. 

"The outlet of any transfer pipette must be of such size that the 
free outflow shall last not more than one minute and not less than the 
following for the respective sizes : 

Capacity (in cc) up to and including. 
Outflow time (in seconds) 











"(c) Burettes and Measuring Pipettes. — Only those emptying through 
a nozzle permanently attached at the bottom are accepted for test. 

"So-called "Shellbach" burettes — that is, those having amilk-glas 
background with a colored center line — will not be accepted for test. 

"The distance between the extreme graduations must not exceed 
65 cm on burettes nor 35 cm on measuring pipettes. 

"The rate of outflow of burettes and measuring pipettes must be 
restricted by the size of the tip and for any graduated interval the 
time of free outflow must not be more than three minutes nor less than 
the following for the respective lengths: 



Time of 


Time of 

graduated, cm 

outflow, sec 

graduated, cm 

outflow, sec 

























"The upper end of any measuring pipette must be not less than 10 cm 
from the uppermost mark and the lower end not less than 4 cm from 
the lowest mark. 

"On a burette the highest graduation mark should not be less than 
5 cm nor more than 15 cm from the upper end of the burette. 

"(d) Burette and Pipette Tips. — Burette and pipette tips should be 
made with a gradual taper of from 2 cm to 3 cm, the taper at the extreme 
end being slight. 

"A sudden contraction at the orifice is not permitted and the tip 
must be well finished. 



"In order to facilitate the removal of drops and to avoid splashing 
when the instrument is vertical, the tip should be bent slightly. 

"The approved form of tips for burettes, measuring pipettes, and 
transfer pipettes is shown in Fig. 59. 

"Special Rules for Manipulation. — These rules indicate the essential 
points in the manipulation of volumetric apparatus which must be 
observed in order that the conditions necessary to obtain accurate 
measurements may be reproduced. 

"(a) Test Liquid. — Apparatus will be tested with water and the 
capacity determined will, therefore, be the volume of water contained 
or delivered by an instrument at its standard temperature. 

Fig. 59. — Tips of burette and pipettes, approved by the Bureau of Standards. 

"(b) Method of Reading. — In all apparatus where the volume is 
limited by a meniscus the reading or setting is made on the lowest point 
of the meniscus. In order that the lowest point may be observed 
it^is necessary to place a shade of some dark material immediately 
below the meniscus, which renders the profile of the meniscus dark 
and clearly visible against a light background. A convenient device 
for this purpose is a collar-shaped section of thick black rubber tubing, 
cut open at one side and of such size as to clasp the tube firmly. 

" (c) Cleanliness of Apparatus. — Apparatus must be sufficiently clean 
to permit uniform wetting of the surface. 

"(d) Flasks and Cylinders. — In filling flasks and cylinders the entire 
interior of the vessel will be wetted, but allowed a sufficient time to 




drain before reading. Before completely filling to the capacity mark 
flasks should be well shaken to completely mix the contents. 

"Flasks and cylinders when used to deliver should be emptied by 
gradually inclining them until when the continuous stream has ceased 
they are nearly vertical. After half a minute in this position the mouth 
is brought in contact with the wet surface of the receiving vessel to 
remove the adhering drop. 

"(e) Pipettes and Burettes. — In filling pipettes and burettes excess 
liquid adhering to the tip should be removed when completing the 

"In emptying pipettes and burettes they should be held in a vertical 
position, and after the continuous unrestricted outflow ceases the tip 
should be touched with the wet surface of the receiving vessel to com- 
plete the emptying. 

"Stopcocks, when used, should be completely open during emptying. 

"Burettes should be filled nearly to the top, and the setting to the zero 
mark made by slowly emptying. 

"While under normal usage the measurements ordinarily are from the 
zero mark, other initial points may be used on burettes of standard 
form without serious error. 


(a) Flasks 

Capacity (in cc) less than 

Limit of error, cc 

and including 

If to contain 

If to deliver 

























(b) Transfer pipettes 

Capacity (in cc) less than and including 

Limit of error, cc 

















(c) Burettes and measuring pipettes 

Capacity (in cc) of total 

graduated portion less 

than and including 

Limit of error 

(in cc) of total or partial 



Measuring pipettes 












"Further, the error of the indicated capacity of any ten consecutive 
subdivisions must not exceed one-fourth the capacity of the smallest 

Calibration. — For accurate work the apparatus as supplied 
by the makers should never be regarded as correctly graduated 
until it has been tested (calibrated) by the user. Some manu- 
facturers use great care in graduating, especially with such ap- 
paratus as must pass inspection by one of the national standardiz- 
ing bureaus. Others are less careful and pieces are often found to 
have large errors in their graduation. Two general methods are 
in use for calibrating instruments for capacity. In the first 
the quantity of pure water or other liquid of known density 
which would exactly occupy the desired volume at the stated 
temperature is measured by weighing. The position of the 
meniscus is then compared with the mark upon the apparatus. 
If the latter was previously unmarked the position of the menis- 
cus is then marked. In the second method the capacity of the 
instrument is determined by allowing water or another liquid 
to flow into it from a previously standardized piece and the 
capacities are compared. In the first method temperatures 
must be accurately noted and corrections made for any departure 
from the standard temperature, also for air displacement. In 
the second method such corrections have already been made 
when the standard piece was calibrated and we have merely a 
comparison to make of two instruments of capacity. The second 
method is, therefore, shorter in point of time. Errors may occur 
if proper attention is not given to certain details of manipulation. 

Calibration by Weighing. — Water is the most conveniently 
used liquid for this purpose. Since water solutions are generally 


to be used in volumetric analysis water possesses a seeo: 
advantage in that the form of meniscus is most nearly that of 
solutions later to be measured. The problem in calibratin 
in the case of flasks or other pieces having but one or two mar 
and therefore easily remarked, is to determine the correct po: 
tion of the mark when the piece contains the rated quantity o\ 
liquid. For this reason, in laboratories where all of the apparati 
purchased is regularly calibrated it is best and cheapest 
purchase flasks unmarked, though conforming to certain spe< 
cations as to dimensions and shape. 

If the true liter is to be taken as a basis it will first be necessa: 
to determine the apparent weight of this volume of water in air, 
correcting also for the expansion due to the difference in tempera- 
ture between that at which the apparatus is calibrated and used 
and the temperature upon the basis of which the liter is defined 
(4°). From the table on page 181 it is seen that the density of 
water at 20° working temperature is 0.99823. 

One liter of water would therefore weigh 998.23 gm in a 
vacuum. In air both water and weights are apparently lighter 
than in a vacuum. If the density of these were equal the effect of 
air upon them would be equal and no error would be introduced, 
Since the density of the weights is greater and their volume less, 
the buoyant effect is greater upon the water and the apparem 
weight of the water in air is less than the true weight. One lite 
of air, at 760 mm pressure and at 20° and having a humidit; 
of 50 percent, weighs 1.19 gm. This is the buoyant effect upon 
the water. Analytical weights, whether plated or not, are 
usually constructed of brass. The density of brass may be taken 
as 8.4. The weight of air displaced by 998.23 gm of brass weights 
is then 

0.99823X1.19 n v. 
^ = 0.14 gm. 

This is the buoyant effect upon the weights. The difference 
1.19 — 0.14 = 1.05 gm is the apparent loss in weight of the liter 
of water when weighed in air. 998.23-1.05 = 997.18. There 
fore one liter of water at 20° and in air apparently weighs 997.18 
gm. The weight of fractions of a liter will be calculated from 




Density of pure water free from air, by tenths of degrees from 0° to 40° and under standard 



Tenths of degrees 













0.999 8681 










+ 59 












+ 41 












+ 24 












+ 8 


1.000 0000 










- 8 


0.999 9919 










- 24 












- 39 












- 53 












- 67 












- 81 












- 95 






























































0.998 9705 








































































0.997 8019 
















































0.996 8158 




































0.995 9761 
















































0.994 7325 




































0.993 7136 
























0.992 9960 




































0.991 8661 

1 Travaux et Memoires, Bur. Intern. Poids Mesures, 13 (1907.) 



Calibration at Other Than the Standard Temperature. — 

Carelessness with regard to the temperature of the water used 
in calibrating may lead to serious errors. The apparent weight 
of one liter of water at 25° is 996.04 gm and at 15° it is 998.05 
gm, instead of 997.18 gm. This gives a mean variation of 
±0.20 gm for each degree of variation from the standard tem- 
perature of 20°, within the limits of 15° and 25°. This error 
is partly compensated by the change in the actual capacity 
of the glass apparatus, due to contraction or expansion with 
change in temperature. The average coefficient of cubical 
expansion of the glass used for volumetric apparatus is 0.000025 
for each degree of change in temperature. The correction is, 
of course, similar in sign to the temperature variation from 
20°, while the correction to be applied to the required weight of 
water is opposite in sign to the temperature variation. That is, 
if the temperature is above 20° a smaller weight of water should 
be taken for a true liter; but a glass vessel is actually larger 
and should have, from this standpoint alone, a greater weight 
of water to indicate a. mark which can be used at 20°. For 
this reason the net correction represents the difference between 
the two. 

The mean correction of 0.20 gm in the apparent weight of a 
liter of water is not sufficiently accurate for a range of more than 
a degree or two variation. The following table gives the apparent 
weight of a liter of water for various temperatures between 15° 

Temperature, degrees 

Apparent weight of 1 1 
of water in air, grams 

Weight to be taken 
for calibration 





















996.99 . 







24 • 








Fig. 60. — Calibrating device of the Bureau of Standards. 



and 25°, also the corrected weight of water to be taken for cali- 
brating flasks that are to be used at 20°. The latter weight 
includes the correction for expansion and contraction of glass. 
This table is to be used only in case it becomes difficult to main- 
tain the laboratory temperature at 20°. 

Calibration by Standardized Bulbs. — Any bulb or tube that is 
to be used as a comparison standard for calibrating by the second 


Fig. 61. — Morse-Blalock calibrating bulbs. 

method must be of such a form that it will drain well upon 
emptying. It must also have the graduated portion small 
enough to make possible accurate readings at this part. The 
apparatus must be quite rigid so that varying pressure may have 
only an inappreciable effect upon the volume. The apparatus 
shown in Fig. 60 is that used by the Bureau of Standards. 


In Fig. 61 are shown the standard bulbs devised by Morse 
and Blalock. 1 The three pieces shown provide a means for 
calibrating vessels of all of the various capacities common to 
the analytical laboratory, if proper combinations are made. 
In calibrating these bulbs it is necessary to determine the 
capacity from the single mark to the first stem division, also 
the capacity of the stem for the smallest subdivision. If the 

Fig. 62. — Morse-Blalock bulb arranged for calibrating flasks. 

water used is kept at the standard working temperature no 
correction for this factor need be introduced and the value 
997.18 gm, previously deduced as the apparent weight of a 
liter of water, may be used without change. The bulbs are 
supported in such a manner that they may be readily filled from 
a reservoir of distilled water at 20°. The water from the bulb 
is carefully weighed and its volume calculated. That from 
the entire graduated portion of the stem is then weighed in a 
smaller vessel and the calculated volume is divided by the total 

1 Am. Chem. J., 16, 479 (1894). 



number pf stem divisions and the result recorded as the value 
of one division. In order to use the bulbs for the calibration of 
vessels that are to contain a specified volume of liquid, the 
vessel to be calibrated must first be cleaned as hereafter directed. 
The proper bulb is then placed in such a position that it may- 
drain directly into the instrument being calibrated and the 
latter is marked at the meniscus. Instruments to be calibrated 
to deliver, such as burettes, are better calibrated by weighing 
the water'delivered. If, however, it is desired to use the standard 
bulbs for this purpose the burette is so connected that it may 
empty into the bulb from below. The details of manipulation 
will be made clear in the exercises that follow. 

Cleaning Solution. — Prepare a cleaning solution by dissolving 5 gm 
of powdered commercial sodium dichromate in 500 cc of commercial 
sulphuric acid. The solution may be kept in a bottle having a wide 
mouth, such as those in which dry chemicals are purchased. Burettes 
may be inverted and left standing in the bottle the solution then being 
drawn up by suction and held in the burette by closing the cock. For 
cleaning flasks the solution may be allowed to remain in the flask for some 
time or a small amount may be warmed and the flask rinsed with it. 
The chromic acid produced by the interaction of sulphuric acid an 
sodium dichromate oxidizes all organic matter and leaves the glas; 
thoroughly free from it. 

Exercise: Calibration of the Standard Bulbs. — Clean the bulb an 
set it up in a manner similar to that shown in Fig. 60. Distilled wate: 
must be used. Place the bulb so that the graduated stem extends down- 
ward. A glass stopcock must be used for controlling the outflow, sine 
pinchcocks with rubber connections would involve an uncertain change 
in volume of the apparatus. The tip of the outflow tube should be 
contracted to restrict the outflow, according to the size of the bulb, as 
follows : 

Size of bulb, cc 

Time of outflow, sec 



Fill the bulb to the upper mark. With the stopcock wide open allow 
the water to flow into a previously dried and weighed flask until the 
first division (zero) on the lower stem is reached. After 15 seconds 


adjust the water level to coincide with the mark, put the stopper in the 
flask (a glass stopper is desirable) and weigh. Place a weighing bottle 
under the delivery tube and allow the stem to drain to the last mark, 
stopper the bottle and weigh. If the water was at 20° divide its weight 
in grams by 0.99718 and record the result as cubic centimeters capacity 
of bulb and of stem. Record also the value of each stem division and the 
division to be used as the mark for the rated capacity of the bulb. If 
the temperature was not 20° determine from the table on page 182 the 
weight of water that should be used at the observed temperature 
for calibrating bulbs to be used at 20°. 

In the preceding exercise as ■ in those that follow the bulb 
and reservoir may also be set up as in Fig. 62. The chief 
objection to this method of assembling lies in the fact that 
the water entering the bulbs must, for reasons already explained, 
pass through a glass stopcock, which is necessarily lubricated 
with some kind of grease. The result is that no matter how 
well the bulb may have been previously cleaned it acquires, at 
the first filling, a film of oil that absolutely prevents the proper 
draining at the next stage in the experiment. If the first arrange- 
ment is used and the bulb is filled from above a rubber connec- 
tion and pinchcock may be used and this annoyance avoided. 

Exercise : Calibration of Flasks by the Standard Bulbs. — With the 
proper bulb in the position used in the preceding exercise, and with the 
dried flask to be calibrated placed under the delivery tube, allow water 
to flow into the flask until the proper mark on the lower stem is reached, 
exactly following the directions as to time of outflow as given in the pre- 
ceding exercise. Avoid handling either bulb or flask at the parts con- 
taining the water as the temperature is thereby raised. It has already 
been explained that after the bulbs have been standardized they may be 
used for further calibrations without regard to the temperature of the 
water provided only that the temperature does not change during the progress 
of the experiment. 

To mark the flask cut a strip of gummed label, long enough to reach 
around the neck and about 1/4 inch wide. Carefully paste this with the 
original straight edge at the level of the meniscus, where the mark is to 
be made. Melt a small quantity of paraffin and brush a thin layer over 
the label and over a space of about 1 inch on either side of it. Using 
the point of a knife or of a sharpened piece of wood trace the straight edge 
of the label around the neck of the flask, making a mark sufficiently wide 
to be easily visible. The label here merely serves as a guide, making a 


regular line possible. Using a small feather as a brush apply a few drops 
of hydrofluoric acid and allow this to remain on the flask for two or three 
minutes, after which the acid may be washed off and the paraffin 
removed by warming. 

In case the flask already has a graduation and the calibration shows 
this mark to be incorrectly placed it is desirable to indicate the new 
mark by making a small, well-defined arrow with the point resting 
exactly upon the new mark. The operator's initials may be placed 
beside the arrrow and if this is done carefully, no interference will 

If the flask contains no inscription etch the side like that shown in 
Fig. 52, page 169. 

Exercise: Calibration of Flasks by Weighing. — This is, in many 
respects, the most satisfactory method of calibration although more 
time is required. Have the flask clean and quite dry. Place on a 
balance of capacity sufficiently great to carry the filled flask. Counter- 
poise, then add weights to the right pan at the rate of 997.18 gm for each 
liter. Remove the flask from the balance and fill with recently boiled 
distilled water at 20°, nearly to the point where it is thought that the 
mark will be placed. Remove drops from the inside of the neck, above 
the level of the water, using a roll of filter paper. Replace the flask 
upon the balance pan, then carefully drop in water from a pipette until 
the balance is in equilibrium. Mark as directed in the preceding 

In both of these exercises the flasks have been calibrated to 
contain the rated quantity and this is the only way in which 
flasks will be used in this work. 

Exercise : Calibration of Burettes by Weighing. — The marking of a 
burette is too complex to be easily changed and the calibration will 
therefore consist of finding what, if any, corrections must be applied 
to the existing graduations. 

First inspect the burette to determine whether it conforms to specifi- 
cations, especially with respect to outflow time. Clean the burette with 
cleaning solution and distilled water. Fill with distilled water at 20°. 
Weigh accurately a 25 cc weighing bottle to the third decimal then 
measure 5 cc of water into it from the burette, and reweigh. Add 
another 5 cc and weigh, continuing until the bottle is full. Empty the 
bottle, reweigh and continue the process until the water from the entire 
graduated portion of the burette has been weighed. Repeat the process 
in order to have a check upon the work. Calculate the true capacity 
of each of the ten portions, using the weight 0.99718 gm for 1 cc of 


water. Record as follows, the capacities in the last two columns being 
recorded only as far as the second decimal place. 

Mark Weight of water, 

iviarx. each interval. 

True capacity, each True total capacity, zero to end 
interval. of interval. 

Construct a curve showing the true reading at all points. In case 
any marked irregularity is observed at any part of the burette so that 
corrections taken from the curve would be inaccurate, recalibrate this 
portion, using 1 cc at a time. 

Exercise: Calibration of Burettes by the Standard Bulbs. — Set up 
the apparatus as in Fig. 63. The reservoir must be higher than the 
top of the burette and this, in turn, must be placed so that the lowest 
graduation is higher than the bulbs. The tubing leading from the reser- 
voir to the burette may be of well-cleaned rubber. That between the 
stopcock a and the burette and bulbs respectively must be of glass, the 
necessary connections being of heavy rubber tubing with the glass 
tubes pushed together until they touch inside the rubber. 

With the three-way cock b closed, open the cock a and fill the burette 
with water. Close a and open b so that the 2 cc and 3 cc bulbs may 
fill, then drain the burette to the zero mark and the bulbs to that mark 
on the stem of the 2 cc bulb which represents exactly 2 cc. This 
leaves the bulbs moistened as they will be throughout the experiment. 
Leave the burette cock open. Turn the cock b and measure 5 cc of 
water from the burette into the small bulbs. Observe the position of 
the meniscus upon the stem of the 3 cc bulb and calculate the true 
capacity of the first portion of the burette, using the values for the stem 
divisions as determined in the calibration of the bulbs. Repeat this 
process for the other nine portions of the burette and record as follows: 

Mark Bulh rpadinc- True capacity, each True total capacity, zero to end 

*• interval. of interval. 

For a more nearly complete calibration the burette may be calibrated 
2 cc at a time, using the 2 cc bulb alone or 1 cc or any fraction at a time 
using the standard tube, Fig. 64. Such calibration is necessary if the 
bore of the burette is found to be very irregular. 

Calibration of Pipettes. — Pipettes which are graduated in 
small subdivisions from zero to full capacity (" measuring pi- 
pettes") are calibrated in the same manner as are burettes. 
Transfer pipettes are best calibrated by the method of weighing. 

Exercise : Calibration of Transfer Pipettes. — Determine whether the 
time of outflow conforms to the requirements as set forth on page 176. 



If not alter the tip of the pipette before calibrating. Provide a weigh- 
ing bottle having a capacity of 10 cc, also a larger one having a capacity 
equal to that of the pipette. Cut a strip of paper about 2 mm wide 

Fig. 63. — Morse-Blalock bulb set up for calibrating burettes. 

and 5 cm long and carefully rule this in divisions of centimeters, im 
ing from to 5, and subdivisions of millimeters, using fine lines. Del 
mine the approximate location of the capacity mark on the pipet 
by a rough experiment, unless the pipette is already marked. Pa 



the paper strip on the stem of the pipette with the division 2.5 at the 
supposed place for the capacity mark and with the zero toward the point 
3f the pipette. Having cleaned the pipette with chromic acid solution 
;t is drawn full of distilled water which is at a temper- 
iture of 20°, and the water is allowed to flow out 
jntil the zero mark is exactly reached. The pipette 
must be held in a vertical position and the eye must . 
oe in the same horizontal plane as is the meniscus. 
The pipette tip is now touched against the side of 
the beaker to remove the last drop. The finger is 
ohen removed from the top of the pipette and the 
water is allowed to flow, at full speed, into the larger 
weighing bottle, which has already been weighed. 
The tip is immediately touched to the side of the 
weighing bottle to remove the hanging drop. The 
weighing bottle is then stoppered and weighed. Cal- 
ulate the volume of the water from the observed 
(weight and record this as the capacity of the pipette 
io the zero mark. 

Using the small weighing bottle determine in a 
similar manner the capacity of the pipette stem 
between and 5. Divide this capacity by 50 
in order to obtain the value of the smaller sub- 

From the capacities so determined calculate the 
number of stem divisions to be added to the zero in 
Drder to obtain the rated capacity of the pipette. 
Mark the point so determined, using the method 
directed for marking flasks. 

Calculation of the Results of 
Volumetkic Analysis 

Fig. 64. — Morse 
Blalock tube. 

Although the first exercises in volumetric 
analysis will necessarily have to do with the 
making of solutions and with their standardiza- 
tion and adjustment to desired concentration, 
it will be simpler to deal first with the calcula- 
tion of the results of the analysis. When making the determi- 
nation of silver by the gravimetric method, a definite amount 
of the silver compound was [weighed, dissolved in water and a 
'slight but somewhat indefinite excess of hydrochloric acid was 


added, thus precipitating all of the silver as silver chloride, the 
following reaction taking place: 

AgN0 8 +HCl->AgCl+HNOs. 

The silver chloride, representing the entire amount of silver 
present in the compound of unknown composition, was then 
filtered, washed, dried, and weighed and from this observed 
weight and the known weight relations between silver chloride 
and silver the percent of the latter was calculated. The formula 
AgCl expresses the fact that for each 143.34 parts, by weight, of 
silver chloride, there was involved 35.46 parts of chlorine arid 
107.88 parts of silver. In other words, in any given weight of 

silver chloride, 14 » 04 of this weight is chlorine and ., , ' , is: 

silver. If the weight of silver chloride found in the analysis 1 

1 cyj 00 

is multiplied by the fraction .. . * , and the result divided by] 

the weight of sample taken, the quotient will be, when multiplied^ 
by 100, the percent of silver in the sample. If W = the weight 
of silver chloride found, and S = the weight of sample taker 
this would be expressed shortly as follows: 

107.88 IF X 100 , .. 

-iAoo± o = percent of silver in sample. (: 

Instead of adding the hydrochloric acid in slight but indennil 
excess to the solution of the silver salt, one might add exactl 
the amount required to complete the reaction, but no moi 
This would involve the use of some method for determining 
when the reaction is exactly completed (the "end point")? sucb 
as noting when another drop of hydrochloric acid solution fails; 
to produce any further precipitation of silver chloride. Suppose 
also, that the concentration of the hydrochloric acid solution, H 
grams per cubic centimeter, were very accurately known. W< 
should then have the following data : 

(a) Weight of silver salt taken, 

(6) Volume of hydrochloric acid required to react with silver 

(c) Concentration of hydrochloric acid. 

Just as the formula for silver chloride expresses the weigl 
relations between silver, chlorine and silver chloride, so ih 


jquation for the reaction between silver salt and hydrochloric 
Lcid expresses the weight relations between all of the elements 
,nd compounds involved. We are here particularly concerned 
14th the relations between silver and hydrochloric acid, and we 
liote that for every 107.88 parts, by weight, of silver, we require 
!l6.468 parts of hydrochloric acid for complete precipitation of 
he silver as silver chloride. Conversely, if the reaction has been 
exactly completed, for every 36.468 parts of hydrochloric acid 
jised there will have been present 107.88 parts of silver. The 
Iveight of pure hydrochloric acid used is found by multiplying 
|,he number of cubic centimeters by the concentration in grams 
per cubic centimeter, i.e., 

VC = Wt. HClused, 

inhere F = cc of acid solution used and C = gm hydrochloric 

107 88 
acid in 1 cc. If this weight is multiplied by the fraction „ fi ' 

jbhe result will be the weight of silver in the sample. Expressed 

107.88 VCX 100 

36.46 S 

= percent silver. (II) 

This is the most general expression for the calculation of the 
results of a volumetric analysis. T A comparison of expressions 
(I) and (II) will show that the volumetric calculation differs 
from the gravimetric calculation in two respects only: (1) A 
weight of a substance is obtained indirectly by measuring the 
volume of a solution of known concentration, instead of directly 
by weighing the substance. (2) The substance whose weight is 
desired is one which reacts with the substance being determined 
instead of one which is produced by this substance. The gravi- 
metric factor for the ratio of the weight of one substance to the 
weight of another substance which contains the first becomes the 
volumetric factor for the ratio of the weight of one substance to 
the chemically equivalent weight of another substance which 
does not contain the first .7 

The substance which is a visible indication of the end point of 
the reaction is called the "indicator." Indicators will be 
discussed at length in a later section.^ The solution whose con- 



centration is accurately known and of which we measure the 
volume required, is called a " standard solution" because it is 
actually a standard by which the quantity of the substance under 
investigation is measured. The process of running in the stand- 
ard solution until the end point is reached is called " titration." 
The examples given below will serve to illustrate the principle 
above outlined: 

1. 0.5436 gm.of a silver salt was dissolved and titrated by 
standard solution of hydrochloric acid, 1 cc of which containe 
0.00304 gm of the pure acid. 27.2 cc of the standard was used. 
Required, the percent of silver in the salt. 

27.2 X 0.00304 = gm of pure acid used; 

107 88 

qfi A(\ Xg; m of hydrochloric acid = gm of silver present. 

_ , 107.88X27.2X0.00304X100 AKnt% 
Therefore 36.46X0.5436 = 45.00 = percent sil- 
ver in the original sample of silver salt. 

Use of Aliquot Parts. — It often happens that, in order to elimi- 
nate the error due to the lack of uniformity of a sample being 
analyzed, or for reasons of convenience, a larger quantity than is 
necessary for the. titration is weighed and this dissolved in 
definite quantity of solvent and an aliquot part taken for th< 
titration. In such a case the final calculation must include ai 
expression of this fact. Thus in\Example*(l) instead of weighing 
0.5436 gm of the silver salt, suppose|that 10.8720 gm was 
weighed, dissolved, and the solution diluted to 1000 cc, 50 cc 
being then titrated. The statement would then be 
27.2X0.00304X107.88X20X100 : . ' . ., , 
10.8720X36.46 "= Percent silver m the sampl( 

2. In a sample of undried, but otherwise pure, sodium hydrox 
ide, the percent of the base was to be determined. For this? 
purpose 5.5310 gm of the sample was weighed and dissolve( 
and diluted to 250 cc. Portions of 25 cc each were measurec 
and titrated by a standard solution of hydrochloric acid, the 
average amount of acid solution required for complete neutrali- 
zation being 45.1 cc; 1 cc of the standard acid containec 
0.00960 gm of the pure acid. Required, the percent of sodiui 
hydroxide. The solution of the problem is as follows: 


45.1 X 0.0096 = gm of hydrochloric acid used; 
L ' Xgm of hydrochloric acid = gm of sodium hydroxide 

■SO . 40 

■n 25 cc of solution. 


t^Xgm of sodium hydroxide = gm in 5.5310 gm of sample. 

The condensed expression is 

= 85.89 = percent sodium 


The "equivalent weight" of a substance is the number of£ — 
might-units chemically equivalent to eight weight-units of oxygen. 
This definition is sufficiently broad to apply to any system of 
weights, although there are very few cases in scientific work . 
where "grams" might not be substituted for "weight-units," 
ince the metric system is quite universally accepted and used 
n scientific work. The result of this substitution is the "gram- 
jquivalent." In the effort to determine what is the equivalent 
weight of a substitute it is always necessary to inspect the 
aquation for the reaction that occurs in that particular case. In 
the reaction 

HCl+NaOH-*NaCl+H 2 0, 

it is easily seen that the equivalent weights for all of the elements, 
radicals, ions or compounds are the atomic, radical, ionic or 
molecular weights, respectively. In the reaction: 

H 2 S0 4 +BaCl 2 ->BaS0 4 +2HCl, 

the equivalent weights of the compounds are seen to be one- 
tialf of their molecular weights, with the exception of hydro- 
chloric acid, whose equivalent weight is its molecular weight. 
This will be more easily understood if we first determine the 
"hydrogen equivalent" of each substance, this being the number 
of atoms of hydrogen chemically equivalent to one molecule, atom, 
radical, etc., of the substance under consideration, as denoted by the 
equation for the reaction that has taken place. The hydrogen 
equivalent of sulphuric acid is 2 because it reacts by substituting 
another element for two hydrogen atoms. That of barium 
chloride is 2, because one atom of a bivalent element gives place 


to hydrogen or because 2 atoms of a univalent element 
place to one radical which is bivalent. That of barium sulphai 
is 2 and that of hydrochloric acid is 1, for similar reasons. Sin< 
one atom of oxygen is chemically equivalent to two atoms 
hydrogen, it follows that any body that is equivalent tc 1.1 
weight units of hydrogen will be equivalent to 8 weight units 
oxygen and therefore the equivalent weight is the molecular (atomi 
etc.) weight divided by the hydrogen equivalent. 


Find the equivalent weights of the substances whose formulas are in boh 
face in the following equations : 

4. 2HCl+Na 2 C0 3 ^2NaCl+H 2 C0 3 . 

5. HCl+NaHC0 3 ^NaCl+H 2 C0 3 . 

6. HCl+Na 2 CO,-»NaHC0 8 +NaCl. 

7. (NH 4 ) 2 C 2 04+CaCl 2 -*CaC204+2NH 4 CL 

8. (NH 4 ) 2 C 2 04+HC1-^NH4C1+NH 4 HC 2 04. 

9. (NH4) 2 C 2 4 +2HC1-*2NH 4 C1+H 2 C 2 04. 

10. FeCl 2 +2AgN0 3 -»Fe(N0 3 ) 2 +2AgCl. 

11. 2FeCl 2 +Cl 2 -»2FeCl 3 . 

12. H 2 S0 4 +Zn->ZnS0 4 +H 2 . 

13. H 2 S0 4 +2H— >S0 2 +2H 2 0. 

14. CuCl 2 +2AgN0 8 ->Cu(NO») 2 +2AgCl. 

15. 2CuCl 2 +Fe->2CuCl+FeCl 2 . 

Calculation of the Weight of One Substance, Chemically Equn 

alent to a Stated Weight of Another. — In the solution of nearl 

all problems of quantitative chemistry there is involved a calci 

lation of the weight of one substance chemically equivalent to i\ 

given weight of another. In the examples just considered thi 

is a calculation of the weight- of silver or of sodium hydroxk 

equivalent to the weight of hydrochloric acid that is contain* 

in the quantity of standard solution used. In example (1) t] 

107 88 
was expressed as -^-^-FX 0.00304 and in example (2) 

„' *n VX 0.0096. It is easily seen that both of these a 

culations involve the multiplication of the weight of hydn 
chloric acid by a ratio of equivalent weights. From this fol 
lows the rule that to find the weight of one substance chemicall 


I equivalent to a stated weight of another, multiply the stated weight by 
I the fraction: 

equivalent weight of substance calculated 
equivalent weight of substance given 

! This is a simple and very useful rule and its application will 
! obviate the use of the more cumbersome rule of proportions. 
(Applied to gravimetric analysis the fraction given above is the 


16. What weight of oxygen is equivalent to 0.3460 gm of hydrogen, 
I direct oxidation to water being understood? 

17. What weight of carbon dioxide is equivalent to 0.5693 gm of carbon, 
I direct oxidation being understood? 

18. Calculate the weight of ferric chloride and of iron equivalent to 
10.5243 gm of chlorine, the following reaction taking place: 

2Fe+3Cl 2 ->2FeCl 3 . 

19. Calculate the weight of potassium hydroxide equivalent to 1.7521 
Igm acetic acid, assuming complete neutralization. 

20. In problems (4) to (15) calculate the weights of the substances whose 
formulas are in bold face, equivalent to 3 gm of the substances reacting 
with them. 

21. A solution of hydrochloric acid of specific gravity 1.05 contains 10 
percent by weight of the pure acid. What volume of the solution is required 
to precipitate the silver from 0.75 gm of silver sulphate? 

22. 0.4321 gm of impure potassium sulphide was oxidized to potassium 
sulphate and precipitated by barium chloride. 0.8035 gm of barium 
sulphate was produced. What was the percent of sulphur in the sample? 

23. What weight of tartaric acid is equivalent to 3.52 gm of sodium 
hydroxide, the following reaction taking place? 

H 2 C 4 H406+2NaOH->Na 2 C4H 4 06 + 2H 2 0. 

Use of a Standard Solution for the Titration of but One Sub-^ — 
stance. — When a standard solution is to be used for the titra- 
tion of but one substance the calculations will all involve the 
constants representing (1) equivalent weight of the active sub- 
stance in the standard - solution, (2) equivalent weight of the 
substance to be determined and (3) concentration of the standard 
solution. If the standard solution of example (1) is to be used for 
f , . . ,. . ., :. . 107.88X0.00304 
the determination of silver, the expression: 3fi~4fi 


contains quantities that are constants for all such determinations. 
These constants should then be combined in a single constant: 
0.00899. From what was said in the preceding paragraph this 
is seen to be the weight of silver equivalent to 1 cc of this par- 
ticular standard solution of hydrochloric acid. All such calcu- 
lations would therefore be made by the expression 

FX0.00899X100 .. /TTTX 

— o = percent silver (III) 

This is obviously a very simple calculation and such simplifica- 
tion is possible and should be made whenever a given standard 
solution is to be used for a considerable number of determinations 
of a single substance. 

Burette Reading a Direct Percentage Reading. — If some care 
were exercised in adjusting the weight of silver salt used for 
analysis where statement (III) is to enter the calculation, s( 
that exactly 0.8990 gm of sample were used, the expressioi 
would become 


q 899Q - - = percent silver 

whence, V = percent silver. 

In this case the volume of standard solution used is the percenl 
of silver in the sample. From this follows the rule: To maki 
the burette reading a percentage reading first calculate the iveigh 
of the titrated substance thai is equivalent to 1 cc of the standarc 
solution, then use 100 times this weight of sample. 


24. What weight of soda ash must be used for analysis in order that 1 c 
of the hydrochloric acid solution containing 0.0031 gm shall be equivalen 
to 1 percent of sodium carbonate, assuming complete decomposition? 

25. A standard solution of sulphuric acid contains 40.2 gm in 1000 cc 
What weight of potassium hydroxide must be taken so that each cubi 
centimeter of the standard acid required shall indicate 0.1 percent oi 
potassium hydroxide? 

26. A standard solution of barium hydroxide contains 20.35 gm in 100C 
cc. What weight of vinegar is necessary in order that 1 cc of bariu 
hydroxide solution shall indicate 0.1 percent of acetic acid in the vinegar? 

No System. — In the examples given above there was no definite 
basis for the choice of the concentration of the standard solution 


Ball that was required being an accurate knowledge of the existing 
IJconcentration. Thus, in the first example the standard hydro- 
Ijchloric acid contained 0.00304 gm in 1 cc, while in the second 

it contained 0.00960 gm in 1 cc. There is no connection, ap- 
[Jparent or real, between these concentrations; they were chosen, 

at least to a certain extent, at random, upon the assumption 
lithat the determination of the concentration (standardization) 

iwas carried out with the greatest possible accuracy but that 
in making the solution no particular care was exercised. The 
method of calculating analyses made by means of such standard 
solutions would in all cases be analogous to these examples 
and any substance can be determined by means of such solu- 
tions, provided that the reaction involved is definite, complete 
and well understood, and that the end point can be determined 

Normal System. — The calculations of volumetric analysis 
may be considerably shortened by the proper adjustment of 
the concentration of the standard solution. 

The reaction between hydrochloric acid and silver nitrate is 
3xpressed by the equation: 

HCl+AgN0 3 -+HN0 3 +AgCl, 

and expression (II) was deduced for the calculation of the per- 
cent of silver in a salt that had been titrated by a standard 
solution of hydrochloric acid. This expression was 

107.88 VCX 100 

■ .g = percent silver. 

By using the proper indicator in each case we might have 
completed such reactions as the following for the titration of 
the substances indicated : 

HC1 + NaOH-*NaCl + H 2 0, (a) 

HC1 + KOH->KCl + H 2 0, (b) 

HC1+NH 4 0H-^NH 4 C1 + H 2 0, (c) 

2HCl+Na 2 C0 3 ->2NaCl+H 2 0+C0 2 , (d) 

HCl+NaHC0 3 ->NaCl+H 2 0+C0 2 , (e) 

2HCl+Ba(OH) 2 ->BaCl 2 +2H 2 0, (f) 

2HCl+CaC0 3 -+CaCl 2 +H 2 0+C0 2 . (g) 



Many other substances may also be titrated by this sam 
standard solution and in each case the expression for the pe 
cent of the substance to be calculated would be the same as 
(II) with the exception that, for the equivalent weight (combininj 
weight) of silver (107.88) we should substitute the equivalen 
weight of the substance to be calculated; we should then have 

40.008 VCX 100 

36.46 S 
56. 108 VCX 100 

36.46 S 
35.05 7CX100 

36.46S ' 
17.03 FCXIOO 

36.46 S 
53 VCX 100 

36.46 S 
84.008 7CX 100 

36.46 S 
85.693 VCX 100 

36.46 S 
50.045 VCX 100 

36.46 5 

percent NaOH (a') ' 

= percent KOH (b' 

= percent NH 4 OH (c') 

= percent NH 3 (c ; 

= percent Na 2 C0 3 (d') 

= percent NaHC0 3 (e' 

= percent Ba(OH) 2 (f 

= percent CaCO 3 (g ; 

In all of these expressions for the percent of the varioi 
substances as titrated by a single standard solution, the onb 
difference lies in the equivalent weight of the substance. Th< 
volume of standard required will depend, among other things, 1 
upon the purity of the sample and, since this is unknown, the 
volume required cannot be predicted. The concentration of 
the standard is under control and may be arbitrarily fixed atf 
any desired figure. The equivalent weights concerned are 
constants, in any given case, and the weight of sample may be 
made whatever is desired. 

If the standard solution is made of such strength that the 
tumber of grams contained in 1000 cc will be represented by 
the equivalent weight (in the case of hydrochloric acid 36.46 gm) 
the concentration in grams per cubic centimeter will then be 



3.03646, and the fraction nc Aa , which is involved in all of the 


[expressions, will become o fi 4ft =0.001, so that we shall then 

0.04008 7X100 


0.03505 7X100 


percent NaOH (ai) 

percent KOH (bi) 

percent NH 4 OH (ci) 

land so on. 

The standard solution of hydrochloric acid thus made, con- 
taining 1 gram-equivalent of the active substance in 1000 cc 
of solution, is a solution of general application and the calcula- 
tion of the results of analyses of various substances is simplified 
by this choice of concentration. Such a solution is called a 
" normal solution" which will be defined as a solution containing 
1 gram-equivalent of the active substance in 1000 cc. 

From the foregoing discussion the following deductions may be 

1. 1 cc of any normal solution is equivalent to one-thousandth 
of one gram-equivalent ( = one milligram-equivalent) of any 
substance. This is because 1 cc of any normal solution contains 
one-thousandth of one gram-equivalent of the active substance. 

2. If the milligram-equivalent of the substance titrated in a 
given determination is multiplied by the number of cubic centi- 
meters of a normal solution used for the titration, the result is 
the weight of the former in the sample. This follows as a result 
of (1). 

3. 1 cc of any normal solution is equivalent to 1 cc of any 
other normal solution. This also follows as a result of (1). 

4. The relative volumes of various standard solutions equiva- 
lent to each other are inversely as the respective normalities of 
these solutions. Thus 6 cc of a fifth-normal solution is equivalent 
to 26 cc of a tenth-normal solution. 

These principles are very important and their intelligent ap- 
plication will serve to shorten many of the calculations of 
volumetric analysis. 


It frequently happens that the normal solution is too con 
trated or too dilute for convenient use in a given analysis, 
this case the advantage of the normal solution may be retai 
by making the concentration of the solution to be some sim 
multiple of the concentration of the normal solution, such 

2, 3, -St -r-xi -^., 77^:, etc. This factor must then be introduced 
5 1U 5U 10U 

into the calculations involving the solution. Solutions made of 

normal or a simple multiple of normal strength are said to be 

made in the " normal system" and are, for the sake of brevity, 

N N N 
designated as N, 2N, y^, ^, y^, etc. 



27. 1 cc of normal acid is equivalent to what weight of ammonium c 
bonate, assuming complete decomposition? 


28. 1.1256 gm of a silver alloy is dissolved and titrated by in potassium 

thiocyanate solution according to the following equation : 

KCNS + AgN0 3 ->KN0 3 + AgCNS. 

35.2 cc of standard solution is required. What is the percent of silver i 
the alloy? 


29. 0.5 gm of limestone was dissolved in 50 cc of -^ acid. The unused 

excess of acid was titrated by 16.2 cc of j~ base. What was the percent 

of calcium carbonate in the limestone? What percent of calcium? 


30. 0.4 gm of soda ash was titrated by 20.9 cc of -^ acid. What was 

the percent of sodium carbonate in the sample ? 

31. 0.5 gm of an ammonium salt was decomposed by sodium hydroxide 

and the resulting ammonia distilled into 50 cc of -= acid solution. The 

unused excess of acid was titrated by 29.3 cc of ^ base. What percent 

of ammonia in the salt? 

Decimal System. — Instead of using the normal system, a 
further simplification may be made by adjusting the standard 
until each cubic centimeter shall be equivalent, not, as in the 
normal system, to a decimal fraction of a gram-equivalent of the 


(substance to be titrated but to a decimal or simple fraction of 
a gram of the substance. For example, a solution of hydro- 
chloric acid would be made with each cubic centimeter equiva- 
lent to 0.0100 gm, 0.0010 gm, 0.0050 gm, etc., of silver. 
This results in a very much simplified calculation and still more 
(time is saved if the weight of sample used bears a definite and 
■simple relation to the equivalence of the standard. ' • 

Such solutions as these are frequently made for technical 
iwork in industrial laboratories, where large quantities of standard 
Isolutions are often required for the titration of a single con- 
stituent of a large number of samples. Mention may be made of 
the use of potassium permanganate or potassium dichromate 
Isolutions for the titration of iron in ores, sodium thiosulphate 
solutions for the determination of the available chlorine in bleach- 
ing powder, potassium ferrocyanide solutions for the determina- 
tion of zinc and hydrochloric acid solutions for the determina- 
tion of hardness of water. 

The method of calculation of the necessary concentration of 
a solution to be made in the decimal system is the reverse of 
the method for calculating the equivalence of a solution of given 

Example: What must be the concentration of a solution of 
potassium hydroxide in order that each cubic centimeter shall 
be equivalent to 0.001 gm of sulphuric acid? The reaction 
involved is: 

2KOH+H 2 S0 4 -+K 2 S0 4 +2H 2 0. 

The equivalent weight of potassium hydroxide is 56.108 and that 
of sulphuric acid is 49.02. Each cubic centimeter must con- 
tain ' n9 X 0.001 gm of potassium hydroxide. This is 
0.00114 gm. 


32. Calculate the concentration ( ) of standard solutions of hydro- 
chloric acid such that 1 cc=o=the following weights of other substances: 
0.002 gm of silver; 0.005 gm of silver chloride; 0.010 gm of potassium 
hydroxide; 0.005 gm of sodium hydroxide; 0.002 gm of sodium; 0.002 
gm of ammonia. 


33. Calculate the concentration of a nitric acid solution such that 1 cc 
the following weights of substances: 0.040 gm of potassium hydroxide 
0.005 gm of calcium carbonate; 0.001 gm of nitrogen as ammonia. 

34. What is the concentration of a potassium hydroxide solution o 
which 1 cc=c= 0.010 gm of potassium acid tartrate? 

Choice of System. — Summarizing, it has been shown that 
volumetric analysis may be carried out by the use of standai 
solutions made in "no system," in the "normal system" or in 
the "decimal system," and that for any of these systems a defi 
nite, precalculated weight of sample may be taken so that the 
burette reading in cubic centimeters will indicate directly the 
percent or simple fractions of percent of the constituent being 
determined. Which of these systems shall be selected in prac- 
tical work will be determined by the circumstances. If but a 
few titrations are to be made with a given standard solution the 
time saved in simplified calculations will not justify the expendi- 
ture of time required for adjusting the concentration to the nor- 
mal or decimal system. If many titrations are to be made, one 
of the latter two systems will be used. The normal system is 
most useful for standard acids and bases because their application 
is more general and a solution so made will give simplified cal 
culations for the titration of many other substances. There are 
many standard solutions which are not to be used so generally but 
which are made for the titration of but one substance. In 
such instances the decimal system will always be used. 

Temperature Correction for Standard Solutions. — It is often 
difficult to control the temperature of the laboratory within 
close limits and errors may thereby be introduced into volumetric 
determinations, due to changes in the density of standard solu 
tions and in the capacity of measuring instruments when the 
temperature varies from 20°. The following table, adapted 
from tables published by the Bureau of Standards, 1 indicates 
the corrections for water and for two concentrations of most of 
the common standard solutions of acids and bases. Approxi- 
mately these corrections will apply also to most other solutions 
of similar concentrations. 

1 Bur. Stand., Circ. 19, table 33. 




cc per observed liter to give 


volume at 20° 


Y^ Solutions 

"2 Solutions 




+ 1.0 





































It will be seen that a variation of 2°, either way, from 20° 
nvolves an error of 0.04 percent in measurements of tenth- 
normal solutions, or 0.04 to 0.05 percent for half-normal solu- 
tions. This may be ignored for much of the routine work of the 
industrial laboratory but for more exact work the corrections 
should be applied. 

The use of standard acids and bases provides a means for the 
quantitative determination of practically any acid or base and 
of many salts. This is an extremely useful department of work, 
in view of the fact that no gravimetric method will serve to de- 
termine the essential constituent of acids and bases, the ionizable 
i hydrogen and hydroxyl. For example, from potassium hy- 
jdroxide potassium may be determined as chlorplatinate or 
I perchlorate, but this gives no information concerning the per- 
! cent of potassium hydroxide since potassium from any salt 
i present is also precipitated and weighed. By using the proper 
| indicator salts of strong bases with weak acids or of strong acids 
1 with weak bases may also be titrated. Thus sodium carbonate 
may be titrated by standard hydrochloric or sulphuric acid if 
methyl orange is used as indicator. 

Adjustment to Exact Concentration. — In most of the exercises 

< of the following pages the student is directed to adjust his 

standard solutions to the exact stated concentrations. If a 


solution is desired to be tenth-normal, or perhaps of such cor 
centration that 1 cc is equivalent to 0.1 mg of phosphorus orO 
mg of tin, etc., it is first made to this approximate concentratioi 
It is then standardized and finally adjusted to the require( 
strength, with or without an additional standardization to con 
firm the accuracy of the dilution. The only exceptions to th 
rule are in cases of solutions of unstable substances which chang 
on standing. 

Of course this process of adjustment is, in some cases, a some- 
what tedious procedure and there is a too common custom of 
omitting the final adjustment, using a correction factor in the 
calculations of analyses made by means of this solution, this 
factor having been found by the first standardization. Thus, 

if the first standardization showed a solution to be 1.0359-^, 


all calculations of titrations would be made as though the solu- 
tion were fifth-normal, except that the factor 1.0359 would enter. 
This figure is known as a " normality factor." 

To the inexperienced analyst it may seem that this is, after all, 
the best method of dealing with the question — -that the use of 
such a factor requires much less work than is involved in the 
process of exact dilution and restandardization. But hen 
again the question must be resolved with respect to the way i] 
which the solution is to be used, as was done in the matter oi 
choice of system. The work involved in adjusting the standan 
solution is balanced against the labor that is saved by simplifie( 
calculations. But the latter quantity is, of course, to be mul- 
tiplied by the number of determinations that will be made before 
the solution deteriorates and requires restandardization, or before 
the supply is exhausted. Obviously, this means that if any 
considerable use is to be made of a given standard solution it 
should be adjusted to the exact desired concentration. It 
may also be remarked that if adjustment is not to be made there 
is no logic in trying to work to either the normal or the deciim 
system. The method described on page 197 is far simpler ii 
this case. 


In a broad sense the word "indicator" applies to all substances 
which, by undergoing any visible change, indicate the end point 
of reactions. When the indicators are inorganic the reactions 
are usually definite and well understood. The indicators used 
in acidimetry and alkalimetry are organic and the direct cause 
of color change is, even now, not thoroughly understood. Many 
of the organic dyes show, in acids, a color different from that in 
bases. The color change is generally reversible an indefinite 
number of times. The molecular structure of the dye is often 
very complex and it is not easy to follow the changes in structure. 

Simple Ionization Theory. — Most or all of the indicators of this 
class are known to possess, in certain conditions at least, the prop- 
erties of acids or bases. The acid or basic nature is usually 
weakly emphasized. From this Ostwald deduced a theory as to 
the cause of color change. 1 According to this theory these dyes 
are, when uncombined with other acids or bases, weak electrolytes 
and largely in the molecular state. If a base is added to a weakly 
acid indicator, the salt is formed and this is highly ionized, ac- 
cording to the general rule. The molecule possesses one color (or 
is colorless), while the anion shows a different color. The result 
of the addition of a base is therefore a color change. If another 
acid is now added to the ionized salt the weak acid is reformed, 
the molecule reappears and the color change is reversed. The 
added acid has a further effect upon the indicator acid in suppress- 
ing the already small ionization. Similar reasoning would apply 
to basic indicators. Phenolphthalein is in the presence of acids, 
a derivative of phthalic anhydride and phenol having the 
following constitution: 


C 6 H 4 <^ ^>0 

C = (C 6 H 4 OH) 2 . 

1 Scientific Foundations of Analytical Chemistry, 118. 



According to the theory of Ostwald this is a very weak acic 
giving a small concentration of ions thus : 

HPh ±=f H + Ph, 

the symbol Ph representing the negative radical of the compounc 
The ionization constant is very small and equilibrium occur 
with an inappreciable concentration of the anion. Upon the 
addition of a base the ionized acid is neutralized, equilibrium is 
disturbed and the ionized salt is produced, hence the color of the 
anion (red) appears. Methyl orange is known to have, under 
certain conditions, the structure 

(CH3)2N-C 6 H4-N = N-C 6 H 4 S03H. 

This is an acid, the red molecular form predominating in acid 
solutions, and the yellow anion appearing in basic solutions. 

(CH 3 )2N-C 6 H4-N-=N-C 6 H 4 S03H<=± 

(CH 3 ) 2 N-C 6 H4-N = N-C 6 H4S03 + H. 

Theory of Chromophors. — This explanation is not sufficient 
in itself for several reasons. The silver salt of phenolphthalein : 
is intensely purple, even when dry, and the dry salt cannot b< 
highly ionized. Ethers of tetrabromphenolphthalein have beei 
prepared; 1 these are non-ionizable but colored. The monoetlr 
ether is 


C 6 H 4 <^ 

X^-C 6 H 2 Br 2 OH 
^C 6 H 2 Br 2 = 

Litmus is known as both blue and red in the dry state, when it 
must be chiefly molecular, no matter what the color may be. 
Also the studies of recent years upon the constitution of organic 
dyes have shown that in many cases a change of molecular 
structure takes place upon the addition of an acid or a base 

Phenolphthalein is known to have the structure shown abov 
but in basic solutions there is a salt of. a carboxyl acid which is 
quinone derivative. The phenol derivative is then in equilibriu 

1 Nietzki and Burckhardt: Ber., 30, 175 (1897). 


with the quinone derivative and this equilibrium is. disturbed 
in one direction or the other by the addition of an acid or a base. 


C 6 H 4 <^ ^>0 <=> C 6 H 4 <^ +H 

C = (C 6 H 4 OH) 2 C = C 6 H 4 = 

\C 6 H 4 OH 
If an acid is added the first (colorless) molecular form is produced 
because suppression of ionization results from the increase in 
concentration of hydrogen ions. If a base is added the ionized 
form is neutralized, forming ionized salt and water, and thus 
the new structure predominates. When methyl orange changes 
from the sulphonic acid to one of its salts a change of structure 
also takes place. The structure peculiar to the non-ionized body 
(present when an acid is added) is not that of an azo compound 
but one containing the quinone ring. There is then equilibrium 
between the two forms: 
(CH3)2N-C G H 4 -N = N-C G H 4 S0 3 H<=± 

Yellow, predominates in basic solution. 

(CH 3 ) 2 N = C 6 H 4 = N-NH-C 6 H 4 S0, 

! . l 

Red, predominates in acid solution. 

The acid, by suppressing the ionization of the first form, causes 
the second, a lactonic form, to predominate. A base, by forming 
salt and water from the sulphonic acid, causes the first form to 
predominate. With the acid the quinone ring gives a red color. 
With the base the azo group gives a yellow color. The quinone 
ring, =C 6 H 4 = , is one of a class of groups known as "chromo- 
phors ,; because, wherever they appear in any compound they give 
rise to color. Other well characterized chromophors are the azo 
group — N = N — , the nitro group, — NO2, and the dicarbonyl 
group,— CO— CO— .! 

Classification of Organic Indicators. — Ostwald classified the 
indicators according to their supposed dissociation constants into 
three groups : 

(a) Very weak bases and relatively strong acids. 

(b) Moderately strong acids and bases. 

(c) Very weak acids and relatively strong bases. 

1 Vide Hantzsch: Ber., 32, 575 (1899), and Stieglitz: J. Am. Chem. Soc, 
25, 1112 (1903). 



Since he explained the color change upon the basis o 
formation it would necessarily follow that the relative sensitive 
ness would vary in the three classes. In class (a) the indicato: 
would be highly sensitive to bases but not easily affected by 
acids, except by very strong ones. The indicators of class (b) 
would be moderately sensitive to both acids and bases, while in 
class (c) they would be highly sensitive to acids and to none but 
strong bases. While some of these indicators are here called 
" relatively strong" acids or bases, it must be remembered that, 
compared with the strongest electrolytes, all are weakly ionized 
and all lie in the class of weak electrolytes. 

While the theory of color change by ionization must be 
regarded as based upon incorrect assumptions, the above classi- 
fication is still a convenient one since the same relative sensi- 
tiveness would follow from the application of any of our theories. 
Phenolphthalein may be taken as an illustration. According 
to the simple ionization theory there is to be considered merely 
the following system in equilibrium : 


where Ph is understood to mean the negative radical 

C 6 H 4 (CO) 2 C 6 H 4 OH.C 6 H40. 

Phenolphthalein falls in the class of very weak acids and it is 
consequently chiefly molecular unless a strong base be present, 
a weak base forming an easily hydrolyzed salt. Even weak 
acids can decompose the salt, therefore phenolphthalein will be 
easily affected by acids but will not be highly sensitive to bases. 
According to the view that color is due to the existence of 
chromophors the equation 


C 6 H 4 < / ^0 ?t C 6 H 4 / +H 

C = (C 6 H 4 OH) 2 C = C 6 H 4 = 

\C 6 H 4 OH 

is an expression of equilibrium between the non-ionized colorless 
form and the ionized form containing a chromophor. Here 
again a strong base will be required to produce the ion containing 



the chromophor, while a weak acid will reform the colorless 
molecule and for the same reasons that are given above. 

The classification according to ionization loses much of its 
significance when it is remembered that each indicator exists in 
;at least two forms. On this account we shall rather lay stress 
upon the sensitiveness of the indicator toward acids and bases 
and shall so classify a few of the more commonly used indicators. 


Class I 

Highly sensitive to 

acids, less sensitive to 


Class II 

Moderately sensitive to 
acids and bases 

Class III 
Highly sensitive to 
ess sensitive to 

Rosolic acid 




Methyl orange. 
Ethyl orange. 
Methyl red. 

This classification shows that an indicator which is highly 
sensitive to acids is, in a corresponding degree, weakly sensitive 
to bases. For this reason the only generally serviceable standard 
acid or base is a highly ionized one. Such a standard may be 
used for the titration of either weak or strong electrolytes, the 
selection of indicator being made with reference to the substance 
titrated rather than to the standard. 

Many of the indicators undergo a color change upon the addi- 
tion of certain salts, as well as of acids or bases. In all such cases 
the salt is one derived from an acid and a base of unequal degree 
of ionization. The result of the partial hydrolysis of such a salt 
is the production of an excess of ions of either hydrogen or 
hydroxyl, according to whether the acid or the base is the more 
strongly ionized. An indicator of the proper sensibility will be 
affected exactly as though the solution were that of an acid or a 
base. Thus sodium carbonate in solution yields, by partial 
hydrolysis, sodium hydroxide and bicarbonate. The difference 
in degree of ionization of these two electrolytes is so great that all 
indicators show the color that is ordinarily exhibited in basic 
solutions. On the other hand the hydrolysis of ferric chloride 
yields a strong acid and a weak base and here again the difference 
in ionization is sufficiently large to give an "acid reaction" with 


most indicators. The exact point at which the indicator will 
change color depends upon the relative strength of acid and base 
and also upon the sensibility of the indicator itself. If a solution 
of sodium carbonate containing methyl orange is titrated by 
hydrochloric acid the color change does not occur until the 
sodium carbonate is completely decomposed, according to the 
following equation : 

Na 2 C0 3 +2HC1^2NaCl+H 2 C0 3 . 

This is because methyl orange is practically insensible to such a 
weakly ionized acid as carbonic acid and a slight excess of hy- 
drochloric acid is necessary in order to affect the indicator. If 
phenolphthalein, an indicator of high sensibility to acids, is used 
instead of methyl orange the color change occurs when one-half 
of the reaction represented above is completed: 

Na 2 C0s+HCl->NaHC0s+NaCL 

In other words sodium bicarbonate, yielding upon hydrolysis 
two equivalents of carbonic acid for one of sodium hydroxide, 
is a neutral body to an indicator that is easily affected by even 
weak acids and only with difficulty by bases. Orthophosphoric 
acid may be taken as a final example. The molecule of this sub- 
stance contains three atoms of ionizable hydrogen. Two of these 
atoms are ionized to but a small extent. If the acid is titrated 
by a solution of a strong base the point at which the color change 
occurs (" end-point") will depend upon the indicator used. If 
the indicator is litmus the color changes from red to blue gradually 
instead of suddenly and this change comes after one-third and 
before two-thirds of the hydrogen is neutralized. In the pres- 
ence of methyl orange the 'color changes at the completion of 
the reaction: 

NaOH + H 3 P0 4 -*NaH 2 P0 4 + H 2 0. 

In the presence of phenolphthalein the end point is somewhat in- 
definite but occurs at the neutralization of two-thirds of the acid : 

2NaOH+H 3 P0 4 ->Na 2 HP0 4 +2H 2 0. 

Description of Indicators 

Following is a brief discussion of the preparation and properties 
of the indicators named in the table on page 211. 


Phenolphthalein. — The chemical nature and changes of phenol- 
phthalein have already been discussed. The compound is pre- 
; pared by heating together 5 parts of phthalic anhydride, 10 
parts of phenol and 4 parts of concentrated sulphuric acid for 
several hours at a temperature between 120° and 130°. The 
le mass is then boiled with water, filtered and the residue dissolved 
in dilute sodium hydroxide solution. The solution is filtered and 
neutralized with hydrochloric acid. Phenolphthalein precipi- 
tates and is purified by recrystallization from alcohol. The 
pure substance is a yellowish-white crystalline powder, practically 
insoluble in water but soluble in alcohol. For use in volumetric 
analysis a solution of 5 gm in 1000 cc of 50 percent alcohol is used. 

Rosolic Acid (Corallin). — Commercial rosolic acid is a mix- 
ture of aurine, C19H14O3, oxyaurine, Ci 9 Hi 6 6 , methylaurine, 
C20HK3O3 and pseudorosolic acid, C20H16O4. Each of these sub- 
stances contains the quinone ring. It is prepared by heating 
together phenol, sulphuric acid and oxalic acid. The changes 
occurring as acids or bases are added are not thoroughly under- 
stood. The solution as used in the laboratory is a 1 percent 
solution in 60 percent alcohol. The indicator is red with bases 
and yellow with acids. It is highly sensitive to acids and can be 
used for the titration of weak acids. 

Litmus. — Litmus is obtained by the action of ammonium 
hydroxide and potassium hydroxide upon certain species of 
plants, followed by fermentation. The essential constituent of 
the indicator is azolitmin, C7H7NO4, of unknown constitution. 
It is colored red by acids and blue by bases. Its sensitiveness 
toward both acids and bases is only moderate and this fact makes 
it a very valuable indicator for general qualitative purposes but 
of little use for quantitative analysis. A 10 percent solution 
in water is used in the laboratory. 

p-Nitrophenol. — This indicator is prepared by the action of 

/N0 2 (4) 
nitric acid upon phenol. Its formula, C 6 H 4 <^ 

X OH (1), 
is indicated by its name. It is yellow with bases and colorless 
with acids. It is applicable to the same uses as is litmus. The 
indicator solution is a 0.02 percent solution in water. It should 
not be kept in a closely stoppered bottle. 


Methyl Orange. — The constitution and chemical propertie 
have already been discussed. The substance as obtained in 
commerce is usually the sodium salt. This is a yellow powder 
soluble in water. A solution containing 0.5 gm in 1000 cc 
used. In presence of ammonium salts the indicator is not very 
sensitive. Also the color is destroyed by iron or aluminium 

Ethyl Orange. — This substance is analogous in constitution t 
methyl orange, it being the diethyl ester of amidoazobenzen 
sulphonic acid. Its properties and uses are the same as those o: 
methyl orange. 

Cochineal. — Cochineal is the dried female insect Coccus cacti 
Linne. The essential coloring matter is carminic acid, C11H12O7 
whose constitution is not definitely established. The solution 
for use as an indicator is made by digesting 3 gm of the dried, 
unpowdered insects with 250 cc of 75 percent alcohol until 
the coloring matter is extracted, then decanting. The bottles 
containing the solution should not be closely stoppered. The 
indicator is violet with bases and red with acids. Its sensitive- 
ness is not diminished by ammonium salts. 

Lacmoid. — Lacmoid is prepared by heating together resorcin, 
sodium nitrite and water. It is a deep blue dye of unknown 
constitution; the molecular composition is probably represented 1 
by the formula C12H9NO4. It is soluble in alcohol and less so 
in water. In solution it is colored blue by bases and red by 
acids. Its sensitiveness toward both acids and bases is moderate 
but it finds an application in the titration of carbonates in boiling 
solution, carbonic acid being decomposed as fast as it is formed. 
Litmus may be used in the same way but is not so sensitive in hot 
solutions. The indicator solution should contain 2 gm of purified 
lacmoid in 1000 cc of 50 percent alcohol. 

Erythrosine (lodeosine). — -This dye is tetraiodofluorescein, 
a synthetic derivative of fluorescein, and it has the constitution 
represented by the formula 


C 6 H 4 <f>0 C 6 HI 2 OH 


C 6 HI 2 OH. 


The substance is therefore a phthalein. It is almost insoluble 
in water but is soluble in hot alcohol and ether. A solution 
of 0.5 gm of the sodium salt in 1000 cc of recently boiled water 
lis used as indicator. It is colored red by bases and light yellow 
'by acids. It is highly sensitive to bases and is therefore appli- 
cable to the titration of the alkaloids. Its sensitiveness to acids 
is correspondingly small so that it may be used for the titration 
of carbonates of the alkali and alkaline earth metals without 
boiling to expel carbonic acid. On account of its limited solu- 
bility titrations are made by adding 10 to 20 cc of the solution 
in ether to the titrating solution, shaking after each addition 
of acid. 

Methyl Red. — This dye is p-dimethylaminoazobenzene-o- 
carboxylic acid: 

(CH 3 ) 2 N-C 6 H4-N = N-C 6 H 4 C0 2 H. 

The indicator solution is prepared by dissolving 1 gm of the solid 
in 100 cc of 95 percent alcohol. The solution is pale yellow in 
basic solutions and violet red with acids. It belongs to Class III 
and is especially good for the titration of ammonium hydroxide 
and the alkaloids, all being weak bases. It cannot be used 
if much carbonic acid is present, hence is useless for the titration 
of carbonates. 

All indicator solutions must be adjusted to exact neutrality before 


Thus far we have dealt with only the calculation oi the result < 
of analyses, assuming that the standard solution was ready f<>r 
use in the experiment. The determination of the concentration 
of the standard solution is called " standardization." The 
details of the experimental work will be treated later and will be 
mentioned here only so far as they may serve to illustrate the 
methods of calculation. 

Standardization may be accomplished by one or more of four 

Direct Weighing. — The active substance of the solution 
is accurately weighed and dissolved to make a definite volume 
of solution. This method is applicable to only those substances 
that can be obtained in a pure state or in a state of uniform an 
accurately known composition. Most of such substances ar 
crystallized salts or acids, or soluble gases. 

Weighing a Substance Produced by a Measured Volum 
of the Solution. — Sulphuric acid solution may be standardized 
by precipitating a measured volume by adding an excess of 
barium chloride. From the weight of barium sulphate found 
the weight of sulphuric acid may be calculated by the method 
given on page 196. Similarly hydrochloric acid solution may be i 
stanotardized by precipitating as silver chloride. 

Measuring the Volume of Solution Required to React with I 
a Known Weight of a Substance of Known Purity. — An acid l| 
may be allowed to react with a pure carbonate and the required 
volume noted. Sodium thiosulphate solution may likewis 
be titrated against a weighed quantity of iodine (or indirectly 
against a weighed quantity of arsenic trioxide. 

Titration Against Another Solution Which has Already Bee 
Standardized. — This is a very common expedient. 



Primary Standards. — -It will be noticed that in each of these 
cases there is some substance of known composition that is 
measured or weighed and the solution is somehow compared 
with this for standardization. This substance of known com- 
position is called the "primary standard," whether it be the 
substance dissolved in the solution, something produced by the 
solution or something reacting with the solution. 

The following examples will illustrate the methods of calcula- 
tion in each of the cases discussed. 

(1) The method of calculation for the first method of stand- 
ardization is self-evident. The normality is equal to the ratio 
of the number of grams dissolved in 1000 cc to the number of 
grams in 1000 cc of a normal solution. That is, 

grams per 1000 cc 

normality = ; — f—, r-rr- 

J equivalent weight 

(2) A solution of hydrochloric acid was standardized by 
precipitating the chlorine from 40 cc, as silver chloride. The 
weight of silver chloride found was 0.6327 gm. Required, the 
normality of the solution. 

1 cc acid solution— ' — gm silver chloride. 

1 cc normal acid solution— 0.1433 gm silver chloride. 

6327 6327 

Therefore normality = ' 4Q -^0.1433= 4o^o^U^ = - 1107N - 

To make the solution decinormal 1000 cc would be diluted to 
1107 cc. 

(3) A similar solution was standardized by titration of pure 
sodium carbonate in presence of methyl oraQge, the following 
reaction being completed: 

Na 2 C03+2HCl->2NaCl+H 2 C0 3 . 
It was found that 32.2 cc acid — 0.1638 gm of the primary 
standard, sodium . carbonate. Required the normality. 

