WORKS OF DR. H. W. SCHIMPF

PUBLISHED BY

JOHN WILEY & SONS, Inc.

Essentials of Volumetric Analysis.

An introduction to the subject adapted to the needs of Students of Pharmaceutical Chemistry. Third Edition, rewritten. xiii + 366 pages, Si by 8. 6 1 figures. Cloth, $2.50 net.

A Systematic Course of Qualitative Chemical Analysis of Inorganic and Organic Sub- stances.

With Explanatory Notes. Third Edition, ix + 187 pages. 6 by 9. Cloth, $1.75.

ESSENTIALS

OF

VOLUMETRIC ANALYSIS

AN INTRODUCTION TO THE SUBJECT, ADAPTED TO THE NEEDS OF STUDENTS OF PHAR- MACEUTICAL CHEMISTRY

EMBRACING THE SUBJECTS OF ALKALIMETRY, ACIDIMETRY, PRECIPI- TATION ANALYSIS, OXIDIMETRY, INDIRECT OXIDATION, IODOM- ETRY, ASSAY PROCESSES FOR DRUGS, ESTIMATION OF AL- KALOIDS, PHENOL, SUGARS, THEORY, APPLICA- TION AND DESCRIPTION OF INDICATORS

BY

HENRY W. SCHIMPF, Pn.G., M.D.

Professor of Analytical Chemistry in the Brooklyn College of Pharmacy

-y

. THIRD EDITION— REWRITTEN AND ENLARGED TOTAL ISSUE, SEVEN THOUSAND

NEW YORK

JOHN WILEY & SONS, INC. LONDON: CHAPMAN & HALL, LIMITED

COPYRIGHT, 1903, 19 n, 1917

BY HENRY W. SCHIMPF

OF

. BRAUNWORTH & CO.

4/22 BOOK MANUFACTURERS

BROOKLYN, N. Y.

PREFACE TO THE THIRD EDITION

THE exhaustion of the second edition and the appearance of the new United States Pharmacopoeia as well as the increas- ing demand for the book have necessitated a complete revision of " Essentials of Volumetric Analysis," and its issuance as a third edition.

The book has been considerably improved by the introduc- tion of many new assay processes, among which may be men- tioned the assays of mercuric salts, phosphates and hypophos- phites by means of standard sulphocyanate solution; assays of chlorates, perborates, chloral, acetone, resorcinol, phenyl- sulphonates, arsenates, and alkali cacodylate.

The " c.c." has been replaced by " mil " and in other re- spects changes are made to accord with the new Pharmacopoeia.

HENRY W. SCHIMPF.

NEW YORK CITY, October, 1917.

iii

PREFACE TO THE FIRST EDITION

TEE growing need for a short text-book which will make the principles of volumetric analysis readily available without going too deeply into detailed and discursive description has led to the preparation of this elementary treatise.

In the following pages the aim is to present the principles of this interesting and important subject in a form readily intelligible to students and available for lecture-room and laboratory work. The essential points are condensed within the limits of a small book with the intention of furnishing an outline which may serve as a practical guide as well as an introduction to the more advanced and voluminous works on the subject.

If presented in a suitable manner volumetric analysis rarely fails to prove interesting to the student, because it gives him a clear conception of the quantitative significance of chemical equations and thus affords practical proofs of chemical laws; it furthermore trains the student to make careful obser- vations, to form habits of accuracy in manipulation, and since the processes are easily carried out, enables him to arrive readily at a definite numerical conclusion.

The subject-matter in this book is systematically arranged as far as can be, and treated as concisely as is consistent with clearness of expression. The processes are grouped under five headings: Neutralization, Precipitation, Oxidation, Indirect Oxidation, and lodometrv. The principles underlying each

y

vi PREFACE

group are definitely indicated, and their application illustrated by numerous practical examples. Other subjects treated include methods of calibration and of the accurate reading of graduated instruments, the calculation of the results of analyses, the preparation and standardization of volumetric solutions. The indicators, their selection for special cases and the ionic theory regarding their action, as well as assay process for phenol, sugars and vegetable drugs also receive special treatment. The author hopes that he has prepared a book which will prove serviceable to those for whom it was written and that it will be as generously received as were the four editions of his Text-book of Volumetric Analysis.

HENRY W. SCHIMPF.

CONTENTS

PAGE

LIST OF ELEMENTS WITH THEIR ATOMIC WEIGHTS xii

TABLE OF MULTIPLES OF ATOMIC WEIGHTS AND COMBINATIONS. . . xiii

CHAPTER I INTRODUCTION i

CHAPTER II

GENERAL PRINCIPALS OF CHEMICAL COMBINATION UPON WHICH VOLUMETRIC ANALYSIS is BASED 4

CHAPTER III

VOLUMETRIC OR STANDARD SOLUTIONS 7

To Titrate. Residual Titration.

CHAPTER IV

INDICATORS 17

The lonization Theory. The lonization Theory of Indicators. A Guide for the Selection of Indicators.

CHAPTER V APPARATUS USED IN VOLUMETRIC ANALYSIS 29

CHAPTER VI

ON THE USE OF APPARATUS 40

On the Reading of Instruments. Calibration of Instruments.

vii

Viii CONTENTS

CHAPTER VII

PAGE

METHODS OF CALCULATING RESULTS 50

Table of Normal Factors, etc., of Alkalies, Acids and Alkali Earths. On Stating Results. Table of Molecular Weights and Normal Factors for the most Common Oxids.

CHAPTER VIII

ANALYSIS BY NEUTRALIZATION 58

Alkalimetry. Preparation of Standard Acid Solutions. Esti- mation of Alkali Hydroxids. Estimation of Alkali Carbonates. Mixed Alkali Hydroxid and Carbonate. Estimation of Alkali Bicarbonates when Mixed with Carbonates, Sodium Borate. Sodium Cacodylate. Sodium Glycerophosphate. Estimation of Organic Salts of the Alkalies. Table of Normal Factors, etc., of the Organic Salts of the Alkalies. Estimation of Salts of the Alkali Earth Metals. Estimation of Mixed Hydroxids and Carbonates of Alkali Earths. Acidimetry. Estimation of Acids. Preparation of Standard Alkali Solutions. Table Showing Quantity to be taken for Analysis in Direct Percentage Estimations.

CHAPTER IX

ANALYSIS BY PRECIPITATION 113

Preparation of Decinormal Silver Nitrate, Decinormal Sodium Chlorid, and Decinormal Sulphocyanate. Estimation of Soluble Haloid Salts. Mohr's Method with Chromate Indicator. Ti- tration without an Indicator. Estimation of Haloid Acids. Assay of Phosphoric Acid. Estimation of Cyanogen. Esti- mation of Silver Salts. Estimation of Metallic Silver and Silver Alloys. Assay of Mercuric Salts. Table of Substances Estimated by Precipitation.

CHAPTER X

ANALYSIS BY OXIDATION AND REDUCTION 139

Preparation of Decinormal Potassium Permanganate. Volumetric Analyses by Means of Potassium Permanganate. On the Use of Empirical Permanganate Solution. Typical Analyses by Permanganate. Direct Titrations. Estimation of

CONTENTS ix

PAGE

Ferrous Salts. Estimation of Metallic Iron in Ferrum Reduc- tum. Estimation of Oxalic Acid and Oxalates. Estimation of Calcium. Estimation of Hydrogen Dioxid, Barium Dioxid, and Sodium Perborate. Estimation of Ferric Salts. Estima- tion of Nitrous Acid and Nitrites. Residual Titrations. Esti- mation of Hypophosphorous Acid and Hypophosphites. Esti- mation of Calcium Salts. Estimation of Lead Acetate and Subacetate. Estimation of Manganese Dioxid. Estimation of Nitrates, Chlorates, Chromates, and Chromic Acid. Estima- tion of Tin. Estimation of Copper.

Volumetric Analysis by Means of Potassium Dichromate. Preparation of Decinormal Potassium Dichromate. Estima- tion of Ferrous Salts. Table of Substances Estimated by Permanganate or Dichromate.

Analysis by Indirect Oxidation. Preparation of Decinormal lodin. Starch Solution. On the Use of Sodium Bicarbonate in Titrations with lodin. Estimation of Arsenous Compounds. Estimation of Antimony Compounds. Estimation of Sul- phurous Acid and Sulphites. Estimation of Sodium Thiosul- phate. Hydrogen Sulphid and Sulphids. Table of Substances Estimated by Means of lodin Solution.

Estimation of Substances Readily Reduced. lodometry. Esti- mations Involving the Use of Sodium Thiosulphate V.S. Prep- aration of Decinormal Thiosulphate V.S. Estimation of Free lodin. Indirect lodometric Estimations. Estimation of Free Chlorin or Bromin. Estimation of Available Chlorin. Esti- mation of Hydrogen Dioxid. Distillation Methods. Estima- tion of Manganese Dioxid. Chromic Acid and Chromates. Estimation of Alkali lodids. Digestion Methods. Estimation of Chlorates, Bromates and lodates. Estimation of Ferric Salts. Estimation of Chromates, Arsenates, Antimonic Salts, Copper, and Mercury.

Chlorometry, Reduction Methods, Involving the Use of Arsenous Acid Solutions. Preparation of Standard Arsenous Acid V.S. Iodized Starch Test Paper. Estimation of Free Halogens. Estimation of Available Chlorin. Chlorometric Assay of Manganese Dioxid.

CONTENTS

PAGE

Reduction Methods Involving the Use of Stannous Chlorid V.S. Estimation of Iron by Means of Stannous Chlorid. Estima- tion of Mercuric Salts.

PART II CHAPTER XI

ESTIMATION OF ALKALOIDS 251

Table of Factors for Alkaloids. Gordin's Modified Alkali- metric Method for Titrating Alkaloids.

CHAPTER XII

ASSAYING OF VEGETABLE DRUGS AND THEIR PREPARATION 259

Separation of Alkaloids and Use of Immiscible Solvents. Kebler's Modification of the Keller Method. Assay of Galenical Preparations. Lloyd's Method. Katz's Method.

CHAPTER XIII

ESTIMATIONS INVOLVING USE OF DECINORMAL BROMIN V.S 271

Preparation of Decinormal Bromin Solution. Assay of Phenol. Resorcinol. Phenylsulphonates.

CHAPTER XIV

SOME TECHNICAL METHODS FOR FATS, OILS AND WAXES 277

The Acid Value. The Saponification Number. Volatile Fatty Acid Number. Reichert's Number. The Reichert- Meissl Number. HiibPs Number. Hanus' Number. Acid Number for Resins.

CHAPTER XV

ESTIMATION OF SUGARS 291

Preparation of Fehling's Solution. Determination of the End-point. Estimation Starch after Inversion. Estimation of Maltose in Malt Extracts. Estimation of Diastasic Value of Malt Extract.

CONTENTS xi

CHAPTER XVI

PAGE

ESTIMATION OF FORMALDEHYDE 299

The Ammonia Method. The Ammonium Chlorid Method. Oxidation Method by Means of Hydrogen Dioxid. The lodo- metric Method. The Cyanid Method. Assay of Paraformal- dehyde and Acetone.

CHAPTER XVII

ESTIMATION OF ALCOHOL IN TINCTURES AND BEVERAGES 308

Alcoholometric Table.

PART III

CHAPTER XVIII

THE NITROMETER 313

The Law of Charles. The Law of Boyle.

CHAPTER XIX

ASSAY OF NITRITES 318

Spirit of Nitrous Ether. Amyl Nitrite. Sodium Nitrite. Nitric Acid in Nitrates.

CHAPTER XX

HYDROGEN DIOXID 323

Use of Nitrometer. Use of Urea Apparatus. The Hypo- chlorite Method. The Hypobromite Method. Table Showing Weight in Milligrams of Hi02 corresponding to one cc. of Moist Oxygen.

CHAPTER XXI

ESTIMATION OF SOLUBLE CARBONATES BY THE USE OF THE NITROM- ETER 328

CHAPTER XXII

ESTIMATION OF UREA IN URINE 329

The Doremus' Ureometer. The Hinds-Doremus Ureometer. Squibb 's Urea Apparatus.

APPENDIX

DESCRIPTION OF INDICATORS ALPHABETICALLY ARRANGED 334

A LIST OF THE MORE COMMON ELEMENTS WITH THEIR SYMBOLS AND ATOMIC WEIGHTS

Name.

Atomic Weight* based on O=i6.

Atomic Weight based on H = i .

Approximate Atomic Weight.

... Al.

27.1

26 o

27 O

... Sb.

1 2O 2

I IQ 3.

1 2O O

... As.

74.06

74 3

7e o

... Ba.

137 37

136 4

1 36 o

Bismuth .............

... Bi.

208.0

206 4

206 o

Boron

... B.

II .0

IO.Q

II .0

... Br.

70-02

70 36

' 80 o

Cadmium

. . . Cd.

112 40

in 6

III O

Ca.

40.07

-7Q g

Carbon

... C.

12.005

II .QI

12 O

Chlorin

... Cl.

35.46

35 .18

-2f t

Chromium . .

. Cr.

52. O

r i >7

c 2 O

Cobalt

. . . Co.

58.07

<?8 56

5**w.

fg o

. Cu.

63 57

63 i

67 o

Fluorin

... F.

IQ.O

18 o

IQ O

Gold

Au.

IO7 2

105 7

1 06 o

Hydrogen

... H.

I 008

I OOO

I O

lodin

... I.

126.92

125 .0

126 o

Iron . .*...

.. Fe.

rc 84

trr t

r6 O

Lead

Pb

2O7 2O

2O5 75

206

Lithium ....

. Li.

6 Q4

6 08

7 O

Ivlagnesium

..Mg.

24 32

24.18

24 o

Manganese

...Mn.

CA Q-l

^4.6

cc .0

M^ercury

. Hg.

2OO 6

108.5

2OO.O

Molybdenum

...Mo.

96 o

05 .3

QtJ .0

Nickel

. . Ni.

c8 68

g8.i

t;8 o

Nitrogen

... N.

14 01

1^ .O^

14.0

Oxygen

O.

16 oo

i«.88

16 o

Pho sp ho ru s

p

3O 77

•51 O

Platinum

Pt

IQ3 .3

104 O

Potassium

... K.

^yo • •

3O I

38.86

30- O

Silver

. Ag.

ov • A 107 88

107.12

107.0

Sodium

. . . Na.

22.88

21> ,O

Strontium

... Sr.

87 67

86.94

87.0

Sulphur .... .........

... S.

31.83

32.0

Tin

Sn

TlR 7

118.1

118 o

Zinc

... Zn.

6*\ 37

64.9

65 .0

International Atomic Weights, 1916.

TABLE OF MULTIPLES OF SOME ATOMIC WEIGHTS AND COMBINATIONS IN FREQUENT USE.

i

2

3

4

5

6

7

8

9

H

1.008

2.016

3.024

4-032

5.040

6.048

7.056

8.064

9:072

O

16.000

32.

48.

64-

80.

96.

112.

128.

144-

OH

17.008

34-016

51.024

68.032

85.04

102.048

119.056

136.064

153.072

H20

18.016

36.032

54-048

72.064

90.08

108.096

126. 112

144.128

162.144

N

14.01

28.02

42.03

56.04

70.05

84.06

98.07

1 1 2. O8

126.09

NH3

17.034

34.064

51 . 102

68.128

85.170

102.204

119.238

136.272

153.306

NH4

18.042

36.084

54-126

72.168

9O. 2IO

108.252

126.294

144.336

162.378

NO2

46.01

Q2.02

138.03

184.04

230.05

276.06

312.07

388.08

414.09

NQ3

62.01

I24.O2

186.03

248.04

310.05

372.06

434-07

496 . 08

558.09

C

I2.OO

24.

36.

48.

60.

72.

84.

96.

108.

C02

44.00

88.

132.

176.

220.

264.

308.

352.

396.

C03

6O.OO

120.

180.

240.

300.

360.

420.

480.

540.

CN

26.01

52.O2

78.03

104.04

130.05

156.06

182.07

208.08

234.09

Cl

35-40

70.92

106.38

141.84

177-3

212.76

248.22

283.68

319.14

Br

79-92

159.84

239.76

319.68

399-6

479-52

559-44

639.36

719.28

I

126.92

253.84

380.76

507.68

634-6

761.52

888.44

1015.36

1142.28

S

32.07

64.14

96.21

128.28

160.35

192.42

224.49

256.56

288.63

SO3

80.07

160.14

240.21

320.28

400 . 35

480.42

56o . 49

640 . 56

720.63

S04

96.07

192.14

288.21

384-28

480.35

586.42

672.49

768.56

864.63

P04

95-04

IQO.OS

285.12

380. .16

475-20

580. 24

665.28

760.32

855.36

P206

142.08

284. 16

426.24

568.32

710.40

852.48

994-56

1136.64

1278.72

Na

23-00

46.

69.

92.

US-

138.

161.

184.

207.

K

39.10

78.20

117.30

156.40

195.50

234.60

273.70

312.80

351-90

Li

6.94

13.88

20.82

27.76

34-70

41.64

48.58

55-52

62.46

Ca

40.07

80. 14

I2O.2I

160.28

200.35

240.43

280.49

320.56

360 . 63

Ba

137-37

274-74

412. II

549.48

686.85

824.22

761.59

898.96

1036.33

Ag

107.88

215.76

323.64

431.52

539-40

647.28

755.16

863.04

970.92

Fe

55.84

III. 68

167.52

223.36

279.20

335.04

390.88

446.72

502.56

Hg

200.6

401.2

601.8

802.4

1003.0

1203.6

1404.2

1604.8

1805.4

X111

ABBREVIATIONS AND SIGNS

cc cubic centimeter

gm gramme, 15.43235 grains

gr grain

at. wt atomic weight

V.S volumetric solution

T.S test solution.

U. S. P United States Pharmacopoeia.

N

— normal.

- decinormal.

10

N

— centmormal.

100

N

— semi-normal.

2

2

^ or 2N double-normal.

mil.. ..milliliter. xiv

THE ESSENTIALS OF VOLUMETRIC ANALYSIS

CHAPTER I INTRODUCTION

IN a chemical analysis the aim is to determine the nature of the chemical substances contained in a given compound or to ascertain their quantities. In the former case the analysis is a qualitative, in the latter a quantitive, one.

The quantitive analysis of a substance may be made either by the gravimetric or the volumetric method.

The Gravimetric Method consists in separating and weighing the constituents either in their natural states or in the form of new and definite compounds, the composition of which is known to the analyst. From the weights of these new compounds the analyst can calculate the quantities of the original constituents.

Example. To determine the quantity of silver in a solu- tion by the gravimetric method we proceed as follows :

Ten grams of a solution containing silver in the form of silver nitrate (AgNO3) is placed into a beaker, and, after slightly acidulating with nitric acid, is treated with hydro- chloric acid, drop by drop, until no further precipitation occurs. The precipitate which consists of silver chlorid (AgCl) is then separated by filtration, thoroughly washed, dried and weighed. Its weight is found to be 0.69 gm. The calcu- lation is then made as follows: 143.34 gms. of silver chlorid

2 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

represents 107.88 gms. of silver or 169.89 gms. of silver nitrate, as the equation shows:

AgN03 + HC1 = AgCl + HN03.

169.89 143-34

Therefore, 0.69 gm. of silver chlorid will represent

- —

gm. of silver, 143-34

or '- — X 0.60 = 0.8 1 7 gm. of silver nitrate.

143-34

The Volumetric Method. This method depends upon the use of solutions (standard solutions) which are of known strength and paying attention to the valume of such a solution which must be added to the substance under analysis to perform with it and complete a certain reaction. Thus, if we conduct an analysis by means of such a solution, and can express by a chemical equation the reaction which takes place, we can readily and accurately calculate the quantity present of the substance under analysis.

Example. If a silver solution is to be analyzed by this method it is treated with a standard solution of sodium chlorid, added slowly from a burrette until no more silver chlorid is precipitated. Each mil of this standard solution will precip- itate a certain weight of silver as silver chlorid, and hence by noting the number of mils used to complete the precipitation, the weight of the silver in the solution analyzed is easily ascertained.

N The — sodium chlorid solution is generally used for this

purpose. It is made by dissolving iV of the molecular weight of the salt (in grams) (5.846 gms.) in sufficient water to make 1000 mils. 1000 mils of this solution will precipitate TV of the atomic weight of silver (in grams) (10.788 gms.), and

INTRODUCTION 3

hence each mil of the sodium chlorid solution represents 0.010788 gm. of metallic silver, and by multiplying this figure by the number of mils of the sodium chlorid solution used, the quantity of silver in the solution under analysis is ascer-

N

tained. If in the above analysis 100 mils of the — sodium

10

chlorid solution were used, then 0.010788X100=1.0788 gms. of metallic silver.

The reaction is illustrated by this equation :

AgN03 + NaCl = AgCl + NaNO3. 10)107.88 10)5846 N

1000) 10. 788 gms. 1000) 5.846 gms. = 1000 mil — V.S.

10 0.010788 gm. 0.005846 gm. = i mil " "

From the examples given it will be seen that the gravi- metric operations consume no little time, and require the exercise of considerable skill. The washing of the precipitate must be thoroughly performed in order that it be freed from all adhering matter. The drying also is a matter of some consequence and must be performed in such a manner as to prevent the admixture of dust or the decomposition of the precipitate by excessive heat. A very accurate balance is also required.

The volumetric operations, on the other hand, do not require that the substance to be determined be separated in the form of a compound of known composition and weighed in the dry state; in fact, the substance may be accurately estimated when mixed with many others. It therefore obviates the necessity for the frequent separations and weighings which the gravimetric method demands, and enables the analyst to do the work in a very short time.

The instruments needed for volumetric work are few and simple, and comparatively little skill is required. Further- more, the results obtained are in most instances more accurate.

CHAPTER II

GENERAL PRINCIPLES OF CHEMICAL COMBINATION UPON WHICH VOLUMETRIC ANALYSIS IS BASED

i. WHEN substances unite chemically the union always takes place in definite and invariable proportions. Thus when silver nitrate and sodium chlorid are brought together, 169.89 parts (by weight) of silver nitrate and 58.46 parts (by weight) of sodium chlorid will react with each other, producing 143.34 parts of a curdy white precipitate (silver chlorid).

These substances will react with each other in these pro- portions only.

If a greater proportion of silver nitrate than that above stated be added to the sodium chlorid, only the above pro- portion will react, the excess remaining unchanged.

The same is true if sodium chlorid be added in excess of the above proportions. For instance, if 200 parts of silver nitrate be mixed with 58.46 parts of sodium chlorid, 169.89 parts only will react with the sodium chlorid, while 30.11 parts of silver nitrate will remain unchanged. Again, when potassium hydroxid and sulphuric acid are mixed potassium sulphate is formed, 112.2 parts of potassium hydroxid and 98.1 parts of sulphuric acid being required for complete neutralization. These two substances unite chemically in these proportions only.

The equation is

2KOH + H2SO4 = K2SO4 + 2H2O.

GENERAL PRINCIPLES OF CHEMICAL COMBINATION 5

In other words, 112.2 parts of KOH will neutralize 98.1 parts of H2SO4, and consequently 98.1 parts of H2SO4 will neutralize 112.2 parts of KOH.

Oxalic acid and sodium carbonate react upon each other in the proportions shown in the equation

H2C2O4 • 2H2O + Na2CO3 = Na2C2O4 + CO2 + 3H

126.05 106

126.05 parts of crystallized oxalic acid are neutralized by 106 parts of anhydrous sodium carbonate.

2. Definite chemical compounds always contain the same elements in exactly the same proportions, the proportions being those of their atomic weights, or some multiple of these weights.

Thus sodium chlorid (NaCl) contains 23 parts of metallic sodium and 35.46 parts of chlorin, these being the atomic weights of sodium and chlorin respectively.

Potassium sulphate (K2SO4) contains twice 39.1 = 78.2 parts of potassium, 32.07 parts of sulphur, and four times 16 = 64 parts of oxygen.

Potassium hydroxid (KOH) contains 39.1 parts of potas- sium, 16 parts of oxygen, and one part of hydrogen. Hydro- chloric acid (HC1) contains one part of hydrogen and 35.46 parts of chlorin.

Upon these facts the volumetric methods of analysis are based.

It has been shown that 98.1 gms. of sulphuric acid will neutralize 112.2 gms. of potassium hydroxid. It is therefore evident if a solution of sulphuric acid be made containing 49.05 gms. of the pure acid in 1000 mils, that one mil of this solution will neutralize 0.0561 gm. of potassium hydroxid. In estimating alkalies with this acid solution, the latter is

6 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

added from a burette, in small portions, until the alkali is neutralized, as shown by its reaction with some indicator.

Each mil of the acid solution required before neutralization is complete indicates 0.0561 gm. of KOH, and the number of mils used multiplied by 0.0561 gm. gives the quantity of pure KOH in the sample analyzed.

One mil of the same solution will neutralize 0.04 gm. of sodium hydroxid (NaOH), 0.053 gm- °f anhydrous sodium carbonate (Na2CO3), etc.

If a solution of crystallized oxalic acid be made by dis- solving 63.02 gms. in sufficient water to make icco mils, we will have a normal solution, the neutralizing power of which is exactly equivalent to the above mentioned normal sulphuric acid solution.

The strength of acids is estimated by alkali volumetric solutions. A normal solution of potassium hydroxid containing 56.1 gms. in the liter will neutralize exactly i liter of the normal acid solution; imil of this normal alkalfrvill neutralize 0.03646 gm. of HC1, 0.06362 gm. of H2C2O4,or 0.04905 gm. of H2SO4, etc.

CHAPTER III VOLUMETRIC OR STANDARD SOLUTIONS

ANY solution employed in volumetric analysis for the purpose of estimating the strength of substances, that is, any solution the chemical power or titer of which has been deter- mined, is designated a standard or volumetric solution. Such a solution is said to be "titrated" (French litre = title or power), and is sometimes also called a set solution or a stand- ardized solution. It may be normal, decinormal, empirical, or of any strength, so long as its strength is known.

When volumetric analysis first came into use the solutions were so made that each substance to be estimated had its own special volumetric solution, and this was usually of such strength as to give the result in percentages. Thus a certain strength of standard acid was employed for potash, another for soda, and a third for ammonia, and in testing the acids, each had its own special standard alkali. These solutions were known as normal solutions; they are still to some extent in use, and since solutions now designated as normal are of an entirely different character, it is important that no miscon- conception should exist when a normal solution is spoken of.

Normal Solutions are those which contain one liter (1000 mils), the molecular weight of the active reagent expressed in grams, and reduced to the valency corresponding to one atom of replaceable hydrogen or its equivalent. In other words, those which contain in 1000 mils in any given reaction the chemical equivalent of one gram of hydrogen. The now ac- cepted basis for the atomic weights O = i6 makes a slight

7

8 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

change in the quantity of reagent in a normal solution, i.e., it will contain in 1000 mils the exact equivalent of 8 gms. of oxygen.

Thus in univalent or monobasic compounds the full molec- ular weight in grams is contained in a liter of the normal solu- tion.

Example. Hydrochloric acid, HC1, having one replaceable hydrogen atom, its normal solution would contain the full molecular weight, 36.46 gms. in a liter. A normal solution of potassium hydroxid should contain 56.1 gms. of KOH in a liter, while that of sodium hydroxid should contain 40 gms. of absolute NaOH.

Normal solutions of bivalent or dibasic compounds contain in 1000 mils one-half the molecular weight in grams. Thus, oxalic acid, H2C2O4 + 2H2O = 126.05, having two replaceable H atoms, one-half of its molecular weight in grams = 63. 02 is contained in a liter of its normal solution. For the same reason a liter of a normal solution of sulphuric acid contains

——=49.05 gms., and a liter of a normal solution of sodium

carbonate Na2CO3 contains — = 53 gms., while in the case

of trivalent or tribasic compounds one-third of the molecular weight in grams is contained in a liter of the normal solution. Thus it will be seen that one mil of any normal acid solution will neutralize one mil of any normal alkali solution, because one molecule of. a univalent acid will neutralize one molecule of a univalent alkali, or a half a molecule of a bivalent alkali. This is shown by the equations

HC1 + NaOH = NaCl + H2O,

36.46 46

2HC1 + Na2C03 = 2NaCl + H2O + CO2. 2)72.92 2)106

36.46 ""53

VOLUMETRIC OR STANDARD SOLUTIONS 9

The value of a reagent as expressed by its hydrogen equivalent is readily seen in the case of acids and alkalies by reference to the chemical formula, but in such standard solutions as potassium dichromate, potassium permanganate, sodium thiosulphate, and others, the particular reaction in any given analysis must be taken into account in making a normal solution; for instance, when K2Cr2O7 is to be used as a precipitating agent, its reaction is as follows :

2Ba(C2H3O2)2 + K2Cr2O7 + H2O = 2BaCrO4 + 2KC2H3O2

+ 2HC2H3O2.

It is thus seen that one molecule of K2Cr2C>7 will cause the precipitation of two atoms of barium in the form of chromate. Each atom of barium is chemically equivalent to two atoms of hydrogen; therefore one-fourth of a molecule of K2Cr2O7 is equivalent to one atom of hydrogen. And therefore a normal solution of this salt, when used as a precipitating agent, must contain in one liter one-fourth of

204.2

its molecular weight in grams; -- = 73.55 gms.

4

If K2Cr2O7 is to be used as an oxidizing agent, the three atoms of oxygen which it yields for oxidizing purposes must be taken into account. When this salt oxidizes it splits up into K2O + Cr2O3 + O3. The three atoms of oxygen combine with and oxidize the salt acted upon, or they combine with an equivalent quantity of the hydrogen of an acid and liberate the acidulous part, which then combines with the salt. As the equations show,

6FeO + K2Cr2O7 = K2O + Cr2O3 + 3Fe2O3 or (Fe6O9) ; 6FeSO4 + K2Cr2O7

7H2S04 + K2Cr207 - 3SO4 + 7H2O + K2SO4 + Cr2(SO4)3.

10 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

Each of these atoms of oxygen are equivalent to two atoms of hydrogen. Thus Os is equivalent to He.

Hence a liter of a normal solution of K2Cr2O7, when used as an oxidizing agent, contains one-sixth of its molecular weight in grams.

The same may be said of potassium permanganate when used as an oxidizing agent.

2KMnO4 has five atoms of oxygen which are available for oxidizing purposes, and each of these is capable of taking two atoms of hydrogen from an acid and liberating the acidulous part. The hydrogen equivalent of this salt may therefore be said to be one-tenth of the weight of 2KMnO4, and a normal solution of this salt contains 31.606 gms. in a liter.

Sodium Thiosulphate (Hyposulphite), Na2S2O3, is another instance. The molecule of this salt has two atoms of sodium, which have replaced two atoms of hydrogen of thiosulphuric acid. Thus it would seem that a normal solution should contain one-half of the molecular weight in grams. But the particular reaction of this salt with iodin is taken into account.

One molecule reacts with one atom of iodin, as seen in the equation

2Na2S2O3 + 12 = 2NaI + Na2S4O6.

Since iodin is univalent, a molecule of the salt is equivalent to one atom of hydrogen.

A normal solution of this salt therefore contains the molecular weight in grams in a liter.

N Decinormal Solutions, — , are one- tenth the strength of

normal solutions.

N Centinormal Solutions, — , are one-hundredth the strength

of normal solutions.

VOLUMETRIC OR STANDARD SOLUTIONS H

N Seminormal Solutions, — , are one-half the strength of

normal solutions.

2

Double-normal Solutions, ^, are twice the strength of

the normal.

Empirical Solutions. A solution which does not contain an exact atomic proportion of the reagent may be employed as a volumetric solution after its strength or titer has been determined. Such a solution is said to be empirical, and solutions of this sort are very frequently used. To prepare solutions of exactly normal strength is a tedious process and often inconvenient. If the solution is approximately normal and its strength accurately determined, it may be used as it is. Again, in the case of standard solutions of the caustic alkalies, wJiich, when not kept with all precautions, deteriorate readily through absorption of carbon dioxid from the air, as well as through their action upon the glass containers. To restore the titer of such solutions by the introduction of more of the alkali is an unnecessary waste of time, inasmuch as it is only necessary to determine its exact strength and then use it as it is. For instance, if an approximately normal solution of potassium hydroxid is on hand, its strength is determined as follows s

Ten mils of an exactly normal oxalic or other acid solution are put into a beaker, and after diluting with a little water, and adding three or four drops of methyl orange, the empirical potassium hydroxid solution is run in from a burette until the color of the solution changes from red to yellow; the number of mils required is then noted.