I' 0.1638 
1 cc «acid— » 2 2 gm sodium carbonate and 

1 cc normal acid— 0.053 gm sodium carbonate. 
Therefore normality = g^x ~6~05Z = ° * 9598 N * 


(4) Another acid solution was standardized by titration against 
a measured volume of standard potassium hydroxide solution 
in presence of methyl orange according to the equation: 

HC1+K0H->KC1+H 2 0. 

1 cc of the primary standard contained 0.00468 gm of potas- 
sium hydroxide. It was found from the titration that 50 cc 
of potassium hydroxide solution— 43.5 cc of hydrochloric aci 

The weight of potassium hydroxide in 50 cc of solution = 
50X0.00468 gm. Since this weight was equivalent to 43.5 
cc of acid, the potassium hydroxide equivalent to 1 cc acid = 


gm. The normality of the hydrochloric acid solu- 

50X0.00468 ftAAe1kT 
tlOn = 43.5X0 .0561 = Q -° 95N - 

In case the primary standard is a solution already standardized 
in the normal system the normalities of the solutions are inversely 
as the respective volumes that are equivalent to each other. 

(5) 30.0 cc of yt^ sodium thiosulphate solution is found by 

titration to be equivalent to 29.8 cc of iodine solution. The 
normality of the latter is required. 

_ . . 30.0 N 1 ._ N 
ThlS1S 29:8 10 = 1007 10' 

If solutions are to be standardized in the decimal syste 
the calculations involve nothing more than finding the weight of 
the substance in terms of which the standardization is to be ex- 
pressed, equivalent to 1 cc of the solution which is being stand- 
ardized, always using as the starting point the known weight of 
the primary standard and following the method explained on 
page 196. In many cases the standardization is to be expressed 
in terms of the primary standard itself. For example, iodine 
solution is to be standardized against pure arsenic trioxide and 
expressed in terms of the same substance. Here we have the 
very simple method of weighing a suitable amount of arsenic 
trioxide, then dissolving and titrating by the iodine solution. 

Then . ... ,i. gmAsA 

1 cc iodine solution -" 

cc I-solution 


Other familiar examples of this class of methods are the 
tandardization of permanganate solutions against elementary 
ion or antimony for obtaining the weights of these elements 
quivalent to one cubic centimeter of the solution. 

The following example will serve to illustrate the first case 
just discussed : 

(6) A solution of potassium permanganate was standardized 
against sodium oxalate as follows: 2.5340 gm of sodium oxalate 
was dissolved and the solution was diluted to 1000 cc. 25 cc 
portions were titrated and gave an average of 24.25 cc of potas- 
sium permanganate solution equivalent to the oxalate solution 
used. Required the weight of iron and of calcium equivalent 
to 1 cc of permanganate solution. 

25 cc of the oxalate solution contained 0.025X2.5340 gm 

ji c • , i *• • ■ -i ** 0.025X2.5340 
and 1 cc of permanganate solution is equivalent to oTTok 

gm of sodium Oxalate. This weight, multiplied by the ratio 
of the equivalent weight of iron or of calcium to that of sodium 
oxalate, will give the weights of these substances that are 
equivalent to 1 cc of the standard solution. Then 

. .. 0.025X2.5340X55.88 nnnoio 
lcc solution- 24.25X67.005 = - 00218 S m Fe 


0.025X2.5340X 20.035 _____ 

24.25X67.005" " = - 00078 ^ Ca ' 


35. 30.0 cc of sulphuric acid solution yields 0.3625 gm of barium sul- 
phate. Calculate the normality and the dilution necessary to make the 
solution tenth-normal. 

36. 44.6 cc of silver nitrate solution yields 1.2870 gm of silver chloride. 
Calculate the normality and the dilution necessary to make the solution 

37. 39.7 cc of barium chloride solution yields 2.5346 gm of barium 
chromate. Calculate the normality and dilute to make the solution 

38. An acid solution is standardized by titrating pure sodium carbonate, 
using methyl orange as indicator. 45.1 ceo 2.4085 gm of sodium car- 
bonate. Calculate the normality and dilution necessary to make the solu- 
tion exactly normal. 


39. 50 cc of nitric acid solution is added to 0.4530 gm of pure calcium' 
carbonate. The unused excess of acid is titrated by a solution of a base 
and 6.15 cc of the latter is required. The base is then titrated against 
the acid in order to compare their concentrations and 21.3 cc of acid is 
found to be equivalent to 19.2 cc of base. Calculate the dilution necessary 
to make each solution fifth-normal. 

40. A sulphuric acid solution is standardized by titrating a sample ol 
potassium bicarbonate which contains 98.45 percent of the pure compound. 
35 cc acid =0=0.0391 gm of the sample. Calculate the normality and the 
dilution necessary to make the solution hundredth-normal. 

41. 32.9 cc of potassium hydroxide solution exactly neutralizes 0.3118 
gm of pure potassium acid tartrate, KHC4H4O6. Calculate the dilution 
necessary to make the solution twentieth-normal. 

42. 45.9 cc of sodium hydroxide solution was added to a solution con- 
taining 0.25 gm of crystallized oxalic acid, H2C2O4.2H2O, the indicator 
being phenolphthalein. The excess of base was titrated by 1.3 cc of acid. 
4.9 cc of the acid is found to be equivalent to 50 cc of base. Calculate 
the normality of each solution. 

43. A solution of barium hydroxide was standardized by titration against | 
succinic acid, H2C4H4O4, in presence of phenolphthalein. 20.9 cc of barium 
hydroxide solution neutralized a solution of 1.22 gm of succinic acid. Cal- 
culate the normality. 

44. 38.1 cc-of a sodium bicarbonate solution exactly neutralizes 36.7 
cc of a tenth-normal solution of hydrochloric acid. What is the normality 
of the first solution? 


Standard Acids. — It has already been shown that the most 
serviceable acids or bases for standard solutions are those that 
belong to the class of strong_electrolytes. For standard acids 
this practically limits one to the use of hydrochloric, sulphuric 
or nitric acid. On account of the ease with which it may be re- 
duced, nitric acid is not to be recommended for standard solu- 
tions of general application. Of the first two acids named, hy- 
drochloric acid is usually to be. preferred because it is monobasic 
and cannot form acid salts. 

Materials for Standardization. — For standardizing an acid 
use may be made of any method which involves a definite re- 
action with a pure substance or which produces a precipitate 
or gas that may be weighed or measured. This makes either 
volumetric or gravimetric methods available. Since most strong 
bases are hygroscopic and also combine readily with carbon di- 
oxide, purification and weighing are difficult and these substances 
are unsuitable for use as primary standards for standardizing 
acid solutions. We have left the alkali and alkaline earth car- 
bonates for this purpose. Calcium carbonate and sodium car- 
bonate are suitable and both substances may be obtained nearly 
pure. Precipitated calcium carbonate may be used but it is 
seldom free from other salts and an analysis must be made be- 
fore it is used for standardizing. The form most often used is 
the natural crystallized calcite known as Iceland spar, which is 
often nearly 100 percent calcium carbonate, although its purity 
should not be assumed without analysis. The chief disadvan- 
tage in the use of any form of calcium carbonate for standardizing 
acids lies in the fact that it is insoluble in water and an excess 
of acid must be used in order to hasten the process of solution. 
In such a case a direct titration cannot be made but the excess 
of acid must be determined by titration with a solution of a 




base, whose concentration as compared with the acid must b 

Several of these difficulties may be obviated by the use o 
sodium carbonate, which is soluble in water. As this substanc 
is obtained in commerce it contains variable quantities of wat 
and salts incident to the process of manufacture, such as sodiu 
sulphate and sodium chloride. This is largely due to the com- 
paratively large solubility of the salt and the consequent diffi- 
culty in purifying it by crystallization. At 20° a saturate 
solution contains sodium carbonate to the extent of 17.7 pe 
cent of its weight. The pure salt is best obtained from the 
bicarbonate by heating. Sodium bicarbonate dissolves to the 
extent of 8.8 percent of the weight of solution at 20°. It may 
therefore be more readily purified by fractional crystallization, 
especially if the purest obtainable commercial salt is used for 
the purpose. When heated sodium bicarbonate decomposes 
according to the equation 


2NaHC0 3 -+Na 2 C0 3 +H 2 0+C0 2 . 

The dissociation tension of carbon dioxide from sodium bicar 
bonate is as follows: 1 


Tension, mm of mercury 

55 c 

60 c 





The tension of carbon dioxide in the atmosphere is less than 1 
mm and, in consequence, sodium bicarbonate will slowly de- 
compose at temperatures below 55°. At 100° the decomposition 
is fairly rapid and the bicarbonate is completely changed into 
normal carbonate by heating for a short time at about 300°. 
At still higher temperatures the normal carbonate will yield 
some sodium oxide and carbon dioxide. 

Gravimetric Standardization. — Acids may also be standardized 
gravimetrically in case insoluble salts can be produced. Such 
method will apply to hydrochloric or -sulphuric acid but not 
nitric acid, since no insoluble nitrate is known. A point fr 
quently overlooked is that this method is really a standardiz 
tion with respect to the negative radical and is an acid standard 

l Lescoeur: Ann. chim. phys. [6] 28, 423 (1892). 


nation only in case no salts of the acid are present. Even in 
the purest commercial acids ammonium salts are often present 
knd the weighing of silver chloride or of barium sulphate will 
|:hus not give a basis for the correct calculation of acid strength. 
Standardization by Direct Weighing. — It is possible to weigh 
:he active substances directly in the exact amount necessary 
? or a solution of desired strength only in case the substance is 
available in pure form. This is not the case with most of the 
inorganic acids 1 and with comparatively few salts. It - then 
becomes necessary to calculate the approximate quantity re- 
quired to make a solution somewhat stronger than that desired, 
to standardize the solution so made and dilute exactly to the 
required strength. Such dilution may be accomplished with 
accuracy in case the water to be added may be measured in a 
burette, i.e., if less than about 10 cc is required. The final 
volume obtained by dilution is the sum of the initial volume 
and that of the added water only in case no expansion or contrac- 
tion occurs upon mixing. This is practically the case if the solu- 
tion is already dilute and the relative amount of water to be added 
is small. Dilution may then be accomplished by measuring a 
specified amount of solution in a flask, running in the calculated 
^mount of water from a burette and mixing directly in the gradu- 
ated flask. The "neck of the flask must be capable of easily 
holding the required water above the graduation. This fact, 
together with considerations of volume changes already men- 
tioned, places a practical limit upon the amount of dilution 
that may be accurately made in one process. If more than 10 
cc of water must be added to 1000 cc of solution it is necessary 
to dilute to nearly the required amount, restandardize and 
redilute to the exact value required. For example, if it is found 

that a solution is 1.3462 X tt: and it is necessary to dilute the 

solution to make it exactly tenth-normal, each 1000 cc must be 
diluted to a volume of 1346.2 cc. The addition of 346.2 cc 
of water could not be accurately made because there is no gradu- 
ated vessel capable of accurately measuring this quantity. 
Such an addition might also cause an appreciable volume change. 

1 Vide, Moody: J. Chem. Soc, 73, 658 (1898), and Acree: Am. Chem. J., 
36, 117 (1906). 


The correct procedure would be to add first to each 1000 c 
of solution about 335 cc of water, measured in a graduate 
cylinder, restandardize and then carefully complete the dilutio 
in the manner already explained. 

Exercise: Preparation of Pure Sodium Carbonate. — Use the best 
grade of sodium bicarbonate that is obtainable. Make qualitative teste 
for sulphates, chlorides and potassium, using an approximately 5 percent 
solution of the material for these tests. If impurities are found, purify 
by recrystallization as follows: Make a saturated solution of sodiu 
bicarbonate by warming the purest obtainable salt with distilled water 
Decant from any undissolved matter remaining and evaporate th 
solution in a large porcelain or platinum dish, at a temperature not 
higher than 40°, until crystals begin to separate. Allow to cool and 
stand, uncovered but in a place which is protected from dust, unti 
about 25 gm of crystals have formed, pour off the solution, wash th 
crystals once with cold water and press between filter paper. Dry at 
100°, powder and heat in a platinum crucible at a temperature betwee 
270° and 300° until the weight is constant. The product should be pur 
sodium carbonate but a test for sulphates and chlorides should be made 
Preserve in a closely stoppered weighing bottle. 

Exercise: Preparation of Tenth-normal Hydrochloric Acid. — Deter 
mine the specific gravity (if not already known) and from this the per 
cent of hydrochloric acid in the concentrated solution found in th 
laboratory. From the data so obtained calculate the weight or volum 
necessary to make 2500 cc of tenth-normal solution. Measure 2 percen 
more than this amount into a 1000 cc graduated flask and fill to the mar 
with water. Empty into a bottle having a capacity of about 2500 cc 
and add 1500 cc more of water. Stopper and mix thoroughly by 
shaking. Since the solution has been warmed by the reaction between 
acid and water it should be allowed to stand until the temperature of the 
room is attained before standardizing. 

Calculate the weight of sodium carbonate necessary to make 250 c 
of a tenth-normal solution. Weigh this quantity of the prepared pur 
material on counterpoised glasses, then brush into a beaker. Dissolv 
the weighed carbonate in distilled water and carefully rinse into a 250 cc 
graduated flask. Fill to the mark and mix thoroughly. Imperfeci 
mixing is often found to be the source of discrepancies in titrations with 
the acid solution. The solution will not remain constant in concentra- 
tion and should not be kept for more than one day. Fill a burette with 
the solution and another with the acid solution. Before making th 
titrations practice reading the color change as follows: Place 100 c 
of distilled water in a beaker and add a drop of methyl orange and 0.5 



|)f sodium carbonate solution. Drop in the acid solution until the last 
llrop changes the tint from yellow to pink. Now drop in sodium car- 
bonate solution until the yellow color reappears. Repeat the alternate 
Additions of carbonate and acid until the color change can be observed 
|vhen but one drop of either solution is added. It will aid in the next 
process if this solution is preserved and another prepared, the two 
[showing the two colors of methyl orange. These may be set aside for 
Comparison. Measure 40 cc of the carbonate solution into another 
beaker or Erlenmeyer flask, dilute to about 100 cc and titrate with the 
jicid solution in presence of a drop of methyl orange. Calculate the 
normality of the solution, also the volume of water to be added to each 
J1000 cc to make exactly tenth-normal. If water to be added is more 
than 10 cc add nearly the required amount to each liter of the acid, 
mix, and restandardize. If the quantity to be added is less than 10 
fcc the acid is diluted as follows: Fill a dry 1000 cc graduated flask 
lo the mark with the acid solution. This flask should be capable of 
lolding the required amount of water above the mark. From a burette 
tdd the calculated quantity of water directly to the solution in the 
lask and mix thoroughly. Pour into a dry bottle and make a second 
iter of diluted solution in the same manner, having first rinsed and 
pied the graduated flask. Check the accuracy of the dilution by 

Record upon the label of the bottle your name, the name of the 
Standard solution, its normality and the date of standardization, thus: 

Fig. 65. — Form of label for standard solution. 

Soda Ash 

The commercial grade of sodium carbonate known as "soda 
ish" contains, in addition to sodium carbonate, considerable 
water, some potassium carbonate and varying quantities of other 
sodium and potassium salts, such as sulphates and chlorides. A 
complete analysis would include the determination of all radicals 




but on account of the fact that soda ash is used in many indus- 
tries because of its basic properties its valuation generally in- 
cludes a determination of basicity and of water with other 
impurities. The first determination may be made by direct 
titration by a standard acid in the presence of methyl orange, 
while the last may be determined directly or taken by difference, 
in which case the percent of water includes all other impurities. 
In view of the fact that the basicity of the alkali carbonates 
toward methyl orange is due to the hydrolysis of a salt of a strong 
base and a weak acid, it is obvious that any salt derived similar! 
will likewise be basic and that therefore the titration by acid 
really a method for determining the radical of salts of weak aci 
and is not a basis for the calculation of any particular salt, such 
as sodium carbonate. Thus, for example, if a mixture of sodium 
carbonate, potassium carbonate, sodium bicarbonate, sodium 
silicate, and sodium borate were being titrated these salts would 
all be decomposed by the standard acid before an end point witl 
methyl orange would be reached. In the absence of a mon 
extended analysis it would be impossible to calculate the percenl 
of any one of these radicals or compounds. Since all are basi( 
in solution and since all would serve for most purposes whe 
sodium carbonate is used industrially it is customary to report t 
percent, arbitrarily calculated, of either sodium carbonate 
sodium oxide (regarded as being combined), assuming that 
basicity of soda ash is due to sodium carbonate. It should a 
be noted that sodium carbonate could contain either al 
hydroxides or bicarbonates but not both in the same samp 
Ordinarily no attention is given to either. 

On account of the lack of uniformity of most commercial so 
ash it is necessary to select a rather large sample, dissolving i 
water and measuring an aliquot part for the titration. Th 
directions for sampling on pages 9 to 14 should be carefull; 
followed but exposure to air should not be unduly prolonged. 

Determination. — Fill a 20 cc weighing bottle with soda ash, prope 
sampled. Assuming that the sample is pure sodium carbonate, calc 
late the approximate weight that should be contained in 500 cc 
solution so that 25 cc shall require about 35 cc of tenth-normal ac 
for its titration. Weigh this quantity of soda ash into a gradual 
500 cc flask, dissolve and dilute to the mark with water. Mix the 


oughly by inverting the flask and shaking. Measure out 25 cc by means 
of a calibrated pipette, allowing this portion to run into 200 cc beakers 
or Erlenmeyer flasks. Add just enough methyl orange to tint the solu- 
tion and then titrate with the tenth-normal acid. Make at least two 
more titrations and calculate the percent of sodium carbonate, also of 
sodium oxide, assuming that the basicity of the substance is entirely 
due to the sodium oxide combined as carbonate. 

"Pearl ash" (crude potassium carbonate) may be evaluated 
in a similar manner. 

Mixtures of Carbonates and Bases 

The use of the two indicators, methyl orange and phenol- 
phthalein, provides a means for the determination of carbonates 
and bicarbonates when in mixture, and also of carbonates and 
soluble bases. Bases and bicarbonates (acjd salts) cannot occur 
in the same mixture. If phenolphthalein is added to a solution 
containing sodium -carbonate and sodium hydroxide and the 
solution is titrated by a standard acid, the end point is reached 
when the sodium hydroxide is neutralized and the sodium car- 
bonate is changed to bicarbonate: 

NaOH+HCl-+NaCl+H 2 0, (1) 

Na 2 C03+HCl->NaHCOs+NaCL (2) 

Since the solution has now become colorless, methyl orange may 
jbe added and the titration continued until the red tint appears, 
'when the sodium bicarbonate has been completely decomposed: 

NaHC0 3 +HCl-*NaCl+H 2 C0 3 . (3) 

(Represent by A the cubic centimeters of acid used in completing 
Ithe first titration and by B that used in the second. 

BX normality X0. 106 = gm of sodium carbonate, 
(A— B) X normality X 0.040 = gm of sodium hydroxide. 

'It must be noted that here the equivalent weight of sodium carbcn- 
!ate is 106 instead of 53, since by equation (3) lHCl^lNaHC0 3 
and by equation (2) lNaHC0 3 =o=lNa 2 C0 3 . 

The quantitative conversion of sodium carbonate into sodium 
bicarbonate by means of an acid can take place only when care is 


taken to prevent the escape of carbon dioxide. At the point 
where the acid enters the solution there is at first complete con- 
version of a part of the carbonate into the normal sodium salt 
of the added acid. If carbon dioxide escapes from the solution 
at this point more acid will be required to produce the first 
end point than would otherwise be the case. The escape of gas 
may be prevented by keeping the solution at a temperature near 
0° and by adding the standard acid very slowly and while stirring 
vigorously. 1 It is very difficult to avoid the escape of carbonic 
acid unless the relative amount of carbonate is small. On this 
account the method is best suited to bases, in which carbonate 
occurs as an impurity, rather than to materials in which carbonate 
is an essential constituent. 

Determination of Hydroxide and Carbonate in Commercial "Caustic 
Soda" or "Caustic Potash." — Arbitrarily assuming that the sample is 
pure sodium or potassium hydroxide, calculate the approximate quan- 
tity necessary to dissolve and dilute to 1000 cc so that 50 cc shall requir 
about 40 cc of tenth-normal acid. From a stoppered weighing bott 
weigh this amount of well-mixed sample into a 1000 cc graduated flas! 
Dissolve in 500 cc of water and cool to 20°. Dilute to the mark, mi 
measure out 50 cc by means of a pipette, add a drop of phenolphthalei 
cool to 0° in ice water and titrate to the disappearance of the pink colo 
Add a drop of methyl orange and continue the titration to the nex 
color change. Calculate the percent of sodium carbonate and of sodiu 

Mixtures of Carbonates and Bicarbonates 

If a mixture of a carbonate and bicarbonate is to be invest 
gated the procedure is the same as in the preceding exercise. In 
this case the phenolphthalein changes color at the completic 
of the reaction 

Na 2 C0 3 +HCl->NaHC0 3 +NaCl. 

When the color change of methyl orange occurs the sodium 
carbonate so produced, as well as that originally present, 
been decomposed. If A is the acid used in the first titration ai 
B that used in the second titration 

A X normality X0.106 = gm of sodium carbonate, and 
(B — A) X normality X 0.084 = gm of sodium bicarbonate. 
1 Kuster: Z. anorg. Chem., 13, 127 (1897). 



These calculations are based upon the same arbitrary assumption 
regarding the presence of other salts as is noted in the discussion 
of the valuation of soda ash. 

Determination of Sodium Carbonate and Bicarbonate in a Mixture. — 

Proceed as in the preceding exercise, except that the calculation is to be 
made as above indicated. A thermometer must be placed in the solu- 
tion and the temperature lowered to 0°. The standard solution is added 
very slowly. 

Hardness and Alkalinity of Water 

The difference between "hard" and "soft" water lies in the 
fact that the former contains various inorganic compounds 
which form insoluble soaps when used with soap for cleansing. 
The salts of calcium and magnesium are chiefly responsible for 
this action. When ordinary soap is dissolved in such a water 
there is at once formed a precipitate of calcium and magnesium 
salts of the fatty acids. 

Temporary Hardness. — Bicarbonates of the metals named 
above are quite common in ground waters. When the water is 
boiled the excess of carbonic acid is expelled and normal car- 
bonates are formed: 

Ca(HC0 3 ) 2 -^CaC0 3 +H 2 0+C0 2 ; 
Mg(HC0 3 )2-+MgC03+H 2 0+C0 2 . 

Because of the very limited solubility of the normal carbonates 
a precipitate is formed and the hardness of the water is diminished 
to this extent. Such hardness as is removed in this manner by 
boiling is known as "temporary hardness" and the water so 
treated is partly "softened." 

Permanent Hardness is due to the slight amount of normal 
carbonates (about 0.013 gm per liter at 20°) that remains to 
saturate the water, and to non-decomposable, soluble salts of 
calcium, magnesium, iron or other metals that may form in- 
soluble soaps of fatty acids. The most common of such com- 
pounds are chlorides, sulphates and nitrates. 

Alkalinity. — The "alkalinity" of a natural water represents its 
content of carbonate, bicarbonate, borate, silicate, phosphate 
and hydroxide, or such of these as may be present. It is obvious 
that hydroxides and bicarbonates cannot exist in the same solu- 


tion. Also a water which contains sodium or potassium carbon- 
ate is not likely to contain calcium, magnesium or iron as bi- 
carbonates, because reactions like the following would occur: 

Ca(HC0 3 )2+Na 2 C03-^CaC03+2NaHC03. 

Methods for the Determination. — If none of the other salts 
mentioned as causing alkalinity were present a simple titration 
with standard acid in presence of methyl orange or erythrosine 
would serve as a determination of bicarbonate hardness. Also i 
one could neglect the fact that carbonates are not entirely precipi 
tated when formed by heating bicarbonates, such a determination 
of bicarbonate hardness could be reported as of temporary 
hardness. Because neither of these conditions is fulfilled the 
most accurate method for the determination of temporary hard- 
ness is by making the titration in presence of methyl orange 
before and after boiling. 

Permanent hardness equals non-carbonate hardness plus hard 
ness due to a saturated solution of calcium carbonate, mag 
nesium carbonate, etc., as already explained. If a water shoulc 
have no temporary hardness its non-carbonate hardness am 
permanent hardness would be identical. Non-carbonate hard 
ness is best determined by the use of "soda reagent," a standarc 
solution of equal weights of sodium hydroxide and sodium 
carbonate. A measured quantity of the water is boiled tc 
decompose bicarbonates and to this is added the standard solu 
tion in excess. Reactions like the following take place: 

CaS04+Na 2 C03->CaC03+Na 2 S04; 
MgCl 2 +Na 2 C0 3 -+MgC0 3 +2NaCl. 

After filtering to remove insoluble carbonates the solution is 
titrated by standard acid in presence of methyl orange and the 
hardness is calculated from the amount of soda reagent tha 
has been used by reactions like those represented above. 

The use of sodium hydroxide in the standard solution appear 
to diminish the solubility of the normal carbonates that arc 

Clark's method for the determination of total hardness is basec 
upon the use of a standard soap solution, added until a permanen 


lather is produced. It was formerly extensively used but it 
is inaccurate and is now little used. 

Expression of Results. — The small percent of dissolved salts 
usually found makes desirable a method of expressing results 
which is different from that used in other connections. Instead 
of percent it is customary to report parts per hundred thousand 
of water, parts per million or grains per gallon. The Imperial 
English gallon of water weighs 70,000 grains while the United 
States gallon weighs in air at 15.5° and 760 mm pressure, 58335 + 
grains. This gives at least four different systems which have 
been at various times and in various countries commonly used 
for the expression of the results of water analysis. Hardness 
itself has also been expressed in Clark's degrees (grains of calcium 
carbonate per Imperial gallon), German degrees (parts of calcium 
oxide per 100,000 parts of water) and French degrees (parts of 
calcium carbonate per 100,000 parts of water). This has re- 
sulted in much confusion but there is now a general tendency 
toward the practice of expression in parts per million or, more 
exactly, milligrams per liter, although in industrial operations 
the report is often made as grains per United States gallon. 
Upon the assumption that one liter of water, at the usual working 
temperature, weighs 1000 gm, every milligram of dissolved 
matter will represent one part per million of water. This as- 
sumption is not quite correct and "milligrams per liter" is a 
better expression than "parts per million." 

The following exercises may be performed at this point by 
students who will not carry out a more extensive analysis of 
water later. (See page 3-39.)^ | 

Determination of Alkalinity. — Prepare a fifth-normal solution of 
sulphuric acid, following the general directions given on page 
224 for tenth-normal hydrochloric acid and standardizing against pure 
sodium carbonate in presence of methyl orange. Dilute 100 cc of this 
fifth-normal solution to 1000 cc with recently boiled and cooled distilled 
water. The work of standardization and dilution must be done very 
carefully in order to avoid the necessity for restandardization of the last 

Potassium acid sulphate may be substituted for sulphuric acid in this 
determination. As this substance usually contains considerable water 
due allowance should be made when calculating the required weight for 
the fifth-normal solution. 


Measure 100 cc of the water sample in a volumetric flask and rinse 
into a porcelain dish or casserole with recently boiled water; or measure 
the sample directly into the porcelain dish by means of a pipette. Add 2 
drops of methyl orange indicator and titrate with the fiftieth-normal acid 
solution already prepared. Calculate the alkalinity in the conventional 
way as milligrams of calcium carbonate per liter. 

Determination of Temporary Hardness. — Determine the alkalinity 
of the original sample as already directed. Measure a second portion of 
200 cc of the sample and rinse into a 500 cc Erlenmeyer flask. Boil 
gently for 10 minutes, cover and cool to room temperature and immedi- 
ately rinse into the 200 cc volumetric flask. Dilute to the mark with 
recently boiled and cooled distilled water and mix thoroughly. It 
is important to have distilled water which is quite free from carbonic 
acid for this determination, as otherwise part of the precipitated carbon- 
ates will be redissolved and later titrated as alkalinity. On this account 
the wash bottle must not be used in the ordinary manner by blowing 
into it. 

Stopper the flask containing the boiled sample and allow to stan 
until the precipitate has settled, then remove 100 cc of the clear liqui 
by means of a pipette and redetermine the alkalinity. The differem 
(if any) between the two titrations represents the permanent hardness 

While temporary hardness is due to bicarbonates of calcium, mag 
nesium or iron (and sometimes other metals of the earth or alkaline 
earth groups) it is customary to calculate it in terms of calcium carboi 
ate, the most common of the products of boiling hard water. Th 
convenience of a fiftieth-normal acid for this use should be noted. Sim 
the equivalent weight of calcium carbonate is almost exactly 50, 1 cc o 
fiftieth-normal acid is equivalent to 1 mg of calcium carbonate. Repoi 
the temporary hardness in milligrams of calcium carbonate per liter 

Determination of Non-carbonate Hardness. — Prepare an approxi 
mately tenth-normal (to methyl orange) solution of soda reagent, using 
equal weights of sodium carbonate and sodium hydroxide, standardizing 
against tenth normal hydrochloric acid and using recently boiled wate 
for dilutions during the titration. 

Fill a dry 200 cc volumetric flask (or one that has been rinsed with tl 
sample) with the sample of water. Rinse this into a 500 cc Erlenmeye 
flask of resistance glass, using 200 cc of distilled water. Boil for U 
minutes to expel carbon dioxide then add exactly 25 cc of standarc 
soda reagent from a pipette. Boil for 10 minutes, rinse into the 200 
volumetric flask and dilute to the mark, using recently boiled and cook 
distilled water. Filter through a dry filter and discard the first 50 cc 
of filtrate. From the filtrate that is subsequently obtained measure 


50 cc portions by means of a pipette into Erlenmeyer flasks and titrate 
at once with fiftieth-normal sulphuric acid or potassium acid sulphate, 
using methyl orange or erythrosine. If erythrosine is used, add 1 cc 
of the indicator solution and 5 cc of neutral chloroform, titrating until 
the pink color just disappears from the chloroform when violently 

Calculate the non-carbonate hardness in the conventional manner by 
assuming the typical calcium sulphate as the hardness-diving agent 
and reporting milligrams of this compound per liter. In waters of the 
"alkali" type, containing carbonates of sodium or potassium, a negative 
value will be found for non-carbonate hardness, i.e., more standard 
acid will be required after the treatment above outlined than is equiva- 
lent to the soda reagent added. 

Standard Bases. — Standard solutions of bases are subject to 
change in basic concentration and must be frequently restand- 
ardized. This is because glass -is appreciably soluble in bases 
and the accumulation of alkali silicates in the solution gives an 
increase in basicity toward all indicators. Bases also absorb 
carbon dioxide when exposed to the air and this results in a 
decreased basicity toward indicators of the class of phenol- 
phthalein. For this reason it is desirable to avoid unnecessary 
contact with the air after standardization. 

For the preparation of standard basic solutions free from 
carbonates one may either use a substance whose carbonate is 
insoluble or one which contains little carbonate and whose 
carbonate may be precipitated by the addition of another sub- 
stance. For the first method barium hydroxide is generally 
used. This base always contains some barium carbonate but this 
remains undissolved when the solution is made. If, thereafter, 
carbon dioxide is absorbed by the solution, an equivalent amount 
!of barium carbonate is precipitated and the solution remains 
J free from carbonate, although it must be restandardized. Sodium 
(hydroxide or potassium hydroxide may be obtained nearly 
| free from carbonates by dissolving in alcohol, decanting or 
! filtering from undissolved carbonate and evaporating the 
;alcohol in an atmosphere that is free from carbon dioxide. 
■Bases so prepared may now be obtained from the manufacturers 
[and should be used for the preparation of standard solutions 
jwhenever possible. 



For the second method potassium hydroxide or sodium hy- 
droxide may be used and a slight excess of barium chloride added 
to the solution, barium carbonate being thereby precipitated. 
The material used should be the sticks that have been crystallized 
from alcohol. As sodium and potassium carbonates are only 

Fig. 66. — Burette with three-way stopcock, connected with reagent bottle. 

slightly soluble in alcohol the bases from this solution are nearb 
free from carbonates. In either case reabsorption of carbon 
dioxide is prevented by passing entering air through a soda lime 
tube, removing the solution through a siphon directly to the 
burette, as in Fig. 66. Covering the solution with a layer of 


toluene or of any oily substance is not to be recommended 
because this results in fouling the burette. 

It should be understood that these precautions are unneces- 
sary in most titrations. While it is true that alkali carbonates 
have not the same basicity toward some of the indicators as have 
the alkali hydroxides, a proper correction is made by standardiz- 
ing in presence of the indicator that is to be used in the determina- 
tions. This is a principle that must always be observed. One/ 
must not, for instance, use the standardization in presence of methyl 
orange as a basis for the calculation of determinations made in 
presence of phenolphthalein. 

Selection of Base for Standard Solutions. — The proper base 
for a standard solution will depend upon the nature of the titra- 
tion to be made. For reasons discussed on page 211 one will 
generally select a highly ionized substance and from this stand- 
point sodium hydroxide and potassium hydroxide are about 
equally good. On account of the relative cheapness of the former 
it should be given preference, wherever possible. However, in 
certain cases (see analysis of edible oils, beginning on page 345) 
a standard solution of a base is used for the saponification of 
oils. On account of the greater solubility of potassium soaps, 
potassium hydroxide is best for this purpose. 

Standardization. — The standardization of solutions of bases 
may be accomplished by indirect methods, titrating against a 
previously standardized acid solution, or by direct methods, 
titrating against a weighed substance of known purity. In 
general the first method is to be recommended because the availa- 
ble solid acids of uniform purity are limited to the organic 
acids. This makes necessary the use of phenolphthalein in the 
standardization, which limits the use of solutions so standardized 
to determinations that can be made by means of this indicator. 
Some solid substances that may be used as primary standards 
are oxalic acid, H 2 C 2 04.2H 2 0, benzoic acid, HC 7 H 5 2 and succinic 
acid, H 2 C 4 H 4 04. No direct gravimetric method can be used 
because there is no precipitating reagent for the hydroxyl radical 
and a determination of the metal would not give the basic strength 
because of the invariable presence of salts of the same metal. 

Exercise : Preparation of Tenth-normal Sodium Hydroxide Solution. 

— Select the best grade of sodium hydroxide obtainable, that purified 


by alcohol being preferable. Calculate the weight necessary for 2500 cc 
of tenth-normal solution and weigh out the quantity with 2 percent 
added to compensate for carbonates, water, and other impurities. Dis- 
solve in a 2500 cc bottle having a solid rubber stopper. Fill the bottle 
with recently boiled and cooled distilled water, mix thoroughly and 
allow to cool to the temperature of the room. Titrate portions of about 
30 cc each against the tenth-normal acid, using methyl orange as 
indicator. Continue these practice titrations until the results agree 

Dilute to make the solution exactly tenth-normal, using recently 
boiled and cooled water. Restandardize against the acid, using methyl 
orange, then use, in turn, phenolphthalein, methyl red, cochineal and 
lacmoid, as well as such other indicators as are available. The indi- 
cators of the first class (page 211) will indicate a weaker base than those 
of the third class, on account of the presence of small amounts of 
carbonates in the standard solution. Record the normality of the basic 
solution according to each indicator and use the proper figure for sub- 
sequent titrations, according to the indicator there used. 

Determination of the Concentration of the Laboratory Acids. — 
Use the "dilute" acids, either sulphuric, hydrochloric, nitric or acetic.-' 
The sample should not be weighed in the analytical balance and it is' 
better to determine the specific gravity and then take a measured volume 
Determine the specific gravity with an accurate hydrometer. From tl 
approximately known percent of acid in the solution calculate the dih 
tion required to make a solution approximately equivalent in concenti 
tion to the standard base. Make 500 cc of solution, measuring acci 
rately the volumes used. Titrate the finally diluted solution against tl 
standard base, using methyl orange for any of the acids above name 
except acetic acid. For this use phenolphthalein. Calculate the 
cent of acid, by weight, in the original solution. If tables are at han< 
compare the results of the experiment with the percent correspondii 
to the specific gravity found. 

Determination of the Purity of Citric Acid. — Weigh about 3 gm o; 
commercial citric acid, dissolve and dilute to 500 cc. Titrate portions 
of 30 cc to 40 cc by the standard base, using phenolphthalein. Calculate 
the percent of the tribasic acid H3C6H5O7.H2O assuming that no othe] 
acid is present. 


Vinegar contains from 3 to 6 percent of acetic acid, in ad 
tion to coloring matter, dissolved solids and sometimes unfer 
mented sugar or alcohol. Cider vinegar contains also fron 



0.08 to 0.16 percent of malic acid. Vinegar is sometimes adul- 
terated with other added acids, notably sulphuric acid. The 
complete analysis will include the determination of the sub- 
stances just mentioned and others that serve to characterize the 
vinegar with respect to its origin or quality. The determination 
'of total acidity alone is of value in determining the "strength" 
of the vinegar. This is practically due to acetic acid alone in 
pure vinegars other than those made from cider. In cider 
jvinegar the determination of malic acid is also of importance as 
[indicating its origin. 

1 Determination of Total Acidity of Vinegar. — Weigh a clean, stoppered 
jweighing bottle of at least 25 cc capacity. Add about 25 cc of vinegar, 
stopper and reweigh with an accuracy of 1 mg. Carefully transfer this 
jsample to a 100 cc volumetric flask, using a stirring rod. Rinse the 
jbottle and rod with recently boiled water and dilute the vinegar and 
[rinsings to the mark on the flask. Mix thoroughly. By means of a 
Ipipette measure two or three separate portions to Erlenmeyer flasks, 
liilute to 50 cc, add a drop of phenolphthalein and titrate with standard 
base. Calculate the total acidity as percent of acetic acid, using the 
normality of the base as determined in presence of phenolphthalein. 

Boric Acid 

Boric acid is one of the weakest of all inorganic acids, having 
i percentage ionization not far above that of hydrocyanic acid. 
It is therefore not possible to titrate it by standard bases in 
brdinary solution because no indicator will give a sudden color 
phange at any point in its neutralization. It can be determined 
>y titrating in presence of glycerine, phenolphthalein being used 
jis indicator. 1 Hydroxylated organic compounds form ester- 
like compounds with the various boric acids, the result being 
substances more strongly ionized than boric acid. Tartaric 
jicid forms such substances, with a resultant change of its optical 
•otatory power. With glycerine and orthoboric acid the com- 
)ound C 3 H 5 (OH).HB03 is probably formed because the substance 
•eacts as a monobasic acid. Boric acid combined as salts may be 
iletermined by first adding, in presence of methyl orange, enough 
iiydrochloric or sulphuric acid to completely decompose the 
i Thomson: J. Soc. Chem. Ind., 12, 432 (1893). 


borate. In case carbonates are present carbonic acid is also 
produced and this must be expelled by boiling because phenol- 
phthalein is to be used in the last titration. If the more common 
biborates are thus decomposed they yield orthoboric acid, methyl 
orange indicating the end-point at the completion of the reaction : 

Na 2 B40 7 +5H 2 0+2HCl->2NaCl+4H 3 B03. 

The addition of glycerine. produces the reaction: 

C3H 5 (OH)3+H3B03-^C3H 5 OH.HB0 3 +2H 2 0. 

The monobasic acid is then neutralized in the titration: 

C 3 H50H.HB03+KOH-^C3H 5 OH.KB03+H 2 0. 

Commercial borax is essentially crystallized sodium biborate, 
Na 2 B 4 O 7 .10H 2 O. It loses water rapidly and seldom corresponds 
to this formula. 

Determination. — Crush and mix a sample of a borate quickly. Weigh 
about 10 gm, dissolve and dilute to 500 cc. Titrate measured portions 
o f 20 cc with tenth-normal acid in presence of methyl orange, first diluting 
the solution to 50 cc. Now add to other diluted portions of the borate 
solution the exact amount of acid indicated by the titration but no indi- 
cator. Boil for a few minutes, then cool. Add a drop of phenolphthal- 
ein and 30 cc of glycerine. The glycerine must be neutral. Titrate 
with standard base to the appearance of a pink color. Add 10 cc more 
of glycerine and if the pink disappears add more base. Continue the 
addition of glycerine and standard base until a permanent end-point 
is attained. Calculate the percent of boron trioxide, B2O3, in the 
original borate. Reference to the formula lor the complex of boric acid 
and glycerine will show that the equivalent weight of boron trioxide 
must be one-half its molecular weight. 

Use of Two Standards. — There are cases where a direct titra- 
tion cannot be made conveniently and where two standard 
solutions may be used with advantage. If the titrated substance 
is nearly insoluble in water and dissolves only when the standard 
solution is added a direct titration is a very tedious operation. 
This is due to the fact that it would not be permissible to add more 
standard than is chemically equivalent to the substance with 
which the standard reacts, so that the active mass of the standard 
is, necessarily, always small, particularly toward the end of the 
titration, where it approaches zero. The velocity of solution 


would therefore become extremely low, this also approaching 
zero toward the end of the titration. 

For this reason it is much simpler to add an excess of the 
[standard solution, the total quantity used being measured 
laccurately. The active mass of the standard is then relatively 
large throughout the operation and solution of the sample pro- 
ceeds fairly rapidly. After solution is complete the unused 
excess of standard is titrated by means of a second standard 
'solution. This gives sufficient data for the calculation. The 
volume used of the second standard is calculated to the equivalent 
volume of the first. This figure is then subtracted from the total 
volume used of the first standard and the remainder is^the volume 
used by the titrated substance. 

A familiar example of the use of two standards is in the titra- 
tion of the acid-neutralizing power of limestone for agricultural 
purposes. The first standard is hydrochloric acid and the 
second is either sodium hydroxide or potassium hydroxide. 
A tenth-normal solution of acid is not active enough for this 
purpose and a fifth-normal solution should be prepared. The 
preparation and standardization of fifth-normal solutions will 
follow the same lines as that of tenth-normal solutions, with 
which the student is already familiar. 

"Alkalinity" of Limestone for Agricultural Purposes. — Have the 
sample finely powdered. Weigh exactly 0.5-gm samples on counter- 
poised glasses and brush into Erlenmeyer flasks. Moisten with water 
then pipette 75 cc of fifth-normal hydrochloric acid into each flask. 
After effervescence has become slow, connect with a reflux condenser 
and heat to boiling. Boil for 1 minute to expel carbon dioxide, then 
cool and rinse down the condenser, the stopper and the upper part 
of the flask. Add a drop of methyl red or phenolphthalein and titrate 
the unused excess of acid with fifth-normal base. (In standardizing 
the base against the acid, use must be made of the same indicator as 
is chosen for the determination.) 

Calculate the percent of calcium carbonate in the limestone. This 
will give a fictitious value for dolomitic limestones, which contain 
magnesium carbonate having an equivalent weight of 42, as com- 
pared with 50 for calcium carbonate. In such cases the calculated 
percent of calcium carbonate may be more than 100. 


Reactions of neutralization constitute a very important group 
in quantitative chemistry. Of no less importance is that group 
which is composed of reactions of oxidation and reduction. 
These do not involve the use of the organic indicator but of some 
inorganic substance, this being, in some cases, the standard 
reagent itself. While the quantitative relations between the 
reacting substances are here different from those of neutraliza- 
tion, the same principles will serve for the calculation of the 
results of the titration. The equivalent weights will be deter- 
mined by dividing the molecular weights by the hydrogen 
equivalents but the latter will be determined, not by the valence 
of the reacting parts but by the change of valence in the reaction, 
because oxidation always involves a change of valence of tl 
oxidizing and reducing agents. In the reaction: 

BaCl 2 +H 2 S0 4 ^BaS04+2HCl 

the hydrogen equivalent of each radical is plainly represented 
its valence, for it is in this measure that it may enter into tl 
exchange of double decomposition. In the reaction: 

2KMnO 4 +10FeSO4+8H 2 SO4->K 2 SO4+2MnSO4+ 
5Fe 2 (S0 4 ) 3 +8H 2 

the hydrogen equivalents of potassium permanganate and fei 
sulphate cannot be so represented because something that 
different from ordinary double decomposition has taken pi 
Manganese has left a negative radical and entered a positive 01 
and has thereby changed its valence and has been reduced, 
change of valence will represent its power as an oxidizing age 
Iron has taken on more of the radical with which it was alreac 
combined, has increased its valence and has been oxidized. 



lis not difficult to determine the change of valence of manganese 
if its state of oxidation in the two compounds is first determined. 
Apparent Valence. — For the purposes of this inspection the 
question of the actual valence need not be considered. There 
jhas been much difference of opinion regarding the structural 
formulas that should represent inorganic compounds and the 
real valence of an element in an oxide is indicated by the structural 
jformula of that oxide. For example, the dioxide of manganese 
might be represented by the 

formula Mn^ , in which the valence of manganese, is 4, or by 


Mn^ Jjin which the valence is 2. If the element is regarded as 

| always being in direct combination with the oxygen the latter 
| may be taken as a measure of the apparent valence and the loss 
of oxygen or its equivalent in other elements, as the result 
of a reaction, will be measured by the apparent change of valence. 
I This change will therefore be the hydrogen equivalent of either 
oxidizing or reducing agent. 

The apparent valence of an element which is in the negative 
radical of an oxyacid or of its salt is found by subtracting the 
total valence of the positive radical from the total valence of the 
other element of the negative radical and dividing the result by 
the number of atoms of the element in question. Thus the 
apparent valence of manganese in potassium permanganate is 
(4X2) —1 = 7; that of chromium in potassium dichr ornate is 
(7X2)-2 = 6 

When the element in question forms a simple radical its valence 
is easily seen from an inspection of the formula. The valence of 
chromium in chromium chloride is 3, that of manganese in man- 
ganous sulphate is 2, that of iron in ferric sulphate is 3, etc. 

In the reaction involving the reduction of manganese from a 
permanganate to a manganous salt, the apparent valence changes 
from 7 to 2 and the hydrogen equivalent of the compound is 
therefore 5. 

The equivalent weight of potassium permanganate, containing 



one atom of manganese in each molecule, is one-fifth of its molecu- 
lar weight or 31.606. Iron is oxidized in the reaction. Its appar- 
ent change of valence is from 2 to 3, its hydrogen equivalent is 
therefore 1 and its equivalent weight is its atomic weight. 
In the reaction: 

K 2 Cr 2 7 + 6FeCl 2 + 14HC1->2KC1 + 2CrCl 3 + 6FeCl 3 + 7H 2 0, 

potassium dichromate is the oxidizing agent. Closer examina- 
tion shows that the element chromium is the real cause of the 
•oxidizing action, since it leaves the negative radical by changing 
to a base-forming lower oxide. The change of valence (reduction) 
of chromium is from 6 to 3, its hydrogen equivalent is 3, and that 
of potassium dichromate, containing two atoms of chromium in 
the molecule is 6. The equivalent weight of potassium dichro- 
mate is then one-sixth of its molecular weight. 

The assignment of apparent valence to the oxidizing and reduc- 
ing elements is a valuable aid in balancing oxidation and reduc- 
tion equations. For example the equation last given is to be 
balanced. The empirical equation is 

K 2 Cr 2 7 +FeCl 2 +HCl->KCl+CrC] 3 +FeCl 3 +H 2 0. 

First inspect the equation to determine the oxidizing and reducing 
elements, which are seen to be chromium and iron. Determine 
the changes in apparent valence and from these their hydrogen 
equivalents and those of the compounds containing them. Write 
these above the respective compounds thus: 


K 2 Cr 2 7 +FeCl 2 +HCl->KCl+CrCl 3 +FeCl 3 -f-H 2 0. 

The hydrogen equivalent of the reducing agent is to be the 
coefficient of the oxidizing agent, and vice versa. These coeffi- 
cients will not thereafter be changed. From these will follow the 
coefficients of potassium chloride, chromium chloride and ferric 
chloride and from all of these will be calculated the coefficient of 
hydrochloric acid From the latter will follow the coefficient of 
water. The balanced equation is then': 

K 2 Cr 2 7 +6FeCl 2 +14HCl->2KCl+2CrCl3+6FeCl 3 +7H 2 0. 



45. Determine the equivalent weights of the oxidizing and reducing agents 
in the following equations and balance the latter: 

(a) KMn0 4 + MnS0 4 +KOH->K 2 S0 4 +Mn0 2 +H 2 0. 

(b) K 2 Cr 2 7 +SnCl 2 +HC1-*KC1 +SnCl 4 +CrCl 8 +H 2 0. 

(c) H 3 As03+l2+H 2 0-^H 3 As0 4 +HI. 

(d) HgCl 2 +SnCl 2 ->HgCl+SnCl 4 . 

(e) HgCl 2 +SnCl 2 ->Hg+SnCl 4 . 

46. How much potassium permanganate must be contained in 1000 cc 
of solution so that 1 cc shall be equivalent to 0.002 gm of iron? To 0.002 
gm of manganese? 

47. How much potassium dichromate must be contained in 1000 cc of 
solution so that 1 cc shall be equivalent to 0.005 gm of iron? 

48. A solution of potassium permanganate contains 2.83 gm in 1000 cc. 
What weight of pure ferrous ammonium sulphate, FeS0 4 (NH 4 ) 2 S0 4 .6H 2 0, 
should be taken for standardization so as to require 30 cc of the permanga- 
nate solution? 

49. A solution of potassium dichromate is of such concentration that 1 cc 
=c= 0.005 gm iron. A solution of potassium permanganate contains 3.26 
gm of the salt in 1000 cc of solution. To what volume should 1000 cc of 
the stronger solution be diluted to make the two equivalent to the same 
weight of iron? 

50. What weight of iodine is required for 1000 cc of a tenth-normal oxidiz- 
ing solution? 

51. What weight of arsenic trioxide is equivalent to 1 cc of tenth-normal 
iodine solution? 

Potassium Permanganate. — This substance is a very con- 
venient one for use as a standard oxidizing agent, not only because 
it is readily and quantitatively reduced by many reducing agents, 
but because it serves as its own indicator. The color of even a 
dilute solution is quite intense, while-4he^xeduced salt, man- 
ganous sulphate, forms colorless solutions. The least excess, 
after the oxidation is finished, is made evident by the color of 
the unchanged permanganate. Practically, standard potassium 
permanganate is most often used for the determination of iron, 
less frequently for the determination of manganese, and occa- 
sionally for the determination of tin or antimony and indirectly 
for the determination of phosphorus, calcium and many thero 
elements. Theoretically it could be used for the direct or indirect 
determination of any reducing agent or oxidizing agent but 
other methods are frequently better for substances other than 
those named. 


A solution of potassium permanganate, if made by simp 
dissolving the salt in water, decomposes slowly and must be f: 
quently restandardized. Morse 1 has shown that this is due to th 
presence of manganese dioxide, traces of which may have bee 
contained in the original salt or it may have been formed 
reduction of potassium permanganate by dust or organic impur 
ties. The work of Morse shows that if the solution is filtere 
through asbestos, and is kept in clean, glass-stoppered bottl 
its concentration will remain constant almost indefinitely. 


Iron may be readily and quickly titrated by a solution of po- 
tassium permanganate. For this purpose the standard solutior 
should be made in the decimal system. The normal system L< 
not adapted to this determination because the potassium per- 
manganate will not often be used for the determination of sub 
stances other than iron and the decimal system provides more sim- 
ple calculations. If the iron compound, as an ore, is taken ir 
such quantity that the weight of sample is a simple multipk 
of the iron equivalent of the standard, the burette reading is 2 
direct percentage reading. Thus, if 1 cc of standard is equivalei 
to 0.002 gm of iron and if 2 gm of iron ore is taken for analys 
each cubic centimeter of standard represents 0.1 percent of in 

Reduction of Permanganate by Chlorides. — One serioi 
disadvantage in the use of potassium permanganate for tl 
titration of iron lies in the fact that if hydrochloric acid is pre 
ent, it also is oxidized by the standard, the result being a net 
tious value for the percent of iron. But most iron ores dissob 
best in hydrochloric acid, the solubility in sulphuric acid being 
very slight. After solution of the ore is accomplished it is neces- 
sary to remove the hydrochloric acid or to find some method for 
avoiding its reducing action upon the permanganate. To 
move the acid, the solution of the ore may be evaporated wil 
sulphuric acid. 

Use of Manganous Sulphate.— Instead of removing tl 
hydrochloric acid it is possible to minimize its reducing actk 
by the addition of manganous sulphate to the solution. 

1 Am. Chem. X, 18, 401 (1896); 20, 521 (1898). 


irder to understand the remarkable action of manganous salts 
a preventing the oxidation of chlorides in dilute solutions it 
3 necessary to examine more closely the reaction between 
•otassium permanganate and ferrous chloride. This reaction, 
ccurring in an acid solution, is usually represented as follows: 

KMn0 4 +5FeCl 2 +8HCl->KCl+MnCl 2 +5FeCl 3 +4H 2 0. (1) 

'his reaction really takes place in at least two stages: 

KMn0 4 +4FeCl 2 +8HCl->KCl+MnCl 3 +4FeCl3+4H 2 0, (2) 
MnCl 3 +FeCl 2 ->MnCl 2 +FeCl 3 . (3) 

Equations (2) and (3) are seen by inspection to indicate the same 
nd products as equation (1) and if no side reactions of any sort 
ook place the ordinary calculations would not be affected by the 
•ccurrence of two stages in the reaction. But manganese 
richloride is a very unstable compound and if not reduced at 
>nce by ferrous salts in the immediate vicinity (a condition that 
specially maintains near the end of the titration) this compound 
>roceeds to decompose spontaneously: 

2MnCl 3 ^2MnCl 2 +Cl 2 . (4) 

\.gain, if this liberated chlorine could react quantitatively with 
errous chloride: 

■ 2FeCl 2 +Cl 2 -+2FeCl 3 (5) 

here would be no error involved in the titration for equations (2), 
4) and (5), in sequence, are still equivalent to equation (1). 
rhis condition does not obtain. Reaction (5) is slow and un- 
certain and one of the reacting bodies (chlorine) is a gas which 
•eadily escapes. Therefore, according to Barneby, 1 the function 
)f manganous salts, added in large quantities, is to increase the 
concentration of the manganous ion, thus obstructing the progress 
)f reaction (4). By this means the spontaneous decomposition 
)f manganese trichloride is prevented and it is reduced only by 
contact with ferrous chloride, with which it reacts very readily. 

Observation of End Point. — As ferrous chloride becomes 
oxidized a yellow color appears in the solution, deepening to 
*ed as the concentration of the ferric salt increases. Experience 

1 J. Am. Chem. Soc, 36, 1429 (1914). 


shows that this renders difficult or impossible the observatioi 
of the first faint pink of the permanganate anion, whicl 
constitutes the indication of the end point of the reaction. Then 
is no corresponding color in the case of sulphates of iron and tl 
goes to show that the color of ferric chloride solutions is not 
to the ferric ion but rather to hydrolysis. Partial or compl( 
hydrolysis may yield any or all of the products indicated ty 
these equations. 

FeOHCl 2 + H 2 0->Fe (OH) 2 C1 + HC1, 
Fe(OH) 2 Cl+H 2 0-+Fe(OH) 3 +HCl. 

The basic chlorides or ferric hydroxide form colloidal solutions 
(sols) which are colored. As these reactions are reversible 
the addition of an excess of hydrochloric acid hinders hydrolysis 
and lightens the color but it has already been shown that a 
large excess of this acid is undesirable because of its reducing 
action upon permanganate. Reinhardt 1 showed that the 
formation of color can be prevented by the addition of phos- 
phoric acid to the solution. Evidently, hydrolysis is here 
prevented by diminishing the concentration of one of the re- 
acting ions (ferric) through the formation of the phosphate, 
which has a very small ionization. ^ 

FeCl 3 +H 3 P0 4 ->FeP0 4 +3HCl. 

The " Zimmermann-Reinhardt solution' ' contains manganous 
sulphate, phosphoric acid and sulphuric acid and it is added to 
the largely diluted ferrous chloride solution just before titration 
of the latter by standard permanganate solution. 

Primary Standards. — Potassium permanganate cannot be 
obtained in a sufficiently high state of purity to make possible 
direct weighing of the substance for standardization. Primary 
standard reducing agents of known and uniform purity must be 
used. Some reducing agents that maf be used for this purpose 
are iron, ferrous ammonium sulphate, oxalic acid and ammonium 

Iron. — Pure iron is not obtainable commercially and can be 
prepared only with great difficulty. Iron wire may be obtained 

» Chem. Zg., 13, 323 (1889). 


Javing a fairly high degree of purity. A common fallacy has 
c fcund acceptance by many chemists, to the effect that the so- 
Jalled " piano wire," purchased for standardizing, is uniformly 
19.6 percent or some other stated percent of iron. The purity of 
iuch wire varies widely and this material cannot be used as a 
1 Irimary standard for accurate work unless it has been carefully 
) analyzed. Such analysis is best accomplished by electrolysis. 
The gravimetric determination of iron by precipitation as ferric 
lydroxide and weighing as ferric oxide is not accurate unless made 
vith extreme care, on account of the difficulty experienced in 
mrification of the precipitate. If iron is to be used for standard- 
zing solutions, weighed portions of the previously analyzed 
naterial are dissolved in dilute sulphuric acid. The necessary 
imitations as to accuracy render the use of iron as a primary 
tandard decidedly unsatisfactory, especially since it has become 
possible to obtain iron salts in a high state of purity. 

Ferrous Ammonium Sulphate. — Ferrous ammonium sulphate 
("Mohr's salt") may be purified very readily by crystallization 
and preserved without oxidation, if kept dry. Many analyses 
conducted upon high-grade samples have shown that the salt 
may be obtained in a state of practical purity, the percent of 
iron being almost exactly that calculated from the formula. 
Oxidation in air does not readily occur, so that the standard 
sample may be preserved almost indefinitely if kept in a stop- 
pered bottle. 

Sodium Oxalate. — This is one of the best primary standards 
now available for standardizing oxidizing solutions. The 
reaction with potassium permanganate is represented thus : 

5Na 2 C204+2KMn04+8H 2 S04->K 2 S04+2MnS04+5Na 2 S04 

+ 10CO 2 +8H 2 O. 

The salt may be purified by ordinary methods of recrystallization 
from water solutions or by the addition of alcohol to rather 
concentrated water solutions. By the latter method crystal- 
lization is more rapid and the. crystals are consequently finer. 
At temperatures between 240° and 250° the crystals may be dried 
completely and without decomposition. 1 The solid does not 
contain water of crystallization. The standardization of potas- 
1 Bur. Stand., Circ. 40. 


sium permanganate solutions is made by titration of weighed, 
portions in presence of sulphuric acid, at temperatures between 
80° and 90°. 

Standard, certified samples of sodium oxalate may be obtained 
from the Bureau of Standards and these may be kept indefinitely, 
if in a stoppered bottle. Stock solutions cannot be preservec 

Oxalic Acid. — Oxalic acid and ammonium oxalate may be puri- 
fied by recrystallization to the condition required for primary 
standards. If previously analyzed they are nearly as satisfac- 
tory as ferrous ammonium sulphate. The student is cautione( 
against the practice of assuming the purity of any of his primary, 
standards without an analysis as the basis for the assumption. 

Reduction of the Iron Solution.— Iron exists in the ferric con- 
dition in most ores or other minerals. In order to reduce the 
solution of ferric salt either stannous chloride, zinc, sulphurous 
acid or hydrogen sulphide may be employed. The first two an 
the only ones now commonly used. 

Stannous chloride, in solution, possesses the advantage of 
instantaneous action if added to the hot solution of ferric chloride. 
If the iron is to be reduced by stannous chloride an addition oi 
this salt to the ore during the process of solution will materially 
hasten the action. For the final reduction the stannous chloride 
solution may be added from a pipette, the disappearance of th( 
red color of basic ferric chloride providing an approximate indica- 
tion of the end-point. 

In the analysis of iron ores there is occasionally trouble at this 
point unless certain precautions have been taken. In the first 
place, many iron ores contain appreciable quantities of organic 
matter and this serves to produce a yellow color when the ore is 
dissolved. As color due to this cause does not disappear when 
the iron has been reduced it is not possible to determine when the 
correct amount of stannous chloride has been added. This 
trouble may be avoided by igniting the weighed sample for a 
short time in a porcelain crucible, before dissolving. 

The second cause of irremovable color lies in fusion of insoluble 
residues in platinum crucibles. The pyrosulphate which is used 
as a flux dissolves traces of platinum and this, with stannous 
chloride, forms a yellow solution containing a complex of tin and 


platinum. This interference is avoided by the substitution of 
porcelain crucibles for those of platinum. 

After a slight excess of stannous chloride has been added the 
solution is cooled and a considerable excess of mercuric chloride 
is added, the unused stannous chloride being thereby oxidized: 

2HgCl 2 +SnCl 2 -+SnCl4+2HgCl. 

Mercuric chloride will not oxidize ferrous chloride and hence may 
be left in the solution. If an insufficient excess of mercuric 
ichloride is used, or if it is added too slowly, free mercury may be 
produced : 

HgCl 2 +SnCl2->SnCl 4 +Hg. 

JThe indication of such action is the appearance of a gray precipi- 
tate of mercury instead of the characteristic white silky crystals 
of mercurous chloride. If mercury is so produced the deter- 
mination is ruined because this mercury will itself reduce some of 
standard oxidizing solution during the process of titration of the 

Stannous chloride cannot be used to reduce ferric solutions 
previous to titration by potassium permanganate unless the 
interference of chlorides is to be prevented by the addition of 
manganous sulphate. Instead, pure zinc or zinc of predeter- 
mined iron content may be used to reduce the iron. For this 
purpose an approximately weighed quantity of granular zinc or 
jzinc dust may be added directly to the solution and dissolved in 
I it, or the ferric solution may be passed through a reductor. The 
'latter is a tube having the dimensions of a burette, filled with 
amalgamated zinc and having a dropping funnel fixed in the top. 
| The acidified solution is passed through the tube once or twice 
[and the iron is thereby reduced. Most zinc contains iron or other 
reducing matter and if it is to be dissolved in the iron solution, 
as above described, a blank determination should be made to 
determine the amount of standard that will be required to oxidize 
the reducing matter of the zinc. 

On account of the slow reducing action of zinc, stannous chloride 
is much to be preferred, where conditions will permit its use. 

Sulphurous acid and hydrogen sulphide reduce iron solutions 
quickly but the disadvantages involved in their preparation 
prevent their extensive use for this purpose. 



Exercise: Preparation of Standard Potassium Permanganate Sol 
tion. — Calculate the weight of potassium permanganate required fc 
2500 cc of solution, each cubic centimeter to be equivalent to 0.005 
ot iron, using for this purpose the method shown on page 196. We 
2 percent more than the calculated weight and dissolve in about 1( 

cc of distilled water. Allow to stand at leg 
an hour, then filter through a double asbesl 
filter arranged as in Fig. 67, all vessels hi 
ing been previously cleaned by chromic a< 
mixture. Dilute the solution to 2500 cc ai 
standardize as follows: 

Calculate the weight of ferrous ammonii 
sulphate that will be approximately equh 
lent to 35 cc of the permanganate sol 
tion. Weigh three such portions into Erl( 
meyer flasks of 250 cc capacity. Dissob 
each portion, immediately before titratii 
in 50 cc of distilled water and add 10 cc 
dilute sulphuric acid. Titrate at once 
the potassium permanganate solution, tl 
first appearance of a permanent pink ti 
being taken as the end-point. If sodii 
oxalate is used as the primary standard 
proceed as follows: Warm the acid solutioi 
of oxalate to 90° and titrate, stirring vigor- 
ously and continuously. The permanganate 
must not be added more rapidly than 10 tc 
15 cc per minute and the last 0.5 to 1 cc 
must be added drop-wise, with particular 
care to allow the color from each drop tc 
disappear before another drop is added. A1 
the end of the titration the temperature 
must not be below 60°. A thermometei 
should be kept in the solution during the 
entire experiment. 

Calculate the value of the permanganate 
solution in terms of iron and dilute two liters of the solution to the exacl 
equivalence of 0.005 gm of iron per cubic centimeter. From this iror 
value calculate also the calcium and manganese values (see pages 252 and 
253), and record all upon the label of the bottle and in the record book. 
Determination of Iron in an Ore. — Sample the ore and grind the 
last selection to pass the 100-mesh sieve Weigh exactly 0.5 gm of ore 
on the counterpoised glasses, brushing into each of three porcelain cru- 

Fig. 67. — Apparatus for 
filtering potassium per- 
manganate solution. 


jibles. Heat the crucibles without covers for 5 minutes, using the 

frdinary desk burner, then allow the crucibles to cool, place in casseroles 

Lnd add to each 25 cc of concentrated hydrochloric acid. If method 

b) is to be used for reducing the iron add also at this point 5 cc of 5 

Percent stannous chloride solution. Cover and warm until solution 

k complete or until no further action appears to take place. If the 

jesidue is not colored, proceed, without filtration, by one of the methods 

a) or (b) given below. If the residue is colored it may contain iron. In 

[his case filter on a small paper and wash the paper free from iron solu- 

iion with hot water. Set the filtrate and washings aside and burn the 

paper at a low temperature in a porcelain crucible. If the residue is 

! mail in amount and apparently contains little silicious matter it may be 

I lecomposed by fusing with potassium pyrosulphate. Cool and dissolve 

[he mass in hot water, adding the solution to the former filtrate. 

pow proceed by one of the methods given below. 

(a) Add 4 cc of concentrated sulphuric acid to the iron solution and 
tkvaporate by holding the casserole over a free flame, keeping in constant 
[notion to hasten evaporation and to prevent bumping. Evaporate 
ijintil the characteristic white fumes of sulphuric acid appear, this being 
the point at which all water and hydrochloric acid have been expelled, 
pool and dilute to 50 cc, rinsing the solution into a 250-cc Erlenmeyer 
llask. Add 2 gm of granular zinc, as free as possible from iron, place a 
funnel with a short stem in the flask and warm until the zinc is dissolved, 
rhe iron solution should now be quite colorless or faintly green. Cool 
and titrate at once with standard potassium permanganate solution. 
Make a blank determination without the ore, to determine the amount 
pf iron and other reducing matter in the zinc, by dissolving 1 gm of zinc 
In 25 cc of dilute sulphuric acid, under the same conditions as noted 
above, and titrating. Express the result of the blank experiment as 
phe number of cubic centimeters of potassium permanganate reduced 
by 1 gm of zinc. The proper value will then be subtracted from the 
Lrolume of permanganate used in the iron titration and the percent of 
ron then calculated. 

I If the zinc is nearly pure it will dissolve very slowly. Solution may 
be hastened by dropping a coil of platinum wire into the flask, and keep- 
ing it in contact with the zinc. 

(b) Concentrate the iron solution, if necessary, to about 50 cc and 
transfer to a 1000 cc Erlenmeyer flask. While the solution is nearly boil- 
ing add, drop by drop from a pipette, a 5 percent solution of stannous 
chloride until the ferric chloride has just been reduced, this being made 
svident by the disappearance of the red color. Add two drops more of 
stannous chloride solution then cool quickly by immersing the flask 
in running water. When cool add. all at once, 25 cc of a 5 percent 


solution of mercuric chloride and mix well with the solution. The pr 
cipitate should be pure white mercurous chloride without a trace of gray 
mercury. Dilute to 500 cc and add 50 cc of a solution containing 1 
gm of phosphorous pentoxide, 245 gm of sulphuric acid and 67 gm 
crystallized manganous sulphate in each liter of solution. Titrate 
once with standard potassium permanganate solution and calculate the 
percent of iron in the ore. 



Calcium may be precipitated as oxalate, filtered and washed 
free from ammonium oxalate, then dissolved in hot, dilute sul- 
phuric acid and the resulting oxalic acid titrated with potassium 

CaC 2 4 + H 2 S0 4 ->CaS0 4 + H 2 C 2 4 , 
2KMn0 4 +5H 2 C 2 4 +3H 2 S0 4 ->K 2 S0 4 +2MnS0 4 + 

10CO 2 +8H 2 O. 