Assuming that 10.4 mils were required to neutralize

the 10 mils of normal acid, hence its strength is — or

104

0.9615, that of a strictly normal solution, and the number of

12 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

TOO

mils used in any estimation must be multiplied by —

or 0.9615, and then by the normal factor for the substance analyzed.

It is a good plan to have the factor marked on the label of the bottle containing such an empirical solution. In this case it would be X 0.961 5 = normal.

Standard solutions for use in volumetric analysis are usually solutions of acids, bases, or salts, and in two cases elements, namely, iodin and bromin.

A- standard solution of a base is usually used for the esti- mation of free acids.

A standard solution of an acid is usually used for the estimation of a free base, or the basic part of a salt, the acid of which can be completely expelled by the acid used in the standard solution. Example, carbonates.

A standard solution of a salt may be used as a precipitant, or it may be used as an oxidizing or reducing agent.

That part of the reagent in a standard solution which reacts with the substance under analysis is the active con- stituent of the solution. As Ag in AgNO3 is the active con- stituent of the standard solution of silver nitrate,

AgNO3 + NaCl = AgCl + NaNO3,

or Cl in NaCl, is the active constituent of the standard solution of sodium chlorid.

If the reagent is a base, as KOH, the basic part K is the active constituent. If the reagent is an acid, the active constit- uent is the acidulous part, as SO4 in H2SO4.

If the action of the reagent is oxidizing, then that part of the reagent which produces the oxidation is the active constituent.

The valence of an acid is shown by the number of replace-

VOLUMETRIC OR STANDARD SOLUTIONS 13

able hydrogen atoms it contains. 'Thus HC1 is univalent, H2SO4 is bivalent, which means that a molecule of HC1 is chemically equivalent to one atom of hydrogen, and a molecule of H2SO4 is chemically equivalent to two atoms of hydrogen.

The valence of a base is shown by the number of hydroxyls it is combined with. As KOH is univalent, Ca(OH)2 is bivalent.

The valence of a salt is shown by the equivalent of base which has replaced the hydrogen of the corresponding acid.

Thus NaCl, in which Na has replaced H of HC1, is uni- valent.

K2SO4, in which K2 has replaced H2 of H2SO4, is bivalent.

Preparation of Volumetric Solutions. In preparing volu- metric solutions it must be borne in mind that most salts when dissolved in water cause a condensation in volume, through reduction of temperature, while some substances, as for instance sulphuric acid and alkali hydroxids, cause a rise in tempera- ture and a consequent increase in volume. It is therefore necessary, after making a solution, to set it aside for a short time, in order to allow it to attain the proper temperature before measuring it.

It is always the best plan to take a weighed quantity of the salt, slightly greater than that required by theory, and to dissolve it in less water than is needed for the finished solution. This solution is titrated, its strength determined and then diluted to the proper measure.

After dilution it should be again carefully titrated and its titer verified.

All volumetric solutions should be made with distilled water and with reagents which are of a high degree of purity.

Standard Temperature. A cubic centimeter is the volume occupied by one gram of distilled water at its maximum density 4° C. (39° F.). This, however, is not the cubic centimeter

14 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

used in volumetric analysis. It is convenient to use in analyses of this sort a cubic centimeter which represents the volume of one gram of distilled water at a temperature which is easily attained and maintained at any season of the year.

The temperature at which measuring instruments are graduated is the temperature at which volumetric solutions should be prepared, and at which all volumetric operations should be conducted. 4° C. is a temperature at which it is obviously impossible to work except during one or two months of the year. For this reason the temperature of 16° C. (60.8° F.) has been taken as the standard. The cubic centimeter used in volumetric analysis, under this standard, is the volume occupied by one gram of distilled water at the latter tempera- ture.

The employment of this standard of temperature, though long in vogue, is justly criticized as too low. In order to obtain accurate results the temperature of the atmosphere in which the titration is performed must not be too much at variance with the temperature at which the instruments are graduated and the solutions made. A temperature of 16° C. is one which is exceedingly difficult to maintain in the warm months of the year, therefore it has been suggested to take a higher temperature as the standard.

The U.S.P. VIII recommends the employment of 25° C. (77° F.) as the standard. This, while better than the lower temperature, is regarded by many as being too high and the use of 20° C. (68° F.) as the standard is now being very favorably considered, this being nearer the average temperature of the atmosphere in laboratories throughout the year. What ever temperature is adopted, it is at this temperature that the whole set of measuring instruments must be graduated, and all titrations carried out. It would be obviously improper to use a burette graduated at 16° C. and a flask or cylinder

VOLUMETRIC OR STANDARD SOLUTIONS 15

graduated at 25° C., or to employ solutions at a temperature which is different from that at which they are made.

Mil. The term cubic centimeter (cc.) has been replaced by the word mil in the U. S. P. IX, and in many other recent books. The United States Bureau of Standards declared that the term cubic centimeter was a misnomer, there being a slight difference between the thousandth part of a liter and the cubic centimeter, as one liter was determined to be equiv- alent to 1.000027 cubic decimeters. Hence the word mil, an abbreviation of milliliter, was adopted and represents the thousandth part of a liter.

The U. S. P. liter is the volume occupied by 996.04 gms. of distilled water, weighed in air with brass weights at a tem- perature of 25° C. and barometric pressure of 760 mm.

To Titrate a substance means to test it volumetrically for the amount of pure substance it contains. The term is ^ised in preference to " tested " or " analyzed," because these terms may relate to qualitative examinations as well as quantitative, whereas titration applies only to volumetric analysis.

Residual Titration, Re-titration, sometimes called Back Titration, consists in treating the substance under examina- tion with standard solution in a quantity known to be in excess of that actually required; the excess (or residue) is then ascertained by residual titration with another standard solution.

Thus the quantity of the first solution which went into combination is found.

N Example. Ammonium carbonate is treated first with —

H2SC>4 in excess, and the excess then found by titration with

16 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

N The quantity of the — KOH used is then deducted from

N the quantity of — H2SO4 added, which gives the quantity of

the latter which was neutralized by the ammonium carbonate.

Titrations may be carried out in flasks, beakers, or in white porcelain evaporating dishes. Flasks of the Erlenmeyer pattern, see Fig. 24, having a short narrow neck and a broad flat bottom, are very desirable for this purpose; they admit of shaking their contents without danger of loss, and permit ready observation of color changes. If a flask is used the tip of the burette should extend well into its neck in order to prevent any loss of the reagent. The flask should be rotated after each addition of the reagent, and when the end of titration is near, any of the solution adhering to the sides of the flask should be washed down with distilled water. A white porcelain tile, or a sheet of white paper placed under the flask or beaker, aids materially in the observation of the color change.

CHAPTER IV INDICATORS*

IN volumetric analysis the substance to be analyzed in the state of solution is placed in a beaker and the standard solution is added from a burette until a certain reaction is produced. The exact moment when a sufficient quantity of the standard solution has been added is known by certain visible changes, which differ according to the substance analyzed and the standard solution used. When such a visible change occurs the " end-reaction " is reached.

The end-reaction manifests itself in various ways, as follows :

1. Cessation of precipitation.

2. First appearance of a precipitate.

3. Change of color.

In some cases, however, the addition of the standard solution to the substance under analysis does not produce either a precipitate or a change of color; in such cases we must resort to the use of an indicator.

An indicator is a substance which is used in volumetric analysis, and which indicates by change of color, or some other visible sign, the exact point at which a given reaction is complete.

Generally the indicator is added to the substance under examination, but in a few cases it is used alongside, a drop of the substance being occasionally brought in contact with a drop of the indicator.

* A more detailed description of the individual indicators is given in the Appendix.

17

18 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

Thus in estimating an alkali with an acid volumetric solu- tion the alkali is shown to be completely neutralized when the litmus tincture which was added becomes faintly red or the phenolphthalein colorless. Again, when haloid salts are estimated with nitrate of silver solution, chromate of potas- sium is added as indicator. A white precipitate is produced as long as any halogen is present to combine with the silver, and when all is precipitated the chromate of potassium acts upon the silver nitrate, forming the red silver chromate, this color thus showing that all the halogen has been precipitated.

INDICATORS COMMONLY USED

The principle indicators used are:

Tincture of Litmus, which shows acidity by turning red and alkalinity by becoming blue.

Phenolphthalein Solution, which is colorless in acid solu- tions and red in alkaline solutions, but is not reliable for alkaline phosphates, bicarbonates or ammonia.

Methyl-orange Solution turns red with acids and yellow with alkalies. It is not affected by carbonic acid, and is therefore adapted for the titration of alkaline carbonates.

Rosolic-acid Solution is yellow with acids and violet-red with alkalies. It is very sensitive to ammonia.

Tincture of Turmeric turns brown with alkalies, and the yellow color is restored by acids.

Cochineal Solution turns violet with alkalies and yellowish with acids. It is used chiefly in the presence of ammonia or alkali earths.

Eosin Solution is red by transmitted light, and shows a strong green fluorescence by reflected light. Acids destroy this fluorescence and alkalies restore it.

Brazil-wood Test-solution turns purplish-red with alkalies and yellow with acids.

INDICATORS 19

Fluorescein Test-solution shows a strong green fluorescence by reflected light in the presence of the least excess of an alkali.

Neutral Potassium-chromate Test-solution is used in the titration of haloid salts with silver-nitrate solution. It indicates that all the halogen has combined with the silver by producing a red-colored precipitate (silver chromate).

Potassium-ferricyanide Test-solution is used in the estima- tion of ferrous salts with potassium-dichromate solution. It gives a blue color to a drop of the solution on a white slab as long as any iron salt is present which has not been oxidized to ferric.

THE IONIZATION OR DISSOCIATION THEORY

When a soluble salt dissolves in water, its molecules split up or dissociate more or less completely into parts called ions. This behavior of substances, on going into solution, is known as electrolytic dissociation or ionization.

Ions are electrically charged atoms or groups of atoms and are of similar composition to the substances formed from the compound when an electric current is passed through the solution. The electro-positive ions migrate to and collect around the negative pole (cathode) and hence are called cathions, while the electro-negative ions are called anions, because they concentrate around the positive pole or anode. The dissociation of a compound into its ions when an electric current is passed through its solution, although spoken of as electrolytic dissociation, is really not caused by the electric current, since the dissociation into ions occurs at once upon dissolving the substance in water and without the aid of an electric current, the action of the current being the transportation of the separated ions to the poles.

20 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

The dissociation of compounds into ions when dissolved in water is illustrated in the following list:

Sodium chlorid into (Na + ) and (Cl — )

Potassium nitrate into (K + ) " (NO3-)

hydroxidinto (K+) " (OH-)

acetate into (X + ) " (C2H3O2-)

Sulphuric acid into (H + ) " (HSO4 - )

or(H + ) " (H + )and(S04)

The extent of this dissociation depends upon the nature of the substance and the degree of dilution; the greater the dilution the more complete the dissociation. Furthermore, strong acids and bases dissociate readily, even in compara- tively concentrated solutions, while the weaker acids and bases are more or less undissociated when dissolved, i.e., they are not readily split up into ions. Their salts, however, are immediately and completely ionized. Therefore, upon neu- tralizing a weak acid or base, an ionizable salt is formed. According to the theory of Arrhenius, the reactions of analytical chemistry are chiefly reactions between ions and not between atoms.

Strong acids, bases and salts exist in solution, not as molecules but chiefly in the form of ions. The formation of silxer chlorid by the reaction between silver nitrate and sodium chlorid takes place according to the following equation:

Ag/NO3Aq. +Na/ClAq. = AgCl( solid) + Na/NO3Aq.

The state of dissociation being denoted by the vertical

line between the ions of the molecules.

+ +

This theory also explains why K/CIO3 with Ag/NOa does

not form AgCl, in that the reaction involves the ion C1O3 and not the atom Cl.

INDICATORS 21

Theories of Indicators. In connection with the use of indicators in neutralization analyses, the question as to the cause of the color changes is one of considerable interest.

Two distinct views are held. Of these the lonization Theory of Ostwald has received almost universal acceptance. According to this theory the color changes are ascribed to a change in the indicator from a molecular to an ionic condition. As exemplified in the case of phenolphthalein the colorless molecule

OCOC6H4C-(C6H4OH)2 .... (I)

i I

is dissociated into the red negative ion

OCOC6H4C.(C6H4OH)C6H46. . . . (II)

In the other and less known view on this subject (the Chromophoric Theory)* the sensitiveness of the indicators and its color changes is ascribed to a change in the consti- tution of the indicator, involving a chromophoric group, under the influence of hydrogen and hydroxyl ions. According to this view the color change is due (in the case of phenol- phthalein) to a change of constitution from the colorless lactoid (I) with no chromophoric group, to the colored quinoid (III) with a quinone chromophore.

(NaOOC-C6H4)(HOC6H4)C:C6H4:0, . . (Ill)

and that the ionization of the sodium salt is merely a coin- cidence and not the cause of the color change.

Whichever of these views is the correct one, remains for future investigations to prove. That of Ostwald, being most generally accepted at the present time, is treated more fully below.

* See Julius Stieglitz, Jour. Am. Chem. Soc., XXV, 1112 (1903).

22 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

The lonization Theory of Indicators.* The indicators used in alkalimetry and acidimetry are compounds of feeble acid or basic character, and hence not prone to dissociation in solution, but when neutralized the salt formed ionizes the instant of its formation; the ions so liberated give rise to colors which differ from those of the original compounds.

Any feeble acid or base may be utilized as an indicator if its ions have a color different from that of the un-ionized compound. Strong acids or bases are not suited as indicators, because they ionize while in a free state on dilution, and thus give no color when neutralized.

A solution in which H ions predominate has an acid reac- tion, while one in which OH ions predominate reacts alkaline.

Phenolphthalein is a feebly acid indicator, and in its undissociated state is colorless. It does not dissociate readily unless neutralized, but when sodium hydroxid is added to its solution, a sodium salt of phenolphthalein is formed which immediately ionizes and the ions liberated impart to the solution a brilliant red color. If now some acid is added the sodium salt is decomposed and the acid phenolphthalein again set free, and the solution becomes colorless.

If a few drops of phenolphthalein solution be added to an acid solution, ionization of the former is prevented by the presence of the stronger acid; if now some sodium hydroxidr solution is added, the OH ions of the latter unite with the H ions of the acid, and when the acid is completely neutralized the first drop of excess of alkali unites with the phenolphtha- lein, forming a salt which immediately ionizes and produces the characteristic red color which shows the end of the reaction.

In the titration of a feeble acid the end-point is often

* See Ostwald's ''Lehrbuch der Allgemeinen Chemie," 1891, and "Scientific Foundations of Analytical Chemistry," 1900; also Walker's "Introduction to Physical Chemistry," 1899.

INDICATORS 23

indistinct and is lacking in sharpness; this is because the indicator used has a greater tendency to ionize than the acid itself. In this case the H ions present just before the com- pletion of the reaction are not in sufficient amount to fully retard the ionization of the indicator, and hence the latter dissociates partly before neutralization is complete and gives rise to an indefinite end-reaction. Therefore it is necessary when titrating a feeble acid that an indicator should be selected whose alkali salt ionizes with the production of a distinct color change, and whose tendency to ionize is less than that of the acid. Phenolphthalein is a suitable indicator in this case, provided a strong alkali be used for titrating.

Fixed alkalies readily yield ionizable salts with phenol- phthalein, but ammonia does not. The latter being too weak a base to yield with so feeble an acid, a salt which can with- stand the hydrolytic action of the water in dilute solutions, and as a consequence a larger excess of the ammonia must be used to overcome this. Thus is accounted for the imperfect color change of phenolphthalein when ammonia or its salts are present and why the color becomes visible only after a large excess of the alkali is added.

Paranitrophenol is also an acid indicator; it exists in solution in the form of undissociated colorless molecules, yet its electro-negative ion is intensely yellow in color. This com- pound has a slight tendency to dissociate in dilute solutions, but the presence of a trace of a stronger acid will overcome this tendency and the solution remains colorless. If an alkali is, however, added, a salt of paranitrophenol is formed which immediately ionizes and exhibits the intense yellow color of its liberated ion.

Other indicators exhibit a color in both the ionized and the non-ionized state, but the colors in both conditions are different, as in the case of litmus, lacmoid and methyl orange.

24 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

Methyl orange is both an acid and a base and will form ionizable salts with either acids or alkalies; its indicator characteristics are, however, due essentially to its basic char- acter.

The salt which methyl orange forms with acids dissociates into red ions; this, upon the addition of an alkali, returns to its undissociated state, which is yellow. Because of its weak basic character its compound with acids is readily decomposed by alkalies, but it takes a strong acid or a relatively large quantity of a feeble acid to dissociate it in its non-ionized state, hence this indicator is very sensitive to alkalies, and much less so to acids.

With reference to the acid character of the indicator the explanation of its action is that the non-ionized indicator is red, while its ion is yellow. Acids lessen its ionization and produce a red color, while alkalies produce a highly ionizable salt and hence a yellow color. When a weak but slightly ionizable acid is added to the methyl orange solution, the H' ions of the acid given up in excess of the amount required for neutralization are not sufficient to yield enough of a non- ionizable salt to produce a decided red color, hence a large quantity of such a weak acid is required to give an acid indication. This would explain why methyl orange is not suitable as an indicator for weak acids, and why it is very sensitive to alkalies.

Litmus is an acid indicator which slightly dissociates in solution. Its non-ionized molecules are red, but its negative ions are blue. If it is added to an alkali, a salt is formed which at once ionizes and gives a blue color. If added to an acid, ionization is prevented and the red color of the non- ionized molecules appears.

From the above explanations it will be seen that indicators cannot be indiscriminately used, and that no one indicator

INDICATORS 25

will be suitable for every titration. Hence the indicator must be selected to suit each case. This selection is facilitated by reference to the classification of indicators, according to F. Glaser, Ztchr. f. a. Chem., 1899, 273 + .

Group I. Indicators Forming Fairly Stable Salts. The members of this group comprise such indicators as are (i) of a strong acid character and which react readily with weak bases, or (2) of a feeble basic character and which require a strong acid to form a stable salt. Hence they will be found to be very sensitive to alkalies, and are useful in the titration of weak bases, as ammonia and the amine bases, as well as strong bases and acids. The indicators of this group are the following, arranged in the order of their sensitiveness towards alkalies:

(i) lodeosin, Resazurin; (2) Tropaeolin OO, Luteol; (3) Methyl and Ethyl Orange; (4) Congo Red; (5) Cochineal; (6) Lacmoid.

Group II. Indicators Possessing Faint Acid Properties and Yielding Salts which are Very Unstable. These are readily decomposed by relatively feeble acids, and are in consequence very sensitive towards acids, slightly so towards alkalies.

They are: (i) Rosolic acid; (2) Curcuma; (3) Phenol- phthalein, Flavescin; (4) Alpha-naphtholbenzein.

Group III. Indicators Occupying a Place Midway between the Other Two Groups. They are somewhat stronger acids than those of Group II, but feebler than those of Group I.

They are fairly sensitive towards both acids and alkalies, but are more sensitive towards acids than those of Group I, and less so towards alkalies. They are:

(i) Fluorescein, Phenacetolin; (2) Haematoxylin, Gallem, Alizarin; (3) Litmus; (4) Paranitrophenol.

This division of indicators into groups, as above, facilitates the selection of an indicator suitable for the work in hand.

26 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

For instance, for titrating weak acids, a glance at the groups will show that the members of Group II are best adapted for this purpose. Again, weak bases will be best titrated by the indicators of Group I. Strong acids or bases may be titrated by means of any of the indicators.

The quantity of indicator taken in a titration is a matter of considerable moment. The smallest quantity which will produce a distinct color should be taken, but it is equally important that the quantity be not too small for the volume of liquid; for in high dilutions the hydrolytic action of water asserts itself, and intermediate tints will result, which interfere with the sharpness of the end color.

If too much of the indicator is used, the sensitiveness is lessened, because acid or alkali must be added to convert the indicator into a salt, or when formed to decompose it; i.e., a minimum of excess of the titrating fluid would react with a small portion only of the indicator and intermediate tints would be produced, until sufficient of the titrating solution has been added to neutralize all of the indicator present. This is especially true when using centinormal solutions. 20 drops of litmus added to 10 mils of water require from 10 to 14 drops

N

of acid or alkali solution to produce a change of color.

100

Thus the indicator itself takes up some of the standard solution, and hence the necessity for using as small a quantity of the indicator as possible; usually from j to 5 drops of the indicator may be taken to each 50 or 100 cc. of the fluid titrated.

The degree of dilution of the substance titrated is also a matter of considerable moment. In very concentrated solu- tions ionization does not so readily occur, while too great a dilution diminishes the reactive ability of the ions because of their greater separation, and also because of the hydro- lytic dissociation of water itself into H* and OH ions which

INDICATORS 27

react acid and alkaline respectively, and which brings about a premature dissociation of the indicator.

The Requirements of a Good Indicator, according to H. A. Cripps, are:

I. The end-reaction should be marked by a prominent change of color.

II. The smallest possible quantity of the reagent should be required to effect this change.

III. High tintorial power, which of itself assists in the fulfilment of the second requirement, less of the indicator being required.

IV. The change of color should not be affected by the impurities commonly present in the substance under examina- tion, nor by the products of the reaction.

In addition to these requirements it is a distinct advantage if the color reaction is equally decided in alcoholic as in aque- ous liquids.

A GUIDE FOR THE SELECTION OF INDICATORS

For Hydroocids and Carbonates

Indicators not affected by C02 Indicators affected by CO2

(Cold Titrations) (Hot Titrations)

Methyl Orange, Gallem, Phena- Phenolphthalein (useless in cetolin, Congo Red, lodeosin, presence of NH3 or its salts), Cochineal. Lacmoid, Rosolic Acid, Resazurin.

FOR AMMONIA (NH3) FOR AMMONIUM CARBONATE

Rosolic Acid, Methyl Orange, Same indicators as for ammonia, Congo Red, Litmus, Gallem. also Phenacetalin and Phenol-

phthalein.

28 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

FOR INORGANIC ACIDS

t, HCl, HN03.

Phenolphthalein, Litmus, Rosolic Acid, Methyl Orange. H3PO<.

Phenolphthalein (neutralized to Na2HPO4).

Methyl Orange and Cochineal (each neutralized to NaH2PO4).

Rosolic Acid and Methyl Orange. H3B03.

Phenolphthalein (after addition of glycerin), Litmus and Turmeric paper.

FOR ORGANIC ACIDS

Phenolphthalein (all) Rosolic Acid (except acetic, citric and tartaric), Gallei'n.

CHAPTER V

APPARATUS USED IN VOLUMETRIC ANALYSES

The Burette is a graduated glass tube which holds from 25 to 100 mils and is graduated in fifths or tenths of a mil

and provided at the lower end with a rubber tube and pinch-cock. The use of this instrument is to accurate- ly measure quantities of standard solutions used in an analysis. It is in an upright position when in use, and the flow of the solution can be regulated so as to run out in a stream or flow in drops by pressing the pinch-cock between the thumb and forefinger. The quantity of solution used can be read from the graduation on the outside of the tube. This is the simplest and most common form of burette, and is known as Mohr's (Fig. i).

The use of the pinch-cock in Mohr's burette may be dispensed with by introducing into the rubber tube a small piece of glass rod, which must not fit too tightly. By firmly squeezing the rubber tube surrounding the glass rod a small canal is opened, through which the liquid escapes. A

29

FIG. i.

30 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

Closed

48-

very delicate action can in this way be obtained, and the flow of the liquid is completely under the control of the operator. (See Fig. 2.)

The greatest drawback to this burette is that it cannot be used for permanganate or other solu- tions that act upon the rubber.

This defect can be overcome by the use of a burette having a glass stop-cock in place of the rubber tubing and pinch- cock. This form has the additional advan- tage of being capable of delivering the solution in drops while both hands of the operator are disengaged (Fig. 3).

Another good arrangement is that in which the tap is placed in an oblique position, so that it will not easily drop out of place

(Fig. 4).

These glass stop-cock burettes should be emptied and washed immedi- ately after use, especially if soda or potassa solution has been used; for these act upon the glass and often cause the stop- per to stick so firmly that it cannot be turned or removed without danger of breaking the instrument.

The most satisfactory form of glass stop-cock is that shown in Fig. 5.

When a number of estima- tions are to be made in which the same volumetric solution is employed, the arrangement shown in Fig. 6 is very serviceable.

50-

FIG. 3.

FIG. 4.

APPARATUS USED IN VOLUMETRIC ANALYSIS 31

A T-shaped glass tube is inserted between the lower end of the burette and the pinch-cock and connected by a rub- ber tube with a reservoir containing the volumetric solution.

The tube which communicates with the reservoir is provided with a pinch-cock, which, when open, allows the solution to flow into and fill the burette in so gradual a manner that

FIG. 5.

FIG. 6.

no bubbles are formed. The burette is emptied in the usual manner.

E. &> A. Automatic Burette (Fig. 7). This is used for the same purpose as the foregoing. It is provided with a side tube for connection with reservoir, and has an overflow cup which prevents its being filled to above the zero mark. The three-way stop-cock is so arranged that if turned one

32 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

way the inlet is opened and the liquid from the reservoir flows into and fills the burette. If turned the other way the inlet is closed and the outlet is opened and the burette may be

FIG. 7.

FIG.

emptied. If the handle of the stop-cock is turned half-way round, both openings are closed.

There are many other forms of automatic burettes.

When working with solutions which are readily altered by contact with air, as for example, stannous chlorid, potas-

APPARATUS USED IN VOLUMETRIC ANALYSIS 33

slum, sodium, or barium hydroxid or ammonia, an arrange- ment like that depicted in Fig. 8 is very serviceable. In this

the upper end of the burette is connected with the reservoir by means of a rubber tube, thus making an air-tight combination between the burette and the reser- voir. Its utility may be further enhanced by providing the reser- voir with a soda-lime tube or some other suitable absorption tube.

Another form of apparatus is shown in Fig. 9. In this both the burette and the reservoir are provided with tubes containing soda-lime to insure a protection against the admission of CO2 and moisture from the air.

Pinch-cocks used with Mohr's burettes are of various kinds. See Figs. 10, n and 12.

The screw pinch-cock,* Fig. 12, is a very useful device; it may be used like the ordinary pinch-cock by pressure with the fingers upon a-a, when a rapid flow is desired, or the nut-screw (b) may be so adjusted as to FIG. 9. allow a slower flow or to deliver

the solution in drops, thus giving the operator the freedom of both hands for other work.

Burette supports are of various forms. One of the best

* W. v. Hergendorf.

34 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

for one or two burettes is shown in Fig. 13. It is made of iron, can stand firmly upon an uneven surface, and does not easily tip over. The burettes are fastened to it by means of clamps, illustrated in Figs. 14 and 15.

A revolving burette-holder for eight burettes is shown in

(6)

FIG. ii.

Fig. 1 6. Burrette-supports are also made with white porcelain base, which enables the operator more readily to see the change of color in the liquid titrated.

Pipettes are of two kinds — those which are marked to deliver one quantity only, and those which are graduated on the stem like burettes. Their use is to measure out portions of solutions with exactness.

APPARATUS USED IN VOLUMEjC TRANALYSIS 35

Pipettes are filled by applying the mouth to the upper end and sucking the liquid up to the mark, then, by closing the upper opening with the forefinger, the liquid is prevented from running out, but may be delivered in drops or allowed to run out to any mark by lessening the pressure of the finger over the opening.

In using the pipettes of the first class (Fig. 17) the finger

FIG. 14.

FIG. 13.

FIG. 15.

is raised and the instrument allowed to empty itself entirely. A drop or two, however, usually remains in the lower portion of the instrument, which may be blown out, though this is not considered good practice. By inclining the pipette and placing the point against the side of the vessel which is to receive the liquid, the instrument may be emptied more satisfactorily.

Pipettes of the second class (Fig. 18) are never emptied completely when in use. The flow of the liquid is regulated by the pressure of the finger over the upper opening, and stopped at the desired point.

36 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

A very convenient form of pipette is one which has attached to its upper end a piece of rubber tubing, into which a short piece of glass rod has been inserted. By squeezing the

FIG. 16.

rubber surrounding the glass rod firmly between the fingers, a canal is opened and the liquid can be drawn up into the pipette by suction with the lips and run out again. By re- moving the pressure the canal closes and the flow of the liquid is stopped at any point (Fig. 19).

APPARATUS USED IN VOLUMETRIC ANALYSIS 37

The Nipple Pipette is very convenient for measuring small quantities of liquids, such as i or 2 mils (Fig. 20).

When a volatile or highly poisonous solution is to be measured it is not advisable to suck it up with the mouth.

50 mils

10 mils

FIG. 17. FIG. 18. FIG. 19.

The pipette in this case is filled by dipping it into the liquid contained in a long, narrow vessel, until the liquid reaches the proper mark on the pipette, then closing the upper opening and withdrawing. When this is done the liquid which adheres to the outside of the pipette should be dried off before the measured liquid is delivered.

A French firm has introduced pipettes provided with suction

38

THE ESSENTIALS OF VOLUMETRIC ANALYSIS

pumps, shown in various forms by Fig. 21, which possess the advarltage over the ordinary forms provided with a com- pressible rubber bulb, that the liquid can with perfect ease be drawn up to the desired point on the scale, and with absolute accuracy maintained at the same height as long as may be desired.

The Measuring-flask is a vessel made of thin glass having

FIG. 20.

FIG. 21.

a narrow neck, and so constructed as to hold a definite amount of liquid when filled up to the mark on the neck. These flasks are of various sizes, holding 100, 250, 500 1000 mils, etc., but are generally called " Liter Flasks." (Fig. 22.)

Liter flasks are used for making volumetric solutions.

Those which have the mark below the middle of the neck are to be preferred, because the contents can be more easily shaken.

APPARATUS USED IN VOLUMETRIC ANALYSIS 39

Liter flasks are sometimes made with two marks on the neck very near together; the lower one is the liter mark. If the flask is required to deliver a liter, it must be filled to the upper mark, the difference between the two measures being

FIG. 22.

FIG. 23.

FIG. 24.

the equivalent of the liquid which remains in the flask adhering to the sides.

The Test-mixer, or Graduated Cylinder (Fig. 23) is for measuring and mixing smaller quantities of solutions. They are made of different sizes, holding TOO, 250, 500 and 1000 mils, and graduated in fifths or tenths of a mil.

Titration Flasks. Titrations may be carried out in flasks of any usual shape, or in beakers, or evaporating dishes, but the flask illustrated in Fig. 24 is to be preferred.

CHAPTER VI ON THE USE OF APPARATUS

IT is important that all apparatus used in volumetric analysis be perfectly clean. Even new apparatus should be cleansed by passing dilute hydrochloric acid through them and then rinsing with distilled water.

If the burette, pipette, or other instrument is even slightly greasy, the liquid will not flow smoothly, and drops of the liquid will remain adhering to the sides, thus leading to inaccurate results.

Greasiness may be removed with dilute soda solution. If this fails the instrument should be allowed to remain for some little time in a solution containing sulphuric acid and potassium dichromate, which will radically remove all traces of grease.

The burette or other measuring instruments should never be filled with volumetric solution without first rinsing, even if the burette be perfectly dry.

It is well to wash the inside of the instrument with two or three small portions of the solution with which it is to be filled.

The burette may be filled with the aid of a funnel, the stem of which should be placed against the inner wall of the burette, so that the solution will flow down the side and thus prevent the formation of bubbles.

The burette should be filled to above the zero mark, and the air-bubbles, if there are any, removed by gently tapping with the finger.

40

ON THE USE OF APPARATUS 41

A portion of the liquid should then be allowed to run out in a stream so that no air-bubbles remain in the lower part of the burette. In the glass tap burette it can be easily seen if any air is present, but in the pinch-cock burette it is some- times necessary to take hold of the rubber tube between the thumb and forefinger and gently stroke upward. Or the glass nip at the lower end of the burette may be pointed upward, and the pinch-cock opened wide so that a stream of the liquid will pass through and force out any air that may be inclosed.

If the titration is to be conducted at a high temperature, as in the estimation of carbonates, when litmus is used as the indicator, or in the estimation of sugar by copper solution, a long rubber tube should be attached to the lower end of the burette. The boiling can then be done at a little distance, and the expansion of the liquid in the burette avoided. The pinch-cock is fixed about midway on the tube.

Hart calls attention to the fact that if the fluid in a burette or pipette be run out rapidly at one time and slowly at another, different amounts of fluid are obtained.

This is due to the adhesion of the fluid to the inner sides of the instrument, and reading before it has settled down. It is therefore advisable always to deliver burettes slowly, as more constant results are then obtained.

Solutions which are measured by means of pipettes should be dilute, since concentrated solutions adhere to glass with different degrees of tenacity, and hence the amount of fluid delivered is slightly less than that measured.

The temperature of the solutions measured should be taken into account, since all liquids are affected by change of temperature, expanding and contracting as the tempera- ture is increased or reduced.

This change of volume in the case of standard solutions does not exactly correspond to that in pure water; in fact?

42 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

some of them differ widely. The correction of the volume of a standard solution for the temperature by the expansion coefficient of water is not entirely satisfactory, but in the case of very dilute solutions this may be done.

Casamajor (C. N., xxxv, 160) gives the following figures showing the relative contraction and expansion of water below and above 15° C.:

Degrees C. Degrees C.

8— .000590 1 7 + .000305

9-. 000550 18 + . 000473

10— .000492 19 + .000652

1 1 — .OOO42O 2O + .00084 1

12 — .000334 2 1 +.OOIO39

13— .000236 22 +.001 246

14— .OOOI24 23 + .001462

15 — normal 24 +.00 1686

By means of these numbers it is easy to calculate the volume of liquid at 15° C. corresponding to any volume observed at any temperature between 8° C. and 25° C. If 25 cc. of solution had been used at 20° C., the table shows that i cc. of water passing from 15° to 20° is increased to 1.000841 cc. Therefore, by dividing 25 cc. by 1.000841, the quotient, 24.97 cc. is obtained, which represents the volume at 15° C. corre- sponding to 25 cc. at 20° C.

These corrections are of value only for very dilute solutions and for water, but useless for concentrated solutions. Slight variations of atmospheric pressure may be disregarded.

ON THE READING OF INSTRUMENTS

In narrow vessels the surface of liquids is never level. This is owing to the capillary attraction exerted by the sides of the vessel upon the liquid, drawing the edge up and forming

ON THE USE OF APPARATUS

43

a saucerlike concavity (Fig. 25). All liquids present this concave surface except mercury, the surface of which is convex.

This behavior of liquids makes it difficult to find a distinct level, and in reading the measure either the upper meniscus (a) or the lower meniscus (b) may be used (Fig. 26).

The most satisfactory results are obtained if the lowest point of the curve (b) is used, especially with light-colored

FIG. 25.

FIG. 26.

FIG. 27.

solutions. But if dark-colored or opaque solutions are measured it is necessary to use the upper meniscus (a) for reading.

In all cases the eye should be brought on a level with the surface of the liquid in reading the graduation.

The eye is very much assisted by using a small card, the lower half of which is black and the upper half white, This card is held behind the burette, the dividing line between white and black being about an eighth of an inch below the surface of the liquid. The eye is then brought on a level with it, and the lower meniscus can be distinctly seen as a sharply defined black line against the white background (Fig.

27).

44 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

Erdmann's Float, Fig. 28, is an elongated glass bulb, which is weighted at its lower end with mercury, to keep it in an upright position when floating. It is of such diameter that

FIG. 28.

FIG. 29.

FIG. 30.

it will slide easily up and down inside of a burette. There is a ring at the top by which it can be lifted in or out by means of a bent wire. Around its center a line is marked. At this line instead of at the meniscus the reading is taken.

ON THE USE OF APPARATUS 45

These floats are sometimes provided with a thermometer, and they then register the temperature as well as the volume. Others are provided with projecting points along the sides,

FIG. 31. FIG. 32.

the object of which is to prevent it from adhering to the walls of the burette. (See Fig. 29.)

For the purpose of facilitating the reading, special forms of burettes are constructed which are provided with a dark longitudinal stripe on a white enameled background (Fig. 30); the reflection of the dark stripe with the meniscus pro- duces the peculiar appearance shown in Fig. 3 1 . The narrowest point is at the middle of the meniscus, and by reading from this point very accurate measurements are obtained, .The

46 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

same effect can be produced by holding behind an ordinary burette a white flexible card having a heavy black longitudinal stripe, about one-eighth inch in width.

Another form of burette designed for the purpose of facilitating reading is that provided with white enameled sides, leaving a strip of clear glass in front and back (Fig. 32). This form is especially adapted for use with dark- colored liquids, such as iodin and permanganate.

CALIBRATION OF INSTRUMENTS

Burettes are made from tubes of nearly uniform width. They are filled with distilled water at 15° C* (59° F.) to the o mark, and then 25, 50 or 100 cc. run out, and another mark made at the surface of the liquid. The distance between these two marks is then divided into 25, 50 or 100 equal parts, and the spaces again subdivided into fifths and tenths. Now it is very rarely possible to obtain tubes of exactly the same caliber throughout, and the divisions made as above do not always represent exactly what they are intended to.

If the tube is wider at one point the divisions at that point will contain more, and if it is narrower they will contain less than they should.

Hence before using a new burette, or in fact any other measuring instrument, it is essential that the error, if any, should be determined. This is done as follows:

Fill the burette to the o mark with distilled water at 15° C. (59° F.) and run out 10 cc. at a time into a small weighed flask, and weigh after each addition of 10 cc.

Each 10 cc. should weigh exactly 10 gms., and every

* Instead of 15° C. (59° F.) the temperature 25° C. (77° F.) is recommended because this more nearly approaches the ordinary temperature of the atmosphere in temperate climes.

ON THE USE OF APPARATUS

47

deviation found should be noted and taken into consideration in using the instrument.

Example

Flask weighed 20.0000 grams.

" + iocc. "" 30.1005 " " + 2occ. " 40.0499 " " +30 cc. " 49.8000 " " +40 ec. " 59-97CO " " + 500:. " 70.0100 "

Thus the ist 10. cc weighed 10.1005 grams.

2d 10 cc. 9-9494 "

3d 10 cc. 9-7501 "

4th 10 cc. 10.1700 "

5th 10 cc. 10.0400 "

Having obtained these data, a table like the following may be constructed and kept in some convenient place where it can be readily consulted whenever the burette it represents is being used. It is not necessary to carry the figure beyond the second decimal place.

No. of cc. as read on burette.

No. of cc. as corrected.

No. of cc. as read on burette.

No. of cc.

as corrected.

No. of cc. as read on burette.

No. of cc. as corrected.

I

1. 01

14

14.06

27

26.79

2

2. 02

IS

«5-°S

28

27.76

3

3-°3

16

16.04

29

28.73

4

4.04

17

«7-°3

3°

29.70

5

5-°5

18

18.02

31

30-71

6

6.06

19

19.01

32

3I-72

7

7.07

20

20.00

33

32.73

8

8.08

21

20.97

34

33-74

9

9.09

22

21.94

35

34-75

10

IO.IO

23

22.91

36

35-76

ii

11.09

24

23.88

37

36.77

12

12.08

25

24.85

38

37-78

J3

*3-°7

26

25.82

39

38.79

48 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

There should be no greater deviation than 0.15 cc. A burette which deviates more is best not used. In the foregoing table there is a deviation of 0.30 cc. at one point.

In order to test the accuracy of a pipette, fill to the mark with distilled water at 15° C. (59° F.); empty into a previously weighed flask, weigh again and thus determine the weight of the water measured. One gram is equal to i cc.

Liter flasks are tested as follows:

The flask, perfectly dry and clean, is counterpoised on a balance capable of turning with .005 when carrying about 2000 grams; it is then filled to the mark with distilled water at 15° C. (59° F.) and the increase in weight should be exactly the number of grams as the cc. indicated at the mark.

When very accurate determinations are required, various factors should be taken into account, namely, atmospheric pressure, the temperature of the air, and that of the water, which should be the same. Atmospheric humidity should be ot a definite degree, and the weights used should be made of specified material.

The calibration of a burette, which is graduated in mils, at 25° C., is conducted as follows:

Fill the burette to the o mark with distilled water at 25° C. The air should be half saturated with moisture, and of the same temperature as the water. The barometric pressure should be 760 mm. The weights used should be of brass, and the counterpoise of the flask should be of the same kind of glass as that of which the flask is made. Run out 10 mils at a time into the small counterpoised flask, and weigh each 10 mils.

Each 10 mils should weigh exactly 9.9604 gms., and every deviation found should be noted and taken into consideration in using the burette.

ON THE USE OF APPARATUS

49

Example ist 10 mils weighed 9.9610 gms. (deviation negligible)

9.8972 "

10

3d I0 4th 10 5th 10

9-9543 10.0864

9-9753

From these data a table like the following may be con- structed, and consulted whenever the burette it represents is being used.

No. of Mils as Read on Burette.

No. of Mils as Corrected.

No. of Mils as Read on Burette.

No. of Mils as Corrected.

No. of Mils as Read on Burette.

No. of Mils as Corrected.

I

I

18

17-944

35

34.983

2

2

19

18.937

36

35-995

3

3

2O

19.930

37

37-ooS

4

4

21

20.929

38

38.020

5

5

22

21.928

39

39-033

6

6

23

22.927

40

40 . 046

7

7

24

23.926

4i

41 . 046

8

8

25

24.925

42

42 . 046

9

9

26

25.924

43

43 • 046

10

10

27

26.923

44

44-046

ii

10.993

28

27.922

45

45 • 046

12

11.986

29

28.921

46

46 . 046

13

12.979

30

29.920

47

47-047

U

13.972

31

30.932

48

48.047

15

14-965

32

31-945

49

49.047

16

iS.958

33

32.957

So

50.047

17

16.951

34

33-970

CHAPTER VII

METHODS OF CALCULATING RESULTS

N EACH mil of a — univalent volumetric solution contains

of the molecular weight in grams of its reagent, and will neutralize y^Vfr of the molecular weight of a univalent substance, or -^-^ of the molecular weight of a bivalent substance.

N Each mil of a — bivalent volumetric solution contains

•oVfr of the molecular weight in grams of its reagent, and will neutralize or combine with innnr of the molecular weight of a bivalent salt, or ^^ of the molecular weight of a univalent salt.

N A — is only ^V the strength of a normal solution and will

neutralize only ^ the quantity of salt, etc.

Normal and decinormal solutions of acids should neutralize normal and decinormal solutions of alkalies, volume for vol- ume. Decinormal solution of silver nitrate and decinormal solution of hydrochloric acid or sodium chlorid should combine, volume for volume, etc.

Rules for Direct Percentage Estimations: i. With normal solutions TV or ^ (according to its atomicity) of the molec- ular weight in grams of the substance is weighed for titration, and the number of mils of the V.S. required to produce the desired reaction is the percentage of the substance whose molecular weight has been used,

50

METHODS OF CALCULATING RESULTS 51

Thus, if sodium hydroxid (NaOH) is to be examined by titration with a normal acid solution TV of its molecular weight in grams, 4 gms. is weighed out, and each mil of the acid solution required represents one per cent of the pure salt.

If sodium carbonate (Na2CO3) is to be titrated jV of its molecular weight in grams, 5.3 gms. is taken.

2. With decinormal solutions T^ or yj-fr of the molecular weight in grams of the substance to be analyzed is taken, and the number of mils will, in like manner, give the percentage.

The following equations will serve to explain more fully:

TV Sodium hydroxid with — sulphuric acid:

2NaOH + H2S04 = Na2SO4 + 2H2O.

2X40=80 2)98

10)40 4g = looo mils

4.0 gms. = 100 mils

N Sodium carbonate with — sulphuric acid:

Na2CO3 + H2SO4 = Na2SO4 + H2O + CO2.

2)98

2o)io6__ 49 _ I000 miis

5.3 gms. = 100 mils

N

With — sulphuric acid: 10

2NaOH + H2SO4 = Na2SO4 + 2H2O.

2X40=80 2)98

100)40 49 _ I000 miis

0.40 gm. = 100 mils

In the case of a trivalent substance as citric acid fa of the molecular weight in grams is taken for analysis when a normal solution is employed and ^fo when a decinormal solution is used.

In other words, when it is desired that each mil of the

52 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

standard solution should represent one per cent of the substance upon which it acts, the rule is to take for analysis as much of the substance as is represented by 100 mils of the standard solution.

In the case of substances whose percentage of purity is high, it is advisable to take smaller quantities, in order to avoid the use of excessive quantities of standard solution. Thus sulphuric acid, which contains 92.5 per cent of absolute sulphuric acid, would require under the above conditions 92.5 mils of normal alkali solution.

In the case of this acid, if 4.9 gms. are taken for analysis, each mil of a normal alkali solution would represent one per cent of H2SO4.

If half of this quantity, i.e., 2.45 gms. are taken for analysis, each mil of the normal alkali will represent two per cent of H2SO4, and thus less of the standard solution will be required. Again, if 0.49 gm. be taken, each mil of the standard alkali will represent 10 per cent of E^SO^

In the case of liquids where it is not always convenient to weigh off the exact quantity required for titration by the direct percentage method, the liquid is diluted to a convenient degree with water, and then a quantity of this diluted liquid (representing the weight required of the substance) is measured for analysis.

Example. A sulphuric acid solution of specific gravity 1.826 is to be analyzed. Two mils are accurately measured and diluted to 100 mils and then 13.41 mils of this solution (representing 0.49 gm. of the acid) are taken for analysis.

N Each mil of — NaOH V.S. required in the titration rep-

N

resents 10 per cent of absolute H2SO4. If — NaOH V S

10

is employed, each mil will represent one per cent To de-

METHODS OF CALCULATING RESULTS 53

termine the amount of the diluted liquid to be taken we proceed as follows :

Two mils of sulphuric acid, specific gravity 1.826, weigh 3.652 gms., therefore the 100 mils of diluted acid contain this weight, and i mil of the same contains 0.03652 gm.

If 0.03652 gm. are contained in i mil, then 0.49 gm. are contained in how many mils ?

gm. mil gm.

.03652 : 1:10.49 : x,

#=13.41 mils.

Factors or Coefficients for Calculating the Analyses. It frequently occurs that from the nature of the substance, or from its being in solution, this percentage method cannot be conveniently followed.

The best way to proceed in such a case is to find the factor.

The first step in all cases is to write the equation for the reaction which takes place between the substance under analysis and the solution used.

For instance, a solution of caustic potash is to be examined,

N a — solution of sulphuric acid being used.

2KOH + H2S04 = K2S04 + 2H20.

2)H2 2)98 N

56 49 = 1000 mils — acid.

-, N 0.56 gm. .049 = i mil — acid.

N The factor for KOH when — solution is used is .056 gm.,

N that being the quantity neutralized by each mil of the — acid.

N If — acid were*, used the factor would be .0056 gm.

54 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

The number of mils of the acid used to produce the desired result, when multiplied by the factor, gives the quantity in grams of KOH in the solution taken.

Example. If 10 grams of caustic potash solution were

N taken, and 40 mils of — acid were required, the 10 gms. of

solution contained .056 gm.X4o = 2.24 gms. of pure KOH.

To find the percentage the following formula may be used:

W

Q =the quantity of pure substance found by calculation; W = weight of substance taken.

If the above example is taken, we have

2.24100

*~ =22-4%-

Or the calculation may be made by proportion.

The quantity of the substance taken is always the first term, and the quantity of pure substance found, the second term.

The following rule is easily remembered: As the quantity taken is to the quantity found, so is 100 to x, the percentage of pure substance in the sample.

Three terms of the equation being given, the fourth is found by multiplying the means and dividing the product by the given extreme. By applying this rule to the above case we have

10 : 2.24 : : 100 : x. 00=22.4%.

METHODS OF CALCULATING RESULTS

55

TABLE SHOWING THE NORMAL FACTORS, ETC., OF THE ALKALIES, ALKALI EARTHS, AND ACIUS.

Substance.

Formula.

Molecular Weight.

Normal Factor.*

Quantity of Substance to be taken t so that each

mil of ^V.S.

will indicate i per cent.

Sodium hydroxid

NaOH

4O

o 040

Grams 4 o

Sodium carbonate

Na2CO3

106

o.os ;

C.2

Sodium bicarbonate . .

NaHCO3

84

o 084

8 4

Potassium hydroxid

KOH

^6 i

o 0561

•> 61

Potassium carbonate

K2CO3

I3S.2

o .0691

6.01

Potassium bicarbonate

KHCO3

IOO I

O IOO I

IO OI

Ammonia (gas)

NH3

17 .03

o . 01703

i .703

Ammonium carbonate, normal Lime

(NH4)2C03 CaO

96.08

<6 i

0.04804 o 02805

4.804

2 805

Calcium hydroxid

Ca(OH)2

74. 1

0.0370?

3 .701?

Calcium carbonate

CaCO3

IOO I

o o^o

e

Nitric acid

HNO3

63 01

o 063

6.3

Hydrochloric arid

HC1

36.46

0.03646

3 .646

Sulphuric acid

H2SO4

98 08

o 04004

4Q

Oxalic acid, crystallized Acetic acid

H2C2O4-2H2O HC2H3O2

1 26 . 05

60 .03

0.063

o . 0603

6-3

6 .03

* This is the coefficient by which the number of mils of normal solution used is to be multiplied in order to obtain the quantity of pure substance present in the material examined.

t This is the quantity of substance to be taken in direct percentage estimations.

N Each mil of — acid or alkali V.S. employed will then indicate i per cent; in the case

of many of these substances it will, however, be better to take smaller quantities so that less of the standard solution be required. Thus if one-half the quantity be

N taken each mil of the — V.S. will represent 2 per cent, if one-tenth of the quantity be

taken each mil will represent 10 per cent. If, however, — solutions be used and one- tenth of the quantity indicated in the table be taken each mil will indicate i per cent.

On Stating Results. In reporting the results of volu- metric work, it is customary to state the quantity of pure substance found; thus in the case of salts, the quantity of the anhydrous salt is reported. It is, however, often required

56 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

to state the results according to the dualistic formulae of Berzelius, that is, the metals are reported as oxids and the acids as anhydrids. Thus if sodium carbonate is analyzed, a statement of results by this method will give the quantity of Na2O, spoken of as the base, soda. If we look upon sodium carbonate as Na2OCO2, we can readily see 106 gms. of the anhydrous salt contain 62 gms. of Na2O.

By reference to the following equations we see that 98.08 gms. of sulphuric acid will neutralize 62 gms. of Na2O or 106 gms. of Na2OCO2.

Na2O + H2SO4 = Na2SO4 + H2O.

2)62 2)98.08

31 gms. 49.04 gms. =1000 mils — V.S.

Na2OC02+H2S04 = Na2S04+H20 + C02.

2)^o6_ N

53 gms. = to 1000 mils — V.S.

N Thus one mil of — H2SO4 will represent 0.031 gm. of

Na2O and 0.053 gm. of Na2OCO2.

In the case of sodium bicarbonate (NaHCO3) two molecules contain one molecule of the base, soda, as here shown.

2NaHC03 = Na20,H20(C02)2.

According to this 62 gms. of Na2O represent two molecular weights (168 gms.) of NaHCOs. In the case of ferrous sul- phate, one molecule (FeSO4) contains, according to this system, FeO and SO3. In the same way, ferric salts contain Fe2O3. In stating the results of analyses of acids according to this system, the quantity of acid anhydrid found is reported, not the quantity of the whole acid. Thus if sulphuric acid is analyzed, the quantity of SO3 is reported. In the case of phosphoric acid the quantity of P2Os is stated, etc.

METHODS OF CALCULATING RESULTS

57

TABLE SHOWING THE MOLECULAR WEIGHTS AND NORMAL FACTORS FOR THE MOST COMMON OXIDS.

Name.

Formula.

Molecular Weight.*

^ Factor.

Soda

Na2O K2O CaO MgO Li2O SrO BaO ZnO PbO As203 Sb,O, Hg20

FeO Fe203 Ag20 S03

N2055 C02

62.0 94.2

40.32 29.88 103.63

J53-37 81.37 223.1 197.92 288.4 416.0 2 16.0 71.82 159.64 231.76 80.07 142.08 108.02 44.0

0.031 gm. 0.0471 " 0.028 "

O.O2O "

0.0149 "

0.0518 "

0.0767 " 0.0407 " 0.1115 "

0.0495 "

0.0721 "

0.208 "

o . 108 " 0.0718 " 0.0798 " 0.1159 " 0.040 " 0.02368 " 0.054

0.022 "

Potash

Lime

M^agnesia

Lithium oxid .

Strontium oxid

Barium oxid

Zinc oxid

Lead oxid ...

Arsenous oxid

Antimonous oxid

M^ercurous oxid

M^ercuric oxid

Ferrous oxid

Ferric oxid

Silver oxid .

Sulphuric anhydrid

Phosphoric anhvdrid

Nitric anhvdrid

Carbonic anhydrid . ...

* Approximate.

CHAPTER ANALYSIS BY NEUTRALIZATION

THIS is based upon the fact that when an acid and an alkali react each loses its individuality and a neutral salt is formed, i.e., a body which has neither the character of an

acid nor that of an alkali. This mutual neutralization of

+ acid and alkali is the result of a union of the H* ions of the

acid and the OH' ions of the alkali, forming non-ionized water (HOH).

An acid is a compound which in aqueous solution disso- ciates (ionizes) into positive and negative ions, the positive

ion being H. Thus hydrochloric acid in an ionized state is

+ + +

H- + Ch Sulphuric acid ionizes into H' + H' + (SO4) .

An alkali is a basic compound which ionizes into positive and negative ions, and in which the negative ion is (OH).

The reaction between hydrochloric acid and potassium hydroxid. in accordance with this theory, is illustrated by the following equation:

An acid is generally recognized as such by its color reac- tions with certain substances known as indicators; for example, it turns blue litmus red, and decolorizes a red phenolphthalein solution. Alkalies are recognized by their turning red litmus

58

ANALYSIS BY NEUTRALIZATION 59

blue, and by producing a deep red color with phenol- phthalein.

The strength of an acid solution is ascertained by noting the quantity of alkali that is required to neutralize it. The stronger the acid, the more alkali is required. The strength of an alkali is estimated by the quantity of acid which is required to neutralize it. The estimation of the strength of acids is called acidimetry, while the estimation of alkalies is called alkalimetry.

The principal alkaline substances which may be estimated by means of standard acid solutions are the hydroxids and carbonates of sodium, potassium, lithium and ammonium, and the hydroxids and oxids of calcium, barium and strontium and the alkaloids.

When an acid is brought in contact with an alkali, a reaction takes place in which a neutral salt is formed. This is known as neutralization, and takes place between definite and invariable proportions of the reacting bodies; thus, if 1 12. 2 parts of potassium hydroxid are mixed with 98.08 parts of absolute sulphuric acid, the alkali as well as the acid will be exactly neutralized. If only 80 parts of the acid have been added the mixture would still be alkaline, for it requires 98.08 parts of the acid to neutralize 112.2 parts of potassium hydroxid. If more than 98.08 parts of the acid have been added, the mixture would be acid, and would consist of potassium sulphate and free sulphuric acid. The reaction is thus illustrated:

2KOH + H2SO4 = K2SO4 + 2H2O.

2K = 78.2 2H= 2.016

20 = 32.0 8=32.070

2H= 2.016 40 = 64.00

112.216 98.086

Sodium hydroxid will unite with oxalic acid in the proper-

60 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

tion of 80.016 parts by weight of the former and 126.048 parts by weight of the latter, as the equation shows.

2NaOH + H2C2O4 • 2H2O = Na2C2O4 + 4H2O.

6 6H'= 6.048

20 = 32 20 = 24.000

2H'= 2.016 60 = 96.000

80.016 126.048

Ammonia water unites with hydrochloric acid as per the equation,

NH4OH + HC1 = NH4C1 + H2O.

35.05 36.46

Sodium carbonate with hydrochloric acid,

Na2C03 + 2HC1 = 2NaCl + H2O + CO2.

106 72.92

Upon a careful perusal of the foregoing equations it will be seen that since definite weights of acids neutralize definite weights of alkalies, the quantity of a certain alkali in solution can be easily determined by the quantity of an acid solution of known strength required to neutralize it, and vice versa.

Referring to the first equation we see that 98.086 gms. of H2SO4 neutralize 112.216 gms. of KOH. If we prepare a normal solution of H2SO4 we take half the molecular weight, 98.086 = 49.043 gms., to 1000 mils. Half the molecular weight is taken because sulphuric acid is a bivalent acid. 1000 mils of this solution will neutralize 56.108 gms. of KOH; hence i mil will neutralize 0.056108 gm. of KOH.

Thus if 10 gms. of a solution of KOH be treated with the above normal solution of H2SO4, and it is found that 25 mils of the acid solution are required to neutralize the alkali solution, the latter contains 25X0.0561=1.40+ gm. of pure KOH.

Since the acid and alkali as well as the neutral salt which

ANALYSIS BY NEUTRALIZATION 61

is formed are colorless, and no visible change takes place during the reaction, it is necessary to add some substance which by change of color will show when the neutralization is complete. Such a substance is known as an indicator.

In the case of sodium hydroxid with oxalic acid (see the second equation) we find that 126.048 gms. of crystallized oxalic acid neutralizes 80.016 gms. of NaOH. Oxalic acid, like sulphuric, is bivalent, therefore a normal solution of it contains half the molecular weight in grams,, i.e., 63.024 gms. in 1000 mils.

1000 mils will neutralize 40 gms. of NaOH; i mil will neutralize 0.040 gm. of NaOH.

The neutralizing power of all normal acids is exactly the same, because they all contain in 1000 mils the molecular weight in grams of the acid in the case of univalent acids, and half of the molecular wieght in grams of bivalent acids.

Thus i mil of any normal acid will neutralize 0.0561 gm. of KOH or 0.040 gm. of NaOH or x^Vo of the molecular weight of any other univalent alkali, or 2J00 of the molecular weight of an alkali earth, the latter being bivalent. In like manner all decinormal solutions have a like neutralizing power, their neutralizing equivalence is one-tenth that of normal solutions.

Thus i mil of a decinormal acid will neutralize 0.00561 gm. of KOH or 0.0040 gm. of NaOH, etc.

Alkalimetry

Preparation of Standard Acid Solutions. It is possible to carry out the titration of most alkalies by means of one standard acid solution, but the same standard acid is not equally applicable in all cases; furthermore, the standard

62 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

acids are frequently employed for other volumetric operations than neutralization, and therefore it is advisable to have a variety.

The standard oxalic acid solution is preferred by some, because of the ease with which it may be prepared, provided a pure oxalic acid is to hand. It does not, however, keep very long, is unreliable for use with methyl orange, and is inapplicable for the titration of alkali earths, because it forms insoluble compounds with these metals. Standard hydrochloric acid is the most desirable for alkali earths, because it forms soluble compounds with them; its disadvantage, however, is in its volatility and its consequent uselessness in hot titrations. Standard sulphuric acid is preferred by most analysts as being the best general standard. A pure acid can be gotten without difficulty and the standard solution made from it is unaffected by boiling, and can therefore be used in hot as well as in cold titrations; it reacts sharply with the indicators and it keeps its titer indefinitely. It is, however, not suited for the titration of alkali earths, because it forms insoluble compounds with them, which precipitate and are very annoying to the operator. Iji the preparation of standard solutions the greatest care should be exercised in order that the product be absolutely accurate. The slightest inaccuracy in the strength of a standard solution will result in relative errors in the analysis. It is customary to prepare one standard solution, and then from this to adjust various others. For example, a normal oxalic acid may be made first, and by means of this a normal alkali solution, which in turn may be utilized for the adjusting of other standard acid solu- tions.

Normal Oxalic Acid V.S. (H2C2O4-2H2O - 126.05, 7- V.S. = 63.025 gms. in 1000 mils). Dissolve 63.025 gms. of purified

ANALYSIS BY NEUTRALIZATION 63

oxalic acid* in enough distilled water to make, at or near 25° C, exactly 1000 mils.

/N Tenth-Normal Oxalic Acid V.S. I— V.S.= 6.3025 gms.

in 1000 mils.)

Dissolve 6.3025 gms. of purified oxalic acid in sufficient distilled water to make, at or near 25° C., exactly 1000 mils of solution. Or, better: dissolve 6.45 gms. of pure crystal- lized oxalic acid in sufficient distilled water to measure 1000 mils. Then into a flask accurately measure 25 mils of a freshly standardized tenth-normal potassium hydroxide V.S., dilute with an equal volume of distilled water, add 3 to 5 drops of phenolphthalein T.S. and heat it to boiling. From a burette gradually add the oxalic acid solution (which is too concen- trated) until the red tint of the alkali solution fails to re- appear after vigorous shaking and boiling. Note the num- ber of mils of the oxalic acid solution consumed, and then dilute it so that equal volumes of this and of the tenth- normal potassium hydroxid V.S. neutralize each other at standard temperature (25° C.). This solution deteriorates on standing, hence must be frequently renewed or restand- ardized.

This solution is in every respect equivalent in neutralizing power to any other tenth-normal acid V.S. with either litmus or phenolphthalein T.S. as indicator, but not with methyl orange. Its specific use in the U. S. P., however, is in stand- ardizing or determining excess of tenth-normal potassium permanganate V.S.

* Purified oxalic acid is in the form of colorless transparent, unefHoresced clinorhombic crystals, soluble in 10 parts of cold water, in about 3 parts of boiling water and in 2.5 parts of alcohol. Ten gms. when ignited upon plati- num foil leaves not more than 0.005 gm. of residue. If more residue is left the acid should be purified by recrystallization.

64 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

Normal Hydrochloric Acid V.S. (HC1 = 36.47; V.S.

= 36.47 gms. in 1000 mils). Mix no mils of hydrochloric acid of sp.gr. 1.155 witn enough water to measure, at or near 25° C., 1000 mils.

FIG. 33-

Of this liquid (which is still too concentrated) measure carefully, at 25° C., into a flask or beaker 10 mils, add 20 mils of distilled water and a few drops of methyl orange T.S., and then gradually add from a burette sufficient recently pre-

N pared and standardized — potassium or sodium hydroxid

ANALYSIS BY NEUTRALIZATION 65

at the same temperature to just produce a permanent faint

yellow tint.

N Note the number of mils of — alkali solution consumed

and then dilute the acid solution so that equal volumes of

N it and — alkali neutralize each other. It is usually advisable

to make two or three titrations, as just described, before dilution, taking an average of the results.

Example. Assuming that the 10 mils of the acid solution

N required 12 mils of the — alakli, each 10 mils of the acid

must be diluted to 12 mils, or the whole of the remaining acid in the same proportion.

After the dilution a new trial should be made. 10 mils of the acid V.S. should required exactly 10 mils of the alkali.

This method is fairly satisfactory if an accurately stand- ardized normal alkali hydroxid solution is at hand; the latter, however, always contains a small qunatity of carbonate, hence methyl orange is most desirable as an indicator.

Standardization by Means of Sodium Carbonate. Pure anhydrous sodium carbonate may be obtained by heating to dull redness a few grams of pure sodium bicarbonate for about thirty minutes. The resulting carbonate is practically free from impurity.

The sodium bicarbonate loses on ignition one-half of its carbonic acid gas :

2NaHCO3 + Heat = Na2CO3 + CO2 + H2O.

The bicarbonate should, however, be tested before igniting, and if more than traces of chlorid, sulphate, or thiosulphate, are found, these may be removed by washing a few hundred grams, first with a saturated solution of sodium bicarbonate, and afterward with distilled water.

66 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

0.53 gm. of the pure anhydrous sodium carbonate is accurately weighed and dissolved in about 200 mils of dis- tilled water in a flask and a few drops of methyl orange T.S. added as indicator. The acid to be " set " of " standardized " is then run into the sodium carbonate solution until a per- manent light-red color is produced. It should require exactly

N

10 mils of the — acid solution, i

If 8 mils of the acid solution are consumed to bring about the required result, then every 8 mils must be diluted to 10 mils, or the whole of the remaining solution must be diluted in this proportion :

Na2C03 + 2HC1 = 2NaCl + H2O + CO2. 2)106 2)72.9

53 gms. 36.45 = 1000 mils — V.S.

0.53 gm. = 10 mils I

Instead of attempting to weigh exactly 0.53 gm. of the an- hydrous sodium carbonate, it is better to take a larger quantity (about 2 gms.) weighed accurately. Dissolve this in 100 mils of distilled water, add 2 or 3 drops of methyl orange T.S. and run into it, little by little, from a burette the acid solution to be standardized, reducing the flow to drops toward the end until a pink color is obtained. Make 2 or 3 trials, and take the average number of mils consumed.

One gm. of pure anhydrous sodium carbonate requires for exact neutralization 18.868 mils of normal acid V.S.

Assuming that 2.3 gms. of the anhydrous carbonate were taken, and this required 38.2 mils of the trial acid. Then

2.3 X 18.868 = 43.3964 mils. 43.3964 : 38.21:100 : x. #=88.26.

Now adjust the acid by measuring 882.6 mils and dilut- ing with distilled water to make 1000 mils.