In order to determine the equivalent weight of oxalic acid it is 
necessary here, as in other cases, to note the change which it 
undergoes as it becomes oxidized. The products of oxidation are 
water and carbon dioxide. By using the method explained on 
page 241 it will be found that the apparent change of valence of 
carbon is 1. Since each molecule of oxalic acid contains two 
atoms of carbon the hydrogen equivalent of oxalic acid as a 
reducing agent is 2 and its equivalent weight is one-half its mo- 
lecular weight. When oxalic acid reacts as an acid its hydrogen 
equivalent is also 2, although there is no necessary connection 
between the two cases. Since one molecule of oxalic acid is 
formed by the decomposition of one molecule of calcium oxalate, 
containing one atom of calcium, the equivalent weight of calcium 
is also one-half its atomic weight. The calcium equivalent of the 
potassium permanganate solution may be calculated from the 
iron equivalent by the expression: 

eq. wt. of calcium J . , . . 

— =-: Xwt. iron = wt. of calcium. 

eq. wt. ot iron 


52. 1 cc of a solution of potassium permanganate is equivalent to 0.002 
gm of iron. Calculate the weight of calcium, calcium oxalate and oxalic 
acid equivalent to 1 cc. 


53. What weight of potassium dichromate is equivalent to 4.75 gm of 
otassium permanganate, both being used for the oxidation of iron? 

54. What weight of crystallized oxalic acid, H2C2O4.2H2O, is equivalent to 
.56 gm of crystallized ferrous ammonium sulphate, Fe(NH 4 ) 2 (S0 4 )2 6H 2 0? 

55. WTiat is the normality of a solution of potassium permanganate, 1 cc 
f which is equivalent to 0.005 gm of iron? 

56. 30 cc of potassium permanganate solution oxidizes 0.1905 gm of 
rystallized oxalic acid. How dilute 1000 cc of the solution to make 
;xactly tenth-normal? 

Determination. — Use samples of not more than 0.2 gm of the calcium 
;ompound. Dissolve, precipitate, filter and wash the calcium oxalate 
iccording to the method already learned in an earlier exercise (page 81). 
Pierce the point of the filter paper and wash the precipitate into a beaker 
arith the least possible quantity of hot water. Thoroughly wash the 
paper in the funnel with successive portions of 5 cc of hot dilute sulphuric 
icid until any remaining precipitate shall have been dissolved. Again 
«rash the paper with hot water, then warm (but not boil) the acid and 
precipitate in the beaker until the precipitate is dissolved. Titrate 
with standard potassium permanganate, inserting a thermometer and 
keeping the solution at 80° to 90° (refer to page 250). Calculate the 
percent of calcium in the sample. 


In acid solutions potassium permanganate is always reduced 
to the form of manganous salts, manganese being thereby reduced 
to its lowest state of oxidation, corresponding to the monoxide. 
In basic solutions the reduction goes only so far as to produce 
manganese dioxide. If the reducing agent is a manganous salt 
it is also oxidized to manganese dioxide : 

2KMn0 4 +3MnS04+2H 2 0-^K 2 S04+5Mn0 2 +2H 2 S04. (a) 

This reaction may be made the basis of a determination of 
manganese by titration with potassium permanganate. 1 The 
manganese ore is dissolved in hydrochloric acid, manganous 
chloride being formed. The chloride is converted into sulphate 
and the hydrochloric acid removed by evaporation with sul- 
phuric acid. When the titration is made the solution must be 
only feebly basic. If a strong base is present a reduction of 

1 Volhard: Chem. News, 40, 207 (1879). 


potassium permanganate to potassium manganate occurs, tl 
reducing agent being manganese dioxide or manganese hydroxic 

2KMn0 4 +4KOH+Mn0 2 ^3K 2 Mn0 4 +2H 2 0, (1 

4KMn04+6KOH+Mn(OH) 2 -^5K 2 Mn0 4 +4H 2 0. 1 

These undesirable reactions may be prevented by having presei 
a considerable excess of a weak base, such as is produced by shs 
ing an excess of zinc oxide with water, this giving a suspension 
zinc oxide in a saturated solution of zinc hydroxide. The latte 
is so dilute and so weakly ionized that the formation of potassiui 
manganate does not take place. It does provide, however, sui 
cient base to neutralize the sulphuric acid produced by reaction 
(a), because the excess of zinc oxide keeps the solution saturated 1 
with zinc hydroxide. 

Manganese dioxide possesses, to a slight extent, acid-forming 
properties, since it is able to produce a class of salts that are 
theoretical derivatives of manganous acid, H 2 MnOs( = H 2 O.Mn0 2 ). 
This acid is not known in the free state but certain manganites, as 
those of calcium and zinc, CaMn0 3 and ZnMn0 3 , are produced 
when manganese dioxide is formed in presence of soluble calcium 
or zinc compounds. If no such metal is present at the moment 
of oxidation of manganous salts to manganese dioxide, manganoi 
manganite, MnMn0 3 , is precipitated, thus removing a certs 
amount of unoxidized manganese from the solution. An ern 
would thereby be introduced but this is prevented by havii 
zinc hydroxide present. The saturated solution of manganes 
dioxide can have but a small concentration of manganous acid ai 
this is at once precipitated as zinc manganite. The addition 
zinc oxide therefore serves a double purpose. It maintains 
feebly basic solution throughout the titration and also prevents 
the precipitation of manganous manganite. 


57. From the equation for the reaction between potassium permangam 
and manganous sulphate, calculate the equivalent weight of potassk 
permanganate and of manganese in manganese sulphate. 

58. What weight of potassium permanganate must be contained in 1( 
cc of a solution, 1 cc of which will oxidize 0.002 gm of manganese? • 

59. If a solution of potassium permanganate is fifth-normal with respe 


to iron in acid solution what is its normality with respect to manganese in 
basic solution? 

60. 1 cc of a solution of potassium permanganate is equivalent to 0.010 
gm of iron. What weight of manganese will be oxidized by 1 cc? 

Determination. — Calculate the approximate weight of pyrolusite 
or other manganese compound that is necessary to reduce about 40 cc of 
the standard potassium permanganate solution already made, arbitrarily 
assuming a state of purity for the sample and also that one-fifth of the 
sample is finally to be titrated. Dry the sample to constant weight at 
120° and weigh the calculated quantity of dried material, placing in a 
casserole. Dissolve by warming with concentrated hydrochloric acid. 
When solution is complete or when no further action is apparent filter 
and wash the residue and paper, preserving the filtrate and washings. If 
the residue contains any dark material, burn the paper in a platinum 
crucible and fuse with 1 to 2 gm of sodium carbonate. If manganese is 
present the fusion will be colored green by sodium manganate. Dis- 
solve the fused mass in hot water and add to the main solution. 

Remove hydrochloric acid by evaporating with sulphuric acid (about 
2 cc) to the appearance of fumes of sulphur trioxide. Redissolve, adding 
a little nitric acid if necessary, wash into a 250 cc volumetric flask, dilute 
to the mark and mix. Treat 50 cc portions as follows : Measure into an 
Erlenmeyer flask of 1000 cc capacity and neutralize by the addition of 
zinc oxide suspended in water, shaking and continuing the addition until 
any iron is precipitated as ferric hydroxide and a sufficient excess of zinc 
oxide is present to maintain a milky appearance throughout the subse- 
quent titration. Dilute to about 300 cc, heat nearly to boiling and 
titrate, adding the potassium permanganate solution 5 cc at a time until 
a permanent color is produced. Treat a second 50 cc portion of the 
[manganese solution in a similar manner but adding, all at once, 5 cc 
less of standard solution than was added to the first portion then adding 
1 cc at a time. To a third portion add the entire quantity used in the 
isecond, less 1 cc, then complete the titration by adding the standard 
solution 0.1 cc at a time. Treat the fourth portion in the same manner 
'as the third. From the last two titrations calculate the percent of 
j manganese in the ore. 

Available Oxygen 

Manganese dioxide is used not only as an ore of manganese but 
jalso as an oxidizing agent in various laboratory processes and in a 
j commercial way, as, for example, in the production of chlorine 
sfrom hydrochloric acid. In such cases the percent of manganese 


is not as important as is that of oxygen available for oxidation 
processes. In most cases the available oxygen may be calculate! 
with sufficient accuracy for commercial requirements from the 
percent of manganese. This can be accurately done only in case 
no other manganese compound and no other peroxide is present. 
Generally no other peroxide is present, although manganese fre- 
quently occurs in pyrolusite in small quantities as other com 
pounds than the peroxide. If an accurate determination of availa- 
ble oxygen is required it may be made by reducing a weighed sampl 
of the manganese dioxide by a measured amount of a reducin 
agent, titrating the excess of the latter by standard potassiu 
permanganate. The reducing agent may be any of those already 
discussed in connection with standardization of potassium per- 
manganate. Ferrous ammonium sulphate or oxalic acid is to 
be preferred. The reaction between manganese dioxide and these 
reducing agents in presence of sulphuric acid is represented by 
the following equations, which should be balanced by the student 
as an exercise in calculation of hydrogen equivalents. Determine 
also what fraction of the total oxygen of manganese dioxide is 
" available." 

Mn0 2 +FeS0 4 +H 2 S0 4 ^MnS0 4 +Fe 2 (S0 4 ) 3 +H 2 0, 
Mn0 2 +H 2 C 2 4 +H 2 S0 4 ->MnS04+C0 2 +H 2 0. 


Another method for determining the available oxygen 
described on page 265. 

Determination. — Dry the sample of either pyrolusite or commercial 
manganese dioxide to constant weight at 120°. Calculate the weight 
that would be equivalent to approximately 40 cc of the standard potas- 
sium permanganate already made, arbitrarily assuming that the sample 
is pure manganese dioxide. Weigh samples of the calculated weight 
into 250 cc Erlenmeyer flasks. Calculate the weight of crystallized 
ferrous ammonium sulphate or oxalic acid that would reduce approxi- 
mately 50 cc of the standard solution of potassium permanganate and 
add this quantity to each flask. Calculate the approximate volume of 
dilute or concentrated sulphuric acid necessary to enable the reactions 
to proceed and add three times this volume to each flask. Add 50 cc 
of water, warm to 70° and titrate immediately with standard potassium 
permanganate solution. Calculate the percent of available oxygen in 
the sample. 


Potassium permanganate solution is also useful for the titra- 
tion of hydroferrocyanic acid and hydroferricyanic acid. Ferro- 
cyanides are oxidized in acid solution: 

10H 4 Fe(CN),+2KMnO 4 +3H^O^10HJF , e(CN) i +K 2 SO4 
J +21II1S04+8HJO. 

1 Perricyanides may be reduced in basic solution by ferrous sulphate 
land then titrated by potassium permanganate. The reaction 
e ts as follows: 
?JK J Fe(CN) i +FeS04+3KOH-^K4Fe(CN) i +Fe(OH),4'K 2 S04. 

I Potassium Dichromate. — The equation for the reaction of 
r- potassium dichromate with ferrous salts is given on page 242. 
ojrhis substance possesses several advantages over potassium per- 
se nanganate as a standard oxidizing agent. It is relatively more 
>y (table and therefore may be obtained in a state of uniform purity. 
itlThis makes it possible to standardize solutions by direct weighing 
k w hen the degree of purity of the salt has been established by analy- 
i Wl The relative stability is the same with solutions and the 
1 tandard solution can be kept almost indefinitely without chang- 
ng its concentration. Potassium dichromate may also be used 
e titration of iron and other reducing agents in presence of 
liydrochloric acid or chlorides, without oxidation of the latter 
place. This is a very decided advantage in the determina- 
of iron since it makes possible the use of stannous chloride as 
ucing agent without the addition of manganous sulphate 
phosphoric acid. There is no indicator that can be added 
to the solution which is being titrated by potassium 
mate and the color of the standard solution is not sum- 
intense to be of any use for this purpose. The indicator 
is commonly used is potassium f erricyanide, placed in drops 
white porcelain "spot plate." Drops of the solution are 
from time to time by means of a stirring rod and allowed 
the drops of f erricyanide. So long as ferrous iron is 
the blue of ferrous f erricyanide is apparent on the spot 
When the last trace of iron has been oxidized there is 
ced on the plate only the light brown ferric ferricyanide. 
being nothing in the appearance of the solution of the iron 
o indicate the approach to the end-point, the titration is 


necessarily somewhat tedious unless a system is devised fc 
rapid readings. Such a system has been used in connection wit 
the determination of manganese and is indicated in the ne? 
exercise and this removes the last objection to the use of pota 
sium dichromate for the titration of iron. 


61. A solution of potassium permanganate contains 25.38 gm in 1000 c 
What must be the concentration of a potassium dichromate solution in ord 
that it shall have the same oxidizing power toward iron? 

62. Balance the following equation and calculate the equivalent weigh 
of tin. 

K 2 Cr 2 7 +SnCl 2 +HC1->KC1 +CrCl 3 +SnCl 4 +H 2 0. 

Exercise: Preparation of Standard Potassium Dichromate Solu- 
tion. — The solution should be of such concentration that 1 cc is equivalent 
0.005 gm of iron. Calculate the weight of potassium dichromate neces- 
sary for 2000 cc of such a solution. If the salt is known to be pure, 
weigh exactly the calculated weight and omit further standardization. 
If it is not pure but its oxidizing power known from previous determina- 
tions, calculate the weight of impure sample required and use this weight.; 
If nothing is known of the purity use 1 percent more than the weight oi 
pure salt required for 2500 cc of solution and standardize the solution 
directed below. In any case dissolve the weighed salt and dilute to t 
proper volume. In case titration for standardization is to be omitt 
and direct weighing is to be made the basis for standardization, 2000 
of the solution should be accurately made and poured into a dry bottl 

Standardization, if this should be necessary, is accomplished by titn 
tion against ferrous ammonium sulphate or iron wire, the first meth 
being preferable. Write and balance the equation for the oxidation 
ferrous sulphate by potassium dichromate in presence of sulphuric acL 
referring, if necessary, to the equation for the oxidation of the chloride 
page 242. Calculate the weight of crystallized ferrous ammoniuir 
sulphate necessary to reduce 35 cc of the dichromate solution. Weigr 
five portions of exactly this weight into 250 cc beakers and dissolve eacl 
in 50 cc of recently boiled and cooled water just before titrating. Pre- 
pare a 0.01 percent solution of potassium ferricyanide and place a droj 
in each of the depressions of a white porcelain spot plate. Add to th( 
solution of ferrous ammonium sulphate three times the calculated amoum 
of sulphuric acid necessary, as indicated by the equation, and titrate ai 
once, as follows: To the first solution add the dichromate solution 5 c< k 
at a time, stirring well after each addition, and test by removing a droj 
by means of the stirring rod and touching to a drop of potassium ferri- 


cyanide solution on the spot plate. The end-point is reached when a 
blue color is no longer produced on the plate, after standing for 2 
minutes. Dust or reducing gases will interfere by reducing traces of ferric 
chloride. Titrate the second solution by adding 5 cc less than the 
amount of dichromate solution used in the first, then adding 1 cc at a 
time. Titrate the third solution by adding 1 cc less than the total 
used in the second, then adding 0.1 cc at a time. Titrate the fourth 
and fifth solutions in the same manner and take the average of the 
last three titrations for permanent record. Calculate the value of the 
solution in terms of iron. Dilute to make 1 cc equivalent to 0.005 gm 
of iron. 

Instead of weighing five portions of ferrous ammonium sulphate a 
standard solution may be made by dissolving ten times the required 
amount, adding the necessary sulphuric acid and diluting to 500 cc. 
Portions of 50 cc are then measured and titrated. The solution oxidizes 
upon exposure to air and the method of weighing separate portions and 
dissolving just before titration is preferable. 


Determination of Iron. — Prepare a sample of iron ore by grinding 
Ito pass a 100-mesh sieve. Weigh five portions of exactly 0.5 gm each, 
(using the counterpoised glasses and brushing the ore into porcelain 
crucibles. Heat the inclined crucibles for 5 minutes over the desk 
mrner, cool, place in casseroles and dissolve in hydrochloric acid, with or 
rithout the addition of stannous chloride. Reduce each solution just 
>efore titration, following the directions given for dissolving and reduc- 
[ng by method (b) of the permanganate method. Do not add the solu- 
tion of manganous sulphate and of phosphoric and sulphuric acids. 
>ilute to 100 cc. The titration is carried out exactly as directed for 
[standardizing potassium dichromate solution. Calculate the percent 
)f iron in the ore. 


The most important ore of chromium is a compound of iron 

md chromium known as "chromite," having a composition 

(orresponding with the formula FeO.Cr 2 03. Although chro- 

ium is here in its lowest state of oxidation the substance is thought 

be a salt of a hypothetical chromous acid, H 2 Cr 2 4 . Chromite 

mnot be dissolved in acids nor is it possible to decompose it 

isily by fusion with alkali carbonates. Fusion with sodium 

;roxide decomposes it, oxidizes the iron to ferric oxide and the 


chromium to chromium trioxide, forming then sodium chromate 
Upon dissolving in water and filtering, ferric oxide is removed 
The addition of acid produces sodium dichromate. This can 
then be reduced by adding an excess of a standard reducing agent, 
such as ferrous ammonium sulphate, titrating the excess by stand- 
ard potassium dichromate or permanganate. The reactions are 
expressed by the following equations, which should be balanced 
by the student. 

FeCr 2 04+Na 2 2 ^Fe 2 3 +Na 2 Cr04+Na 2 0, 
Na 2 Cr0 4 +HCl-+Na 2 Cr 2 7 +NaCl+H 2 0, 

Na 2 Cr 2 7 +FeS0 4 +HCl->NaCH-Fe 2 (S04)3+ 

FeCl 3 +CrCl 3 + H 2 
Iodine and Sodium Thiosulphate. — Iodine and thiosulphate* 
react quantitatively, forming sodium iodide and sodium tetra 
thionate : 

2Na 2 S 2 3 +l2^Na 2 S40 6 +2NaI. 

This is an oxidation of sodium thiosulphate by iodine, which ii 
itself reduced. The solution may be originally neutral or acid 
or alkali bicarbonates may be present. Normal carbonates 03 
hydroxides should not be present since they also combine witl 
iodine : 

2NaOH + I 2 -+NaIO + Nal + H 2 0, 
2Na 2 C0 3 +I 2 ->NaIO+NaI+C0 2 . 

The color of dilute solutions of iodine is sufficiently intense t< 
serve as a fairly accurate indicator. Much more accurate r© 
suits are obtained by the use of starch as an indicator, mere trace 
of iodine producing a visible blue or rose red color with starch 
Because of the fact that iodine is an excellent oxidizing agent fo 
many substances when a bicarbonate is present, and that hydri 
odic acid is oxidized by many oxidizing agents when an acid i 
present, free iodine being liberated, the two standard solution 
of iodine and sodium thiosulphate form a most useful pair fo 
volumetric analysis. As an example of their use the reactions o 
arsenic may be noticed. Arsenious acid or an arsenite is o: 
dized by free iodine thus: 

H 3 As0 3 +I 2 +H 2 O^H 3 As0 4 +2HI, 

Na 3 As0 3 +I 3 +H 2 0«=>Na 3 As0 4 +2HI. 


This reaction does not take place quantitatively but is reversible. 
If, however, sodium bicarbonate is present in excess the hydri- 
odic acid is neutralized as fast as it is formed and the reaction is 
completed. Standard iodine solution may, in this way, be used 
for the titration of arsenious acid. On the other hand arsenic 
acid is reduced by hydriodic acid: 

H 3 As0 4 +2HI^H 3 As0 3 +H 2 0+I 2 . 

This is seen to be the reverse of the reaction expressed above 
and it would follow that it also is incomplete unless one of the 
products is removed. This may be done by adding sodium 
thiosulphate to remove the iodine, in which case the standard 
sodium thiosulphate indirectly titrates the arsenic acid. In prac- 
tice hydriodic acid is not kept as a reagent because of its insta- 
bility but potassium or sodium iodide and hydrochloric acid are 
used instead, hydriodic acid being thus made available in the 

Standardization. — By properly purifying iodine standard 
solutions may be made by direct weighing. Commercial iodine 
is usually not sufficiently pure for this purpose and must be 
analyzed if it is. to be used in this way. Iodine solutions will not 
remain constant in oxidizing power because of interaction between 
iodine and water, and it is usually not advisable to attempt to 
dilute solutions to a definite concentration because they must be 
restandardized after short intervals of time. For this reason 
standardization by direct weighing is. not practicable and the 
iodine need not be purified before dissolving. The solution may 
then be standardized by titrating against any standard reducing 
solution. The best substances for this purpose are sodium thio- 
sulphate and arsenious oxide. 

Iodine does not dissolve easily in water but is readily soluble 
in a solution of potassium iodide or sodium iodide. Such a solu- 
tion probably contains an iodide having the formula KI 3 or Nal 3 . 
Many organic liquids are good solvents for iodine. Examples are 
the alcohols and acetic acid. These will be discussed in the sec- 
tion dealing with the analysis of fats and oils. When starch and 
iodine are brought to gether a deep, indigo-blue color is produced 
and this se r ves as a very delicate test for either starch or iodine. 
The natureof the blue substance has long been the subject of 


investigation and discussion. It is probably a so lid solutionof 
iodine ip stfl/rnh . 

Sodium thiosulphate may sometimes be obtained in a suffi- 
ciently pure condition to allow standardization by direct weigh- 
ing. It is better to make the solution somewhat more concen- 
trated than that desired and to standardize and dilute to a 
definite concentration. For standardization the solution may be 
directly titrated against standard iodine solution or indirectly 
against potassium dichromate or a salt of copper. It has already 
been stated that potassium .dichromate may be obtained in a 
state of uniform purity. If to a standard solution of this salt 
potassium iodide and hydrochloric acid are added, iodine is 
liberated as follows: 

K 2 Cr 2 7 +6KI-r-14HCl->8KCl+2CrCl 3 +7H 2 0+3I 2 . 

The liberated' iodine may be titrated by sodium thiosulphate and 
the latter thus standardized. The solution of potassium dichro- 
mate used for iron determinations may be used also for this 
purpose. It was standardized in the decimal system, however, 
and it will be necessary to calculate its value in the normal system 
because the solution of sodium thiosulphate is to be used for the 
determination of several different substances. The following ex- 
ample will illustrate the method of calculation of standardization. 

Example. — 40 cc of a solution of potassium dichromate liberates 

iodine equivalent to 22 cc of sodium thiosulphate solution. 1 cc of 

potassium dichromate solution is equivalent to 0.005 gm of iron. What 

is the normality of the thiosulphate solution? 

1 cc of sodium thiosulphate solution is equivalent to ^ c c of potas- 

sium dichromate f solution and to ^ X 0.005 gm of iron. A normal 

solution would be equivalent to 0.05584 gm of iron. Therefore the 

.J . 40X0.005 rt1MOW 
normality is 22XO05584 = ai ° 28 N ' 

The standardization against a salt of copper is also an excellent 
method. . This is described on page 268 in connection with th 
determination of copper. 

Sodium thiosulphate is quite stable in solution and may be 
kept for months without appreciable change in concentration if 
the water contains no trace of acid. Even carbonic acid causes 



decomposition and free sulphur is deposited from the solution, 
sulphurous acid being formed. This is because thiosulphuric 
acid is very unstable and rapidly decomposes : 

Na 2 S 2 3 +H 2 C0 3 ->Na 2 C03+H 2 S 2 03, 
H 2 S 2 03 — >H 2 S03-f-S. 

Even a very small amount of carbonic acid is sufficient to start 
the decomposition by liberating some thiosujphuric acid. As 
sulphurous acid accumulates it aids the decomposition which is 
thus progressive: 

Na 2 S 2 03+H 2 S03->Na 2 S03+H 2 S 2 3 , 
H 2 S 2 03 — >H 2 S03-f-S. 

In order to avoid starting this series of reactions the water 
should be boiled and cooled before making the solution. 


63. A solution of potassium dichromate contains 4.95 gm of the salt in 
1000 cc. What weight of sodium thiosulphate is equivalent to 1 cc? 

64. A solution of potassium dichromate contains 6.235 gm in 1000 cc. 
30 cc of this solution is equivalent to 42.9 cc of sodium thiosulphate solution. 
What is the normality of the latter? 

65. What weight of potassium dichromate must be dissolved in 250 cc to 
make a solution, 25 cc of which is equivalent to 35 cc of sodium thiosulphate 
solution containing 13.65 gm of the crystallized salt in 1000 cc? 

66. 1 cc of potassium dichromate solution is equivalent to 0.005 gm of 
iron. What is the iodine equivalent? 

67. 25 cc of iodine solution is equivalent to 0.125 gm of potassium dichro- 
mate. To what volume should 1000 cc be diluted to make the solution 

68. 20 cc of potassium dichromate solution oxidizes 0.0240 gm of oxalic 
acid, H2C2O4.2H2O. 1 cc of the same solution oxidizes the same weight of 
iron as does 1.2 cc of potassium permanganate solution. What is the nor- 
mality of the latter solution? 

Exercise : Preparation of Tenth-normal Sodium Thiosulphate Solu- 
tion. — Calculate the weight of crystallized sodium thiosulphate, 
Na 2 S 2 03.5H 2 0, required for 2500 cc of tenth-normal solution. Crush 
the salt and dissolve 2 percent more than this weight in cold, recently 
boiled water and dilute to 2500 cc. Keep the bottle well stoppered, 
and out of direct sunlight. 

Make 200 cc of a solution containing 30 gm of potassium iodide. 
The starch solution is made as follows: Moisten 1 gm of starch with 


cold water to make a thick paste. Heat 200 cc of water to boiling 
and pour it into the starch paste. Boil, with constant stirring, for one 
minute. The solution does not keep well and should be made each 
day as required. Standardize the sodium thiosulphate solution against 
potassium dichromate. If the solution used in iron determinations is 
at hand, use this, otherwise make 250 cc of exactly tenth-normal 
solution by weighing the salt, dissolving and diluting to the required 
volume. Measure 35 cc of either solution into an Erlenmeyer flask, add 
40 cc of potassium iodide solution and 10 cc of concentrated hydro; 
chloric acid. Titrate at once with sodium thiosulphate solution, 
deferring the addition of starch as long as possible. If starch is added 
before the iodine is nearly all reduced a precipitate of starch iodide 
will form, free iodine being, in this way, removed from the possibility 
of reacting with thiosulphate. A false end point is then obtained. 

The solution of chromium chloride, formed by reduction of potassium 
dichromate, is green. The solution has an amber tint as long as much 
free iodine is present. Upon the addition of starch the solution acquires 
a blue-green color and the change to pure green at the end point may be 
difficult to detect at first trial. With a little experience the difficulty will 
disappear. Make at least three titrations and calculate the normality 
of the sodium thiosulphate solution. Dilute to make tenth-normal. 

Make a blank test upon the potassium iodide, omitting the potassium 
dichromate but adding the hydrochloric acid. If iodine is found, 
correct the observed volume of sodium thiosulphate before calculating 
its concentration. 

Oxidizing Power op Peroxides 

Such peroxides as those of manganese, barium, lead, and hy- 
drogen readily oxidize hydriodic acid and liberate iodine. The 
titration of the latter by standard sodium thiosulphate solution 
constitutes an indirect determination of the oxidizing power, or 
"available oxygen" of the peroxide. In practice it is sometimes 
not found convenient to add an iodide and hydrochloric acid 
directly to the peroxide because the solution is usually colored by 
impurities dissolving as chlorides. In such a case hydrochloric 
acid is added to the peroxide and the liberated chlorine is distilled 
into potassium iodide solution. In the case of manganese perox- 
ide the reactions may be represented thus: 

Mn0 2 +4HCl-+MnCl2+2H 2 0+Cl 2 , 
C1 2 +2KI-+2KC1+I 2 . 




69. Calculate the equivalent weight of manganese dioxide and of available 
oxygen and find the weight of each that is equivalent to 1 cc of tenth-normal 
sodium thiosulphate solution. 

These reactions are analogous to those occurring with other 
peroxides and the determination of available oxygen in manga- 
nese dioxide is the most frequently made of all. The apparatus 
for carrying out the decomposition and distillation should have 
ground glass joints and should not allow contact of iodine or 
chlorine with organic matter. The modified apparatus of Bun- 
sen, Fig. 68, may be used. The receiver must be kept cold in 
order to avoid loss of iodine. 

Fig. 68. 

■Modified Bunsen's apparatus for the determination of available 

Determination. — Dry 2 to 4 gm of either commercial manganese 
dioxide or pyrolusite at 120° until the weight is constant. The sample 
already used for the determination of manganese may; be used for this 
determination and the results obtained by the two methods compared. 
Weigh enough sample to be equivalent to about 35 cc of the standard 
sodium thiosulphate solution and place in the flask of a Bunsen distil- 
ling apparatus or of some other suitable type. Place in the receiver 2 
gm of potassium or sodium iodide, dissolve this in water and dilute until 
the bend is just sealed when the apparatus is in the proper position. 
Immerse the receiver in ice water, then add to the flask containing the 
manganese dioxide 30 cc of concentrated hydrochloric acid and quickly 
nsert the stopper carrying the delivery tube. Warm the acid gently, 
distilling the chlorine into the potassium iodide solution. Raise the 


temperature gradually until the acid is boiling and boil for five minutes 
after action is completed. While the burner is still under the flask lower 
the receiver until the delivery tube is entirely out of it, then remove the 
burner. Remove the delivery tube from the flask and rinse it inside 
and outside, the water flowing back to the receiver. Rinse the whole 
iodine solution into an Erlenmeyer flask and titrate with sodium thio- 
sulphate solution. 

Make a blank test on the iodide used, as follows: 

Weigh out the same amount as was used in the determination o 
available oxygen, the weighing being accurate to centigrams. Dissolve 
in 100 cc of distilled water and add 5 cc of concentrated hydrochlori 
acid. If a yellow color appears, indicating the presence of free iodine 
titrate with sodium thiosulphate, using starch at the end. Deduct the 
thiosulphate used in the blank from that used in the determination o 
available oxygen and calculate the percent of available oxygen, als( 
the theoretical percent of manganese and of manganese dioxide, 
the sample is the same as that used for the direct determination o 
manganese and of available oxygen by potassium permanganate ai 
interesting comparison of results of different methods may be made 
although the calculation of available oxygen from the percent of man 
ganese may not check with the direct determination, for reasons already 

Sodium thiosulphate may be used to titrate the iodine producec 
by the action of almost any oxidizing agent upon a solution o 
potassium iodide and hydrochloric acid. Peroxides have alread; 
been discussed. Other substances that may be determined are 
free halogens (chlorine and bromine being allowed to displace 
iodine from potassium iodide), easily reducible oxyacids and their j 
salts, as the halogen oxyacids, nitrous acid and persulphuric 
acid, oxysalts of metals that exist in acid radicals, as dichromates, | 
chromates, permanganates and manganates, and salts of metals I 
that possess more than one valence, as iron, copper, mercury and 

While sodium thiosulphate may be used for the determinatioi 
of almost any oxidizing agent it is not necessarily true that tl 
provides the best method for all such materials. In many case 
other methods will be found to give better results or to be moi 
conveniently applied. 



70. Complete the following equations, balance and determine the equiva- 
lent weights of each of the oxidizing agents. 

Br 2 +KI— > 


KBr0 3 +KI+HCl— > 


K 2 Cr 2 7 +KI+HCl-> 

KMn0 4 +KI+HCl-> 


CuCl 2 +KX-^ 


The gravimetric determination of copper may be made by 
precipitating as cupric hydroxide, heating and weighing as 
cupric oxide, or by precipitating as cupric sulphide, heating with 
sulphur and hydrogen and weighing as cuprous sulphide. Both 
methods are difficult of execution and are subject to considerable 
errors. Electrolytic methods are more accurate and more easy 
of accomplishment.^Copper may be determined volumetrically 
by several methods, one of the best being Low's "iodide method." 1 
This method depends upon the insolubility of cuprous iodide and 
the instability of cupric iodide. If to a solution of a cupric 
salt, containing no highly ionized acid and no other oxidizing 
agent, potassium iodide is added there is an immediate precipita- 
tion of cuprous iodide with liberation of iodine : 

Cu(C 2 Hs02)2+2KI->CuI+2KC a H,0 2 +I. 

The iodine may be titrated by standard sodium thiosulphate 
solution and copper calculated. 


71. Calculate the equivalent weight of copper and the weight which is 
equivalent to 1 cc of tenth-normal sodium thiosulphate solution. 

72. What weight of a copper ore should be taken for analysis in order 
that 1 cc of fifth-normal thiosulphate solution should indicate 1 percent of 
copper in the ore? 

1 J. Am. Chem. Soc, 18, 458 (1896); 24, 1082 (1902). 
See also a comparison of methods by Fernekes and Koch: Ibid., 27, 
1224 (1905). 


If sodium thiosulphate solution is to be standardized against 
pure copper, the metal is dissolved in nitric acid, most of the 
nitrogen oxides are expelled by boiling and any remaining trace 
of nitrous acid is oxidized by bromine. The excess of nitric 
acid is then neutralized by ammonium hydroxide, acetic acic 
and potassium iodide are added and the free iodine titrated al 

If a copper ore or crude copper is to be analyzed all metals 
whose iodides are insoluble or whose salts will oxidize potassium 
iodide must first be removed. The addition of metallic alu- 
minium to the solution containing sulphuric acid will precipitate 
copper and leave in solution all other metals of higher decomposi- 
tion potentials as well as those soluble in sulphuric acid, providing 
that nitric acid be absent. The latter can be removed by evapo- 
rating with sulphuric acid. This treatment also precipitates 
lead as sulphate, which may be removed by filtration. After 
the copper is precipitated by aluminium it is removed by filtration 
washed, dissolved in nitric acid and determined as in the stand- 
ardization of sodium thiosulphate. 

During the process of filtration and washing, copper oxidizes 
and dissolves to some extent in the sulphuric acid. This woul< 
occasion a loss and to prevent this a solution of hydrogen sulphide 
is used for the wash liquid. Any small amount of copper tha 
might be dissolved is reprecipitated as cupric sulphide. If 
brown color appears in the filtrate below this is an indication o 
incomplete precipitation by the aluminium or of resolution during 
filtration. In either case the determination is spoiled unless 
the copper can be recovered by refiltration. 

Exercise: Standardization of Sodium Thiosulphate Solution. — The 
standard solution already prepared may be used and the copper equiva- 
lent calculated. In case it is desired to standardize against copper 
or a copper salt, proceed by one of the following methods: 

(a) Standardization against Metallic Copper of Known Purity. — Weigl 
sufficient copper to require about 35 cc of sodium thiosulphate solu- 
tion. Place in a 250 cc flask and dissolve by warming with 5 cc 
a mixture of equal volumes of concentrated nitric acid and watei 
Dilute to 25 cc and boil to expel nitrogen oxides. Add 5 cc of brc 
mine water and boil until all excess bromine is removed. Cool and adc 
strong ammonium hydroxide until a clear blue solution is obtained, thei 


I boil until copper hydroxide begins to precipitate. Acidify with acetic 
acid and boil, if necessary, to dissolve any precipitated cupric hydroxide. 
Cool, add 3 gm of potassium or sodium iodide and titrate the liberated 
iodine with sodium thiosulphate solution. Calculate the copper equiva- 
lent of the solution. 

(b) Standardization against a Copper Salt. — Weigh the proper amount 
of cupric sulphate of known purity, dissolve in 25 cc of water, make 
slightly basic with ammonium hydroxide and from this point proceed 
as in (a). 

Determination. — Dissolve 0.5 gm of ore in a covered casserole with 
10 cc of concentrated hydrochloric acid and 5 cc of concentrated nitric 
: acid, boiling if necessary to aid solution. If the sample is an alloy use 
| enough to contain 0.2 to 0.4 gm of copper and dissolve in 10 cc of nitric 
j acid, 1:1 (specific gravity 1.2). Add 7 cc of concentrated sulphuric 
I acid and evaporate until dense white fumes of sulphur trioxide appear. 
I Cool, add 25 cc of water and boil to dissolve the sulphates. Filter to 
; remove lead sulphate and gangue, allowing the filtrate to run into a 
I beaker. Wash the residue and paper and dilute the filtrate and washings 
! to 75 cc. 

Cut a strip of sheet aluminium about 2.5 cm wide and 14 cm long, 
i bend into a triangle and stand on its edge in the solution. Cover and 
i boil until all copper is precipitated and the solution is colorless or green 
n from ferrous sulphate. If this condition cannot be attained it is because 
t nitric acid was not completely removed when evaporating with sul- 
phuric acid. When all copper is precipitated, wash down the sides of 
I the beaker with a jet of hydrogen sulphide solution, pour the solution 
■ into a filter paper and filter quickly. 

Transfer the copper to the filter, washing the aluminium with hydro- 
gen sulphide solution while still in the beaker. Wash thoroughly with 
\\ hydrogen sulphide solution and then place a clean flask under the filter. 

I Add to the beaker containing the aluminium 6 cc of nitric acid, sp. gr. 
1 1.2. Boil shortly to dissolve adhering copper then pour the acid slowly 

over the filter to dissolve the copper on the paper. When all copper 
'seems to be dissolved pour over the paper 5 cc of bromine water. 
l\ Wash beaker and paper thoroughly with hot water then open the paper 
f and wash into the flask any particles of copper that have escaped the 
! action of the acid. Boil until all bromine is removed, add strong 
i | ammonium hydroxide until a deep blue is obtained, boil until copper 

I I hydroxide begins to precipitate and from this point proceed as in the 
; standardization of. sodium thiosulphate solution. Calculate the percent 
| of copper in the ore. 


Bleaching Powder 
When gaseous chlorine is passed over slaked lime it is absorbe 
with formation of an unstable compound that is easily made t 
yield chlorine under certain conditions and the compound pro 
vides a convenient means for storing and transporting chlorine t 
be used for bleaching, disinfecting, etc. This compound, know 
as " bleaching powder," is a double salt of calcium with hydro 
chloric and hypochlorous acids and may be represented by th 


formula Ca ( . When dissolved in water it is probably 

X C10 

ionized in the manner characteristic of both acids. When any 
acid, even carbonic acid, is added to bleaching powder chlorine 
is liberated: , 

CaCl.C10+H 2 C0 3 -»CaC0 3 +H 2 0+Cl 2 . 

This is due to the fact that when hydrochloric acid and hypochlo- 
rous acid come together, even in dilute solutions, they act upon 
each other with the formation of chlorine : 

HC1+HC10^H 2 0+C1 2 . 

Because of the easy decomposition of bleaching powder by car- 
bonic acid it rapidly deteriorates when exposed to air, chlorin 
escaping. Loss of efficiency also occurs through loss of oxygen 

2CaCl.C10-+2CaCl 2 +0 2 , 

and through a decomposition such that calcium chlorate 
formed : 

6CaCl.C10->Ca(C10 3 ) 2 +5CaCl 2 . 

The decompositions represented by the last two equations resu 
in the formation of chlorine compounds in which the chlorine 1 
not liberated upon acidification. A determination of total chk 
rine would therefore be of little value as an estimate of the usefu 
ness of bleaching powder. " Available chlorine" is bette 
determined by a volumetric process. For this purpose the acidi- 
fied solution may be treated with potassium iodide and the liber- 
ated iodine titrated with standard sodium thiosulphate solution, 
or the solution may be titrated directly by a standard solution of 
sodium arsenite. For the last titration the indicator is a paste of 
starch and potassium iodide used on a porcelain plate or absorbed 


by filter paper and dried. This method of reading the end point 
is inconvenient and the first method of titration is the better one. 

If calcium chlorate is present in bleaching powder and a strong 
acid is used for liberating the chlorine, the chlorate will be de- 
composed, though but slowly. This is because chloric acid is 
formed by the reaction of chlorate with added acid and chloric 
acid is slowly reduced by hydrochloric acid, liberating chlorine: 

HC10 3 +5HC1->3H 2 0+3C1 2 . 
During the titration the effect of these reactions is seen in an 
uncertain end point. As sodium thiosulphate is added the blue 
color of starch iodide disappears and then returns and deepens. 
As the addition of thiosulphate is continued the blue finally 
permanently disappears, but this end point does not represent 
the titration of chlorine really available in bleaching processes 
because that which comes from calcium chlorate is evolved too 
slowly to be of much use. This interference with the titration 
may be almost entirely averted by using a weak acid instead of 
a strong one for the decomposition of the chlorohypochlorite. 
The concentration of chloric acid does not then become suffi- 
ciently large to cause more than slight oxidation of hydriodic acid. 
The most suitable acid for the purpose is acetic acid. 

Determination. — If bleaching powder were pure calcium chloro- 
hypochlorite, CaCl.CIO, it would contain about 56 percent of chlorine. 
For reasons already discussed the amount of available chlorine is much 
less than this and in the average commercial product it is not much more 
than 25 percent. Upon this basis calculate the weight that should be 
(used when 50/1000 of the weighed sample is to be taken for the final 
titration. Weigh from a closed weighing bottle into a 1000 cc graduated 
flask. Fill to the mark with water and agitate until the powder is thor- 
oughly disintegrated and all soluble matter is in solution. Measure 
50 cc portions into flasks, add 5 gm of potassium iodide and 25 cc of 
10 percent acetic acid to each and titrate with sodium thiosulphate 
solution. Calculate the percent of available chlorine in the powder. 

Standard Iodine Solution. — Iodine- solutions do not maintain 
a constancy of oxidizing power and standard solutions must be 
restandardized frequently for accurate work. 

It has already been explained (page 261) that free iodine is 
usually dissolved with the aid of an iodide and that a molecular 
compound with a formula such as Nal 3 is present in such solu- 


tions. Two-thirds of the iodine in this compound (better repre 
sented as NaI.I 2 ) is so loosely bound that it behaves as fre 
iodine. The formula would indicate a necessary ratio of 150 : 254 
sodium iodide to iodine, or 166:254, potassium iodide to iodine 
However, in practice it is found necessary to use a much highei 
ratio (2 : 1) in order to dissolve the iodine readily and to preserve 
the solution in a fairly stable condition. The tenth-norma 
solution is convenient for the following determination. 

Arsenical Insecticides 

Two of the most important insecticides containing arsenic are 
London purple and Paris green. The former is a waste product 
of certain aniline dye industries and contains much dye in addi- 
tion to a fairly large quantity of arsenic. Paris green is a fairly 
definite compound of cupric arsenite and cupric acetate, repre- 
sented by the formula: Cu3(As03)2.Cu(C 2 H 3 02)2. This com- 
pound is decomposed by boiling with sodium hydroxide, pre- 
cipitating cuprous oxide and forming sodium arsenate and arsenite 
in solution. The formation of cuprous oxide is due to the reducing 
action of sodium arsenite, forming sodium arsenate. If the solu- 
tion is to be titrated for the determination of total arsenic this 
arsenate must first be reduced. For this purpose the solution is 
concentrated, then hydrochloric acid and potassium iodide 
are added and the resulting free iodine is removed by sodium 

Na 3 As04+2HCl+2KI±^Na 3 As03+2KCl+H 2 0+l2, 
2Na2S 1 0,+Ia-*Na 8 S 4 0e+2NaL 

The exact removal of iodine must be determined without the 
aid of starch. In strongly acid solutions starch is partly inverted, 
dextrine being one of the intermediate products and dextrine 
forms with iodine a deep red color which is not later removed and 
which interferes in the titration of iodine solution. 

The first equation above represents a reaction that can be 
quantitatively reversed at will. The complete reduction of 
pentavalent arsenic has just been accomplished in acid solution, 
one of the products (iodine) being removed. If the solution is 
now made basic, thus removing one of the products (hydro- 
chloric acid or hydriodic acid) of the reverse reaction and if 


standard iodine solution is added a quantitative oxidation of 
trivalent arsenic occurs. The addition of a strong base is not 
permissible because this will combine with iodine : 

2KOH + I 2 -+KIO + KI+ H 2 0. 
Neither is it possible to use a normal carbonate, for similar rea- 
sons. Alkali bicarbonates may be present and are used in prac- 
tice for neutralizing the acid. 

The method that was formerly given as "official" by the Asso- 
ciation of Official Agricultural Chemists 1 is based upon the prin- 
ciples just discussed. Several difficulties which are experienced 
in carrying out the method have led to a change to a distillation 
method as the official one. One of the principal difficulties is the 
formation of a yellow colloidal solution of arsenious iodide when 
potassium iodide and hydrochloric acid are added to reduce the 
arsenic solution. This color makes impossible the exact removal 
of iodine by sodium thiosulphate. If the analysis is performed 
carefully, as described below, this difficulty will disappear. 

Determination of Total Arsenic and of Copper in Paris Green. — To 

2 gm of Paris green in a 250-cc flask add about 100 cc of a 2-percent 
solution of sodium hydroxide. Boil until all of the green compound has 
been decomposed and only red cuprous oxide remains. Cool, filter into 
a 250-cc volumetric flask, washing the paper well, dilute to the mark and 
mix well. Reserve the cuprous oxide on the filter for the copper 

Measure two or three portions of 50 cc of the solution into 250- 
cc flasks and concentrate by boiling to about half the original volume. 
Cool to 60°, add 10 cc of concentrated hydrochloric acid and 1 gm 
of potassium iodide. Mix and allow to stand for about ten minutes. 
From a burette carefully add sodium thiosulphate solution until the 
iodine is all reduced. Starch should not be added but care should be 
exercised in reaching the end point. Allow to stand for 5 minutes longer 
and if iodine color reappears carefully add more thiosulphate solution. 
Immediately add, as rapidly as can be done without loss by effervescence, 
15 gm of sodium bicarbonate, free from lumps. Titrate at once with 
standard iodine solution, deferring the addition of starch until near the 
end point. Calculate the percent of total arsenic, expressed as arsenious 
oxide, in the Paris green. 

The residue of cuprous oxide is treated on the filter with 5 cc of nitric 
acid, specific gravity 1.2, the solution being caught in a 250 cc flask. 

l Bm. Chem., Bull. 107. 



Wash the paper well with hot water and proceed as directed for the 
standardization of sodium thiosulphate against metallic copper, page 

268, beginning with " Dilute to 25 cc and boil ." Calculate the 

percent of copper in the Paris green. The result may also be expressed 
as cupric oxide, if desired. 

If preferred, the solution of cuprous oxide in nitric acid may be 
diluted and electrolyzed, after boiling to expel nitrous acid. This 
determination is described on page 156. 

The present official method 1 for the determination of total 
arsenic in Paris green is based upon the volatility of arsenic 
trichloride with steam. The sample is dissolved in concentrated 
hydrochloric acid in a distilling flask. Cuprous chloride is added 
to reduce any pentavalent arsenic. 

The distillate containing the arsenous chloride and hydrochlo- 
ric acid is absorbed in cold water. The excess of acid is then 
neutralized with sodium hydroxide, sodium bicarbonate is added 
in excess and the arsenic is titrated with standard iodine solution. 
The reactions involved have already been discussed. 

Determination of Total Arsenic. Official Method.— Prepare the 
following standard solutions: 

(a) Arsenous Acid. — Dissolve 2 gm, accurately weighed, of pure 
arsenous oxide in a beaker by boiling with about 200 cc of water and 
10 cc of concentrated sulphuric acid. Cool the solution, transfer to a 
500 cc volumetric flask, dilute to the mark and mix well. Keep in a 
stoppered flask or bottle. 

N v. 

(b) Iodine Solution, 20 . — Mix by grinding in a porcelain mortar 6.35 

gm of pure iodine with 12.5 gm of potassium iodide. Dissolve in 
water, filter into a 1000 cc volumetric flask, dilute to the mark and mix 
well. Standardize against solution (a) as follows: 

Pipette 50 cc of the arsenous acid solution into a 1000 cc Erlenmeyer 
flask, add 400 cc of water, then gradually add 10 gm of sodium bicar- 
bonate. Mix and titrate at once with the iodine solution, adding 5 cc 
of starch solution when the slow disappearance of the iodine color indi- 
cates that the end point is nearly reached. Calculate the value of the 
iodine solution in terms of arsenous oxide, As 2 3 . 

For the determination of total arsenic in Paris green the apparatus 
shown in Fig. 69 is necessary. 

A is a distilling flask having a capacity of 250 cc and fitted with a 50 
cc dropping funnel. The capacities of the Erlenmeyer flasks B, C and 

1 J. Assoc. Off. Agr. Chem., Vol. II, No. 1, Pt. 2, p. 5. 



D are 500 cc, 1000 cc and 100 cc, respectively. B and C are surrounded 
by cracked ice and contain 40 cc and 100 cc, respectively, of water. D 
contains 50 cc of water which serves as a trap. The upright tube of 
flask B reaches to the bottom of the flask and acts as a safety valve, 
preventing liquid from drawing back from B when the distillation 
slackens or stops. 

Place 5 gm of cuprous chloride in the distilling flask. Calculate the 
theoretical weight of Paris green that would be equivalent to 250 cc of 
the standard iodine solution. Weigh this amount and rinse into the 
distilling flask by means of 100 cc of concentrated hydrochloric acid. 
Distill until only about 40 cc of liquid remains in the distilling flask, 

Apparatus for arsenic distillation. 

then add 50 cc of concentrated hydrochloric acid through the dropping 
funnel and distill. Continue this process until 200 cc of distillate has 
been obtained. Stop the distillation and rinse down the condenser and 
all connecting tubes into the flasks. Rinse the contents of all three of 
the receiving flasks into a 1000 cc< volumetric flask. Allow the solution 
to attain the temperature of the room then dilute to the mark and mix 
well. Measure 400 cc of this solution into a 1000 cc Erlenmeyer flask, 
add two or three drops of phenolphthalein solution and nearly neu- 
tralize with a saturated solution of sodium hydroxide, leaving the solu- 
tion slightly acid. Add 10 gm of sodium bicarbonate and titrate the 
arsenic with standard iodine solution as directed for the standardization 
of the solution. Calculate the percent of total arsenic in the sample, 
expressing as arsenous oxide. 



The completion of the reactions of neutralization depends 
upon the small ionization of one of the products (water). The 
completion of reactions of oxidation and reduction depends upon 
the relative potentials of oxidizing and reducing agents. Certain 
other reactions are made the basis of volumetric determinations, 
completed because of the formation of a precipitate. In some 
cases an indicator is added while in others the cessation of pre- 
cipitation with further addition of standard solution is the 


An example of titration without an added indicator is to be 
found in Gay-Lussac's 1 method for silver. This method is one 
of the oldest of those analytical methods that have survived to the 
present day and, while it is not now extensively used because it 
is somewhat troublesome in the matter of execution, it is one of 
the most exact of all known volumetric processes. It depends 
upon the titration of the solution of a silver salt by a standard 
solution of sodium chloride. The very small solubility of silver 
chloride renders the reaction practically complete. The con- 
verse of this method may be used for the determination of 
chlorine, bromine, or iodine in soluble halides. 

Exercise : Preparation of Standard Solutions. — Calculate the weight 
of pure sodium chloride that is equivalent to 5 gm of silver, weigh this 
quantity, dissolve in distilled water and dilute to 1000 cc in a volu- 
metric flask. Make a second solution by diluting 100 cc of this solu- 
tion to 1000 cc. Record the silver equivalent of 1 cc of each of these 

Determination. — Silver may be determined in any alloy that contains 
no other metal forming insoluble chlorides but the approximate percent 

1 Instruction sur 1' essai des matieres d' argent par la voie humide. 
Paris, 1832. 



of silver should be known. A silver coin may be used. United States 
silver coinage contains approximately 90 percent of silver. Weigh 
enough of the alloy to give 0.5 gm of silver, place in a 250 cc flask having 
a ground glass stopper and dissolve in 10 cc of a mixture of equal volumes 
of water and concentrated nitric acid. Both water and acid must be 
tested and found free from chlorine. Boil to expel oxides of nitrogen, 
assisting this action by drawing air through the flask by means of a 
filter pump. Add to the solution in the flask exactly 99 cc of the more 
concentrated standard salt solution, stopper and shake until the pre- 
cipitated silver chloride flocculates and settles readily. Add from a 
second burette the more dilute standard solution, 0.5 cc at a time, 
allowing the solution to run down the sides of the flask and observing 
whether turbidity is produced. Shake the flask if more silver chloride is 
formed and continue the addition of the dilute standard solution until 
the last 0.5 cc fails to produce a visible precipitate in the clear, superna- 
tant liquid. Do not use the last 0.5 cc in the calculation. 

It may sometimes happen that the percent of silver in the alloy is not 
known with sufficient accuracy and either too much or too little of the 
more concentrated solution is used. In the first case the first addition 
of the dilute solution fails to produce a precipitate while in the second 
case an unduly large quantity of the dilute solution is required to reach 
the end point. In either case the determination should be begun again, 
the proper alteration being made in either the weight of sample taken 
or the volume of concentrated standard solution. From the results 
of the titration calculate the percent of silver in the alloy. 

In the determination of silver by the method of Volhard 1 an 
inorganic indicator is added to the solution. The silver should 
be in the form of nitrate, a solution of a ferric salt, acidified 
to suppress hydrolysis, is added and the silver is titrated by a 
standard solution of potassium thiocyanate or ammonium thio- 
cyanate. Silver is precipitated as silver thiocyanate: 

AgN0 3 +KCNS-+AgCNS+KN0 3 . 

When all of the silver is removed from the solution an additional 
drop of the standard solution of thiocyanate produces the red 
color of soluble ferric thiocyanate: 

Fe(NOa)«+3KCNS->Fe(CNS) 8 +KN0 8 . 

Mercury thiocyanate is insoluble in dilute nitric acid and 
mercury must therefore be absent. The color of salts of copper, 

1 J. prakt. Chem., [2] 9, 217 (1874). 


nickel and cobalt obscures the end point and these metals, 
should be absent although as much as 60 percent of copper may 
be present. 

The converse of this method may be used for the determina 
tion of the thiocyanate radical. 

Exercise : Preparation of Solutions. — Make a solution of silver nitrate 
1 cc of which contains 0.005 gm of silver. Standardize gravimetric 
ally by precipitating and weighing silver chloride, or by Gay-Lussac' 
volumetric method. 

Make 1500 cc of a solution of potassium thiocyanate or ammonium 
thiocyanate by weighing 2 percent more than the calculated quantity 
of salt required to make 1 cc equivalent to 0.005 gm of silver. 

Make 100 cc of a solution (saturated without heating) of ferric 
ammonium sulphate, adding enough nitric acid to remove turbidity 
and to cause the red color to give place to pale yellow. 

Standardize the thiocyanate solution as follows: Measure 35 cc 
of the silver nitrate solution into a beaker or Erlenmeyer flask, dilute 
to about 75 cc, add 1 cc of ferric ammonium sulphate solution and 
titrate with the thiocyanate solution until a permanent red tint is 

Determination. — Weigh not more than 0.25 gm of a silver alloy con- 
taining no mercury, nickel or cobalt and not more than 60 percent of cop- 
per and place in a 250 cc flask. Dissolve in 10 cc of a mixture of equal 
volumes of concentrated nitric acid and water, boiling to expel oxides 
of nitrogen. Cool, dilute to about 75 cc and titrate exactly as in th 
standardization of the thiocyanate solution. Calculate the percent o 
silver in the alloy. 

Halogens and the Cyanide Radical 
Volhard's method also applies to the determination of th 
halogen hydracids and cyanogen. A measured excess of stand 
ard silver nitrate solution is added, precipitating all of the 
chlorine, bromine, iodine or cyanogen. The excess of silver 
nitrate is determined by titration by standard thiocyanate 
solution by the method already described. In the original 
method the precipitated silver halide was not removed by filtra- 
tion before titration of the excess of silver. Rosanoff and Hill 
have shown 1 that the silver chloride reacts with the red soluble 
ferric thiocyanate, which is produced at the end point, as follows : 
!J. Am. Chem. Soc, 29, 269 (1907). 


This occurs to an appreciable extent, even though the solu- 
bility of silver chloride is less than that of silver thiocyanate. 
Rosanoff and Hill found that as much as 43 percent of ammonium 
thiocyanate is changed in two minutes by reaction with silver 
chloride. It is therefore necessary to remove the precipitate by 
filtration before the final titration. 

Determination, — Use the standard thiocyanate and silver nitrate 
solutions prepared for the preceding exercise. Weigh" enough of a 
soluble chloride, bromide, iodide or cyanide to be equivalent to about 
40 cc of the silver nitrate solution. Dissolve in a small amount of 
water, acidify with nitric acid and add 50 cc of the standard solution 
of silver nitrate. Filter and wash thoroughly and titrate the excess of 
silver nitrate by standard thiocyanate solution. Calculate the percent 
of halogen or cyanogen in the sample. 

A method for the direct titration of the halogens by standard 
silver nitrate solution is described on page 397 in the discussion 
of water analysis. 


Ferrocyanide Method. — The ferrocyanide titration of zinc has 
been practised for a long time and many modifications 1 of the 
details of the method have been published. Concerning the 
accuracy of this method there has been considerable controversy, 
particularly as it applies to the determination of zinc in ores. 
However, most of the errors have been traced to methods used 
in separating zinc from interfering metals, preceding the titra- 
tion. The modified Waring method 2 is described below. 

Most of the important zinc ores are soluble in acids, although 
the aluminates require fusion with potassium pyrosulphate. In 
the acid solution silica is rendered insoluble by evaporation, as 
otherwise zinc silicate might precipitate. After evaporation 
with sulphuric acid to expel nitric acid the solution is boiled 
with a piece of aluminium, which precipitates lead, copper, cad- 
mium and bismuth, or such of these as are present, these metals 
all lying below aluminium in the electrochemical series. Iron 
is not precipitated but is reduced to the ferrous condition. In 

l J. Ind. Eng. Chem.,-4, 468 (1912). 

2 J. Am. Chem. Soc, 26, 4 (1904) and, 29, 262 (1907). 


the clear solution sulphuric acid is neutralized by sodium bicar- 
bonate and then formic acid is added in slight excess. In pres- 
ence of this weakly ionized acid zinc is precipitated by hydrogen 

Zinc sulphide, so separated from interfering metals which 
would form insoluble ferrocyanides, is redissolved in hydrochloric 
acid and titrated with a standard solution of potassium ferro 
cyanide, the .following reaction first taking place : 

2ZnCl 2 +K 4 Fe(CN) 6 ->4KCl+Zn 2 Fe(CN) 6 . 

Zinc ferrocyanide, so formed, does not flocculate readily but a 
more ferrocyanide is added to the hot solution a potassium zinc 
ferrocyanide is precipitated: 

3Zn 2 Fe(CN) 6 +K 4 Fe(CN) 6 -^2K 2 Zn 3 [Fe(CN) 6 ] 2 . (2) 

Equations (1) and (2) may thus be combined: 

3ZnCl 2 +2K 4 Fe(CN) 6 -^6KCl+K 2 Zn 3 [Fe(CN) 6 ] 2 . (3) 

As indicator, a solution of uranium acetate or nitrate or of 
ammonium molybdate is used on an outside test plate. The 
yellow or brown color that is produced by a slight excess of 
ferrocyanide with ammonium molybdate is of unknown composi- 
tion. When uranium salts are used a brown ferrocyanide o: 
uranium is formed. 

The ferrocyanide titration of zinc is easily performed and i 
fairly accurate if the details of the standardization of the solutio 
and of the determination are watched closely. All things con 
sidered the gravimetric method, weighing zinc as pyrophosphate, 
is to be preferred. This method is described on page 507. 

Determination in Zinc in Ores. — Weigh 0.5 gm of powdered ore and 
brush into a 250-cc casserole. Add 5 cc each of concentrated nitric and 
hydrochloric acids, cover and heat to decompose the ore. Finally boil 
to expel all red oxides of nitrogen, then remove the cover and rinse this 
and the sides of the casserole. Cool and add 5 cc of concentrated sul- 
phuric acid. Evaporate until sulphur trioxide fumes are freely evolved, 
the casserole being held in the hand and agitated during the process of 

Cool, add 50 cc of water and warm until only insoluble gangue 
remains. Bend a strip of heavy aluminium foil, 2.5 cm wide and 14 
cm long, into a triangle and place in the solution. Boil for 10 minutes, 
which should remove all color except a faint green due to ferrous sul- 


phate. Filter through a paper containing a piece of aluminium, into a 
500-cc flask containing a rod or strip of the same metal, and wash with 
hot water. 

Add a drop of methyl orange and neutralize with a solution of sodium 
bicarbonate (about 5 percent). Barely restore the pink color by 
adding a 20-percent solution of formic acid, a drop at a time, then add 
an excess of 5 drops. Dilute to 200 cc and then add 4 gm of ammonium 
thiocyanate, unless iron is known to be present in only a small amount. 

Remove the rod of aluminium and insert into the neck of the flask a 
rubber stopper, through the single hole of which passes a tube leading 
to the bottom of the flask. Connect the tube with a Kipp generator 
for hydrogen sulphide and heat the solution to boiling. With the 
stopper loosely fitted, pass a stream of gas through the boiling solution 
until most of the zinc has been precipitated, then push the stopper in 
so that the pressure of the gas from the generator will aid absorption. 

When the white zinc sulphide settles readily remove the stopper 
and rinse. Filter through paper and wash with hot water, but without 
attempting to remove adhering precipitate from the flask. Finally 
place the paper and precipitate in this flask and add 10 cc of concentrated 
hydrochloric acid and 50 cc of water, allowing the acid to act upon any 

1 precipitate adhering to the tube. Warm until all of the sulphide in the 
i flask is in solution and boil to remove hydrogen sulphide. Drop in a 
bit of litmus paper and neutralize the acid with ammonium hydroxide, 
then add 3 cc excess of concentrated hydrochloric acid. 

Rinse the solution into a 500-cc beaker, dilute to about 250 cc, 
heat nearly to boiling and titrate with standard potassium ferro- 
cyanide, using a drop of nearly saturated uranium nitrate or acetate 
solution or of a 2-percent ammonium molybdate solution, on a white 
test plate as indicator. Near the last the solution is stirred and tested 
after each drop of standard solution is added. When a yellow or brown 
color finally appears it will be found that two or three of the tests 
immediately preceding this one will develop a color after standing for a 
few minutes. When this is the case the burette reading corresponding 
to the earliest positive test for ferrocyanide is taken as the amount of 
j solution equivalent to the zinc. 

Calculate the percent of zinc in the ore. 

The solution of potassium ferrocyanide should be made so that 
1 cc is equivalent to 0.005 gm of zinc. It must be standardized against 
zinc, zinc oxide or zinc sulphate of known purity, following exactly 
the same method that has been outlined for the ore, beginning with the 
dissolving of zinc sulphide. Record the concentration of the solution 
in terms of zinc equivalent to 1 cc of ferrocyanide solution. 



In most of the exercises in the preceding portion of this book 
determinations have been made of single constituents of various 
substances and interfering substances have usually been either 
absent or capable of being removed with comparative ease. 
Standard methods have been employed and attention has been 
centered upon the chemical principles underlying the method and 
the proper manipulation. In the pages that follow the student 
will become acquainted with the application of these and other 
determinations to the testing and analysis of some materials 
which are of importance to our industrial life. Such materials 
are often quite complicated in composition and most varied 
procedures are necessary in a determination of their industrial 
value. The chemist will then find it necessary to have at his 
command all of the chemical principles and methods of analysis 
that have already been learned and to apply these to an intelli- 
gent study of the material under examination. He will also be 
prepared to take up other methods of testing. , Some of the tests 
are purely physical but they are, in industrial practice, applied 
by the chemist and not by the physicist because the former is 
usually engaged in the analysis of the same or similar materials. 
Other analytical determinations are empirical, rather than exact, 
in their nature but must be made with the same degree of care 
and attention as the determinations involving definite elements 
or compounds. 




Carbonate Minerals 

The most important and abundant of the carbonate minerals 
are the calcites and the dolomites. The calcites consist essen- 
tially of calcium carbonate and the dolomites of double carbon- 
ates of calcium and magnesium but these compounds seldom or 
never occur in a pure state in nature. Iceland spar is one 'of the 
best-known examples of a nearly pure natural variety of calcium 
carbonate, yet in many samples of Iceland spar substances other 
than calcium carbonate occur in appreciable amounts. For pur- 
poses of geological investigation there is usually required a com- 
plete analysis with the utmost accuracy that can be attained 
For technical purposes this is not the case. The mineral is to be 
used for a given industrial purpose where the essential constituen 
is the one of chief importance and where impurities are importan 
only to the extent that they may reduce the percentage of th< 
essential constituent or that they exert an undesirable influence 
in the industrial operation to which the mineral is to b( 

The particular application of the mineral to the industrial proc- 
ess will determine which impurities are of considerable and which 
are of minor importance. Those of minor importance are fre- 
quently grouped, with no attempt at separation, into certain arbi- 
trary classes. For example a limestone may be used as a source 
of quick lime, as a flux in iron smelting, as a paving material, as 
a building stone, as a raw material for hydraulic cements, or for 
any one of a variety of other purposes. All limestone contains 
more or less of material insoluble in acids, consisting chiefly of 
various silicates and of quartz. For the first purpose named 
these substances are important only as they act as diluents of the 
essential calcium carbonate, unless they occur in relatively large 



quantities. For such a purpose the analysis would be so made as 
to include all such materials as simply " insoluble" or "silicious 
matter," no separation of the components being made. If the 
limestone were to be used as a flux in the smelting of iron ore, 
the nature of this insoluble material should be more exactly 
determined, since it not only reduces the actual percent of cal- 
cium carbonate but also may contain substances that themselves 
require a flux or that may even add very objectionable impurities, 
such as sulphur or phosphorus, to the iron itself. For paving or 
building material the physical properties of the mineral are of 
great importance and the chemical analysis might be considerably 

As another example of such empiricism in analysis, may be men- 
tioned the usual report on calcium. This element is usually pre- 
cipitated as the oxalate. It will readily be understood, however, 
that if barium or strontium is present and not previously sepa- 
rated it will also precipitate and will be included in the finally 
weighed oxide. Unless it is known that barium or strontium is 
present in more than very small amounts the percent of " cal- 
cium" alone is made a part of the report for technical purposes, 
strontium or barium serving the same purpose as does calcium. 
This, obviously, involves a slight error, not only in the naming of 
the element but in the percent as well, because the factors for 
these three metals in their oxides are all different. For exact 
scientific purposes the separation and determination of all ele- 
ments or radicals may be necessary while for technical purposes 
the analysis will be ordered according to the use to which the sub- 
stance is applied. This is an example of the so-called "proxi- 
mate" analysis, as distinguished from the "ultimate" analysis. 
It is important to note that the term " proximate " does not imply 
carelessness in working or neglect of sources of error. It should 
not even convey the idea of inexact figures, but merely grouping 
together of more than one substance to be reported by one 
generic term, as, for example, " insoluble matter " above. The 
proximate analysis of coal will include the determination of 
percents of "volatile combustible matter," "fixed carbon," 
"ash," and "moisture," yet each one of these terms covers many 
substances which are all determined together with no attempt at 
a separation into the ultimate constituents, simply because the 


figures so determined serve a useful purpose in fixing a valuation 
on the coal. 

It was formerly the custom to report the analysis of acids, bases 
and salts, not as radicals but as anhydrides. Calcium carbonate 
would be reported as calcium oxide and carbon dioxide, sulphuric 
acid as water and sulphur trioxide, etc. This custom has now 
largely fallen into disuse in most lines of analytical chemistry but 
has been retained in the analysis of minerals. 