ANALYSIS BY NEUTRALIZATION 67

This method may be employed as well for the standardization of sulphuric or oxalic acid.

Other Methods for standardizing hydrochloric acid V.S. are: (a) by means of silver nitrate (gravimetrically and volu- metrically); (b) by means of borax; (c) by means of the specific gravity; (d) by means of calc-spar.

N Normal Sulphuric Acid V.S. (H2SO4 = 98.09; - V.S.=

49.045 gms. in 1000 mils). Mix carefully 30 mils of pure concentrated sulphuric acid (sp.gr. 1.835) w^h enough water to make about 1050 mils, and allow the liquid to cool to about 25° C.

Titrate 10 mils of this liquid in the manner described under

N

- hydrochloric acid, and dilute it so that equal volumes of

the acid and the alkali will neutralize each other.

The standardization of the normal sulphuric acid solution may also be effected by the use of pure anhydrous sodium carbonate, as described under normal hydrochloric acid V.S., and by various other methods, among which are: (a) the iodometric; (b) the specific gravity method; (c) the borax method; (d) by precipitation with barium chlorid (gravi- metrically) .

Standard acid solutions are used in other strengths besides

N N

normal, namely, Half -normal — , Fifth-normal — , Tenth-nor-

N N N

mal — , Twentieth-normal — , Fiftieth-normal — , and Hundredth- 10' 20' 50'

7 N

normal .

100

Estimation of Alkali Hydroxids

Potassium and sodium hydroxids are usually titrated with

N

- sulphuric or hydrochloric acid; they are, however, so prone

to absorb carbon dioxid out of the air that they are seldom free from carbonate, and hence the selection of aai indicator

68 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

is a matter of some importance. Phenolphthalein or litmus may be employed, but it is then advisable to boil the solu- tion while titrating, in order to drive off the liberated carbon dioxid, because the latter gives an acid reaction with phenol- phthalein and litmus and thus causes an end-reaction tint to appear before neutralization is complete. It is better, usually, to employ an indicator which is not affected by carbon dioxid. Methyl orange is mostly preferred; cochineal and Congo red are also useful. These indicators are especially serviceable in the presence of carbonates in that they are not affected by carbon dioxid, and can therefore be employed in direct titra- tions without the use of heat.

The quantity of carbonate in a recent sample of sodium or potassium hydroxid is so small usually that it is customary to disregard it and to report the total alkalinity as hydroxid.

A definite quantity of the sample (from 0.5 gm. to i gm. of the solid or an equivalent of a solution) is taken for analysis, dissolved in 36 to 50 mils of water in a white porcelain dish or a beaker placed over a white surface, and a few drops of a suitable indicator added.

The vessel is then placed beneath a burette containing the standard acid solution and the latter run in, drop by drop, until the last drop just causes the color to change. The solution should be rotated or stirred after each addition of the standard acid.

The alkali hydroxids are so exceedingly hygroscopic that they take up water from the air while being weighed; it is therefore difficult to make a direct weighing with any degree of accuracy.

The best way is to take- a small piece of the sample (about i gm.), place it immediately in a tared stoppered flask and take the weight accurately. It is then dissolved in water, transferred to the porcelain dish or beaker and titrated.

ANALYSIS BY NEUTRALIZATION 69

Potassium Hydroxid (KO H = 5 6 . i ) . An accurately weighed portion (preferably less than i gm.), is placed in a small beaker, dissolved in 50 mils of water, three drops of methyl orange

N added, and the titration begun with — sulphuric acid and

continued until the yellow color of the solution is changed to pale red. Then the burette is carefully read to see how much of the acid solution was used. The number of mils of the latter are multiplied by the normal factor for KOH (0.0561) and the result is the quantity of pure KOH in the sample taken for analysis.

The following equation illustrates the reaction:

2KOH + H2SO4 = K2SO4 + 2H2O. 2)112.2 2)98.09 N

56.1 gms. 49.045 gms., quantity in 1000 mils of — acid V.S. 0.0561 gm. (the factor for KOH), quantity neutralized by

•i *N i mil of — acid.

i

N Thus 1000 mils of — H2SO4 V.S. containing 49.045 gms,

of absolute H2SO4 will neutralize 56.1 gms. of KOH. There.

N fore each mil of — H2SO4 V.S. will neutralize 0.0561 gm. oj

pure KOH.

Example. In the above analysis let it be assumed that 0.915 gm. of potassium hydroxid were taken and that 15.3 mils of the standard acid were required to neutralize it, then 0.0561 gm.Xi5.3 = o.8583 gm., the quantity of pure KOH in the 0.915 gm. taken.

The percentage is then calculated in this way:

0.915 : 0.85831:100 : x] #=93.8+ per cent.

0.8583X100

=03.8.

0.915 ^

70 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

The fixed alakli hydroxids are rarely free from carbonate, through absorption of carbon dioxid from the air.

In the foregoing method, the carbonate is calculated as hydroxid, hence the assay is fallacious. The method of the U. S. P. IX separates the carbonate by precipitating with barium chlorid, filtering, and titrating the filtrate with normal hydro- chloric acid V.S., using phenolphthalein as indicator. The use of standard sulphuric acid V.S. is inadmissible in this method because it reacts with barium chlorid, producing a precipitate of barium sulphate.

The U. S. P. Method is as follows:

Take about 10 gms. of potassium hydroxid, accurately weighed in a glass-stoppered weighing bottle, dissolve in 250 mils of distilled water, which has been previously boiled and cooled (to expel carbon dioxid), add 30 mils of barium chlorid T.S. and then dilute with distilled water to 500 mils. Thor- oughly agitate the liquid and filter it through a dry filter into a dry flask. Reject the first 20 mils, and titrate 200 mils of the clear filtrate with normal hydrochloric acid V.S., using phenolphthalein T.S. as indicator.

Sodium Hydroxid (NaOH = 4o). This is estimated in exactly the same manner as described for potassium hydroxid, the following equation being applied:

2NaOH + H2SO4 = Na2SO4 + 2H2O.

2)80 2)98.09 N

40 gms. 49.045 gms. = looo mils — V.S.

N .040 gm. i mil — V.S.

The factor.

The official solutions of potassium and of sodium hydroxid are estimated in this same manner, 10 mils may be taken for analysis, diluted with 20 mils of water.

Ammonia Water (NH3-H2O). Three mils of ammonia water are put into a stoppered weighing bottle and the weight

ANALYSIS BY NEUTRALIZATION 71

taken. Forty mils of water are then added and the solution

N titrated with — sulphuric acid. As indicator, litmus, methyl

orange or rosolic acid may be used. Phenolphthalein is use- less for titrating ammonia, and even methyl orange and rosolic acid are unsuitable in the presence of much salts of ammo- nium. Because of the volatile character of ammonia its solu- tions readily lose strength upon exposure. It is therefore best to measure a quantity into a weighing bottle and find its weight as directed for potassium hydroxid. If the specific gravity of the ammonia solution is known, the weight of a given volume is easily calculated, it being only necessary to multiply the volume in mils by the sp.gr. Thus, if the sp.gr. of an ammonia solution is 0.9585 and the volume taken is 3 mils, the weight of the 3 mils is 3 Xo.Q585 = 2.8755 Sms-

N In the titration with — sulphuric acid each mil of the

latter represents 0.017 gm. of NH3, as shown by the equation

2NH3 • H2O + H2SQ4 = (NH4) 2SO4 + 2H2O. 2)34 2)98-09 N

17 gms. 49-045 gms. = iooo mils — V.S.

.017 = i mil -V.S.

Factor.

N If 16.9 mils of — acid were required in the above assay,

then 0.017 gm.X 16.9 = 0.2873 gm., the quantity of pure NH3 in the 3 mils (2.8755 gmsO °f ammonia water taken.

The percentage is found as follows :

If 3 mils of ammonia water weighing 2.8755 gms. contain 0.2873 gm. of NH3, 100 gms. of ammonia water will contain Xgm. of NH3,

0.2873X100

= 9-99/ per cent.-

2.8755

72 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

Stronger ammonia water and spirit of ammonia may be estimated in the same manner.

Estimation of Alkali Carbonates

When carbonates are treated with acids carbonic-acid gas is liberated. This gas shows an acid reaction with most indicators, and the reaction will seem to be completed before the alkali is entirely neutralized.

To avoid this, the titration may be conducted at the boiling temperature (hot way) in order to drive off the carbon dioxid. The standard acid being added until two minutes' boiling fails to restore the color indicating alkalinity. If the titration is conducted at a boiling temperature, it is advisable to attach to the lower end of the burette a long rubber tube with a pinch -cock fixed about midway on the tube.

The boiling can then be done at a little distance from the burette and the expansion of the standard solution therein thus prevented.

Another method is to add to the carbonate a measured excess of the standard acid, and then after boiling to drive off the carbon dioxid, an indicator is added, and the excess of standard acid determined by titration with a standard alkali (residual titration way). The quantity of the latter, deducted from the quantity of the standard acid taken, gives the quantity of the latter which reacted with the carbonate. Still another method is to titrate the carbonate direct, without heat (cold way), using an indicator which is not affected by caibon dioxid. The best of the indicators which are not so affected is methyl orange; others are cochineal and Congo red. When employing methyl orange as an indicator standard oxalic acid solution should not be used, as the end-reaction is very indefinite and unreliable.

The end-reaction with this indicator is at all events not a clearly marked one, and considerable practice and an eye for color is required to detect the point at which yellow changes ft?

ANALYSIS BY NEUTRALIZATION

73

pale pink. It is a good plan to have on the bench two vials, one containing an acid and the other an alkali tinted with methyl orange, with which comparisons can be made.

FIG. 34

Potassium Carbonate (K2COa= 138.2). Weigh carefully one gram of the salt, previously dried to constant weight at 180° C. Dissolve in 25 mils of distilled water in a beaker or flask, add a few drops of methyl orange T.S., and titrate with normal sulphuric acid until a faint orange-red color appears.

K2CO3 + H2SO4 = K2SO4 + H2O + CO*

2)138.2 2)98.09 N

69.1 gms. 49.045 gms. = iooo mils — V.S.

N Each mil of — H2SO4, therefore, reoresents 0.0691 gm.

of pure potassium carbonate.

If 14.3 mils of the normal acid are required the salt contains

74 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

14.3X0.0691 gm. = 0.98813 gm. of pure K2CO3 or 98.813 per cent. If it is desired to use litmus or phenolphthalein, it will be necessary to boil the solution as described above.

Other alkali carbonates are estimated in exactly the same manner as described for potassium carbonate.

Potassium Bicarbonate (KHCO3 = 100.1).

2KHC03 + H2S04 - K2S04 + 2H2O + CO2.

2)200.2 2)98.09 N

100.1 gms. 49.045 gms. = iooo mils — V.S.

N Each mil of — acid V.S. = o.iooi gm. of KHCO3.

Previous to weighing, this salt should be dried to constant weight in a desiccator over sulphuric acid.

Sodium Carbonate (crystallized) (Na2CO3- ioH2O = 286.16).

Na2CO3-ioH2O + H2SO4 = Na2SO4 + nH2O + CO2.

2)286.l6 2)98.O9 -jy,

143.08 gms. 49-045 gms. = looo mils — V.S.

N

Each mil of — acid = 0.143 gm. crystallized sodium car- bonate.

Sodium Carbonate (anhydrous) (Na2COs = io6).

Na2CO3 + H2SO4 = Na2SO4 + H2O + CO2. 2)106 2)98.09 N

53 gms. 49-045 gms. = 1000 mils — V.S.

N Each mil of — acid =0.053 gm- Na2CO3.

Sodium Bicarbonate (NaHCO3 = 84).

2NaHCO3 + H2SO4 = Na2SO4 + 2H2O + 2CO2. 2)168 2)98.09 N

84 gms. 49-045 gms. = 1000 mils — V.S.

N Each mil of — acid = 0.084 gm- NaHCO3.

ANALYSIS BY NEUTRALIZATION 75

Previous to weighing, the salt should be dried to constant weight in a desiccator over sulphuric acid. Lithium Carbonate (Li2CO3 = 73.88).

Li2CO3 + H2SO4 = Li2SO4 + H2O + CO2.

2)73-88 2)98.09 N

36.94 gms. 49-045 gins. = TOGO mils — V.S.

N Each mil of — acid = 0.03694 gm. Li2CO3.

The U.S. P. IX recommends residual titration for this salt. Previous to weighing it should be dried to constant weight at 100° C. 1.5 gm. is dissolved in 50 mils of normal sulphuric acid V.S. and the solution titrated with normal potassium hydroxid V.S., using methyl orange T.S. as indicator.

Ammonium Carbonate (N3HnC2O5 = 157.03). Normal ammonium carbonate has the formula (NH4)2CO3, but the normal salt loses upon exposure NH3 and H2O. The commer- cial salt, therefore, generally is a mixture of bicarbonate and carbamate.

(NH4)2C03 - NH3 = NH4HC03; (NH4)2C03 - H20 = NH4NH2C02.

The commercial carbonate is therefore generally expressed thus:

NH4HC03.NH4NH2C02 or N3HnC2O5.

This salt may be estimated by direct titration with normal or decinormal acid, using rosolic acid or methyl orange as an indicator.

Two grams of the salt are taken, dissolved in about 50

N mils of water and titrated with — H2SO4 V.S. The reaction

is as follows:

2N3HiiC2O5 + 3H2SO4 = 3(NH4)2SO4 + 4CO2 + 2H2O.

52.34 gms. 49-045 gms. = 1000 mils — acid V.S.

76 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

N Each mil of — acid V.S. represents 0.052 gm. of N3HnC2O5

or 0.017 gm. of NHs. The U.S. P. salt should contain between 30 and 32 per cent of NH3.

If in this titration 37.3 mils of the standard acid are required then the two grams of ammonium carbonate contained 0.052 gm.X 37.3 = 1.939 gms. of the salt.

1.939X100

- =96.95 per cent.

If rosolic acid is used as indicator heat must be applied to expel carbon dioxid. The estimation of the carbonic acid may be effected by precipitating a definite weight of the salt with barium chlorid, collecting the precipitated barium car- bonate, dissolving it in a measured excess of normal hydro- chloric acid and retitrating with normal alkali.

The method usually employed by skilled analysts (the resid'ual titration method), is to add a measured excess of the standard acid solution, and thus convert the ammonium car- bonate into the less volatile ammonium sulphate; then gently boil to get rid of CC>2, and titrate back with a standard alkali V.S. (using methyl orange as an indicator) until the excess of acid is neutralized. The quantity of free acid thus found, when deducted from the amount of acid first added, gives the quantity which was required to neutralize the ammonium carbonate.

Thus 2 gms. (in solution) of ammonium carbonate are

N treated with 50 mils of — H2SO4 V.S., which is more than

sufficient to neutralize it; the solution is then gently boiled to drive off CO 2, a few drops of litmus tincture added, and

N then titrated with — KOH V.S. until the litmus no longer

shows an acid reaction and the solution is neutral. If methyl orange is used as an indicator here, boiling is not necessary.

ANALYSIS BY NEUTRALIZATION 77

N Let us assume that 12.7 mils of the — KOH V.S. were

N used. By deducting the 12.7 mils from the 50 mils of — acid

first added, we find 37.3 mils of the acid went into combination with the ammonium salt, the calculation is then made as described above.

Mixed Alkali Hydroxid and Carbonate

If it is desired to ascertain the proportion in which these exist in a mixture, we proceed as follows:

First determine the total alkalinity by means of normal hydrochloric acid, using methyl -orange as an indicator. Then dissolve a like quantity of the mixture in 150 mils of water and add sufficient barium chlorid to precipitate all of the carbonate as barium carbonate, and then add water to make 200 mils and set aside to settle. When the supernatant liquid is clear take one-fourth (50 mils) of it, and titrate with normal hydrochloric acid, using phenolphthalein as indicator.* The number of mils multiplied by 4 will be the quantity of normal acid required by the caustic alkali. The difference between this and the number of mils representing the total alkalinity is calculated as carbonate.

Example. Assuming that we are analyzing a mixture of sodium hydroxid and carbonate.

Two grams of the substance are dissolved in water and titrated with normal acid solution. 43.2 mils of the latter are required. Another 2 gms. is dissolved, treated with barium chlorid as directed, and one-fourth of the clear solution titrated

* The slight error which occurs in this method because the volume of the precipitate is included in the measured liquid, may be overcome by using the entire quantity of liquid, including the precipitate (instead of taking one- fourth of it), and titrating with oxalic acid V.S. in the presence of phenol- phthalein. Oxalic acid in very dilute solutions does not react with alkali earth carbonates.

78 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

with normal acid. 5.6 mils are required; then 5.6X4=22.4 mils, representing the sodium hydroxid.

43.2 mils = total alkalinity; -22.4X0.040 = 0.896 gm. sodium hydroxid.

20.8X0.053 = 1.1024 gms. sodium carbonate.

Another way is to filter the mixture after barium chlorid has been added, titrate the filtrate with normal acid to find the quantity of hydroxid, then dissolve the precipitated barium carbonate in normal hydrochloric acid in excess, and re titrate with normal alkali, thus ascertaining the amount of carbo- nate.

When the alkaline carbonate is present in very small quantities the method of Lunge may be employed.

A few drops of phenacetolin solution are added to impart a scarcely perceptible yellow to the liquid. Normal acid solution is then run in until a pale rose tint appears, indicating that all the alkali hydroxid is neutralized; the volume of acid is noted, and the titration continued; the red color is inten- sified, and when the carbonate is entirely decomposed a golden-yellow color results.

Considerable practice is required with solutions of known composition to accustom the eye to the changes of color.

Mixed Alkali Bicarbonates and Carbonates

Thompson's Method. Take 2 grams of the salt and dissolve in 100 mils of water. Divide the solution into two equal parts and titrate one portion with normal acid solution, using methyl orange as indicator, and note the quantity required. We will assume 13 mils.

Then treat the second portion with a measured excess (say 25 mils) of normal sodium hydroxid solution free from

ANALYSIS BY NEUTRALIZATION 79

CO2. This converts the bicarbonate into carbonate. Now add an excess of pure neutral barium chlorid solution in order to precipitate all the carbonate as barium carbonate, and then titrate with normal acid, using phenolphthalein as indicator, to determine the excess of sodium hydroxid. 15 mils are required. Thus

25 -15 = 10 mils, the equivalent of bicarbonate; and 13 -10 = 3 m^sj the equivalent of carbonate; 10 X. 084 = .840 gm. sodium bicarbonate; 3 X. 053 = .159 gm. sodium carbonate.

Sodium Borate (Borax) (Na2B4O7 + ioH2O = 382.16). Dis- solve 5 gms. of the salt in 100 mils of distilled water and titrate the solution with normal hydrochloric acid V.S., using methyl orange T.S. as indicator.

Na2B4O7 • ioH2O + 2HC1 = 2NaCl + 4H3BO3 + 5H2O.

382.16

Each mil of normal HC1 V.S. corresponds to 0.19108 gm. of crystallized borax or to o.ioi gm. of anhydrous borax.

Sodium Cacodylate (Na(CH3)2AsO2 = i6o.oi). This salt is assayed by titration with normal HC1 V.S. in the presence of methyl orange. Each mil of the acid V.S. corresponds to o.i 60 gm. of the salt.

This salt is occasionally slightly acid in reaction; if so, it should be carefully neutralized to phenolphthalein before titration is begun.

Sodium Glycerophosphate (Na2C3H5(OH)2PO4 = 216.1). This salt is titrated with half -normal hydrochloric acid V.S. in presence of methyl orange.

N Each mil of — acid V.S. corresponds to 0.10805 gm. oi

the salt.

80 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

Estimation of Organic Salts of the Alkalies

The tartrates, citrates and acetates of the alkali metals are converted by ignition into carbonates, the whole of the base remaining in the form of carbonate.

Each molecular weight of a normal tartrate gives when ignited one molecular weight of carbonate:

Every two molecular weights of an acetate or an acid rate give one molecular weight of carbonate:

Every two molecular weights of a normal citrate give three molecular weights of carbonate :

These reactions are taken advantage of in volumetric analysis, and the tartrates, citrates and acetates of the alkalies are indirectly estimated by calculating upon the quantity of carbonate formed by burning them, the quantity of carbonate being found by titration in the usual manner.

The Process. Before igniting, the salt to be examined should be thoroughly dried in a desiccator over calcium chlorid or in a drying oven, the latter only for such salts as have no water of crystallization in their composition. If the weight is taken before and after, the amount of moisture present is determined. One or two grams of the dried salt is weighed accurately, placed in a porcelain crucible, and heat applied gradually, until dull redness is reached and white fumes cease to be given off. Upon applying heat to the salt, the latter swells, fuses, and then boils, and if the heat is applied too rapidly at this point, there is apt to be a considerable loss

ANALYSIS BY NEUTRALIZATION 81

of material through sputtering. The flame of the burner must not come in contact with the carbonized mass. The completion of the ignition is known to be reached when the black contents of the crucible is dry and crisp. The crucible is then allowed to cool, its contents disintegrated with the aid of a stout glass rod and then treated with boiling water to dissolve out the alkali carbonate, and the solution filtered through a small, wetted filter into a flask or beaker. The filtrate should be perfectly colorless. If it has a yellow or brownish color it indicates incomplete ignition and should be rejected, and a fresh quantity of the salt subjected to ignition. The contents of the crucible and the filter should be washed with several small portions of water until the washings no longer show an alkaline reaction. The filtrate mixed with the wash water is now titrated with standard sulphuric or hydro- chloric acid, using methyl orange as the indicator. From the quantity of carbonate found in the filtrate, the equivalent amount of the organic salt may be calculated. The quantity of standard acid employed is multiplied direct by the factor for the original salt. The residual titration method, using an excess of half-normal sulphuric acid V.S., boiling, and retitrating with half-normal potassium hydroxid V.S., is recommended in the U.S.P.

In tne case of organic salts of the alkali earths,* residual titration should always be resorted to. The residue in the crucible being dissolved in standard hydrochloric acid, and retitrated with standard alkali.

Lithium salts, because of the sparing solubility of the carbonate in water, should also be titrated by the residual method.

Potassium Tartrate (K2C4H4O6= 226.23). Two grams of

* Organic salts of the alkali earths subjected to ignition as above are -eHuced partly to oxids.

82 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

the salt is placed in a platinum or porcelain crucible and heated to redness in contact with the air until completely charred; that is to say, until nothing is left in the crucible but carbonate and free carbon.

The crucible is now cooled, and its contents treated with boiling water, which dissolves the potassium carbonate, the carbon being separated by filtration. In order to obtain every trace of carbonate it is well to wash the crucible with several small portions of hot water, and add the washings to the rest of the filtrate through the filter.

If the salt is completely carbonized the filtrate will be colorless, but if the carbonization is not complete the solution will be more or less colored and should be .rejected, and a fresh quantity of the salt subjected to ignition.

To the filtrate, which contains potassium carbonate, add

N a few drops of methyl orange, and titrate with — sulphuric

acid V.S. until a light orange-red color appears and the car- bonate is neutralized.

The following equations will explain the reactions:

K2C4H4O6 = K2CO3 + C2 + CO + 2H2O

226.2 138.2

then

K2C03 + H2S04 = K2S04 + H20 + CO2

138.2 98.07

therefore

K2C4H406 = K2C03 = H2S04

2)226.2 2)138.2 2)98.07 ,,

113.1 gms. = 69.; gms. = 49.03 gms. = iooo mils — V.S.

N and each mil of — H2SO4 represents 0.1131 gm. of potassium

tartrate.

Example. Two grams of potassium treated as described

ANALYSIS BY NEUTRALIZATION 83

N above require 16.3 mils of — H2SO4. It therefore contains

0.1131X16.3=1.8435 gms.

1.8435X100

— = 92.17 per cent.

Potassium and Sodium Tartrate (KNaC4H4O6.4H2O

= 282.22) (Rochelle Salt). This salt is treated in exactly the same way as described for potassium tartrate.

When ignited the double tartrate is converted into a double carbonate of potassium and sodium :

KNaC4H4O6 = KNaCO3 + etc.;

2IO.I 122. 1

then

KNaC03 + H2SO4 = KNaSO4 + CO2 + H2O

therefore

KNaC4K406 = KNaC03 = H2SO4

2)210.1 2)122.1 2)98.07 ^,

105.05 61.05 49.03 = 1000 mils— V.S.

N and each mil of — H2SO4 represents 0.10505 gm. of

KNaC4H4O6.

Example. If one gram of rochelle salt treated as above

N described requires 7 mils of — sulphuric acid, it contains

0.10505X7=0.7353 gm. = 73.53 per cent.

According to the U.S.P. IX this salt is assayed by residual titration, using half-normal sulphuric acid V.S., as follows:

Heat two gms. of the salt in a porcelain crucible, until thoroughly carbonized. Allow the carbonized mass to cool, dis- integrate it with the aid of a stout glass rod and transfer the mass and crucible to a beaker. Add 50 mils of distilled water and 50 mils of half-normal sulphuric acid V.S., cover the beaker

84 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

with a watch glass and boil the contents for thirty minutes. Then filter the solution and wash the residue with hot distilled water until the washings cease to redden blue litmus paper. Now determine the residual acid in the cooled filtrate by titra- tion with half-normal KOH V.S., using methyl orange as indicator.

Subtract the number of mils of half-normal alkali V.S. used from 50. mils (the quantity of half-normal acid V.S. taken). The remainder represents the quantity of half-normal sulphuric acid V.S. which reacted with the carbonate in the

N charred mass. This, multiplied by : factor for rochelle

salt, represents the weight of the latter in the amount taken for analysis. The foregoing equations show the normal factor for anhydrous rochelle salt to be 0.10505 gm., hence the half normal factor for the same is 0.05253 gm.

Hence, if in the above assay 22 mils of half -normal KOH V.S. were used, this quantity deducted from 50 mils, leaves 28 mils, the quantity of half- normal sulphuric acid which was taken up by the alkaline carbonized mass; then 0.05253X28 = 1.470 gm. of KNaC4H4Oe in the 2 gms. of rochelle salt taken for analysis =73. 5 per cent.

Potassium Bitartrate (KHC4H4O6=i88.i) (Cream of Tar- tar). The estimation of this salt is affected in the same way as the tartrate.

The bitartrate having but one atom of potassium in its molecule, it takes two molecules to form one molecule of car- bonate.

2KHC4H406 = K2C03 + H2S04

2)376.2 2)98.07 N

188.1 gms. 49.03 gms. = 1000 mils — V.S.

Each mil of - H2SO4 V.S.=o.i88i gm. of KHC4H4O6.

ANALYSIS BY NEUTRALIZATION 85

Another way of estimating bitartrate is to dissolve a weighed

N quantity of hot water and titrate with — potassium hydroxid

until neutral, and thus the amount of tartaric acid existing as bitartrate is found. The bitartrate is acid in reaction. In detail the method is as follows :

Two grams of the bitartrate are dissolved in 100 mils of hot water, a few drops of phenolphthalein T.S. added, and

N then titrated with — KOH V.S. until a faint pink color indi-

cates that all of the acid has been neutralized. Not less than 10.6 mils of the normal alakli should be required, corresponding to 99.6 per cent of pure salt.

The following equation will show the reaction:

KHC4H406 4 KOH •• K2C4H4O6 + H2O. 188.1 56.1 = 1000 mils of - KOH V.S.

N Each mil of — KOH V.S. represents 0.1881 gin. of

KHC4H4O6.

If 10.6 mils are required for neutralization, then 10.6X0.1881 = 1.993+ gms-:

1.993X100

-=99.6 per cent.

Potassium Citrate (KsCeHsOy = 306.3) .

2K3C6H507 - 3K2C03 = 3H2S04.

6)612.6 6)414.6 6)294.18 ^j

102. i gms. 69.1 = 49-03 gms. = looo mils — acid.

N Thus each mil of — acid represents 0.1021 gm. of pure

N potassium citrate, and each mil of — acid represents 0.0510 gm.

Potassium Acetate (KC2H3O2 = 98. i ) . In estimating potas- sium acetate the salt is ignited and the residue treated in

86 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

exactly the same manner as in the estimation of the citrates and tartrates before mentioned.

2KC2H3O2 = K2CO3.

2)196.2 N

98.1 gms. = 1000 mils — H2SO4.

N Each mil therefore of — H2SO4 corresponds to 0.0981

gm. of potassium acetate.

Sodium Acetate (NaC2H3O2.3H2O = 136.09). 2 (NaC2H302 . 3H20) = Na2CO3,

N Each mil of — H2SO4 V.S. represents 0.06804 gm. of crys-

tallized sodium acetate.

Sodium Benzoate (NaC7H5O2 = 144.05). 2NaC7H5O2

N Each mil of — H2SO4 V.S. represents 0.07202 gm. of sodium

benzoate.

Sodium Salicylate (NaC7H5O3= 160.05). 2NaC7H5O3 =

N Each mil of — H2SO4 V.S. represents 0.08002 gm. of sodium

salicylate.

Lithium Citrate (Li3C6H5O7 = 209.92). As stated before, the organic salts of lithium and those of the alkali earth metals are best titrated by the residual method, after ignition, because the carbonates formed are insoluble in water. It is likewise best to use standard hydrochloric instead of standard sulphuric acid. The process for lithium citrate here given exemplifies the method.

Two grams of the salt is thoroughly ignited in a porcelain crucible as described for potassium tartrate. The residue of lithium carbonate is then dissolved out of the crucible by add-

ANALYSIS BY NEUTRALIZATION 87

N ing 50 mils of — hydrochloric V.S. and filtering. The crucible

and filter are washed with several small quantities of water and the washings added to the acid filtrate. Three drops of methyl orange are now added, and the solution titrated with

N

^- potassium hydroxid V.S. until the yellow color appears.

Assuming that 8 mils of the standard alkali were required, then 50-8 = 42 mils, the quantity of half-normal hydrochloric acid which reacted with the lithium carbonate. This quantity multiplied by the half -normal factor for lithium citrate, 0.03498, gives the weight of pure salt in the 2 gms. taken.

0.03498X42 = 1.473!+ gm- or 73-65 Per cent.

Calcium Lactate (Ca(C3H5O3)2 + 5H2O = 308.23). This is assayed like the foregoing, half normal hydrochloric acid V.S.

Each mil of half normal acid V.S. corresponds to 0.05454 gm. of CaCCsHsOa^-

Strontium Salicylate (Sr(C7H5O3)2 + 2H2O = 397.74). This is assayed like the foregoing, but care must be taken not to allow the temperature to exceed red heat during the carbon- ization. Each mil of half normal hydrochloric acid V.S. cor- responds to 0.099435 gm. of Sr(C7H5O3)2 + 2H2O.

Pulvis Effervescens Compositus (Seidlitz Powder). The weight of the mixture in the blue paper should be not less than 9.5 gms. nor more than 10.5 gms. It should contain not less than 23 per cent nor more than 27 per cent of sodium bicarbonate, and not less than 73 per cent nor more than 78 per cent of rochelle salt.

The Assay for Sodium Bicarbonate. Dissolve 2 gms. of the contents of the blue paper in 80 mils of distilled water, add 20 mils of half-normal sulphuric acid V.S., boil the solu- tion until the volume is reduced to about 50 mils and titrate the excess of acid with half-normal potassium hydroxid V.S.,

88 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

using phenolphthalein T.S. as indicator. Subtract the number of mils of half-normal potassium hydroxid V.S. used from 20 mils (the quantity of half-normal sulphuric acid taken) and the remainder will be the quantity of the latter which reacted with the sodium bicarbonate. The calculation is as follows:

2NaHCO3 + H2SO4 = Na2SO4 + 2H2O + 2GO2.

2)168.02 2)98.09

2) 84.01 2)49.045

42.005 24.522 gms. in looo mils of half -normal V.S.

Each mil of half-normal acid V.S. = 0.042005 gm. of NaHC03.

Assuming that 8 mils of half-normal potassium hydroxid were consumed in the assay then 12 mils of half-normal sul- phuric acid reacted with, and hence represent the sodium bicarbonate present.

Hence 0.042X12 = 0.504 gm. or 25.2 per cent.

The Assay] for Rochelle Salt. Take 2 gms. of the contents of the blue paper used in the preceding assay, place it in a porcelain crucible and carbonize it as described under potas- sium and sodium tartrate. The difference between the num- ber of mils of half-normal sulphuric acid V.S. consumed in this assay and in the preceding assay for sodium bicarbonate multiplied by 0.07055 represents the rochelle salt.

N Thus, 50.0 mils of — H2SO4 V.S. taken

N 16.5 mils of — KOH V.S. required for neutralization

N 33.5 mils of — H2SO4 consumed

33.5 mils -12 mils =2 1. 5 mils, representing rochelle salt; 0.07055X21.5 = 1.516 gm. or 75.8 per cent of crystallized rochelle salt.

ANALYSIS BY NEUTRALIZATION 89

Ammonium Benzoate (NH4C7H5O2= 139.08) and Ammo- nium Salicylate (NELtCyHsOs^SS-oS). It is evident that because of the volatility of ammonia, these salts cannot be assayed by the method employed in the case of organic salts of the fixed alkalies. The official assays of these salts depends upon the liberation of the organic acid, by the addition of sulphuric acid, and its extraction with an immiscible solvent. The solution so obtained is evaporated and the organic acid residue dissolved in neutralized diluted alcohol and titrated with tenth-normal barium hydroxid V.S. The method in detail is as follows: Dissolve about 0.5 gm. of the salt pre- viously dried in a desiccator over sulphuric acid, and accurately weighed, in 10 mils of distilled water, in a separator. Add to the solution 5 mils of diluted sulphuric acid and extract the liberated organic acid by shaking out with three successive portions of 25, 15, and 10 mils respectively of chloroform, passing the chloroform solution through a filter previously moistened with chloroform, and removing any of the organic acid adhering to the stem of the funnel with a few mils of chloroform. Evaporate the chloroform solution at a very low temperature to 5 mils, add 25 mils of diluted alcohol,

N which has been previously neutralized with — KOH V.S. in

the presence of phenolphthalein.

N Titrate this solution with — barium hydroxid V.S., using

phenolphthalein as indicator.

N Each mil of the — Ba(OH)2 V.S. represents benzoic acid,

o.oi2205~gm.; salicylic acid, 0.013805 gm.; ammonium benzoate 0.013908 gm., and ammonium salicylate 0.015508 gm.

90 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

TABLE SHOWING THE NORMAL FACTORS, ETC., OF THE ORGANIC SALTS OF THE ALKALI METALS.

Substance.

Formula.

Molecular Weight.

Equivalent Weight in Carbonate.

Norma^ Factor.

Lith urn benzoate

LiC7H5O2

127 98

•26 04

O 12798

citrate

Li3C6H5O7

2OQ Q2

02 3S

0.06997

' salicylate Sod um acetate

LiC7H503 NaC2H3O2 3H2O

143.98 1 36 O7

36.94 e-2 o

0.14398 o 13607

benzoate ' salicylate Potassium acetate

NaC7H502 NaC7H603 KC2H3O2

144.04 i 60 . 04

08 12

53-o 53-o 60 i

o. 14404 0.16004 o 09812

bitartrate ' citrate ^

KHC4H4O6 K3C6H5O7 H2O

I88.I4 2 24. 36

69. i 207 3

0.18814

o 10812

' tartrate

K2C4H4O6|H2O

235. 24

138.2

o. 11762

' and sodium tartrate

KNaC4H406.4H20

282. 2O

122. I

0.1411

Estimation of the Salts of the Alkali Earths

Standard solution of hydrochloric or of nitric acid is pre- ferred by many operators for the titration of hydroxids or carbonates of the alkali earths. These acids possess the advantage over most other acids of forming soluble salts. The hydroxids may be estimated by any of the indicators, but as they readily absorb CC>2 out of the air they are generally contaminated with more or less carbonate, and the residual method should be used, i.e., a known excess of standard acid should be added, the mixture boiled to expel any trace of CC>2, and titrated with standard alkali.

The carbonates are of course estimated in the same way, as are also the organic salts of the alkali earths, after ignition. As an example:

One gram of calcium carbonate is mixed with 5 mils of water. An excess of normal hydrochloric acid V.S. is now added, and the solution boiled to drive off the CO2. Then a few drops of phenolphthalein T.S. are added, and

ANALYSIS BY NEUTRALIZATION 91

1 N r

titrated with — alkali V.S. until a faint pink color is ob- tained.

N Note the quantity of — alkali used, and deduct this from

N the quantity of — acid first added, and the remainder will

represent the amount of acid which combined with the calcium.

N Each mil of — acid V.S. represents 0.05 gm. of CaCO3.

CaCO3 + 2HC1 = CaCl2 + H2O + CO2.

2)100 2)72.92 N

50 gms. 36.46 gms. or 1000 mils — acid V.S.

N Assuming that 30 mils of — HC1 V.S. were added to the i

N

gm. of CaCO3, and that n mils of — KOH V.S. were re- quired to bring the mixture back to neutrality, then 19 mils

N of — HC1 were actually required to saturate the CaCOs.

Therefore 0.050 X 19 = 0.950 or 95 per cent.

The hydroxids and carbonates may also be estimated by direct titration with standard hydrochloric acid (in the cold) using methyl orange as indicator. A better plan, however, would be to add the standard acid in slight excess, and then standard alkali until a distinct yellow color appears; the slight excess of alkali is then determined by adding standard hydrochloric acid until the red color reappears. A much more distinct color reaction is thereby obtained. The quantity of the standard alkali used is deducted from the total quantity of standard acid added.

Soluble salts of calcium, barium and strontium, such as chlorids, nitrates, etc., may be readily estimated as follows:

92 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

A weighed quantity of the salt is dissolved in water, cautiously neutralized if it is acid or alkaline, phenolphthalein is added, the mixture heated to boiling, and standard solution of sodium carbonate delivered in from time to time, with constant boiling until the red color is permanent.

This process depends upon the fact that sodium carbonate forms with soluble salts of these bases insoluble neutral car- bonates.

CaCl2 + Na2CO3 = CaCO3 + 2NaCL Ba(NO3)2 + Na2CO3 = BaCO3 + 2NaNO3.

As an example of the process: Take of calcium chlorid one gram, dissolve it in a small quantity of water, neutralize the solution if it is acid or alkaline, heat to boiling, add a

N few drops of phenolphthalein, and titrate with - - sodium

carbonate, delivered cautiously while boiling until the red color is permanent.

CaCl2 + Na2CO3 = CaCO3 + 2NaCl.

2ll_i3_ 2)106 N

56.5 gms. 53 gms. or 1000 mils — V.S.

N Each mil .of — Na2CO3 V.S. represents 0.0565 gm. of

CaCl2. If 17 mils are used the salt contains 0.0565 gm.Xiy = 0.96 gm. or 96 per cent.

Normal Sodium Carbonate V.S. (Na2CO3 = io6) contains 53 gms. in T liter. This solution is made by dissolving 53 gms. of pure sodium carbonate (anhydrous) previously ignited and cooled, in distilled water, and diluting to i liter at 25° C.

If a pure salt is not at hand the solution may be made as follows :

About 85 gms. of pure sodium bicarbonate, free from thiosulphate, chlorid, etc., are heated to dull redness (not to

ANALYSIS BY NEUTRALIZATION 93

fusion) for about thirty minutes to expel one-half of the CO2; it is then cooled under a desiccator. When cool, 53 gms. are dissolved in distilled water to make i liter at 25° C. (77°

N F.). This solution should neutralize — acid V.S. volume

for volume.

The alkali earths may also be estimated by dissolving them in water, precipitating the base as carbonate, with an excess of ammonium carbonate and some free ammonia. The mixture is then heated for a few minutes, and the car- bonate separated by filtration, thoroughly washed with hot water till all soluble matters are removed, and then titrated with normal acid V.S. as directed for carbonate.

CaBr2 = CaCO2 = H2SO4.

2)198.52 2)99.35 2)98 N

99.26 gms. 49-675 gms. 49 gms. or 1000 mils — V.S.

i

N Each mil of — acid thus represents 0.09926 gm. of CaBr2.

Another method for the estimation of salts of the alkali earths consists in precipitating them as oxalates out of an ammoniacal or weak acetic acid solution. The precipitated oxalate is then, after thorough washing, titrated with tenth- normal potassium permanganate V.S. Or the excess of oxalic acid in the filtrate may be titrated with standard 'permanganate V.S. This method is, however, especially applicable to calcium estimations because of the completeness with which this metal may be precipitated as oxalate.

Magnesium Carbonate. One gm. of magnesium carbon- ate is dissolved in 30 mils of normal sulphuric acid V.S., and the excess of the latter determined by titration with normal potassium hydroxid V.S., using methyl orange as indicator. Magnesium Oxid and Magnesium Hydroxid are estimated in

94 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

exactly the same manner. The absence of calcium oxid must, however, be assured, or, if present, its quantity determined.

This is done by precipitation with ammonium oxalate in the presence of ammonium chlorid and ammonium hydroxid. The precipitate is incinerated and the residual calcium oxid weighed.

Alkalimetric Assay of Zinc Salts. Zinc Oxid. Digest about 1.5 gm. of zinc oxid, accurately weighed, with 50 mils of normal sulphuric acid V.S., until solution is complete. Then titrate the excess of sulphuric acid with normal potassium hydroxid V.S., using methyl orange as indicator.

Zinc carbonate and zinc stearate are assayed in the same manner. In the case of the latter, boiling for ten minutes may be necessary in order to completely dissolve the salt. Each mil of normal sulphuric acid V.S. corresponds to 0.040685 gm. of zinc oxid.

The Estimation of Mixed Hydroxids and Carbonates of Alkali Earths. This may be done as described under esti- mation of mixed alkali hydroxids and carbonates, page 77, except that in this case it is unnecessary to precipitate the carbonate by barium chlorid in that the alkali earth carbonates are already insoluble.

Assay of Chloral Hydrate. Dissolve an accurately weighed quantity (about 4 gms.) of chloral hydrate in 10 mils of dis- tilled water, add 30 mils of normal potassium hydroxid V.S., and let the solution stand about two minutes. Then add phenolphthalein T.S., and at once titrate the excess of alkali by means of normal sulphuric acid V.S.

This reaction must take place in the cold or at least at ordinary temperature, otherwise the alkali will attack the chloroform which is formed and yield too high results. The

ANALYSIS BY NEUTRALIZATION 95

chloroform liberated produces a slight turbidity which may be dissipated by swinging the flask for a minute or two. The reaction is thus expressed :

165.4

Each mil of normal alkali consumed corresponds to 0.1654 gm. of C2HC130+H20.

Acidimetry

The Estimation of Acids by Neutralization. In the preceding pages it has been shown how alkalies are estimated by the

FIG. 35. FIG. 36.

use of acid solutions of known neutralizing power. In the estimation of acids, which will now be described, the order

96 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

is reversed, alkaline solutions of known power being used in determining the strength of acids and of acid salts. Thus the procedure is analogous to that of the alkalimetric methods. The choice of the indicator, whether litmus, phenolphthalein, or methyl orange, depends upon the particular acid to be estimated. Phenolphthalein is employed for the organic acids and boric acid and is preferred for phosphoric acid; while methyl orange and litmus are usually employed in the titration of the inorganic acids.

The standard alkali used may be either an hydroxid or a carbonate, the former is, however, usually preferred, because the carbonate when brought in contact with an acid gives off carbonic acid gas (CC>2) which interferes to a great extent with most indicators. On the other hand, it must be remem- bered that the alkali hydroxids are very prone to absorb carbon dioxid from the atmosphere, therefore their solutions often contain some carbonate, the presence of which even in small quantities will occasion errors when used with most indicators, especially with litmus and phenolphthalein. It is therefore advisable, when using these indicators or others which are affected by carbon dioxid, to employ gentle heat toward the close of each titration, in order to drive off the liberated gas. Methyl orange is not affected by this gas, and therefore heating is not necessary when this indicator is used. In fact, it is imperative that heat should not be employed with this indicator.

In acrdimetrical operations when methyl orange is used as indicator, residual titrations may be advantageously done, because the change of color from yellow to red which is brought about by the acid is much more readily seen than that from red to yellow.

In the U. S. P. standard solutions of both potassium and sodium hydroxid are official. The former, however, is pref-

ANALYSIS BY NEUTRALIZATION 97

erable, because it attacks glass less energetically, and also foams much less than does the sodium hydroxid solution. The neutralizing power of both is, however, the same. Standard solutions of alkali hydroxid should be preserved in small vials, provided with well-fitting rubber stoppers, or better still, they should be provided with tubes filled with a mixture of soda and lime, which absorbs CC>2 and prevents its access to the solution. A vessel of this description is illustrated in Fig. 35.

An improvement upon this is shown in Fig. 36, since it allows of the burette being filled without removing the stopper, and consequently without any access of CC>2 whatever.

Where a series of titratio'ns of the same kind have to be made with the same alkali standard solution, the arrangement shown in Fig. 9 may be used, both the reservoir and the burette in this case being provided with soda-lime tubes.

Preparation of Standard Alkali Solutions

Normal Potassium Hydroxid V.S. (KOH = 56.11; y V.S.

= 56.11 gms. in 1000 mils). Potassium hydroxid being prone to absorb carbon dioxid out of the air the pure article is not readily obtained in commerce. If pure potassium hydroxid were easily obtained it would only be necessary to dissolve 56.1 gms. in sufficient water to make 1000 mils. But since it always contains some CO 2 and H2O, it is necessary to take a slight excess and dilute the solution to the proper volume after having determined its strength.

The standardization may be effected by means of any of the standard acid solutions.

A satisfactory method for the preparation and standard- ization of this solution is as follows :

Dissolve 75 gms. of potassium hydroxid in sufficient recently

98 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

boiled distilled water to make about 1050 mils at 25° C. (77° F.), and fill a burette with a portion of this solution.

Dissolve 0.63 gm, of pure oxalic acid in about 10 mils of water in a beaker or flask, add a few drops of phenolphthalein T.S., and then carefully add from the burette the potassium hydroxid solution, agitating frequently and regulating the flow to drops towards the end of the operation until a permanent pale pink color is obtained. Note the number of mils of the alkali solution consumed, and then dilute the remainder so that exactly 10 mils of the diluted liquid will be required to neutralize 0.63 gm. of oxalic acid. Instead of weighing off 0.63 gm. of the acid, 10 mils of its normal solution may be used.

Example. Assuming that 8 mils of the stronger potassium hydroxid solution had been consumed in the trial, then each 8 mils must be diluted to 10 mils, or the whole or the remaining solution in the same proportion. Thus if 8 mils must be diluted to 10 mils, icoo mils must be diluted to 1250 mils.

8 : 10 :: 1000 :x #=1250 mils.

It is always advisable to make another trial after diluting. Ten mils should then neutralize 0.63 gm. of pure oxalic acid. Standardization by Means of Potassium Bitartrate.

This method is based upon the reaction

KHC4H406 + KOH = K2C4H406 + H2O. 188.14 56.1

N 1000 mils of — KOH contains 56.11 gms. of KOH and

will react with 188.14 gms. of potassium bitartrate. 25 mils

N of — KOH will therefore react with 4.7035 gms. of potassium

bitartrate.

A solution of potassium hydroxid, 75 gms. in 1050 mils, is

ANALYSIS BY NEUTRALIZATION 99

prepared and titrated against pure potassium bitartrate, using phenolphthalein as indicator.

Into a flask of about. 300 mils capacity introduce 4.7035 gms. of purified dry potassium bitartrate,* followed by 15 mils (accurately measured at 25° C.) of the potassium hydroxid solution which is being prepared and 80 mils of distilled water. Heat the solution to boiling, add from 3 to 5 drops of phenol- phthalein T.S. and then cautiously add from a burette further portions of the potassium hydroxid solution. Agitate the flask frequently, boil the liquid toward the end of the operation and reduce the flow of the potassium hydroxid solution to drops until the red color produced by its influx no longer disappears on shaking and the liquid is not deeper in color than a pale pink. Note the number of mils of the potassium hydroxid solution consumed and then dilute the remainder of the solution so that exactly 25 mils of the diluted liquid at standard temperature shall be required to neutralize the 4.7035 gms. of potassium bitartrate. If only 22 mils are consumed in the trial, then each 22 mils must be diluted to 25 mils or the whole of the remaining solution in the same proportion.

* Purified potassium bitartrate for standardizing caustic alkali volumetric solutions may be obtained as follows: 100 gms. of the salt are placed in a beaker, together with 85 mils of water and 25 mils of 10 per cent hydrochloric acid, the beaker is covered and heated on a boiling water-bath, stirring occasionally for three hours. The liquid is then quickly cooled, decanted, and the residue washed first by decantation with 100 mils of cold water, then again washed after it has been transferred to a plain filter, using cold water until the filtrate ceases to become opalescent when acidified with a few drops of nitric acid upon the addition of silver nitrate solution. The precipitate is then dissolved in the smallest possible quantity of boiling water (about 1500 mils) filtered, and the filtrate stirred constantly while cooling. When cold, the crystalline precipitate is collected upon a filter washed with 300 mils of cold water, run through it in small portions at a time. It is then allowed to drain; and §nally dried in an air oven at 120° C. until its weight is constant,

100 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

Standardization by Means of Potassium Bi-iodate.*

Potassium bi-iodate is an acid salt and may be directly titrated with potassium hydroxid, using phenolphthalein as indicator.

One molecule of the bi-iodate is equivalent to one molecule of potassium hydroxid, as shown by the equation,

389.94 56.11

To standardize a potassium-hydroxid solution, weigh accu- rately 3.8994 gms. of potassium bi-iodate, dissolve it in about 25 mils of water, add a few drops of phenolphthalein, and then run into this, from a burette, the hydroxid solution which is to be standardized, until a pale pink color appears. Note the number of mils used and dilute the solution so that exactly 10 mils of it will neutralize 3.8994 gms. of the bi-iodate.

Example. Assuming that 8.2 mils had been consumed, then each 8.2 mils must be diluted to 10 mils, or the whole of the remaining solution in the same proportion.

The advantages of this salt as an ultimate standard are (i) that it may be procured in the market in a state of absolute purity; | (2) that it is permanent, being neither deliquescent nor efflorescent; (3) that it can be dried at 110° C. without decomposition; (4) that the results obtained with it are quite accurate, and (5) that it may be employed for standardizing most of the volumetric solutions commonly found in the laboratory.

* See Meinecke, Chem. Ztg., XIX. 2. Also Caspari, Proc. A. Ph. A., 1904, 389.

t According to Caspari, the salt may be readily prepared as follows: See A. Ph. A., 1904, 390. Potassium bicarbonate is mixed in solution with an equivalent amount of iodic acid, and to the neutral solution is added an amount of iodic acid equal to the quantity first used. The solution is then evaporated until crystallization begins, and the first crop of crystals rejected. Those which separate after the solution has cooled to 50° C. are almost pure and will be rendered absolutely pure if recrystallized.

ANALYSIS BY NEUTRALIZATION 101

Standardization by Means of Normal Acid V.S. 20 mils of a strictly normal acid V.S. are placed into a beaker, 2 drops of methyl orange T.S. are added and the potassium hydroxid solution delivered into it until the red liquid just turns yellow. If the alkali hydroxid solution is strictly normal, there will be consumed exactly 20 mils. If less is consumed the solution is too strong and must be so diluted with distilled water that equal volumes of it and the normal acid will exactly neutralize each other. Thus if 18 mils of the alkali are con- sumed, then each 18 mils must be diluted to 20 mils.

N Normal Sodium Hydroxid V.S. (NaOH = 4o; — V.S. = 40

gms. in 1000 mils). Dissolve 54 gms. of sodium hydroxid in enough recently boiled distilled water to make about 1050 mils of solution, fill a burette with a portion of this, anci cnecK it with normal acid, or a weighed quantity of oxalic acid or potassium bitartrate, in exactly the same manner as described for normal potassium hydroxid.

Other strengths of standard alkali V.S. are half-normal

/N\ /N\ /N\

( — ], Fifth-normal, ( — ), Tenth-normal (— ), Twentieth-

/N\ /N\ /N\

normal (— ), Fiftieth-normal I— I, Hundredth-normal ( — I. \2o/ \5<V \ioo/

These are all prepared by properly diluting the normal V.S. and then checking the strength of the product.

Other standard alkali solutions in frequent use are normal sodium carbonate, normal and other strengths of ammonia, and tenth-normal barium hydroxid.

Tenth Normal Barium Hydroxid V.S. (Ba(OH)2 + 8H2O =

N 315.51 — V.S. = 15. 776 gms. in 1000 mils). Dissolve about 18

gms. of crystallized barium hydroxid in 1000 mils of recently boiled distilled water, and, if necessary, filter the solution quickly.

102 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

Introduce 25 mils of tenth-normal hydrochloric acid V.S. into a flask, add 3 drops of phenolphthalein T.S. and run in the barium hydroxid solution from a burette, provided with a soda-lime tube, reducing the flow to drops toward the end until a pale pink color is obtained which does not disappear on shaking the liquid for ten seconds. Note the number of mils consumed and if the solution is too concentrated, dilute the remainder with freshly boiled and cooled distilled water, so that equal volumes of the barium hydroxid solution and of tenth-nor- mal hydrochloric acid V.S. at 25° C. exactly neutralize each other.

NOTE. This solution absorbs CC>2 from the air very rapidly, and thereby loses its titer. It should therefore be preserved in rubber-stoppered bottles provided with a soda-lime tube. (See Fig. 35.) This solution must be standardized each time before using.

Estimation of the Inorganic Acids

To weigh off directly a definite quantity of a fluid acid, is not a very easy matter. It is always a better plan to measure a small quantity of the acid and weigh it accurately in a tared and stoppered weighing flask (Fig. 37), then to add water and titrate with the standard alkali in the presence of a suitable indicator. If the specific gravity of the acid is known or can be easily taken, it is sufficient to measure a certain quantity of it by means of a pipette, and then determine its weight by multiplying the volume in cubic centimeters by the specific gravity. It must be remembered, however, that the liquid must be measured at the same temperature at which the specific gravity was taken. This method is applicable to the diluted acids as well as to the concentrated acids of commerce, as hydrochloric, nitric and sulphuric.

ANALYSIS BY NEUTRALIZATION

103

In the case of very volatile acids, i.e., such as evolve acid vapors at ordinary temperatures, the determination of the weight by means of the specific gravity is inadmissable. Such acids should be weighed in a Lunge pipette, Fig. 38, or in a simple bulb pipette provided with a glass stop-cock, Fig. 39, or in a Grethan's pipette, Fig. 40.

The Lunge pipette is used by producing a vacuum in the bulb (a), the air-tight glass mantle (c) is then removed,

FIG. 38.

FIG. 39.

FIG. 40.

and the tip of the tube (d) sunk into the acid which is drawn up into the bulb, upon opening the cock (6); when sufficient of the acid has been drawn into the apparatus the cock is closed, the tip of the pipette wiped, the glass mantle put in place, and the whole weighed. The weight of the empty pipette deducted gives the weight of the acid taken up. The pipettes shown in Figs. 39 and 40 are filled by applying direct suction with the lips, the operator protecting himself against inhalation of harmful vapors by attaching an absorption tube containing soda-lime, caustic soda, or similar substance.

104 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

The quantity of acid to be taken (in most cases) should be such as will require for neutralization from 20 to 50 mils of the standard alkali. In the case of concentrated inorganic acids, 2 or 3 gms. may be taken, while in the case of the dilute acids, from 6 to 8 gms.

Any of the indicators may be employed for the inorganic acids, but because of the usual presence of carbonate in the standard alkali, methyl orange is preferred.

Hydrochloric Acid (HC1 = 36.47). About 2 mils of hydro- chloric acid (sp.gr. 1.155) are introduced into a tared weighing flask and its weight accurately taken. (The weight is found to be 2.098 gms.) Fifty mils of water are now added, followed by 2 drops of methyl orange, and the solution carefully titrated with normal potassium hydroxid until the reddish color of the solution is changed to yellow.

Assuming that 18.4 mils were required, then 18.4 milsX 0.03647 gm. = o.67i gm. of absolute hydrochloric acid in the 2.098 gms. taken.

To find the per cent apply the proportion

2.098 gms. : 0.671 gm. :: 100 : x. #=31.9 per cent. The equation is :

HC1 + KOH = KC1 + H2O.

N 36.47 gms. = 56.1 gms. = 1000 mils — V.S.

N .03647 gm. = i mil — V.S.

Sulphuric Acid (H2SO4 = 98.09). About i mil of the con- centrated acid is weighed in a tared weighing flask and found to weigh 1.8 gms. Fifty mils of water are added and then 2 drops of methyl orange, and the titration with normal potas- sium hydroxid begun, and cautiously continued until the

ANALYSIS BY NEUTRALIZATION 105

yellowish color of the solution indicates the completion of the operation. Note the number of mils of alkali solution used and apply the equation

H2SO4 + 2KOH = K2SO4 + 2H2O.

2)98.09 2)112.2 N

49.04 gms. = 56.1 gms. = 1000 mils — V.S.

Thus each mil of normal KOH V.S. represents 0.04904 gm. of pure H2SO4.

Aromatic Sulphuric Acid. This is an alcoholic solution containing free sulphuric acid and ethyl sulphuric together with a volatile dil. The latter as well as the alcohol present neutralizes alkali hydroxid, and must therefore be removed by heating before titration of the acid is undertaken. The U.S.P. method is as follows: Pour about 10 mils of aromatic sulphuric acid into a tared flask, stopper and weigh, transfer to a small flask using 60 mils of distilled water to rinse the weighing-flask, and, having connected the flask with a reflux condenser, boil the liquid for six hours; when cold, dilute it with distilled water to about 100 mils, and titrate with normal potassium hydroxid V.S., using methyl orange as indicator. Calculate as in preceding assay.

Phosphoric Acid (H3PO4 = 98.064). In the assay of phos- phoric acid by direct neutralization with standard KOH, the acid is converted into, first, KH2PO4, then K2HPO4, and finally into the normal K3PO4. We have no indicator which reliably shows the completion of the neutralization, i.e., the formation of the tribasic K3PO4. Litmus cannot

kbe used as indicator here for the dipotassic or disodichydric phosphate (K2HPO4 or Na2HPO4) which is formed is slightly alkaline towards litmus ; the same is true of most other indicators. Thompson, however, has demonstrated that this acid may '•

106 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

be accurately titrated by standard alkali when using either methyl orange or phenolphthalein, or both, successively.

If methyl orange is used the color changes upon the com- pletion of the formation of monobasic phosphate, KH2PO4, as per the following equation :

H3P04 + KOH = KH2P04 + H20.

If phenolphthalein is used, the color changes upon the com- pletion of the formation of the dibasic phosphate, KH2PO4,

As an example: A weighed quantity of the acid is diluted with water to measure 20 mils and sufficient pure sodium chlorid added to saturate the solution. Four drops of methyl

N orange are then introduced and the titration with — KOH

begun and continued until the red color changes to yellow, indicating the formation of the monobasic phosphate

H3PO4 + KOH = KH2PO4 + H2O. 98.064 gms. 56.1 gms. = 1000 mils — V.S.

Each mil of y KOH V.S. = 0.098064 gm. of H3PO4.

The use of sodium chlorid in this assay is to decrease the ionization of the acid salts produced in the reaction.

Another portion of the acid is treated in like manner, adding sodium chlorid and titrating, but using phenolphthalein as the indicator. The titration is continued until a faint permanent pink color appears. It is advisable to use heat, or better still, a standard alkali solution which is quite free

ANALYSIS BY NEUTRALIZATION 107

from CO 2- The end-reaction in this case marks the formation of the dibasic phosphate

H3PO4 + 2KOH = K2HPO4 + 2H2O.

2)98.064 2)112.2 N

49.032 gms. 56.1 gms. = 1000 mils — V.S.

N Each mil of — KOH V.S. = 0.049032 gm. of H3PO4.

N Just twice as much of the — KOH V.S. will be taken in

this assay as in the foregoing.

The two assays may be combined as follows: A weighed quantity of the acid is diluted with water saturated with sodium chlorid and titrated with the normal alkali V.S., using methyl orange as indicator until the red color of the solution changes to yellow. The number of mils is noted and multiplied by 0.098064 gm. A few drops of phenol- phthalein solution are now added and the titration continued until a pale-red color appears. The total number of mils of normal alkali used in the double titration is then multiplied by 0.049032 gm.

The U.S. P. method for the assay of phosphoric acid is given on page 127.

Hypophosphorous Acid (HPH2O2 = 66.06).

HPH2O2 + KOH = KPH2O2 + H2O.

N Each mil of — alkali represents 0.06606 gm. of HPH2O2.

Nitric Acid (HNO3 = 63.02).

= KNO3+H2O.

N Each mil of — alkali represents 0.06302 gm. of HNO3.

Hydrobromic and hydriodic acids may be estimated in

108 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

the same way as the foregoing, but it is usually preferred to estimate them by precipitation analysis. This is also true of phosphoric acid. Sulphurous acid is best assayed by oxida- tion with iodin.

Boric Acid (H3BO3 = 62.02). This acid is estimated by

N neutralization with — NaOH in the presence of a large

quantity of glycerin. (Thompson's Method, J. S. C. I., XII, 432). The addition of sufficient glycerin to a boric acid solution, so that no less than 30 per cent be present throughout the titration, develops the acidity of boric acid with regard to phenolphthalein to a great degree, and enables one to titrate direct with standard soda solution. One gm. of boric acid is dissolved in 50 mils of water; to this is added an equal volume of glycerin, then a few drops of phenolphthalein, and the titration with normal sodium hydroxid begun and con- tinued until a pink color appears.

N Each mil of — NaOH =0.06202 gm. of

H3BO3 + NaOH = NaH2BO3 + H2O.

62.02 40 gms. in 1000 mils — V.S.

N 0.06202 gm. = i mil — V.S.

Estimation of the Organic Acids

As the individual organic acids require different indicators, the table on page 28 should be consulted in the selection of an indicator for a particular organic acid. Phenolphthalein is, however, the most suitable for organic acids generally.

Acetic Acid (HC2H3O2 = 6o.o3). Mix 3 gms. of the acid with 50 mils of water, add a few drops of phenolphthalein T.S., and titrate with normal potassium hydroxid V.S. until

ANALYSIS BY NEUTRALIZATION 109

a permanent pale pink color is obtained, and apply the fol- lowing equation:

HC2H3O2 + KOH = KC2H3O3 + H2O. 60 56.1

N Thus 1000 mils of — KOH V.S. will neutralize 60 gms.

N of acetic acid; therefore each mil of — KOH V.S. represents

0.060 gm. of acetic acid.

If 1 8 mils are required to neutralize 3 gms. of the acid, it contains 18X0.06= 1.08 gms. of absolute acetic acid.

1.08X100

— = 36 per cent. o

Tartaric Acid (H2C4H4O6= 150.05). Dissolve 2 gms. of tartaric acid in 50 mils of distilled water, add a few drops of phenolphthalein and then pass into the solution from a burette

N

— potassium hydroxid V.S. until a faint pink tint is acquired

by the solution, and apply the equation

H2C4H406 + 2KOH - K2C4H406 + 2H2O.

2)£So_ N

75 gms. = 1000 mils — V.S.

Citric Acid (H3C6H5O7 = 192.06).

+ 3KOH =

3)192-06 N

64.01 gms. = looo mils — V.S.

N Each mil of — KOH represents 0.06401 gm. of citric acid.

Trichloracetic Acid (CC13COOH= 163.39).

CC13COOH + KOH - CC13COOK +H20. 163.39

110 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

N Each mil of -- KOH V.S. represents 0.16339 gm. of tri-

chloracetic acid.

The other organic acids are assayed in exactly the same manner as that described for the foregoing with the exception of benzoic and salicylic acids.

Benzoic Acid (HC7H5O2) and Salicylic Acid (HC7H5O3). These acids are assayed as follows: Dissolve 0.5 gm. of the acid, previously dried to constant weight in a desiccator over sulphuric acid, in 25 mils of diluted alcohol which has been previously neutralized with tenth-normal potassium hydroxid V.S., phenolphthalein T.S. being used as indicator. Titrate this solution with tenth-normal barium hydroxid V.S.,* using phenolphthalein as indicator.

The reactions are

2HC7H5O2 + Ba(OH)2 = Ba(C7H5O2)2 + 2H2O.

2)244.10

_

1 2. 205 = to 1000 mils — Ba(OH)2 V.S.

N Each mil of — Ba(OH)2 V.S. represents 0.012205 gm. of

benzoic acid and 0.013805 gm. of salicylic acid.

Lactic Acid (HC3H5O3 = 90.05). A definite weight is treated with an excess of normal potassium hydroxid V.S. and boiled for twenty minutes. The boiling solution is then titrated with normal sulphuric acid V.S. The quantity of the latter consumed deducted from the quantity (in mils) of normal potassium hydroxid V.S. originally added, gives the quantity of the alkali solution which represents the lactic acid being assayed.

*See page 101.