The ultimate analysis of carbonate minerals is exhaustively 
and scientifically treated in a bulletin of the U. S. Geological 
Survey and only reference to this will be made. 1 The exercise 
to follow will deal with the analysis made with industrial ends 
in view. This exercise will be the student's introduction to 
separations in quantitative analysis. Heretofore the work with 
the solution has terminated with the filtration and the removal of 
the precipitate. The filtrate could contain nothing but impuri- 
ties and by-products of the reaction and therefore could be of 
no further importance to the analyst. In the next and in many 
later exercises the filtrate must be carefully conserved because 
it contains substances still to be determined. The quantity of 
wash liquid must be made as small as possible, not merely to 
minimize its solvent action upon the precipitate but also because 
the washings must be added to the chief filtrate and the total 
bulk must not be excessive for subsequent operations. Even 
with the exercise of great care in this regard an occasional con- 
centration of the solution by evaporation is necessary in order 
to reduce its volume to a workable value. 

Another point that will here appear for the first time is that 
many of the elements or radicals that must be separated and 
determined are present in the mineral in relatively small quanti- 
ties. The student has been accustomed to a rapid appearance 
of a considerable quantity of precipitate and if this should not 
appear when the appropriate reagent is added he is likely to 
conclude that none of the substance is present and to pass to the 
next determination. None of the constituents ordinarily present 
in a given mineral or other complex material should be assumed 
to be absent. The reagent should be added and sufficient time 
allowed for the precipitation to become completed, remembering 

1 U. S. Geol. Surv., Bull. 422, by W. F. Hillebrand. 


that precipitation starts and proceeds slowly from very dilute 
solutions. Even when no precipitate is finally visible it is the 
safest plan to filter, wash and ignite the paper in a weighed cruci- 
ble, when a small amount of precipitate will often be detected, 
when otherwise it would have been weighed with the next 
precipitate to be produced. 

Analysis of Carbonate Mineral. — Read again the discussion of sam- 
pling on page 9 and apply this to the preparation of a sample of limestone, 
dolomite or other carbonate mineral, for analysis. The small sample 
finally used should weigh about 10 gm and should pass a sieve having 
100 meshes in each linear inch. 

Carbon Dioxide. — Determine carbon dioxide exactly as directed on 
page 134, noting that if dolomite is under investigation solution will 
proceed rather slowly while the acid is cold. It is obvious that hydro- 
chloric acid must be used since, considerable quantities of calcium are 
present and the solubility of calcium sulphate is not large. 

Silica or Insoluble Matter. — The residue from the carbon dioxide 
determination may be used for this determination but it is better to use 
new samples. Weigh duplicate portions of 0.5 gm each into casseroles. 
Dissolve in 5 cc of concentrated hydrochloric acid, covering the cas- 
serole while the mineral is dissolving. Rinse down <the cover glass and 
the sides of the casserole and evaporate to dryness on the steam bath. 
Heat the dry material at dull redness until the organic matter has been 
oxidized, leaving the residue white or reddish brown from iron oxide. 
Cool, add 5 cc of concentrated hydrochloric acid and warm until all 
soluble matter has passed into solution. Dilute to about 50 cc with 
hot water, boil and filter into a Pyrex beaker, using particular care in 
removing all of the residue to the filter paper, since the white casserole 
makes this process somewhat uncertain. Wash the residue free from 
chlorides with hot water, collecting the washings and filtrate in the same 
beaker. The total volume, after filtration and washing, should not be 
greater than 100 cc. If it is greater than this amount it should be con- 
centrated by evaporation. Place the paper and residue in a weighed 
platinum crucible, burn the paper and then ignite for 10 minutes over 
the blast lamp. Report the percent of silicious matter. If silicious 
matter amounts to more than 0.5 percent it should be separated into 
its constituents. In this case add to the residue in the crucible 2 gm of 
sodium carbonate and fuse over the blast lamp until the silicate is com- 
pletely decomposed, as shown by the cessation of effervescence. Cool, 
place the crucible in a casserole containing 50 cc of water and warm until 
the material is completely dissolved or disintegrated. Carefully add 


to the covered casserole concentrated hydrochloric acid until efferves- 
cence no longer occurs, then remove the crucible and rinse. Evaporate 
the solution and heat at about 120° for 10 minutes. Add 5 cc of con- 
centrated hydrochloric acid and warm until soluble matter is dissolved, 
then dilute with 5 cc of water and filter on an extracted paper. Wash 
with hot water until chlorides are completely removed, adding the 
nitrate and washings to the' original solution of the mineral. Ignite the 
residue and paper in a weighed platinum crucible, weigh and report as 

Iron and Aluminium. — If the solution has a volume greater than 100 
cc it should be evaporated to concentrate to about this volume. Drop 
into the solution a very small bit of litmus paper and then add dilute 
ammonium hydroxide, stirring, until the solution is distinctly basic, 
avoiding undue excess of ammonium hydroxide. Boil for 5 minutes 
or until the odor of ammonia is faint. Filter through an extracted 
paper and wash until free from chlorides, adding the washings to the 
filtrate. Remove the paper from the funnel, fold and burn in a weighed 
porcelain crucible. Burn the paper at a low temperature in presence 
of an excess of air, inclining the crucible to facilitate oxidation. Weigh, 
and if the amount of the oxides is not greater than 0.5 percent report as 
aluminium oxide and iron oxide. If more than this amount is present a 
separation is usually made. In this case add' 1 gm of potassium acid 
sulphate to the crucible containing the oxides of iron and aluminium. 
Fuse at a relatively low temperature until violent effervescence has 
ceased then heat to redness until the oxides have been completely 
dissolved. Cool the crucible and dissolve the mass in hot water. 
Reduce the iron and titrate with standard potassium dichromate or 
potassium permanganate according to the methods already learned. 

Manganese. — Add bromine water to the nitrate and washings from 
iron and aluminium until a yellow color is produced, then boil. If 
manganese is present it will precipitate as brown manganese dioxide. 
The quantity is usually small but it must not be disregarded. Filter, 
wash free from chlorides and ignite. Weigh the oxide Mn.304 and 
calculate as MnC>2, assuming that the manganese was originally pres- 
ent in this form. (This is an arbitrary assumption because manganous 
carbonate is of common occurrence.) 

Calcium. — Acidify the solution with hydrochloric acid and concen- 
trate to about 100 cc, boiling until all bromine is removed. Add a 
bit of litmus paper, then ammonium hydroxide until basic. Heat to 
boiling and add, drop by drop with stirring, 10 cc of a saturated solu- 
tion of ammonium oxalate, or enough to precipitate all of the calcium. 
Determine the calcium as directed on page 81 or 253, with the following 
addition, designed to complete the separation of magnesium: Filter 


Ithe precipitate of calcium oxalate and wash once with hot water but 
without making any attempt to transfer completely to the filter. Place 
ithe beaker under the paper and add to the precipitate on the filter 
! enough concentrated hydrochloric acid to dissolve all calcium oxalate. 
2 cc should be sufficient. Wash the paper thoroughly with hot water, 
j precipitate the calcium once more and determine as already directed. 
; Add the filtrate from the second filtration to that from the first. Calcu- 
late the percent of calcium oxide in the sample. 

Magnesium. — Determine the magnesium in the nitrate and washings 
as was done in the case of a magnesium salt. The determination is 
discussed on page 112. Report the percent of magnesium oxide. 

Sodium and Potassium. — These elements are not often present in 
(more than traces in the carbonate minerals and their determination is 
not often required for industrial purposes. If such a determination is 
to be made, a new sample of mineral should be used. Follow the pro- 
cedure directed under the analysis of silicate minerals, page 294. 

If the analysis has been made with care and every substance present 
|has been determined the sum of the percents of all of the constituents 
of the mineral should be 100. This will serve as a check upon the 
accuracy of the work but the sum will rarely be exactly 100. The 
omission of the determination of other substances present in small 
quantity will give rise to a negative error, while imperfect washing and 
other experimental errors will summate as a positive error, so that the 
sum of all percents may be either greater or less than 100. The student 
should be able to work so that the sum of all the errors should not be 
greater than ± 1 percent. 

Silicate Minerals 

Silica, as a constituent of various simple and complex silicates, 
is distributed widely in the earth's crust. Associated with other 
j minerals or in a nearly pure form silica itself is also to be found. 
These minerals are only slightly soluble in acids or bases and their 
[analysis requires a preliminary decomposition by some agent 
which will react at elevated temperatures. When silicon dioxide 
| or a silicate is heated with an alkali carbonate to the point of 
[fusion the corresponding alkali silicate is produced, carbon diox- 
ide is evolved and whatever heavy metals may have been origi- 
inally present as silicates are left in the form of oxides. The alkali 
I silicates are soluble in water (as colloids) and most of the metallic 
| oxides so produced are soluble in hydrochloric acid. The pre- 



viously insoluble mineral is, by this means, obtained in solution, 
and the ordinary analytical processes will henceforth apply. 

All of * the natural silicates may be regarded as being derived 
from silicon dioxide, the anhydride A of the various silicic acids. 
These acids are not known in the free state but their existence ma^ 
be supposed from the composition of the salts. Thus H 2 Si0 3 
H 2 O.Si0 2 , H 4 Si0 4 = 2H 2 O.Si0 2 , H 6 Si 2 7 = 3H 2 0.2Si0 2 , H 4 Si 2 0, 
2H 2 0.2Si0 2 ,H 4 Si 3 8 = 2H 2 0.3Si0 2 ,H 2 Si 2 O c =H 2 0.2Si0 2 . These 
may be taken as the acids from which various natural 
silicates are derived. Kaolin, the essential constituent of the 
various impure clays, is Al 2 Si 2 07.2H 2 0, a salt of H 6 Si 2 7 . The 
felspars are double silicates derived from the acid H 4 Si 3 8 . As j 
examples of the felspars may be mentioned orthoclase, KAlSi 3 8 , 
and albite, NaAlSi s 8 . 

Fusion of orthoclase with sodium carbonate causes reactions 
which may be simply represented thus: 

2KAlSi 3 8 +6Na 2 C0 3 -^K 2 Si0 3 +5Na 2 Si0 3 +2NaA10 2 +6C0 2 . 

The completion of the reaction is assured by the presence of a 
considerable excess of sodium carbonate. According to the 
terms of the mass law the reaction should be completed by simply 
heating at a sufficiently high temperature to decompose com- 
pletely any carbonates that may be formed by the decomposition. 
The mass of silicates resulting from the fusion may be decom- 
posed by hydrochloric acid but if this is not preceded by disin- 
tegration and solution of the water soluble parts by hot water 
the result of such treatment will be to form a protective coating 
upon lumps of the fusion, thus retarding the action of the acid. 
Upon addition of hydrochloric acid to the mixture of substances 
after treatment with water, oxides of earth and alkaline earth 
metals form soluble chlorides, while the alkali silicates are de- 
composed with formation of alkali chlorides and silicic acid. 
The first separation occurs in the removal of the silicic acid which 
must first be converted into the less soluble silicon dioxide. This 
conversion is partly, but imperfectly, accomplished by evapora- 
tion to dryness and heating to about 120°. Rehydration readily 
occurs and the silica partly redissolves because of its marked tend- 
ency toward the formation of hydrosols. This tendency is 
diminished by long heating at high temperatures since such treat- 
ment results in incipient fusion and change into the irreversible 


I colloid, silicic acid, the production of the latter being promoted by 
| the presence of strong acids. It is practically impossible to com- 
41 pletely separate silica by one evaporation and filtration, a small 
i\ proportion invariably returning to the solution. By evaporating 
jj the filtrate, heating, and again filtering, all but a trace of silica 
i may be removed. The residue of silica is never pure but con- 
ji tains small amounts of oxides of iron, aluminium and calcium. 
j! In order to correct the error arising from this cause the precipitate 

II is treated by hydrofluoric acid, which converts silica into the 
gaseous silicon tetrafluoride. After volatilization of this and of 
the hydrofluoric acid the residue is -weighed and its loss reported 
as silica. 

After the separation of silica the metals will be determined by 
the usual methods, such as those used in the analysis of carbon- 
ate minerals. It will readily be seen that, after the fusion of the 
silicate with sodium carbonate or potassium carbonate, a deter- 
mination of the alkali metals in this portion of the sample will be 
without significance. Other methods must be employed for 
decomposing the silicate. Two such methods are in general use. 

J. L. Smith Method. — The method of J. Lawrence Smith 1 
depends upon the action of calcium chloride upon silicates at 
about 800°, resulting in the formation of alkali chlorides and sili- 
cates of calcium and other metals. The sample is intimately 
mixed with ammonium chloride and precipitated calcium carbon- 
ate, and gently heated. The reaction that occurs might be repre- 
sented as follows, assuming orthoclase to be the silicate. 

2KAlSi 3 08+6CaC03+2NH 4 Cl->2KCl+Al 2 03+6CaSi03+H 2 

+2NH 3 H-6C0 2 . 

This can be only an approximate representation of what has 
really happened during the heating. Ammonium chloride dis- 
sociates at about 450° into ammonia and hydrochloric acid: 


The ammonia escapes while the hydrochloric acid combines with 
calcium carbonate: 

CaC0 3 +2HCl-*CaCl 2 +H 2 0+C0 2 . 

^m. J. Sci., [3] 1,269 (1871). 


That the decomposition of silicates is, in a large measure, due t 
calcium chloride is undoubtedly true. That calcium carbonate 
as such, also plays an important part in the reactions would b 
inferred from the above interpretation of the reaction, since in 
this reaction only one-third of the calcium carbonate could form 
the chloride. The significance of this equation is lessened by 
the fact that it is a representation of a series of reactions tha 
cannot well be tested. The silicate resulting from the decomposi 
tion is probably not calcium metasilicate alone but is much mor 
complex than this. 

After the decomposition is- complete the mass is treated wit 
hot water which dissolves chlorides of sodium, potassium an 
calcium, as well as those of other metals present. After filtra 
tion the most of the calcium, iron and aluminium is precipitate 
by ammonium carbonate and ammonium oxalate; the solution i 
then evaporated in a platinum dish and heated to expel am 
monium salts. If desired, sulphuric acid may be added t 
convert sodium, potassium and ammonium salts into sulphate 
thus providing less liability of loss of sodium and potassium durin 
the heating. In this case the Gladding modification of the Lind 
method must be used or else the sulphates must be converte 
into chlorides by precipitating with barium chloride. If su 
phuric acid is not added the volatilization of ammonium salt 
must be conducted with greater care, since the chlorides of sodiu 
and potassium are appreciably volatile at a temperature of bright 
redness. Read again the discussion of the determination of 
sodium and potassium on page 96. 

Hydrofluoric Acid Method. — Another method for treating 
silicates without the addition of sodium or potassium carbonate 
is that of decomposing by means of hydrofluoric acid. The finely 
powdered silicate is moistened with concentrated sulphuric acid 
and then hydrofluoric acid is added. The silica is volatilized, 
upon warming, as silicon tetrafluoride. After evaporation of the 
excess of hydrofluoric acid and of sulphuric acid the residue is 
dissolved in water and the solution analyzed by practically the 
same procedure as is followed in the Smith method. 

Analysis of Silicate Mineral. — To insure complete decomposition of 
the silicate it must be ground much more finely than is necessary for 
most other minerals. After it is pulverized to pass a 100-mesh sieve 


about 3 gm of the sample is ground in an agate mortar until it will pass 
a 200-mesh sieve. 

Moisture. — Weigh about 0.5 gm of the silicate into a platinum cru- 
cible, and dry for one hour at 100° to 105°, the loss being calculated as 
hygroscopic moisture. If combined water is also to be determined this 
can be done by heating to a temperature above 200° in a combustion 
tube in an atmosphere of dried carbon dioxide, absorbing the moisture 
in weighed tubes filled with calcium chloride. 

Silica. — Mix with the sample about 5.0 gm of sodium carbonate, 
approximately weighed. Place the cover on the crucible but slightly 
toward one side so the contents may be observed. Heat gently at first 
in order to avoid violent effervescence and consequent loss by spatter- 
ing, gradually raising the temperature until the full heat of the burner 
fails to cause more than slight action. Place the crucible over the blast 
lamp and heat for 15 minutes after carbon dioxide has ceased to be 
evolved. Remove the lamp, lift the crucible by means of the tongs, 
using care to avoid contact of the latter with the fusion, and slowly 
rotate the crucible in such a manner that the fused mass will be spread 
over the sides of the crucible as it solidifies. 

After cooling place the crucible and its contents in a Pyrex beaker, 
a platinum dish or a porcelain casserole and cover with hot distilled 
water. Digest until the entire mass has become disintegrated and re- 
moved from the crucible. Cover the beaker and add, gradually, from 
a pipette, concentrated hydrochloric acid until the carbonates are com- 
pletely decomposed. Remove the crucible and cover by means of a 
stirring rod, rinsing well. Crucible tongs should not be used for this 
purpose unless they are tipped with platinum. The solution is now 
evaporated to dryness. If speed is essential and if a casserole has been 
used the greater part of the liquid may be evaporated by holding over 
a free flame, giving a rotary motion to the casserole to prevent spatter- 
ing or bumping. If other work may be carried on at the same time the 
solution should be evaporated over the steam bath. ■ When completely 
dry heat for 15 minutes at about 120°, cool and add 5 cc of concentrated 
hydrochloric acid and 50 cc of water, warm until soluble salts are dis- 
solved and filter the residue of silica at once, washing until free from 
chlorides, adding the washings to the filtrate. Evaporate the solution 
again to dryness, heat at about 120° and repeat the treatment with 
hydrochloric acid and water, filtering in a different paper. 

Place both papers with their residues in a platinum crucible, dry 
and burn the papers in the usual way, then ignite over the blast lamp 
for 20 minutes, cool and weigh. Small quantities of metal salts will 
have been retained in the residue. In order to correct for their presence 
the residue is moistened with one or two drops of sulphuric acid and 



about 5 cc of hydrofluoric acid is added. The silicon tetrafluoride being 
volatilized, the acids are evaporated and the residue, which should be 
small, is heated over the burner. The loss in weight is taken to repre- 
sent silica. The residue of oxides is dissolved by warming with a few 
drops of hydrochloric acid and is added to the nitrate from the silica. 
Iron, Aluminium, Manganese, Calcium, and Magnesium. — Concen- 
trate the nitrate, if necessary, to about 100 cc then determine iron, 
aluminium, manganese, calcium and magnesium exactly as directed for 

the analysis of carbonate minerals, cal 
culating each as the oxide. 

Sodium and Potassium. — Grind the 
sample in an agate mortar to pass 
200-mesh sieve and weigh 0.5-gi 
samples on counterpoised glasses 
Place 0.5 gm of resublimed ammoniui 
chloride (free from sodium and potas- 
sium) in the clean agate mortar, adc 
the weighed sample and grind together 
Add 3 gm of powdered calcium car- 
bonate and grind again. Place a gram 
of powdered calcium carbonate on the 
bottom of the crucible and brush in 
the mixture already prepared. Adc 
another gram of calcium carbonate to 
the mortar, grind to remove the las 
traces of sample from the mortar an( 
brush this in on top of the charge 
This should not entirely fill the 

Adjust the crucible in a hole in 
piece of asbestos board in such 
manner as that the board comes jus 
to the upper level of the charge. The upper part of the crucible wil 
then act as a condenser for any vapors of alkali chlorides that migh' 
form in the lower part of the crucible. 

Heat the lower part of the crucible gradually until the evolution o: 
ammonia becomes slow, then more strongly with the nearly full flame 
of two flat burners. This heating should be finally applied to all but the 
upper end. After 45 minutes of such heating the crucible is allowed to 
cool and the charge emptied into a platinum dish. The materials will 
usually be sintered together in a mass that has shrunken away from the 
crucible, the calcium carbonate in the bottom of the crucible serving 
to prevent sticking of the mass to the crucible. Actual fusion into a 

Fig. 70. — Smith's crucible for 
the determination of alkali metals 
in silicates. 


slag. that is not decomposed by water is an indication that the tempera- 
ture was too high or that more calcium carbonate should be used. 

Add 75 cc of water to the materials in the dish and warm until the 
mass is thoroughly disintegrated. Boil for a minute, allow to settle and 
filter. Add 50 cc more of water to the residue, boil, allow to settle and 
filter through the same paper, transferring all of the residue to the 
paper. Wash with hot water until a few drops of the washings show 
only a faint test for chlorides. 

To determine whether decomposition was complete in the crucible, 
dissolve the washed residue in hydrochloric acid. Any gritty residue 
must be retreated as before, as this indicates insufficient heating. 
Evaporate the combined filtrates and washings to a volume of 75 cc 
or less, then add 5 cc of dilute ammonium hydroxide and sufficient 
ammonium carbonate solution to precipitate all iron, aluminium and 
calcium present. Digest at a temperature just below the boiling point 
until the precipitate settles readily, then filter and wash once. Dissolve 
the precipitate in dilute hydrochloric acid and repeat the precipitation, 
washing the precipitate with hot water. Evaporate the two filtrates 
and all washings in a platinum dish. When dry, carefully heat over 
the free flame of the burner until all ammonium salts are volatilized, 
but without, at any time, allowing the dish to become red. Cool, dis- 
solve in not more than 10 cc of hot water, add a few drops of dilute 
ammonium hydroxide solution and 1 cc of saturated, recently prepared 
ammonium oxalate solution. Digest until the small amount of calcium 
oxalate settles, then filter and wash the paper, collecting filtrate and 
washings in the platinum dish. Add 1 cc of concentrated hydro- 
chloric acid and evaporate to dryness. Heat carefully to volatilize 
ammonium chloride. Cool in a desiccator and weigh as sodium chloride 
and potassium chloride. Dissolve in a few cubic centimeters of hot 
water and transfer to another dish. Dry, ignite and weigh the first 
dish with any matter that was not soluble and subtract this weight 
from the one already observed. The difference is sodium chloride and 
potassium chloride. 

Determine the potassium in this mixture as directed on page 102, 
omitting the washing with ammonium chloride solution, since we are 
here dealing with chlorides instead of sulphates of sodium and potassium. 

From the weight of potassium chlorplatinate found, calculate (a) 
the percent of potassium oxide in the silicate and (b) the weight of 
potassium chloride in the mixed chlorides. Subtract the latter weight 
from the total sodium and potassium chlorides and from the remain- 
ing sodium chloride calculate the percent of sodium oxide in the silicate. 

The sum of all oxides determined should approximate 100 percent. 
Since experimental errors accumulate and since small quantities of other 


substances than those named will generally remain undetermined, the 
sum of all will usually be less, rather than more than 100 percent. It is 
customary with many chemists to report this discrepancy as " undeter- 
mined" but it will be remembered that such a percent is not merely 
that of undetermined matter but that it also includes the accumulated 
errors of all of the other determinations. 



As might be expected from a knowledge of its origin, coal is 
made up of a large number of organic and some inorganic com- 
pounds. The attempt to separate and identify these compounds 
has long engaged the attention of chemists and geologists but 
comparatively little progress has been made in this direction. 
The reason for the lack of success in this attempt is that most 
chemical methods of analysis involve the breaking up of organic 
substances and the disappearance of the original forms. Such 
an analysis would prove extremely useful from the stand- 
point of geology and would, no doubt, be of service in indus- 
trial applications. In the absence of adequate methods for 
this purpose, the examination of coal is made with one or 
more of three ends in view: (1) To determine the geological 
origin of the coal; (2) to determine its adaptability to various 
industrial uses, such as steaming, heating, manufacture of pro- 
ducer gas or illuminating gas, coke, tars, etc.; or (3) to determine 
its fuel value in heat units per unit weight of coal. The 
methods used are classified under the head of proximate analysis 
or of ultimate analysis. The proximate analysis of materials 
may be defined as "the determination, not of elements or radicals 
but of groups of compounds falling within approximate limits 
of composition and having similar properties." The ultimate 
analysis is, as the word indicates, a determination of the ele- 
mentary composition. It is made with greater difficulty, 
involves more expensive apparatus and requires longer time 
than the proximate analysis, often gives no more useful informa- 
tion than is given by the proximate analysis and is, for this 
reason, less frequently made. 

Methods for the analysis of coal were standardized by a com- 



mittee of the American Chemical Society in 1899. 1 A joint com 
mittee of the American Chemical Society and the America 
Society for Testing Materials has since made two reports, 
revising, in many respects, the methods of the original committee 

Sampling. — The correct sampling of coal is a very difficul 
process. In a substance showing such a lack of uniformity 
composition it is obvious that the sample must be selected wit 
extreme care if the results of the analysis are to express th 
average composition. For a thorough discussion of this matter 
and especially with regard to the selection of samples from cars, 
heaps, mines, etc., the student is referred to a paper by Bailey. 3 
When the laboratory sample is finally selected it should be 
sealed in an air-tight container to prevent changes in the moisture 
content. This container may be a tin or galvanized can with 
a screw top, sealed with a rubber gasket and adhesive tape, 
or fruit jars, tightly sealed. 

In the chemical laboratory this sample must be further pul 
verized and divided in order finally to obtain a sample which i 
small enough in quantity and fine enough to serve for the actu 
analysis, and which shall also be representative of the origin; 

When the sample reaches the laboratory it is treated i 
the usual progressive crushing and quartering until the prope 
fine sample is obtained. During this process there is a continuec 
loss of moisture so that the fine sample finally obtained has no 
the same percentage composition as the sample as receivec 
If the coal as received is taken as a basis its composition will no 
be the same as that of another coal of different moisture conte 
but otherwise identical with the first. The only scientificall 
correct method is to calculate all other percents to a dry coal 
basis, reporting the moisture in the sample as received. It 
is sometimes necessary, however, to base all calculations upo 
the coal as received. In this case moisture must be determine 
at once, as well as after crushing and preparing the sample f< 
analysis, as already mentioned. Loss of moisture occurs at a 
times when the coal is exposed. The sample is ground fi 

* J. Am. Chem. Soc, 20, 281 (1898); 21, 1116 (1899). 
2 J. Ind. Eng. Chem., 6, 517 (1913); 9, 100 (1917). 

* J. Ind. Eng. Chem., 1, 161 (1909). 

FUELS 299 

enough to serve for the analysis and a determination of moisture 
is made upon this sample. The difference between the percent 
of moisture in the original coarse sample and the fine sample 
provides a basis for the calculation of the analytical results, 
to the original coal basis. 

Proximate Analysis. — The proximate analysis of coal as usually 
carried out includes the determination of moisture, volatile 
combustible matter, coke, " fixed carbon" and ash. The figures 
thus obtained give considerable information as to the geological 
age of the coal, determine its 'fitness for industrial uses and 
provide a basis for an approximate calculation of fuel value. 

Moisture. — The accurate determination of moisture in coal is a 
difficult operation. If the coal is heated to 100° or above this 
temperature there is danger of loss of volatile constituents other 
than moisture, whether these were originally present or were 
formed upon heating. Oxidation also takes place during heating. 
This may result in either gain or loss in weight, according to 
whether the greater part of the oxygen is retained in solid com- 
pounds or is lost as volatile ones. The usual tendency is toward 
a gain in weight so that this error to some extent compensates 
ithe loss first mentioned. Such, compensation is, however, not 
to be depended upon and such a determination is therefore not 
fcdeal. A method that removes some of these objections is that 
pf drying at ordinary temperatures over sulphuric acid, under 
diminished pressure. The partial pressure of oxygen being 
reduced to a negligible quantity, oxidation is entirely prevented. 
[There is also no tendency toward breaking down of the non-vola- 
pile organic constituents with the production of simpler volatile 
bnes. The method seems to give low and somewhat variable 
results, however, due to the low rate at which moisture is lost 
by the coal toward the last of the drying process. 

Volatile Combustible Matter, Fixed Carbon and Coke. — When 
:|;oal is subjected to dry distillation out of contact with air variable 
ihuantities of volatile products are expelled and a residue of 
norganic matter and non-volatile carbon is left. This residue is 
[he "coke," the carbonaceous portion of the coke being called 
t* fixed carbon." It should be understood that the original coal 
;lid not consist of free carbon and volatile organic compounds, but 
lhat heating at high temperatures resulted in the formation of 


such substances by a decomposition of the non-volatile bitumens 
composing the coal. Such a decomposition of complex com- 
pounds into simpler and more volatile compounds is technically 
known as "cracking." Cracking is a somewhat indefinite proc 
ess and the products depend to a large extent upon the tempera 
ture and time of heating. The volatile products of the distilla 
tion of coal have a very important application in the industries 
The non-volatile coke is also an industrial material, being exten- 
sively used as a fuel for reducing ores, for the manufacture o 
producer gas and water gas and for many other purposes. Th< 
relative quantities of volatile matter, fixed carbon and mineral 
matter are of importance in the determination of the fitness of 
various coals for various industrial uses. 

On account of the variation in results obtained by variation in 
the manner of heating it becomes difficult or impossible to make 
an intelligent comparison of different coals unless a standard 
method is adopted for their examination. The control anc 
measurement of the temperature at which they are heatec 
involves the use of a pyrometer and an easily controlled furnace 
It is also difficult entirely to exclude air during the heating and a 
variable oxidation occurs. The student is again reminded tha 
volatile combustible matter, fixed carbon and coke are arbitrarj 
classifications of substances produced by an arbitrary methoc 
and that they have little scientific meaning except as they provid 
a basis for comparisons. 

Ash. — Most of the inorganic matter of coal is left behind ai 
"ash" when the coal is burned but the compounds so remainin; 
are not identical with those of the original coal. Oxidizabl 
substances are oxidized, sometimes to a variable extent depending 
upon the degree of excess of oxygen in the atmosphere in which 
the coal is burned. Decomposable compounds are also change 
at the furnace temperatures. The first class of changes i 
illustrated by iron disulphide, the essential Compound of iror 
pyrites. This is oxidized to ferric oxide and sulphur dioxide 
the latter escaping with the other volatile products of com 
bustion. Carbonates of the earth and alkaline-earth metals 
and hydrated silicates are examples of compounds which an 
decomposed by heating. Such carbonates are nearly or com 

FUELS 301 

pletely converted into oxides and the hydrated silicates (such as 
clay) are partly deprived of their water of hydration. 

Because of such changes as these the percent of "ash," as 
determined by burning a weighed sample of coal and weighing 
the residue, must not be regarded as indicating directly the 
percent of inorganic matter or " non-coal' ' in the original ma- 
terial, but rather as the percent of non-combustible matter that 
would be left after the combustion of the coal in air. This in- 
accuracy of expression affects the report on fixed carbon as well 
as that of ash, since the former is obtained by subtracting the 
sum of ash, moisture and volatile combustible matter from 100. 
In the majority of cases the error is comparatively small and 
the correction is scarcely worth making. But if coal carries 
unduly large amounts of pyrite or of carbonates such a correction 
may be necessary, if a report on fixed carbon is to have any 

When iron disulphide is burned in air the sulphur is lost as 
sulphur dioxide and iron is left as ferric oxide : 

4FeS 2 +1102->2Fe 2 03+8S0 2 . 

Upon the assumption that total sulphur represents the sulphur 
of iron pyrite, the loss on burning each molecule of iron disul- 
phide results from the substitution of three atoms of oxygen for 
four of sulphur. That is, 

(4X32) -(3X16) =80. 

80 5 
Therefore ^qq or c °f the total sulphur percent should be added 

to the ash percent as a correction for use in obtaining the correct 
percent of fixed carbon. 

Such carbonates as limestone lose carbon dioxide when 
strongly heated. It is not by any means certain that the con- 
version to calcium oxide, etc., is complete at the temperatures 
employed in burning the coal and if a correction is to be made 
it is better to add a few drops of sulphuric acid to the ash, evapo- 
rating the excess of acid and heating gently before a weighing 
is made. The ash which has been treated with sulphuric acid 
now contains sulphates instead of carbonates, and it is heavier 
on this account. Upon a separate sample of coal a determination 


of carbon dioxide of carbonates is made by any of the standarc 
methods (see pages 128 to 137). 

The gain in weight through sulphation is represented by th 
difference between the molecular weights of sulphur trioxid 
and carbon dioxide. That is, 

80-44 = 36. 

36 9 
Therefore jt or tt °f the percent of carbon dioxide of car 

bonates should be subtracted from the percent of sulphatec 
ash to obtain the original carbonated inorganic matter. 

It should be noted that the term "ash" is a misnomer when 
applied to the percent corrected as above described. The latte 
should more properly be called "inorganic matter" or "non 

Fusing Point of Ash. — The composition of coal ash has a 
important bearing upon the industrial uses of the coal. Thi 
is particularly true with regard to the fusibility of the ash 
which is the principal property upon which depends the forma 
tion of clinker. If the inorganic matter of coal is of such com 
position as to produce an easily fused ash, a slag will form ir 
the furnace before combustion is complete, with the result thai 
more or less coal matter is so glazed and protected from oxida- 
tion as entirely to prevent its combustion. Thus there is 
troublesome clinker formed which clogs the grates and interferes 
with proper stoking, and also a variable waste of combustible 
matter. This action is, of course, more pronounced in cases 
of high furnace temperatures. Clinker will often contain 
very high percent of combustible matter, the latter being deter 
mined by grinding the solid clinker and burning at temperature 
under the fusing point of the ash. 

In most high grade coals the inorganic matter consists very 
largely of silicates of aluminium, all more or less refractory 
Clay is typical of these. • If substances capable of yielding basic 
oxides upon ignition are present, slag formation is promotec 
and clinker forms. The most abundant of such base formers 
are iron pyrite and calcium carbonate. On this account a high 
sulphur content is usually indicative of ability to form a fusible 
ash. A similar indication is provided by a red ash, since mos 
of such color comes from ferric oxide. The presence of lime 

FUELS 303 

stone is betrayed by neither high sulphur nor red ash. Con- 
sequently a coal may contain little sulphur and form a white ash 
and yet clinker badly, the calcium carbonate (or oxide) readily 
uniting with refractory silicates to produce complex silicates 
which fuse or soften at the prevailing furnace temperatures. 

On account of the objectionable character of clinker a deter- 
mination of fusibility of ash often becomes highly significant. 
Of course such a complex mixture as that which composes coal 
ash can have no very definite fusing point, so that observations 
of the behavior of the ash at high temperatures must be made 
according to somewhat arbitrary standards. The ash is ground, 
mixed with water and (usually) some organic binder, and molded 
into pyramids similar to the well known "Seger cones," which 
have long been used for measuring furnace temperatures. The 
"cone" is really a triangular pyramid having a base perpen- 
dicular to one of the plane sides. When such a piece is heated 
to its softening temperature it leans toward the vertical side. 
The temperature at which the apex curves down to touch the 
supporting plane is taken as the "fusing point." 

Fieldner and Hall 1 and Fieldner and Feild, 2 working in the 
Bureau of Mines, found that the fusing temperature so observed 
varies somewhat according to the size, shape and inclination of 
the test piece, the fineness of grinding the ash before molding 
and the atmosphere in which the pyramid is heated. The effect 
of the surrounding atmosphere is to cause changes in the composi- 
tion of the ash, the latter reacting with oxidizing or reducing 
gases. The nearest imitation of furnace conditions was found 
in a mixture of equal volumes of hydrogen and water vapor. 
This keeps the iron of the ash in the ferrous condition, in which 
form it is usually found in furnace clinker. 

This determination is described with great detail in the papers 
already cited. If the results are to have any great value the 
determination must be made by a carefully standardized method 
and equipment, including the special furnace there described, 
or a similar one. 

Preparation of Sample, (a) When Coal Appears Dry. — If the sample 
is coarser than 4-mesh and larger in amount than 10 pounds quickly 

1 J. Ind. Eng. Chem., 7, 399 and 474 (1915). 

2 Ibid., 7, 742 and 829 (1915). 


crush it with a jaw crusher to pass a 4-mesh sieve and reduce it on a 
riffle sampler (page 14) to 10 pounds, then crush at once to 20- 
mesh by passing through rolls or an enclosed grinder and take, without 
sieving, a 60-gm total moisture sample, immediately after crushing. 
This sample should be taken with a spatula from various parts of the 
20-mesh product and should be placed directly in a rubber-stopperei 

Thoroughly mix the main portion of the sample, reduce on the smal 
riffle sampler to about 120 gm and pulverize to 60-mesh by any suitabl 
apparatus without regard to loss of moisture. Mix and divide the 60 
mesh sample on the small riffle until it is reduced to 60 gm. Preserve 
this in a rubber-stoppered bottle. 

Determine moisture in both the 60-mesh and the 20-mesh samples 
by the methods given under the head of moisture. 

(b) When Coal Appears Wet. — Spread the sample on weighed pans, 
weigh and dry for 12 hours at room temperature or in a special drying 
oven through which air circulates freely at 10° to 15° above the room 
temperature. Reweigh and continue the drying until the loss in weight 
is not more than 0.1 per cent per hour. Complete the sampling as 
with dry coal and calculate the percent of moisture by air drying. 

To find the total moisture in wet coal, as received, compute 
as follows: 

Let a = percent of moisture by air drying, 

M 2 o = percent of moisture in 20-mesh coal. 

Then j-jrx -°+ a = total moisture as received. 

The percents of the various constituents of the coal are de 
termined on the 60-mesh coal. To compute these percents t 
the dry-coal basis or the as received bases, proceed as follows 

Let M 6 o = percent moisture in 60-mesh coal, 

M r = total percent moisture in coal as received, 
Peo = percent of any constituent in 60-mesh coal, 
Pq = percent of any constituent in dry coal, 
P r = percent of any constituent in coal as received. 

_ /100-Ma _ / 100 -AfA 
^ \ 100 / io ~\100-M 60 / 


FUELS 305 

Proximate Analysis. — Procure a sample as directed in the preceding 

Moisture. — The oven for drying the 20-mesh and 60-mesh samples 
must be so constructed as to provide a uniform temperature in all parts 
and a minimum of air space. Provision must be made for renewing 
the air in the oven at the rate of two to four times a minute, with air 
dried by passing through concentrated sulphuric acid. 

A convenient form of crucible for the moisture determinations and 
one which allows the ash determinations to be made on the same sample 
is a flat bottomed porcelain crucible. A fused silica crucible of similar 
shape may be used. In either case a well fitted aluminium cover should 
be provided. Glass capsules, with covers ground on, may also be used 
but ash determinations will then require different samples. 

Sixty-mesh Sample. — Heat the empty crucible under the conditions 
under which the coal is to be dried, cover, cool in a desiccator over con- 
centrated sulphuric acid for 30 minutes and weigh. Place approxi- 
mately 1 gm of the sample in the crucible, cover and reweigh. Remove 
the cover and place the crucible in the oven, which is maintained at 
104° to 110°. Heat for 1 hour then cover the crucible, cool over sul- 
phuric acid for 30 minutes and weigh. Calculate the loss as moisture. 
Twenty-mesh Sample. — Use 5-gm samples weighed with an accuracy 
of 2 mg and heat for l}4 hours. The procedure is otherwise the same 
as for the 60-mesh sample. 

Permissible differences in duplicate determinations: 

Same analyst Different analysts 

Moisture under 5% 0.2% 0.3% 

Moisture over 5 % 0.3 0.5 

Ash. — Place the uncovered porcelain crucible containing the dried 
60-mesh coal in a cold muffle furnace through which air may be drawn 
and gradually raise the temperature to between 700° and 750°. When 
combustion appears to be complete remove the crucible, cover, cool in a 
desiccator, and weigh. Repeat the heating for 15-minute periods until 
the change in weight is not greater than 1 mg. Calculate the percent 
of ash. 

Permissible differences in duplicate determinations: 

Same analyst Different analysts 

No carbonates present 0.2% 0.3% 

Carbonates present 0.3 0.5 

Carbonates and pyrite present 

and more than 12% ash 0.5 1.0 



Volatile Combustible Matter. — For this determination an electri- 
cally heated, vertical-tube furnace is desirable, although a muffle 
furnace, heated by gas or electricity may be used. If the tube furnace 
is available the vertical tube should be about l}i inches in diameter 
and 6 inches deep. The junction of a thermo-couple connected with 
a pyrometer meter should be placed immediately below the bottom of 
the crucible. The furnace is to be kept covered during a determination. 

If the determination of volatile combustible matter is not an essentia 
part of the specifications under which the coal is bought a No. 4 M6kei 
burner may be used for the heating. 

The platinum crucible should have a capacity between 10 and 20 cc 
a diameter between 25 and 35 mm and a depth between 30 and 35 mm 
The cover must fit closely. 

Weigh the crucible then add, as nearly as possible without careful 
adjustment, 1 gm of 60-mesh coal, cover the crucible and reweigh. If 
a furnace is used for the heating it should be brought to a temperature 
of 950° (±20°), and the crucible is placed on a support of platinum or 
of nickel-chromium in the furnace. If a Meker burner is used, adjust 
the flame so that its extreme height is 15 cm and support the crucible 
on a platinum or a nickel-chromium wire triangle so that the bottom is 
1 cm above the top of the burner. 

After the luminous flame above the crucible has disappeared, tap 
the crucible cover lightly in order to seal more perfectly and thus 
guard against the entrance of air. Heat for exactly 7 minutes then 
remove the crucible from the furnace or flame without disturbing 
the cover. Cool in a desiccator and weigh. Calculate the loss as 
percent of total volatile matter, including moisture. Subtract tht 
percent of moisture found in the 60-mesh sample and report the percent 
of volatile combustible matter. 

Sub-bituminous coal, lignite or peat are given a preliminary gradua 
heating for 5 minutes to expel part of the large quantity of volatile 
matter. They are then heated as above directed for 6 minutes. 

Permissible differences in duplicate determinations: 


Bituminous coals 


Same analyst 

Different analysts 





Fixed Carbon. — Subtract from 100, the sum of moisture, volatile 
combustible matter and ash, corrected to the basis of either dry coal 
or as received. The remainder is the percent of fixed carbon on the 
basis considered. 

FUELS 307 

Ultimate Analysis. — The complete ultimate analysis of coal 
will include the determination of all elements. The com- 
plete analysis is not often made, the determination of sulphur, 
carbon, hydrogen and nitrogen being all that is usually required. 

Sulphur. — Sulphur may exist in coal in one or more of four 
forms: elementary sulphur, inorganic sulphides (principally 
iron pyrite) inorganic sulphates and organic compounds. The 
accurate determination of the amount present in the different 
forms is difficult and not often required. 

Three methods have found considerable use for determining 
total sulphur in coal. The powdered sample may be gently 
heated with sodium carbonate (Atkinson's method 1 ) or with a 
mixture of sodium carbonate and magnesium oxide (Eschka's 
method 2 ) ; or it may be fused with sodium peroxide. In the first 
two cases air is the oxidizing agent while in the last sodium 
peroxide performs this function. The ultimate result in all 
three is that the organic portion of the coal is burned while 
sulphur is oxidized, either dioxide or trioxide being retained by the 
basic constituents of the added reagents, soluble sulphites or 
sulphates being produced. In the Eschka method magnesium 
oxide serves to provide a porous mixture into which air readily 

It is necessary to provide against absorption of sulphur oxides 
from the surrounding atmosphere. For this reason ordinary 
illuminating or coal gas cannot be used for direct heating of the 
mixture unless it is first purified. Alcohol burners are suitable 
for this purpose but a muffle furnace, heated by electricity, is 

When the oxidation of the coal is complete the mixture is boiled 
with water and bromine, the latter to oxidize to sulphates any 
sulphites that may have been formed. After filtration the solu- 
tion is acidified and boiled, whereby bromates or hypobromites 
are decomposed and the bromine expelled from the solution : 
NaBr0 3 +HCl-+NaCl+HBr0 3 , 
NaBrO + HCl->NaCl + HBrO, 
HBr0 3 +5HBr^3H 2 0+3Br 2 , 
HBrO+HBr-+H 2 0+Br 2 . 

l J. Soc. Chem. Ind., 6, 154 (1886). 
2 Chem. News, 21,^261 (1870). 



It is necessary thus to decompose salts of oxyacids of bromin 
because of the tendency shown by barium sulphate to occlud 
these compounds at the moment of precipitation. 

Nitrogen. — Nitrogen may be determined by combustion or 
by the Kjeldahl, the Gunning or the Kjeldahl-Gunning methods. 
The principles underlying the latter are discussed on page 513 
and following. In the combustion 
method the coal is mixed with fine cupric 
oxide and is heated in a tube closed at 
one end. Oxidation occurs and the gases 
are passed through more heated cupric 
oxide to complete the oxidation of carbon 
monoxide or of gaseous hydrocarbons, 
then over heated copper, the latter to 
reduce oxides of nitrogen to elementary 
nitrogen. The mixture of carbon dioxide, 
water vapor, sulphur dioxide and nitro- 
gen is passed into sodium hydroxide 
solution where all gases except nitrogen 
are absorbed. The latter is collected in 
a eudiometer and measured, the weight 
being then calculated. A special form 
of eudiometer called the " nitrometer," 
is useful for this purpose. This is shown 
in Fig. 71. In order to force the gases 
out of the combustion tube a short space 
in the closed end is filled with sodium 
bicarbonate. Carbon dioxide is evolvec 
at will by heating this and it causes the 
expulsion of other gases from the tube. 
Carbon and Hydrogen. — The deter- 
mination of carbon and hydrogen is made 
by combustion and absorption of water 
vapor and carbon dioxide in weighed tubes containing ap- 
propriate absorbents. The method is the same as that used in 
other connections for organic substances containing sulphur and 
nitrogen. For the absorption of water vapor calcium chloride or 
sulphuric acid may be used, the latter giving more nearly complete 
absorption while the former is more conveniently used. For 

Fig. 71.— Schiff' 

FUELS 309 


carbon dioxide, solid soda lime or a solution of potassium hydroxide 
is used, any of the standard forms of absorbing tubes or bulbs 
serving as containers. A combustion tube of glass, silica or 
porcelain is used for the decomposition of the coal. The relative 
merits of these materials were discussed in an earlier section 
(pages 37 and 42). For the present purpose the silica tube is 
most satisfactory and any suitable combustion furnace may be 
used. The tube should be 95 cm long and should project 10 to 
15 cm beyond the furnace at each end. The combustion is 
» effected by heating the coal in an atmosphere of oxygen. Since 
this also produces volatile organic compounds and some carbon 
monoxide, it is necessary to pass the gases through a solid oxidiz- 
ing agent in order to complete the oxidation. Cupric oxide is 
used for this purpose in the analysis of most organic substances, 
but when sulphur is present, as in all coals, sulphur oxides are 
not completely retained by cupric oxide and are therefore ab- 
sorbed in bulbs containing potassium hydroxide. In the com- 
bustion of such materials lead chromate is substituted for part 
of the cupric oxide. Lead sulphate is formed and this is not decom- 
posed by heating. A lower temperature is used to avoid fusing 
the material into the tube. 

Oxygen for the combustion may be made from manganese 
dioxide and potassium chlorate and stored in a large " gasometer " 
or it may be purchased in steel cylinders. Oxygen made by the 
first method named always contains chlorine oxides and it must 
be purified by passing through a concentrated solution of potas- 
sium hydroxide. Most of the commercial oxygen formerly 
obtainable in compression cylinders contained chlorine oxides, 
carbon dioxide and hydrogen, and such oxygen must be properly 
purified before it enters the combustion tube. Since the develop- 
ment of the manufacture of oxygen from liquid air it is possible 
to obtain gas having a high degree of purity. It is then only 
necessary to pass the oxygen through potassium hydroxide or 
soda lime in order to remove traces of carbon dioxide. In order 
to be able to control the flow a " gasometer" should be filled from 
the high-pressure cylinder and this used as the supply for the 
combustion or else a special high-pressure control valve should 
be used on the tank. 
Oxygen. — The direct determination of oxygen is difficult and 


is seldom attempted. The percent is sometimes estimated by 
subtracting the sum of percents of all other elements, water and 
ash, from 100. This is but a rough approximation but is usually 
all that is required. 

Ultimate Analysis of Coal: Sulphur. Eschka Method. — Thoroughly 
mix two parts of light calcined magnesium oxide and one part of anhy- 
drous sodium carbonate. These materials should be as free as possible 
from sulphur. Place 3 gm of the mixture on a sheet of white glazec 
paper and then add 1 gm of 60-mesh coal, accurately weighed. Mix 
well, then transfer to a porcelain, silica or platinum dish, 2 inches in 
diameter and 1 inch deep. Cover with about 1 gm of the Eschka 

A flame of illuminating gas cannot be used for heating on account 
of absorption of sulphur dioxide from the flame by the basic mixture. 
A gas or electric muffle furnace is best for the purpose but an alcohol, 
gasoline or natural gas flame may be used. 

If a flame is applied directly, heat the dish in a slanting position 
on a triangle over a low flame until most of the volatile matter is driven 
off, then gradually increase the temperature and heat for 30 minutes 
or more, stirring occasionally, until all black particles have disappeared. 

If a muffle furnace is to be used, place the crucible in the cold furnace 
and gradually raise the temperature so that about 1 hour is required 
to reach 870° to 925°. Maintain this temperature for 1% hours, then 
cool in the furnace or in air that is free from gases containing sulphur 

After either treatment rinse the material into a 200-cc beaker, add 
100 cc of nearly boiling water and digest on the steam bath for 30 min- 
utes, stirring frequently. Filter and wash the insoluble matter thor- 
oughly with hot water. The filtrate and washings should total about 
250 cc. Add 20 cc of saturated bromine water to the solution and stir, 
then make slightly acid with dilute hydrochloric acid (shown by the 
failure of more acid to cause effervescence) and boil until all bromine 
is removed and the solution is colorless. Add two or three drops of 
methyl orange or methyl red and neutralize with 10 percent sodium 
hydroxide solution, then add 1 cc of approximately normal hydrochloric 
acid, or an equivalent volume of a solution of any other normality. 

Heat to boiling and add, dropwise and with stirring, 20 cc of 5 percent 
solution of barium chloride. Digest on the steam bath until the pre- 
cipitate settles readily and then filter and wash free from chlorides with 
hot water. Fold the paper, place in a weighed platinum, porcelain, 
silica or alundum crucible and burn the paper at a low temperature 

FUELS 311 

and with free access of air (see page 93), finally heating to dull red- 
ness for 10 minutes. Cool and weigh. Calculate the percent of sul- 
phur in the 60-mesh coal. 

The residue of magnesium oxide, ash, etc., should be dissolved in 
hydrochloric acid and tested for sulphur. If any is found this must be 
determined quantitatively and added to the percent already found. 
. A blank experiment should be performed, using all of the reagents of 
the regular experiment but omitting the first heating. Any sulphur 
that is so obtained from the reagents is to be subtracted from that found 
in the analysis of the coal. 

Permissible differences in duplicate determinations: 

Sulphur less than 2 % . 
Sulphur more than 2 % , 

le analyst 

Different analysts 





Carbon and Hydrogen. — A tube combustion furnace of any of the 
approved types and about 75 to 80 cm long is necessary. The combus- 
tion tube may be of hard glass, silica or porcelain. It should be long 
enough to project for at least 10 cm at each end of the furnace, in order 
to prevent heating of the rubber stoppers that must be inserted in 
the ends. The internal diameter of the tube should be 12 to 15 mm. 
Since coal always contains sulphur, lead chromate must be used in the 
combustion tube and this is prepared by fusing, cooling and crushing 
about 100 gm. The largest pieces should be small enough to easily enter 
the tube. By sifting the crushed material, using a 40-mesh sieve, a finer 
grade will be obtained and this is used for mixing with the powdered 

The combustion tube should have well rounded ends. It is filled 
according to the following directions, assuming that the length of the 
furnace is 75 cm and that of the tube is 95 cm. 

Into one end of the tube insert a closely fitting roll of copper gauze, 
5 cm long, and push this in until a space of 10 cm is left at the end 
of the tube. Into the other end pour the coarsely crushed lead chromate 
until a space 50 cm long is filled. The material should be well settled 
but not packed in such a way as to obstruct the passage of gases. Insert 
another roll of copper gauze like the first, to hold the lead chromate in 
place. Another roll of copper gauze, 10 cm long, is inserted in such a 
way as to leave a space of 10 cm at the end of the tube, and a 
space of 5 cm between the two rolls. The latter space is for the boat 
containing the coal. The ends of the tube are closed by rubber stoppers 
carrying short glass tubes for connecting with the rest of the apparatus. 
The method of filling the tube is shown in Fig. 72. Fig. 73 shows the 











method of assembling the complete apparatus. Entering oxygen anc 
entering air are passed through cylinders A and A', re- 
spectively, containing a good quality of soda lime. The gases 
are next passed through U-tubes, B and B' ', containing fused, 
granular calcium chloride. These tubes are connected with 
the combustion tube by means of a three-way stopcock 
Gases leaving the combustion tube first pass through two 
U-tubes C and D (preferably glass stoppered) containing 
calcium chloride, then through the carbon dioxide absorption 
bulbs E containing potassium hydroxide solution and calcium 
chloride. Following these tubes is a guard tube F containin 
^ calcium chloride, also an aspirator G. For the detailed direc- 
8 tions for filling and connecting the various absorption tubes 
*S and aspirator, refer to the discussion of the determination o; 
•§ carbon dioxide in carbonates. 

^ When the combustion tube and all parts of the apparatus ar 
§ in order start a slow but steady current (three bubbles pe 
Jd second in the bulbs) of air by means of the aspirator, then 
© heat gradually the entire length of the combustion tube. The 
£ drying tubes C and D and the potassium hydroxide bulbs 
g need not be in the train because this preliminary heating is 
*■§ for the purpose of thoroughly drying the contents of th 
rg combustion tube and oxidizing any organic matter with which 
o the tube might be contaminated. They are therefore re 
g moved, carefully wiped clean and dry, and are then close 
*JJ and placed in the balance case. After they have stoo 
I? for 15 minutes, if ready to proceed with the blank test 
jQ these pieces are weighed. The temperature of the part of 
the tube containing lead chromate must not be higher than 
£2 is indicated by dull redness although other parts may be 
6 heated to any temperature under the softening point of the 
£ tube. 

When moisture has been expelled from the tube to the 
extent that no condensation is noticed on the forward end the 
calcium chloride tubes C and D and the potassium hydroxide 
bulbs are weighed and placed in the absorption train. The 
flow of air is continued for 20 minutes, when the aspirator is 
stopped and the absorption tubes are again removed, stop- 
pered, placed in the balance case and weighed after standing 
for 15 minutes. If there is a gain of more than 0.5 mg in 
either the weight of the potassium hydroxide bulbs or the 
combined weights of the two calcium chloride tubes the entire 
operation must be repeated until there is no greater gain than 0.5 mg. 

FUELS 313 

When this is the case, that part of the combustion tube which is at 
the left of the lead chromate is allowed to cool. 

Provide a porcelain or platinum boat, about 5 cm long and of the 
proper width for insertion into the combustion tube. In the bottom 
of this place a layer of powdered lead chromate 1 mm deep and then 
weigh into the boat about 0.5 gm of powdered coal which has been 
properly sampled and dried. Mix with a platinum wire. Remove 
the rubber stopper at the left; end of the tube and quickly remove the 
roll of copper gauze (now largely oxidized to cupric oxide) by means of 
a wire hook and insert the boat, pushing the latter in until it touches 
the roll of cupric oxide which confines the lead chromate. Replace the 
first roll and the rubber stopper as quickly as possible, start a current of 
oxygen through the tube and gradually heat the cooled portion of the 

■J B' 


C D E F 


Fig. 73. — Diagram of connections for combustion apparatus. 

tube. Volatile matter will escape from the coal but this will be com- 
pletely oxidized by the lead chromate in the forward part of the tube. 
Backward diffusion of volatile combustible matter will occasion no loss 
by condensation because the roll of cupric oxide behind the boat will 
serve to oxidize a small quantity of such gases. 

When all glowing of the coal has ceased, turn the three-way stopcock 
so that air is drawn into the tube, gradually lower the temperature of 
the left end so as to avoid cracking and continue the passage of air 
until about 1000 cc more of water has run out of the aspirator. Remove 
the absorption tubes and bulbs, close and allow to stand 15 minutes and 
then weigh. From the total gain in weight of the two calcium chloride 
tubes, due to absorption of water vapor, calculate the percent of hydro- 
gen in the coal. From the weight of carbon dioxide absorbed in the 
potassium hydroxide bulbs calculate the percent of carbon in the coal. 

Duplicate determinations should be made. If many samples are to 
be analyzed much economy of time will result from the use of two boats 
and two sets of absorption tubes. As each experiment is finished 
another can be started and the combustion will proceed while the first 
set of tubes is standing in the balance and being weighed. 



Fuel Value. — After a decision has been reached as to the class 
of coal that is most suitable for a given industrial purpose th< 
inquiry that is next in importance concerns the number of hea 
units that can be obtained from unit weight of the various grades 
of coal entering into that class. The custom of purchasing coa 
upon a tonnage basis at a contract price, with nothing more thar 
the variety of coal named, is rapidly being displaced, by larg* 
consumers, by the method of purchasing upon a heat unit basis 
A contract price is made, based upon a specified number of hea 
units per pound or ton and any deviation from this fuel value 
involves a corresponding alteration in the price. 

In scientific work the fuel value is always calculated as calories 
per gram or kilogram of fuel, while in industrial work it is gener 
ally calculated as " British thermal units" per pound of fuel 
The calorie is the quantity of heat required to raise 1 gm of wate 
1° C. in temperature. The British thermal unit (B. T. U.) is the 
quantity of heat necessary to raise 1 lb of water 1° F. in tem- 
perature. The calculation of fuel values in either system in- 
volves equal weights of water and fuel and the relation 

cal per gm . . . .• ,,, _ _ 

p rp tj vr is therefore the relation of the centigrade degree 

to the Fahrenheit degree. That is, cal per gmX •= = B. T. U. per 

lb, and B. T. U. per lbXq = cal per gm. 

This may be demonstrated as follows : 
Let a = temperature equivalent of 1° C, 

b = temperature equivalent of 1° F., 

c = weight equivalent of 1 gm, 

d = weight equivalent of 1 pound. 

Then B. T. U. perlbX^ = B. T. U. per gm; 

IB. T. U.=— cal, 

therefore B. T. U. per lb X- = cal per gm, 


or B. T. U. per lbXq = cal per gm. 

Conversely cal per gmX- R -B. T. U. per lb. 

FUELS 315 

Calculation of Fuel Value from the Ultimate Analysis. — The cal- 
culation from the ultimate analysis is based upon the assumption 
that the heat of oxidation of a compound is equal to the sum of 
the heats of oxidation of the elements composing it. . The ele- 
ments of coal that are oxidizable are carbon, hydrogen, and sul- 
phur. Nitrogen is mostly evolved in the free state and inorganic 
matter other than sulphides is of little or no value for producing 
heat. Water is incombustible and absorbs heat in becoming 
vaporized, thus reducing the available heat energy. Oxygen 
is not only incombustible but, because of the fact that it is 
already combined with carbon and hydrogen it reduces the per- 
cent of these elements still available for combustion and therefore 
reduces the amount of heat that is available when the fuel is 
burned. The fuel value of elementary carbon is 8080, that of 
hydrogen is about 34,500 and that of sulphur is 2162 calories per 
gram, if it is understood that the products of combustion are 
cooled to ordinary temperature after combustion. This assump- 
tion is not realized in practice but it is customary to make the 
ssumption in calculating fuel values. There is, of course, no 
way of knowing the manner in which the elements are combined 
in the organic compounds making up the coal substances. The 
arbitrary assumption is, however, made that all oxygen that is 
ilready contained in the dry coal is in combination with hydrogen. 
One-eighth of the percent of oxygen would then be the percent 
}f hydrogen not available as fuel. The complete statement for 
3alorific value, based upon these assumptions, is given in the 
'olio wing formula of Du Long : 

34500 (h-J0) +8080 C+2162 S 

q = - i0Q - cal per gm, 

vhere H, O, C and S represent percents of hydrogen, oxygen, 
:arbon and sulphur respectively. 

However this formula is hardly an accurate statement of avail- 
able heat. Before any compound can be burned to the oxides of 
he constituent elements the compound must be dissociated into 
ts elements. This change will involve absorption or liberation 
usually absorption) of heat energy and the amount of energy 



change will depend upon the method of combustion. This 1 
not known for coal and the correction cannot be applied. Ther 
is also a fuel value (positive or negative) for the mineral matte 
contained in the coal, since the ash left upon burning is not th 
same as the original matter. This heat must also be omittec 
from the calculation because its quantity is not known. 

The changes taking place during the conversion of wood into 
the various forms of coal are represented by the following approxi 
mate figures for the composition of the combustible portior 
disregarding small quantities of elements other than those des 
ignated, as well as ash. 











Bituminous coal 


Anthracite coal 



The gradual loss of volatile matter involves a relatively large loss 
of oxygen. This fact explains the steady rise in fuel value 
the coals progress toward the anthracite, although the decreaj 
in the ratio of hydrogen to carbon remaining would lead one 
expect a fall in fuel value. 

Calculation of Fuel Value from the Proximate Analysis.- 
Many attempts have been made to devise a formula that willj 
serve for calculating fuel value from the results of the proximate 
analysis. Such a formula must necessarily be purely empiric* 
and must fail in many cases because coals of practically identic* 
proximate composition may vary widely in ultimate compositioi 
and constitution. The chief value of all such formulas lies 
making possible an approximate valuation of the fuel whei 
neither the ultimate analysis nor calorimetric determinations cai 
be obtained. Following are a few examples of these formulas 
The results of such calculations must be used with great cautioi 

Formula of Haas: 

156.75 [100- (% ash+% S+%H 2 O)]+40.5X % g 
B.T.U. per lb. 



Formula of Goutal: 1 

82 C + AM = cal per gm when C = % fixed carbon, M = % 
volatile combustible matter and A = a coefficient whose value is 
fixed by the volatile combustible matter as follows: 

M = 2-15 15-30 30-35 35-40 
A = 130 100 95 90 

Formula of Gmelin: 

[100- (% H 2 0+% ash)] 80-6CX% H 2 = cal per gm, 

in which C is a coefficient which varies with the percent of mois- 
ture as follows: 

Percent moisture 



- 4 


+ 6 


+ 12 


+ 10 


+ 8 


+ 6 


+ 4 

If fixed carbon is calculated upon a basis of true coal, dry and 
ash free, the following table may be used: 


B.T.U. per lb 


B. T. U. per lb 





























Determination of Fuel Value by Means of the Calorimeter. — 
The best laboratory method for the determination of fuel value 
is by the use of one of the standard calorimeters. Practically 
,11 of these depend upon the measurement of rise in tempera- 
ure of water, caused by the combustion of fuel within a closed 
or "bomb" immersed in the water. Combustion is best 

Compt. rend., 135, 477 (1902). 


effected by electrical ignition in an atmosphere of compressec 
oxygen since oxidation is complete and no reactions can occu 
other than those of ideal combustion. The original bomb calo 
rimeter of Berthelot 1 has been improved and changed in such 
manner as to make it a practical instrument for industrial as wel 
as for purely scientific laboratories. Successful modifications are 
those of Mahler 2 and Emerson. 3 Parr 4 has also perfected a fue 
calorimeter in which the oxidizing agent is sodium peroxide. The 
advantages of this instrument for industrial testing are chiefly 
due to the fact that it dispenses with the use of compressed oxygen 
and much of the accessory apparatus for filling the bomb 
Charging and firing become a comparatively simple matter and 
the instrument may be operated by persons who have limited 
scientific training. The great fault of this and similar instru- 
ments comes from the reaction of the products of combustion 
with the excess of sodium peroxide or with sodium monoxide 
formed. Such reactions as the following occur: 

H 2 0+Na 2 0-*2NaOH, 

C0 2 +Na 2 0->Na 2 C0 3 , 
S0 3 +Na 2 0->Na 2 S0 4 . 

Materials in the ash also react and form sodium salts. All 
these reactions involve heat liberation or absorption and this 
cannot be exactly calculated because the composition of the coa 
is never exactly known. The best that can be done is to deter 
mine experimentally approximate corrections which will apply to 
different classes of coal. The sum of these corrections may often 
be as large as 10 percent of the total rise in temperature during 
determination and the uncertainty is so great as to render the 
instrument of questionable value for any but the most approxi- 
mate determinations. The instruments using compressed oxy- 
gen, while usually more expensive, are best for accurate fue 
testing, even for the works laboratory. 

Fuel Value. — If a calorimeter is available the heat units 
should be determined experimentally. Calculation may also 

1 Ann. chim. phys., [5] 23, 160 (1881); [6] 10, 433 (1887). 

2 Chem. Zentr., 63, 889 (1892). 

3 J. Ind. Eng. Chem., 1, 17 (1909). 

4 J. Am. Chem. Soc, 22, 646 (1900); 29. 1606 (1907). 



be made from the analytical results, using the formulas already- 
given as well as others that have been proposed. Comparison 
with calorimetric data will indicate the degree of usefulness of the 

Following is a description of the Emerson calorimeter and also 
directions for making the determination of fuel value. 

Bomb. — The bomb is made of steel, consisting of two cups 
joined by means of a heavy steel nut. The two cups are machined 

Fig. 74. — Emerson's calorimeter. 

at their contact faces with a tongue and groove, the joint being 
made tight by means of a lead gasket inserted in the groove. The 
lining is of sheet nickel, platinum or gold, spun in to fit. The 
bomb is made tight with a milled wrench or spanner. The pan 
holding the combustible is of platinum or nickel, and the support- 
ing wire of nickel. 

Calorimeter. — The jacket is a double walled copper tank, 
the space between the walls being filled with water. The calo- 
rimeter can is made as light as is possible, of sheet brass, nickel 

Stirring Device. — The stirrer is directly connected to a small 
motor and is enclosed in a tube to facilitate its action in circulating 



the water. The stirrer is mounted on a post on the calorimete 
jacket as is the thermometer holder. 

Fig. 75. — Section of Emerson calorimeter. 

Ignition Wire. — Unless ignition of the fuel requires a very hig 
temperature a platinum resistance wire is suitable. For ignition 
of such substances as are used in determining the water equivalen 
of the calorimeter (naphthalene, cane sugar, etc.) or of anthracit 

FUELS 321 

coal an iron wire is more certain in its action because it burns 
and produces a higher temperature. When iron wire is used a 
correction of 1600 calories per gram of wire is subtracted from 
the total calories obtained from the fuel combustion. This 
is the heat of oxidation of the iron. 

Formation of Nitric Acid. — When coal is burned in air prac- 
tically all of the nitrogen is liberated in the elementary form. 
On account of the high concentration of oxygen in the calorimeter 
bomb a considerable portion of the nitrogen is oxidized and the 
products dissolve in the water which is formed by the combustion 
of hydrogen. A dilute solution of nitric acid is thereby formed. 
This gives rise to a positive error in the observation of fuel value, 
the magnitude of the error depending upon the extent to which 
nitric acid is formed. As a rule the error is small and it may be 
ignored for ordinary fuel testing but if a correction is to be made 
the nitric acid is titrated by standard base, at the end of the 

The heat of formation and solution of nitric acid from elemen- 
tary nitrogen is 230 calories per gram. It is convenient to use a 
tandard solution of base, 1 cc of which is equivalent to 5 calories. 
The normality of such a solution is 

= 0.3450 N. 


The number of cubic centimeters of base required to titrate the 
litric acid in the bomb after the combustion is multiplied by 5, 
:he product being subtracted from the observed calories. 

Formation of Sulphuric Acid. — A similar error results from the 
ormation and solution of sulphuric acid. In ordinary combus- 
tion of coal, organic sulphur as well as the sulphur of pyrites is 
>xidized to sulphur dioxide, which leaves the furnace as a gas. 
3n the other hand, in the calorimeter some sulphur trioxide is 
ormed and this dissolves as sulphuric acid. The higher degree 
)f oxidation of sulphur, as well as the solution of the acid that is 
ormed, yield additional heat and this also should be subtracted 
rom the observed calories. 