ANALYSIS BY NEUTRALIZATION

111

TABLE SHOWING QUANTITY OF SUBSTANCE TO BE TAKEN FOR ANALYSIS IN DIRECT PERCENTAGE ESTIMATIONS.

Molec- ular Weight.

Quantity to be taken so that , ., . N each mil of —

V.S. will rep- resent I %.

Percentage strength of Official Substance.

Ac

Ar Ar Ar Li Li

PC So

id acetic HC2H3O2.

60.03 62.02 210.08 80.93

36.47 127.93 66.04 90.05 63.01 126.05 98.06

6.03 gn

6.2

7-0 8.093

3-647 12.793 6.604 9-0 6.301 6-3

9! 806'" 4-903 4-903 7-5 1-6338 5-25 i-7 i-7 3-704 3-69 9-397 9.812

IO.OI

18.81 6.91 10.81 5-6i 5-6i

14.10 13-6 14.4 8.4 5-3

is.

36.

99-5 99-5

10.

32. 10. 30. 85- 68.

85-

94- 99-5 99.

10.

28. o. 14 98.5 98.5 99- 99.

99-5 99. 99. 85- 4-S

99. 99-5 99- 99.

boric H3BO3

citric H3C6H5O7+H2O

hydrobromic, HBr. Dil

hydrochloric HC1

hydriodic HI Dil

hypophosphorous, HPH2O2 lactic HC3H5O3

nitric HNO3

oxalic H2C-O4+2H2O

phosphoric, H3PO4 phosphoric with methyl orange

phosphoric with phenolphthalein sulphuric H2SO4

98.09 150.05 163.38 I57.II 17.03

I7-03 74.09 73-88 281.. 92 98.12 IOO. I

188.1 138.2 324-36 -56.1 56-1

282.2 136.0 144.0 84.0 106.0

tartaric H^HjOe

trichloracetic HC2C13O2

Qmonium carbonate N3HnC205

nmonia water NH3

nmonia water stronger. NH3.

me water Ca(OH)2

thium carbonate Li2CO3

" citrate LiaCeHsOy-f^^O

tassium acetate KC2H3O2

bicarbonate KHCO3

bitartrate KHCjH^Oe

carbonate K2CO3

citrate, K3C6H5O7+H2O hydroxid KOH.

hydroxid liquor, KOH

sodium tartrate KNaC4H4O6-t-4H2O d um acetate, NaC2H3O2+3H2O

' benzoate NaCvH5O2

' bicarbonate NaHCO3

' carbonate Na2CO3

Referring to the table it will be seen that if the quantities indicated are taken for analysis, the amount of standard solution required for substances of high percentage strength will be very large (in some cases over 99 mils), while for sub-

112 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

stances of low percentage strength, as for instance lime water, so small a volume of standard solution is required as to be unreadable (0.14 mil). It is therefore advisable to take for analysis a smaller quantity of high percentage substances and a larger quantity of such substances as contain a low percent- age. It is usually best to so adjust it that no less than 10 nor more than 30 mils of the standard solution be required. For example: In the case of citric acid, instead of taking for analysis 7 gms. it will be better to take one-fourth of this quantity, then each mil of the standard solution used will represent 4 per cent, and only one-fourth as much will be required, i.e., 24.9 mils instead of 99.5 mils. Again, in the case of lime water, if 37.04 gms. are taken instead of 3.704

N gms., 1.4 mils of the — standard solution will be required,

which is better than 0.14 mil, but in this case it will be still

/N\ better to use a decinormal ( — ) solution, then 37.04 gms. of

lime water would require for neutralization just 14 mils of

the — acid V.S. 10

If half the quantity indicated in the table is taken, then each mil of the standard solution would represent 2 per cent. If one- tenth the quantity is taken each mil will represent 10 per cent. If double the quantity is taken each mil will represent 0.5 per cent, etc.

CHAPTER IX ANALYSIS BY PRECIPITATION

THE general principle of this method* is that the deter- mination of the quantity of a given substance is effected by the formation of a precipitate, upon the addition of the standard solution to the substance under examination. There are three ways of determining the end-reaction in precipi- tation analyses:

1. By adding the standard solution until it ceases to produce any more precipitate, as in the estimation of silver by standard sodium chlorid, and the estimation of haloid salts and acids by means of standard silver nitrate. The application of this ending is almost limited to the above estimations, because in these only can accurate results be obtained. The silver halids formed are not only quite insoluble, but they have a tendency to curdle closely upon shaking (especially in acid solutions), and thus leave a clear supernatant liquid in which any further precipitation can readily be seen. Most of the other precipitates, such as barium sulphate, calcium oxalate, etc., although heavy and insoluble, are so finely divided and powdery that they do not readily subside.

2. By the use of an indicator, as in the estimation of haloid salts by means of standard silver nitrate solution, using neutral potassium chromate as the indicator. The latter is added to the haloid solution (which must be neutral), and the silver nitrate V.S. delivered into the mixture until a permanent red color (silver chromate) is produced. Silver nitrate reacts by preference with the halogen, and does not react with the chro-

113

114 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

mate until the halogen has been entirely precipitated. Hence the production of a permanent red color in the precipitate marks the completion of the precipitation of the halogen.

Another illustration is in the estimation of silver by sul- phocyanate solution, using ferric alum as indicator. The sulphocyanate produces with the silver a white precipitate of silver sulphocyanate, but when the precipitation of silver is complete the sulphocyanate reacts with the ferric alum present and a red ferric sulphocyanate appears and marks the end-point. On the other hand, the indicator may be used externally, i.e., alongside of the liquid being analyzed, a drop of the latter being brought in contact with a drop of the indi- cator at frequent intervals in the course of the titration, as in the estimation of phosphoric acid by means of uranium nitrate solution, in which potassium ferrocyanide is used as indicator.

3. By adding the standard solution until the first appearance of a precipitate, as in the estimation of cyanogen by silver nitrate solution, and the estimation of chlorin by mercuric nitrate V.S. In these estimations the standard solution is added to the solution of the substance under analysis until a precipitate appears.

Preparation of Decinormal ( — j Silver Nitrate (AgNOa

N = 169.89; — V.S. = 16.989 gms. in 1000 mils). Dissolve 16.989

gms. of pure silver nitrate* in sufficient water to make, at or near 25° C. (77° F.), exactly 1000 mils. One liter of this solution thus contains ytr of the molecular weight in grams of silver nitrate. It is therefore a decinormal solution.

If pure crystals of silver nitrate are not readily obtainable, and pure sodium chlorid is at hand, a solution of the silver

* This should be pulverized and dried at 1 20° C. for half an hour in a covered crucible before weighing.

ANALYSIS BY PRECIPITATION 115

nitrate may be made of approximate strength, a little stronger than necessary, and then standardized by means of the sodium chlorid, as follows: 0.11692 gm. of sodium chlorid is dissolved in distilled water, and a burette filled with the solution of silver nitrate to be standardized. The silver solution is now slowly added from the burette to the sodium chlorid solution contained in a beaker until no more precipitate of silver chlorid is produced.

If neutral potassium chromate is used as an indicator, the end of the reaction is shown by the appearance of yellowish- red silver chromate. This indication is extremely delicate. The silver nitrate does not act upon the chromate until all of the chlorid is converted into silver chlorid.

In the above reaction 20 mils of silver nitrate should be required. But since the silver nitrate solution is too strong, less of it will complete the reaction, and the solution must be diluted so that exactly 20 mils will be required to precipitate the chlorin in 0.11692 gm. of Nad.

Thus if 17 mils are used, each 17 mils must be diluted to 20 mils, or each 170 mils to 200 mils, or the entire remain- ing solution in the same proportion.

After dilution a fresh trial should always be made.

Silver nitrate solution should be kept in dark amber- colored, glass-stoppered bottles, carefully protected from dust.

The U.S.P. IX directs the use of pure metallic silver (foil, wire, or powder) for preparing this solution. Into a flask con- taining 10.788 gms. of pure silver introduce gradually about 30 mils of nitric acid (sp.gr. about 1.403 at 25° C.) or sufficient for complete solution. Thoroughly rinse the neck of the flask and the funnel which is used. Evaporate the solution to dryness, carefully protecting it from dust. Then, after drying in an air oven at about 120° C., for ten minutes, dissolve the

116 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

residue in sufficient distilled water to measure exactly 1000 mils at 25° C.

N N

Decinormal — Sodium Chlorid (NaCl = 58.46; — V.S.

= 5.846 gms. in 1000 mils). Dissolve 5.846 gms. of pure sodium chlorid in enough water to make exactly 1000 mils at the standard temperature.

Check this solution with decinormal silver nitrate. The two solutions should correspond, volume for volume.

Pure Sodium Chlorid may be prepared by passing into a saturated aqueous solution of the purest commercial sodium chlorid a current of dry hydrochloric acid gas. The crystal- line precipitate is then separated and dried at a temperature sufficiently high to expel all traces of free acid.

N The method of standardizing — NaCl solution is as follows :

0.33978 gm. of silver nitrate is dissolved in 10 mils of dis-

N tilled water, and the solution carefully titrated with — NaCl

V.S. until precipitation ceases. Twenty mils of the standard solution should be required.

AgN03 + NaCl = AgCl + NaNO3.

10)169.89 10)58.46 N

16.989 gms. 5.846 gms., or 1000 mils — NaCl V.S.

Each mil of the standard solution represents 0.016989 gm. of pure AgNO3.

0.016989X20=0.33978 gm. -

— - — =100 per cent. 0.33978

This solution may also be standardized by residual tira- tion with Volhard's solution.

ANALYSIS BY PRECIPITATION 117

N Decinormal — Potassium Sulphocyanate (Volhard's Solu-

N tion) (KSCN = 97.18; — V.S. = 9.718 gms. in 1000 mils).

Dissolve 10 gms. of pure crystallized potassium sulphocyanate (thiocyanate) in 1000 mils of water.

This solution, which is too concentrated, must be adjusted so as to correspond exactly in strength with decinormal silver nitrate V.S. For this purpose introduce into a flask 20 mils

N of — AgNO3 V.S., 3 mils of ammonioferric sulphate solution,

and 5 mils of diluted nitric acid (10 per cent and free from nitrous compounds).

Dilute the liquid with 75 mils of distilled water, and titrate it with the sulphocyanate solution.

At first a white precipitate of silver sulphocyanate is produced, giving the fluid a milky appearance, and then as each drop of sulphocyanate falls in it is surrounded by a deep brownish-red cloud of ferric sulphocyanate, which quickly disappears on shaking, as long as any of the silver nitrate remains unchanged.

When the point of saturation is reached and the silvei has all been precipitated, a single drop of the sulphocyanate solution produces a faint brownish-red color, which does not disappear on shaking.

Note the number of mils of the sulphocyanate solution used, and dilute the whole of the remaining solution so that equal volumes of this and of the decinormal silver nitrate will be required to produce the permanent brownish-red tint. (The same tint of brown or red to which the volumetric solution is adjusted must be attained when the solufion is used in volumetric testing.)

Assuming that 19 mils of the sulphocyanate solution were

118 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

required to produce the reaction, then each 19 mils must be diluted to make 20 mils, or the whole of the remaining solution in the same proportion.

Always make a new trial after the dilution to see if the

N solutions correspond, e.g., 50 mils of — silver nitrate are taken,

and 5 mils of ammonioferric sulphate, 5 mils of pure nitric acid and 200 mils of water are added, and there should be required exactly 50 mils of the potassium sulphocyanate solu- tion. The same depth of reddish-brown tint should be ob- tained in all assays by this method, as is obtained in standard- izing the solution.

Estimation of Soluble Haloid Salts

The estimation of these salts is based upon the powerful affinity existing between the halogens and silver, and the ready precipitation of the resulting chlorid, bromid and iodid. Standard solution of silver nitrate is used for this purpose, and for the sake of exactness and convenience, is made of decinormal strength. In some cases it is advisable to use centinormal solutions.

Mohr's Method with Chromate Indicator. This method is the best to use, if the haloid salts are in neutral solution, and salts of lead, bismuth, barium or iron are absent. If the solution is acid the indicator is inadmissable, in that acids have a solvent action upon silver chromate and thus prevent the end-reaction from being clearly and accurately observed. If the above-mentioned metals are present, the indicator is likewise useless, as these bases form insoluble, highly colored compounds with the chromate. The neutral potassium chro- mate (yellow chromate) which is used as the indicator must be

ANALYSIS BY PRECIPITATION 119

free from chlorid * and should be used in the form of a 10 per cent solution.

In the volumetric analysis of soluble haloid salts (chlorids, bromids and iodids) 0.5 gm. of the well-dried salt is dissolved in 40 mils of water in a beaker. This is placed upon a white surface and a few drops of the chromate indicator (or suffi- cient to give the solution a pale yellow tint), added. The

N decinormal -- silver nitrate solution is then added cautiously

from a burette, stirring constantly until a permanent red tint is produced. The red tint is due to the formation of silver chromate, which does not appear permanent until the last trace of halogen has been precipitated. The reactions are as follows :

NaCl + AgN03 = AgCl + NaNO3 and

K2CrO4 + 2 AgNO3 = Ag2CrO4 + 2KNO3.

If the solution to be estimated is acid it should be accu- rately neutralized with ammonia, or sodium or calcium car- bonate. If it is alkaline in reaction it should likewise be neutralized, using acetic acid for this purpose. •

In the estimation of bromids and iodids it must not be forgotten to take into account the invariable presence of chlorids as an impurity.

The method in detail is exemplified in the following assay :

Estimation of Sodium Chlorid. One gm. of the well-dried sodium chlorid is dissolved in sufficient distilled water to measure 100 mils. Of this solution 10 mils (representing o.i

"The presence of chlorid in the chromate solution may be determined by adding a small quantity of silver nitrate solution, and then some nitric acid. If the red precipitate dissolves completely and leaves a clear solution, chlorid is absent. If it is found to be present it may be removed by the addition of a few drops of silver nitrate solution, and filtering, without using any nitric acid,

120 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

gm. of the salt) is taken, a few drops of neutral potassium

N

chromate solution added, and then the - - silver solution

10

delivered from a burette with constant stirring or shaking until the chlorid is entirely precipitated, as evidenced by the formation of a permanent red color (silver chromate). The equation is

NaCl + AgN03 = AgCl + NaNO3.

10)58.46 10)169.89 N

5.846 16.989 gms. = looo mils — V.S.

N Thus each mil of — V.S. represents 0.005846 gm. of NaCl.

If in the above assay 17 mils of the silver solution were re- quired, then 17 Xo.oo5846 gm. =0.099382 gm. or 99.382 per cent.

0.099382 X 100

— = 99.382 per cent.

Titration without an Indicator — Gay-Lussac's Method. In

this method no indicator is used, the standard solution being added until it ceases to produce any further precipi- tation. This method is applicable to acid solution of the haloid salts, and to the haloid acids — hydrochloric, hydro- bromic and hydriodic; also to the estimation of silver by standard solution of sodium chlorid. The method is carried out in hot solutions, slightly acidulated with nitric acid, in order to facilitate the precipitation of the silver halid. The haloid acids are neutralized with an alkali and then slightly acidulated with nitric acid before the titration is begun. The calculations are precisely like those in the foregoing assays.

Volhard's or Sulphocyanate Method. This method depends upon completely precipitating the halogen in the presence of nitric acid, by a measured excess of standard silver nitrate solution, and then estimating the excess of silver by retitrating

ANALYSIS BY PRECIPITATION 121

with standard sulphocyanate solution, using ferric alum as an indicator.

The sulphocyanate has a greater affinity for silver than it has for iron, and therefore, so long as any silver is in solu- tion, the sulphocyanate will combine with it and form a pre- cipitate of silver sulphocyanate.

As soon as the silver is all taken up, the sulphocyanate will combine with the ferric alum and strike a brownish-red color.

The sulphocyanate solution is to be made of such strength that it corresponds with the silver solution, volume for volume.

The difference between the volume of silver solution origi- nally added and the volume of sulphocyanate solution used, will give the volume of silver solution equivalent to the haloid salt present.

This method has the advantage over the direct method for haloids with chromate indicator, in that it may be used in the presence of nitric acid. It thus enables one to estimate the haloids hi the presence of phosphates or other salts which precipitate silver in neutral but not in acid solutions, and also in that the presence of barium, bismuth, lead, iron and other metals do not interfere, as they do with the chromate in Mohr's method. The presence of mercury, however, exerts a disturbing influence upon the end-reaction. The nitric acid acidulates the solution and thus facilitates the precipitation of silver by the halogens, and prevents its precipitation by other substances. The quantity of nitric acid employed is of no great importance, except in the case of iodids (because silver iodid is slightly soluble in nitric acid). Usually suffi- cient of the acid is added to just remove the color produced by the indicator. A very large excess of the acid would, however, interfere with the proper determination of the end-

122 ANALYSIS BY OXIDATION AND REDUCTION

reaction, in that it to a slight extent prevents the formation of ferric sulphocyanate. In the estimation of iodids by this method, the nitric acid should be added after the standard silver solution, while in the case of the other haloid salts the acid may be added before.

The indicator also should be added after the standard silver solution, when estimating iodids, because being a ferric salt it is, like nitric acid, capable of liberating iodin.

The solutions required for this method are: (I) Decinormal Silver Nitrate (page 114); (II) Decinormal Potassium Sulphocyanate (page 117);

(III) Ferric Alum Solution. (The indicator.)

This is a 10 per cent aqueous solution of ferric-ammonium sulphate, FeNH4(SO4)2 + i2H2O.

(IV) Nitric Acid (C.P.). This must be free from nitrous acid. If it or any of the lower oxids of nitrogen are present they may be removed by diluting with one-fourth part of water and boiling until colorless.

The process is exemplified in the following assays :

Estimation of Potassium Bromid. Dissolve about 0.4 gm. of the salt accurately weighed in 25 mils of distilled water in a glass-stoppered flask. Add 50 mils of tenth-normal silver nitrate V.S., then add 2 mils of ammonio-ferric sulphate T.S. and 2 mils of nirtic acid, shake well and finally determine the excess of the silver nitrate V.S. by titrating with tenth-normal potassium sulphocyanate V.S. until a permanent reddish tint pervades the supernatant liquid.

The difference between the number of mils of tenth-normal silver nitrate V.S. added and the number of mils of tenth- normal potassium sulphocyanate used, multiplied by 0.011902 gm., gives the weight of pure KBr in the quantity of salt taken for analysis,

ANALYSIS BY PRECIPITATION 123

Assuming that in the foregoing analysis 16.9 mils of the sulphocyanate solution were used, then 50-16.9 = 33.1 mils of the silver nitrate solution reacted with the potassium bromid. Hence 0.011902X33.1=0.3939 gm. of KBr, or 98.5 per cent. The reactions are

KBr + AgNO3 = AgBr + KNO3.

10000)119.02 £*jj

.011902 gm. = i mil of — AgNOs V.S.

10

lodids are estimated in exactly the same manner.

When this method is used for the estimation of chlorids, however, the precipitated silver chlorid must be removed by nitration, because of the action of ferric sulphocyanate upon silver chlorid, which causes the results of the analysis to be too high.

In the case of silver bromid or iodid no such reaction takes place, or if it does, the reaction is so slow as not to interfere in the least with the getting of accurate results. Therefore, when this method is used for the determination of bromids or iodids, there is no need for filtering to remove the pre- cipitate.

The procedure in detail is exemplified by the following:

Estimation of Sodium Chlorid. Dissolve about 0.2 gm. of the salt accurately weighed (or a near equivalent of its solution) in 25 mils of distilled water in a 2oo-mil graduated flask. Add 50 mils of tenth-normal silver nitrate V.S. to the solution, and after adding 5 mils of nitric acid, add sufficient distilled water to make 200 mils and mix well. Then filter through a dry filter into a dry flask, and reject the first 20 mils of the filtrate. Collect the next 100 mils of filtrate, add 2 mils of ammonio-ferric sulphate T.S. and titrate it with tenth-normal potassium sulphocyanate until a permanent red- dish supernatant liquid is obtained. Multiply the number of

124 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

mils of the potassium sulphocyanate solution required by 2 and subtract this figure from 50 (the number of mils of tenth- normal silver nitrate V.S. added). The difference multiplied by the decinormal factor of sodium chlorid, 0.005046 gm., gives the weight of pure NaCl in the amount of salt taken for analysis.

Estimation of Haloid Acids

These acids, namely, hydrochloric, hydrobromic and hydriodic, may be estimated by Gay-Lussac's method above described, or they may be estimated by Mohr's Method. using neutral potassium chromate as an indicator. In this case it is necessary to carefully neutralize the acid with

N ammonia and then titrate with — silver nitrate solution,

using a few drops of chromate as indicator, in the manner described in • the foregoing assays. They may also be esti- mated by Volhard's Method, in which an- excess of the standard silver nitrate solution is used, in the presence of nitric acid, and the amount of the excess determined by residual titration with potassium sulphocyanate, using ferric alum as the indi- cator. This method is especially useful for iodids and hydri- odic acid, in that the nitric acid need not be added until after an excess of silver nitrate solution is used, and thus liberation of iodin by the nitric acid avoided.

The estimation of the haloid acids may also be effected by neutralization with standard alkali, in the same way as other acids, but since hydrobromic and hydriodic acids are now frequently prepared by the method of Fothergill, in which potassium bromid or potassium iodid (according to the acid to be made) is brought in contact with tartaric acid (as shown in the equation), an excess of the latter acid is

ANALYSIS BY PRECIPITATION 125

unavoidably present, and hence the neutralization method is inapplicable.

KI + H2C4H406 = KHC4H406 + HI

Potassium Tartaric acid Potassium Hydriodic

lodid Bitartrate Acid

KBr " " HBr

Assay of Hydrobromic Acid, Using Chr ornate as Indicator.

Ten gms. of hydrobromic acid are diluted with sufficient dis- tilled water to make 100 mils. Ten mils of this solution, representing i gm. of the acid, is exactly neutralized with diluted ammonia water (using litmus solution as indicator) ; 3 drops of neutral potassium chromate solution are added,

N

and then the — silver nitrate run in from a burette until the 10

solution acquires a permanent red tint. The following equa- tion is then applied :

HBr + AgNO3 = AgBr + HNO3.

10)80.92 N

8.092 gms. = 1000 mils — V.S.

10

If the assay is to be made by the direct percentage method, 8.09 mils of the solution (10 gms. in 100 mils) (representing 0.809 gms. of the acid) should be taken, in which case each mil of the standard silver solution consumed will at once indi- cate i per cent.

Hydriodic and hydrochloric acids may also be assayed by this method.

Assay of Hydriodic Acid by the Sulphocyanate Method. Introduce into a 2oo-mil stoppered flask 2 gms. of the acid, add 50 mils of distilled water and 25 mils (accurately measured) of decinormal silver nitrate, shake thoroughly, and then add 3 mils of nitric acid, and heat on a water-bath until the pre-

126 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

cipitate has a bright yellow color. Then cool and add 2 mils of the ammonio-ferric sulphate T.S., and finally the decinormal potassium sulphocyanate run in slowly from a burette, until a permanent reddish-brown tint is produced. Note the num- ber of mils of sulphocyanate solution employed.

Deduct this from the 25 mils of silver solution added, and multiply the remainder by the factor for HI, which is 0.012793.

I. HI + AgNO3 = Agl + HNO3.

10)127.93 10)169.89 N

12.793 gm^. 16.989 gms. = 1000 mils — V.S.

N

0.012793 gm. of HI = i mil — V.S. 10

II. AgN03 + KSCN = AgSCN + KNO3.

III. 2FeNH4(S04)2+6KSCN

= 2Fe(SCN)3 + (NH4)2S04 +3K2S04.

The reddish-brown color which marks the end-reaction is due to the formation of Fe(SCN)3 ferric sulphocyanate.

Assuming that in the above titration 10.2 mils of decinormal sulphocyanate were employed, then 25 mils -10.2 = 14.8 mils.

0.012793X14.8 = 0.1893 gm.

0.1893X100

— = 9.46 per cent.

Assay of Syrup of Hydriodic Acid. Twenty-five mils of the syrup are introduced into a stoppered flask, 100 mils of distilled water are added, followed by 40 mils of decinormal silver nitrate and the mixture thoroughly shaken. Five mils of diluted nitric acid and 3 mils of the ferric alum solution are now added, and after again shaking the mixture it is titrated with decinormal potassium sulphocyanate until a per- manent reddish-brown tint appears.

ANALYSIS BY PRECIPITATION _ 127

OTHER ASSAYS BY VOLHARD'S METHOD

Assay of Syrup of Ferrous lodid U.S.P. Dilute about 10 gms. of the syrup, accurately weighed, with 30 mils of dis- tilled water, add 50 mils of tenth-normal silver nitrate V.S. and 5 mils of nitric acid and heat on a water-bath until the precipitate of silver iodid is yellow. Cool, add 2 mils of ferric ammonium sulphate T.S. and determine the residual silver nitrate by titration with tenth-normal potassium sulphocyanate V.S. A permanent reddish-brown tint marks the end point. Not less than 16 mils nor more than 19.3 mils of the latter should be required.

The quantity of tenth-normal sulphocyanate so used de- ducted from 50 mils (the quantity of tenth-normal silver nitrate V.S. taken in the assay) gives the quantity of the latter (in mils) which reacted with the ferrous iodid.

Each mil of tenth-normal silver nitrate V.S. corresponds to 0.015484 gm. of FeI2.

FeI2 + 2AgN03 = 2AgI + Fe(NO3)2. 2)300.68 10)154.84 N

15.484 gms. = 1000 mils — AgNO3 V.S.

Assuming that in the foregoing assay 16 mils of the tenth- normal sulphocyanate were used, then 50 mils - 16 mils = 34 mils.

0.015484X34 = 0.520456 gm.

0.520456X100

= 5.20 per cent.

10

Assay of Phosphoric Acid U.S.P. The neutralization method for the assay of phosphoric acid has been discarded in the U.S.P. IX, and the precipitation with standard silver nitrate is recommended. An excess of the latter solution is added to

128 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

the neutralized phosphoric acid, and the residual silver solution determined by titration with standard sulphocyanate.

The procedure is as follows : An accurately weighed quantity of the acid (about o.i gm.) is diluted with 10 mils of distilled water, in a loo-mil flask. One drop of phenolphthalein T.S. is added, and the solution carefully neutralized with special potassium hydroxid solution, i.e., one free from chlorid.

Fifty mils of tenth-normal silver nitrate V.S. are now added, the mixture agitated and zinc oxid (free from chlorid) gradually added in small portions at a time, until the liquid is neutral to litmus paper. It is then diluted with sufficient distilled water to measure 100 mils, again thoroughly agitated and filtered through a dry filter, collecting 50 mils of filtrate. To this 50 mils of filtrate is added 2 mils of nitric acid and 2 mils of ammonio-ferric sulphate T.S. and then titrated with tenth- normal potassium sulphocyanate V.S. to the production of a permanent red color.

The number of mils of the latter multiplied by 2 is sub- tracted from 50 mils (the quantity of tenth-normal silver nitrate V.S. added) and the difference is the quantity of the latter which reacted with the phosphoric acid.

Each mil of tenth-normal silver nitrate V.S. represents 0.0032687 gm. of H3PO4.

NOTES. The neutralization of phosphoric acid by potas- sium hydroxid in the presence of phenolphthalein is repre- sented by the equation

H3P04 + 2KOH = K2HP04 + 2H20.

An excess of KOH must be avoided, as this reacts with silver nitrate.

The reaction between silver nitrate and the dipotassic phosphate is usually represented by the following equation,

ANALYSIS BY PRECIPITATION 129

which, for the purposes of this assay, may be considered suf- ficiently accurate :

K2HPO4 + 3AgN03 = Ag3P04 + 2KNO3 + HNO3.

Normal silver phosphate being precipitated.

It is very probable that the supernatant liquid in this reac- tion contains also some free phosphoric acid, but that seems to be immaterial.

3K2HPO4 + 6AgNO3 = 2 Ag3P04 4- 6KNO3 + H3PO4.

The zinc oxid is added in this assay for the purpose of neutralizing the acid or acids liberated in the interaction. This is necessary because silver phosphate (Ag3PO4) is soluble in dilute nitric and phosphoric acids. Too large an excess of zinc oxid must be avoided, as it acts upon silver nitrate, precipitating Ag2O if left in contact too long. It is advisable for this reason to filter as soon as possible after the addition of zinc oxid.

If, as is claimed,* that phosphoric acid is liberated by the interaction of dipotassic phosphate and silver nitrate, the zinc oxid will form with it, zinc phosphate which in a neutral solution is transpqsed by the silver nitrate into silver phosphate, and zinc nitrate, thus converting all of the phosphoric acid into Ag3PO4.

Zinc oxid appears to be the most suitable substance avail- able for this neutralization, because its action upon silver nitrate is very slow and insignificant, if not in large excess.

Assay of Sodium Phosphate. This may be performed by the same method.

Into a loo-mil graduated flask, introduce an accurately- weighed quantity of the salt (about 0.4 gm.), dissolve it in

* J. Rasin, J. A. C. S., Vol. 33, page 1103.

130 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

10 mils of distilled water, add 50 mils of tenth-normal silver nitrate V.S. and shake well. Then add gradually pure zinc oxid in small portions, until the liquid is neutral to litmus. Dilute the mixture with distilled water to make 100 mils. Shake thoroughly and filter through a dry filter. Reject the first 20 mils of filtrate. Collect 50 mils, add to it 2 mils of nitric acid and 2 mils of ammonio-ferric sulphate T.S. and titrate with tenth-normal potassium sulphocyanate V.S. to the production of a permanent red color.

Each mil of tenth-normal silver nitrate V.S. used corre- sponds to 0.004735 gm. of Na2HPC>4. The calculations are as given in preceding assay.

Assay of Hypophosphites. This is performed by the method just described for phosphoric acid and sodium phosphate. The hypophosphites being first oxidized to phosphates. Dis- solve about i gm. (accurately weighed) of the salt in 10 mils of distilled water and 10 mils of nitric acid and evaporate the solution to dryness, on a water-bath. Then add 5 mils of nitric acid and again evaporate it to dryness on a water- bath. Dissolve the residue in 20 mils of distilled water, add a drop of phenolphthalein T.S., and sufficient special potassium hydroxid T.S. (free from chlorids) to produce a pale pink color, and then dilute the mixture to exactly 100 mils, with distilled water. Transfer 10 mils of this solution to a 100- mil flask, add 50 mils of tenth-normal silver nitrate V.S. and proceed from this point as directed under assay of sodium phosphate.

Each mil of tenth-normal silver nitrate V.S. corresponds to

0.002836 gm. of Ca(PH2O2)2; 0.003472 gm. of KPH2O2; 0.0035357 gm. of NaPH2O2.

ANALYSIS BY PRECIPITATION 131

Estimation of Cyanogen

Titration with Standard Silver Solution to First Appearance of a Precipitate— Liebig's Method. This gives fairly accurate results. The cyanogen must be in the form of an alkali salt and in an alkaline solution. If hydrocyanic acid is to be estimated, it must be made alkaline by the addition of potas- sium or sodium hydroxid. The standard silver solution is then added cautiously and with constant stirring until a per- manent precipitate of silver cyanid is produced. When silver nitrate is added to an alkaline solution of a cyanid, the precipitate which at first forms redissolves on stirring and a soluble double cyanid (AgCN,KCN or AgCN,NaCN, depending upon the alkali used) is formed, and when all of the cyanid has been taken up, the further addition of silver nitrate causes a decomposition of this soluble double salt and the formation of a permanent precipitate of silver cyanid. Therefore, the iqxrst appearance of this precipitate affords a delicate proof of the completion of the reaction.

These equations illustrate the reactions:

2NaCN + AgNO3 = AgCN,NaCN + NaNO3.

Double cyanid of silver and sodium

AgCN,NaCN+AgNO3 =2AgCN+NaNO3.

Silver cyanid

According to these equations it is seen that the end- reaction is reached when two molecules of the alkali cyanid have reacted with one molecule of silver nitrate. The slightest excess of silver nitrate above this quantity brings about a decomposition of the double salt and a precipitation of the silver cyanid, as above stated.

This double combination is so firm that if the estimation is done in the presence of a halogen, no permanent precipitate

132 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

of silver halid is formed until after all of the cyanogen present has been converted into a double salt. This fact is taken advantage of in the processes for hydrocyanic acid and alkali cyanid in which potassium iodid is employed as indicator in. the presence of ammonia water. The latter prevents the precipitation of silver cyanids and thus allows the silver iodid to precipitate alone.

N i mil of — AgNO3 V.S. =0.005204 gm. CN;

0.005404 gm. HCN; 0.0098 gm. NaCN; 0.0130 gm. KCN.