The application of this correction is not so simple as that for 
xidation of nitrogen. A determination of total sulphur, such 
is is ordinarily made in the analysis of coal, does not give any 



data as to the amount existing in the coal as inorganic sulphate 
which, obviously, do not develop heat through oxidation. Be 
sides, not all of the remainder of the sulphur is oxidized to th 
highest form. As the correction is usually small it is scarceb 
advisable to attempt any calculations upon such an uncertaii 

Radiation. — Radiation or absorption of heat by the calorim 
eter may be avoided by making the calorimeter "adiabatic. 
This may be done in a number of ways, three of which will be 

(1) The water in the surrounding jacket may be heated by 
electrical means, so as to keep pace with the rise in temperature 
of the calorimeter water. This is the most satisfactory method, 
although somewhat complicated and expensive apparatus is 

(2) The water in the jacket may be warmed by chemical 
action. By Richards' method a basic solution is used to fill the 
jacket and an acid is run in from a burette at a rate which 
depends upon the rate of change in temperature of the calorim- 
eter water and upon the concentration of the acid. The acid 
solution may be standardized in terms of the number of calories 
liberated by the action of each cubic centimeter upon the base, i: 
which case the proper rate of addition is more easily determine 

(3) The jacket of the calorimeter may be evacuated, on th 
principle of the Dewar flask, the transfer of heat outwardly the 
being limited to that which occurs through conductivity of th 
glass of the jacket. This would appear to be the least troubl 
some method but it has not worked well in practice. 

Radiation Corrections. — If adiabatic conditions cannot be 
maintained several methods for making radiation corrections 
are available. 

(1) The combustion may be begun as far below atmospheri 
temperature as it is to end above it. By this means absorptio 
of heat in the first half of the experiment would appear to balan 
radiation during the last half. 

This is the roughest sort of approximation and it would n 
serve for ordinarily accurate work. 

(2) The rate of change of temperature may be observed for 
certain period before firing and for another period after th 

FUELS 323 

calorimeter water has absorbed all of the heat from the bomb. 
The average of these rates is then considered to be the mean rate 
of absorption or radiation of heat for the entire experiment and 
if this is multiplied by the time elapsing between the firing and 
the maximum absorption the net gain or loss during the entire 
observation period is given. 

This method is very commonly employed and it gives a very 
[close approximation to the true correction. 

(3) Observations are made in the same way as in method (2). 
In addition the time, a, required for six-tenths of the total rise in 
temperature is observed, also the time, b, for the remaining rise. 
Instead of averaging the two radiation (or absorption) rates the 
preliminary rate, Ri, is multiplied by a and the final rate, R 2 , 
by b. The corrected rise is then 

T+R 1 a+R 2 b, 

[where T = total rise, and Ri and R 2 are regarded as positive for 
tailing temperatures and negative for rising temperatures. 

The observation of the time, a, is subject to some uncertainty 
Jwhen the temperature is rising rapidly and on this account 
pie method is not so easily applied as is method (2). It will 
[rarely be found that the difference between the corrections 
fcalculated by these two methods will differ by more than 0.2 
Ipercent and as this is well within the permissible variation 
Inethod (2) is recommended for all but the most refined work. 
I (4) The Regnault-Pfaundler method approaches theoretical 
[accuracy more nearly than any of the methods already de- 
scribed. . For a discussion of this method, see White: Gas and 
iFuel Analysis (International Chemical Series) page 224. 

Time -temperature Curves. — Three types of time-tempera- 
llure curves are produced, according to whether the experiment is 
ma) begun and finished below room temperature, (b) begun below 
jknd finished above or (c) begun and finished above. These 
Jtypes are illustrated in Fig. 76. The relative slopes of the ends 
I)f the curves represent Ri and R 2 . 

It will be observed that these slopes are easily determined in 
fcurve (6) but that it is especially difficult to decide as to what 
llemperature should be taken as the maximum produced by the 
luel combustion, in the experiment represented by curve (a). 



Conditions represented by curve (6) are to be obtained whe 

Determination: Heat of Combustion of Solid Fuels. — Place th 
lower half of the bomb in the holder, and the fuel pan in the wire suppor 
after having wired the fuse wire according to Fig. 75. 

Extend the wire across the pan, allowing it to dip sufficiently to be 
contact with the fuel, which is later to be placed in the pan. The wii 
must in no case touch the pan. The fuse wire should be placed in seri( 
with two 32-c.p. lamps in parallel when the 110- volt power circuit is use 
for firing. 


Tirr>3, Minutes 
Fig. 76. — Time-temperature curves. 

The fuel used is sampled and powdered according to directions already 

Fill a weighing bottle with the prepared sample, and weigh accurate 
to one-tenth of a milligram. Pour from this into the pan in the bom 
until the pan is approximately half full. Weigh the bottle again, an 
the difference between the above weighings gives the net quantity of 
the fuel in the bomb. This weight should be greater than 0.5 gm a 
not more than 1.2 gm. For hard coal the maximum charge should 
not greater than 1 gm. Hard coal should not be as finely divided as so: 

The upper half of the bomb is now placed in position and the nut 
screwed down as far as may be by hand, care being taken not to cr< 
the threads. The shoulder on the upper half of the bomb, over whi 

FUELS 325 

the nut makes bearing contact, should be lubricated with oil. Extreme 
care should be taken that no oil or grease is deposited on the lead gasket. 

The bomb is now ready to be filled with oxygen. The nipple is 
coupled to the oxygen piping by means of the attached hand union and 
after the connection of the bomb to the oxygen piping is accomplished 
the hand set screw on the table is tightened. In handling the bomb, 
care should be taken not to tip or jar it, as fuel may be thrown from the 

The spindle valve on the bomb is opened one turn and then the valve 
on the oxygen supply tank is very cautiously opened. The pressure 
gauge should be carefully watched and the tank valve so regulated 
that the pressure in the system shall rise very gradually. When the 
pressure reaches 300 lb per square inch, the tank valve is closed and 
the spindle valve immediately afterward. The bomb should be im- 
mersed in water immediately to detect any possible leaks. The bomb 
is now ready for the calorimeter, which is prepared as follows: 

1900 gm of distilled water, weighed or measured in a calibrated flask, 
is placed in the calorimeter can at a temperature about 1.5° below the 
jacket temperature (which should be in the proximity of the room 
temperature). The bomb is then placed in the calorimeter and the 
stirrer and thermometer are lowered into position as indicated by 
Fig. 75. The thermometer is immersed about 3 inches in the water. 
The bulb of the thermometer should not touch the bomb. 

The terminals of the electric circuit used for firing should now be 
attached. Care should be taken that neither the bomb nor the stirrer 
touches the sides of the can. The stirrer is now started and allowed 
to run 3 or 4 minutes to equalize the temperature throughout the 

Readings of the thermometer are now taken for 5 minutes (reading 
to 0.001° or 0.002° every minute) at the end of which time the switch 
is turned on for an instant only, which will be found sufficient to fire 
the charge. In course of a few seconds the temperature begins to rise 
rapidly and approximate readings are taken every minute until the rise 
becomes slow, more accurate readings then being taken. After a 
maximum temperature is reached and the rate of change of temperature 
is evidently due only to radiation to or from the calorimeter, the readings 
are continued for an additional 5 minutes, reading every minute. These 
readings before the firings and after the maximum temperatures are 
necessary in the computation of the cooling correction. The time 
elapsed from the time of firing to the maximum temperature should be, 
in no case, more than 6 minutes. 

When through with the run, replace the bomb in the holder and allow 
the products of combustion to escape through the valve at the top of 


the bomb. Unscrew the large nut and clean the interior of the bomb. 
The inside of the nut should be kept greased, also the threaded part 
at the top of the lower cup. 

Immediately after each run, the lining of the bomb should be washec 
out with a cloth moistened with a dilute solution of ammonium hydrox- 
ide and then with water. When the apparatus, after using, is to 
be left for several hours or more before making another test, the linings 
should be removed and the inner surface of the bomb slightly coatee 
with oil. This oil under the linings should be removed when nexi 
preparing the bomb for use, as an excess of it may be ignited with 
possible resulting injury to the linings. 

Heavy Oils, Coke, Hard Coal, Etc. — The determination of the hea 
of combustion of heavy oils, such as crude petroleum, and also of coke 
and extremely hard coals, is best made by mixing with a ready burning 
combustible, such as a high-grade bituminous coal or pure carbon 
This auxiliary combustible facilitates the complete combustion of the 
whole mixture in the case of coke and hard coal, and with the heavy oil it 
acts as a holder and prevents rapid evaporation of the oil. The auxiliary 
combustible should be placed at the bottom of the pan and the cok( 
coal or oil sprinkled over it. The carbon or other auxiliary combus 
tible should be dried with extreme care and carefully standardized as 
the resulting rise in temperature per gram in the calorimeter when 
completely burned. 

Calculation. — First plot a smooth curve, using temperature 
as ordinates and time as abscissas. Use only the straight portioi 
of the ends of the graph for calculating Ri and R 2 . 

The difference between the temperature at maximum and th 
temperature at firing gives directly the total rise in temperatur 
in the calorimeter. To this rise a cooling correction must b 
applied, which is computed as follows : 

The change in temperature during the preliminary 5 minute 
of reading, divided by the time (5 minutes) gives the rate of 
change of temperature per minute, due to radiation to or froi 
the calorimeter, and also any heating due to stirring, etc. Thi 
factor is Ri and in like manner the readings taken after temper* 
ture change has become uniform give R2. The two rates 
change of temperature give the existing conditions in the calorii 
eter at the start and at the finish of the run. The algebrai 
signs of Ri and R 2 will be (+) for falling temperatures and ( 
for rising temperatures. Therefore, the algebraic sum of the fr 

FUELS 327 

rates, divided by two, will give the mean value of the rate of 
change of temperature during the entire run, due to radiation or 
absorption by the calorimeter. This value multiplied by the time 
from firing to maximum will give the total cooling correction. 
The cooling correction is thus expressed : 

— o X time from firing to maximum temperature. 

This quantity is either added to, or subtracted from, the appa- 
rent rise taken from the data of the run, according to its 

The corrected rise of temperature divided by the weight of 
fuel used, will give directly the rise per gram of fuel. 

This rise per gram is multiplied by the weight of water plus the 
" water equivalent." This figure is furnished by the manufac- 
turers or it may be determined by use of a standard fuel, as 
naphthalene or cane sugar. The product is calories per gram of 
fuel, which is the result to be obtained. The result in calories 
per gram of fuel, multiplied by the factor 1.8 gives B.T.U. per 
pound of fuel. 

The final expression for fuel value is then 

(T + ^^.timeJaQOO+e) 

— q— — = cal per gm, 

where T = total rise from firing temperature to maximum, 
S = gm of coal, 
e = water equivalent of calorimeter, 

pi and R 2 having the significance already mentioned. In using 
this formula it must always be remembered that Ri and R 2 are 
regarded as radiation rates and that if the temperature is rising 
;hey must be given negative signs. 

Since cal per gmX 1.8 = B.T.U. per lb, if the latter quantity 
Is desired in most cases the product (1900+e)Xl.8 should be 
calculated at the beginning. 
j Permissible differences in duplicate determinations : 

Same analyst Different analysts 

0.3% 0.5% 


Gas Mixtures 

The separation and exact determination of gases may 
accomplished by using various gravimetric and volumetri 
methods. Certain gases may be absorbed in suitable reagents 
the absorption product being precipitated and determined gravi 
metrically. Examples of this class of methods have been met 
the determination of the halogens (page 122). Sulphur dioxi< 
may be absorbed in a basic solution, oxidized by bromine an 
precipitated as barium sulphate. Carbon dioxide has already 
been determined by absorption in a weighed solution of potassium 
hydroxide. Numerous other examples will suggest themselves 
On the other hand many gases can be absorbed by reagent 
in which they can be determined volumetrically. For example 
chlorine may be absorbed by a solution of potassium iodide and 
the liberated iodine titrated by standard sodium thiosulphate 
(page 264) ; carbon dioxide may be absorbed in a standard solu 
tion of a base and the solution titrated by a standard acid i 
presence of phenolphthalein ; sulphur dioxide may be absorbe 
and titrated by standard iodine solution, etc. 

For commercial mixtures of gases these methods are not ofte 
used because the time required for a complete analysis is too lonj 
The analysis of such mixtures as illuminating gas, natural gas 
producer or water gas, or chimney or mine gas must be made b 
more rapid methods even at a sacrifice of a degree of accuracy 
The gases from a measured volume of the original mixture ar 
absorbed in suitable reagents and the volume loss is measurec 
The results of the analysis are computed in percents by volume 

A standard type of apparatus for gas volumetry and one that 
is to be found in most laboratories is that of Hempel. 

Gas Burette. — The gas burette, in which the gas mixture is 
measured, is shown in Fig. 77. The measuring tube (a) is con- 
nected with a levelling tube (b), the gas being confined over water. 
In making a reading the water is brought to the same level in 
the two tubes so that the gas is measured at atmospheric pressure 
A complete analysis may usually be completed in a time sui 
ciently short that no serious error is caused by barometri 
changes. Changes in temperature during the course of an anah 
sis constitute the most serious sources of error. To make tl 



method even commercially accurate great care must be exercised 
in this regard. A quiet room in which no other work is being 
performed should be used. The operator must at all times avoid 
touching the burette or levelling tube directly with the hands or 


Fig. 77. — Gas burette with levelling tube. 

breathing upon them more than is necessary. Sometimes the 
burette is enclosed in a water jacket to guard against any but 
very slow changes. 

The burette may have a simple rubber tip and pinch cock at 



the top or it may be closed by a glass cock. The latter is desir 
able but is liable to become stuck by contact with basic reagent 
It must be well lubricated and frequently used or loosened. Th 
glass three-way cock at the bottom of the tube is not often use 
and is not required. 

Absorption Pipette. — The apparatus in which the gases ar 
absorbed is known as an " absorption pipette." The simples 
form of the Hempel pipette is illustrated in Fig. 78. The reagen 
fills the lower bulb, the bent tube and the capillary tube. Th 
latter is connected with the gas burette by means of a short bei 
capillary tube and when the gas mixture is forced into the pipett 

Fig. 78. — Hempel's simple 
absorption pipette. 

Fig. 79. — Hempel's double or "com- 
pound" absorption pipette. 

the absorbent fills the upper bulb. Some reagents are rapid 
altered and rendered inefficient by contact with air. Protectio 
from such action is afforded by the compound pipette (Fig. 79 
the second pair of bulbs being filled with water. It is sometimes 
necessary to insert solid reagents, such as sticks of yellow phoi 
phorus, copper wires for reducing cupric chloride, etc., or rolls 
iron gauze or glass tubes for giving greater absorbing surface 
the reagent. A pipette for solids and liquids has an opening 
the bottom of the first bulb for the insertion of such materia 
(Fig. 80.) 

In order to increase the rate of absorption modifications of th 
original pipette have been introduced. The gas is caused to 



bubble through the reagent instead of being forced down over 
the latter. After absorption has been completed the remaining 
gas is drawn from the top by turning the three-way cock to com- 
municate with the upper part of the bulb. (See Fig. 81.) 

When transferring gases from the burette to the pipette it is 
necessary to avoid mixing the water of the burette with the 
reagent in the pipette because the latter 
is thereby diluted. It is still more im- 
portant that the entrance of reagents 
into the burette should be prevented 
because such contamination of the water 
would cause premature absorption of 
gases. In order that such mixing may 


80. — Hempel's double pipette, modified 
to admit solids. 

Fig. 81. — Bubbling absorp- 
tion pipette. 

be avoided it is necessary that there be a neutral zone in the con- 
necting tubes, into which neither water nor reagent shall enter. 
If this part of the tube has any but a very small capacity 
there will be an appreciable error, due to the gas that is left in 
the tube each time. For this reason the connecting tubes are of 
capillary dimensions. 

One of the most serious disadvantages in the use of Hempel 
j pipettes comes from the necessity for connecting and discon- 
necting each pipette in turn as the different gases are absorbed. 
To obviate this inconvenience many modifications have been 



made in the direction of a composite apparatus that does no 
require the interchange. The most important feature of sue] 
forms of apparatus is a permanent connection of the burette 
with the several absorption pipettes, communication being estab 
lished with each in turn by special forms of stop cocks. Thi 
usually involves the use of longer capillary tubes and this in 
creases the error already mentioned as inherent in connectin 

In apparatus designed for the analysis of chimney gases the 

feature of permanent connection 
must be combined with porta- 
bility because the analysis must 
usually be conducted at the 
plant. A modification of the 
Orsat apparatus is here illus- 
trated (Fig. 82). 

Solubility of Gases in Rea- 
gents. — When water is used as 
the confining liquid in the gas 
burette and water solutions are 
used as absorbents in the absorp 
tion pipettes it is impossible t< 
avoid small errors, due to th 
solubility of the components o 
the gas mixture in water, 
the gases are taken into 
burette containing pure water, 
each gas dissolves and the 
volume is diminished after the 
total volume has been read. In 
order to avoid the disappear- 
ance of a part of the gases in this 
way, the water must have been previously saturated by allowing 
the gas to bubble through it. This does not entirely obviate the >; 
error because, as the mixture is drawn back into the burette for 
measurement after the removal of each constituent, the partial 
pressure of that constituent being reduced to zero, a part passe 
out of the solution in the burette and mixes with the remainin 
gases, the total observed volume being rendered too large. 

Fig. 82. — Orsat's apparatus (modi- 
fied) for analysis of chimney gases. 

FUELS 333 

illustrate this action, suppose that a mixture of oxygen, carbon 
dioxide and carbon monoxide is being analyzed. The water in the 
burette is first saturated with the mixture but the amount of each 
dissolved is a function of its partial pressure (concentration) in 
the mixture. The measured gases are passed into a pipette con- 
taining potassium hydroxide where the carbon dioxide is com- 
pletely absorbed, its partial pressure in the gases being reduced 
to (practically) zero. Upon passing the mixture of carbon 
monoxide and oxygen back into the burette a certain amount 
of dissolved carbon dioxide will be given up by the water and 
the volume will be somewhat larger than the sum of the volumes 
of the other two gases. Also where the mixture was confined 
over potassium hydroxide solution the latter dissolved small 
amounts of carbon monoxide and oxygen and some of these 
gases may be given up to mixtures later being analyzed. The 
calculation of the amount of error may thus become a compli- 
cated matter. The error is negligible, from the industrial stand- 
point, if the analysis is completed within a short period of time. 

Fuel and Lighting Gases.— In illuminating gas the following 
constituents are determined: carbon dioxide, ethylene and its 
homologues, oxygen, carbon monoxide, hydrogen, methane and 
nitrogen. They are absorbed, in the order named, one after 
another, and the contraction in volume noted after each absorp- 
tion. Hydrogen and methane are determined by combustion 
and nitrogen is computed by subtracting the sum of the other 
gases from 100. The various absorbents for these gases will be 

Carbon Dioxide. — A solution of any of the strong bases may 
be used for absorbing carbon dioxide. Potassium hydroxide 
possesses the advantage of large solubility and rapid absorption 
of gas and is almost always used for this purpose in gas analysis. 
Also potassium carbonate, formed by absorption of carbon diox- 
ide in potassium hydroxide, is more soluble in the basic solution 
than is sodium carbonate. A solution made by dissolving solid 
potassium hydroxide in twice its weight of water (about 33 per- 
cent, by weight) is suitable, this being the same strength as that 
employed for gravimetric determinations. 100 cc of a 33 per- 
cent solution will absorb about four liters of carbon dioxide 
before it becomes inefficient. The potassium hydroxide used 


should not be that which has been purified from alcohol solu-i 
tion, because traces of alcohol are retained in the solid base and 
alcohol vapor or other organic vapors are given up to the gas. 
The solution may be used in either the single or double Hempel 
pipette or in any of the modified pipettes. If the Hempel pipette 
is used it should contain rolls of iron gauze in order to increase 
the surface of solution exposed. As the solution is forced down, 
leaving the gauze exposed, the film of solution retained upon the 
surface of the wires greatly increases the rate of absorption. 

Hydrogen sulphide will be included in the fraction absorbed 
by potassium hydroxide unless it has been otherwise removed. 
Its quantity is usually small. 

Carbon Monoxide. — The most conveniently used absorben 
for carbon monoxide is a solution of cuprous chloride. Thi 
salt is only slightly soluble in water and must be dissolved in 
either hydrochloric acid or ammonium hydroxide. Either solu- 
tion absorbs carbon monoxide with the formation of a rather 
unstable compound whose exact nature is unknown. The acid 
solution is made as follows: Mix 86 gm of cupric oxide and 1 
gm of finely divided copper and slowly add to 1000 cc of a mix- 
ture of equal volumes of concentrated hydrochloric acid an 
water. Stir until the solid matter has dissolved, then place i 
bottles having bundles of copper wire reaching from top to bo 
torn. Stopper the bottles and allow to stand until colorles 
Cupric chloride, formed by dissolving cupric oxide in hydro 
chloric acid, is reduced by copper to cuprous chloride : 

CuO+2HCl->CuCl 2 +H 2 0, 
CuCl 2 +Cu-»2CuCl. 

100 cc of this solution will efficiently absorb about 400 cc of 
carbon monoxide. Absorption takes place slowly and the gas 
must be shaken with the solution for some time or be allowed 
to bubble through it. The double pipette must be used because 
cuprous chloride is readily oxidized in contact with air, cupric 
chloride being formed. 

Oxygen. — Oxygen is absorbed by a solution of potassium 
pyrogallate or by yellow phosphorus. The former solution is 
prepared by dissolving 120 gm of potassium hydroxide in 80 cc 



of water, cooling and placing in the double absorption pipette 
then adding 15 cc of a 25 percent solution of pyrogallic acid and 
mixing. Potassium hydroxide purified by alcohol should not be 
used. The solution will readily absorb about 200 cc of oxygen 
for each 100 cc of solution. It does not act rapidly at tempera- 
tures below 15°. 

Yellow phosphorus may be used in the form of sticks which are 
placed in the double pipette for solids and liquids and kept 
covered with water. This absorbent possesses the great advan- 
tage of retaining its capacity for absorbing oxygen until the sticks 
have become completely used up. The product of the union of 
phosphorus and oxygen, phosphorus pentoxide, dissolves in 
water so that the surface of the sticks is always fresh. Absorp- 
tion becomes slow below 15°, and traces of unsaturated hydrocar- 
bons of the ethylene series partially inhibit the absorption. For 
the latter reason phosphorus is not suitable for use in those 
forms of assembled apparatus for the analysis 
of chimney gases in which the ethylene hydro- 
carbons are not determined at all. 

Ethylene and Its Homologues. — These gases 
give higher illuminating power to the mixture 
of methane, hydrogen and carbon monoxide, 
gases which burn with a non-luminous flame. 
For this reason they are collectively known as 
"illuminants." Fuming sulphuric acid or 
bromine water may be used. Fuming sul- 
phuric acid reacts with members of the 
ethylene series of hydrocarbons, forming 
addition products as well as condensation 
products. These are either liquids or soluble 
solids and are therefore removed from the 
gas mixture. The absorption is not rapid and the acid should 
be shaken with the gas if the Hempel pipette, or one similar to 
it, is used. The single pipette is used, since a water seal in the 
second bulbs is inadmissible. In order to increase the contact 
of gas with acid the pipette contains a third small bulb which 
is filled with glass beads. Contact of the acid with rubber 
connections must be avoided. 100 cc of fuming sulphuric acid 
will absorb about 800 cc of ethylene. 

Fig. 83. — Fuming 
sulphuric acid pi- 
pette for unsatu- 
rated hydrocarbons. 
Gill's modification. 



Bromine water absorbs ethylene and its homologues with 
formation of bromine addition compounds: 

C 2 H 4 +Br 2 ->C 2 H 4 Br 2 . 

It is somewhat more convenient to use than fuming sulphuric 
acid but does not absorb with great readiness. If excess 
bromine is placed in the pipette the absorbing power is un< 
minished until all of this bromine has been dissolved. 

Hydrocarbon Vapors. — Gases formed by distilling coal often 
contain vapors of liquid hydrocarbons, chiefly benzene. Thes 
are partly absorbed by fuming sulphuric acid but may not 
entirely removed. They may be absorbed in absolute alcohc 
and so determined. The absorbing power of absolute alcohol 
not large and gases coming from the ordinary gas burette, being 
saturated with water vapor, soon diminish the efficiency of the 

alcohol by imparting moisture 

to it. Dennis and O'Neill 

suggested 1 the use of a soli 

tion of nickel sulphate in am 

monium hydroxide. Neith( 

this solution nor absolut 

alcohol is an entirely sati 

factory absorbent. The d< 

termination of hydrocarbo 

vapors is frequently omittec 

these vapors then being ab 

sorbed along with unsaturate 


Hydrogen. — The determi- 
nation of hydrogen is made 
by burning with oxygen, measuring the resulting contraction in 
volume, or by absorption in palladium. The combustion may be 
carried out by exploding the mixture of hydrogen and oxygen 
over mercury in a suitable pipette or the burning may be made 
to proceed more slowly. 

If 9 mixture of hydrogen with an excess of oxygen or air is 
burned the resulting water vapor condenses and only the excess 
of oxygen or air remains as gas. From the equation 
2H 2 +0 2 ->2H 2 
» J. Am. Chem. Soc, 25, 503 (1903). 

Fig. 84. — Hempel's explosion pipette. 



it is seen that two-thirds of the volume of the disappearing gas 
is that of hydrogen. Therefore, two-thirds of the contraction 
measured after cooling equals the volume of hydrogen. 

The explosion pipette is shown in Fig. 84. The confining 
liquid should not be water since larger quantities of gases will 
dissolve in it at the moment of explosion because of the momen- 
tary increase in pressure. Mercury is substituted for water and 
the pipette is so constructed as to permit altering at will the 
difference in level between the mercury in the two bulbs. If 
the ordinary single pipette were used it would be impossible to 
force gas into the pipette because of the great density of mercury 
and the consequent back pressure. Ignition is effected by con- 
necting with the secondary of an induction coil. 

Fig. 85.- 

-Pipette for the preparation 
of hydrogen. 

Fig. 86. — Pipette for slow 

If pure hydrogen is mixed with pure oxygen and burned the 
explosion is too violent for safety. If the gas to be burned is 
rich in hydrogen it is mixed with air instead of with oxygen. 
One volume of hydrogen requires more than two and one-half 
volumes of air for complete combustion, allowing a small excess. 
Dilution with air is not necessary if the residual gas is poor in 
combustibles. On the other hand it may be necessary to enrich 
the gas, before burning, by adding a measured volume of pure 
hydrogen. This is conveniently generated from zinc and sul- 
phuric acid in a special pipette (Fig. 85). 

Combustion may also be effected by passing the mixture with 
oxygen through a heated capillary tube or by exposing the mix- 




ture to a glowing platinum wire in the pipette arranged f 01 
slow combustion (Fig. 86). In using this pipette either of two 
methods of procedure may be followed: The hydrogen is placed 
in the pipette, the wire made to glow by the passage of a current 
and a measured volume of oxygen led in, or the hydrogen and 
oxygen are mixed in the burette and slowly brought into th< 
pipette, in which the wire is glowing. In either case combustioi 
occurs without explosion. 

Hydrogen may be separated from nitrogen and methane by 
absorption in palladium sponge which has been superficially 
coated with oxide. Absorption readily takes place at 100° and 
the hydrogen may be later removed by passing oxygen through 

Fig. 87. — Palladium tube. 

the palladium, the hydrogen being thereby oxidized and palla 
dium oxide again formed on the surface. A tube of the fora 
shown in Fig. 87 is used. The enlarged part is filled with asbesto 
which has been coated with spongy palladium and the tube i 
connected directly with the burette at one side and with a 
pipette filled with water at the other. Upon passing the gases 
through two or three times the hydrogen is quantitatively ab- 
sorbed, a small amount being oxidized by the trace of palladious 
oxide, and a certain amount is also burned by oxygen of air 
which was already in the tube. Except for this small amount 
of oxygen, the shrinkage in volume gives directly the volume of 
hydrogen. The amount of air in the tube must be known. This 
may be determined by connecting with the gas burette and 
measuring the expansion between two temperatures. One- 
fifth of the total volume is taken as the contraction due to con- 
tained oxygen. 

The absorption of hydrogen by palladium is hindered by traces 
of hydrochloric acid. On this account the ammoniacal solution 

FUELS 339 

of cuprous chloride should be used for the absorption of carbon 
monoxide if this is to be followed by palladium absorption of 

The explosion pipette gives fairly accurate results and is not 
difficult to manipulate but requires a battery and an induction 
coil. It is subject to the disadvantage that only a small amount 
of gas may be used, on account of the relatively large volume of 
air that must be mixed with it in the pipette, so that the error in 
reading volumes is relatively large. Gill has devised a pipette 1 
which overcomes this objection. The bulb in which the explosion 
is to take place is large enough to hold the entire residue from 100 
cc of gas, together with the necessary oxygen for the combustion, 
and is made of quite heavy glass. Both the slow combustion 
pipette and the palladium tube permit the use of larger quantities 
of gas. 

Methane. — Methane is determined by combustion, the pro- 
cedure being the same as for hydrogen. In the analysis of 
natural gas and illuminating gas, as well as many other com- 
mercial gas mixtures, hydrogen and methane will both occur in 
the residue after other gases have been absorbed. They must 
therefore be burned together unless hydrogen is to be absorbed 
by palladium. According to the equation 

CH 4 +20 2 -^C0 2 +2H 2 

i one volume of methane with two volumes of oxygen will produce 
one volume of carbon dioxide, the rest of the oxygen disappear- 
ing as condensed water vapor. The contraction is therefore 
twice the volume of the methane. Since a volume of carbon 
dioxide equal to that of the methane is produced a measurement 
I of the former by absorption in potassium hydroxide will give a 
1 direct determination of the volume of methane. For the residue 
I of hydrogen and methane, therefore, the procedure is as follows: 
An excess of air or oxygen is mixed with the gases and the mix- 
jture exploded. The gases are cooled and measured in the 
[burette, the contraction being noted. Carbon dioxide is then 
I determined by absorption in potassium hydroxide. The volume 
[ of carbon dioxide is equal to the volume of methane. Twice this 

1 J. Am. Chem. Soc, 17, 771 (1895), ' 


volume is the contraction due to the combustion of methane 
This contraction subtracted from the total contraction leaves 
the contraction due to the combustion of hydrogen. Two- 
thirds of this contraction is equal to the volume of hydrogen 
The volumes of hydrogen and methane so determined, multi- 
plied by the ratio of the total residue to the volume taken for 
explosion, gives the volumes of hydrogen and methane in the 
original gas. The following example will illustrate the calcula- 
tions involved: 

100 cc of illuminating gas gave, after all absorbable gases had 
been removed, 65.2 cc of residue, this consisting of hydrogen, 
methane and nitrogen. 15 cc of the residue was mixed with air, 
the total volume then being 90.5 cc. After explosion the volume 
was 71.0 cc. Carbon dioxide was absorbed, the volume of the 
remaining gases being then 66.6 cc. 


Volume methane = volume carbon dioxide = 71.0— 66.6 = 

4.4 cc. 
Contraction due to combustion of methane = 2X4.4 = 

8.8 cc. 
Total contraction = 90.5 -71. 0=19.5 cc. 
Contraction due to combustion of hydrogen = 19.5 — 8.8 

10.7 cc. 

Volume of hydrogen = oX 10.7 = 7.1 cc. 

Volume of methane in original. gas = -y^-X4.4= 19.1 cc. 

65 2 
Volume of hydrogen in original gas = -r^~X 7. 1 = 30.9 cc. 

{volume of sample 
— sum of volumes 
of all other gases. 

Since 100 cc of gas was taken for analysis, the volume of the con- 
stituents will also be their percents by volume. 


Analysis of Illuminating Gas. — For this exercise the Hempel appara- 
tus may be used or any of the modified pipettes or burettes may be 
substituted. The method of manipulation is not essentially different 
for the different forms of apparatus except in minor details and such 
variations will readily suggest themselves. Throughout the analysis 
avoid touching the body of the burette or the bulbs of the pipettes with 

FUELS 341 

the hands, or breathing upon them. Allow the sides of the burette to 
drain thoroughly each time before reading. 

Prepare water for the gas burette by allowing the gas to bubble 
through it for ten minutes. Fill the burette with this water, raise the 
levelling tube until the water flows out of the top of the burette, then 
close the upper cock. Place a rubber tube on a gas cock and allow 
gas to escape through it until all air is displaced. With the gas still 
running connect the tube with the top of the burette, open the burette 
cock and fill with gas until the 100 cc mark has been passed. Close 
the upper cock and detach from the gas supply. It is desirable that ex- 
actly 100 cc of gas be taken, measured at the prevailing pressure of the 
atmosphere. In order to do this allow the water to drain down the sides 
of the burette for 1 minute then raise the levelling tube, compressing 
the gas until the 100 cc mark is exactly reached. . Now close the cock 
at the bottom of the burette or close a pinch cock which is placed on the 
rubber connecting tube. Open the upper cock momentarily and close 
again. This permits gas to escape until the pressure within the burette 
is the same as that of the atmosphere. 

Hydrocarbon Vapors. — Place the pipette filled with absolute alcohol 
on the stand by the burette and connect with the burette by a bent 
capillary tube, having previously caused the alcohol to fill the lower bulb 
and the capillary up to a point near the top. Force the gas into the 
pipette, detach the latter* and shake for 1 minute. Return the gas 
to the burette, allow the water to drain down the sides, adjust the 
levelling tube to provide atmospheric pressure and measure. Record 
the difference as "hydrocarbon vapors." 

Carbon Dioxide {and Hydrogen Sulphide). — Attach the burette to the 
pipette containing potassium hydroxide solution, pass the gas into the 
pipette and directly back again. Measure and record the percent of 
carbon dioxide (including also hydrogen sulphide if present). 

Illuminants. — Determine illuminants by absorption in fuming sul- 
phuric acid or bromine water, drawing back to the burette at once. 
Avoid the entrance of any water into the fuming sulphuric acid. 

Oxygen. — Absorb oxygen by yellow phosphorus, allowing three 
minutes, or by potassium pyrogallate, shaking the pipette for three 
minutes. If pyrogallate is used in a pipette containing rolls of iron 
gauze the shaking may be omitted. 

Carbon Monoxide. — Absorb carbon monoxide in either acid or basic 
solution of cuprous chloride. The gas should be shaken with the cu- 
prous chloride solution for three minutes, then passed into the pipette 
containing potassium hydroxide to absorb vapors of hydrochloric acid. 

Hydrogen, Methane and Nitrogen— Pass all of the gas residue into the 
cuprous chloride pipette for storage, pour out the water from the 


burette and replace with water that has been saturated with air. Dete* 
mine hydrogen and methane by one of the following described methods 

Combustion by Explosion. — Return 10 to 12 cc of the gas to th 
burette, measure accurately, then draw air into the burette until j 
total volume of nearly 100 cc is obtained. Do not attempt to obtaii 
exactly 100 cc as there is danger of loss of gas during the adjustment o 
volume. Measure, then transfer the mixture of air and gas to the ex 
plosion pipette, allowing water from the burette to enter and fill th( 
capillary of the explosion pipette. Close the rubber connecting tub* 
(which should have thick walls and be securely, wired in place) with a 
screw clamp. Place the mercury reservoir bulb so that the mercury 
is at the same level as inside the explosion bulb, then connect the ter- 
minal wires with the secondary of an induction coil and cause, a spark to 
pass. A flash will pass across the bulb and mercury will almost imme- 
diately begin to flow into the bulb, on account of the contraction of gas 
volume resulting from the combustion. At all times when the gas is 
in the explosion pipette the mercury must be so adjusted in level that a 
pressure much greater or less than that of the atmosphere is avoided. 
Return the gas to the burette, allow the water to drain down the side 
then measure. Absorb the carbon dioxide and remeasure. In ore 
to be sure that an excess of oxygen was present the gas should be pas 
into the phosphorus or pyrogallate pipette. If no oxygen is foui 
the explosion must be repeated with another sample of gas, using 
larger proportion of air. Calculate the percents of hydrogen, methai 
and nitrogen by the method already discussed. Repeat with anothe 
portion of the residue in the cuprous chloride pipette. 

Slow Combustion. — Use the pipette shown in Fig. 86. Measui 
about half of the residue which is stored in the cuprous chloride pipett 
and transfer this to the combustion pipette. If the residue is known 
to be chiefly methane not more than 25 cc should be used. If it is 
chiefly hydrogen more may be taken since hydrogen requires for com- 
bustion only half its own volume of oxygen. Fill the burette with pure 
oxygen, and measure accurately. Connect the terminals of the plati- 
num wire of the pipette with a current source and heat the coil to bright 
redness. Pass the oxygen into the combustion pipette but not so rapidly 
as to cause an explosion. When the combustion is completed transfer 
the entire gas mixture to the burette and record the volume and con- 
traction. Determine carbon dioxide, test for excess of oxygen and 
calculate exactly as in the case of explosion. 

Absorption of Hydrogen by Palladium, Followed by Combustion oj 
Methane. — If the palladium 'tube is to be used for absorption of hydra- 
gen the solution of cuprous chloride in ammonium hydroxide must have 
been used for absorption of carbon monoxide. The entire gas residue 



s used. Connect the palladium tube with the burette on one side and a 
hipette filled with water on the other. The palladium tube should 
lip into a beaker of water which is kept nearly boiling. Pass the gases 
hrough the tube and back, repeating two or three times. Replace the 
lot water with water at the temperature of the room and again pass the 
;as through the tube to cool it. Determine the internal volume of the 
palladium tube as already directed and subtract one-fifth of this volume 
j'rom the total contraction. The remainder is the volume of hydrogen. 
Determine methane by combustion by either of the methods already 

Fig. 88. — Aspirator for sampling chimney gases. 

Chimney Gases. — Ideal combustion of fuel gases or of coal 
should yield waste gases containing only carbon dioxide, water 
vapor, and nitrogen. In practice complete combustion is not 
secured without a considerable excess pf air, and oxygen is there- 
fore found in the chimney gases*. The presence of carbon 
knonoxide is an indication of imperfect draught and incomplete 


combustion while a large excess of oxygen shows that heat ha 
been wasted in raising the temperature of unused air. Fo 
control work the determination of oxygen, carbon dioxide an 
carbon monoxide is sufficient and the portable form of apparatu 
(Fig. 82) is used. It contains a burette and three pipettes fo 
these determinations. Many modifications of this apparatu 
will be found illustrated and described in the scientific journa 
and trade catalogues. 

To obtain the sample for analysis a porcelain or iron tube 
inserted into the stack at the proper point. An aspirator w 
caused to draw a continuous stream of gas from the stack, the 
sample being removed by the burette as often as desired. ' The 
determination of the three gases is made as with the Hempel 
apparatus. Potassium pyrogallate should be used for the ab- 
psortion of oxygen because of the possible presence of traces of 


Burning Oils 

The chemist's examination of fuel oils usually has more to do 
■with the determination of certain physical constants than with 
[the actual analysis. Petroleum products are cheaper than animal 
pr vegetable oils and are, consequently, seldom adulterated with 
|bhe latter. Animal and vegetable oils are rarely used for burning. 
[The examination of the fuel oil, therefore, usually resolves itself 
Into a determination of the fitness of the oil for the purpose for 
which it is to be used. The determinations may include specific 
gravity, flash point, burning point and fractional distillation. 

Specific Gravity. — The relation between the specific gravity 
and the volatility of petroleum fractions is fairly definite, so 
that it is often possible to secure the correct oil by specifying 
only the specific gravity. This may be determined by means 
of a Westphal balance or a floating hydrometer. The latter is 
most conveniently used and is sufficiently accurate for most 
purposes. The specific gravity may be expressed in relation 
to water or in degrees Baume*. The system of Baume* is much 
tased in commercial testing. In this system two scales are used, 
pne being for liquids lighter than water, the other for liquids 
[heavier than water. The first is applicable to all petroleum 
[products and to most other oils and fats. 

In the original Baume* scale for liquids heavier than water the 
[point to which the hydrometer sinks in a solution of sodium 
'chloride, 15 percent by weight and at 15° C, was taken as 15°. 
The corresponding point for pure water was taken as 0° and all 
other points were located by these two. For liquids lighter than 
|water the scale has the point 10° for the density of pure water at 
Il5° C. and the point 0° corresponds to the density of a 10 percent 
solution of sodium chloride. 



Several modifications of these scales have come into use anc J 
much confusion has resulted thereby. As the system is at 
present used in many industrial laboratories the following for- 
mulas may be used for converting specific gravity into Baume* 
degrees and vice versa. 

For liquids heavier than water: 

Q 145 A 

B= 145-^ 


where B = degrees Baume and S =-specific gravity at =-= Vo* 


For liquids lighter than water : 

Q 14Q A 

S= 130+B and 

B-^-130. j 

On account of the complexity of this system and the fact that it 
is entirely unnecessary it is unfortunate that it has become so 
generally used in chemical industries. 

Flash Point. — The "flash point" is the temperature at which 
the oil gives off vapor rapidly enough that the mixture with air 
becomes explosive and will flash if a small flame is brought into 
the mixture. This is one of the most important tests to be 
applied to burning oils because it determines the degree of safety 
attending the use of the oil in enclosed vessels, such as lamps and 
burners of various kinds. In most of our States the lower 
limits of flash and burning points are specified^for kerosene by 
legal restriction. 

The location of the flash point depends to a great extent upon 
the manner of confining and heating the oil. The mixture of 
vapor with air is explosive at any temperature if the concentra- 
tion of vapor is sufficiently great. Under ordinary circumstances 
the vapor is evolved so slowly that it escapes by diffusion before 
an inflammable mixture is obtained and it is only when the tem- 
perature is raised that rapid evolution of vapor results in the 
production of a mixture that will ignite. From this it will 


[readily be seen that the flash point is lowered by rapid heating, by 
Confinement of the vapor by covering the tester, as well as by too 
Iplose contact of the test flame with the surface of the oil. It has 
(therefore become necessary to regulate by law not only the tem- 
peratures of the flash point but also the exact form of the tester 
[and the manner of heating. The following extract from the 
jlndiana law of March 11, 1901, is an illustration of the manner in 
which these tests are governed by law. 

" The test shall be made in a test cup of metal or glass, 

cylindrical in shape, two and one-quarter inches in diameter and four 
pnches deep (both measurements being made inside the cup) and this 
cup shall be filled to within one-quarter of an inch of the brim with the 
oil or other substance to be tested. The cup shall be placed in a water 
bath sufficiently large to leave a clear space of one inch under the cup 
and three-eighths of an inch around it, and in such a manner as to 
project about one-quarter of an inch above the water bath. The space 
I between the cup and the water bath shall be filled with water of medium 
j temperature and shall be heated by an alcohol lamp, with its flames so 
graduated that the rise in temperature, from 60 degrees Fahrenheit to 
the highest test temperature, shall not be less than two degrees per 
minute and shall, in no case, exceed four degrees per minute. A 
Fahrenheit thermometer shall be suspended in such a manner that 
the upper surface of its bulb shall be, as near as practicable, one-quarter 
of an inch below the surface of the oil undergoing the test. As soon as 
the temperature reaches the point of ninety-eight degrees Fahrenheit, 
the lamp shall be removed from under the water bath, and the oil shall 
then be allowed to rise to the temperature of one hundred degrees 
Fahrenheit by the residual heat of the water, and at that point the first 
test for flash shall be made as follows : A taper (hereinafter described) 
shall be lighted and the surface of the oil shall be touched with the flame 
of the taper (and it shall be lawful to apply this flame either to the center 
of the oil surface or to any or all parts of it) but the taper itself shall 
not be plunged into the oil. If no flash takes place at the temperature 
of one hundred degrees Fahrenheit, the lamp shall be placed under the 
water bath, and the temperature raised to one hundred and three degrees 
Fahrenheit, when the lamp shall be again withdrawn and the oil allowed 
to rise to one hundred and five degrees by the residual heat of the water, 
when the test shall be made by again applying the flame of the taper as 
hereinbefore specified; if no flash occurs the test shall be repeated as 
often as the oil gains five degrees in temperature, three degrees with 
the lamp under the water bath, and two with the lamp removed. These 



tests shall be repeated until a flash is obtained. The one making 
test shall further test the oil by applying the taper at every two degre 
rise without removing the lamp or stirring; but if a flash is obtained 
this means, by a less rise in temperature than five degrees herein 
quired, he shall at once remove the lamp, stir the oil, and immediatel 
apply the flame. The taper used for testing may be of any wo( 
giving a clear flame, and it shall be made as slender as possible, ai 
with a tip no more than one-sixteenth of an inch in thickness, 
taper or match with sulphur on it shall be used, unless the sulphur 
first removed before lighting. When a taper is first lighted, it shall 
be applied to the oil immediately (that is to say, before an ash or c( 
has had time to form on the end of the taper beyond the end of the flamt 
and the flame shall be made to touch the oil, but the taper itself shall 
not be brought in contact with the oil; provided, that if the taper be so 
brought in contact with the oil, but not held there longer than for the 
space of one second, and the oil flashes, the test shall not thereby be 
vitiated, but the Supervisor of Oil Inspection shall immediately remove 
the lamp, and again test the oil by the flame without allowing the body 
of the taper to touch the oil. No oil or other substance, which, by the 
test herein described, flashes at any temperature below one hundred 
and twenty degrees Fahrenheit, shall be allowed to be sold, offered for 
sale, or consumed for illuminating purposes in this State. And it sha 
be lawful to sell for illuminating purposes any oil or oils herein describee 
to be consumed within this State, which shall bear a flash test of on 
hundred and twenty degrees Fahrenheit, as shown by said apparatus 

The Indiana law is not specific in the matter of covering th 
tester and the inference is that the open tester is permitted. 

Burning Point.— The burning point ("fire test") is the tern 
perature at which vapor is evolved with sufficient rapidity t< 
sustain a continuous flame. It is determined by removing th 
cover, if one was used during the flash test, and continuing th 
heating after the flash point has been passed, applying the tes 
flame until a temperature is reached where continuous flam 
results. The thermometer bulb is immersed in the oil and th 
temperature is always noted just before the application of th 
test flame, which should be as small as possible. 

Examination of Kerosene. — Determine the specific gravity at 15° C 
with a hydrometer float or a Westphal balance, reporting in the usua 
units and also in degrees Baum6, using the formula given on page 34 
for calculation of degrees Baume\ 


Determine the flash and burning points, using, preferably, the tester 
I specified by the law of the state in which the oil was sold and following in 
; detail the directions furnished with the instrument. If no such tester 
[is available construct one as follows: Upon a small sand bath place 
f a 3-inch porcelain dish, pressing the dish into the sand until the latter 
| is within 1/4 inch of the top of the dish. Fill to the same height with 

the oil to be tested and suspend in the middle of it a thermometer. 
I Cover the dish with a watch glass having a perforation for the ther- 
i mometer and a notch at the side for the application of the test flame. 

Heat the oil so that the temperature shall rise at the rate of about 2° 
i per minute. When the temperature has reached 85° F. begin testing 
(and test for each two degrees rise in temperature by inserting a small 

flame (a gas flame 1/4 inch long) and immediately withdrawing it. 
i The experiment should be performed where the light of the room is not 
'strong and in a place free from air currents. The flash point is reached 
; when a flash passes entirely across the dish. Remove the cover and 
(continue the heating and testing until a permanent flame is sustained. 
iThis temperature is the "fire point" or "burning point. ,, 

The method used with the form of apparatus just described will not 
•I give the same flash point as that obtained by another form of apparatus 
I and cannot be used as a legal check where another form of tester is 
i specified by law. It is here described because it will afford practice 

in the determination when no other tester is available. 

Elliott Tester. — This instrument is also known as the New York 
Board of Health tester. It may be used either open or closed for 
the flash test but the closed test is preferable unless otherwise 

The outer cup is filled with water to the mark placed on the inner 
'surface. If no mark is found, entirely fill the cup with water then push 
the oil cup in as far as it will go, the excess water overflowing. Remove 
the oil cup and take out 10 cc more of water. This allows room for 
I expansion during heating. 

Flash Test. — Place the tester in a room which is free from air currents 
iand in which the light is not bright. Insert the oil cup and carefully 
j fill with the kerosene sample to within 5 mm of the shoulder, but allowing 
' no oil to splash on the shoulder. Replace the glass cover, in which the 
thermometer is fixed at such height as to bring the bulb just beneath 
the surface of the oil. Heat the water in the bath without stirring, 
fast enough to cause the thermometer in the oil to indicate a rise of 2° 
to 3° F. per minute. When the temperature reaches 85° F. begin the 
tests. The test flame should be a gas jet, not over 5 mm long. In 


making a test the flame is inserted at the notch in the cover so that it 
passes well into the cup but without, at any time, touching the surfac 
of the oil. Withdraw the flame immediately. Repeat the test a 
every 2° rise in temperature, always reading the thermometer immediatel 
before applying the flame. When a flash passes across the cup the 
flash point is reached. 

Fire Test. — Remove the cover and suspend the thermometer in 
the same position that was used during the flash test. Continue the 
heating at a rate not exceeding 10° per minute, making the tests as 
before. When the vapor finally burns with a continuous flame th< 
burning point is reached and the temperature indicated by the ther 
mometer just before the final test is called the fire test. 

Fractional Distillation. — Any fraction of petroleum now ap- 
pearing in commerce includes many different chemical com- 
pounds and can itself be separated by fractional distillation into 
other fractions having boiling points within still more narrow 
limits. The determination of the amount distilling between 
certain specific limiting temperatures yields information regard- 
ing the composition of the mixture. The results have little 
significance, however, unless the distillation is conducted in 
standard apparatus and by a standard method. 

Lubricating Oils 

For purposes of lubrication either mineral, animal or vege 
table oils or mixtures of these are used. Such an oil should have 
the proper viscosity for the purpose, should be free from acidity 
should produce the minimum of gumming under continued use 
and, if to be used as a lubricant for cylinders of internal com 
bustion engines, it must be capable of undergoing distillatio] 
without the deposition of more than a very small percent of free 
carbon. This is analogous to the " fixed carbon" of coal. In 
many cases specifications provide against the presence of more 
than small amounts of animal or vegetable oils or even agains 
any quantity, because of the gumming action that occurs by 
oxidation and because of the development of acids through par- 
tial hydrolysis of the oil. 

Viscosity. — Viscosity is usually expressed either as a specific 
property with the viscosity of water considered as unity, or in 
terms of an arbitrary scale of one of the standard instruments. 



The exact determination of viscosity is a difficult process. For 
commercial purposes an approximate determination is all that 
is necessary. The various instruments that are used for the 
determination of viscosity of oils do not give the same results 
but when the arbitrary scale of a given instrument is used, com- 

Fig. 89. — Engler's viscosimeter. 

parative results are obtained for different oils. The Engler 
[viscosimeter 1 is illustrated in Fig. 89. The principle used in 
I this and many other viscosimeters is that of measuring the time 
: required for a given quantity of oil to flow through a standard 

1 Z. angew. Chem., (1892) 725; J. Soc. Chem. Ind., 12, 291 (1893). 



Determination. — If the oil is not perfectly free from suspended solids, 
filter through muslin. If the oil is not dry, decant after long standing. 
Pour the oil into the inner cup until the points marking the required 
level are reached. Fill the outer cup with water to the mark on the 
inside and heat by the ring burner until both water and oil are at the 
desired temperature (25° unless otherwise specified). « 
wooden plug closes the gold-lined orifice in the bottoi 
of the oil cup. When this is lifted note the time on 
stop watch and allow the oil to flow out until 200 cc is 
measured in the graduated flask, noting the time when 
the graduation is reached. The viscosity is the number 
of seconds required for 200 cc of oil to flow out. The 
instrument is standardized by measuring the time 
necessary for 200 cc of water to flow at 20°. This 
should be 50 to 52 seconds. The relative viscosity is the 
ratio of the time required for the oil to that for water at 
the same temperature. 

Specific Gravity. — Determine as with burning oils 
unless the viscosity is too high to permit the use of either 
of these methods. In the latter case a picnometer is to 
be used or the specific gravity is determined at higher 
temperatures. The special hydrometer designed by 
Sommer 1 may also be used for the determination of the 
specific gravity of highly viscous oils. This is illustrated 
in Fig. 90. The brass cup has a capacity of exactly 10 cc. 
It is filled with the oil at 20°, the cap is screwed on and 
the cup is then suspended from the hydrometer float, 
which is placed in pure water at 20°. The specific 
gravity is read on the stem of the float, at the position 
of the meniscus. 

Acidity. — Shake a small amount of oil in a test-tube 
with warm water and test the water with litmus. If 
acidity is shown a weighed sample of oil is shaken with 
alcohol and the acids titrated with a standard alcoholic 
solution of a base which is preferably tenth-normal 
potassium hydroxide. 


Fig. 9 0.— 

Sommer's hy- 
drometer for 
asphalt and 
viscous oils. 

Separation of Saponifiable from Mineral Oils. — The method 
of separation depends upon the difference in chemical nature 
between mineral oils and those of animal or vegetable origin. 
The former are mostly hydrocarbons while the latter are esters 

1 J. Ind. Eng. Chem., 2, 181 (1910). 


Herived from glycerine and small quantities of other higher alco- 
hols with fatty acids. The esters are saponifiable by bases and 
the resulting soaps are soluble in water while the unsaponified 
[mineral oils easily dissolve in petroleum ether. 

! Determination. — Weigh a 100 cc Erlenmeyer flask, add about 10 
|rai of the oil and weigh again. Add 50 cc of an approximately half- 
i normal solution of potassium hydroxide in alcohol, place in the neck 
Ipf the flask a funnel having a stem not more than 5 cm long and warm 
Ipn the steam bath for 30 minutes. Remove the funnel and evaporate, 
.frequently blowing out the vapor, until the odor of alcohol disappears. 
wThe evaporation of alcohol may be hastened by inserting a glass tube 
n the flask so that the end is four or five centimeters above the liquid, 
attaching a pump and drawing air through the flask. The tube must 
pe slanted downward outside the flask in order to prevent condensed 
jilcohol from returning to the flask. Cool, add 50 cc of petroleum 
pther, stir thoroughly with the soap and rinse into a separatory funnel 
Jvith petroleum ether, disregarding any soap that may adhere to the 
Bask. Add to the ethereal solution in the separatory funnel an equal 
j/olume of water, shake and allow to separate completely. The water 
will dissolve the soap that was produced from animal or vegetable 
nils while the petroleum ether containing the mineral oil will form 
phe upper layer. Separate and discard the water solution and then 
rinse the ethereal solution into the flask in which saponification was 
accomplished, having previously washed and dried the flask. Evapo- 
rate the petroleum ether by placing the flask in a steam bath from 
which the flame has been removed. The evaporation may be hastened 
by the same device as was used in evaporating alcohol from the soap, 
piter all ethereal odor has disappeared the flask is cooled and weighed. 
JThis gives directly the percent of mineral oil in the sample, and this 
[percent subtracted from 100 gives the percent of saponifiable oil. The 
method gives somewhat high results for saponifiable oils because some 
loss of mineral oil occurs during the extraction of the soap. 

Chill Test. — The chill test is the determination of the tempera- 
lure at which turbidity appears because of the formation of 
crystals. This is the temperature at which the oil would tend to 
plog oil holes in bearings. It has little significance except where 
paponifiable oils are present because mineral oils do not crystallize 
[upon cooling. 

Determination. — A 4-ounce bottle having a wide mouth is half 
Blled with the oil and a thermometer placed in it. The bottle is placed 



in a freezing mixture and stirred continuously with the thermometer. 
When the liquid ceases to be perfectly clear the temperature is noted 
as the " chill test." 

Cold Test. — This is the determination of the temperature at 
which the oil ceases to flow freely and at which it will therefore 
fail to be delivered to bearings from the oil cup. From a con- 
sideration of the composite nature of all oils it will be seen that 
both chill and cold tests can give results which are only approxi- 
mately constant. They are of service, however, in forming a 
basis for judging the fitness of oils for use within known tempera- 
ture ranges. 

Determination. — The bottle containing the oil that was used in the 
chill test is placed in the freezing mixture and cooled until the oil becomes 
solid. It is then removed and allowed to warm by contact with the air. 
being stirred with the thermometer meanwhile. At intervals of two 
degrees rise in temperature the bottle is inverted. When the oil has 
become sufficiently fluid to flow from one end of the bottle to the othei 
the temperature is noted as the "cold test." 

Edible Fats and Oils 

For an interesting discussion of the fat and oil industries 
reference may be made to an address by Lewkowitsch. 1 

Composition. — The chief constituents of animal and vegetable 
oils are esters derived from fatty acids and the triatomic alcohol 
glycerine. Of the former the most important are palmitic, stearic 
and oleic acids, the first two being saturated acids, the last ar 
unsaturated acid. The glycerides of these acids are respective!} 
known as palmitin, stearin and oleiri and they have the folio win 

CsH5(Ci6H3i02)3 C3H5(Ci8H3502)3 C3H5(C]8H3302)3 

Palmitin Stearin Olein 

In addition to these are esters of higher alcohols other thar 
glycerine and of other saturated and unsaturated fatty acids 
also in certain cases small amounts of free higher alcohols. Th( 
chief differences in properties of different oils are caused by varia 
tions in the proportions of the constituent esters. Vegetable 
oils contain much palmitin while stearin predominates in anima 
1 Bull. soc. chim., [41 5, 1 (1909); Am. Chem. J., 43, 428 (1910). 


oils. The more liquid oils contain more olein and esters of acids 
having smaller molecular weights. 

The true waxes differ chemically from the oils and fats in that 
they are not glycerides but are esters of mono- or diatomic 
alcohols with the higher fatty acids. These alcohols are either 
aliphatic or aromatic. Some examples of such esters are as 
follows: Cetyl palmitate, derived from palmitic acid and cetyl 
alcohol, Ci 6 H 33 OH; this is the chief constituent of spermaceti. 
Ceryl palmitate, the chief constituent of opium wax, is derived 
from palmitic acid and ceryl alcohol, C27H55OH. Myricyl 
palmitate occurs in beeswax. It is an ester of palmitic acid and 
myricyl alcohol, C30H61OH. Ceryl cerotate is the chief constitu- 
ent of Chinese wax. It is an ester of cerotic acid, C25H51COOH, 
and ceryl alcohol. The most important aromatic alcohols occur- 
ring in waxes are the isomeric alcohols cholesterol and phytos- 
terol, C26H43OH. These are found as esters of palmitic, stearic 
and oleic acids. 

Notwithstanding the differences in composition the task of 
separating and determining the percent of different oils in a 
mixture is a difficult and often impossible one, because of the 
fact that the same general compounds constitute the greater 
proportion of all fats and oils. The chemist must usually be 
satisfied if he can recognize single oils or, with the nature of a 
single oil known, determine the approximate extent and nature 
of adulteration. The differences in molecular weight and degree 
of saturation, the presence and percent of free alcohols or acids 
and the occasional occurrence of traces of unusual substances, 
characteristic of certain oils, constitute the bases of the tests used 
in the effort to identify an oil. The examination becomes there- 
fore not an analysis, in the usual sense, but a series of tests applied 
in order to gain information regarding the identity of a pure oil 
land, so far as is possible, the composition of a mixture. Certain 
physical and chemical " constants" are determined and compared 
with the constants obtained from oils of known purity. The 
I chief obstacle to the use of such figures lies in the fact that, for a 
given kind of oil they are actually variable within certain limits. 
These limits may be very narrow, but since they do include a 
certain range it sometimes happens that the ranges for two or 
more oils overlap. Thus olive oil from Italy is not chemically 


identical with olive oil from California. The soil, climate, 
variety of plant and method of expressing from the olive have 
their influence upon the properties of the various glycerides 
and other substances present in the oil. It is only when the 
ranges of variation do not overlap that it is easy to determine 
the identity of a single oil, although it often happens that while 
overlapping occurs with a single constant it does not occur with 

The significance of the various constants and their methods of 
determination will be described. 

Specific Gravity. — In a general way the specific gravity of oils 
increases with the percent of (a) glycerides of unsaturated acids, 
(b) glycerides of soluble acids and (c) free fatty acids. Old 
oils also usually have higher specific gravities than the normal, 
on account of oxidation. The specific gravity of the waxes 
and of solid fats is usually higher than of liquid oils. These 
rules do not hold in all cases and the determination of specific 
gravity, like that of the other constants of oils, is made for com- 
paring with recorded data for the purpose of identification more 
often than for throwing light upon the chemical constitution 
of oils of known purity. 

Unfortunately there has been a great lack of uniformity in 
selecting conditions and modes of expression for specific gravities 
of oils as they are recorded in the literature. Temperatures of 
15.5°, 20°, 25°, 40°, 60°, 100° and others are commonly used. 
In favor of the higher temperatures it may be said that the fats 
and waxes are all liquid at these temperatures so that determina- 
tions may readily be made. The specific gravity has been vari- 
ously expressed as the weight of oil at t° divided by the weight 
of the same volume of water at either f, 0°, 4° or 15.5°. These 
differences make the compilation of comparison tables difficult. 
However it has been found 1 that a fair degree of approximation 
may be made in correcting the specific gravity to another tem- 
perature by using the coefficient 0.0007 as the change for each 
centigrade degree. This is the average value for a considerable 
number of oils between temperatures of 15.5° and 98°. Of course 
this does not remedy the lack of uniformity of expression, noted 

■ Wright: J. Soc. Chem. Ind., 26, 513 (1907). 


For the determination use a picnometer, a Westphal balance or an 
accurately calibrated hydrometer. If a Westphal balance is used the 
plummet should be accurately calibrated at the temperature at which 
the balance is to be used. The thermometer in the plummet should be 
compared with a standard thermometer. The picnometer method is 


Determination at «^o. — Use a 25 cc specific gravity bottle (picno- 
meter). Clean with chromic acid, followed by distilled water, then rinse 
with alcohol and dry in an oven at 100°. Cool in the balance case (in 
which the air should be at a temperature not above 20°) and weigh. 
Fill with distilled water which has been recently boiled to expel dissolved 
gases and cooled to a few degrees below 20°. Insert the stopper and 
nearly immerse the stoppered bottle in a distilled water bath which is 
kept at exactly 20°. After 30 minutes take off the drop of water from 
the tip of the stopper, remove the bottle and wipe perfectly dry with a 
clean towel but without warming the bottle to above 20°. Place in the 
balance case and weigh after 15 minutes. Calculate the weight of 
contained water. 

Empty and dry the bottle inside and out, then fill with oil and ma- 
nipulate as before, calculating the weight of contained oil. This weight 
divided by the weight of contained water gives the specific gravity of 

a -i . 20 °- 

the oil at oao 

If the specific gravity has been determined at any other temperature 
jor if it is desired to calculate the specific gravity at any temperature 
from the determination at 20°, these changes may be made with a fair 
degree of accuracy by the use of the following formula: 

G = G'+ 0.0007 (T'-T), where 

G = specific gravity at temperature T, 

G' = specific gravity at temperature T\ 

20° 20° 

i Determination at -jo~. — Multiply the specific gravity at 2 n6 by the 

jdensity of water at 20°, as shown in the table on page 181. The 

jproduct is the specific gravity of the oil at -jo~' 

Determination at the Temperature of Boiling Water. — Fill a 25 cc 
picnometer, dried and weighed as above described, with freshly boiled 
not water. Nearly immerse in a bath of briskly boiling water and leave 
tor 30 minutes, replacing evaporated water with boiling distilled water. 
Insert the stopper, previously heated to 100°, remove the picnometer 


from the bath, wipe dry, cool to room temperature and weigh. Cal 
culate the weight of contained water. 

Fill the flask, dried at 100°, with the dry, hot, freshly filtered fat or 
oil, which must be entirely free from air bubbles. Keep in the boiling 
water bath for 30 minutes then insert the stopper, which has been heatec 
to 100° wipe dry, cool to room temperature and weigh. Calculate the 
weight of contained oil and from this and the weight of water containec 
at boiling temperature calculate the specific gravity of the oil at the 
temperature of boiling water. 

This determination is necessarily less accurate than the one at 20° 
on account of the difficulty involved in keeping the bath at any constan 
temperature. Superheating may easily occur with distilled water am 
less pure water will have a boiling point above 100°. Variation in] 
barometric pressure will also change the temperature of the bath so 
that it becomes necessary to carry out both parts of the experiment 
at the same atmospheric pressure. However the determination is, 
sanctioned and has been made official by the Association of Official 
Agricultural Chemists. 1 

The specific gravity at any temperature other than 20° may be 
determined by the method outlined for this temperature or it may be 
calculated from the determination at this temperature, using the formula 
given above. It should be understood that the figure desired for 
purposes of identification is the specific gravity at the temperature fo 
which data may be found in the literature. 

Index of Refraction. — The measurement of index of refractioi 
is a valuable addition to the list of tests for oils. While not in 
all cases characteristic it will frequently serve to distinguish 
between certain possibilities when other tests are not conclusive 
The refractive index increases with increasing molecular weighi 
of the combined acids and with increasing unsaturation. I: 
free fatty acids are present in an oil the refractive index will b< 
lower than the normal value for the oil. In consequence of the 
latter fact one may expect to find abnormally low indices foi 
old or rancid fats or oils. 

The selection of standard temperatures for the determination 
is highly desirable in order to make comparison data useful 
Temperatures of 20° for oils and 40° or 60° for fats and waxes 
are suitable in most cases. For calculating the index of refrac- 
tion at any temperature from experimental results at another 

* J. Assoc. Off. Agr. Chem., Vol. II, No. 3, Pt. II, p. 299. 


the formula of Tolman and Munson 1 may be used: 
R = R' + 0.000365(T'-T), where 
R = index of refraction at temperature T, 
R' = index of refraction at temperature T" 


Fig. 91. — Abbe's refractometer. 

The coefficient 0.000365 is the average change of refractive 
■ndex for 1° for a large number of common oils. 

The index of refraction is determined by the use of any of the 
Standard instruments, such as the Abbe, Pulfrich, Zeiss butyro- 
Irefractometer or the immersion refractometer. 2 Of those named 

(' i J. Am. Chem. Soc, 24, 754 (1902). 
2 For a discussion of the theories of refraction and of the various types of 
efractometers, see Shook: Met. Chem. Eng., 12, 572 and 630 (1914) and 




the Abbe refractometer is probably the most generally usefi 
instrument for the laboratory because it may be used with either 
solids or liquids covering a wide range of refractive indices anc 
because it does not require the use of monochromatic light 
This instrument is shown in Fig. 91. A layer of the oil is enclosec 

between two prisms in such a manne] 
that light rays enter it at an angle differ 
ent from the normal, refraction resultinj 
(Fig. 92) . The instrument measures th 
angle of total reflection of the ray emerg 
ing from the oil, the field being a dividec 
light and dark one. Dispersion is com 
rected by a " compensator " consisting 
of two similar Amici prisms, of direct 
vision for the D-line and rotated simul- 
taneously, though in opposite directions, 
around the axis of the telescope by means 
of the screw head. In this process of 
rotation the dispersion of the compen- 
sator passes through every value fro 
zero (when the refracting edges of th 
two prisms are parallel and on differen 
sides of the optical axis) to double th 
amount of dispersion of a single Ami 
prism (the refracting edges being parall 
and on the same side of the optiqal axis 
The dispersion produced by the oil ifl 
the refractometer may thus be annulled 
by rotating the screw head of the com- 
pensator until the latter produces a dis- 
persion equal to that of the oil but in 
the opposite direction. The border lin 
between light and dark fields then becomes sharp and distinct 
even when white light is used for illumination of the refrac 
tometer prisms. 