Titration with Standard Silver Solution, Using Chromate Indicator — Vielhaber's Method. This method is especially recommended for the assay of weak solutions containing hydrocyanic acid, as bitter almond oil, bitter almond water, cherry laurel water, etc., but it may also be employed for alkaline cyanids.

A sufficient quantity of an aqueous suspension of mag- nesium hydroxid * to make the solution opaque and distinctly alkaline is added; this is followed by a few drops of potassium

N

chromate indicator and then the — silver nitrate delivered

10

into the mixture from a burette until a permanent red tint appears, as in the titration of haloid salts. The method is a very satisfactory one if chlorids are absent.

The reactions in this method are the same as in the fore- going, but the end-reaction (the production of silver chromate) does not occur until the double cyanid is completely decom-

* Calcined magnesia triturated with water.

ANALYSIS BY PRECIPITATION 133

posed, at which point the addition of another drop of silver solution reacts with the chromate and produces the red pre- cipitate (silver chromate).

The equations are as follows: Sodium is used in the equations instead of magnesium in order to make the explana- tion clearer.

(a) 2NaCN+AgNO3 = AgCN,NaCN + NaNO3 = (2HCN);

(b) AgCN,NaCN + AgNO3 = 2 AgCN + NaNO3.

These equations show that it requires two molecules of silver nitrate to completely precipitate two molecules of cyanid. 169.89 gms. of AgNO3 is equivalent to 27.01 gms. of HCN, while by Liebig's method 169.89 gms. of AgNO3 is equivalent to 54.02 gms. of HCN.

N i mil — AgNO3 V.S. = 0.002602 gm. CN;

0.002702 gm. HCN; 0.004902 gm. NaCN; 0.006511 gm. KCN.

Example. 1.35 gms. of the diluted acid is mixed with enough water and magnesia to make an opaque mixture of about 10 mils. Add to this 2 or 3 drops of potassium chromate solution and then from a burette deliver the decinormal silver nitrate V.S. until a red tint is produced which does not dis- appear by shaking.

This method is exemplified in the U.S.P. IX Assay of Oil of Bitter Almonds.

Titration with Standard Silver Solution, Using Potassium lodid as Indicator. This method, employed in the Assay of Dilute Hydrocyanic Acid, U.S.P. IX, is conducted as follows:

Into a loo-mil flask containing 25 mils of distilled water and 5 mils of potassium hydroxid T.S., which has been tared,

134 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

pour about 5 mils of the acid, and weigh it accurately. Then add 3 drops of potassium iodid T.S. and titrate with tenth- normal silver nitrate V.S. to the production of a slight per- manent precipitate.

Each mil of tenth-normal silver nitrate V.S. used, corre- sponds to 0.005404 gm. of HCN. The slight precipitate which is produced in this assay and which marks the end reaction, is silver iodid (Agl).

The precipitate of silver iodid does not appear until the cyanid has been entirely converted into the double salt (AgCN.KCN).

Estimation of Potassium Cyanid (KCN = 65.11). One gm. of potassium cyanid is dissolved in sufficient distilled water to make 100 mils, then 65.11 mils of this solution mixed with 5 mils of ammonia water and 3 drops of potassium iodid

N solution are titrated with — AgNOs V.S. until the appearance

of a permanent precipitate. Each mil indicates 2 per cent, i mil of — AgNO3 V.S. =0.0130 gm. KCN.

Estimation of Silver Salts

Soluble silver salts may be estimated by direct titration with standard sodium chlorid, the process being exactly the converse of the precipitation methods for halogens. The stand- ard sodium chlorid solution is added to the solution of the silver salt until precipitation ceases; or the titration may be done in the presence of chromate indicator, the end-point being then known, to be reached when the red color of the silver chromate disappears. The first of these methods is im- practicable, too much time being consumed in waiting for the precipitate to settle so as to render the supernatant liquid

ANALYSIS BY PRECIPITATION 135

sufficiently clear to recognize whether a precipitate is produced in it by the further addition of the standard solution.

If chromate indicator is used, the end-point is easily over- stepped, because of the slow decomposition of the silver chromate by the chlorid. It is better to add an excess of sodium chlorid solution and then retitrate with standard silver nitrate solution until the red color appears.

Silver salts may also be titrated by means of standard sulphocyanate solution, using ferric alum as indicator.

Assay of Silver Nitrate by Means of — Sulphocyanate.

This method, as applied to the assay of halogen compounds, is described in the preceding pages. The great advantage which this method presents over the others, is that the presence of most other metals does not interfere. The only metal which does materially interfere with the determination of silver is mercury.

Example. Weigh accurately about 0.8 gm. of the silver salt, and dissolve it in 50 mils of distilled water, add 2 mils of nitric acid and 2 mils of ferric-ammonium sulphate T.S. and titrate the mixture with tenth-normal potassium sulpho- cyanate V.S. until a permanent reddish color of ferric sulpho- cyanate appears.

The following equation explains the reactions:

AgNO3 + KSCN = AgSCN + KNO3.

10)169.89 10)97.18

16.989 gms. 9.718 gms. or 1000 mils standard V.S.

Thus each mil of the standard V.S. represents 0.016989 gm. of pure silver nitrate, or 0.010788 gm. of metallic silver.

Assay of Metallic Silver and Silver Alloys. A quantity of the metal, weighing about 0.5 gm., is dissolved in 10 mils of nitric acid, and after complete solution is attained it is

136 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

heated sufficiently to drive off all traces of nitrous acid. The solution is then diluted with about 100 mils of distilled water and assayed by one of the methods described under the assay of silver nitrate. The sulphocyanate method is the preferred one.

Assay of Silver Oxid. About 0.5 gm. is dissolved in 3 mils of nitric acid and 50 mils of distilled water, 2 mils of ferric- ammonium sulphate T.S. are added and the mixture titrated with a standard sulphocyanate solution.

Assay of Mercuric Salts by Direct Titration with Sulpho- cyanate Solution. This method is applicable to most mercuric .compounds in the absence of halogens (especially chlorin). It consists in titrating the mercuric salt in the presence of nitric acid and a large volume of distilled water with a standard sulphocyanate solution.

Ferric ammonium sulphate is used as indicator. The assay is carried out as follows :

Assay of Mercuric Oxid. Dissolve an accurately weighed quantity (about 0.5 gm.) of the mercuric oxid in 10 mils of distilled water and 5 mils of nitric acid. Then further dilute the solution with 150 mils of distilled water. To this add 2 mils of ferric-ammonium sulphate T.S. and titrate with tenth- normal potassium sulphocyanate V.S. to the production of a permanent yellowish-red color. Each mil of the sulphocyanate solution used corresponds to 0.01083 gm. of HgO.

This process, while not strictly a precipitation analysis, is for the sake of convenience included under this heading. The mercuric sulphocyanate formed in the process is so sparingly soluble in water that if a high dilution of the salt being titrated is not made, a heavy white precipitate would be produced.

ANALYSIS BY PRECIPITATION

137

TABLE OF SUBSTANCES ESTIMATED BY PRECIPITATION

Name.

Formula.

Molecular Weight,

Standard Solution Used.

Factor.*

Acid hydrobromic

HBr

go Q-I •

N — AeNO,

" hydrochloric

HC1

•?6 4.7

10 ngiN^3

N — AeNO,

" hydrocyanic it «

HCN HCN

27.02

10 g 3

N -AgN03

without indicator

N

AaNO

0.005404

U «(

HCN

27 O2

10 Ag Us

with chromate indicator

N

A<rNO

" hydriodic

HI

I0 AgiNU3

with iodid indicator

N — AffNOo

o . 005404

Allyl-iso-thiocyanate

CSNC3H5

127.93 00- 12

10 wpy*

(i

0.004956

Ammonium bromid

NH4Br

07 06

(i

0.009796

" chlorid . . .

NH4C1

C7. eo

<t

o . 00535

" iodid

NHJ

144.96

«

0.014496

Calcium, bromid

CaBr2

IQO Q3

«

o . 009996

" chlorid

CaCl2

Ill OI

u

0.00555

Ferrous bromid '.

FeBr2

215 66

<«

0.010784

" iodid

Felo

•JOG 66

u

0.015484

Lithium bromid .

LiBr

86 86

U

0.008686

Potassium bromid

KBr

119.02

U

0.011902

" chlorid

KC1

1 A. C6

u

o 0074^6

cyanid

KCN

6c 1 1

N — AgNO,

o 013022

« «

KCN

6c ii

10 Agi>W3

without indicator

N — AeNO,

o 006511

« tt

KCN

6c 1 1

10 rtgi>^3

with chromate indicator

N — AeNO,

o 013022

10 g 3 with iodid indicator

* This is the coefficient by which the number of cc. used of the decinormal solu- tion is to be multiplied in order to obtain the quantity of pure substance in the sample analyzed. It represents the weight of the substance precipitated by i cc. of the decinormal solution.

138 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

TABLE OF SUBSTANCES ESTIMATED BY PRECIPITATION—

Continued

Name.

Formula.

Molecular Weight.

Standard Solution Used.

Factor.*

Potassium iodid

KI

166 02

— AeNO,

o 016602

sulphocyanate . . Silver (metallic)

KSCN Ae,

97.18 2X107 88

10 ^S^^3

«

N - NaC1 or

0.009718 o 010788

" nitrate

AeNO,

169 89

10

— KSCN

10

o 016989

" oxid

AR,O

211 76

«

o 011588

Sodium bromid

NaBr

IO2 92

N - A.eNO,

o 010292

" chlorid

NaCl

q8 46

10

o 00^846

" iodid

Nal

140 02

«

o 014002

Strontium bromid

SrBr2

247 47

«

o 012^7^

" iodid

Sri,

341 47

«<

o oi7<;7'?

Zinc bromid

ZnBr2

22C 21

«

o 011260

" chlorid

ZnCl2

I "?6 2O

M

o 006814

" iodid *.

ZnI2

3IQ 21

M

o 011:060

CHAPTER X ANALYSIS BY OXIDATION AND REDUCTION

AN extensive series of analyses are made by these methods with extremely accurate results; in fact, the results are generally more accurate than those obtained by gravimetric methods.

The principle involved is exceedingly simple. An oxidizing agent is employed for the estimation of an oxidizable sub- stance, and likewise a reducing agent is employed for the estimation of a reducible substance. Oxidizing agents are always reducible and reducing agents always oxidizable. An oxidation and a reduction take place at the same time, i.e., the oxidizing agent is itself reduced in the operation and the reducing agent is at the same time oxidized.

Thus substances which are capable of absorbing oxygen or are susceptible of an equivalent action may be accurately estimated by subjecting them to the action of an oxidizing agent of known power, and from the quantity of the latter required for complete oxidation the weight of the oxidizable substance is ascertained.

Example. Ferrous oxid (FeO), an oxidizable substance, is ever ready to take up oxygen, while potassium permanganate and potassium dichromate are always ready to give up some of their oxygen. When potassium permanganate gives up its oxygen in this way it is reduced and decolorized, while the ferrous oxid in taking up oxygen is oxidized to ferric oxid (Fe2O3). The decolorization of the permanganate here spoken of is taken advantage of in volumetric analysis for the deter- mination of the completion of the oxidation. The perman- ganate in the form of a standard solution being slowly delivered

139

140 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

from a burette, until it is no longer decolorized, the iron salt is known to be completely oxidized, when the permanganate is no longer reduced. The reaction is as follows:

Ferrous oxid Ferric oxid

5

Fe

f K2O.

The oxidation of ferrous oxid by potassium dichromate is shown by the following equation:

As before stated, an oxidation is always accompanied by a reduction, the oxidizing agent being itself reduced in the operation. As shown in the above equations, the manganic compound is reduced to a manganous, and the chromic to a chromous, while the ferrous salt is oxidized to a ferric condition.

In the same way any substance which readily yields oxygen in definite quantity or is susceptible of an equivalent action which involves its reduction to a lower quanti valence, may be estimated by ascertaining how much of a reducing agent of known power is required for its complete reduction.

Example. The available chlorin in bleaching powder may be accurately ascertained by treating it with a standard solu- tion of arsenous oxid, and from the volume of the solution required for the complete reduction of the chlorin, the quantity of the latter present is found, or in other .words, from the quantity of arsenous oxid (As2Os), oxidized to arsenic oxid (As2O5) the weight of the chlorin present is ascertained.

The principle substances which are used as oxidizing agents in volumetric analysis, are potassium permanganate, potassium dichromate and iodin. The latter contains no oxygen, but it abstracts hydrogen from accompanying water and liberates the oxygen which does the oxidizing, hence iodin Is known as an indirect oxidizing agent. The other two

ANALYSIS BY OXIDATION AND REDUCTION 141

contain available oxygen which they readily give up when brought in contact with an oxidizable substance.

The principal reducing agents or deoxidizers which are used in volumetric analysis are, sodium thiosulphate, sul- phurous acid, oxalic acid, arsenous oxid, stannous chlorid, ferrous oxid, hydriodic acid, hydrosulphuric acid, metallic zinc, and magnesium.

Preparation of Decinormal f — j Potassium Permanganate

N

(2KMnC>4 = 316.06; — V.S. = 3.1606 gms. in i liter). Abso- lutely pure potassium permanganate cannot be obtained, therefore the preparation of a decinormal solution of this salt cannot be effected by simply dissolving the requisite proportion of the molecular weight in the water. The presence of oxidizable matter in the water used, the contact of dust and exposure to light, have a tendency to decompose the salt and hence weaken the standard solution. It is therefore advisable to use boiling distilled water, and to preserve the solu- tion in amber glass bottles, provided with ground-glass stoppers. It will then retain its. strength for several weeks, but should nevertheless be checked by titration immediately before using. It is not necessary, and it is usually undesirable, to make the solution an exact decinormal one. It is preferable to fix the titer of the solution and employ it as it is.

Place 3.5 gms. of pure crystallized potassium perman- ganate in a flask, add 1000 mils of distilled water, and boil until the crystals are dissolved; put a plug of absorbent cotton in the mouth of the flask and set it aside for two days so that any suspended matter may deposit. After the lapse of this time pour off the clear solution into a glass-stoppered bottle, and when wanted for use standardize by either of the following methods:

142 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

Standardization by Means of Iron. Thin annealed binding- wire, free from rust, is one of the purest forms of iron.* o.i gm. of such iron is placed in a flask which is provided with a cork through which a piece of glass tubing passes, to the top of which a piece of rubber tubing is attached, which has a vertical slit about one inch long in its side, and which is closed at its upper end by a piece of glass rod (this arrangement is known as the " Bunsen Valve "). (See Fig. 41.) Diluted sulphuric acid is added and gentle heat applied. The iron dissolves and the steam and liberated hydrogen escape through the slit under slight pressure. The air is thus pre- vented from entering and the ferrous solution protected from oxidation.

A better form of apparatus in which to dissolve the iron and avoid oxidation through admission of air is shown in Fig. 42. A loo-mil flask is fitted with a rubber stopper and

a | 1 shaped glass tube; into this

flask is placed 20 mils of diluted sulphuric acid (1:5) and then 2 or 3 crystals of

FIG. 41.

FIG. 42.

pure sodium carbonate; this causes an evolution of carbon dioxid which expels the air from flask. The o.i gm. of iron

* This contains 99.6 per cent of iron.

ANALYSIS BY OXIDATION AND REDUCTION 143

wire above described is now introduced, the stopper inserted, and a beaker containing a solution of pure sodium carbonate placed in position so that the tube will dip into the solution. Gentle heat is applied until the iron is wholly dissolved, and only a few minute particles of carbon remain (which must not be mistaken for iron). When the flame is withdrawn the cooling of the flask and contents causes a drawing up of the sodium carbonate solution, but the first drops that enter the flask cause an effervescence with evolution of carbon dioxid, which drives the liquid back and at the same time fills the flask with the gas; this is repeated until the flask and contents are cold. Another useful form of apparatus for this purpose is shown in

Fig. 43-

When the iron is completely dissolved a small quantity of cold, recently boiled, distilled water should be used to rinse the lower end of the stopper and the neck of the flask, and the titration with potassium per- manganate at once begun and continued until a faint permanent pink color is produced. If the solution is decinormal, exactly 17.84 mils will be required to pro- duce this result.

The iron is converted by the sulphuric acid into ferrous sulphate, Fe2 + 2H2SO4 = 2FeSO4 + 2H2. This ferrous sulphate is easily oxidized by the air, and therefore it is directed that access of air should be prevented, and the distilled water with

FIG. 43-

144 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

which the solution is diluted previously boiled in order to drive off any dissolved free oxygen.

ioFeSO4 + 2KMnO4 + 8H2SO4 100)558.2 100)316.06 N

5.582 gms. 3.1606 gms. or 1000 mils — V.S.

N This equation, etc., shows that each mil of — perman-

ganate represents 0.005582 gm. of metallic iron.

Standardization by Means of Owalic Acid. 0.12605 gm. of the pure crystallized acid is weighed (or 20 mils of deci- normal oxalic acid carefully measured) and placed in a flask with 3 mils of sulphuric acid C. P. and distilled water to make 100 mils. The solution is warmed to 85° C. and the perman- ganate solution delivered in from a burette.

The action is in this case less decisive and rapid than in the titration with iron, and more care should be used. The color disappears slowly at first, but afterwards more rapidly.

Note the number of mils of the permanganate solution used, and then dilute the remainder so that equal volumes of it and of decinormal oxalic acid solution will exactly corre- spond.

Example. Assuming that 18.5 mils of the permanganate solution are required to produce a permanent pink tint in the above test, then the permanganate solution must be diluted with distilled water in the proportion of 18.5 mils of the permanganate solution and 1.5 mils of water, or 1850 mils to 150 mils.

After dilution a new trial should be made, in which 50 mils of the diluted permanganate solution should require

N exactly 50 mils of — oxalic acid V.S.

ANALYSIS BY OXIDATION AND REDUCTION 145

The reaction between potassium permanganate and oxalic acid is illustrated by the following equation:

2KMnO4 + 5(H2C2O4.2H2O) +3H2SO4

= K2SO4 + 2MnSO4 4- ioCO2 + i8H2O.

Standardization by the lodometric Method. This method, which was proposed by Volhard, is the most accurate and rapid for the standardization permanganate. It is based upon the fact that potassium permanganate reacts with potas- sium iodid in solutions acidulated with either hydrochloric or sulphuric acid, and liberates an equivalent quantity of iodin, which may be estimated by standard solution of sodium thiosulphate. The reactions are illustrated by the equations

(a) 2KMnO4 +8H2SO4 + loKI 316.06

1269.2 \b) I2-\-2 (Na2S2O3 . 5H2O) = 2NaI + Na2S4O6 + ioH2O.

253.84 496.44

Thus it is seen that 2KMn.O4 (316.06 gms.) containing five atoms of available oxygen, has the power of liberating its equivalent of iodin, i.e., 10 atoms or 1269.2 gms. (see equation a) and that 496.44 gms. of sodium thiosulphate will reduce 253.84 gms. of iodin (see equation b). Hence

N 1000 mils of — sodium thiosulphate (containing 25.384 gms.)

will reduce, and therefore be equivalent to 12.692 gms. of iodin, which in turn represents 3.1606 gms. of potassium

N permanganate. Therefore i mil of the — thiosulphate repre-

sents 0.012692 gm. of iodin and 0.0031606 gm. of potassium permanganate, which latter is the quantity of potassium per-

N manganate present in i mil of its — V.S.

146 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

The process is conducted as follows: Into a 2oo-mil flask place about 0.5 gm. of potassium iodid and 10 mils of diluted sulphuric acid, add to this (slowly from a burette) exactly 10 mils of the permanganate solution to be standardized and dilute the mixture (which is brown in color, because of the liberated iodin) with distilled water to about 150 mils. Then slowly titrate (with constant stirring) with an accurately stand -

N ardized — sodium thiosulphate until the color of the solution

is a faint yellow, then add a few drops of starch solution and continue the titration until the color is discharged. Note the number of mils consumed and dilute the permanganate with distilled water so that equal volumes of the two solutions correspond to each other.

Example. If 13 mils of the thiosulphate solution were required, then each 10 mils of the permanganate solution must be diluted to 13 mils.

Standardization with, Ferrous Ammonium Sulphate (Mohr's salt) *(FeSO4- (NH4)2SO4-6H2O). 392.14 gms. of this salt contain 55.82 gms. of iron (3.512 gms. contain 0.5 gm. of iron). 3.512 gms. of the salt are accurately weighed out and dissolved in sufficient recently boiled distilled water to make 250 mils. Fifty mils of this solution containing o.i gm. of iron are transferred to a small flask, 10 mils of diluted sul- phuric acid added, and then the permanganate solution to be standardized is run in slowly until a faint pinkish tint appears. Whatever number of mils is consumed th t number represents o.i gm. of iron, and must be diluted to 17.91 mils to make the solution exactly decinormal.

Volumetric Analyses by Means of Potassium Permanganate When potassium permanganate solution is added to a solu- tion of any readily oxidizable substance strongly acidulated with sulphuric acid, it undergoes reduction, as shown in the

ANALYSIS BY OXIDATION AND REDUCTION 147

equation below. The molecule (2KMnO4) has eight atoms of oxygen which it gives up in the process of oxidation. These eight atoms of oxygen unite with the replaceable hydrogen of an accompanying acid, liberating an equivalent amount of acidulous radical. Three of these atoms of oxygen liberate sufficient acidulous residual to combine with the potassium and manganese of the permanganate, while the other five atoms are available for direct oxidation.

or, for combination with the hydrogen of more acid, more acidulous residual being set free, to combine with the salt acted upon.

5(804) when combined with ioFeSC>4 forms Feio(SO4)i5 or 5Fe2(SC>4)3, ferric sulphate. Thus it is seen that one molecule of potassium permanganate (2KMnO4) has the power of converting 10 molecules of a ferrous salt to the ferric state.

The equation in full is

ioFeSO4 + 2KMnO4 + 8H2SO4

We have seen that 2KMnO4 has 5 atoms of oxygen available for oxidizing purposes, and that each of these will combine with 2 atoms of hydrogen. 2KMnO4 is consequently chemically equivalent to 10 atoms of hydrogen, and a normal solution of this salt when used as an oxidizing agent is one that contains in one liter one-tenth of the weight of 2KMnO4 expressed in grams, and a decinormal solution, one which contains one-hundredth of this weight.

As before stated, when potassium permanganate is brought

148 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

in contact with a ferrous salt or other oxidizable substance, it is decomposed and decolorized. Hence when titrating with a standard solution of this salt it is decolorized so long as an oxidizable substance is present. As soon, however, as the oxidation is completed the standard solution retains its color when added to the substance, and the first appearance of a faint red color is the end-reaction, and the oxidation is known to be completed.

In titrating with potassium permanganate it must be remembered that excess of free acid (preferably sulphuric) should always be present in the solution titrated, in order to keep the resulting manganous and manganic oxids in solu- tion; these, forming a dense brown precipitate, would make it difficult if not quite impossible to recognize the pinkish color of the end-reaction. Sulphuric acid alone, if in large excess, has a reducing effect upon potassium permanganate.

Nitric and hydrochloric acids are prejudicial and should be avoided; they are, however, frequently present in salts which are to be analyzed, and in such event should be removed by converting them into sulphate. By adding a small excess of sulphuric acid and applying heat, until hydro- chloric acid or nitrous vapors are no longer evolved, the chlorid or the nitrate is converted into sulphate, and the deleterious effect of their presence overcome. Hydrochloric acid, unless present in very small quantities, and the titra- tion conducted at a low temperature, will vitiate the analysis through its action upon the permanganate whereby chlorin is liberated,* thus

KMn04 + 8HC1 - KC1 4- MnCl2 + 4H2O + 5C1.

* This decomposition of the permanganate by hydrochloric acid is due to the presence of ferric salt, which latter seems to act catalytically, for oxalic acid may be accurately titrated with permanganate even in the presence of hydrochloric acid, no chlorin being given off. Thus the decomposition of the permanganate is not due to the hydrochloric acid alone.

ANALYSIS BY OXIDATION AND REDUCTION 149

A very convenient way of obviating the irregularities due to the presence of hydrochloric acid is to add a few grams of manganous sulphate * to the solution before titrating it.

Mercuric sulphate { and magnesium sulphate may also be used with satisfactory results.

Potassium permanganate, being so readily decomposed by contact with organic matter, should be protected from such contact. It should never be filtered through paper (glass- wool or guncotton may be used), nor should it be used in a Mohr's burette or in any other apparatus in which it is in contact with rubber or cork. Furthermore, all substances of an oxidizing or reducing nature, aside from that being analyzed, must be excluded from the solution. Among such substances may be mentioned hydriodic acid, sulphureted hydrogen, nitrous acid and the lower oxids of nitrogen, phos- phorous and hypophosphorous acids> thiosulphuric, sulphurous, and all the other acids of sulphur except sulphuric, also ous salts and the metallic suboxids and peroxids.

Burettes and other apparatus which have been used for

* Kessler and Zimmermann suggest using 20 mils of a solution of man- ganous sulphate (200 gms. per liter).

% Cady and Ruediger (J. A. C. S., XIX, 575) concluded from the following general principles that it is possible to titrate iron with permanganate in the presence of hydrochloric acid if an excess of mercuric sulphate be added to the solution. Mercuric halids in solution ionize to an extremely slight extent, while the mercuric salts of oxyacids are readily ionized, since compounds of slight icnization always result when their constituent ions meet; mercuric halids are always produced when a mercuric salt of an oxyacid is added to a solution containing halogen ions. Therefore when mercuric sulphate solution and hydrochloric acid are mixed, ionization of both occurs, and the mercuric ions unite with the chlorin ions and produce mercuric chlorid which is only 'fery slightly ionized. In the presence of a large excess of mercuric sulphate, the mercuric ions resulting from its dissociation diminish the ionization of the mercuric chlorid until it is practically nil. Thus no chlorin ions will be present in the solution to induce decomposition of the permanganate.

150 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

permanganate, should be emptied and rinsed immediately after use, and any manganic oxid which may be adhering to the glass should be removed by means of hydrochloric acid and boiled water.

Not only oxidizable substances but reducible substances may be estimated by means of potassium permanganate.

In the estimation of oxidizable substances the standard potassium permanganate is added directly to the acidulated solution of the substance being analyzed. The completion of the oxidation being then known by the appearance of a faint pinkish tint. This is the direct method.

In the estimation of reducible substances (i.e., oxidizing substances) the indirect or the residual method is employed.

In this an accurately weighed or measured quantity of the substance is brought together with an excess of a third substance having reducing power, and which is similarly effected by the permanganate and by the substance analyzed. After completion of the reaction the excess of the reducing substance is found by titration with standard permanganate. The difference between the quantity so found and that originally added gives the quantity which reacted with the salt under analysis, and from this the calculation is made.

On the Use of Empirical Permanganate Solutions.

A. If the standardization of the solution is done by means of iron, as described on page 142, o.i gm. of iron wire (repre- senting 0.0996 gm. of pure iron) will require 17.84 mils of the permanganate solution if the latter is exactly decinormal. If less than this quantity of solution is used (say 17.5 mils) it indicates that the solution is stronger than decinormal, and may either be diluted so that each 17.5 mils will measure 17.84, or it may be used as it is. This latter is in most cases

ANALYSIS BY OXIDATION AND REDUCTION 151

preferable. The value of i mil of the solution in iron is calculated thus:

17.5 mils. : i mil 1:0.0996 gm. : x. #=0.00569 gin.

If a solution of this strength is to be used for the estima- tion of iron, simple multiplication of the number of mils used by 0.00569 gm. gives the weight of Fe present. If, however, this solution is employed for the titration of other oxidizable substances, the number of mils consumed is multiplied by 0.00569 gm. and then by a fraction in which the numerator represents a quantity of the substance examined, equivalent in grams to an atom of iron in its reaction with permanganate, and the denominator is the atomic weight of iron.

Example. If in a titration we use 40 mils of a perman- ganate solution, the titer of which has been found to be i mil = 0.00569 gm., the calculation would be:

in the case of ferrous sulphate (FeSO.*, 151.89),

151.89 40X0.00569 gm.X — ^—=0.6192 gm.;

00*

in the case of oxalic acid (Rz&zQ*- 21120 = 126.05),

63 40X0.00569 gm.X — ^-=0.256 gm.;

in the case of hydrogen dioxid (H2O2 = 34),

40X0.00569 gm.X -^—=0.0693 gm.

B. Another way. The solution just mentioned, of which 17.50 mils are consumed in titrating o.i gm. of iron wire, is compared with a true decinormal permanganate solution, of which 17.84 mils are consumed in the same reaction. The

152 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

T *7 ft A

strength of the former solution is therefore — — as compared

/ 0

with a decinormal solution.

In titrating with this solution the number of mils consumed

are to be multiplied by - and then by the true decinormal

factor for the substance being analyzed.

Example. Forty mils of the solution are consumed.

In the case of ferrous sulphate (FeSO4= 151.89), the deci- normal factor (i.e., the weight of ferrous sulphate represented by i mil of a true decinormal solution) is 0.015189 gm.

1084

4oX— — Xo.oi5i89 gm. =0.6192 gm. I7S°

In the case of oxalic acid (H2C2O4, 2H2O = 126.05), the deci- normal factor is 0.0063 8m-

1784 40 X-1 — X 0.0063 gm.=o.256 gm.

In the case of hydrogen dioocid (H2 02 = 34), the decinormal factor is 0.001688 gm.

1784 40 X— — Xo.0017 gm. =0.0693 §m-

C. If the standardization is done by means of oxalic acid, as described on page 144, in which 10 mils of a strictly deci- normal oxalic acid solution are titrated with the permanganate solution which is being standardized, exactly 10 mils of • the latter will be consumed if it is of decinormal strength. If in the trial, however, it is found that only 9.6 mils are con- sumed it indicates that the solution is stronger than decinormal;

100

its strength being expressed by —T. If, on the other hand,

90

ANALYSIS BY OXIDATION AND REDUCTION 153

more than 10 mils of the solution are consumed (say 10.4 mils)

the solution is below decinormal strength, namely, .

In using a solution of the first strength the number of mils of it consumed in any titration is to be multiplied by — — and

then by the decinormal factor for the substance examined. In the case of the weaker solution the number of mils con-

100 sumed is multiplied by — and then by the decinormal factor

for the substance being analyzed.

Examples. Ferrous sulphate (FeSO4=i5i.89) is titrated with the stronger solution, 40 mils of the latter being consumed.

TOO

Then 40 X— -7 Xo.oisiSg gm.= 0.628 gm.

90

Oxalic add (H2C2O4 • 2H2O = 126.05), 4° m^s are consumed.

100 Then 4oX-r-Xo.oo63 gm. = o.26o gm.

Hydrogen dioxid (H2O2 = 34), 40 mils are consumed.

100

Then 40 X— 7- X 0.0017 gm. = 0.0703 gm.

90

D. If the checking of the permanganate solution is done by the iodometric method (page 145) and it is found that 10 mils of the permanganate requires the use of 13 mils of decinormal thiosulphate solution, the titer of the solution

is ' expressed with reference to decinormal as — . In using a solution of this strength, the number of mils of it consumed

in an analysis is multiplied by — and then by the decinormal factor for the substance analyzed.

154 THE ESSENTIALS OF VOLUMETRIC ANALYSIS TYPICAL ANALYSES WITH PERMANGANATE

A. Direct Titrations

a. Estimation of Ferrous Sulphate (FeSO4 + 7^0 = 278). One gm. of ferrous sulphate is dissolved in 25 mils of water and the solution strongly acidulated with sulphuric acid. Deci- normal potassium permanganate is then delivered from a burette until a permanent pink tint is obtained, indicating the complete oxidation of the ierrous salt.

The reaction is as follows:

(ioFeSO4 + 7H20) +. 2KMnO4 + 8H2SO4 ' , * * 100)2780 100)316.06 l^

27.80 gms. = 3.1606 gm. = 1000 mils — V.S.

N 2.780 gms. = 100 mils — • V.S.

N

0.0278 gm. = ' i mil — V.S.

10

-= 5Fe2(SO4)3 + K2SO4 + 2MNSO4 + 78H2O.

Thus 316.06 gms. of permanganate = 2 780 gms. of crys- tallized ferrous sulphate, which equals 55.82 gms. of metallic

N

iron. One mil of — permanganate solution therefore repre- sents 0.0278 gm. of FeSO4 + 7H2O or 0.005582 gm. of Fe.

N

In the analysis 35 mils of the — permanganate were con- sumed. The i gm. taken then contains 35X0.0278=0.973 gm. or 97.3 per cent.