The scale is graduated to read directly the index of refractio 
The prisms are enclosed in such a manner that water 
any desired temperature may be circulated about them. Th 
heating arrangement for the water is a special feature of the Zeiss 

Fig. 92.— Path of rays 
the Abbe refractometer. 



instruments. This is shown in Fig. 93. Water is caused to 
pass through the heating spiral and the refractometer under a 
iconstant pressure, the temperature being controlled by regulating 
|the size of the burner flame and the rate of flow through the 
(clamp C. Since the pressure of water in the labo- -& — a, 
jratory mains is not constant the pressure is made 
[independent by fixing the upper and lower levels 
by means of the overflow tubes in the vessels A 
,knd B. 

The Zeiss "butyro-refractometer" is an instru- 
ment which uses the same arrangement of prisms 
las that of the Abbe instrument. It is made espe- 
cially for use in the examination of butter and has 
i purely arbitrary scale. 
^Readings of the butyro-refrac- 
lometer can be converted 
Into indices of refraction by 
Lse of a table furnished with 
fine instrument. The chief 
liisadvantage of this instru- 
Inent is the absence of the 
compensator. The prisms are 
achromatized for pure butter 
fcnd give no dispersion of 
Ivhite light when this fat is 
lised. For all other liquids 
ihe line of division between the light and dark fields is indistinct 
Ind consists of a prismatic series of colors unless monochromatic 
fight is used. 

Fig. 93. — Zeiss' apparatus for heating 
refractometer prisms. 

Determination by Means of the Abbe Refractometer. — Set up the 

[•efractometer in front of a window or a source of sodium light. Con- 
beet the heating apparatus as shown in the figure and adjust the flow 
bf water and the height of the flame until the desired temperature (20° 
for oils, 40° or higher for fats and waxes) is attained. Open the prism 
|>o that the lower half is in a horizontal'position and place two or three 
pops of oil or melted fat upon it, using a glass rod or pipette but 
[ivoiding scratching the prisms. Quickly close and lock the prisms, 
bllow time for the temperature to become constant then adjust the 
compensator until the line of division of the field is sharply defined and 


bring this line to the cross hairs. Read the index of refraction upo 
the scale. 

Clean the prisms by applying a mixture of equal volumes of alcoho 
and ether, using a tuft of absorbent cotton. (Ordinary cotton may con 
tain grit.) 

Melting Points of Fats. — From the fact that fats are not single 
pure compounds it will be seen that they cannot have definit 
and sharp melting points and the observation will, therefor 
be a somewhat arbitrary one. The official method 1 follows. 

Determination. — Prepare an alcohol-water mixture of graduate 
density as follows: Boil, separately, water and 95 percent alcohol fo 
10 minutes to remove dissolved gases. While still hot pour the wate 
into an 8-inch test-tube until it is almost half full. Nearly fill the tub 
with the hot alcohol, pouring down the side of the inclined tube to avoi 
too much mixing. If the alcohol be added after the water has coole 
the mixture will contain so many air bubbles as to be unfit for use. 

Prepare discs of fat as follows: Allow the melted and filtered fat to 
fall a distance of 15 to 20 cm from a dropping tube upon a piece of ice 
or upon the surface of cold mercury. The discs thus formed should be 
1 to 1.5 cm in diameter and should weigh about 200 mg. Since a 
recently melted and solidified fat does not possess its normal melting 
point the discs should stand for 2 to 3 hours before testing. 

Place the test-tube containing the alcohol-water mixture in a t 
beaker, containing ice water, until cold. Drop the disc of fat into t 
tube and it will at once sink to a point where the density of the mixtu 
is exactly equal to its own. Lower an accurate thermometer, whic 
can be read to 0.1°, into the test tube until the bulb is just above th 
disc, stirring very gently with the thermometer. Slowly heat the wate: 
in the beaker, stirring constantly by means of an air blast or some other 

When the temperature of the alcohol-water mixture has risen to" 
about 6° below the melting point of the fat the disc will begin to shrivel 
and roll into an irregular mass. Now lower the thermometer until the 
fat particle is even with the center of the bulb. Rotate the thermometer 
gently and regulate the temperature so that about 10 minutes is require 
for the last increment of 2°. As soon as the fat becomes spherical rea 
the thermometer. This serves as a preliminary observation of melti 
point. Remove the tube from the bath and place in the latter a secon 
tube of alcohol-water mixture. The test-tube is of sufficiently lo 

1 J. Assoc. Off. Agr. Che m ., Vol. II, No. 3, Pt. II, p. 301. 


temperature to cool the bath to the desired point, ice water having been 
used for cooling. Add another disc of fat and regulate the temperature 
so as to reach a maximum of 1.5° above the melting point as already 
determined. Run still another determination, which should agree 
closely with the second. The disc should not be allowed to touch the 
sides of the tube in any determination. 

Melting Point of Fatty Acids ("Titer Test"). — A preliminary saponi- 
fication of the fat and separation and washing of the resulting fatty 
acids is necessary for this determination. The author considers the time 
consumed in the entire experiment to be out of proportion to the value 
of the results, in most cases. If the determination is required it may be 
found described in the official methods. 1 

Iodine Absorption Number. — The iodine absorption number 
is the percent of halogen, expressed as iodine, absorbed by the 
fat or oil when subjected to the action of a halogen solution 
under specified conditions. The absorption takes place because 
of the presence in the oil of glycerides of unsaturated acids which 
contain double or triple bonded carbon atoms. 

This action is analogous to the addition of oxygen, forming 
saturated oxygen compounds which are often hard and resinous 
in their nature. Such absorption of oxygen from the air is known 
as "drying/' although the term is here misapplied since no real 
drying occurs. The determination of halogen absorption number 
is, in a general way, a measure of " drying" properties and 
serves for the distinction between the broad, general classes of 
" drying," " semi-drying " and " non-drying" oils. 

Of the unsaturated acids whose glycerides commonly occur in 
fats and oils the following important members may be mentioned : 

Oleic Acid, C18H34O2. The structure of this acid is sufficiently 
indicated by the formula CH 3 (CH 2 ) 7 CH = CH(CH 2 ) 7 COOH. 
Olein, the triglyceride of this acid, occurs to some extent in all 
oils and fats. Olein is liquid at ordinary temperatures and its 
presence in oils is responsible, in many cases, for their liquid 
character. Either oleic acid or olein will absorb two atoms of 
bromine or one molecule of iodine monochloride or mono- 
bromide, the double bonded carbon atom thus becoming satu- 
rated. Similarly either oleic acid or olein might be expected 
to absorb one atom of oxygen and to give "drying" properties 

1 J. Assoc. Off. Agr. Chem., Vol. II, No. 3, Pt. II, p, 302. 


to a fat or oil but this action does not take place readily an< 
most of the oils of pronounced drying properties are found t 
contain considerable quantities of simple or mixed glycerides ( 
linolic or linolenic acids. 

Linolic Acid, Ci 8 H 32 02, contains two pairs of doubly linkec 
carbon atoms: 

CH 3 (CH 2 ) 4 CH = CH.CH 2 .CH = CH(CH 2 ) 7 COOH. 

This acid or its glyceride, linolin, will absorb four atoms 
halogen or one molecule of oxygen. It gives marked drying 
properties to oils, linolin being abundant in linseed, soy bean and 
poppy seed oils. 
Linolenic Acid, Ci 8 H 3 o0 2 , probably to be represented as 

CH 3 .CH 2 .CH = CH.CH 2 .CH = CH.CH 2 .CH = CH. (CH 2 ) 7 COOH. 

This acid possesses three sets of double bonds and will absorl 
six halogen atoms or three oxygen atoms. It occurs as simp] 
or mixed glycerides in linseed oil and, together with linolic aci( 
plays the most important part in the hardening or "drying 1 
of this oil when it is exposed to the air. An isomer, isolinoleni 
acid, also occurs as a constituent of the glycerides of drying oils 
Clupanodonic Acid, Ci 8 H 28 02, has the following structure: 

CH 3 .CH 2 .CH = CH.CH 2 .CH = CH.CH 2 .CH = CH.CH 2 .CJ 

HOOC(CH 2 ) 4 .CH. 

Having four pairs of doubly linked carbon atoms it is able to ! 
absorb eight halogen or hydrogen atoms or four oxygen atoms. 
This acid will be mentioned in connection with the detection of 
fish oils. 

Ricinoleic Acid, Ci 8 H 34 3 , is hydroxyoleic acid and, like oleic 
acid itself, contains only one pair of doubly linked carbon atom 
It will not readily absorb oxygen from the air and it does 
impart drying properties to an oil. It is, however, an import a 
constituent of castor oil and will be mentioned later in 
discussion of acetyl value. 


The five acids named above serve to illustrate the principle that 
only those unsaturated acids which contain more than one pair 
of doubly bonded carbon atoms are important from the stand- 
point of drying. Also an interesting, although perhaps un- 
expected fact is that trebly linked carbon atoms do not, under 
ordinary conditions, absorb halogens or oxygen to the point of 
complete saturation, only two atoms of halogen or one of oxygen 
adding to each such pair. Thus, the acids of the tariric series, 
C n H 2 n-40 2 , are isomeric with those of the linolic series. Tariric 
acid, C18H32O2, is isomeric with linolic acid. Its unsaturated 
state is shown by the formula, CH 3 (CH 2 ) 10 C=C(CH 2 )4COOH. 
This and other acids of the series, or their glycerides, absorb 
two halogen atoms quite readily but are not oxidized upon ex- 
posure to air. They are therefore quite unimportant as con- 
stituents of the drying oils. After the treble linking has been half 
saturated by the absorption of halogens the remaining double 
bond becomes saturated but very slowly and this property is 
partly responsible for the fact that, to some extent, the iodine 
absorption number is a function of the time allowed for the re- 
action and of the nature of the halogen solution. 

Many of the methods for determining iodine absorption num- 
ber have been open to the objection that they permit more or 
less substitution in saturated compounds as well as addition to 
unsaturated compounds. HubFs 1 method, formerly much used, 
is especially faulty in this respect. Hubl's solution is made by 
dissolving 26 gm of iodine in 500 cc of alcohol and 30 gm of 
mercuric chloride in 500 cc of alcohol, the two solutions being 
then mixed. The resulting solution probably contains 2 some 
mercuric chloriodide and iodine monochloride, the reaction 
being expressed as follows: 

HgCl 2 +I 2 ->HgICl+ICl. 

The latter is the active constituent of the solution but its con- 
centration is relatively small, which accounts for the fact that 
much time is required for the absorption. The oil is dissolved 
in chloroform and allowed to stand with a measured volume of 

1 Dingl. polyt. J., 253, 281 (1884); J. Soc. Chem. Ind., 3, 641 (1884). 
2 Ephriam; Z, angew. Chem. (1895), 254. 


the solution for three hours, after which the excess of iodine i 
titrated. Substitution takes place to a considerable exten 
and the amount of iodine absorbed varies with the time allowec 
When hydrogen in a saturated ester is substituted by iodine 
hydriodic acid is also formed. In the case of palmitin: 

C3H 5 (Ci 6 H3l02)3+l2— > C3H5(Ci 6 H3i02)2Cl 6 H3oI02 + HI. 

By determining the amount of hydriodic acid so formed th 
amount of substitution may be determined but it is better t 
use one of the solutions suggested by Hanus and Wijs becaus 
these cause very little substitution. 

Hanus' 1 solution is made by dissolving iodine in glacial acetic 
acid and adding an equivalent weight of bromine. The active 
constituent is iodine monobromide, IBr. The oil is dissolved in 

Wijs' 2 solution contains iodine monochloride, IC1, and is made 
by adding an equivalent amount of chlorine to a solution of iodine 
in glacial acetic acid. Either chloroform or carbon tetrachloride 
is used as the solvent for the oil. 

Both Hanus' and Wijs' solutions are more active than that of 
Hiibl, the absorption being completed in thirty minutes. Th 
solutions are also more stable and need not be so frequent! 
restandardized. The amount of substitution taking place i 
also much less and is practically zero with many oils. The sol 
tion of Hanus is more conveniently made than that of Wijs an 
the method of Hanus will therefore be described. 3 

Determination. — Prepare an iodine monobromide solution as follows: 
Equivalent quantities of bromine and iodine are dissolved in glacial 
acetic acid in such a ratio as to make a solution somewhat more than 
fifth-normal, referred to total halogen. The glacial acetic acid is first 
tested to insure absence of reducing substances. A drop of sulphuric 
acid and two or three drops of tenth-normal potassium dichromate solu- 
tion are added to 10 cc of the acetic acid and the mixture is warmec 
The yellow color should persist without the appearance of green chromiui 

1 Z. nahr. Genussm., 4, 913 (1901). 

2 Ber., 31, 750 (1898) .J 
8 S3e a comparison of methods by Tolman and Munson: J. Am. Chem. 

Soc, 26, 244 (1903); 26, 826 (1904). 


Dissolve 13.6 gm of powdered iodine in 825 cc of glacial acetic acid, 
warming the flask if necessary. Cool, decant to insure that no particles 
of iodine remain undissolved, and mix. Measure from a burette 25 cc of 
the solution into a 250 cc Erlenmeyer flask, add 10 cc of 15 percent 
sodium iodide solution and 100 cc of water and mix. Titrate at once 
with tenth-normal sodium thiosulphate solution. 

From a small burette measure 3 cc of bromine into 200 cc of glacial 
acetic acid. Mix and titrate 5 cc of the solution against sodium thiosul- 
phate solution, adding sodium iodide and water as in the iodine titration. 
Calculate the volume of tenth-normal thiosulphate solution that would 
be equivalent to 800 cc of iodine solution, then calculate the volume 
of bromine solution that would be equivalent to this volume of thiosul- 
phate solution. Add this quantity of bromine solution to the iodine in 
a glass stoppered bottle and mix well. 

Prepare starch solution, also prepare and standardize a tenth-normal 
solution of sodium thiosulphate according to the directions given on 
page 263. 

Half fill a 20 cc weighing bottle with oil, place in it a piece of glass 
rod and weigh without the stopper. Carefully pour about 0.25 gm of 
the oil into a 500 cc bottle having a ground glass stopper, using the glass 
rod to assist in the transference. Reweigh and prepare two more 
samples in the same manner. 

Dissolve the weighed sample of oil in 10 cc of chloroform then add 
25 cc of iodine monobromide solution, measuring from a pipette. Stop- 
per the bottle, mix and allow to stand for thirty minutes, shaking 
occasionally. The bottle should not be left in strong light. 

At the time that the iodine monobromide solution is measured into 
the oil solution, measure the same amount of solution into two bottles, 
containing the chloroform but no oil. Treat these in exactly the same 
manner as the solution containing oil. 

At the end of the absorption period add 10 cc of the potassium iodide 
solution which was used in standardizing the sodium thiosulphate solu- 
tion. Add 100 cc of water, washing down any iodine that may be on 
the stopper. Titrate the unabsorbed iodine with standard sodium 
thiosulphate, shaking constantly. When only a faint yellow remains 
add 1 cc of starch solution and finish the titration. At the last the bottle 
should be closed and shaken until all iodine remaining in the chloroform 
has been extracted by the potassium iodide. The temperature should 
be kept as nearly constant as possible throughout the experiment. 

From the volume of sodium thiosulphate required for the iodine solu- 
tion alone subtract that required for the oil and iodine solutions. The 
remainder is the volume corresponding to the absorbed iodine. Calcu- 
late the percent of iodine absorbed. 


Acid Value. — Fresh oils sometimes contain small amounts o: 
free fatty acids produced during the process of extraction 
Rancid fats and oils contain free acids as products of hydroly 
sis of the glycerides composing them. The acid value is defined ai 
the number of milligrams of potassium hydroxide required to 
neutralize the free fatty acids in 1 gm of oil or fat. Acidity is also 
sometimes expressed in terms of oleic acid as percent, or as "acic 
degree, " which is cubic centimeters of normal base equivalent to th 
free acids in 100 gm of oil or fat. The determination of acic 
value is made for the purpose of determining the condition of the 
oil and its fitness for a given use, rather than for the purpose o 
identifying it, since the acid value is a variable within rather wide 
limits for any oil. 

Determination. — Weigh 20 gm of oil or fat into a 200 cc flask and add 
50 cc of 95 percent alcohol which has been made neutral to phenolph- 
thalein by a dilute solution of sodium hydroxide. Heat to the boiling- 
point in a steam bath and agitate thoroughly. Titrate with a tenth-nor- 
mal solution of sodium or potassium hydroxide using phenolphthalein 
Shake vigorously during the titration and add the standard solutio 
until the pink color persists. 

Saponification (Kottstorfer) Number. — The saponification 
number 1 is the number of milligrams of potassium hydroxide 
required to saponify 1 gm of oil or fat. Different oils show differ- 
ent saponification numbers because of variation in the molecula 
weight of the esters contained in them, those of relatively lo\> 
average molecular weights requiring more base for the saponifica- 
tion of a given weight of oil than those of relatively higher mo- 
lecular weights. The variation is, however, not as great as is the 
case with iodine absorption numbers and the saponification num- 
ber is consequently not as valuable for use in identifying oils as is 
the iodine number. The comparatively small variation in saponi- 
fication number is due to the relatively small variation in th 
average molecular weight of the esters entering into the compo 
sition of the oils. Of the whole list of the more common oil 
nearly all are chiefly composed of stearin, palmitin and olein, th 
molecular weights of these being 890, 806 and 884, respectivel 
and their equivalent weights are one-third of these numbers 
The saponification numbers of the pure esters would be as follows 

1 Z. anal. Chem., 18, 199 and 431 (1878). 


56 ^fi 

for stearin ^gg X 1000=189, for palmitin ^g X 1000 = 208, 

and for olein ^gg X 1000 = 190. The greatest possible variation in 

the proportions of these three esters could make a difference of but 
19 in the saponification numbers. The occurrence of appreciable 
quantities of esters of lower acids in certain oils causes a much 
greater deviation from the numbers given above. For example, 
butter fat is chiefly composed of the following glycerides in the 
approximate proportions indicated: butyrin, CsHs^HyC^, 7.0 
percent; caproin, CsHsCCeHnC^a, caprylin, CsHsCCsHisCWa and 
caprin, C 3 H 5 (CioHi902)3, 2.3 percent; olein 37.7 percent; palmitin, 
stearin and glycerides of small quantities of other acids 53.0 per- 
cent. The calculated saponification number of butyrin is 554, of 
caproin 434, of caprylin 356 and of caprin 302. The presence of 
these esters of small molecular weight raises the saponification 
number to about 227, a number which serves to distinguish butter 
fat from a large number of other fats, particularly from oleomar- 
garine which has a saponification number of about 195. 

On the other hand the saponification numbers of true waxes are, 
on the whole, considerably lower than those of oils or fats. As 
noted on page 355, the characteristic difference between oils and 
fats, on the one hand, and waxes on the other is that the latter 
consist of esters of higher fatty acids with monohydric and di-v 
hydric alcohols instead of with the trihydric alcohol, glycerol. 
rThese esters have higher equivalent weights than those of the 
glycerides and the saponification numbers are correspondingly 

Thus, spermaceti consists chiefly of cetin (cetyl palmitate), 
P16H33O.CO.C15H31, mixed with smaller proportions of other 
psters and possibly of free alcohols. Cetyl palmitate is an ester 
of palmitic acid and the monohydric cetyl alcohol. Its theoret- 

deal saponification number is 117 ( = T^rXl000) and this gives 

(spermaceti the actual saponification number of 122 to 129, which 
|at once serves to characterize it as a true wax. 

Beeswax may be noticed as another example. This wax is 
largely composed of a mixture of cerotic acid, C25H51COOH, and 
jmyricyl palmitate, C 30 H 6 iO.CO.Ci 5 H3i. The saponification 



number of free cerotic acid is 141 ( = ™~r X1000) and that of 

myricyl palmitate is 83 ( = ^ 7 ^X1000). It might therefore 

be expected that beeswax would have an abnormally low saponi- 
fication number and this also on account of the presence of about 
10 to 15 percent of unsaponifiable hydrocarbons. The actual 
saponification number is found to be about 94. 

Water solutions of potassium hydroxide act upon oils very 
slowly because of the small solubility of most oils in water. 
Hot alcohol dissolves oils more readily and alcoholic solutions 
of potassium hydroxide are therefore used for the saponification. 
Commercial alcohol contains aldehydes which are changed by 
potassium hydroxide into resinous bodies, a dark red solution 
being produced and the basic concentration being diminished 
The alcohol should therefore be purified by first heating with a 
stick of potassium hydroxide in a flask fitted with a reflux con- 
denser and placed'on a water or steam bath, then distilling. 

Insoluble Acids (Hehner Value) and Soluble Acids. — The 
determination of the saponification number may be conveniently 
combined with the determination of soluble acids and insoluble 
acids. Among the most important of the acids of smaller 
molecular weight than oleic acid combined as glycerides ar 
butyric, caproic, caprylic and capric acids, discussed above 
These acids are soluble in water, the solubility decreasing as th 
molecular weight increases, so that, while butyric acid is infi 
nitely soluble, capric acid dissolves to the extent of 1 part in 1000 
parts of boiling water. The next acid in the series, lauric acid, 
is almost insoluble while the next member, myristic acid, is 
practically insoluble. An approximate separation of the lower 
acids from the higher ones may be accomplished by saponifying 
the oil, decomposing the resulting soap with sulphuric acid and 
washing the fatty acids with water. The percent of insoluble 
acids is called the Hehner value. l 

An inspection of the formula for a typical triglyceride, as tha 
of palmitin, C 3 H 5 (OCi 6 H3iO)3, shows that the acid residn 
comprises the greater part of the compound. Also since th 
variation in the molecular weights of the three acids, palmitic, 

1 Z. anal. Chem., 16, 145 (1877). 



i stearic and oleic, which make the greater part of the acids of 
'most oils and fats, is small as compared with the molecular 
I weights themselves, it is not to be expected that there would be a 
' large variation in either the Hehner value or the percent of soluble 
acids. The former has an average value of about 95 and the 
latter of considerably less than 1. Therefore these numbers are 
without any great significance in most cases and their determina- 
tion will give little assistance in the task of identifying most 
oils. A few exceptions to this statement .should be noticed. 

Butter has already been mentioned as containing unusually 
large quantities of butyric, caproic, caprylic and capric acids. 
Consequently its Hehner value falls to 88-90 and its percent of 
soluble acids rises to about 5. Other notable exceptions are 
\cocoanut, palm nut, croton and porpoise oils. Practically, it is 
in these cases only that the determination of soluble and insoluble 
acids will be of any great use. 

Determination. — Prepare a tenth-normal solution of sodium or potas- 
sium hydroxide in boiled and cooled water, using phenolphthalein in 
standardizing. Purify two liters of alcohol by heating on a steam bath 
for thirty minutes with about 10 gm of sodium hydroxide, using a reflux 
condenser. Distill and make 1000 cc of a solution of 40 gm of potassium 
hydroxide in the alcohol. The potassium hydroxide should be as nearly 
free from carbonate as is possible. Allow the solution to stand until 
the small amount of potassium carbonate that is always present has 
settled out, then decant into another bottle. The concentration does 
I not remain constant for long and the solution need not be standardized 
until it is used for saponifying the oil. 

Prepare also a half-normal solution of hydrochloric acid in water. 

Saponification Number. — Select two ordinary flasks of 250 cc capacity 
having, if possible, necks of slightly larger diameter at the top than at 
I the bottom, though this feature is not essential. Clean with alcohol. 
| Weigh into each flask about 5 gm of oil or fat, using a small bottle and 
'glass rod as in the determination of iodine number. Add to each flask 
50 cc of the alcoholic solution of potassium hydroxide from a calibrated 
pipette or burette, place in the neck of the flask a funnel having a short 
stem and warm on the water bath until the alcohol boils, though it 
should not be evaporated more than is necessary. The oil is usually 
saponified in about thirty minutes. A homogeneous solution must be 
produced, so that no separation will occur when boiling is interrupted. 
Measure 50 cc of the alcohol solution of potassium hydroxide into each 





of two other flasks, for standardization. While saponification of the 
is proceeding titrate these solutions with the half-normal acid, usin 
phenolphthalein. Cool the flasks in which the oil was saponified, add 
drop of phenolphthalein and titrate the excess of base with half-nor: 
acid, deduct from the volume used for 50 cc of the base in the standa 
ization and calculate the saponification number. Preserve the neu 
solution for the determination of soluble and insoluble acids. 

Soluble Acids. — Evaporate the alcohol by placing the flasks upon 
water bath and drawing air through them as explained on page 352 
When the odor of alcohol has entirely disappeared add enough standan 
acid to make the total amount, including that used in titrating the excess 
of base, 1. cc more than the volume that is equivalent to the 50 cc o; 
potassium hydroxide used in saponifying the oil. It is necessary to b( 
very careful about the removal of all alcohol at this point. The acid; 
that are classed as insoluble are considerably more soluble in alcohol 
Consequently any alcohol that might be left with the soap would cause 
an error in the separation of soluble from insoluble acids. 

Connect a reflux condenser and warm on the water bath until th( 
insoluble fatty acids have melted and separated from the water solution 
Add hot water to bring the liquid within about 2 cm of the top of th( 
neck of each flask, again allow the insoluble acids to separate, then coo 
in ice water. Carefully detach the cake of insoluble acids and pour th< 
cold solution through a filter into a flask of 1000 cc capacity. Repl 
the cake of acids in the flask, fill with hot water, separate as bef 
and filter. Repeat the treatment once more and do not discard 
insoluble acids. 

In some cases the insoluble acids will not solidify, even at 0° 
account of the preponderance of oleic acid, a liquid acid. In su 
cases the entire liquid is poured into an already wet. filter and the flask 
and the insoluble acids on the filter are washed with cold water. 

The combined filtrates now contain the excess of half-normal a 
(1 cc) and the soluble acids of the oil, besides potassium chloride 
glycerine and other alcohols, etc. Titrate the acids with tenth-norma 
base in presence of phenolphthalein. From the volume of standard bast 
used deduct the volume equivalent to 1 cc of, the standard acid and 
culate the percent of soluble acids as butyric acid. The arbitr 
assumption that butyric acid is the only soluble acid present is me: 
a convenience. 

Insoluble Acids (Hehner Value). — Allow the cake of insoluble aci 
dry on the filter paper for twelve hours at the temperature of the roo 
then transfer to a small weighed dish. Wash thoroughly with waJ 
alcohol the paper and the flask in which saponification was accomplished 
allowing the solution to run into the dish. Evaporate the alcohol on 




im bath and dry to constant weight at 100°. From 2 to 5 hours dry- 


ling will usually be required. 

If the insoluble acids did not solidify the filter is pierced and the 
: acids are allowed to run into a weighed dish. The flask and paper are 
!i washed thoroughly with hot alcohol, this running into the dish. Evapo- 
rate and dry as with solid acids. 

Calculate the percent of insoluble acids. 

Reichert Number and Reichert-Meissl Number. — There is no 
sharp line of division between the fatty acids volatile with steam 
knd those not volatile and it is not possible to effect more than a 
very approximate separation by a method of distillation unless 
fchis is continued for a very long time. On the other hand 
fairly constant proportions of acids may be distilled if the method 
is rigidly standardized. In this way figures may be obtained 
that have a value in identifying certain oils and fats. The 
determination is made chiefly in the examination of butter and 
tts substitutes. Pure butter contains volatile acids to the extent 
pf nearly 10 percent of the total fatty acids. 

The saturated acids that have been classed as " soluble" (to 
and including capric acid) are the only ones of the series that may 
be distilled without decomposition. They are therefore known 
as " volatile" acids while the higher acids (above lauric) decom- 
pose when distilled and are therefore called " no n- volatile." 
Lauric acid distills with steam but is slightly decomposed. Al- 
though the volatile acids boil at temperatures higher than 100° 
they can be distilled with steam. The boiling points of the more 
commonly occurring "volatile" acids are as follows: 


Boiling point, degrees 






The method proposed by Reichert 1 and the modifications of 
tiis method by Meissl 2 have been extensively adopted. It 
tiould be understood that neither method gives the correct 
ercent of volatile acids but simply the proportion that will be 

1 Z. anal. Chem., 18, 68 (1879). 

I Dingl. polyt. J., 233, 229 (1879); Chem. Zentr., 10, 586 (1879). 



distilled under certain stated conditions. The Reichert Numbei 
is the number of cubic centimeters of tenth-normal base required 
titrate the acids obtained from 2.5 gm of oil or fat by Reichert 1 s 
distillation process. The Reichert-Meissl number is the same 
as the Reichert number except that 5 gm of oil or fat is used. Ttu 
Reichert-Meissl number is not exactly double the Reicherl 

The Reichert-Meissl number of most oils, fats and waxes is 
less than 1 and the determination will be of little service ii 
identifying these oils. The following oils are exceptional ii 
this respect. 

Oil or fat 

Reichert-Meissl number 

Butter fat 


7 . 








Determination. — Prepare the following reagents: 

(a) Sodium hydroxide solution in water, 50 percent by weight. 

(b) Alcohol, 95 percent, redistilled from sodium or potassiur 

(c) Sulphuric acid, 1 part concentrated acid in 5 parts water. 

(d) Potassium hydroxide, approximately tenth-normal, standardize 
against standard acid, using phenolphthalein as indicator. 

If the sample is either real or imitation butter it will contain wate 
and curd. Melt and keep at 60° until the fat has separated and, i 
necessary, filter the fat through a dry paper placed in a hot-wate 
funnel. If the sample is an oil it may usually be weighed withou 

Ordinary flasks of 200 cc capacity, are cleaned and dried. The oil 
or melted fat is dropped in from a weighed bottle until 5 gni, measured 
to within one drop, is obtained. The oil must not be left on the neck of 
the flask. Record the exact weight. Add 10 cc of alcohol and 2 cc of 50 
percent sodium hydroxide solution, connect with a reflux condenser and 
heat upon the steam bath until the oil is saponified. Remove the con- 
denser and evaporate the alcohol as in the determination of soluble acids 
Add 135 cc of recently boiled water and warm on the water bath until 
solution is complete, then cool. Add two or three pieces of pumice 
stone or about 1 gm of crushed porcelain to prevent bumping, then add 
5 cc of the diluted sulphuric acid. Again attach the reflux condenser 



and heat on the steam bath until the acids form a clear layer. Connect 
the flask with a Kjeldahl or Hopkins distilling tube and a condenser 
and distill over a flame at such a rate that 110 cc shall be obtained in 
approximately thirty minutes. The 
distillate is received in a flask 
which is graduated to contain 110 
cc. Mix the distillate, and filter 
through a dry filter to remove 
traces of insoluble acids carried 
over by the steam, receiving the 
filtrate in a flask graduated to con- 
tain 100 cc. Titrate 100 cc of the 
filtrate with standard potassium 
hydroxide. Make the proper cor- 
rection for the fact that only 100 cc 
of the distillate was used, also cor- 
rect the number of cubic centimeters 
of standard potassium hydroxide 

used, in case this solution was not . ^ 

exactly tenth-normal or in case the Fig. 94.— Kjeldahl's distilling tube, 
sample weight was not exactly 5 
gm. The result is the Reichert-Meissl number. 

Polenske Value. 1 — One of the very important constituents 
of some butter substitutes is cocoanut oil, a pure white vegetable 
fat having a pleasant taste and a consistency which is about the 
same as that of butter. Its Reichert-Meissl number is somewhat 
lower than that of butter, as is shown in the table on page 374. 
Salkowski 2 noticed that the volatile acids obtained from cocoanut 
oil in the Reichert-Meissl distillation contained much larger 
quantities of acids insoluble at 15° than do the volatile acids 
from butter. Butyric acid comprises from 60 percent to 70 
percent of the volatile acids from butter and this acid is soluble 
in water in all proportions. The volatile acids from cocoanut 
oil contain larger quantities of caproic, caprylic, capric and lauric 
acids, 3 these being almost insoluble at 15°. The Polenske 
value (called by its author the "new butter value") is the num- 

1 Z. Nahr. Genussm., 7, 273 (1904); J. Soc. Chem. Ind., 23, 387 (1904). 

2 Z. anal. Chem., 26, 581 (1887). 

3 Elsdon: Analyst, 38, 8 (1913). Percents of acids here reported are 
caproic 2, caprylic 9, capric 10. lauric 45, myristic 2, palmitic 7, stearic 5, 
and oleic 2. 


ber of cubic centimeters of decinormal base required to titrate the 
insoluble acids obtained in the Reichert-Meissl distillation. 

The Polenske value for pure butter varies from 1.5 to 3.0, 
while that for cocoanut oil varies from 16 to 18. 

It is necessary to avoid the use of alcohol in the saponification 
of the fat and therefore the determination of Reichert-Meissl 
number must be modified if the two determinations are to be 
combined. Polenske's modification is essentially as follows: 

Determination. — Saponify 5 gm of the fat by heating in a 300-cc round 
flask, using a reflux condenser. For the saponification use 20 gm of glyc- 
erol and 2 cc of a 50 percent solution of sodium hydroxide in water. 
When saponification is complete dissolve the soap in 135 cc of recently 
boiled water and add 5 cc of dilute sulphuric acid (25 cc of concentrated 
acid in 1000 cc of solution) and a small amount of crushed porcelain 
or pumice. Connect with a condenser by means of a Kjeldahl or Hop- 
kins distilling tube and distill into a flask which is graduated at 100 cc 
and 110 cc; the distillation should proceed at such a rate that 110 cc 
passes over in about 20 minutes. When the distillate reaches the 110 cc 
mark on the flask replace the latter by a 25 cc cylinder and stop the 
distillation. Immerse the flask in water at 15° and allow to remain for 
15 minutes. The level of the water must be above the 110 cc mark on 
the flask. Mix the contents of the flask and pass through a dry, 8 cm 
filter and, if desired, determine the Reichert-Meissl number, using 100 cc 
of the filtrate. Rinse the 110-cc flask but without removing any of the 
insoluble acids adhering to it. Wash the filter three times with 15 cc 
of water, this water having previously been used for washing the con- 
denser, cylinder and flask. Dissolve the insoluble acids from the con- 
denser, cylinder and filter, using three successive portions of neutral 
90 percent alcohol and allowing the solution to run into the 110 cc flask. 
Titrate the alcoholic solution with decinormal potassium hydroxide 
solution, using phenolphthalein, and calculate the Polenske value. 

Acetyl Value. — Compounds containing a hydroxyl group will 
readily combine with acetic anhydride, acetic acid and an acetyl 
compound being produced. This takes place with an oil con- 
taining free higher alcohols or hydroxy-acids, the latter either in 
the form of esters or of free acids. The general reaction may be 
thus shown: 



ROH + (CH 3 CO) 2 0->ROCH 3 CO + CH3COOH. 


For example lanopalmic acid forms acetolanopalmic acid: 

C 15 H3oOHCOOH+(CH3CO)20->Ci5H3oOCH 3 COCOOH+ 


After washing out the excess of acetic anhydride the amount 
absorbed may be determined by saponifying the oil with an 
alcohol solution of potassium hydroxide, evaporating the alcohol, 
adding standard sulphuric or hydrochloric acid to liberate the 
acetic and fatty acids and either distilling the acetic acid or 
washing out with water, then titrating. The reactions illustrated 
by the case of aceto-lanopalmitin are 

C,H 6 (OCi6H8oOCHaCOCO),+6KOH^C,H 6 (OH),+ 


2Ci 5 H3oOHCOOK+H2S0 4 ->2Ci 5 H3oOHCOOH+K 2 S04, 
2CH3COOK + H 2 SO-4->CH 3 COOH + K2SO4. 

Effect of Soluble or Volatile Acids. — It should be noticed that 
whether the distillation or the nitration process is employed, 
the standard base required to finally titrate the acid will include 
that equivalent to acids other than acetic. That is, the distilla- 
tion process will yield a distillate of acetic acid and volatile or- 
ganic acids while the filtration process will yield a filtrate con- 
taining acetic acid and soluble organic acids. The close relation 
between soluble acids and volatile acids has already been dis- 
cussed (page 373). To correct for the presence of these acids in 
the solution containing the acetic acid one may either subtract 
the volume of base used in the determination of soluble (or vola- 
tile) acids, or a different method may be used. As a rule this 
correction will be small but with oils showing a high soluble- 
acid number or Reichert-Meissl number, failure to apply the cor- 
rection may result in serious error. 

Benedikt and Ulzer 1 proposed first saponifying the oil and then 
liberating the fatty acids by the addition of sulphuric acid. After 
washing the fatty acids they are acetylated and the excess of 
acetic anhydride removed. The acids are then titrated in cold 
alcoholic solution, under which circumstances the carboxyl 
alone reacts with the base. The acetylated soap is then heated 

1 Monatsh., 8, 41 (1887). 


with alcoholic potassium hydroxide when the acetyl radical is 
saponified. A titration of the excess of base gives the acetyl 
value. This method avoids the interference of soluble fatty 
acids but, as was shown by Lewkowitsch 1 it is subject to another 
error in the fact that acetic anhydride also reacts, to a small 
extent, with non-hydroxylated fatty acids forming acetic acid 
and fatty acid anhydride: 

2C 15 H3iCOOH+(CH3CO) 2 0->(Ci5H3iCO) 2 0+2CH 3 COOH. 

Palmitic acid Acetic anhydride Palmitic anhydride Acetic acid 

In cold alcoholic solution these anhydrides are not at once saponi- 
fied, part remaining until the treatment with hot potassium 
hydroxide solution, being then saponified and giving rise to a 
positive error in the calculation of acetyl value. 

The most desirable method is to acetylate the oil, wash free 
from acetic acid, saponify, liberate the fatty acids and acetic 
acid from the soap and then either distill or filter, titrating the 
acids of the distillate or filtrate and making the proper correction 
for volatile or soluble acids. 

The " acetyl value" is defined to be the number of milligrams of 
potassium hydroxide required to combine with the acetic acid lib- 
erated from 1 gm of acetylated fat or oil. Certain oils are charac- 
terized by unusually high acetyl values. Castor oil is the most 
noteworthy of these, having a value of about 150. Another class 
of oils having high acetyl values is composed of " blown" or 
" oxidized" oils. By blowing air through oils at somewhat ele- 
vated temperatures (70° to 115°) the viscosity and specific 
gravity are considerably increased and they become suitable for 
use as lubricating oils. The chemical changes that take place 
are not thoroughly understood but oxidation is known to occur, 
This is partly due to combination with unsaturated acids (evi- 
denced by a diminished iodine absorption number) and partly 
to the formation of hydroxyl radicals from hydrogen. The 
latter change results in an increased acetyl value and this may 
even reach a number as great as that for castor oil. 

The large variation in acetyl values recorded in the table on 
page 379, adapted from a similar table by Lewkowitsch, 2 will 

1 Proc. Chem. Soc, 6, 72 (1890). 

2 J. Soc. Chem. Ind., 16, 503 (1897). 



indicate the value of this determination for the identification of 
certain oils and fats. In other cases the determination will 
have little value. 

Oil or fat 

Acetyl value (average) 

Butter fat 












Colza . 

Cotton seed 






Shark liver 

Abnormal Variation in Acetyl Values. — Certain abnormalities 
in acetyl values should be noticed and due allowance made in 
specific cases. 

Since acetic anhydride is absorbed by the hydroxyl radical 
it might be expected that free acids, free alcohols or partially 
hydrolyzed glycerides or other esters would show such absorption 
and that their occurrence in oils or fats would cause these to ex- 
hibit unusually high acetyl values. This is found to be the case 
and, since the three classes of substances named above are the 
direct products of hydrolysis, it follows that rancid oils or fats 
will not give normal acetyl values. For example, hydrolysis 
of stearin will yield free stearic acid, together with distearin, 
monostearin or glycerol, according to the degree of hydrolysis: 

C3H5(OC 1 8H 3 50)3 + H 2 0-^C3H 5 OH(OCl8H350) 2 +Cl 8 H3602, 

C3H 5 OH(OCi 8 H350) 2 +H 2 0->C3H5(OH) 2 OC 18 H350+Ci8H360 2 , 
C 3 H 5 (OH) 2 OCi8H350+ H 2 0-^C 3 H 5 (OH) 3 + Ci 8 H M 0,. 

Each of these reactions produces a hydroxy lated compound, 
which is capable of combining with acetic anhydride. The acetyl 
values of these substances are as follows: 

Compound Acetyl value 


Monostearin .... 



The free acids will also combine with acetic anhydride to a 
varying degree and this property will still further increase the 
acetyl value of rancid materials. 

The measurement of acid value is a convenient method for 
determining this property. In case high acid values have been 
obtained the proper correction should be made in the acetyl 
value as the latter has been determined experimentally. In 
other words it is only in fresh oils that acetyl values can be used 
with certainty for identification purposes. 

The most important application of this determination is in 
the identification of castor oil. This oil is nearly pure ricinolein, 
a glyceride of ricinoleic acid. The latter is hydroxylated oleic 

CH 3 (CH 2 ) 5 CH.OH.CH 2 .CH = CH(CH 2 ) 7 COOH, 

and the glyceride, ricinolein, has a theoretical acetyl value of 
159.1. Its abundance in castor oil gives the latter an actual 
acetyl value of about 148, a value which is far above that of 
any other natural oil, only blown oils approaching it in this 

Lastly may be mentioned the occurrence of certain quantities 
of free alcohols, especially in the waxes which have, on this ac- 
count, appreciable acetyl values. Cholesterol, C27H46OH, in fats, 
oils and waxes of animal origin, and its isomers, the phytosterols, 
in vegetable oils, etc., are the most important of such alcohols. 

The method here described for the determination of acety 
value is essentially that of Lewkowitsch and adopted as a pro- 
visional method by the Association of Official Agricultura 
Chemists. 1 It involves the acetylation of the oil before saponi- 
fication and includes the soluble or volatile fatty acids, if calcu- 
lated according to the " official" method. Much confusion wouli 
be avoided if the true acetyl value were recorded instead of this 
acety 1-soluble acid value. In the following exercises the true 
acetyl value will be calculated. 

Determination. — Place about 20 gm, approximately weighed, o; 
oil or fat in a 100 cc flask, add an equal volume of acetic anhydride 
insert a short-stemmed funnel and boil gently for two hours. Cool an( 
pour into 500 cc of water contained in a beaker. Pass a current of car- 

1 U. S. Dept. of Agr., Chem. Bull. 107, 142. 


>on dioxide into the beaker through a fine orifice of a glass tube and boil 
or 30 minutes. At the end of this time siphon out the water layer and 
epeat the treatment with water and boiling until the water is no longer 
icid, as shown by a litmus test. Separate the acetylated oil in a separa- 
ory funnel, filter in a drying oven and dry. 

Weigh accurately 2 to 4 gm of the acetylated oil into a flask and 
aponiiy according to the method used in determining the saponification 
lumber, measuring the alcohol solution of potassium hydroxide accu- 
ately and running blank determinations for standardization. Evap- 
>rate the alcohol and dissolve the soap in water. Add standard sul- 
furic acid in a quantity exactly equivalent to the potassium hydroxide 
idded, warm to melt the fatty acids and filter through a wet paper. 

i( Wash with boiling water until the washings are no longer acid, testing 
vith litmus paper by barely touching a corner to the bottom of the 
unnel. The combined filtrate and washings are titrated with tenth- 
lormal base. Subtract the volume of base already found to be equiva- 
ent to soluble acids and calculate the true acetyl value according to 

at ;he definition of this number. 

Maumene Number and Specific Temperature Reaction. — All 

>ils and fats react with concentrated sulphuric acid, heat being 
jvolved. The reactions are complex and cannot be expressed 
)y a simple equation but oxidation occurs to a considerable 
legree. The heat evolution varies with different oils and is, to 
some extent, characteristic. The Maumene number 1 is the 
lumber of centigrade degrees rise in temperature caused by mixing 
L0 cc of concentrated sulphuric acid with 50 gm of oil. A small 
variation in the proportion of water in the acid causes a con- 
siderable variation in the heat evolved and to this extent the 
igures recorded by different investigators are not comparable 
)ecause "concentrated sulphuric acid," as obtained commer- 
jially, is not a substance with any definite percent of water. 

In order to eliminate the errors due to variation in water a 
letermination may be made, using the same amount of acid but 
substituting 50 gm of water for the oil. The ratio 

Rise in temperature with oil 

Rise in temperature with water 
s known as the " specific temperature reaction." 2 That this 

1 Compt. rend., 35, 572 (1852). 

2 Thomson and Ballantyne: J. Soc. Chem. Ind., 10, 233 (1891). 



number is not subject to variation as is theJVlaumene number ig 
shown by the following table in which the specific temperature 
reaction is multiplied by 100. 

Kind of 

Sulphuric acid of 95.4 

Sulphuric acid of 96.8 

Sulphuric acid of 99 



Sp. temp, 



Sp. temp, 


Sp. temp 























Determination. — Place a beaker, about 5X1.5 inches, inside one that 
is about 6X3 inches and pack the open space between with wool, 
asbestos or cotton. Cover the beakers with a piece of cardboard 
through which passes a thermometer. Weigh into the inner beaker 50 
gm of oil. Bring concentrated sulphuric acid to the same temperature as 
that of the oil and then add. under a hood, 10 cc of this acid, stirring 
thoroughly with the thermometer. When the acid is all in, place the 
thermometer in the center of the oil-acid mixture and note the highest 
point attained by the mercury. The total rise in temperature is the 
Maumene* number. 

Determine also the specific temperature reaction as follows: Clear 
the inner beaker and introduce 50 cc of water. Add 10 cc of acid 
before and note the rise in temperature. The Maumene* number divide< 
by this rise is the specific temperature reaction. 

The drying oils often develop so much heat that active foaming 
results. Such oils should be first diluted with petroleum oils or olive 
oil and the proper correction made in the temperature rise. 

Qualitative Reactions. — If simple and reliable qualitative 
tests were known for all of the oils, it is not likely that the work 
outlined in the preceding pages would often be carried out. It 
has already been explained that comparatively few such tests 
are known because of the similarity in the composition of th 
various animal and vegetable oils. Aside from the mere vari 
tion in the proportion of the various glycerides, free alcohols an 
free acids, there are certain constituents of certain oils that wi 
give color reactions which are characteristic. A few of thos$ 
that are reliable will be described. In most cases these tests 





[should accompany the determination of the analytical constants, 
jrather than be substituted for them. 

, Resin Oil. — Polarize the oil in a 200-mm tube. If the oil is too dark 
in color for this purpose it may be diluted with petroleum ether and the 
proper correction made in the reading. Resin oil has a polarization in 
ja 200-mm tube of from +30° to +40° on the sugar scale (Schmidt and 
Haensch) while other oils read between +1° and —1°. 
I Cotton Seed Oil: Halphen Test. 1 — Mix carbon disulphide containing 
about 1 percent of sulphur in solution, with an equal volume of amyl 
alcohol. Mix equal volumes of this reagent and the oil and heat in a 
■bath of boiling, saturated solution of sodium chloride for 1 to 2 hours. 
[In the presence of as little as 1 percent of cotton seed oil a character- 
istic red color is produced. Lard and lard oil from animals fed on 
Icotton seed meal will give a faint reaction for cotton seed oil. The 
■unknown constituent which gives the color apparently is assimilated 
[by the animal without change. 

A negative result does not prove the absence of cotton seed oil because 
meating the oil for 10 minutes at 250° renders it incapable of giving the 

Arachis (Peanut) Oil.- — Modified Renard 2 Test. — This test is 
[based upon the fact that about 5 percent of " crude arachidic 
■acid" may be isolated from arachis oil, whereas stearic acid is the 
[highest acid that occurs in any considerable quantity in most 
loils and fats. " Crude arachidic acid" is a mixture of true 
larachidic acid, C20H40O2, and lignoceric acid, C24H48O2. 

The oil is first saponified and the excess of base is neutralized 
[with acetic acid. Lead acetate is then added and the lead 
[soaps of the higher acids are separated from those of the lower 
[acids by washing with ether, in which lead soaps of the soluble 
lacids dissolve. The insoluble soap is decomposed by hydro- 
chloric acid and the resulting palmitic, stearic, arachidic, ligno- 
Iceric and traces of other higher fatty acids are extracted with 
■ether, which is then evaporated. Finally the acids are dissolved 
lin 90 percent alcohol and the solution is cooled to 15°. Stearic 
[and palmitic acids remain in solution and the higher acids 
[crystallize, leaving a saturated solution containing 0.00025 gm 
of " crude arachidic acid" in each cubic centimeter of alcohol. 

1 J. pharm. Chim., [61, 6, 390 (1897). 

2 Z. anal, chem., 12, 231 (1871); Compt. rend., 73, 1330 (1871). 


By applying this solubility correction the approximate weight 
of arachis oil in a mixture can be calculated. The test is made 
as follows: 

Weigh 20 gm of the oil into a 250 cc Erlenmeyer flask. Saponify wit] 
a solution of potassium hydroxide in alcohol as directed in the discussion 
of the determination of saponification number. Add a drop of phenol 
phthalein and exactly neutralize with 5 percent acetic acid and was] 
into a 500-cc flask containing a boiling mixture of 100 cc of water ant 
120 cc of a 20 percent lead acetate solution. Boil for a minute and then 
cool the precipitated lead soap by immersing the flask in water, occa 
sionally giving it a whirl to cause the soap to stick to the sides of the flask 
After the flask has cooled, the water and excess of lead acetate can be 
poured off and the soap washed with cold water and with 90 percent 
(by volume) alcohol. Add 200 cc of ether, cork and allow to stand 
for some time until the soap is disintegrated; heat on the water bath, 
using a reflux condenser, and boil for about five minutes. In the oils 
most of the soap will be dissolved, while in lards which contain much 
stearin, part will be left undissolved. Cool the ether solution of soap 
to from 15° to 17° and let stand until all the insoluble soaps have crys- 
tallized out (about twelve hours). 

Filter upon a Buchner funnel or a folded filter and thoroughly wash 
the insoluble lead soaps with ether, then wash them into a separatory 
funnel by means of a jet of ether from a wash bottle, alternating at the 
end of the operation, if a little of the soap sticks to the paper, with hy 
drochloric acid, 10 percent solution. Add sufficient of this dilute hydr 
chloric acid so that the total volume of the acid layer amounts to abou 
200 cc and enough ether to make the volume of the ether layer 150 to 
200 cc and shake vigorously for several minutes. Allow the layers 
separate, run off the acid layer and wash the ether once with 100 cc 
dilute hydrochloric acid and then with several portions of water un 
the washings are no longer acid to methyl orange. If a few undecom- 
posed lumps of lead soap remain (indicated by solid particles remaining 
after the third washing with water) break these up by running off almos 
all of the water layer and then add a little concentrated hydro chlori 
acid, shake and then continue the washing with water as before. 

Distill the ether from the solution of higher fatty acids and dry th 
latter in the flask by adding a little absolute alcohol and evaporatin 
on the steam bath. Dissolve the dry fatty acids by warming with 100 cc 
of 90 percent alcohol (volume) and cool slowly to 15°, shaking gentl 
to aid crystallization. Allow to stand at 15° for 30 minutes. If arae 
oil has been present arachidic acid will separate from the solution 
crystals. Filter, wash the precipitate twice with 90 percent alcoho 






and then with 10 cc portions of 70 percent alcohol, care being taken to 
keep the acid crystals and the washing alcohol at definite temperature 
in order to be able to apply the solubility corrections given below. 

Dissolve the arachidic acid upon the filter with boiling absolute 
alcohol, evaporate to dryness in a weighed dish and weigh. Add to the 
weight found 0.0025 gm for each 10 cc of 90 percent alcohol used in the 
crystallization and washing. 20 times the corrected weight of acid will 
be the approximate weight of arachis oil in the sample used. The 
crystals should be tested qualitatively by determining the melting 
ipoint, which should be 71° to 72°. The crystals should also be examined 
lunder the microscope. As little as 5 percent of peanut oil may be 
detected by this method. 

Sesame Oil: Baudouin Test. 1 — Dissolve 0.1 gm of finely powdered 
sugar in 10 cc of hydrochloric acid (sp. gr. 1.20), add 20 cc of the oil 
to be tested, shake thoroughly for a minute, and allow to stand. The 
aqueous solution separates almost at once. In the presence of even 
i very small admixture of sesame oil this is colored crimson. Some 
pive oils give a slight pink coloration with this reagent, but they are not 
pard to distinguish if comparative tests with sesame oil are made. 

The color was thought by Villa vecchia to be due to a reaction 
Ipf a constituent of sesame oil with furfurol, the latter being 
produced by the interaction of sugar with hydrochloric acid. 
Furfurol was accordingly substituted for sugar and hydrochloric 
Jicid and the method somewhat modified as follows: 

,< Sesame Oil: Villavecchia Test. 2 — Add 2 gm of furfurol to 100 cc 
If alcohol (95 percent) and mix thoroughly 0.1 cc of this solution, 10 cc 
If hydrochloric acid (sp. gr. 1.20) and 10 cc of oil by shaking them 
together in a test-tube. The same color is developed as when sugar is 
Ised, as in the Baudouin test. Villavecchia explained this reaction on 
jpe basis that furfurol is formed by the action of levulose and hydro- 
chloric acid and he therefore substituted furfurol for sucrose. As fur- 
<brol gives a violet tint with hydrochloric acid it is necessary to use the 
lery dilute solution specified in this method. 

I Fish and Marine Animal Oils in Mixtures with Vegetable 
ej)ils. — Practically all of these oils have very considerable 
J drying" properties, as shown by their iodine absorption 
lumbers. They are characterized by the presence of glycerides 

1 J. Assoc. Off. Agr. Chem., Vol. II, No. 3, Pt. II, 314. 
8 J. Soc. Chem. Ind., 12, 67 (1893); 13, 69 (1894). 



containing the unsaturated clupanodonic acid, whose formu 
and properties are mentioned on page 364. The peculii 
" fishy" odor of these oils is probably duetto the presence | 
this acid. 

Absorption of bromine by unsaturated acids or their glycerid 
produces bromides of limited solubility and high melting poin 
Octobromstearin, obtained from clupanodonin, melts at a high 
temperature (above 200°) and has a lower solubility than hexa 
bromstearin, obtained by brominating linolenin, and this als 
differs in a similar manner from tetrabromstearin, obtained fro 
linolin. Therefore the separation of octobromstearin fro: 
brominated fish and blubber oils provides a means for detecting 
marine animal oils in the presence of vegetable oils. The test is 
performed as follows: 

Dissolve in a test-tube about 6 gm of the oil in 12 cc of a mixture of 
equal parts of chloroform and glacial acetic acid. Add bromine, dro 
by drop, until a slight excess is indicated by the color, keeping ti 
solution at about 20°. Allow to stand for 15 minutes or more a 
then place the test-tube in boiling water. If only vegetable oils are 
present the solution will become perfectly clear, while fish oils will 
remain cloudy or contain a precipitate of insoluble bromides. 

Color Reactions. — A large number of qualitative tests, bas 
upon certain color reactions, have been proposed and considerabl 
used in the past for the detection of various oils. Color reactioi 
produced by adding concentrated nitric or sulphuric ack 
may be mentioned. Almost without exception these have bee 
found to be unreliable and they will not be described here 

The "elaidin test" is worthy of brief mention. This is bag 
upon the conversion of liquid olein into its solid isomer, elan 
by the action of nitrous acid, the change being from oleic aci 
of olein to elaidic acid of elaidin : 

CH3(CH2)?CH CH3(CH2)7CH 

II - II 

HC(CH 2 ) 7 COOH HOOC(CH 2 ) 7 CH 

Linolin, linolenin or clupanodonin are not thus affected and tl 
test serves to distinguish between liquid non-drying oils 


drying oils. It has been used to a considerable extent in testing 
the purity of olive oil but must be performed under strictly 
standardized conditions if it is to have even a qualitative value. 
Examination of an Oil or Fat Whose Identity is Unknown. — 
For the purpose of identifying an oil or fat of unknown character 
the complete chemical and physical examination is made unless 
its identity can be determined unmistakably by a qualitative 
test. With all of the constants determined a comparison is then 
made with all available data contained in published analyses of 
oils and fats and a reasonable agreement with such data will 
generally fix the identity of the unknown oil. In making com- 
parisons the most important figure is the iodine number because 
this will serve to classify the oil at once as a drying, semi-drying, 
or non-drying oil, the approximate ranges for these somewhat 
arbitrary divisions being as follows: 


Iodine Number 


200 and higher to 120 
120 to 95 



95 to 70 and lower 

The choice will, by this means, be narrowed down to a limited 
ist of oils or fats. The remaining constants are then considered, 
pe by one, and each comparison will narrow the choice still 
urther. At the last all available qualitative tests are made in 
prder to confirm the results of comparative tests or to aid in 
baking a final decision. It sometimes happens that the figures 
Obtained in the examination of the unknown oil will not all 
[orrespond, even to a reasonable degree, with the recorded data 
lor any of the common oils. This may be the result of (1) 
jrrors in the determinations, (2) adulteration, or (3) a real abnor- 
mality of the oil which is being examined. The first case should 
le at once excluded by repeating the determination of constants 
li which lack of agreement is observed. A careful inspection 
If all data may serve to indicate certain oils which, by addition 
lo one that most nearly resembles the oil under examination, 
lould change the " constants" in the manner observed. The 
flatter of commercial values should also be considered in this 



Oil Constants 

The following figures represent average values as determined on many samples of the oi 
indicated, and by many different chemists. Exact agreement should not be expecte( 
Specific gravity and index of refraction are calculated for the temperatures chosen, fro 
results reported for various temperatures. 


Sp. gr. 
at 20° 

Ind. refr. 
at 20° 






Hemp seed. . 



Poppy seed. 

Soy bean 

Sunflower. . . 



Anise seed . . 
Cotton seed. 


Grape seed. . 







Hazel nut. . . 




Cod liver. . . . 
Herring .... 
Menhaden. . 
Neat's foot. . 



Sheep's foot. 

Cacao butter. . 
Cocoa nut fat . 
Japan wax. . . . 
Myrtle wax . . . 


Palm nut 

Beef tallow 




Hare . . . : 



Mutton tallow. 


Carnaliba wax. 
Spermaceti. . . . 

Sperm oil 

Wool wax 




at 60° 

























at 60° 

465 (20°) 










































































connection, since any commercial material is adulterated, if at 
all, by a cheaper material. 

After the decision as to the identity of the oil has been made 
or has been limited to two or three possible oils, consult a good 
reference book for a complete discussion of these oils and make 
any additional tests that may be there suggested. For this 
purpose are to be recommended Lewkowitsch's Chemical Analysis 
of Oils, Fats and Waxes, 5th edition, volume 2, and Allen's 
Commercial Organic Analysis, volume 2, part 1. 

The figures in the table on page 388 are given for the purpose 
of comparison. They are gathered from various published 
analyses of the more common oils and fats. Exact agreement 
snould not be expected. For more extensive tables consult 
Lewkowitsch: Chemical Analysis of Oils, Fats and Waxes, and 
the technical journals. 

Hardened Oils. — Under any circumstances the analytical 
investigation of oils and fats offers difficulties that are often 
serious. The problems of the analyst are now increased many 
fold by the large development of the industry of hydrogenation 
of liquid oils. 

It has been seen that the most important difference between oils 
and fats lies in the larger proportion of olein in the former and of 
stearin and palmitin in the latter. Olein differs from stearin 
only in that it contains one unsaturated double bond in each 
oleic acid residue; the problem of saturating this group by the 
insertion of hydrogen, thus forming stearin, is one that has oc- 
cupied the attention of chemists for many years. At the present 
time the hydrogenation of the cheaper liquid oils (e.g., cottonseed, 
orn and peanut) to form edible fats is an industry that has at- 
tained large proportions. While this process changes liquid 
)ils to solid fats, it will also make a corresponding change in any 
analytical constants or tests that depend upon the degree of 
msaturation, as well as in the physical properties of the oil. 
L.inolin, linolenin and clupanodonin will be changed to stearin. 
Consequently the halogen absorption number, drying properties, 
pecific gravity, refractive index and temperature reactions will 
)e materially altered, as will also the odor and the general 
ppearance and consistency. It has been stated that fish oils 
>robably owe their characteristic odor to glycerides containing 


clupanodonic acid, while the somewhat similar odor of linseed 
oil is due to gycerides of linolenic acid. It is interesting to note 
that these odors are entirely lost through hydrogenation and that 
the oils are no longer recognizable by any ordinary tests. Many 
individual tests for other oils, such as the Halphen reaction for 
cottonseed oil and the Renard test for sesame oil, fail in the 
hydrogenated product. 

From one standpoint it might appear that the determination 
of what oils originally formed the raw materials for the " hard- 
ended" product is not a necessary one for the analyst to solve, 
since the properties of the finished product are, after all, the ones 
that have the chief practical interest for us. Yet it may some- 
times happen that the identity of the original oil, or the proof 
that a hydrogenating process has been employed may have a 
legal or other significance and the development of a series of 
suitable tests is very desirable. 

Analytical chemistry has made little progress in this direction. 
The application of delicate tests for metals (nickel, palladium, 
etc.) that are used as catalyzers in the hardening process, may 
sometimes serve to show that the material is a hardened oil 
rather than a natural fat. Other than this one can say ver 
little. But this knowledge of the nature of the changes cause 
by hydrogenation should serve to make the analyst more cautiou 
than he might otherwise be when interpreting the results of h 
tests of oils or fats of unknown origin. 


The chemical examination of water may be made to determine 
its fitness for drinking or for industrial uses, such as steam pro- 
duction, laundering, textile industries, etc. It is not necessary 
that a complete analysis should be made for all of these purposes 
because not all substances occurring in water are equally impor- 
tant in the different applications of the water. Natural waters 
often contain substances that are objectionable if they are to 
be used industrially and these substances are, for the most part, 
inorganic salts and, occasionally, acids. Most of such inorganic 
materials are without appreciable effect upon the human system 
and the examination for potability is rather directed toward the 
detection of pollution by sewage. On this account it becomes 
necessary to treat the subject of water analysis in two distinct 

Industrial Analysis.- — By far the largest industrial consumption 
of water is for the production of steam and for this reason the 
chemist is more often called upon for the analysis of water to 
determine its fitness for steaming than for any other industrial 
purpose. Pure water, however desirable it may be for use in 
the steam boiler, is not a natural product. Water from streams 
and other surface origins contains mineral and organic substances 
derived from the surface soil as well as inorganic compounds 
Iderived from springs which feed the stream. Water from wells 
[contains whatever mineral matter is common to the region 
through which it has flowed. Even rain water contains organic 
piatter and ammonia and may develop organic acids when stand- 
ing. Some of the compounds contained in water are com- 
paratively unobjectionable because their action is slight. It is 
fco be remembered, however, that in steam boilers the tempera- 
ture is higher than 100° because of the increased pressure. At 
a pressure of 100 pounds per square inch the boiling-point of 




water is 164° and at 200 pounds per square inch the boiling-point 
is 194°. At these temperatures the chemical activity of many 
dissolved substances is very much augmented. 

According to their effects upon boiler steel the constituents of 
natural waters may be classified as corrosives, incrustants anc 
foam producers. 

Corrosives. — Any soluble compound that can dissolve iron a 
high temperatures will give rise to pitting of the boiler, especially 
when the steel is not of uniform composition. Corrosives com 
monly occurring in water are chlorides, nitrates, and sulphates 
particularly of the alkaline earth metals, and free carbonic acid 
Free inorganic acids are of rare occurrence and absolutely unfit 
a water for steaming without preliminary treatment. A small 
amount of acid will cause corrosion for an indefinite perio 
because of the ready hydrolysis of iron salts. A cycle of re 
actions takes place as follows: 

Fe+2HCl^FeCl 2 +H 2 , 

6FeCl 2 +30^4FeCl 3 +Fe 2 3 , 

FeCl 3 +3H 2 O^Fe(OH) 3 +3HCl. 

A metal chloride which is easily hydrolyzed will also produce 
continuous corrosion: 

MgCl 2 +2H 2 0->Mg(OH) 2 +2HCl, 
Fe+2HCl->FeCl 2 +H 2 , etc. 

Nitrates are equally injurious, although they seldom occur in 
more than small concentration. Sulphates are somewhat les 
corrosive and free carbonic acid still less so. 

Incrustants. — Any substance that can be precipitated by heat 
ing or evaporation of water is, in a sense, an incrustant. Th 
steam boiler as a power producer is also a machine for continuou 
concentration of water solutions, since fresh, impure water i 
continually added and only vapor is removed. Strictly speakin 
only those substances which adhere to the boiler plate when thej 
are precipitated are classed as incrustants because only these ar 
particularly objectionable. These are carbonates of calcium an( 
magnesium and calcium sulphate. In presence of considerabl 
quantities of these materials certain other compounds, such a 

WATER 393 

silicic acid, iron oxide and aluminium oxide, may be included with 
the scale and then become incrustants. 

Calcium and magnesium carbonates are not dissolved as such 
in water but are present as bicarbonates, having been dissolved 
from the mineral carbonates by carbonic acid. 

CaC0 3 +H 2 C0 3 ->Ca(HC0 3 )2, 
MgC0 3 +H 2 C0 3 -+Mg(HC0 3 ) 2 . 

When the water is heated reactions which are the reverse of 
these take place and the normal carbonates are precipitated: 

Ca(HC0 3 ) 2 -+CaC0 3 +H 2 + C0 2 , 
Mg(HC0 3 ) 2 ->MgC0 3 +H 2 + C0 2 . 

These carbonates adhere to the boiler plate, the greatest amount 
of precipitation occurring over the heating surface. The scale 
jthus formed, although comparatively loose, hinders the trans- 
I mission of heat from the steel to the water and causes local super- 
heating. The result is a loss of efficiency and injury to the 
boiler. Although these substances occur in the water as 
bicarbonates, they are arbitrarily calculated as normal carbonates 
because the latter are precipitated when the water is heated. 

Calcium sulphate precipitates only when continued evapora- 
tion of the water concentrates it to the point of saturation. Pre- 
cipitation then causes the formation of a scale that is much more 
perious in its effects than the scale of carbonates, because it is 
[compact and adheres firmly to the boiler. While carbonate 
pcale can be largely removed by occasionally blowing off the water, 
(calcium sulphate can be loosened only by the use of hammer 
land chisel. On this account calcium sulphate is one of the most 
■objectionable incrustants of all compounds found in natural 

Foam Producers. — Carbonates of sodium and potassium in- 
crease the surface tension of water to such an extent that the 
result is foaming or "priming" as steam is taken from the boiler. 
Borne of the alkali waters of the West contain large quantities 
If these salts. 

Expression of Results. — The systems used in the calculation 
|)f results of water analysis were discussed in connection with the 
fletermination of hardness of water (page 231). It is convenient 



to work with 1000 cc of water or simple fractions of this quantity 
and to express results as milligrams of dissolved substance per liter 
of water. These figures may be changed to grains per gallon by 
multiplying by the factor 0.0583 -f. 

The analysis of the water solution will be made by means o: 
methods which give metals and acid radicals as the result o: 
separate determinations. It was at one time customary to calcu- 
late these as basic and acid anhydrides, as is still done in the 
analysis of minerals. There arises the same difficulty that is 
experienced in mineral analysis, viz.: that salts of hydracids 
cannot be expressed as oxides. A much better rule is to calculate 
all constituents as positive or negative radicals. 

Hypothetical Compounds. — There is still current among 
industrial chemists and engineers a custom of making a second 
calculation of compounds supposed to exist in the water. Most 
natural waters are highly dilute solutions of mineral matter. 
In such a solution most of the compounds are highly ionized an< 
all possible combinations of radicals as compounds are presenl 
to some extent, no matter what compounds were originalh 
dissolved by the water. It is evident, therefore, that any lisl 
of compounds calculated from the results of the analysis will b< 
entirely fanciful, so far as the actual condition of the solution is 
concerned. The basis of such a calculation was formerly the 
supposed affinity possessed by the different radicals for eacl 
other. If the radicals commonly occurring in water are arrangec 
in order of decreasing base and acid character the following series 
will be obtained: 

Positive Radicals 

Negative Radicals 



Based upon the assumption that the combination of these 
radicals will follow from their relative affinities the mathematica 
procedure would be to calculate the maximum amount of potas- 
sium chloride that could be formed, taking the excess of eithe 
potassium or chlorine as combined with the next radical of oppo- 

WATER 395 

site sign (either sodium or the nitrate radical) and so on, down 
the list. If the analysis has been accurately carried out and if 
all substances present have been determined the positive and 
negative radicals should be found in equivalent quantities, with 
a very slight excess of either magnesium or of the carbonate 
radical after the calculation is finished. This excess is the 
result of cumulative errors in the determination of the various 
radicals existing in the water and also of the occasional omission 
of small quantities of radicals other than those above named. 
Silicic acid, iron and aluminium are not included in the calcula- 
tion of hypothetical compounds because the colloidal nature of 
their hydroxides causes nearly complete, though indefinite, hy- 
drolysis of any salts that might originally have been present. 
They are therefore, according to custom, reported as oxides and 
this conventional method is sometimes responsible for the appear- 
ance of a slight excess of negative radicals in the final report. 
If ammonium salts are present in any considerable quantity, 
as in sewage effluents or factory wastes, a failure to deter- 
mine this radical also will result in an apparent excess of negative 

The customary method of calculating hypothetical com- 
pounds is not based upon scientific principles, as has already been 
shown. There is a certain justification for such a calculation, 
on account of the fact that when a water is heated and evapo- 
rated there will be produced the least soluble compounds of 
all that might be formed from the various radicals present. 
Through a certain coincidence, this would leave the radicals com- 
bined in about the same manner as is indicated by the conven- 
tional calculation. Heating in the boiler will produce the maxi- 
mum possible quantities of normal carbonates of calcium and 
magnesium. Which of these carbonates is least soluble at high 
temperatures is not definitely known because of difficulties en- 
countered in the determination of solubility. It is assumed, 
however, that if the radical of carbonic acid is not present in 
quantity sufficient to form carbonates with all of the calcium and 
magnesium, calcium, rather than magnesium, will ultimately 
remain to form the sulphate as the water is evaporated. Calcium 
sulphate is certainly next to the carbonates of calcium and mag- 
nesium with respect to its insolubility and it will precipitate 



when evaporation within the boiler concentrates it to the point of 
saturation. After these three compounds have been formed 
the method of combining the remaining radicals is quite imma- 
terial because they will not precipitate in any form, on account 
of the large solubility of salts of the alkali metals. 

In order to emphasize the real basis for any calculation of 
compounds we shall reverse the order of radicals given on page 
394 and calculate combinations in the following order: 

Positive Radicals 

Negative Radicals 



This will give precisely the same result as the calculations from 
the original order, unless there is found to be an excess of either 
positive or negative radicals. In this case the excess will be 
found to be either potassium or the chloride radical instead of 
magnesium or the carbonate radical, but the excess should be 
small enough to be insignificant in either case. 

On account of the fact that sodium and potassium have ver> 
little significance in most boiler waters (the negative radicals anc 
the metals that take part in scale formation being of chie: 
importance) and because they occur in relatively small amounts 
in most waters, the determination of potassium is frequently 
omitted and calculations are based upon the assumption thai 
sodium is the only alkali metal present. 

The conventional method of calculating hypothetical com' 
pounds is illustrated in the example given below. 

The analysis of a ground water gave the following results 

Milligrams per liter 












Oxides of iron and aluminium 





Chloride radical 

Nitrate radical 

Sulphate radical .' 

Carbonate radical 

WATER 397 

Following is the calculation of compounds : 
3^X24.19 = 59.65 = (C0 3 )"^Mg; 

24.19+59.65 = 83.84 = MgC0 3 . 

160.21 - 59.65 = 100.56 = (C0 8 )" remaining; 

20 03 

^-XIOO.56 = 67.20 = Cao(C0 3 ) ,/ remaining; 

100.56+67.20= 167.76 = CaC0 3 . 
75.41 -67.20 = 8.21 = Ca remaining; 

|^X8.21 = 19.69 = (S0 4 )''=c=Ca remaining; 

8.21 + 19.69 = 27.90 = CaS0 4 . 

32.91 -19.69 = 13.22 = (S0 4 )" remaining; 

S|?X6.02 = 12.56 = (S0 4 )"-Na; 

6.02+ 12.56 = 18.58 = Na 2 S0 4 . 
13.22-12.56 = 0.66 = (S0 4 )" remaining; 

||^X0.66 = 0.53 = Ko (SO <)" remaining; 

0.66+0.53 = 1. 19 = K 2 S0 4 . 

5.26 -0.53= 4.73 = K remaining; 

|f^X0.34 = 0.21=K=c=(NO 3 )'; 

0.34+0.21 = 0.55 = KN0 3 . 
4.73-0.21= 4.52 = K remaining; 

00^X4.52 = 4.10 = Cl^K remaining; 

4.52+4.10 = 8.62 = KCl. 

4.52 — 4. 10 = 0.42 = CI' remaining. This excess of chlorine 
represents experimental errors as is explained above. 

Since " milligrams per liter" multiplied by 0.0583 gives "grains 
per gallon" the complete statement of compounds is as follows: 



Formula for compounds 

Milligrams per liter 

Grains per gallon 

Si0 2 










Fe 2 3 +Al 2 3 


KN0 3 

K 2 S0 4 

Na 2 S0 4 

CaS0 4 

CaC0 3 

MgC0 3 

The use of the conversion factors as in the above illustration 
fits only the analysis that was used as the example. In the case of 
other waters certain factors might be the reciprocals of those used 
above or they might involve different pairs of equivalent weights 
All possible combinations should be calculated as a preliminary 
exercise and the factors with their logarithms recorded with the 
list of factors in the note book. It should be noteci that for 
water analysis the conversion factors are not multiplied by 100, 
since percents are not to be calculated. 


73. Calculate the conversion factors used above into the corresponding 
mixed numbers by carrying out the divisions indicated. Also calculate 
the following additional factors and their logarithms and record all of 
these in the note book. The first factor is given as an example. 











C0 8 

S0 4 
N0 3 



C0 3 




C0 3 
S0 4 
N0 3 

S0 4 



. Na 

co 3 

N0 3 


so 4 
N0 8 







N0 3 





WATER 399 

I Industrial Analysis of "Water. — Measure accurately in a calibrated 
[ilask enough water to give, upon evaporation, 0-5 gm to 1 gm of residue. 
Add 2 cc of concentrated hydrochloric acid and evaporate in a platinum 
pr porcelain dish on the steam bath. " Heat the residue at a temperature 
pelow redness until organic matter is removed. 

i Silicious Matter. — Add 1 cc of concentrated hydrochloric acid to 
ithe residue and then add about 20 cc of hot water. Warm and stir 
until all soluble matter has dissolved then filter on an extracted filter 
paper. Wash well with hot water until the combined filtrate and wash- 
ings amount to about 75 cc. Fold the filter paper and burn in a weighed 
ijplatinum crucible. The residue is reported as "silicious" if it is 
[white and does not weigh more than 5 mg, otherwise it may contain 
appreciable amounts of metal oxides or silicates. If the weight is 
greater than 5 mg, add to the residue a drop of sulphuric acid and then 
volatilize the silica by warming with 1 cc (more if necessary) of hydro- 
fluoric acid. Ignite the residue and weigh. Report the loss as silica. 
Dissolve any remaining residue in concentrated hydrochloric acid and 
| add to the main solution. 

Oxides of Iron and Aluminium. — Drop into the solution a very small 
bit of litmus paper and carefully add dilute ammonium hydroxide until 
| the solution is slightly basic. Boil gently to flocculate the hydroxides 
of iron and aluminium and to remove any unnecessary excess of ammo- 
nium hydroxide. Filter and wash with hot water until free from 
chlorides. Ignite and weigh and report as oxides of iron and aluminium. 

Usually iron is not present in quantity sufficiently large to make its 
separate determination important. If this is desired the method given 
on page 288 may be used and the equivalent amount of ferric oxide sub- 
tracted from the combined oxides, the remainder being aluminium 

Calcium. — Add 1 cc of dilute ammonium hydroxide to the filtrate and 
washings from the iron and aluminium hydroxides and then precipitate 
and determine the calcium with the precautions mentioned in the discus- 
sion on page 81. Report as calcium. 

At this point the procedure will be modified, according to whether 
potassium is to be determined. 

If potassium is to be determined, add to the solution 2 cc of concen- 
trated sulphuric acid and evaporate in a weighed platinum dish to 
dryness. Heat the residue carefully to evaporate excess of sulphuric 
acid and then more strongly to expel ammonium salts, finally heating 
for a short time until the dish is dull red. If sulphuric acid fumes do 
not appear upon heating, excess is not present and more should be added 
before the stronger heating. Weigh and record the weight of sulphates 
of sodium, potassium and magnesium. Dissolve the combined 



sulphates and dilute the solution to 250 cc in a calibrated volumetric 

Magnesium. — Fill a dry 100-cc volumetric flask with the solution of 
sulphates and rinse into a Pyrex* beaker. Determine magnesium 
directed on page 112. Notice that only 0.4 of the original sample w 
used for this determination. 

Potassium. — Use 100 cc of the sulphate solution that was obtain 
just before the determination of magnesium. Evaporate in a platinum 
dish over the steam bath adding the necessary quantity of chlorplatinic 
acid before crystallization of salts begins. Determine potassium by 
the Lindo-Gladding method, page 102. Here also 0.4 of the origin 
quantity of sample was used. 

Sodium. — Calculate the weight of potassium sulphate equivalent to 
the potassium chlorplatinate found and also the weight of magnesium 
sulphate equivalent to magnesium pyrophosphate found. Multiply 
the sum of these weights by 2.5 and subtract the product so obtained 
from the total weight of combined sulphates already found. The 
remainder is sodium sulphate. From this calculate sodium. 

If potassium is not to be determined the weighed residue of sulphati 
is dissolved and the solution is diluted to about 100 cc in a 250 cc beake 
of Pyrex or similar resistance glass. Magnesium is then determined i 
the entire solution. From the weight of magnesium pyrophosphat 
found calculate the weight of magnesium sulphate equivalent to it an 
subtract this from the weight of combined sulphates, already founc 
The remainder is assumed as the weight of sodium sulphate, from whic 
sodium is calculated. 

Sulphates. — Use 100 cc of water unless a qualitative test shows th 
presence of only a small concentration of sulphates, in which case 500 c 
or more should be evaporated to about 100^ cc. Add 0.5 cc of concen 
trated hydrochloric acid and precipitate by barium chloride, carryin 
out the precipitation and treatment of the precipitate as directed oi 
page 95. Calculate milligrams per liter of the sulphate radical. 

Chlorides. — Make 500 cc of a standard solution of pure sodium 
chloride, 1 Cc of which contains 0.001 gm of chlorine. Make 1500 cc 
of a solution of silver nitrate such that 1 cc is calculated to be equivale 
to about 0.000505 gm of chlorine and standardize against the sodiu 
chloride solution as follows: Measure 25 cc of the standard sodium chl 
ride solution into a 4-inch porcelain casserole or into a beaker placed on 
white background. Add 1 cc of a 5 percent solution of potassiu 
chromate (from which chlorides have been precipitated by the additio 
of a slight excess of silver nitrate) and then titrate with silver nitra 
solution until the first permanent red tint of silver chromate appear; 
This is the end point of the reaction between silver nitrate and sodiu 

WATER 401 

chloride. Calculate the dilution necessary to make 1 cc of the silver 
nitrate solution equivalent to 0.0005 gm of chlorine and dilute 1000 cc 
to the required volume by adding distilled water from a burette. 

Determine chlorine in the water by titrating 100 cc or more of water 
by standard silver nitrate, using 1 cc of potassium chromate as the 
indicator. In case chlorides are present in very small concentration it 
may be necessary to use more than 100 cc and to evaporate to this 
volume. The end point of the titration is much more easily observed if 
a rather heavy precipitate of silver chloride is present at the end. If, 
therefore, the concentration of chlorine in the water is small it is ad- 
vantageous to add first a measured volume (10 to 25 cc) of standard 
sodium chloride solution, deducting this volume from the total volume 
of silver nitrate required. Calculate milligrams per liter of the chloride 

Nitrates. — Use one of the methods described on pages 427 and 428 for 
the sanitary examination of water. Calculate milligrams per liter of the 
nitrate radical. The concentration of nitrates in most waters is too 
small to be of consequence in steaming but this is not always the case 
and the determination should not be omitted. 

Carbonates. — Titrate 100 cc of the water by fiftieth-normal hydro- 
chloric acid, using methyl orange as indicator. More or less than 100 cc 
of water may be necessary if the concentration of carbonates is very 
small or very large. Enough should be used to require 25 to 45 cc of 
standard acid. Calculate milligrams per liter of the carbonate radical 
(CO) 3 ". This is based upon the customary conventional assumption 
of normal carbonates instead of bicarbonates. 

The determination of carbonates by titration with standard acid is 
the determination of alkalinity as described on page 231 except that 
hardness is arbitrarily calculated as calcium carbonate while in this 
connection the carbonate radical is calculated^ 

Calculate milligrams per liter of the hypothetical compounds using 
the method illustrated on page 394. Calculate the corresponding grains 
per U. S. gallon. 

Treatment.- — If the chief incrustants of a water are bicarbonates 
of calcium and magnesium these may be largely precipitated by 
heating the feed water by means of the exhaust steam. This 
kind of treatment is limited in its application on account of the 
short time allowed for settling. Other incrustants than those 
mentioned are not removed by this process. 

Calcium hydroxide will precipitate bicarbonates and this is 



the cheapest agent available for this purpose. It also possesses 
the advantage of leaving no by-products in the water: 

Ca(HC0 3 ) 2 +Ca(OH) 2 ->2CaC0 3 +2H 2 0. 
Mg(HC0 3 )2+Ca(OH) 2 -^eaC0 3 +MgC03+2H 2 0. 

Sodium carbonate will also precipitate bicarbonates as well as 
other salts of calcium and magnesium. It, however, leaves in 
the water the corresponding salt of sodium and this is objection- 
able. Examples of the reactions are expressed by the following 

Ca(HC0 3 ) 2 +Na 2 C0 3 ^CaC0 3 +2NaHC0 3 , 

CaS0 4 +Na 2 C0 3 ->CaC0 3 +Na 2 S0 4} 

CaCl 2 +Na 2 C0 3 -^CaC0 3 +2NaCl. 

The best procedure is to treat the water first with the amount o 
calcium hydroxide necessary to react with bicarbonates, allowing 
a small excess, and then with sufficient sodium carbonate to 
precipitate alf remaining calcium and magnesium. The reaction 
should be carried out in tanks large enough to provide wate 
for the plant, allowing the necessary time for settling. Th 
initial cost of the purifying plant is almost the only cost becaus 
of the relatively low cost of the small amount of lime and sods 
ash necessary. 

In calculating the necessary treatment the purity of the lim 
or lime water must be known and also that of the soda ash. Th 
results may be expressed according to custom as pounds of 
reagent per 100,000 gallons of water. The quantity of (C0 3 )" 
expressed as grains per gallon, is multiplied by the fraction 

equivalent weight of CaO tl .. . . . , .. 

— —. — ; — 7 . T , e ,rir\ \t» tne result being grains of lime re- 

equivalent weight of (C0 3 ) 

quired for one gallon of water. One pound avoirdupois contain 
*/w> • mi 1. grains CaO per gallon X 100,000 j 

7000 grains. Therefore - 7000 = pound 

ofjcalcium oxide required for 100,000 gallons of water. Instea 
of J this the calculation may be made to gallons of lime water pe 
100,000 gallons of water, where a saturated solution of calciur 
hydroxide is first made and its concentration determined. 

WATER 403 

The amount of sodium carbonate necessary for the precipita- 
I tion of other salts of calcium and magnesium may be calculated 
j in a similar manner. 

An inspection of the statement of the analysis as given on page 
i 395 will show that calcium and magnesium as carbonates may be 
j precipitated by lime water, while calcium as sulphate may be 
| removed by sodium carbonate. Since no other bicarbonates are 
j present the amount of calcium oxide required may be^calculated 
; directly from the amount of carbonate radical. 

56X160.21X0.0583X100,000 1Q . a 
60X7000 - = 124.6. 

| Therefore 124.6 pounds of calcium oxide will be required to 
I precipitate calcium and magnesium carbonates. 



= 18.1. 

Therefore 18.1 pounds of sodium carbonate will be required to 
remove calcium sulphate from 100,000 gallons of water. 

The remaining salts are classed as corrosives and they cannot 
be removed by chemical treatment. 


74. Calculate the treatment for the water already analyzed in the 

Boiler Compounds. — Many commercial mixtures are now on 
! the market, designed to be mixed with feed water as it enters the 
boiler to prevent the formation of scale or to loosen scale already 
I formed. Most of these mixtures are solutions of sodium carbon- 
late or sodium hydroxide with or without the addition of tannin or 
I some other colloidal organic compound. The base precipitates 
| bicarbonates but this precipitation occurs within the boiler 
! where it would occur if no base had been added. The base is a 
jstrong corrosive and may be highly injurious to the boiler if 
I added in excess. Colloidal bodies, such as tannin, starch or 
dextrine, have a certain loosening effect upon the scale already 
present and to some extent prevent the formation of compact 



scale. This action can produce only temporary relief, an 
preliminary treatment such as has already been described 
much cheaper and better. 

Examination of Water for Sanitary Purposes. — The examina 
tion of water to be used for drinking may be made to determin 
the quantity of mineral salts that have a supposed medicina 
value or to determine its potability. The examination for th 
first purpose will follow the lines of methods already describee 
for boiler water or it may be extended to include other substance 
such as lithium, free carbonic acid, hydrogen sulphide, etc 
It is practically certain that none of the mineral waters that ar 
exploited in a commercial way contains enough of any salt or gas 
to have any appreciable effect upon the human system and that 
the beneficial effects that are noticed by those who take treatment 
at the mineral springs are due largely or entirely to other causes, 
such as enforced dieting, bathing, relaxation from care, good 
exercise and the drinking of plenty of water. This being the 
case it is apparent that the chemist's report on the analysis of 
mineral waters can have little value except as a commercial 
document. If this analysis is demanded the methods already 
given may be used, proper modifications being made in quantity 
of water and reagents according to large variations in the perce 
of mineral matter present. Additional determinations will 

Free Carbonic Acid. — Water containing free carbonic acid wi 
rapidly lose carbon dioxide and it therefore becomes necessar 
to make this determination as soon as the water sample is taken 
or to preserve the sample in tightly closed bottles entirely filled 
with water. The determination depends upon the fact that as 
sodium carbonate is added to carbonic acid in presence of phenol- 
phthalein a color change occurs at the moment that all of the car- 
bonic acid has been used in forming sodium bicarbonate, sodium 
carbonate having a basic reaction toward phenolphthalein 
The end point is shown when the following reaction is completec 

Na 2 C03+H 2 C03->2NaHC0 3 . 

Determination. — Calculate the normality of a solution of sodium ca 
bonate so made that 1 cc is equivalent to 0.001 gm of carbon dioxid 
Make 1000 cc of such a solution, using sodium carbonate prepared b 

WATER 405 

Iheating recrystallized sodium bicarbonate to about 300° until it ceases 
to lose weight. Do not allow the solution to come into contact with air 
more than is necessary. 

Titrate 100 cc of water as rapidly as possible with the standard sodium 
carbonate solution and report milligrams per liter of carbon dioxide. 
The report is often made as volume of gas per unit volume of water at 
some specified temperature as, for example, cubic inches of gas per gallon 
of water, at 60° F. Reference to tables of density of carbon dioxide 
will give the necessary data for this calculation. 

Hydrogen Sulphide. — The determination of hydrogen sulphide 
must be made as soon as the water is taken and for the same 
reasons that apply to carbon dioxide. Titration is made with 
standard iodine solution according to the reaction: 

2I+H 2 S^2HI+S. 

Determination. — Make a solution of iodine, 1 cc of which is equivalent 
to 0.001 gm of hydrogen sulphide. Standardize by titrating with stand- 
ard sodium thiosulphate. Titrate 100 cc of the water with the standard- 
ized iodine solution, using starch solution as indicator. Calculate 
milligrams per liter of hydrogen sulphide, also cubic inches per gallon. 

Iron. — The quantity of iron in water is usually too small to 
make ordinary volumetric methods desirable for its determina- 
tion and a more sensitive colorimetric method is substituted. 
The iron is obtained in the ferric state and is treated with potas- 
sium thiocyanate. The red color so produced is compared in 
tubes with that formed by a standard iron solution. A distinc- 
tion may be made between ferrous and ferric iron by using potas- 
sium ferricyanide instead of potassium thiocyanate. In this case 
a blue color is produced by ferrous iron and no visible color re- 
sults from the reaction of the small amount of ferric iron usually 

Determination of Total Iron. — Prepare the following reagents: 
1. Standard Iron Solution. — Calculate the weight of ferrous am- 
monium sulphate required for 1000 cc of solution, 1 cc of which shall 
contain 0.0001 gm of iron. Dissolve this quantity of the salt in about 
50 cc of distilled water and add 20 cc of dilute sulphuric acid. Warm 


slightly and add potassium permanganate solution until the iron is 
completely oxidized, using the smallest possible excess. Dilute to 
1000 cc. 

2. Potassium Thiocyanate. — Dissolve 20 gm of the salt in 1000 cc of 
distilled water. 

3. Dilute Hydrochloric Acid. — Dilute the concentrated acid with a: 
equal volume of distilled water. The acid must be free from nitri 

4. Potassium Permanganate. — Make 500 cc of a solution appro 
mately fifth-normal. 

Evaporate 100 cc of the sample to dryness, or use the residue le 
after the determination of solids. With silt-bearing waters the quantit 
of iron is sometimes so great that it is necessary to use as little as 10 
cc of the sample. With such waters evaporation should be made in 
the presence of 5 to 10 cc of concentrated hydrochloric acid to effect 
complete solution of the iron. If the sample of water contains much 
organic matter, destroy this by ignition, taking care not to prolong 
the ignition so as to render the iron too difficultly soluble. 

Cool the dish and add 5 cc of dilute hydrochloric acid to moisten the 
whole of the inner surface of the dish. Place the dish on the steam bath 
for two or three minutes and again moisten the whole inner surface by 
allowing the hot acid to flow over it. Add 5 to 10 cc of distilled water 
to rinse down the sides of the dish, and again place on the steam bat' 
for about three minutes. 

The hot acid solution is washed from the dish with distilled wate 
into a 100-cc Nessler tube (see page 417). Filter the sample if necessary 
carefully washing the filter paper with hot water. Add a drop or two o: 
potassium permanganate solution to oxidize the iron to the ferric condi 
tion. The color of the permanganate should persist for at least 
minutes; if not, add more permanganate solution, a drop at a time. 

To the cooled solution 10 cc of potassium thiocyanate solution is 
added, and the volume made up to 100 cc and well mixed. 

Immediately compare the resulting color with that in a series of 
standards prepared side by side with the sample in 100-cc Nessler tubes 
in which amounts of standard iron solution ranging from 0.05 cc to 
4 cc are first diluted with water to about 50 cc. ' 5 cc of dilute hydro 
chloric acid and a drop or two of potassium permanganate are adde 
to each tube of standard solution and all are diluted to 100 cc. The 
number of standards needed is governed by the quantity of iron likely 
to be present in the sample examined. 

Potassium thiocyanate is added to each of the standard solutions at 
the same time that this reagent is added to the samples of water under 



WATER 407 

jjxamination. Comparison of the sample with the standards, which are 
I made up to 100 cc after adding the thiocyanate and mixing, should 
pe made immediately. 

j Potability. — The examination of water to determine its suit- 
ability for drinking (potability) is practically always directed 
foward the question as to whether pollution by sewage has 
occurred. This examination is quite different in principle from 
liny of the processes already studied, in that the substances 
actually determined are nearly or quite harmless and are sig- 
nificant only as they point to the probable presence of patho- 
genic bacteria. The chemical examination goes no farther than 
the determination of certain chemical compounds which always 
accompany sewage and which therefore indicate a danger in the 
[use of the water for drinking because disease-producing micro- 
organisms also generally accompany sewage. 

Since this is the case it might be supposed that the examina- 
tion might be more properly made by the bacteriologist, who 
determines directly the presence or absence of bacteria in ab- 
normal numbers and who also makes a direct test for B. coli 
communis, an organism that practically always accompanies 
faecal discharges and is therefore found in all water polluted by 
waste matter from human organisms. If the results of the 
bacteriological examination were unfailing this examination 
would probably suffice for all cases. It should be noted, however, 
that the bacteriologist is also striving for indirect rather than 
direct results. It is not practicable to make a direct examina- 
tion of the species of every organism found in order to test for 
the presence of actual pathogenic forms and reliance is generally 
placed upon the two factors noted above, i.e., the concentration 
of bacteria (number per cubic centimeter) and the presence or 
absence of B. coli communis. If, for some reason, conditions 
were temporarily unfavorable to the growth of bacteria at the 
time the sample was taken, the number of organisms might be 
so reduced as to cause no suspicion of the real condition of the 
water. It is conceivable that the chance entrance of antiseptic 
substances into sewers or streams or the action of sunlight and 
air should bring about this result and B. coli communis might 
also be entirely absent. This might, for instance, happen as a 
result of mixing with factory wastes such as those from the manu- 


facture of paper and textiles and from plating, bleaching and 
other chemical industries. In such a case the water would be 
passed by the bacteriologist when it should be condemned. Od 
the other hand such influences as those mentioned would nc 
eliminate the chemical products of putrescent sewage and tl 
chemical examination would be likely to show pollution. Froi 
this examination the water would be condemned. The conch 
sion is that neither method of examination is infallible and botl 
should be used wherever much importance attaches to the results 
If one method must be omitted it is preferable that this should 
the bacteriological method, provided that sufficient data are 
hand for the proper interpretation of the results of the chemical 

Interpretation. — In order properly to interpret the results of 
the analysis it is necessary to know the normal condition of the 
particular class of water under examination because all of the 
substances occur normally in practically all waters and their 
proportions vary with the source of the water. For example 
chlorides occur in all ground and surface waters and the amoui 
is governed largely by soil conditions. Near the sea shor 
chlorides occur in ground waters in large quantities. Simila 
conditions obtain for nitrogen in all of its forms and for organi 
matter and total solids. The necessary data for the interpreta 
tion of a single analysis cannot well be collected by individua 
without great expenditure of time and labor. They are gener 
ally obtained as a result of organized efforts of state and citj 
boards of health and of scientific societies. 

The details of manipulation in the analytical work late 
described, as well as the directions for taking samples and makin 
the physical examination, are essentially those recommended ty 
the Joint Committee on Standard Methods of Water Analysis o 
the American Public Health Association, American Chemics 
Society and Association of Official Agricultural Chemists. Th 
report of this Committee forms a most valuable contribution 
the scientific phases of the subject, not only in the matter 
unification of practice but also in the guidance that it affords 
a selection of the best methods now available. The third edi 
tion of the report is printed as a special volume to be obtainec 
from The American Public Health Association. 

WATER 409 

For the directions for determinations other than those here 
described reference should be made to the complete report. 

Collection of Samples: Quantity. — The minimum quantity neces- 
sary for making the ordinary physical, chemical and microscopical 
analyses of water or sewage is two liters; for the bacteriological ex- 
amination, one hundred cubic centimeters. In special cases larger 
quantities may be required. 

Bottles. — The bottles for the collection of samples must be of hard, 
clear, white glass, and they must have glass stoppers. Earthen jugs 
Jor metal containers must not be used. 

! Sample bottles must be carefully cleansed each time before using. 
rThis may be done by treating with sulphuric acid and potassium 
dichromate, or with a basic solution of potassium permanganate, 
and afterward with a mixture of oxalic and sulphuric acidsj then 
[thoroughly rinsing with water and draining. 

When clean, the stoppers and necks of the bottles are protected 
from dirt by tying cloth or thick paper over them. 

For shipment they are packed in cases with a separate compartment 
for each bottle. Wooden boxes may be lined with indented fiber paper, 
felt or some similar substance, or provided with spring corner strips, 
to prevent breakage. Lined wicker baskets also may be used. 

Bottles for bacterial samples, besides being washed, must be sterilized 

with dry heat for one hour at 160° or in an autoclave at 115° for fifteen 

minutes. For transportation' they may be wrapped in sterilized cloth 

or paper, or the necks may be covered with tin foil and the bottles put 

in tin boxes. When bacterial samples must of necessity stand for twelve 

hours before plating, bottles holding more than four ounces must be used. 

The bottles used for chemical samples may be sterilized and the 

! samples so collected used for the bacteriological analysis. When 

i bacterial samples are not plated at the time of collection they are kept 

! on ice at a temperature not higher than 15° and preferably as low as 10°. 

Time Interval between Collection and Analysis. — Generally speaking, 

> the shorter the time elapsing between the collection and the analysis of 

j a sample, the more reliable will be the analytical results. Under many 

j conditions, analyses made in the field are to be commended, as data so 

obtained are frequently preferable to those made in a distant laboratory 

! after the composition of the water has changed en route'. 

The allowable time that may elapse between the collection of a sample 
; and the beginning of its analysis cannot be stated definitely, as it 
! depends upon the character of the sample and upon other conditions, 
j but the following may be considered as fairly reasonable maximum 
! limits under ordinary conditions : 


Physical and Chemical Analysis. 

Ground waters 72 hours 

Fairly pure surface waters 48 hours 

Polluted surface waters 12 hours 

Sewage effluents 6 hours 

Raw sewages 6 hours 

Microscopical Examination. 

Ground waters 72 hours 

Fairly pure surface waters 24 hours 

Waters containing fragile organisms . . Immediate examination 

Bacteriological Examination. 

Samples kept at less than 10° 24 hours 

If sterilized by the addition of chloroform, formaldehyde, mercuric 
chloride, or some other disinfectant, samples for chemical and micro- 
scopical examination may be allowed to stand for longer periods than 
those indicated, but as this is a matter which must vary according to 
local circumstances, no definite procedure is recommended. 

If unsterilized samples of sewage, sewage effluents, and highly polluted 
surface waters are not analyzed on the day of their collection, caution 
must be used in regard to the organic contents, which frequently change 
materially upon standing. 

The gaseous contents of samples, especially dissolved oxygen and 
carbonic acid, should be obtained immediately, in accordance with the 
directions given in connection with each determination. 

Representative Samples. — Care should be taken to secure a sample 
which is truly representative of the liquid to be analyzed. In the case 
of sewages this is especially important, in view of the marked variations 
in composition which occur from hour to hour. Frequently satisfactory 
samples can be obtained only by mixing together several portions col- 
lected at different times or in different places — the details as to collecting 
and mixing depending upon local conditions. 

Physical Examination: Temperature. — The temperature of the 
sample should be taken at the time of collection, and should be expressed 
in Centigrade degrees, to the nearest degree or closer if for any reason 
more exact data are required. For obtaining the temperature of water 
at various depths below the surface the thermophone is recommended. 

Turbidity. — The turbidity of water is due to suspended matter,^ 
such as clay, silt, finely divided organic matter, microscopic 
organisms, etc. The increasing use of filters for the purification 
of water and sewage has made this determination one of great 

WATER 411 

J The standard of turbidity is that adopted by the United States 
peological Survey, namely, a water which contains 100 parts of silica 
ier million in such a state of fineness that a bright platinum wire one 
millimeter in diameter can just be seen when the center of the wire is 
100 millimeters below the surface of the water and the eye of the ob- 
server is 1.2 meters above the wire, the observation being made in the 
Iniddle of the day, in the open air, but not in sunlight, and in a vessel 
|o large that the sides do not shut out the light so as to influence the 
(Jesuits. The turbidity of such water is taken as 100. 
i Preparation of Silica Standard. — Dry Pear's "precipitated fuller's 
ftarth" and sift it through a 200-mesh sieve. 

1 gm of this preparation in one liter of distilled water makes a stock 
■suspension which contains 1000 parts per million of silica and which 
Ihould have a turbidity of 1000. Test this suspension, after diluting a 
fcortion of it with nine times its volume of distilled water, with a wire to 
Ascertain if the silica has the necessary degree of fineness and if the sus- 
pension has the necessary degree of turbidity. If not, correct by adding 
■more silica or more water as the case demands. 

Standards for comparison are prepared from this stock suspension 
py dilution with distilled water. For turbidity readings below 20, 
fetandards of 0, 5, 10, 15 and 20 are kept in gallon bottles made of 
■clear white glass; for readings above 20, standards of 20, 30, 40, 50, 60, 
|70, 80, 90 and 100 are kept in 100-cc Nessler tubes approximately 20 
|mm in diameter. 

Comparison of the water under examination with the standards is 
[made by viewing them sidewise toward the light, looking at some 
lobject and noting the distinctness with which the margins of the object 
lean be seen. 

The standards must be kept stoppered and both sample and standards 
[thoroughly shaken before making the comparison. 

In order to prevent any bacterial or algal growths from appearing in 
the standards, a small amount of mercuric chloride may be added to 

Platinum Wire Method. — This method requires a rod with a platinum 
wire having a diameter of 1 mm inserted in it about 25 mm from the end 
of the rod, and projecting from it at least 25 mm at a right angle. Near 
'the end of the rod, at a distance of 1.2 meters from the platinum wire, 
a wire rin^ is placed directly above the wire through which, with his 
eye directly above the ring, the observer shall look when making the 
examination. The rod is graduated as follows: 

The graduation mark of 100 is placed on the rod at a distance of 
100 mm from the center of the wire. Other graduations are made ac- 
cording to the table on page 412. These graduations are the ones used 



to construct what is known as the U. S. Geological Survey Turbiditj 
Rod of 1902. 

Procedure. — Push the rod down into the water vertically as far sli 
the wire can be seen and then read the level of the surface of the water or 
the graduated scale. This will indicate the turbidity. 
The following precautions should be taken to insure correct results 
Observations should be made in the open air, preferably in th( 
middle of the day and not in direct sunlight. The wire must be kepi 
bright and clean. Waters which have a turbidity above 500 are dilutee 
with clear water before the observations are made, but in case this h 
done the degree of dilution used should be stated and should form a part 
of the report. 

Turbidity, parts 

Vanishing depth 

Turbidity, parts 

Vanishing depth 

per million 

of wire, mm 

per million 

of wire, mm 

































































" 300 

















27.7 • 





















WATER 413 

The wire method is used for testing the degree of fineness of the 

(standard silica, and this degree of fineness shall be such that when added 

i jfco distilled water in an amount equal to 100 parts per million, the wire 

jobserved under standard conditions can be just seen at a depth of 100 

mm below the surface of the water. 

Expression of Results. — The resists of turbidity observations are 

' expressed in whole numbers which correspond to parts per million of 

silica, and recorded as follows: 

Turbidity between 1 and 50, recorded to nearest unit. 

Turbidity between 51 and 100, recorded to nearest 5 

Turbidity between 101 and 500, recorded to nearest 10 

Turbidity between 501 and 1000, recorded to nearest 50 

Turbidity between 1001 and above, recorded to nearest 100 

Coefficient of Fineness. — The number obtained by dividing 
ithe weight of suspended matter in the sample (in parts per 
! million) by the turbidity is called the coefficient of fineness. If 
: jgreater than unity it indicates that the matter in suspension in 
(the water is coarser than the standard; if less than unity, that 
lit is finer than the standard. 

Color. — The color of water may form an important indication 
of pollution. There is little of value to be obtained from a 
(quantitative measurement of color although the determination 
is discussed at length in the report of the Committee on Standard 

Odor. — The observation of the odor of cold and hot samples of 
(surface waters is very important, as the odors are usually con- 
nected with some organic growths or with sewage contamination 
ior both. 

The odor of ground waters is often caused by the earthy con- 
'stituents of the water bearing strata. The odor of a contami- 
nated well water is often decisive evidence of its pollution. 

A study of the organisms of water is an invaluable adjunct to 
ithe physical and chemical examination of water. Certain organ- 
isms can be distinguished by their odor, as, for example, the 
("fishy" odor of Uroglena the " aromatic" or "rose geranium" 
|odor of Asterionella and the "pig-pen" odor of Anabcena. 

Determination. — Observe and record the odor, both at room tempera- 
ture and at just below the boiling-point, as follows: 

Cold Odor. — Shake the sample violently in one of the collecting bottles, 
; when it is about half or two-thirds full and when the sample is at room 



temperature (about 20°). Remove the stopper and test the odor at 
the mouth of the bottle. 

Hot Odor. — Into a 500 cc Erlenmeyer flask pour about 150 cc of the 
sample. Cover the beaker with a well-fitting watch glass, place on 
a hot plate and bring the water to just below boiling. Remove the 
beaker from the plate and allow it to cool for not more than five minutes. 
Then shake with a rotary movement, slip the watch glass to one side 
and test the odor. 

Expression of Results. — Express the quality of the odor by some 
such descriptive epithet as the following, which for purposes of record 
may be abbreviated: 

v vegetable 
a aromatic 
g grassy 
f fishy 
e earthy 
c free chlorine 

m moldy 
M musty 

d disagreeable 

P peaty 

s sweetish 

S hydrogen sulphide 

Express the intensity of the odor by a numeral prefixed to the term 
expressing quality, which may be defined as follows: 




Very faint. . 



Decided.. .. 
Very strong 

Approximate definition 

No odor perceptible. 

An odor that would not be ordinarily detected by the 

average consumer, but that could be detected in the 

laboratory by an experienced observer. 
An odor that the consumer might detect if his attention 

were called to it, but that would not otherwise attract 

An odor that would be readily detected and that might 

cause the water to be regarded with disfavor. 
An odor that would force itself upon the attention and 

that might make the water unpalatable. 
An odor of such intensity that the water would be 

absolutely unfit to drink. (A term to be used only in 

extreme cases.) 

Chemical Examination.- — The following determinations may be 
made: Total solids, chlorine of chlorides, albumenoid nitrogen, 
total organic nitrogen, nitrogen of ammonia or ammoniui 
salts, nitrogen of nitrites, nitrogen of nitrates, total organic 
matter, dissolved oxygen and poisonous metals. Besides th( 
chemical analysis certain purely physical tests may be made sucl 
as temperature, color, odor and turbidity. These determinatioi 
have just been described. 

WATER 415 

Total Solids. — This is taken as the residue obtained when a 
measured volume of water is evaporated. The general character 
of the solids may be sometimes noted, also the amount of loss 
suffered by igniting in air and the odor and amount of charring 
afford an indication as to the quantity and character of solid 
organic matter. 

Determination. — Ignite and weigh a platinum dish, then evaporate 
in it 100 cc of water, using the steam bath. The dish need not be large 
enough to hold the entire 100 cc at one time. A small dish is better. 
Heat the residue at about 103° for one-half hour. Cool in the desiccator 
and weigh. Report the increase in weight as milligrams per liter of 
total solids. Heat the dish at low redness until all organic matter is 
burned. The change in weight is loss on ignition. 

Suspended Matter. — Filter a portion of the sample through a very 
fine, close filter paper or a well-formed asbestos mat in a Gooch crucible, 
rejecting the first 15 to 25 cc. Determine the total dissolved solids in 
the filtrate as already directed. This, subtracted^from the solids of 
the original sample, gives the total suspended matter. 

Chlorine of Chlorides. — Chlorine occurs to some extent in all 
natural waters. It is found to a much larger extent in sewage 
where it enters chiefly as sodium chloride of urine and faeces, 
and of household wastes. Sewage polluted streams or wells, 
therefore, always carry abnormally large quantities of chlorine. 
Also the chlorine content may be increased by ocean vapors 
carried inland by natural deposits or by factory wastes. If the 
normal chlorine content exceeds 20 mg per liter, the determina- 
tion will have little significance from the sanitary standpoint. 

Determination. — Use the method described on page 397. Report 
milligrams per liter of chlorine. 

Nitrogen in Various Forms. — Human faeces contains large 
quantities of nitrogen while urine has a normal content of about 
0.85 percent of nitrogen. The entrance of sewage therefore im- 
parts abnormal concentrations of nitrogen to water. This 
nitrogen is at first practically all in the form of organic com- 
pounds and of urea. Part of the^organic nitrogen is* readily 
converted into ammonia by oxidizing with potassium per- 



manganate in basic solution. This part is known as "albume- 
noid nitrogen" because it is contained in albumenous bodies. 

The action of certain forms of bacteria (chiefly anaerobic) 
causes the putrefaction of organic matter and this cleavage of 
complex compounds results in the formation of ammonia from 
the nitrogen. Part or all of this ammonia may combine with 
acids to form ammonium salts. All such nitrogen is known as 
" nitrogen of free ammonia," whether this be of really free 
ammonia or of ammonium salts. 

Where sewage or water polluted by it is exposed to air and 
sunlight the simpler organic compounds produced by putrefaction 
are subjected to oxidation, this being promoted by other forms of 
bacteria (aerobic). Ultimately the organic compounds are 
completely oxidized. The two processes, putrefaction and 
oxidation, are made the basis of the septic process of water puri- 
fication. In the operations of water analysis the changes in 
the forms of nitrogen are most important. "Free" ammonia is 
oxidized to nitrous acid which usually remains combined as 
nitrites of metals or of ammonium. Further oxidation produces 
nitric acid or nitrates, the final stage in the series. 

The analytical estimation of the nitrogen in different forms in 
water offers a valuable indication, not only as to the probability 
of pollution but also concerning the present condition. The 
presence of abnormal quantities of "albumenoid nitrogen" 
indicates the presence of unchanged sewage and the probable 
presence of dangerous micro-organisms in their most virulent 
condition. "Free ammonia" in considerable quantities shows 
that the raw sewage has become fermented and that it must have 
been largely diluted in the time that has elapsed since the 
entrance of sewage. Nitrites are very readily oxidized and will 
not be found in more than traces unless free ammonia is also 
present. Abnormal quantities of nitrates, unless these are of 
inorganic origin, are the result of complete oxidation of organic 
matter and this must have required time and continued action of 
air and sunlight. If all forms of nitrogen are found in abnormal 
quantities continuous pollution is occurring. All of these figures 
are highly significant in view of the fact that the same influences, 
that promote the decomposition and oxidation of nitrogenous 
organic matter also combine for the partial or complete steriliza- 

WATER 417 

tion of the water. It is not, by any means, to be concluded that 
water which has been polluted by sewage but in which the latter 
has become completely oxidized is necessarily safe for drinking. 
Indeed if the analysis shows pollution, even at a remote source or 
time, the water should be condemned as dangerous. The degree 
of danger is still indicated and the indication will prove of 

As compared with most other substances ordinarily considered 
in quantitative analysis the different forms of nitrogen occur in 
jextremely slight concentrations. Unusually delicate reactions 
i must be used in order to cause the figures to have any value. 
Ordinary gravimetric or volumetric processes are rarely used 
in this connection but very sensitive color reactions are 
Imade the basis for the comparison of the water with color 

Free Ammonia is made evident by the brown color produced 
iwhcn a solution of potassium mercuriodide, K 2 HgI 4 , is added. 
This solution is known as "Nessler's reagent," from the name of 
the discoverer of the reaction. 1 The compound that is produced 
when ammonia is added is a complex substance, thought to 
have the composition Hg 2 NI. It is an intensely colored brown 
substance of small solubility and gives a visible color in water 
containing one part of nitrogen as ammonia in ten million parts 
jof water. 

The process of determining free ammonia is one of comparing 
jthe color produced by adding Nessler's reagent to water with 
ithat produced by the reagent with a standard solution of ammo- 
jiiium chloride. The comparison is made in tubes of colorless 
'glass, the two that are being compared having the same cross 
(section so that the same length of column is placed in the line of 
vision. The color is observed by looking vertically downward 
'through the tubes at a white surface placed at an angle in front 
'of a window so as to reflect the light upward. 

If Nessler's reagent is added directly to water containing 
organic matter, iron or aluminium, a precipitate is produced and 
an accurate color comparison is impossible. One of two pre- 
liminary treatments may be used. The free ammonia may be 
separated by distillation and the distillate then "Nesslerized" 

1 Z. anal. Chem., 7, 415 (1868). 




or reagents may be added to the water sample to precipitate int< 
fering substances and " direct Nesslerization" may be employed 
Distillation is preferable, but where apparatus or time is limite 
direct Nesslerization may be useful. 

Direct Nesslerization. — For precipitating organic matter use 
is made of the power of flocculating colloids for adsorbing this 
material, which is itself chiefly colloidal. Cupric sulphate is 
added to the water which is then made basic by the addition o: 
potassium hydroxide. The precipitating cupric hydroxide so 
clarifies the water that direct Nesslerization is practicable. 
Instead of adding cupric sulphate a solution of magnesium chlo- 
ride may be substituted. Colloidal magnesium hydroxide accom 
plishes the same result as does cupric hydroxide. If the water 
already contains much magnesium it is unnecessary to add even 
magnesium chloride. Boiling the water with potassium hydrox- 
ide will cause the precipitation of magnesium hydroxide. If 
hydrogen sulphide is present it will cause the precipitation of 
mercuric sulphide when Nessler's reagent is added. This inter- 
ference is prevented by the addition of lead acetate before the 
removal of colloids by cupric hydroxide. 

The chief objection to direct Nesslerization is the tendency of 
the precipitating colloids to adsorb small amounts of ammonia. 
However, the process is recommended for raw sewages, sewage 
effluents and highly polluted surface waters. 

Whether direct Nesslerization or distillation processes are used 
for free ammonia either an accurately prepared standard solution 
of an ammonium salt or a standard color solution of a permanent 
nature is required. This must have a very slight concentration 
and it is best made by successive dilutions of a more concentrated 
solution. The solvent used is water that has been shown to be 
free from ammonia by a test with Nessler's reagent. The labora- 
tory supply of distilled water is often free from ammonia. If 
it is not it may be purified by the addition of basic potassium 
permanganate solution and distilling. After the distillate no 
longer gives a test for ammonia it is collected and kept in well- 
stoppered bottles. 

For a permanent color standard mixtures of potassium chlor- 
platinate and cobalt chloride are recommended. By properly 
varying the relative concentrations of the two salts, solutions are, 

WATER 419 

obtained in which the color accurately corresponds with that of 
Nesslerized ammonia solutions of known concentrations. These 
solutions are to be preferred to the standard ammonium chloride 
solutions in laboratories where many determinations are to be 
made, because of their permanency. 

Albumenoid Nitrogen cannot be determined by direct Nessleri- 
zation. It is determined after the distillation of free ammonia by 
adding to the residue a basic solution of potassium permanganate 
and distilling. The organic matter is oxidized and remaining 
nitrogen is distilled and Nesslerized. 

As the ratio of nitrogenous organic matter to the ammonia 
obtained by distillation is decidedly variable in sewages and 
other substances containing much nitrogenous organic matter, 
albumenoid nitrogen results on such materials are less accurate 
than total organic nitrogen, obtained by the Kjeldahl process. 
Therefore in sewage work, including analysis of influents and 
effluents of purification plants and the water of highly polluted 
streams, it is recommended by the joint committee that deter- 
minations of total organic nitrogen be substituted for determina- 
tions of albumenoid nitrogen. For ground waters and surface 
waters containing but little pollution, the albumenoid nitrogen 
!is approximately one-half the organic nitrogen; accordingly the 
| continuance of albumenoid nitrogen determinations for this class 
! of work is approved. 

All determinations of nitrogen must be made in a laboratory 
in which the air is free from ammonia. 

Determination. — Prepare the following reagents: 

1. Nessler's Reagent. — Dissolve 25 gm of potassium iodide in the 
minimum quantity of cold water. Add a saturated solution of mercuric 

; chloride until a slight but permanent precipitate persists. Add 200 cc 
jof 50 percent solution of potassium hydroxide made by dissolving the 
I potassium hydroxide and allowing it to clarify by sedimentation before 
(using. Dilute to 500 cc, allow to settle and decant. This solution 
should give the required color with ammonia within five minutes after 
addition, and should not precipitate with small amounts of ammonia 
within two hours. 

2. Basic Potassium Permanganate. — Pour 600 cc of distilled water 
into a porcelain dish holding 1500 cc, boil 10 minutes and turn off the 
gas. Add 8 gm of potassium permanganate and stir until dissolved. 


Add 400 cc of 50 percent clarified solution of potassium or sodium 
hydroxide and enough distilled water to fill the dish. Boil down to 1000 
cc. Test this solution for albumenoid ammonia by making a blank 
determination. Correction should be made accordingly. 

3. Ammonia-free Water. — Test the laboratory supply of distilled wa 
by rinsing a clean Nessler tube several times, filling to the mark and add 
ing 2 cc of Nessler's reagent. Cover and allow to stand for 5 minutes. 
If the color produced at the end of this time is more intense than that of 
the diluted Nessler's reagent at first, the water must be purified. I 
this case add 10 cc of basic potassium permanganate solution to eac 
1000 cc of distilled water and distill, using a tin or aluminium condenser 
if one is available. After the distillate ceases to give a test for ammonia 
it is collected in a clean, glass-stoppered bottle, which is first rinsed with 
the distillate and the rinsings tested for ammonia. 

4. Standard Solution for Color Comparisons. — Use either (a) or (b). 

a. Ammonium Chloride Solution. — Dissolve 3.82 gm of ammonium 
chloride in ammonia-free water and dilute to 1000 cc. Dilute 10 cc of 
this to 1000 cc with ammonia-free water. 1 cc contains 0.00001 gm 
of nitrogen. 

b. Platinum Solution and Cobalt Solution. — Weigh 2 gm of potassium 
chlorplatinate, dissolve in a small amount of distilled water, add 100 
cc of concentrated hydrochloric acid and make up to 1000 cc. 

Weigh 12 gm of cobalt chloride and dissolve in distilled water; add 
100 cc of concentrated hydrochloric acid and make up to 1000 cc. 

Nitrogen of Free Ammonia. — A 750-cc Kjeldahl digestion flask, 
connected with a tin or aluminium condenser in such a way that the 
distillate may be conveniently delivered from the condenser tube 
directly into the Nessler tubes, is freed from ammonia by boiling dis- 
tilled water in it, until the distillate shows no further traces of free 
ammonia. When this has been done, empty the distilling flask and 
measure into it 500 cc of the sample, or a smaller portion diluted to 500 cc 
with ammonia-free water. Apply heat so that the distillation will be 
at the rate of not more than 10 cc nor less than 6 cc per minute. 

Collect four Nessler tubes of the distillate, 50 cc to each portion; 
these contain the free ammonia to be measured as described below. 

Use only Nessler tubes which do not show a variation of more than 
6 mm (0.25 inch) in the distance which the graduation mark (50 cc) is 
above the bottom. The tubes should be of clear white glass, with pol 
ished bottoms. The residue from the distillation is immediately used 
for the determination of albuminoid nitrogen as described below. 

The measurement may be made either by (1) comparison with Ness- 
lerized solutions containing known quantities of nitrogen as ammonium 

WATER 421 

chloride, or (2) comparison of the Nesslerized distillates with permanent 

Comparison with Ammonia Standards. — Prepare a series of 16 Nessler 
tubes which contain the following numbers of cubic centimeters of the 
standard ammonium chloride solution, diluted to 50 cc with ammonia- 
free water, namely: 0.0, 0.1, 0.3, 0.5, 0.7, 1.0, 1.4, 1.7, 2.0, 2.5, 3.0, 
3.5, 4.0, 4.5, 5.0, and 6.0. These will contain 0.00001 gm of nitrogen 
for each cc of the standard solution used. 

Nesslerize the standards and also the distillates by adding approxi- 
mately 2 cc of Nessler's reagent to each tube. Do not stir the contents 
of the tubes. 

Have the temperature of the tubes practically the same as that of the 
standards, otherwise the colors will not be directly comparable. 

Compare the color produced in these tubes with that in the standards 
by looking vertically downward through them at a white surface placed 
at an angle in front of a window so as to reflect the light upward. Allow 
1 the tubes to stand for at least 10 minutes after Nesslerizing before mak- 
ing the comparison. 

In case the color obtained by Nesslerizing the distillates is greater than 
that of the darkest tube of the standards, mix the contents of the tube 
thoroughly and pour out half of the liquid, making up the remainder 
to the original volume with ammonia-free water, then make the color 
comparison and multiply the result by two. If, after pouring out half 
of the liquid, the color is still too dark, repeat this process of division 
until a reading can be made. 

In case the color of the distillates is too high, this process may be 
shortened by mixing together all of the distillates from one sample 
before making the comparison, subsequently taking an aliquot portion 
for comparing with the standards. 

After the readings have been made and recorded, add together the 
results obtained by Nesslerizing each portion of the entire distillate 
from each sample. Calculate parts per million of nitrogen as free am- 
i monia in the sample. 

Comparison with Permanent Standards. — Prepare standards by put- 
iting varying amounts of potassium chlorplatinate and cobalt chloride 
solutions (page 420) in Nessler tubes, filling up to the mark with distilled 
'water as follows: 



Equivalent volume of standard 

Platinum solution. 

Cobalt solution, 

ammonium chloride, cc 






















































It is necessary to use tubes which have the 50 cc mark not less than 
20 nor more than 22 cm above the bottom. These standards may bt 
kept for several months if protected from dust. The method of cal- 
culating results is practically the same as with the ammonia standards. 

Albumenoid Nitrogen. — Interrupt the distillation (made as already 
described) after the collection of the distillate for free ammonia. Add 
50 cc or more of basic potassium permanganate solution and conduct 
this distillation until at least four portions of 50 cc each and preferably 
five portions of the distillate have been collected in separate tub 
Have enough permanganate solution present to insure the maxim 
oxidation of the organic matter. These distillates contain the alb 
menoid nitrogen as ammonia, measurement of which will be made 
described in connection with nitrogen as free ammonia. 

Nitrogen of Free Ammonia by Direct Nesslerization. — Prepare the 
following solutions. 

1. A 10 percent solution of cupric sulphate. 

2. A 50 percent solution of sodium hydroxide. 
If hydrogen sulphide is present in the water prepare also : 

3. A 10 percent solution of lead acetate. 

50 cc of the sample to be tested, mixed, if necessary, with an equal 
volume of ammonia-free water, is placed in a Nessler tube and a few 

WATER 423 

drops of cupric sulphate solution added. After thorough mixing 1 cc 
of the sodium hydroxide solution is added and the solution again thor- 
oughly mixed. The tube is then allowed to stand for a few minutes, 
when a heavy precipitate should fall to the bottom leaving a colorless 
supernatant liquid. Nesslerize an aliquot portion of this liquid. 

If hydrogen sulphide is present add a few drops of lead acetate solu- 
tion before the addition of potassium hydroxide. 

Organic Nitrogen. — This determination may be conveniently com- 
bined with that of ammonia nitrogen, by the procedure for water and the 
first procedure for sewage, described below. 

Procedure for Water. — Distill the ammonia from 500 cc of the sample 
exactly as already directed for the determination of nitrogen of free 
ammonia, page 420. This will usually involve the loss of 200 cc of the 
sample. Cool somewhat and rinse into a Kjeldahl digestion flask, unless 
such a flask was used for the first distillation. Add 5 cc of concentrated 
sulphuric acid which is free from ammonium sulphate, also a small piece 
of ignited pumice, dropped in while hot. Mix and digest over a free 
flame, using a suitable apparatus for removing the sulphuric acid 
fumes (see Figure 116, page 514). The digestion should be continued 
until copious fumes are evolved and the liquid is finally colorless or 
very pale yellow. Remove from the flame and add to the hot solution 
potassium permanganate crystals until a heavy green precipitate per- 
sists. Cool and dilute to about 300 cc with ammonia-free water. Make 
basic by adding 10 percent sodium hydroxide which has been made free 
from ammon'a by boiling for a short time. Distill the ammonia and 
Nesslerize as directed for the determination of nitrogen of free ammonia. 

First Procedure for Sewage. — Use 100 cc or less of the sample, according 
to the amount of ammonia expected. Dilute to 500 cc with ammonia- 
pee water in the distilling flask and distill the free ammonia, Nesslerizing 
the distillate. Cool somewhat and add 5 cc of nitrogen-free concen- 
trated sulphuric acid and 1 cc of 10 percent copper sulphate solution. 
Digest the solution for a half hour after it has become colorless or pale 
(yellow. Add 0.5 gm of potassium permanganate crystals to the hot 
acid solution, cool, transfer to a 500 cc volumetric flask, dilute to 
Ihe mark and mix. By means of a pipette transfer 10 cc of this solution 
to a Kjeldahl distilling flask and dilute to about 300 cc with ammonia- 
free water. Make basic with 10 percent sodium hydroxide solution, 
jlistill and Nesslerize. (With some samples direct Nesslerization may 
j)e employed.) 

!| Second Procedure for Sewage. — Omit the separation of ammonia 
nitrogen and determine this and organic nitrogen together. Determine 
.he ammonia nitrogen on a separate sample by direct Nesslerization as 


directed on page 422. The difference between these results is the organic \ 

Nitrogen as Nitrites. — It has already been explained that 
nitrites will not normally occur in more than traces in water 
because of the readiness with which they oxidize. In order to I 
give this determination any significance it is necessary to use a 
very delicate test. Use is made of the ready action of nitrous 
acid with aromatic amines, forming diazo compounds, and of 
the latter with naphthylamine, forming azo dyes of intense color- 
ing power. When water containing nitrites is acidified and 
sulphanilic acid (p-amidobenzenesulphonic acid) is added there 
is formed the anhydride of p-diazobenzenesulphonic acid, thus: 1 

/NH 2 /N = N X 

HN0 2 +C 6 H 4 < ->C 6 H 4 < > +2H 2 0. 

NSOsH > S0 3 / 

If to this solution a-amidonaphthalene is added, an azo dye, 
a-amidonaphthaleneazobenzene-p-sulphonic acid, is produced. 

>N = N x .N = NC 10 H 6 NH 2 

C 6 H 4 < > +C 10 H 7 NH 2 ->C 6 H 4 < 

X S0 3 / \SO3H 

This dye possesses a very intense red color and one part of nitro- 
gen as nitrite can be detected in 1,000,000,000 parts of water. 

The amino compounds entering into these reactions are not 
easily soluble and their soluble salts are used. The hydrochlo- 
rides may be employed but the reactions appear to proceed more 
rapidly if the acetates are used. 

Determination. — Prepare the following reagents: 

1. Sulphanilic Acid. — Dissolve 4 gm of the purest sulphanilic acic 
in 500 cc of 5-normal acetic acid (sp. gr. 1.041) or in a mixture of 95( 
cc of nitrite-free water and 50 cc of concentrated hydrochloric acid 
This is a practically saturated solution. 

2. a-amidonaphthalene Acetate or Hydrochloride. — Dissolve 2.5 gm 
solid a-naphthylamine in 500 cc of 5-normal acetic acid or in 1000 c< 
of nitrite-free water- containing 8 cc of concentrated hydrochloric acid 
Filter the solution through washed absorbent cotton or an alundun 

3. Sodium Nitrite, Stock Solution. — Dissolve 1.1 gm of silver nitrit* 
in nitrite-free water. Precipitate the silver with sodium chlorid< 
solution and dilute the whole to 1000 cc. 

WATER 425 

4. Standard Sodium Nitrite Solution. — Dilute 100 cc of solution (3) 
to 1000 cc and dilute 10 cc of the resulting solution to 1000 cc with 
sterilized nitrite-free water. Add 1 cc of chloroform and preserve in 
a sterilized bottle. Calculate and record the weight of nitrogen in 1 cc 
of this last solution. 

5. Fuchsine Solution. — 0.1 gm per liter of water. 
Measure 50 cc or 100 cc of the water to be tested into a Nessler 

tube. These Nessler tubes should be of clear, white glass, with the 
graduation mark not varying more than 6 mm in its distance above the 
bottom. At the same time make a set of standards by diluting various 
volumes of the standard nitrite solution in Nessler tubes to 50 or 100 cc 
with nitrite-free water, for example, 0, 1, 3, 5, 7, 10, 14, 17, 20 and 25 cc. 
Add 2 cc of reagents (1) and (2) to each 100 cc of the sample and to 
each standard. Mix and allow to stand 10 minutes. Compare the 
samples with the standards. Do not allow the samples to stand over 
one-half hour before being compared. Make a blank determination in 
all cases to correct lor the presence of nitrites in the air, the water and 
the reagents. Dilute all samples which develop more color than the 
30 cc standard before comparing. Mixing is important. Calculate 
milligrams per liter of nitrite nitrogen. 

Permanent standards may be prepared by matching the nitrite 
standards as above made against dilutions of the fuchsine solution. 
Fuchsine standards have been found to be sufficiently accurate 
for waters high in nitrite and for sewage. Such standards should 
be checked once a month and should be kept out of bright 
sunlight to avoid bleaching. 

Nitrogen as Nitrates. — When a soluble aromatic sulphonic 
acid is mixed with nitrates and sulphuric acid the nitric acid so 
liberated acts upon the aromatic compound and produces nitro- 
j derivatives which are faintly yellow in most cases. If a base is 
now added the sulphonate is formed and this is much more in- 
; tensely colored. These reactions are applied to the determina- 
j tion of nitrates in water by what was originally known as Spren- 
I gel's method. 1 

The method was further modified by Grandval and Lajoux. 2 

• A measured volume of water is evaporated to dryness, sodium 

! carbonate having been first added if the water contains free acid. 

The dry residue is treated with a small amount of a phenolsul- 

1 Pogg. Ann., 121, 188 (1863). 

2 Compt. rend., 62, 101 (1885). 


phonic acid, the mono-nitro derivative being formed. Th< 
reagent is made by heating phenol with sulphuric acid in the pre 
portions indicated below. These interact with the f ormatioi 
phenol-o-p-disulphonic acid. The reaction of this acid with ni- 
tric acid results in the formation of o-nitrophenol-o-p-disulphonic 

C 6 H8.0H.(S0 8 H) a +HN08->C6H2.0H.N02.(S03H) 2 +H20. 

Treatment with a base produces the highly colored sulpho- 
nate, e.g. y 

C6H 2 .OK.N02.(S0 3 K)2< 

Three important sources of error may render impossible 
determination of ^nitrate by the phenolsulphonic acid method or 
may cause incorrect results to be obtained. These may be 
enumerated as follows: 

Interference of Organic Matter. — If the water contains any 
considerable amount of organic matter, as is always the case 
with surface streams, sewage or waters contaminated by sewage, 
the addition of sulphuric acid will cause a charring of the organic 
matter and the color comparison cannot be made accurately 
because of the resulting brown coloration. Of course the same 
interference will result from any color that may have been in the 
sample before treatment. Organic matter may also cause the 
reduction of the nitrates during evaporation. This interference 
may be prevented by first removing the organic matter by means 
of colloidal aluminium hydroxide. 

Interference of Chlorides. — If the sample contains more tha 
about 30 mg of chlorine as chlorides per liter there is a possibilit; 
of reactions occurring between chlorides and nitrates during evapo- 
ration, resulting in the reduction of nitrates. This may be 
avoided by the addition of silver sulphate solution to the slightly 
acidified sample, silver chloride being precipitated. 

Interference of Nitrites. — Nitrites likewise cause variabl 
results to be obtained unless certain precautions are taken 
During the evaporation of water containing nitrites some of th 
latter will be decomposed and nitrogen lost, while some may b 

1 Chamot and Pratt: J. Am. Chem. Soc, 31, 922 (1909); 32, 630 (1910) ; 
3, 366 and 381 (1911). 


WATER 427 

(oxidized to nitrates. On account of the uncertain extent to 
[which these reactions occur it is necessary either to remove the 
(nitrites entirely or to oxidize them quantitatively to nitrates. 
[The latter is accomplished by treatment of the sample with 
hydrogen peroxide. 

All of these interferences may be avoided by using a method 
based upon the reduction of nitrates to ammonia by nascent 
hydrogen. The sample is made basic and concentrated by boil- 
ing. By this means all of the ammonia, free or combined, as 
well as all of the nitrogen as nitrites is removed. Nascent hydro- 
gen is then generated by adding aluminium to the basic solution. 
The resulting ammonia is later distilled and determined by 

Determination .—For the phenolsulphonic acid method the following 
reagents will be required: 

1. Phenoldisulphonic Acid. — Dissolve 25 gm of pure white phenol in 
150 cc of concentrated sulphuric acid. Add 75 cc of fuming sulphuric 
acid containing 15 percent of "free" sulphur trioxide, stir well and heat 
for 2 hours at about 100°. 

2. Sodium or Potassium Hydroxide Solution. — This solution should be 
| made approximately 12 normal. About 5 cc will then be required to 

neutralize 2 cc of the phenolsulphonic acid. 

3. Standard Nitrate Solution. — Dissolve 0.72 gm of pure recrystal- 
lized potassium nitrate and dilute to 1000 cc with distilled water. 
Evaporate cautiously 10 cc of this solution in a dish placed on a steam 
bath. Moisten the residue quickly and thoroughly with 2 cc of phenol- 
sulphonic acid, dissolve and dilute to 1000 cc. Calculate the weight of 
nitrogen in l*cc of the last solution. 

4. Standard Silver Sulphate Solution— Dissolve 4.397 gm of silver 
sulphate, free from nitrate, in nitrate-free water and dilute to 1000 cc. 
If a good grade of silver sulphate has been used this solution will require 
no other standardization for the purpose for which it is here to be used. 
Otherwise it should be standardized against pure sodium chloride. 
1 cc should be equivalent to ling of chlorine. 

5. Standard Sulphuric Acid Solution.— A solution, N/50, standardized 
by titration against pure sodium carbonate (pages 224 and 231). The 
solution must be free from nitrates. 

6. Aluminium Hydroxide.— This must be freshly prepared. From 
a solution of nitrate-free aluminium sulphate or alum precipitate hy- 
droxide by adding dilute ammonium hydroxide. Filter and wash several 
times. The water and ammonium hydroxide must be free from nitrates. 


The aluminium hydroxide so prepared is used without drying and 
before it has had time to change to the crystalloidal form. 

Measure into an evaporating dish enough sample to furnish not more 
than 0.01 mg of nitrate nitrogen. 100 cc will be suitable for ordinary 
unpolluted waters. Add sufficient N/50 sulphuric acid to make the 
water nearly neutral, as determined by a separate titration with methyl 
orange as indicator. (See the determination of carbonates, page 398.) 
Now add silver sulphate solution in quantity sufficient to precipitate 
all but about 0.1 mg of chlorine. (This determination is described 
on page 397.) This treatment may be omitted if the quantity of sample 
used contains less than about 3 mg of chlorine. If silver sulphate 
has been added or if the water sample is perceptibly colored heat the 
mixture to boiling, add a little aluminium hydroxide, stir, filter and 
wash with small amounts of nitrate-free hot water. Evaporate over 
the steam bath and add 2 cc of phenoldisulphonic acid, rubbing with 
glass rod to insure intimate contact. If the residue becomes packed or 
appears vitreous because of the presence of much iron, heat the dish on 
the steam bath for a few minutes. 

Dilute the mixture with distilled water and add slowly solution (2) 
until the maximum color is developed. Transfer to a Nessler tube, 
filtering if necessary. Compare the color with that of standards 
made by adding 2 cc of the basic solution to various amounts of the 
standard nitrate solution and diluting to the mark in Nessler tubes. 
The following amounts of standard are suggested: 0.5, 1.0, 1.5, 2.0, 
4.0, 6.0, 8.0, 10.0, 15.0, 20.0 and 40.0 cc. These standards may be 
kept for several weeks without deterioration if the tubes are kept corked 
to prevent evaporation or contamination. 

The amount of nitrite nitrogen that will remain after evaporation is 
not sufficient to alter materially the results unless present in excess of 
1 mg per liter. In case such quantity is found the nitrite is oxidized to 
nitrate by repeatedly heating with a few drops of hydrogen peroxk 
which is free from nitrate. Proper correction in the nitrate nitroge 
figure must then be made. 

Calculate milligrams per liter of nitrate nitrogen. 

For the reduction method for nitrates prepare the following reagents 

1. Sodium Hydroxide Solution. — Dissolve 50 gm of the purest ob- 
tainable sodium hydroxide in 250 cc of distilled water, add several 
strips of aluminium foil and leave over night. Evaporate to 200 cc by 

2. Aluminium Foil. — Use strips about 10 cm long, 0.33 mm thick 
and 6 mm wide. 

Place 100 cc of the water in a 300 cc casserole or dish. Add 2 cc 
of sodium hydroxide solution and boil until the volume is 20 cc. Using 

WATER 429 

nitrogen-free water, rinse into a test-tube about 15 cm long and 3 cm 
in diameter. The volume of the solution should now be about 75 cc. 
Add a strip of aluminium foil and close the tube with a rubber stopper 
through which passes a bent glass tube, the shorter end of which is 
flush with the lower end of the stopper, the longer end dipping beneath 
the surface of distilled water in a beaker. This serves as a trap to pre- 
vent the entrance of oxygen. 

Allow to stand for four hours or more. If the supernatant liquid is 
then clear and colorless, Nesslerize at once, otherwise rinse into the 
apparatus used for ammonia determinations, dilute to 250 cc, distill 
and Nesslerize the distillate. Calculate milligrams per liter of nitrate 

Required Oxygen. — Besides the indirect estimation of organic 
matter through the determination of nitrogen in its various 
forms a more direct estimation may be made by oxidizing with 
standard potassium permanganate. It is readily seen that no 
calculation of organic matter can be made as a result of such a 
titration because the great variety of organic substances present 
in polluted water gives rise to a great variety of reactions. On 
the other hand the calculation of the oxygen required to oxidize 
all reducing agents in the water gives a fair, though inexact idea of 
the amount of organic pollution. Differences in procedure will 
cause the reduction of varying quantities of potassium per- 
manganate and it is therefore necessary rigidly to standardize 
the method. It is also necessary to correct the results according 
to the amount of nitrites, ferrous salts and hydrogen sulphide, 
if these are found in considerable quantities. 

The determination is carried out by treating a measured 
volume of water with an excess of standard potassium perman- 
ganate solution, at a specified temperature, and titrating the 
excess after a stated period. If the solution is heated during the 
treatment the excess of permanganate is determined by adding a 
measured excess of a standard solution of oxalic acid or ammo- 
nium oxalate, titrating the excess by standard potassium per- 
manganate solution. If cold treatment is used the excess of 
potassium permanganate cannot be determined in this manner 
because oxalic acid reduces permanganates very slowly unless 
heated to at least 60°. In this case the excess of potassium 
permanganate is reduced by adding potassium iodide and titrat- 


ing the liberated iodine with standard sodium thio sulphate 
solution : 

2KMnO4+10KH-8H2SO4-*6K 2 SO4+2MnSO4+8H 2 O + 10I; 
2Na,S 2 0s+2I->Na 2 S 4 06+2NaL 

The method involving digestion at 100° will be described. 

Determination. — Prepare the following reagents: 

1. Sulphuric Acid. — Dilute the concentrated acid with three volumes 
of distilled water. Add potassium permanganate until a faint pink 
color persists after standing for several hours. 

2. Standard Potassium Permanganate Solution. — Calculate the weight 
of crystallized potassium permanganate required for 1000 cc of a solu- 
tion, 1 cc of which shall be equivalent to 0.1 mg of oxygen. Dissolve 
this weight of salt in 1000 cc of distilled water. 

Standardize as follows: Add 10 cc of this solution and 10 cc of solution 
(1) to 100 cc of distilled water in a flask, immersing the flask in boiling 
water for 30 minutes. This destroys the oxygen -consuming capacity j 
of the distilled water. Add 10 cc of solution (3) and then potassium 
permanganate solution until a faint pink color persists. Now add 10 
cc more of the oxalate solution and titrate with the permanganate solu- 
tion for the standardization of the latter. Calculate its value in milli- 
grams of available oxygen per cubic- centimeter. 

3. Ammonium or Sodium Oxalate Solution. — Use the purest obtain- 
able salt. Make a solution of which 1 cc is equivalent, as a, reducing 
agent, to 0.1 mg of oxygen. 

Measure 100 cc or less of the water into a 200-cc flask, add 10 cc of 
solution (1) and 10 cc (exactly measured) of solution (2) and immediately 
place the flask in a bath of boiling water, the water level of which is 
kept above the level of the contents of the flask. Digest for exactly 
30 minutes. Remove the flask and add exactly 10 cc of solution (3). 
Titrate with standard potassium permanganate solution and calculate 
milligrams per liter of required oxygen. 

If 10 cc of permanganate solution is insufficient for complete oxidation 
repeat the digestion with a larger quantity. At least 5 cc excess of 
permanganate should remain after the digestion. 

Interfering Substances. — If oxidizable mineral substances, 
such as ferrous iron, sulphide or nitrite, are present in appreciable 
quantities corrections should be applied as accurately as possible 
by suitable procedures. Direct titration of the acidified sample 
while cold, using a three-minute period of digestion, serves this 

WATER 431 

urpose quite well for polluted surface waters and fairly well 

or purified sewage effluents. Few raw sewages containing no 

rade wastes need such a correction but raw sewages contaning 

I pickling' ' liquors do need it. If the sample contains bothi oxi- 

dizable mineral compounds and gaseous organic substances the 

latter should be driven off by heating, the sample being allowed 

to cool before applying this test for the correction factor. If 

such corrections are made the fact should be stated, with the 

amount of correction. 


Steel and Cast Iron 

The impurities contained in iron and steel usually form a very 
small portion of the total mass. Wrought iron and steel often j 
contain a total of less than one percent of elements other than 
iron while even pig iron does not often contain as much as ten 
percent of other elements. It is therefore not customary to make I 
determinations of the percent of iron but rather of the small, 
amounts of other elements, which give certain very important 
properties to the metal in which they are contained'. Elements 
occurring in iron and steel and commonly determined are carbon, 
silicon, phosphorus, sulphur, manganese and titanium. In alloy 
steels for special purposes determinations are also made of tung- 
sten, nickel, chromium, molybdenum, vanadium and copper. 

Exact and Rapid Methods. — For the determination of each 
element there are available certain well known methods and these 
are continually being revised and supplemented by other newer 
methods. Considerable experience is therefore necessary if the 
analyst is to be able to ' intelligently select the method best 
adapted to his purpose. There is a certain distinction to be 
made between what may be classed as " exact" methods and 
others that are more properly called " rapid' ' methods. Thus a 
determination of carbon in steel, made by an approved exact 
method, may require at least two hours and sometimes longer 
while a less exact determination might be made by another 
method in ten or fifteen minutes. Conversations with works 
chemists will often give the student the erroneous impression 
that the longer methods are impracticable and are taught in 
colleges but not used in practice, while the rapid methods are 
improperly neglected in the students' college courses. It is true 
that more emphasis is laid upon the exact method, as a rule. If 



the science and careful manipulation involved in the longer 
method are properly appreciated and learned the student will 
have no difficulty in learning the shorter and less exact method 
after he enters his professional career. 

, It is highly important that one should understand the proper 
place of each class of methods. In the steel works samples may 
be taken from the melted iron as it runs from the blast furnace 
or from the steel ladles which receive the product of the steel 
furnaces. These samples are taken directly to the works labora- 
tory where the analysis must be made very quickly in order to 
furnish information which will serve as a guide in mixing charges 
for the steel furnace or for properly disposing of or modifying the 
product of a given furnace. The results of such an analysis do 
not often serve as a guarantee to the steel consumer but rather 
as a check upon the various stages in the process of steel manu- 
facture. For this reason rapid methods are quite suitable for 
the purpose. 

When the steel is placed upon the market as a finished product 
the case is quite different. Modern industrial development has 
created new and rigorous requirements regarding the quality of 
steel entering into machinery and structural work. It becomes 
necessary for the steel manufacturer to guarantee the percents 
of the elements in his steel within very narrow limits and a 
method of analysis that will not give results having a high degree 
of accuracy is quite useless for this purpose. 

Standard Methods. — An inspection of the methods for analysis 
of steel as practised in the various works laboratories will show 
that while the several standard methods are quite universally 
used, many variations have been introduced by individual chem- 
ists. Each laboratory usually has its methods described and 
specified and these must be rigidly followed by all chemists 
working in that laboratory. There is, of course, much difference 
of opinion concerning the relative merits of different methods and 
it is inevitable — indeed it is even desirable — that modifications 
should be made whenever any improvement is seen to be possible. 
It is also true, however, that many modifications of good methods 
have made poor methods because the modifications have been 
made without an adequate knowledge of the scientific principles 
underlying the analytical process. Many chemists have con- 



fidence in their methods when this confidence is based upon little 
more than the ability to obtain close agreement of duplicate 
determinations. The error involved in such conclusions has 
been discussed in an earlier section. Many of the analytical 
methods for iron and steel have been in use for a long time. 
Some of these have been retained in practically their original 
form and still bear the names of the chemists who first proposed 
them. Others have been so modified that they bear little 
resemblance to the original method. 

Sampling. — Analysis may be required of either works samples, 
taken from the metal as it runs from the furnace, or of the finished 
product. Samples of the first class are dipped from the melted 
metal by means of a small ladle and are poured onto a clean iron 
plate or into a small iron mold. The sample is crushed, if a brittle 
product like pig iron, or drilled if steel. Pieces of already solidi- 
fied metal are drilled to obtain a sample for analysis. The outer 
case should be first removed because it may contain iron oxide or 
sand, or the percent of carbon may have been lowered by oxida- 
tion. The drill should be set to make as fine drillings as possible 
and if powder is at the same time produced it should be well 
mixed with the larger pieces before weighing for analysis. 

Solution and Evaporation. — Steel and iron analysis involves 
many operations of dissolving and of evaporating solutions. 
The work of the iron and steel laboratory must usually be done 
as quickly as possible and the analyst must therefore give 
considerable attention to the best manipulation, from the stand- 
point of speed and accuracy. 

Dissolving metals in strong acids is always attended with 
the evolution of disagreeable and often poisonous fumes and 
good draught hoods are absolutely essential for carrying on 
this work. The evolution of gases and the boiling of solutions 
for evaporation will occasion loss of the dissolved matter unless 
proper attention is given to the prevention of such loss. Whei 
evaporation is not to follow solution of the sample an Erlenmeyer 
flask is usually the best vessel for the purpose. Loss by spat 
tering is thus reduced to a minimum. Casseroles should be usee 
if the solution of the sample is later to be evaporated. These 
are covered with glasses until the sample has all dissolved anc 
evolution of gases is completed. The cover glasses are then 


rinsed down and removed and the solution is placed on a steam 
bath or hot plate or it is held over a free flame and rotated con- 
tinuously by hand. 

The choice of method to be used for evaporating will depend 
| upon the requirements of. the case. If time does not press and 
I work may be so fitted together as to carry on a large number 
of analyses at a time, evaporation on the steam bath will prove 
to be a convenient and safe process, except for solution of high 
boiling points such as those containing sulphuric acid. If the 
opposite is true and speed is a matter of prime importance, 
evaporations must be hurried but the resulting danger of loss 
through bumping and spattering requires that the casserole 
should be held in the hand and kept in continuous motion. 
This constant agitation considerably increases the rate of evapora- 
tion, particularly because of the spreading of the solution over 
the sides of the casserole and the consequent increase in the 
effective surface. Of course this process is carried out with 
an uncovered casserole as otherwise little advantage would be 
derived from forced boiling. 

Standard Samples. — In the description of the volumetric 
[determinations that follow, methods are given for the standardi- 
zation of the solutions against suitable primary standards. In 
many of these cases it is also convenient to standardize the solu- 
tion against a steel or iron in which the percent of the element in 
question is accurately known, the standard sample being weighed 
tand treated exactly as is the sample that is being analyzed. 
This procedure possesses a further advantage in that it auto- 
matically corrects for any deviation from the theoretical course 
of the reactions occuring during the preparation of the solution 
lor during the titration itself. 

The steel laboratory may prepare and carefully analyze its own 
jstandard samples. It is also possible to obtain most of the 
necessary standard steels and irons from the Bureau of Standards. 
These standard samples may also be used for checking the accu- 
racy of gravimetric methods and in this way they will serve for 
jcontrol work. 

Carbon.— The most important element occurring in steel is 
jcarbon. This is because it is the element which makes possible 
ithe formation of steel by imparting to iron the capability of being 


hardened by suitable heat treatment. The development of 
alloy steels has lately brought nickel, chromium, vanadium, 
tungsten and other metals into prominence as constituents of 
special steels, but without carbon, alloys of these elements with 
iron would be of little value. The effect of carbon upon iron 
with which it is combined is to increase the tensile strength and 
hardness and to decrease the ductility. Carbon is present in 
steel chiefly as a carbide, FesC, although small quantities may 
occur as free carbon. In cast iron large quantities of carbon are 
free, particularly in gray cast iron. A more extended discussion 
of the properties of steel, as dependent upon the condition of the 
carbon, will be taken up later (page 478). 

The determination of total carbon is the only carbon determina- 
tion that is usually required in steel analysis. In the analysis 
of cast iron determinations also of free and combined carbon may 
be required. The determination of total carbon is generally 
made by a combustion process, the carbon being oxidized and the 
resulting carbon dioxide determined, although oxidation in solu- 
tion has been employed, the oxidizing agent being chromic acid. 
The details of the combustion processes vary widely. Fine 
drillings of steel may be burned directly or a preliminary separa- 
tion of the carbon may be made. The apparatus for combustion 
may be a furnace and combustion tube or a special form of closed 
crucible, through which air and oxygen may be passed. The 
resulting carbon dioxide may be measured at an accurately 
observed temperature and pressure and its weight calculated or 
it may be absorbed in a basic solution and either weighed or the 
excess of base determined by titration. 

Direct Combustion. — Iron or steel may be completely burned 
in oxygen at 900° to 1100°. The method is desirable because 
it avoids a rather tedious process of preliminary solution, filtra- 
tion and washing. For the combustion there is required a tube 
of quartz or porcelain, 24 inches long and having an inside 
diameter of 3/4 inch. It must be protected from contamination 
by iron oxide by placing an alundum cylinder in the section 
which is to hold the boat. The combustion furnace may be 
heated by gas or electricity and should be 12 to 14 inches in 
length. The combustion is carried out in a manner quite similar 
to the combustion of coal (page 311). The long furnace and 


tube there used are not required in this case because no volatile 
hydrocarbons are produced and long contact of the gases with 
cupric oxide is not necessary. The small amount of carbon 
monoxide that may be formed at first is completely oxidized 
by passing the mixture with oxygen over a small amount of 
cupric oxide or through platinized asbestos. The platinum 
black in the latter case catalyzes the combination of carbon 
monoxide with oxygen. 

The train of apparatus necessary for the gravimetric determina- 
|tion is as follows: (1) Oxygen tank, (2) bubble tube of 30 percent 
I potassium hydroxide solution, (3) tube of soda lime to retain 
spray from the potassium hydroxide solution and to complete 
the removal of carbon dioxide, (4) combustion tube containing 
(4a) space of about 3| inches inside the furnace, lined with 
an alundum cylinder to hold the combustion boat, (4b) platinized 
asbestos to the end of the furnace, leaving the projecting ends 
of the tube empty, (5) U-tube filled with granular zinc or with 
glass beads moistened with chromic acid solution for the re- 
tention of oxides of sulphur, (6) U-tube filled with calcium- chlo- 
ride and (7) absorption apparatus for carbon dioxide, discussed 
below. The method of preparing and assembling this apparatus 
should be clear from the discussion of the determination of carbon 
dioxide in carbonates (page 129) and of the determination of 
carbon and hydrogen in coal (page 311). 

Absorption Apparatus. — The nature of the absorption ap- 
paratus (7) will depend upon the method chosen for the final 
determination. This may be either gravimetric or volumetric 
in principle. One of the following variations is recommended: 

(a) Absorption in Weighed Bulbs of the Geissler or a Similar 
Type, Filled with 33 Percent Potassium Hydroxide Solution 
(and Carrying a Prolong Filled with Calcium Chloride. — The 
I manipulation of these bulbs is carried with precautions similar 
to those observed in the determination of carbon dioxide in 
I carbonates. (See pages 133 to 136.) 

(b) Absorption in Weighed Tubes Carrying Soda Lime and 
I Calcium Chloride or Phosphorus Pentoxide. — The Fleming tube 
[(figure 95) is satisfactory for this purpose and a determination 
jean be made very rapidly by this method. In filling, the part 
j(s) is to contain, first, a loose plug of asbestos then alternating 



layers, about one-fourth inch thick, of 20-, 40- and 60-mesh soda 
lime. Part (d) is filled with either phosphorus pentoxide on 
glass wool or calcium chloride. (If phosphorus pentoxide is 
used at this point it must also be substituted for calcium chloride 
in tube (6) of the absorption train.) 

Phosphorus pentoxide should not be used 
after it has become visibly moist. On this 
account the charging of tubes with this sub- 
stance should not be undertaken on days 
when the humidity is high as it is then im- 
possible to work rapidly enough to prevent 
absorption of considerable quantities of 
moisture from the air. To prepare phos- 
phorus pentoxide for this purpose a loose 
ribbon of glass wool is first spread on a 
clean piece of paper and the dry material is 
sifted over the ribbon in an even layer. 
The edges of the glass wool are then turned j 
over and the material is rolled into the 
proper shape to fit loosely into the tube. 
By this means the drying agent is given the 
maximum exposure to gases passing through 
the tube. 

(c) Absorption in Standard Barium 
Hydroxide Solution, the Unused Base being 
Titrated in Presence of Phenolphthalein 
after the Absorption is Finished. — The 
Meyer tube (figure 96) is suitable for this 
purpose and it possesses the important merit 
of being easily emptied and rinsed just before the titration. 
The drying tube (6) should be omitted in this and all other 
volumetric methods as the drying of gases entering the carbon 
dioxide absorption tube is quite unnecessary. 

A saturated solution of barium hydroxide in carbon dioxide-free 
water is prepared and kept as a stock solution. This will contain 
approximately 3.9 gm of barium hydroxide, calculated as the 
anhydrous base, in each 100 cc of solution and it will be approxi- 
mately 0.45 normal. 550 cc of this solution is diluted to 1000 
cc. The resulting solution is then approximately 0,25 normal/ 

Fig. 95. — Fleming 
absorption tube. 


so that 1 cc is equivalent to about 0.0015 gm of carbon. This 
solution should be kept in a bottle which is provided with a 
siphon or with an outlet at the bottom and protected in such a 
manner as that entering air shall be drawn through a tube of 
soda lime or of saturated barium hydroxide solution, to remove 
carbon dioxide. 

As carbon dioxide is absorbed in barium hydroxide solution 
barium carbonate is formed and almost entirely precipitated. 
One liter of a saturated solution of barium carbonate in water 
contains 0.022 gm (corresponding to 0.00134 gm of carbon) 
at 20°. Therefore it is possible to titrate the excess of unused 

Fig. 96. — Meyer absorption bulbs. 

base after the absorption is finished, using a standard acid of 
concentration equivalent to that of the barium hydroxide and 
without filtration, provided that the solution was previously 
saturated with barium carbonate and that the titration is 
carried out rapidly and with stirring, so as to avoid resolution 
of the precipitate by local excess of acid. Phenolphthalein 
is used as the indicator. The first condition, above, is usually 
met by the fact that commercial barium hydroxide always 
contains small amounts of carbonate and that more is formed 
by carbon dioxide in the water that is used for preparing the 
standard solution. 

(d) Absorption in Standard Barium Hydroxide Solution, 
the Unused Excess being Titrated after Filtration and Washing 
with Carbon Dioxide-free Water. — This procedure has no ad- 
vantage over that outlined as (c). Absorption of carbon dioxide 
from the air takes place during filtration and washing of the 
precipitate. Also the addition of wash water that is not 
saturated with barium carbonate causes resolution of traces of 
the precipitate. The second error tends to compensate the 


first but a method that depends for accuracy upon mutual com- 
pensation of variable errors is not ideal. 

(e) Absorption in Saturated Barium Hydroxide Solution, 
the Precipitated Barium Carbonate being Removed by Filtration, 
Washed, Dissolved in an Excess of Standard Acid and the 
Unused Excess of Acid Titrated by Standard Base, with Methyl 
Orange as Indicator. 1 — The barium hydroxide solution that is 
used in this method may be much more concentrated than that 
used in methods (c) and (d) because it is not to be titrated. On 
this account absorption is more certain when the gas is bubbled 
rapidly through the solution and in this lies the only important 
advantage over the other volumetric methods. The sources qf 
error noted under (d) obtain here also. 

Drying tubes are omitted from the absorption train in volu- 
metric methods (c), (d) and (e). 

Combustion, Preceded by Solution of the Iron and Separation 
of the Carbon. — Iron or steel dissolves easily in a solution of the 
double chloride of potassium and copper. The cupric chloride 
is the active agent and the double salt is used only because it is 
more easily purified and preserved. The reactions are : 

Fe+2CuCl 2 ->FeCl 2 +2CuCl and 
Fe 3 C+6CuCW3FeCl 2 +6CuCl+C. 

Free carbon is also left undissolved. The residue contains organic 
compounds, formed during the process of solution, and the total 
residue cannot, therefore, be weighed directly for the determina- 
tion of total carbon. If the solution is not well stirred or if not 
enough cupric potassium chloride is used, copper will separate, 
returning to the solution upon stirring or addition of more of the 
solution of cupric salt: 

Fe+CuCl 2 — FeCl 2 +Cu, 
Cu + CuCl 2 ->2CuCl. 

These reactions illustrate the principle of replacement of one 
metal of a salt solution by another which has a higher solutioi 

The solution of cupric potassium chloride must contain h; 

1 Cain: J. Ind. Eng. Chem., 6, 465 (1914). 


Jdrochloric acid in order to prevent the precipitation of cuprous 
[chloride, a substance having very small solubility in water. If 
too much acid is present carbon may be lost through the forma- 
tion of hydrocarbons during the process of solution. This is 
typified by the hypothetical reaction shown by the following 
fequation : 

2Fe 3 C + 12HCl-+6FeC h + C 2 H 2 + 5H 2 . 

Concerning the choice of methods it may be said that direct 
combustion is much more rapid and is accurate if the metal is in a 
proper state of division. The method of preliminary solution is 
fully as accurate, except for high-speed tool steels, and is safer if 
the nature of the steel is not accurately known. There is little 
to choose between the gravimetric method, weighing the absorp- 
tion bulbs before and after the absorption of carbon dioxide, 
and the volumetric methods. The so-called " moist combustion " 
processes and the methods involving measuring the volume of 
carbon dioxide evolved, while attractive in principle, are trouble- 
some in execution and are subject to large errors unless great 
care is exercised. These are therefore little used. 

Determination by Direct Combustion. — The apparatus train described 
on page 437 is assembled, drying tube (6) being used only in the gravi- 
metric modifications. If a high-pressure oxygen cylinder is used to 
upply this gas a special control valve must be provided. 

The entire apparatus is first tested for leaks. With the carbon dioxide 

bsorption tube temporarily removed the furnace is brought to a 

((temperature of about 1100° while oxygen is passed through slowly for 

one-half hour. This preliminary heating to expel traces of moisture and 

organic matter may be dispensed with if the apparatus has been in use. 

Unless the apparatus has been in continuous use one or more blank 
idetermi nations must be made in order to correct for any constant small 
kain of carbon dioxide from imperfect apparatus or reagents. The 
petermination of carbon in steel follows the blank. The different 
modifications, with blanks, will be referred to by the letters that have 
jalready been used in the discussion of the principles of these methods. 

Method (a). — The potassium hydroxide bulbs are carefully wiped and 
the outlets are closed by short pieces of rubber tubing bearing glass 
plugs. They are allowed to stand in the balance case for 10 minutes, 
[after which the rubber tubes are removed and the bulbs are weighed. 


Crucible forceps tipped with rubber should be used for handling the! 
bulbs in this and all subsequent operations. After weighing, the bulbs 
are inserted in the train and oxygen is passed through at the rate oi' 
about three bubbles per second for 15 minutes. At the end of this 
period the gas flow is stopped and the bulbs are removed, plugged and 
placed in the balance case. After 10 minutes the rubber caps are re- 
moved and the bulbs are immediately weighed. If this blank determi- 
nation gives a change of more than 0. 1 mg in the weight of the bulbs it 
must be repeated until the change drops to zero or becomes constant 
within this limit. When this condition is reached the determination of 
carbon may be made. 

While the blank determinations are running, prepare an alundum 
boat by placing a layer of granular alundum (grade RR) on the bottom 
and sides, then igniting for 5 minutes. After this has cooled about 2 gm 
of steel sample is weighed on a counterpoised glass, brushed into the 
boat and distributed in a uniform layer. 

Insert the weighed bulbs into the train and adjust the flow of oxygen 
to the same rate as that used in the blank. Without stopping this 
flow, open the end of the combustion tube next to the oxygen tank and 
insert the boat containing the sample, pushing it by means of a wire into 
the alundum thimble which is next to the platinized asbestos. Close 
the tube as quickly as possible and carefully twist the stopper into the 
end so that no leak can occur. The combustion of the steel begins 
almost immediately and is usually completed within a very short time. 
Now continue the passage of oxygen for 15 minutes, at the end of whi( 
period remove the absorption bulbs, stopper, place in the balanc 
case and weigh, without the rubber caps, after 10 minutes. 

In the meantime another sample of steel should have been made read] 
in a second boat and another absorption bulb should have been weighed. 
As soon as the first bulb is removed the second is inserted in the train. 
The combustion tube is opened as before, the first boat is drawn out 
and the second is inserted. A continuous series of determinations may 
be made in this way without stopping the flow of oxygen or cooling the 
furnace. Weighings may be made while the combustion is proceeding. 
Bulbs may be used without refilling until exhausted, following the rule 
given on page 133. 

From the weights of samples and of carbon dioxide calculate the per- 
cent of carbon in the steel samples. 

Method (b). — The manipulation is exactly like that of method (a). 
The rate of flow of oxygen is judged by bubbling through the potassium 
hydroxide tube following the oxygen tank. 

Method (c). — A saturated solution of barium hydroxide is first pre- 
pared by warming and stirring the solid base with recently boiled water, 


using a ratio of 70 to 100 gm of base to 1000 cc of water according to the 
purity of the barium hydroxide obtainable. Cool to room temperature 
and siphon into a bottle which is then closed with a rubber stopper. 
Dilute 550 cc of this solution to 1000 cc with distilled water, mix and 
place in a bottle which is provided with a guard tube of soda lime and a 
siphon or similar outlet. 

Prepare a solution of hydrochloric acid, 1 cc of which is equivalent 
to 0.002 gm of carbon, standardizing against pure sodium carbonate in 
presence of methyl orange. 

Rinse the Meyer bulbs with boiled water, then measure into them 
from a burette or an automatic pipette attached to the bottle 50 cc of 
the dilute solution of barium hydroxide, first discarding the few drops 
that are in the outlet of the measuring instrument. Add to the bulbs 
from a graduated cylinder enough cold, carbon dioxide-free water to 
bring the liquid just to the lower edge of the upper bulb when the gas 
is flowing. The quantity necessary should be determined, once for all, 
so that it may be added without delay in subsequent determinations. 
With the furnace already heated, connect the bulbs in place while the 
oxygen is flowing and conduct the blank experiment as in method (a), 
to the end of the absorption period, then disconnect the bulbs without 
stopping the flow and rinse the solution into a 250 cc Erlenmeyer flask 
by means of boiled water, paying no attention to the precipitate. Add 
a drop of phenolphthalein and titrate at once (not too rapidly but stir- 
ring vigorously) with standard hydrochloric acid, to the disappearance 
of the pink color. When this point is reached the volume of required 
acid is read. The pink color will return after standing, due to the 
gradual resolution of barium carbonate, but this is not considered in the 

At the time when the absorption bulbs are removed from the train 
of apparatus another tube, charged like the first, is inserted for use in 
the next blank determination, the barium hydroxide solution having 
been measured just before inserting the bulbs. Run this blank deter- 
mination like the first. 

While the gas is flowing for the second blank the steel sample should 
be weighed and placed in the prepared boat, as in method (a). When 
the blank is finished the first set of absorption bulbs, charged as before, 
is substituted and the boat with the steel sample is inserted into the 
combustion tube. Continue the combustion as directed for method 
(a). While this is proceeding the base from the second blank is titrated. 
This titration should agree with the first to within 0.1 cc of standard 

A continuous series of determinations may be made without stopping 
the flow of oxygen or cooling the furnace. Enough blanks must be 


run to obtain agreement of titrations and at least two determinations 
of carbon should be made on each sample of steel. 

The volume of acid required in the determination is subtracted from 
that used in the blanks and the remainder is multiplied by the carbon 
equivalent of the standard acid. From this and the weight of sample 
calculate the percent of carbon in the steel. 

Method (d). — The saturated solution of barium hydroxide whose 
preparation was described for method (c) is used. About 25 cc is placed 
in the Meyer bulbs and carbon dioxide-free water is added so as to fill 
all but the upper bulb when the gas is flowing. The blank experiment 
and the carbon determination are conducted as in method (c), with the 
following exceptions : 

Barium hydroxide need not be accurately measured, although approxi- 
mately equal volumes should be used in all experiments. A graduated 
cylinder is suitable for measuring the solution but it should first be 
rinsed with boiled water. 

At the end of the absorption the solution is filtered rapidly on a paper 
filter and the Meyer tube, the precipitate and the paper are washed 
several times with cold, carbon dioxide-free water. The precipitate 
is then dissolved by adding exactly 25 cc of tenth-normal hydrochloric 
acid (or the solution described for method (c)) the solution being caught 
in the Meyer tube so that all adhering barium carbonate is dissolved. 
The paper is well washed with hot water and the combined filtrate 
and washings rinsed into an Erlenmeyer flask. The excess of acid is 
then titrated by tenth-normal sodium hydroxide, using methyl orange 
as indicator. 

Determination by Combustion, Preceded by Solution. — Prepare 
solution of cupric potassium chloride containing 500 gm of the crysta 
lized salt and 75 cc of concentrated hydrochloric acid in 1000 cc of solu 
tion. Filter through ignited asbestos into the bottle. 

Weigh 1 gm of the steel or iron drillings, place in a 250-cc beaker and 
add 100 cc of the cupric potassium chloride solution. Stir until the 
metal is all dissolved, warming to about 65°. If many determinations 
are to be made a stirring machine is desirable. 

Filter the solution through a Gooch crucible or a carbon tube (shown 
in Fig. 67, page 250). The asbestos used in the filter must have been 
previously ignited to remove all organic matter. Wash with warm (50 ( 
dilute hydrochloric acid until the washings are free from color, then wi 
cold water until free from chlorides. It is desirable that most of i 
water be removed from the filter and carbon, although the latter need n 
be absolutely dry. After partial drying by means of the pump and 
ing oven the asbestos and carbon are carefully removed and placed in 
combustion boat, using a small pair of forceps. This operation sho 



be performed over a sheet of white, glazed paper. The inside of the 
crucible or carbon tube is carefully wiped clean, using a tuft of ignited 

The combustion and subsequent determination are carried out as in 
the direct combustion process except that directly following the combus- 
tion tube there is inserted a U-tube containing a saturated solution of 
ferrous sulphate, acidified by sulphuric acid. If traces of hydrochloric 
acid are retained by the filter or carbon this is partly oxidized, chlorine 
being produced, and partly carried over without change. Chlorine is 
absorbed and reduced by the ferrous sulphate while small quantities 
of hydrochloric acid are absorbed by the water of the solution. 

Free (Graphitic) Carbon. — When steel or iron containing both 
|free and combined carbon is dissolved in nitric acid of specific 
gravity 1.2 the combined carbon passes into solution as hydro- 
carbons, the graphitic carbon being left as an insoluble residue. 
! The latter may be separated by filtration and used for the deter- 
mination of free carbon. If it is to be weighed directly the silica 
[which is also left must be removed by washing with potassium 
I hydroxide solution, then with water. A 'better method is to 
wash the carbon and silica free from iron salts and acids and then 
| determine by combustion, exactly as in the case of total carbon. 

Graphitic carbon may also be determined by difference, sub- 
tracting the combined carbon from the total carbon. 

Determination. — Weigh 1 gm of pig iron or 10 gm of steel and dis- 
solve in nitric acid, specific gravity 1.2, using 15 cc of acid for each 
gram of sample. Filter through ignited asbestos in a Gooch crucible 
or a carbon tube and wash with dilute hydrochloric acid, then with hot 
I water until free from chloride's. Burn in the combustion tube used for 
j total carbon and determine in the same way. 

Combined Carbon. — The percent of combined carbon may be 
[determined indirectly by subtracting the percent of graphite 
I from that of total carbon. The only reliable method for the 
direct determination of combined carbon is that of Eggertz. 1 
I This method depends upon the fact, noted in the discussion of free 
carbon, that when steel or iron containing combined carbon is 
| dissolved in dilute nitric acid the combined carbon forms soluble 
j organic compounds which impart a color to the solution, theinten- 

*Z. anal. Chem., 2, 433 (1862); Chem. News, 7, 254 (1863). 


sity of which varies with the percent of combined carbon. Tl 
solution is then compared in tubes with the solution of a stan( 
ard steel whose carbon content is known, the unknown pei 
cent being then calculated. It will be seen later, when a moi 
extended study of carbon conditions is taken up, that combine< 
carbon may exist in more than one physical state, althouj 
probably always present as the carbide Fe 3 C This different 
in physical state is influenced by the presence of other elemenl 
and also by the mechanical and thermal treatment which the 
steel has received. The color of the acid solution is affected by 
all of these factors and it therefore becomes necessary to use for a 
standard steel one in which not only the percent of combined 
carbon is known to be approximately the same as that of the steel 
being examined but also one that has nearly the same percents of. 
other impurities and that has been subjected to the same thermal 
and mechanical treatment. All of these factors cannot well be 
known in general testing and the method is therefore of little 
value for this class of work. Its chief value is to the steel works 
chemist who knows in every case the nature of the steel with 
which he is dealing and who is thereby enabled to select his 
standard steel with due regard to all of the variable factors. 

Determination. — Treat the standard steel and the steel being ex- 
amined as follows : Weigh 1 gm of the drillings and dissolve in a beaker 
in 30 cc of nitric acid whose specific gravity is 1.2 and which is free from 
chlorine. Warm the acid toward the end of the process, to complete 
the solution. Filter to separate free carbon and silica, receiving the 
filtrate in a 100-cc volumetric flask. Wash the residue, dilute to the 
mark and mix. Transfer 30 cc of the solution of lighter color to an 
Eggertz tube, which is a tube graduated from 1 cc to 30 cc and having 
an internal diameter of about 1 cm. Add the darker solution to another 
similar tube until the color of the two tubes appears to be equal, viewed 
from above. In case the color is very dark, less solution must be used 
or the color observed from one side, or else the color of the darker 
solution is lightened by dilution. 

Calculate the percent of combined carbon. 

Silicon. — Silicon occurs in all steels, generally in quantities 
less than 0.3 percent. Certain silicon steels contain as much as 
20 percent. Cast iron contains as much as 3 percent of silicon. 
Silicon occurs as a silicide which is probably to be represented by 


the formula FeSi, this forming a solid solution with the remainder 
bf the iron. Inclusions of slag also may contain silicon as silicates 
of iron and manganese. Silicon has little effect upon the me- 
chanical properties of steel but is desired in cast iron because of 
its tendency toward throwing carbon out of its combination with 
iron, thus forming gray iron which has a greater fluidity when 
melted than does white iron, and which is therefore better suited 
for foundry purposes. 

When iron silicide is dissolved in nitric acid the silicon is 
entirely converted into silicon dioxide, largely in the state of 
colloidal silicic acid. If the silicic acid is dehydrated and the 
resulting silicon dioxide made insoluble by heating with acids it 
may be separated by filtration. As obtained from pig iron the 
silicon dioxide so obtained contains all of the free carbon of the 
iron. This is removed by ignition. In the adaptation of this 
>rocess to the quantitative determination of silicon in iron and 
teel the chief difficulties encountered are due to the tendency of 
ilica to change from the gel to the sol and also to incomplete 
washing of the silica. In order to assist in the separation of 
silica in an insoluble condition Drown suggested 1 the addition of 
ulphuric acid to the solution during the evaporation to render 
ilica insoluble. This materially shortens the time required for a 
letermination as otherwise the solution must be evaporated and 
Leated for some time in order to completely separate the silica. 
During the washing of the silica, if pure water is used, iron salts 
lydrolyze and insoluble basic salts are retained by the filter. 
Alternate washing with water and hydrochloric acid will remove 
,11 but traces of iron salts and a correction may be made for these 
>y the common process of volatilization of silica by hydrofluoric 

Determination. — Prepare a mixture of 375 cc of concentrated nitric 
cid, 125 cc of concentrated sulphuric acid and 500 cc of water or as 
iuch of this mixture as is needed. Weigh 1 gm of pig iron or 5 gm of 
3teel and dissolve by warming in a casserole or platinum dish with 25 cc 
of the acid mixture. The solution is evaporated by agitation of the un- 
covered casserole over a flame until pronounced fumes of the sulphuric 
icid appear. Allow the solution to cool, then add 10 cc of dilute hydro- 
chloric acid and 50 cc-of water. Warm until iron salts are dissolved 

1 Trans. Am. Inst. Min. Eng., 7, 346 (1879). 


then filter and wash alternately with hot dilute hydrochloric acid an< 
water until free from iron and chlorides. Ignite in a platinum crucibl< 
cool and weigh. Volatilize the silica by treatment with sulphuric acic 
and hydrofluoric acid (see page 293), and from the loss in weight calcu 
late the percent of the element silicon. 

Sulphur. — Sulphur occurs in iron and steel as ferrous sulphide 
FeS, unless manganese is present, in which case it forms mangan 
ous sulphide MnS. Ferrous sulphide is itself brittle. It als 
shows a tendency toward the formation of envelopes surroundin 
the crystalline grains of steel, reducing their cohesion and result 
ing in " shortness," particularly when hot. Sulphur is therefor 
said to cause "red shortness" of steel. Manganese sulphid 
usually occurs as small rounded masses instead of envelop) 
and it is therefore much less objectionable than ferrous sulphid 
The steel maker therefore relies upon manganese to correi 
largely the bad effects of sulphur, although the latter should nc 
be present in steel in quantities greater than 0.05 percent. In 
the best steel its quantity is much less than this. 

The determination of sulphur may be accomplished by oxida 
tion to sulphuric acid, followed by precipitation as bariu 
sulphate, or by evolution methods; in the latter the metal 
dissolved in hydrochloric acid, ferrous sulphide or manganous su 
phide forming hydrogen sulphide. The latter is distilled int 
some absorbing solution and subsequently determined by gravi 
metric or volumetric methods. 

Oxidation Method. — Steel or iron dissolves more readily 
dilute nitric acid than in the concentrated acid and the former 
therefore used for dissolving the sample for nearly all other dete 
minations. But the dilute acid will not serve for dissolving t 
metal as a preliminary to the gravimetric determination 
sulphur because a part of the sulphur will be evolved as hydrogen 
sulphide and will then escape. Concentrated nitric acid co 
pletely oxidizes the sulphide to sulphate. 

6FeS+24HN03-^2Fe 2 (S0 4 )3+2Fe(N0 3 )3+18NO+12H 2 0. 

The sulphur is then precipitated as barium sulphate. 

The separation from the large amount of iron involves so] 
difficulty. Unless a considerable excess of acid is present basi 
ferric salts, products of hydrolysis, are retained by the precipital 


pf barium sulphate. If too much acid is present the precipitation 
pi barium sulphate is incomplete. For the solubility of barium 
(sulphate in hydrochloric acid, see page 92. Nitric acid must not 
be present at all because of its effect upon the occlusion of iron 
salts by the precipitate. 

Silica must be separated by evaporation and nitration before 
(the precipitation of barium sulphate, and during the evaporation 
iind heating that are necessary for this purpose there is danger of 
loss of sulphur through decomposition of ferric sulphate: 

Fe 2 (S04)3^Fe 2 3 +3S03. 

lln order to prevent loss of sulphur trioxide by this means, a 
small amount of sodium carbonate is added before the evapora- 
tion. This immediately forms sodium nitrate or chloride (hy- 
jdro chloric acid also having been added) and this reacts during the 
evaporation, thus: 

Fe 2 (S0 4 )3+6NaCl-^2FeCl3+3Na 2 S04. 
Imodium sulphate is not decomposed by moderate heating. 

| Determination. — Weigh 5 gm of drillings or powder into a casserole, 
^lace under a hood and add 50 cc of concentrated nitric acid which is 
free from sulphuric acid. Action may not begin at once unless the cas- 
fcerole is warmed but after the metal begins to dissolve the action may 
become violent. In this case the casserole should be placed in cold 
water. In the later stages it may again be necessary to heat the cas- 
serole. Add 1 gm of sodium carbonate, free from sulphate, and evapo- 
rate to dryness, holding the casserole over the flame and giving it a 
jrotary motion to prevent bumping and to hasten evaporation. When 
the residue is dry, heat for 15 minutes at a temperature just below red- 
ness, then add 30 cc of concentrated hydrochloric acid, and again 
evaporate to dryness and heat as before. Cool, add 30 cc of concen- 
trated hydrochloric acid, warm until all iron salts are in solution, then 
Evaporate in the same manner as before until ferric chloride begins to 
brystallize. Add a very small amount of hydrochloric acid to redissolve 
tfhese crystals, then add 25 cc of water, filter and wash with water, 
r.hen with a very small amount of hot, dilute hydrochloric acid, repeating 
;he water and acid washing until the paper, silica and carbon are free 
I jrom the red or brown stains of ferric chloride. Finally_wash with hot 
j-vater until the volume of the filtrate is about 200 cc. If fc this residue 
is large in quantity it will contain an appreciable amount of sulphur.. 



In this case transfer the paper with the residue to a platinum crucibL 
burn until paper and carbon have been removed and fuse with 2 gm o 
sodium carbonate. Cool, dissolve the fusion in dilute hydrochlori 
acid, using no more than is necessary, and wash into the main solution 

Heat to boiling and add, a drop at a time and stirring continuously 
10 cc of 10 percent barium chloride solution. Digest at a temperatur 
near the boiling point for 30 minutes, then allow to stand for 2 hours 
Filter and wash, alternately with dilute hydrochloric acid and watei 
until the filter and precipitate are white and finally with water unt: 
free from chlorides. In this washing use as little hydrochloric acid* a 
possible. Ignite the paper and precipitate in a platinum crucible at i 
low temperature and weigh the barium sulphate. If the ignited precipi- 
tate is not white some iron oxide is contained in it. In this case add 1 gm 
of sodium carbonate, fuse, dissolve the fusion in water and dilute hydro- ; 
chloric acid and precipitate as before. 

Calculate the percent of sulphur in the sample. 

Evolution Method. — The determination of sulphur by evolu-, 
tion depends upon the decomposition of metallic sulphides by 
hydrochloric acid, the resulting hydrogen sulphide being distillec 
and absorbed in another solution. The absorbing solution ma] 
form an insoluble' sulphide with the hydrogen sulphide or it ma] 
oxidize the latter to sulphuric acid which is then determine 
gravimetrically. Absorbents of the first class are basic solutioi 
of salts of lead, cadmium or silver. Absorbents of the secon( 
class are bromine in hydrochloric acid, potassium permanganat 
and hydrogen peroxide. A solution of cadmium chloride in exces 
of ammonium hydroxide is to be preferred. The precipitate 
cadmium sulphide may be washed, dried and weighed, but it k 
better to decompose it with hydrochloric acid and titrate th< 
liberated hydrogen sulphide with standard iodine solution. Th< 
solubility of cadmium sulphide in hydrochloric acid is not large 
unless the resultant hydrogen sulphide is removed, as in this cas 
by oxidation. 

The evolution method may be performed in less time than tl 
oxidation method. It has been shown by Phillips to be inaccu- 
rate, 1 however, for white pig iron because of the formation 
organic sulphur compounds, of which methyl sulphide, (CH 3 ) 2 £ 
was isolated. Such sulphides are difficult to expel from th< 

1 J. Am. Chem. Soc, 17, 891 (1895). 


bvolution flask and require as much as two hours of boiling, 
during which time air, carbon dioxide or hydrogen is drawn 
through the apparatus. Phillips found that the organic sulphides 
pould be decomposed by passing the vapors through a tube, 
heated to redness. The additional time necessary for the ex- 
pulsion of organic sulphides from the evolution flask makes the 
nethod impracticable for white pig iron. For steel and gray 
big iron, containing relatively low percents of combined carbon, 
the method is satisfactory. 

Determination. — Prepare an approximately tenth-normal solution 
pf iodine by dissolving the calculated weight of iodine and twice its 
k'eight of potassium iodide or sodium iodide in 1000 cc of water. Stand- 
ardize just before making the sulphur determination by titrating 
tgainst a standard solution of sodium thiosulphate or of sodium arsenite 
|As 2 3 and half its weight of NaOH). The iodine dissolves rather 
lowly unless well powdered. It is well to decant the solution into 
jinother bottle in order to avoid the possibility of particles of undissolved 
pdine changing the concentration of the solution after standardization. 
J Determination. — Use a 300-cc flask, having a round bottom, for 
he evolution flask. Connect, through a 2-hole rubber stopper, a 100- 
ijc separatory funnel and a short tube, bent at a right angle with the 
Blask. The separatory funnel should reach to the bottom of the flask 
|nd should have the bottom turned up, as in the apparatus for the 
determination of carbon dioxide in carbonates. The short exit tube is 
pnnected with another tube which reaches to the bottom of a cylinder 
[laving a capacity of about 100 cc, in which is placed 50 cc of a solution 
made as follows: Cadmium chloride 5 gm, water 375 cc, concentrated 
Lmmonium hydroxide 625 cc. This cylinder is similarly connected 
■vith. a second cylinder containing the same kind of solution. Both 
I ylinders should stand in a large beaker of cold water. Instead of the 
!|wo cylinders a Meyer tube (figure 96) may be used. 
I Weigh into the evolution flask 5 to 10 gm of steel or iron drillings, close 
pe flask and place 75 cc of hydrochloric acid (1:1) in the separatory 
Mann el. Admit the acid fast enough to cause a rapid evolution of 
Hydrogen. Finally add all but about 5 cc of the acid and warm to assist 
he solution. When the metal is all dissolved boil for 5 minutes at 
rate that will permit absorption of the hydrogen sulphide. This 
oiling should completely expel hydrogen sulphide and hydrogen from 
e flask. 

Disconnect the delivery tube, then remove the source of heat and rinse 
e tubes, allowing the washings to run into the absorption cylinders. 


If the tubes contain any cadmium sulphide wash with dilute hydro chloi 
acid and then with water but do not agitate the solution. Rinse t 
contents of the absorption cylinder into a 500-cc beaker, dissolvi] 
adhering precipitate by means of dilute hydrochloric acid, allowing th 
solution to run immediately into the main body of solution. Usually 
of the hydrogen sulphide is absorbed in the first cylinder and the conten 
of the second need not be used if no trace of yellow cadmium sulphi 
appears in it. Add water until the volume is about 300 cc, then a< 
dilute hydrochloric acid until the liquid is distinctly acid in charact< 
stirring gently meanwhile. The disappearance of turbidity is sufficie 
indication of an acid reaction. Rapid stirring. and undue agitation wi 
cause a loss of hydrogen sulphide. Add 1 cc of starch solution and ti- 
trate at once with decinormal iodine solution. 
Calculate the percent of sulphur in the sample. 

Phosphorus. — The proportion of phosphorus in steel oj 
satisfactory quality is not usually higher than 0.1 percent anc 
is frequently required to be less. Acid open hearth and ack 
Bessemer steel contain larger quantities of phosphorus th* 
steel made by basic processes. Phosphorus occurs in steel 
the phosphide Fe 3 P. Its effect is to cause brittleness of tl 
steel, this being at least partly due to the promotion of coan 

The determination of phosphorus in iron or steel may follow 
either gravimetric or volumetric methods. In any case the fins 
determination must be preceded by separation from the relativel 
large excess of iron. The separation is usually made by eithe 
a modification of the method of Fresenius 1 known as the "acetal 
method/' or the molybdate method of Sonnenschein. 2 

Acetate Method. — This method of separating iron and phos 
phorus depends upon the relatively large solubility of ferroi 
acetate as compared with that of basic ferric acetate and fei 
phosphate. The iron is first reduced entirely to the ferroi 
condition by sulphurous acid, then either a small amount reo3 
dized by bromine or a small amount of ferric chloride is adde( 
The solution is now made slightly basic, then an excess of acel 
acid and water is added'. A precipitate forms, consisting 
ferric phosphate and basic ferric acetate, the latter being presei 

1 J. prakt. Chem., 45, 258 (1848). 

2 Ibid., 63, 339 (1851). 


I in very small quantity. The larger part of the iron has remained 
I in solution as ferrous acetate and is separated by filtration. 

This method necessarily leaves a small amount of iron in the 
i phosphorus precipitate. In order to separate this, advantage 
| is taken of the fact that small quantities of iron are not -pre- 
cipitated by ammonium hydroxide if organic acids are present. 
Either citric acid or ammonium citrate is added and the phos- 
phorus is precipitated by " magnesia mixture" in presence of 
; ammonium hydroxide. The ionization of ferric citrate is so 
small that the solubility product of neither ferric hydroxide nor 
| basic ferric citrate is attained. 

The acetate method is accurate if carefully performed, but is 
1 complicated in detail and is more liable to fail than the next 
method to be described. 

Molybdate Method.- — The molybdate method of separating 
j iron and phosphorus depends upon the insolubility of ammonium 
| phosphomolybdate and the solubility of iron in nitric acid. The 
iron or steel is dissolved in nitric acid, carbon is oxidized by potas- 
I sium permanganate, the solution is nearly neutralized and a solu- 
| tion of ammonium molybdate in nitric acid is added. The pre- 
cipitate of ammonium phosphomolybdate is separated by 
I nitration and is then treated according to the method which has 
| been selected for the final determination. The removal of 
I carbon by oxidation is necessary in order that precipitation shall 
I be complete. 

The determination of phosphorus may now be made (1) by 
; drying and weighing the yellow precipitate of ammonium phos- 
Ipho molybdate, (2) by measuring its volume, (3) by titrating its 
imolybdic oxide by means of a standard base, (4) by reducing 
!its molybdic oxide to molybdenum sesquioxide and titrating by 
! standard potassium permanganate solution, or (5) by dissolving 
(the yellow precipitate in ammonium hydroxide and precipitating 
as magnesium ammonium phosphate. 

If method (1) is to be followed it is necessary that care be 
exercised in precipitating the ammonium phosphomolybdate in 
order that its composition may be constant. Precipitated under 
the conditions later described its composition is represented by 
the formula (NH 4 ) 3 P04.12Mo03. The composition is somewhat 
altered by variation in temperature, excess of ammonium 


molybdate, excess of nitric acid and time of precipitation. Th 
precipitate may contain also small amounts of free molybdic acic 
especially if too much nitric acid is present, or of ammonium 
silicomolybdate if silicon has not been removed. This metho 
of direct weighing is not often followed. 

Method (2) is a rapid but inaccurate method. The precipita 
tion is carried out in a pear-shaped bulb having a graduated stem 
The precipitate is packed into the stem by centrifugal actioi 
and its volume is read and converted into weight percent by 
previously determined factor. 

Method (3) was suggested by Pemberton. 1 In this method the 
yellow precipitate is dissolved in an excess of a standard solution 
of potassium hydroxide or sodium hydroxide, the excess being 
then titrated by a standard acid solution, phenolphthalein being 
used as the indicator. The reaction between the phosphomolyb- 
date and the base is as follows: 

(NH 4 )3P04.12Mo03+24KOH->(NH4)3P04+12K2Mo0 4 + 

12H 2 

Upon the addition of standard acid, phenolphthalein change 
color when the excess of base has been neutralized and th 
following reaction has occurred: 

(NH 4 )3P04+HC1-+(NH4) 2 HP04+NH 4 C1. 

Twenty-three equivalents of base have therefore apparently bet 
used at the end point and in order to express this fact in one equs 
tion the reaction is often represented as follows : 
2(NH4) 3 P04.12Mo03 + 46KOH-^2(NH4) 2 HP04+(NH 4 )2MoO, 


This is seen to be really a direct titration of molybdic acid insteac 
of a titration of phosphorus and it is therefore an indirect estima- 
tion of phosphorus and can be correct only in case the compositior 
of the precipitate is constant. It is also essential that no in 
molybdic acid should be present with the phosphomolybdate. 

There is some difference in opinion concerning the accura< 
of this method. If the precipitation is carefully performed it 
probably as accurate as the gravimetric method (5). 

1 J. Chem. Soc, 15, 382 (1893); 16, 278 (1894). 


Determination by Pemberton's Method. — Prepare the following 
reagents : 

(a) Acid Solution of Ammonium Molybdate. — Dissolve 100 gm of 
molybdic acid in a mixture of 144 cc of ammonium hydroxide (specific 
gravity 0.90) and 271 cc of water. Pour this solution, slowly and with 
vigorous stirring, into a mixture of 590 cc of concentrated nitric acid 
(specific gravity 1.42) and 1148 cc of water. Allow to stand at a tem- 
perature of about 40° for several days and then decant from sediment 
and preserve in glass-stoppered bottles. 

(6) Standard Potassium Hydroxide Solution, 1 cc of which is equiva- 

llent to 0.1 mg of phosphorus. This should be as nearly free from 
carbonates as possible and is made as follows : Dissolve 2 percent more 
than the calculated quantity for 1000 cc, dilute to 100 cc and add 1 cc 
of a saturated solution of barium hydroxide. Stopper the flask and 

'allow to stand until the precipitate of barium carbonate has settled. 
Decant and dilute to 1000 cc. Standardize by titration against solution 
(c), using phenolphthalein. Adjust so that 1 cc is equivalent to 0.1 

j mg of phosphorus. 

(c) Standard Hydrochloric Acid Solution, equivalent in concentra- 
tion to the standard base ; use boiled water. 

(d) Potassium Permanganate Solution, 1.5 gm in 100 cc. 

(e) Potassium Nitrate Solution, 1.0 percent. 

Weigh 2 gm of iron or steel into a 250-cc Erlenmeyer flask and add 
100 cc of nitric acid (specific gravity 1.13) and warm until the sample 
is dissolved (see note on page 457: "Interference of Titanium")- Boil 
to expel oxides of nitrogen, then add 10 cc of solution (d) and boil 
until the combined carbon is completely oxidized and the excess of 
potassium permanganate is decomposed, as is made evident by the 
disappearance of the pink color. Dissolve the precipitated manganese 
dioxide by warming with about 1 gm of ferrous ammonium sulphate. 
[Add ammonium hydroxide with vigorous stirring. The last part of 
this operation must be conducted with care because if much ferric 
hydroxide is allowed to form it will not readily redissolve, even though 
(the solution still contains an excess of acid. Redissolve the precipi- 
jtate by the addition of the least necessary quantity of nitric acid. 
[Place a thermometer in the flask and warm to a temperature of 60° to 
po° by placing the flask in a water bath and then add 40 cc of freshly 
[filtered ammonium molybdate solution, stir well and allow to stand 
|15 minutes. Filter immediately and wash flask and precipitate with 
Solution (e) until the washings are neutral to phenolphthalein but with- 
out attempting to remove all precipitate from the flask. 

Transfer the paper and precipitate to the flask in which the precipita- 
tion was made and add enough standard solution of potassium hydro x- 


ide to dissolve the precipitate. Dilute to about 75 cc with recently 
boiled water, add a drop of phenolphthalein and titrate the excess of 
base with standard acid solution. 

Calculate the percent of phosphorus in the steel. 

Method (4). This also is an indirect method for the determina- 
tion of phosphorus, since it also depends upon reactions of molyl 
denum oxides, rather than of phosphorus. The precipitate oi 
ammonium phosphomolybdate, obtained as in method (3), 
dissolved in ammonium hydroxide, the solution is acidified wit) 
sulphuric acid and zinc is then added. Molybdenum trioxide 
M0O3, is reduced to molybdenum sesquioxide Mo 2 3 , which il 
again oxidized by titration with standard potassium perman- 
ganate solution. 

Method (5). This is one of the most reliable of all methods if 
carefully performed, since its accuracy does not depend, in any 
way, upon the composition of the yellow precipitate. A some 
what less acid solution of ammonium molybdate may be used an 
this keeps better than the solution required for volumetric proc 
esses. The yellow precipitate of ammonium phosphomolybdat 
is dissolved in ammonium hydroxide and the phosphorus is the 
precipitated as ammonium magnesium phosphate by the additio 
of a solution of magnesium chloride. The ammonium mag 
nesium phosphate is ignited and weighed as magnesium pyro 
phosphate. Potassium permanganate cannot be used for th 
oxidation of carbon since it would later form a precipitate 
ammonium manganese phosphate. 

Determination. — Prepare the following solutions: 

(a) Acid Solution of Ammonium Molybdate. — Dissolve 100 gm of 
molybdic acid in a mixture of 144 cc of 'ammonium hydroxide (specific 
gravity 0.90) and 271 cc of water. Pour this solution slowly and with 
vigorous stirring into a mixture of 490 cc of concentrated nitric acid 
(specific gravity 1.42) and 1148 cc of water. Allow to stand at a tem- 
perature of 30° to 40° for several days and then decant and preserve 
in glass-stoppered bottles. 

(b) Ammonium Citrate Solution. — Dissolve 50 gm of citric acid in 
water, add 350 cc of ammonium hydroxide (specific gravity 0.90)_and 
dilute to 1000 cc. 

(c) Ammonium Hydroxide Solution containing 2.5 percent of ammonia. 

(d) "Magnesia Mixture" — Dissolve 55 gm of crystallized magne- 


sium chloride and 140 gm of ammonium chloride in water, add 130 cc 
of ammonium hydroxide (specific gravity 0.90) and dilute to 1000 cc. 

(e) Ammonium Nitrate Solution, 10 percent. 

Dissolve 1 to 2 gm of steel in 20 cc of nitric acid (specific gravity 
1.2) in a casserole, cover and boil until nitrogen oxides are expelled. 
Evaporate to dryness on the steam bath or by agitating over a flame. 
Heat for 15 minutes over the direct flame in order to oxidize organic 
matter, formed from combined carbon. Cool, add 30 cc of concentrated 
hydrochloric acid and heat to dissolve iron oxide. Evaporate with 
stirring until ferric chloride begins to crystallize but do not allow 
salts to dry on the sides of the casserole. Add 10 cc of concentrated 
nitric acid, boil to expel chlorine, dilute to 75 cc and filter into a 250-cc 
flask. (If titanium is present, see note below: "Interference of Tita- 
nium.") Wash the silica and carbon on the paper with 2 per cent nitric 
acid and water until the iron is all removed, as made evident by the 
disappearance of brown stains. 

Dilute the filtrate to about 100 cc and add dilute ammonium hydrox- 
ide solution very slowly and with vigorous stirring until a small amount 
of precipitate remains undissolved. Redissolve this in concentrated 
nitric acid, immerse the flask in water and warm to about 60°. Add 50 
cc of ammonium molybdate solution (a), shake and allow to remain at a 
temperature of 65° for an hour. Filter and wash with solution (e) 
until no brown stains remain on the paper. It is not necessary to re- 
move all of the precipitate from the sides of the flask at this point, but it 
must be well washed. Place the flask in which precipitation was made 
under the funnel and dissolve the precipitate by adding about 25 cc of 
ammonium citrate solution (6). 

Wash the paper thoroughly with hot water. Rotate the flask until 
all of the precipitate is dissolved from the sides then nearly neutralize 
with hydrochloric acid. Transfer to a beaker, dilute to 100 cc, add 
10 cc of magnesia mixture, slowly and with vigorous stirring. After 
the solution has stood for 30 minutes add slowly ammonium hydroxide 
of specific gravity 0.90 in quantity equal to 1/9 of the total volume 
of solution. Allow to stand for two hours, filter and wash with dilute 
ammonium hydroxide solution. Ignite in a platinum or porcelain 
crucible until white and weigh the magnesium pyrophosphate. 

Calculate the percent of phosphorus in the steel. 

Interference of Titanium.- — If titanium is present, as it fre- 
quently is in pig iron and sometimes in steel, the phosphorus 
will not all be recovered by any of the methods alreadyr descibed 
because the action of acids upon iron leaves an insoluble double 


salt of phosphoric acid, titanic acid and iron. In this case the 
residue of silica, carbon, ferric phosphotitanate, etc., obtained by- 
filtration of the acid solution of iron, is ignited in a platinum 
crucible to burn organic matter, the silica is volatilized by mois- 
tening with a drop of sulphuric acid and adding 1 cc of hydrc 
fluoric acid, and the residue is then fused with about 2 gm 
sodium carbonate. Sodium phosphate, ferric oxide and sodim 
titanate, Na 2 Ti0 3 , are formed. Sodium phosphate is dissolve 
in water and added to the principal solution of the iron in nitri 
acid. Sodium titanate and ferric oxide are insoluble in watei 
Titanium. — Titanium is often present in pig iron as an impurity, 
being derived from the iron ores. There is now an increasing 
use of titanium, in the form of iron-titanium " alloys," as an 
agent to promote sound castings and sound steel. Its effect is 
to reduce oxides, combine with nitrogen and sulphur and thus to 
prevent blow holes and flaws by the formation of a solid oxide or 
nitride which enters the slag. If this action were ideal there 
should be no titanium remaining in the metal at the end of the 
process but this is not always the case and determinations of 
titanium may be required. Titanium is now also used to some 
extent as an essential constituent of finished alloy steels. 

It has already been stated that much of the titanium wil 
remain as an insoluble compound with iron and phosphon 
when iron is dissolved in acids. This is freed from carbon b] 
ignition and from silica by treatment with sulphuric acid an( 
hydrofluoric acid. The titanium in the acid solution of the sai 
pie is recovered by neutralizing the excess of acid, reducing the 
iron to the ferrous state by sodium thiosulphate or sulphurous 
acid and precipitating titanic acid by boiling. Titanic acid is 
an irreversible colloid (see page 19) and becomes insoluble 
when its solution is boiled for some time. This precipitate is 
removed by filtration and added to the residue already in the 
crucible. The whole is then fused with sodium carbonate and 
the sodium acid titanate, NaHTi0 3 , and ferric oxide are separated 
from sodium phosphate by dissolving the latter and filtering. 
One of two methods of procedure may now be adopted: 

(a) The insoluble sodium acid titanate is fused with potassium 
pyrosulphate, forming titanic acid. Sulphuric acid and water 
are added, the titanic acid forming a colloidal sol. The iron also 


passes into solution as ferric sulphate. This is reduced to the 
ferrous condition by sulphurous acid or ammonium acid sulphite, 
the solution is largely diluted and boiled, when the sol is floccu- 
lated, titanic acid again passing into the irreversible gel. This is 
separated by filtration, washed and ignited and weighed as tita- 
nium dioxide. 

(6) The sodium titanate in the crucible is dissolved in hot, 
dilute sulphuric acid, transferred to a color comparison tube (a 
Nessler cylinder or similar tube) and treated with hydrogen 
peroxide. Titanium is oxidized by hydrogen peroxide to the 
hexavalent condition and forms an intensely yellow solution. 
The color is compared with that produced by a standard titan- 
ium solution in a similar tube. 

Determination. — Weigh from 2 to 5 gm of iron or steel and place 
in a casserole. Add 50 cc of concentrated hydrochloric acid, cover and 
warm until the metal is dissolved. Filter, wash twice with hot water, 
transfer to a platinum crucible and burn the paper and all carbon. Add 
a drop of sulphuric acid and about 3 cc of hydrofluoric acid and finally 
heat to expel the acids and silicon tetrafluoride. 

To the filtrate containing most of the iron add dilute ammonium 
hydroxide, slowly and with continuous stirring, until a small amount 
of ferric hydroxide remains undissolved. Redissolve this in hydro- 
chloric acid, leaving the solution with a small excess of acid. Add a 
20-percent solution of sodium thiosulphate until the red color of ferric 
chloride disappears and sulphur begins to precipitate. Dilute to about 
400 cc, add 20 gm of sodium acetate and 50 cc of 30-percent acetic acid 
and boil for 15 minutes or until precipitation of titanic acid seems to 
be complete. Filter and wash two or three times with hot, 1-percent 
acetic acid and place the paper and precipitate in the crucible contain- 
ing the main portion of titanium. Burn the paper and carbon then 
add about 5 gm of sodium carbonate and thoroughly fuse. Cool, 
place the crucible in a beaker and cover with hot water. When the 
fusion is entirely disintegrated, filter and wash the sodium titanate and 
iron oxide with 1-percent sodium carbonate solution. Proceed by 
method (a) or (6), below. 

(a) Gravimetric Method. 1 — Return the paper containing the washed 
residue of sodium titanate to the crucible in which the fusion was made. 
Add 10 gm of potassium pyrosulphate and heat gently, avoiding loss by 
I effervescence. Gradually raise the temperature until the crucible is 

1 Blair: The Chemical Analysis of Iron, 8th ed.,^172. 




finally red and keep at this temperature until all the iron oxide is dis- 
solved. Cool, add 15 cc of concentrated sulphuric acid and heat until 
the entire contents of the crucible have becomo liquid. Cool and pour, 
slowly and with stirring, into 400 cc of water contained in a 500-cc 
beaker. If basic ferric salts precipitate, redissolve in hydrochloric 
acid. Add 50 cc of a 20-percent solution of sodium thiosulphate. Filter 
if not clear, nearly neutralize with ammonium hydroxide, redissolve 
any precipitate that may have formed and add a clear solution contain- 
ing 20 gm of sodium acetate and 150 cc of 30-percent acetic acid. Boil 
and filter the titanic acid. Wash three times with 1-percent acetic 
acid, transfer the paper and precipitate to a porcelain or platinu 
crucible and burn the carbon, finally igniting for five minutes over th 
blast lamp. Weigh the titanic oxide, Ti0 2 , and calculate the percent 
of titanium. 

(b) Colorimetric Method. — Dissolve the residue in the crucible by 
heating with dilute sulphuric acid, place the filter paper in a beaker 
and pour the sulphuric acid upon it. Heat until the sodium titanate i 
dissolved then remove the paper, rinse thoroughly and rinse the conten 
of the crucible into the beaker. Transfer to a Nessler tube, filtering 
not clear, and dilute to the mark. Prepare a standard solution < 
titanic acid as follows: Ignite 1 gm of the purest obtainable titan 
acid in a weighed platinum crucible, cool and weigh. Dissolve th 
titanic acid in dilute sulphuric acid, rinse into a 1000-cc volumetri 
flask and dilute to the mark. Mix well, transfer to a dry glass-stop 
pered bottle and record the concentration of the solution. Into fou 
similar tubes, having the same capacitj^ and the mark at the sam 
height as the first tube, containing the titanium from the sample, place 
1 cc, 3 cc, 5 cc and 10 cc, respectively, of the standard titanium selu- 
tion and dilute these to the mark. Add to each of the five tubes so 
prepared 5 cc of hydrogen peroxide. Compare the color of the tube 
containing the sample with that of the four tubes of standard, looking 
vertically downward through the tubes toward a white surface, placed 
near a window. As a result of these comparisons, limits will be found 
for the concentration of the sample tube. Prepare other tubes