If it is desired that each cc. of the permanganate solution should represent a certain percentage of pure salt, a molecular quantity of the salt should be taken for analysis instead of i gm. For example, if 2.78 gms. be taken, each mil of the decinormal solution consumed will correspond to i per cent, because 2.78 gms. is the weight of crystallized ferrous sul-

ANALYSIS BY OXIDATION AND REDUCTION 155

phate which can be oxidized by 100 mils of the decinormal solution. If half of this weight be taken, i.e., 1.39 gms., each mil of the permanganate solution compound will represent 2 per cent of pure salt.

Granulated Ferrous Sulphate (FeSO4 + 7H2O) is estimated in the same way as the foregoing, and should correspond with it in strength.

Exsiccated (Dried) Ferrous Sulphate. This salt is tested in the same manner as the other two sulphates. It contains a larger percentage of ferrous sulphate than the other two, having less water of crystallization. Its composition is approx- imately FeSO4+3H2O.

ioFeSO4 + 2KMnO4 + 8H2SO4

100)1518.9 100)316.06

15.189 gms. 3.1606 gms. or 1000 mils — standard solution.

= sFe2(SO4)3 + K2S04 + 2MnSO4 + 8H2O.

Each cc. of the standard solution represents 0.015189 gm. of anhydrous (real) ferrous sulphate.

b. Estimation of Metallic Iron in Ferrum Reductum. Fer rum reductum (reduced iron) always contains besides metallic iron a varying quantity of oxid. Therefore, in assaying this preparation a method must be employed which will estimate the iron only, which is present as metallic iron. This may be done by means of a solution of mercuric chlorid which reacts with metallic iron only and not with the oxid.

The method is as follows:

Introduce into a loo-mil flask i gm. of reduced iron, pre- viously well triturated. Add 10 gms. of powdered mercuric chlorid and 50 mils of boiling distilled water. Boil for five minutes, shaking it frequently, then fill the flask to the loo-mil mark with recently boiled and cooled distilled water. 2HgCl2 +Fe2 = 2FeCl2 + 2Hg.

156 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

Let it cool to room temperature and again fill to the mark. Shake well, and after a few minutes filter and collect 20 mils of filtrate. Add to this 20 mils of dilute sulphuric acid, and immediately titrate with tenth-normal potassium permanganate V.S. until a permanent pink color is produced.

Each mil of the permanganate solution represents 0.005582 gm. of metallic iron.

loFe SO4 + 2KMnO4 + 8H2SO4

100)558.2 100)316.06 N

5.582 gms. = 3.1606 gms. = 1000 mils — V.S. 0.005582 gm. = 0.0031606 gm. = i mil rr V.S.

= 5Fe2(SO4)3 + K2SO4 + 2MnSO4 + 8H2O.

Titration with .an Empirical Permanganate Solution. A solution of permanganate, which is found upon standardiza- tion to be of a strength in which i mil is equivalent to 0.00512 gm. of Fe, is to be used.

Each mil of this solution is equivalent to the following quantities:

FeSO4 0.01393 gm.

FeSO4 + 7H2O 0.02549 "

FeCO3 0.01062 "

FeCl2 0.01162 "

Fe 0.00512 "

c. Estimation of Oxalic Acid and Oxalates with Potassium Permanganate Solution (H2C2O4 + 2H2O = 126.05; H2C2O4 = 90). The estimation of oxalic acid may be accurately made either by neutralization with a standard alkali or by oxidation with standard permanganate. The latter method is, however, the one to be employed in the case of oxalates.

The oxidimetric estimation of oxalic acid is carried out as follows:

ANALYSIS BY OXIDATION AND REDUCTION 157

One gm. of the acid (accurately weighed) is dissolved in sufficient water to make 100 mils. Of this solution, 10 mils (representing o.i gm. of the acid) is taken for analysis. Two mils of diluted sulphuric acid are added, the solution is heated to between 40° C. and 60° C., and keeping it at about this temperature, is titrated with decinormal potassium perman- ganate, agitating constantly, until a faint rose tint marks the completion of the reaction.

Each mil of the permanganate solution consumed repre- sents 0.0063 gm. of crystallized oxalic acid.

The reaction is as follows: 5 (H2C2O4 + 2H2O) + 3H2SO4 + 2KMnO4

6.30 gms. = 3.1606 gms. = 1000 mils — V.S.

= K2SO4 + 2MnSO4+ i8H2O + ioCO2.

Direct Percentage Titration. 0.63 gm. of crystallized oxalic

N acid is oxidized by 100 mils of — permanganate. Therefore

N if 0.63 gm. of the acid is taken for analysis, each mil of —

permanganate will represent i per cent.

Titrating with an Empirical Solution. If the permanganate is checked with iron, we take into consideration that 2KMnO4 will oxidize 10 atoms of iron (558.2 parts), and on the other hand 5 molecules of oxalic acid (630.25 parts). If the titer of the permanganate be found on experiment to be i mil =0.00569, whatever number of mils of this solution is consumed is to

be multiplied by 0.00569 and then by — —-.

Example. 0.3 gm. of oxalic acid requires for oxidation 40 mils of a permanganate solution whose titer is i mil =0.00569 gm. Fe, the calculation is made as follows:

63

40X0.00569 gm. X — -T- =0.256 gm. *

158 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

0.256 gm. is the quantity of pure crystallized oxalic acid present in the 0.3 gm. taken for analysis. This is 85.3 per cent.

0.256X100

-=85.3. o-3

If the standardization of the permanganate is done by means of a decinormal oxalic acid, or by the iodometric method, the calculation is as described on pages 152-153.

Oxalates are estimated in the same manner; a much larger quantity of sulphuric acid is, however, required. This serves to liberate the oxalic acid from its combination.

The presence of precipitates of sulphates of calcium, barium, or lead does not interfere with the recognition of the end- point.

N Each mil of — potassium permanganate represents

Oxalic acid anhydrous (H2C2O4) 0.0045 &m-

Oxalic acid crystallized (H2C2O4 + 2H2O) . . .0.0063 "

d. Estimation of Hydrogen Dioxid and Barium Dioxid with Standard Potassium Permanganate. Hydrogen dioxid (Hydro- gen peroxid) (H2O2 = 34). Hydrogen dioxid and potassium permanganate, though both oxidizing agents, will, when mixed in an acid solution, reduce each other. The reaction which occurs is probably primarily an oxidation of the H2O2 to a higher oxid (H2O4 (?)) which, however, immediately breaks up with the liberation of oxygen. The method of assaying hydrogen dioxid by means of permanganate is applicable not only to this substance but also to the estimation of barium dioxid and the soluble alkali peroxids. The method is usually carried out by adding the permanganate solution to the dioxid in a solution acidulated with sulphuric acid. Immediate decolorization of the permanganate occurs, as long as any

ANALYSIS BY OXIDATION AND REDUCTION 159

hydrogen dioxid is present. When the latter has been entirely taken up the permanganate is no longer decolorized and a faint pink tint marks the end-point. In the estimation of the pharmacopoeial or commercial dioxid solutions, containing 2 or 3 per cent of H2O2, a measured quantity is taken for analysis. The specific gravity of the solution, being nearly that of water, i cc. is taken to represent i gm. In the case of solutions of hydrogen dioxid of high percentage strength, it is advisable to take a weighed quantity for analysis. If hydrochloric acid is present a small quantity of manganese sulphate should be introduced before titrating.

The assay is conducted as follows :

An accurately weighed quantity (about 2 gms.) of the hydro- gen dioxid solution is diluted with 20 mils of distilled water, and acidulated with 20 mils of diluted sulphuric acid and then

N the — permanganate solution run in from a burette, stirring

after each addition until a permanent faint pink tint appears. The reaction is as follows :

5H2O2 + 2KMnO4 +

i ex?) 1 70 100)316.06 j^-

1.7 gms. 3.1606 gms. = 1000 mils — permanganate V.S.

N 0.0017 gm. = i mil — permanganate V.S.

= K2SO4 + 2MnSO4 + 8H2O + 5O2.

N r~i

Thus each mil of — permanganate represents 0.0017 gm-

of absolute hydrogen dioxid. Assuming that in the above estimation 35.5 mils of the permanganate solution were required, then the 2 gms. taken for analysis contained 0.0017 gm.X35.5, which is 0.06035 §m- °f absolute H2Os. This corresponds to 3.017 per cent.

160 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

The Direct Percentage Method. Ten gms. of the solution is diluted with water to measure 100 mils. Seventeen mils of this diluted solution (containing 1.7 gms. of hydrogen dioxid) is acid- ulated with sulphuric acid and titrated with decinormal perman- ganate, as above described. Each mil of the permanganate solu- tion consumed will represent o.i per cent by weight of H2O2.

Titralion with an Empirical Solution. A permanganate solution is on hand which is found upon standardization with iron to be i mil = 0.00569 gm. Fe. To use this solution as it is, we take into consideration that 2KMnO4= (316.06) = 10 atoms of Fe (558.2) and also 5' molecules of H2C>2 (170). 31.606 gms. KMnO4, =55-82 gms. Fe, =17 gms. H2O2. Whatever number of mils of this permanganate solution is used,

multiplied by 0.00569 gm. and then by — — , will give the

weight of H2O2 present in the sample analyzed.

Estimation of Volume Strength. Let us look at the above equation in a different light.

We see that when potassium permanganate and hydro- gen dioxid react, 10 atoms of oxygen are liberated.

The permanganate itself when decomposed liberates five atoms of oxygen. Therefore of the above ten atoms only five come from the hydrogen dioxid.

2KMnO4 + 3H2SO4 = K2SO4 + 2MnSO4 + 3H2O + 50.

In order to find the factor for volume of available oxygen, see the following equation, etc. :

100)316.06 -^ 100)80

3.1606 gms. or 1000 mils — V.S. = 0.80 gm.

10

i mil — V.S.= 0.0008 gm.

ANALYSIS BY OXIDATION AND REDUCTION 161

N Thus it is seen that each mil of — potassium permanganate

V.S. represents 0.0008 gm. of oxygen. But we require to find the volume of oxygen, not the weight represented by

N i mil of — permanganate.

1000 mils of oxygen at o° C. and 760 mm. pressure, weigh 1.43 gms. Therefore, if 1.43 gms. measure 1000 mils, 0.0008 gm. will measure 0.57 mil.

The factor, then, for volume of oxygen liberated when

N hydrogen dioxid is titrated with — potassium permanganate,

N

is 0.57, and the number of mils of the — potassium perman- ganate consumed in the titration gives the volume of oxygen liberated by the quantity of hydrogen dioxid taken.

N Thus if to mils of the — V.S. were required,

0.57 X 19= 10.83 m^s °f oxygen.

It is convenient to operate upon i mil of hydrogen dioxid solution. Then each mil of potassium permanganate V.S. used will represent 0.57 mil of available oxygen and is neces- sary only to multiply the number of mils by this factor to find the volume of available oxygen.

If any other quantity than i mil of dioxid be taken for analysis, it will be necessary after multiplying by 0.57 to divide the result by the quantity of dioxid solution taken, in order to find volume strength.

Hydrogen dioxid solution may also be volumetrically assayed by Kingzett's method, which is described under lodometry.

The gasometric estimation is also described further on.

162 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

Barium Dioxid (Barium Peroxid) (1^02=169.37). This substance is assayed by treating it with an acid, and then estimating the liberated hydrogen dioxid, as follows:

Weigh off 2 gms. of the coarse powder, put it in a porcelain capsule, add about 10 mils of ice-cold water, then 7.5 mils of phosphoric acid (85 per cent), and sufficient ice-cold water to make 25 mils. Stir and break up the particles with the end of the stirrer until a clear or nearly clear solution is obtained and all that is soluble is dissolved.

Five mils of this solution (which corresponds to 0.4 gm. of barium dioxid) is measured off for assay.

Drop into this from a burette, with constant stirring, deci- normal potassium permanganate until a final drop gives the solution a permanent pink tint.

About 40 mils of the decinormal permanganate should be required to produce this result.

In this process, the first step is the formation of hydrogen dioxid by treating the barium dioxid with phosphoric acid, as illustrated by the following equation:

Ba02 + H3P04 = BaHP04 + H2O2.

169.37 34

The hydrogen dioxid is then estimated with decinormal permanganate, as described above.

5(BaO2) = 5H2O2 + 2KMnO4 + 100)845.85 100)170 100)316.06 _-.

8.4585 gms. = 1.7 gms. 3.1606=1000 cc. — permanganate.

= K2SO4 + 2MnSO4 + 8H2O + 5O2.

Sodium Perborate (NaBO3+4H2O = 154.06). Dissolve an accurately weighed quantity (about 0.25 gm.) of the salt in a mixture of 50 mils of distilled water and 10 mils of diluted sulphuric acid, and titrate the solution with tenth-normal

ANALYSIS BY OXIDATION AND REDUCTION 163

potassium permanganate V.S. until a final drop gives the solution a permanent pink color.

Each mil of the standard permanganate used corresponds to 0.0008 gm. of available oxygen, which should be present in not less quantity than 9 per cent. The reaction involved is represented by the equation:

154.06 Percarbonates may be assayed in the same manner.

B. Residual Titrations a. METHODS IN WHICH AN EXCESS or STANDARD PERMANGANATE is ADDED

AND THE EXCESS DETERMINED BY RESIDUAL TlTRATION WITH STAND-

ARD OXALIC ACID.

Estimation of Nitrous Acid and Nitrites. Nitrous acid, when brought in contact with a potassium permanganate solu- tion acidulated with sulphuric acid, is oxidized to nitric acid. Two molecules of KMnO4 reacting with 5 molecules of HNO2, as the equation shows,

5HNO2 + 2KMn04+3H2SO4

= 5HNO3 + K2SO4 + 2MnSO4 + 3H2O.

In the case of nitrites, as for example sodium nitrite, the oxidation takes place in the same manner, and the process may be applied with equally good results to the salts, as well as to free HNO2. At ordinary temperatures the oxidation proceeds very slowly, but at a temperature of 40° C. (104° F.) rapid reaction occurs. But because of the volatility of nitrous acid in acidulated solutions of its salts it is impossible to accurately estimate them by direct titration with permanganate at a raised temperature.

164 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

It is customary to add the nitrite solution to a measured volume of warm acidulated standard permanganate solution and then to titrate the excess of permanganate with standard oxalic acid solution.

N The assay is conducted as follows: To 50 mils — potassium

permanganate V.S. add 100 mils of distilled water and 5 mils of sulphuric acid. To this solution add by means of a pipette 10 mils of a solution of sodium nitrite (i gm. in 100 mils). In adding the sodium nitrite solution immerse the tip of the pipette beneath the surface of the permanganate mixture. Warm the liquid to 40° C., allow it to stand for five minutes and then

N titrate with — oxalic acid V.S. until the color of the perman-

ganate solution is just discharged. Subtract the number of mils of the oxalic acid solution used from the number of mils of permanganate solution taken, and multiply the remainder

by 0.0034505 gm.

5NaNO2 + 2KMnO4 5X69.01=345.05

= 5NaNO3+K2SO4+2MnSO4+3H2O.

N Each mil of — potassium permanganate V.S. corresponds

to 0.0034505 gm. of NaNO2.

The Assay of Hypo phosphites may be made by the same method.

An accurately weighed quantity of hypophosphorous acid or its salt is dissolved in water, the solution strongly acidulated

N with sulphuric acid, and then a measured excess of — potas-

sium permanganate solution added. The mixture is boiled for fifteen minutes to hasten and facilitate the oxidation and

ANALYSIS BY OXIDATION AND REDUCTION 165

N then the excess of permanganate solution titrated with —

oxalic acid V.S.

Example. Three gms. of the acid are diluted with water to make 60 mils. Of this solution 6 mils (representing 0.3 gm. of the acid) are carefully removed with a pipette and introduced into a flask. Three mils of sulphuric acid are added

N and then 50 mils of — potassium permanganate solution, and

the mixture boiled for fifteen minutes.

The potassium permanganate, in the presence of sulphuric acid, oxidizes the hypophosphorous acid to phosphoric, as the equation shows:

5HPH202 + 6H2S04 + 2(2KMn04)

2)330.2 2)632.12

I(X>)l65.I IOO)3I6.O6 -VT

1.651 gms. 3.1606 gms. or 1000 mils — V.S.

= 5H3PO4 + 6H2O + tK2SO4 + 4MnSO4.

Each mil of the decinormal V.S. represents 0.001651 gm. of absolute hypophosphorous acid. The quantity of per- manganate solution directed to be added is slightly in ex- cess. The excess is then ascertained by retitration with decinormal oxalic acid. Each mil of oxalic acid required cor- responds to i mil of decinormal permanganate which Eas been added in excess of the quantity actually required for the oxidation.

The excess of permanganate colors the solution red, and the oxalic acid V.S. is then added until the red color just disappears, which indicates that the excess of permanganate is decomposed.

If 4.7 mils of decinormal oxalic acid are required, it indi- cates that 50 mils -4.7 mils = 45. 3 mils of decinormal perman-

166 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

ganate were actually used up in oxidizing the hypopnosphorous acid; therefore

0.001651 gm.X45.3 =0.0747 + gm. of absolute hypophosphorous acid, HPH2C>2, or

0.0747 X 100

—=24.7 per cent.

In this process boiling facilitates the oxidation, but if the acid is boiled before sufficient permanganate has been added to completely oxidize it, decomposition will take place. Hence direct titration with permanganate is impossible and the residual method must be resorted to.

b. METHODS INVOLVING A PRECIPITATION BY OXALIC Aero AND THE TITRA- TION OF THE EXCESS OF THE LATTER WITH STANDARD PERMANGANATE.

Assay of Calcium Carbonate and Soluble Calcium Salts.

Dissolve an accurately weighed quantity of calcium carbonate (0.3 to 0.4 gm.) in 10 mils of distilled water and 10 mils of diluted hydrochloric acid, and boil the solution to expel all carbon dioxid. Transfer this solution to a 2oo-mil graduated

N

flask, add 100 mils of — oxalic acid V.S., render it alkaline with 10

ammonia water, shake well and allow it to stand for three hours at from 60° to 70° C. or overnight at room temperature. Cool the mixture, if necessary dilute with distilled water to the 2oo-mil mark, mix it well, filter through a dry filter into a dry flask, reject the first 20 mils of filtrate, and proceed as follows: To 100 mils of the filtrate (representing half of the calcium carbonate taken) add diluted sulphuric acid until of acid reaction, then add 25 mils more of the same acid, warm the

N solution to 60° or 70° C., and titrate with — potassium per-

ANALYSIS BY OXIDATION AND REDUCTION 167

manganate V.S. until a permanent pale pink color is ob- tained.

Deduct the number of mils of the permanganate solution used (say 10.8 mils) from half the volume of the oxalic acid solution taken (50 mils) and multiply the remainder (39.2 mils) by 0.005 gm. This gives the weight of CaCO3 in half the quantity of the salt taken for analysis.

50 -10.8 = 39.2 mils; 39.2X0.005 gm. =0.1960 gm.; 0.1960X100

.2

= 98 per cent.

Calcium Oxid (Calx), CaO = 56oy, is assayed in the same manner.

Estimation of Soluble Calcium Salts. To a weighed quan- tity of the calcium salt dissolved in water, a measured excess of normal oxalic acid is added. The mixture is then made alkaline with ammonia and boiled, to facilitate the separa- tion of the precipitate. It is then cooled and diluted with water to an accurately measured volume, and after nitra- tion an aliquot portion removed, acidulated with sulphuric

N acid, and carefully titrated with — potassium permanganate.

Example. 0.4 gm. of calcium chlorid is dissolved in water, 10 mils of normal oxalic acid added, the mixture made alkaline with ammonia water, and boiled for a few minutes. It is then filtered, the residue and filter washed with water, and after cooling the combined filtrate and washings are diluted to make 100 mils.

Of this solution 50 mils are taken for analysis (representing 0.2 gm. of the salt), acidulated with sulphuric acid, and then

N titrated with — potassium permanganate to a faint rose tint.

168 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

The 50 mils of solution represent 5 mils of normal oxalic acid, which is equivalent to 50 mils of decinormal oxalic acid, so that whatever number of mils of decinormal permanganate solution is required in the titration, that quantity is to be deducted from 50 mils and the difference multiplied by the

N

factor for calcium chlorid to find the quantity of pure

CaCl2 present in 0.2 gm.

N If 14 mils of -- permanganate are employed, then 14 from

50 mils leaves 36 mils, the quantity of decinormal oxalic acid solution which combined with the 0.2 gm. of calcium chlorid. Then

0.00555 gm.X36=o.i998 gm.,

the quantity of pure CaCl2 present in the 0.2 gm., or 99.9 per cent.

Calcium salts to be estimated by this method must be tolerably pure, and free at least from impurities which would react with oxalic acid or which would reduce permanganate.

Many of the less soluble calcium salts may be estimated by this method, but they must be subjected to longer treat- ment with the oxalic acid.

Gold and lead salts may also be estimated by the same method.

Estimation of Lead in the Acetate and Subacetate. Take for assay 0.2 gm. of the salt or 2 gms. of the solution in a beaker and add 20 mils of recently boiled distilled water. Pour this slowly and with constant shaking into a graduated

N cylinder containing 50 mils of — oxalic acid V.S. Wash the

beaker with small portions of distilled water and add the washings to the contents of the cylinder. Then dilute the mixture to 100 mils and set aside to allow the precipitate to

ANALYSIS BY OXIDATION AND REDUCTION 169

settle. Remove 20 mils of the clear liquid (representing 0.04 gm. of the salt or 0.4 gm. of the solution) for titration. Add 5 mils of (i : 10) sulphuric acid, warm to 80° C, and titrate

N with — permanganate until a final drop imparts a permanent

pale pink tint.

The reactions are represented by the following equations:

Pb(C2H3O2)2 + H2C2O4-2H2O

Lead acetate f • '

325.1 126.05

= PbC2O4 + 2HC2H3O2 '•+ 2H2O. Pb2O(C2H3O2)2 + 2H2C2O4.2H2O

Lead subacetate

548.2 2X126.05

= 2PbC2O4 + 2HC2H3O2 + 5H2O.

N Each mil of — oxalic acid represents 0.016257 gm. of

Pb(C2H3O2)2 or 0.013705 gm. of Pb2O(C2H3O2)2, or 0.010355 gm. of Pb.

Calculate as described under assay of calcium carbonate.

The U.S.P. IX Method is essentially the same as the fore- going. It is as follows:

Assay of Lead Acetate. Dissolve an accurately weighed quantity of lead acetate (say 5 gms.) in sufficient recently boiled distilled water to make IGO mils of solution. Mix 10 mils of this solution with 50 mils of tenth-normal oxalic acid V.S. in a 200 -mil measuring flask, agitate the mixture thoroughly for five minutes, then fill the flask to the 2co-mil mark with distilled water; filter, and titrate 100 mils of the filtrate (representing one-twentieth of the amount of lead acetate originally taken) with tenth-normal potassium per- manganate V.S., the filtrate being previously acidulated with 10 mils of sulphuric acid and warmed to 80° C.

170 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

Assay of Lead Oxid (PbO = 223.10). Dissolve about 0.4 gm. of lead oxid in 4 mils of acetic acid and 25 mils of recently boiled distilled water. Add 50 mils of tenth -normal oxalic acid V.S. and dilute with distilled water to make 200 mils of fluid. Mix well and filter. Collect 100 mils of filtrate. Acidu- late it with 20 mils of diluted sulphuric acid, warm the solution to 80° C. and titrate with tenth-normal potassium perman- ganate V.S. as in preceding assay.

Each mil of the oxalic acid V.S. corresponds to 0.011155 gm. of PbO.

c. METHODS INVOLVING A REDUCTION BY MEANS OF OXALIC ACID, AND RETITRATION OF THE EXCESS OF THE LATTER WITH POTASSIUM PER- MANGANATE.

' Estimation of Manganese Dioxid (MnO2). The estima- tion of manganese dioxid depends upon the fact that when it is boiled with oxalic acid in the presence of sulphuric acid definite reaction takes place, as the equation shows :

In the estimation a measured excess of oxalic acid solu- tion is added, together with some sulphuric acid, and the mixture heated until solution is complete.

The excess of oxalic acid is then found by retitration with standard permanganate solution. It is well to use a normal oxalic acid solution and a decinormal permanganate solution.

0.5 gm. of the dioxid is a convenient quantity to operate upon. Each mil of decinormal solution represents 0.004346 gm. of MnO2.

Example. 0.5 gm. of MnO2 is treated with 3 mils of sul- phuric acid and 10 mils of normal oxalic acid solution, which is equivalent to 100 mils of decinormal oxalic acid solution, the mixture is heated in a water- bath to 80° C. and then treated

ANALYSIS BY OXIDATION AND REDUCTION 171

as described above, and upon retitrating 25 mils of decinormal permanganate are required. Thus

100 mils -25 mils = 75 mils.

N

of — oxalic acid went into reaction with the MnC>2. Then 10

75X0.004346 = 0.3259 gm.

d. METHODS INVOLVING A REDUCTION BY MEANS OF A STANDARDIZED SOLUTION OF A FERROUS SALT, AND TITRATION OF THE REMAINING UNOXIDIZED FERROUS SALT, BY PERMANGANATE.

Estimation of Nitrates (Pelouze). This method consists in adding a weighed quantity of the nitrate to an acidulated solution of a ferrous salt of known strength, and, when reaction is complete, estimating the ferrous salt remaining, by titra- tion with permanganate, or in certain cases by means of dichromate V.S. The principle upon which the method is based is, that when nitric acid or a nitrate is brought in con- tact with a highly acidulated solution of a ferrous salt, the former gives off oxygen, which, passing over to the ferrous salt, oxidizes it to the ferric state, while at the same time NO is evolved. The reaction is

2HNO3 + 6HC1 + 6FeCl2 = 3F2C16 + 4H2O + 2NO.

Nitric acid Ferrous chlorid

126.02 760.44

Iron 334.92

Thus one molecular weight of nitric acid (63.01) will oxidize three molecular weights of ferrous salt, or three atoms of iron (167.46).

Either hydrochloric or sulphuric acid may be employed. The former is preferred by most operators, and it is generally

172 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

agreed that in order to attain results of sufficient precision the estimation should be done in the presence of hydrochloric acid only. In using hydrochloric acid, however, where the titrations are to be made with permanganate, certain precau- tions (previously mentioned) must be observed, because of the evolution of chlorin which will otherwise take place and spoil the analysis. This may be obviated by adding to the solution to be titrated an excess of manganese sulphate.

The NO which is produced during the reaction must be removed by boiling before titration with permanganate is begun. Air must be absolutely excluded during the entire process to prevent oxidation of ferrous salt by the atmospheric oxygen, as well as to prevent oxidation of NO to HNOs, which will oxidize more ferrous salt. The exclusion of air may be partially affected by the use on the flask of a Bunsen valve stopper (see Fig. 41), but the best method is to employ an apparatus so arranged that a constant stream of CO2 or H gas may be passed through it (see Fig. 44).

This method, although theoretically perfect, is in practice liable to great irregularities, and will give fairly good results only if the directions, especially those as to exclusion of air,

FIG. 44.

are faithfully carried out. The method of Kjeldahl is to be preferred.

To conduct the process, weigh accurately 1.5 gms. of flower

ANALYSIS BY OXIDATION AND REDUCTION 173

wire * free from rust (the iron content of which is known) , place it an Erlenmeyer flask which is provided with a double perforated stopper fitted with two glass tubes, one of which should reach just to the surface of the liquid in the flask when in place, and the other, which is the outlet tube, should reach no lower than the bottom of the stopper. The first of these tubes is connected with an apparatus generating carbon dioxid or hydrogen, while the outlet tube serves to convey the gas into the air or into another flask containing water or an alka- line solution. 30 to 40 mils of pure fuming hydrochloric acid are added to the iron wire in the flask, gentle heat is applied, and a stream of either CO2 or H passed through the flask and maintained throughout the entire process. When the iron is completely dissolved, the stopper is raised just long enough to introduce a small glass tube open at one end and con- taining the nitrate to be estimated. The quantity or nitrate taken must be equivalent to not more than 0.2 gm. of HNO3. The stopper is then reinserted, heat applied, and gradually increased until the reaction is complete. The free hydro- chloric acid liberates nitric acid from the nitrate and oxidation of a portion of the iron is effected. The ferrous chlorid is oxidized to ferric chlorid, as the equation shows, and the solution becomes at first dark brown through the presence of NO. As the heat is increased, the dark-brown color of the solution is gradually changed to yellow, as ferric chlorid is formed, and increases in intensity until the reaction is complete, then the color remains stationary and indicates com- pletion of oxidation. The solution is now allowed to cool, but the stream of CO2 or H gas is maintained. Forty mils of a solution of manganese sulphate are now added (this is not necessary if sulphuric acid is used instead of hydrochloric),

* Or fine piano-forte wire,

174 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

N and titration with — potassium permanganate solution begun,

in order to determine the quantity of unaltered ferrous salt remaining in the solution. Assuming that 89 mils were re- quired, the calculation is made as follows:

Since one molecule of HNO3 (63.01) reacts with three atoms of iron (167.46) the quantity of iron found to have

been oxidized, if multiplied by -~ — ^, will give the quantity

of nitric acid present.

Example. 1.5 gms. of iron wire, 99.6 per cent Fe= (1.494 gms. of iron), is dissolved in hydrochloric acid, as above described, and 0.6 gm. of potassium nitrate, KNOs (ioi.n), added. After oxidation, 98 mils of decinormal permanganate were

N required. Each mil of — KMnO4 = 0.005582 gm. of Fe.

0.005 582 gm.X 98 = 0.547 gm- °f oxidized iron. 1.494 gms. of iron were originally taken. Therefore,

1.494

— - gm. = the quantity of iron oxidized. 0.947 fc

Then

0.047X63.01

S7 gm. of HN03) which equals

0.947X63.01X101.11

0-371 gm- of

or 95.5 per cent pure.

N N

It is usually advisable to use an — instead of an -- KMnO4

5 10

solution.

ANALYSIS BY OXIDATION AND REDUCTION 175

Assay of Chlorates. The chlorate is made to react with a ferrous salt, which is added to its solution in excess, and which it oxidizes to the ferric state. The unoxidized ferrous salt is then titrated with permanganate. The chlorate present being calculated from the quantity of ferrous salt which it oxidized. The method is applicable to the assay of nitrates and chromates as well.

Dissolve an accurately weighed quantity, say 0.5 gm., of potassium chlorate in sufficient distilled water to make ico mils of solution. Take 10 mils of this solution (representing 0.05 gm.), add 25 mils of acidulated ferrous sulphate T.S* Place the whole in a flask provided with a Bun sen valve (see Fig. 41) and boil for ten minutes. Now cool the solution and add 10 mils of manganous sulphate T.S. and titrate the residual ferrous sulphate with tenth-normal potassium permanganate V.S. At the same time conduct a parallel experiment with another portion of 10 mils of acidulated ferrous sulphate T.S. The difference in the number of mils of the permanganate solution used in the two titrations, multiplied by 0.00204 gm., gives the weight of KClOs present.

The reaction is expressed as follows:

122.56 6X151.91

Example. In the blank experiment 25 mils of the acidulated ferrous sulphate T.S. required 26.2 mils of tenth-normal per- manganate V.S.

The residual ferrous sulphate in the actual assay required 2.4 mils of the permanganate V.S., hence 23.8 mils of the per-

* Dissolve 3 gms. of clear crystals of ferrous sulphate in 90 mils of distilled water, then add sufficient sulphuric acid to make 100 mils of solution.

176 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

manganate V.S. represent the chlorate which reacted with the ferrous sulphate.

0.00204X23.8X100

=97.! per cent.

Chromic Acid and Chromates. Chromic acid oxidizes ferrous salts in the same manner as nitric acid does. The reac- tion is thus expressed:

5Fe=- 334-92 20°

To an accurately weighed quantity of ferrous ammonium sulphate (Mohr's salt) FeSO4 + (NH4)2SO4 + 6H2O (the per- manganate titer of which is known), which is dissolved in a sufficient quantity of diluted sulphuric acid in an Erlenmeyer flask,* add a weighed quantity of the chromate or chromic acid in a concentrated aqueous solution. Warm the mixture on a water-bath, under a constant stream of carbon dioxid until the liquid assumes a clear green color. This occurs in a few minutes, and indicates complete reduction of the chromate.

Now allow the solution to cool, continuing the passage of carbon dioxid through the flask, and transfer the cold solution to a large beaker, and after diluting it to about 300 cc. and strongly acidifying it with sulphuric acid, titrate it

N for unoxidized ferrous salt by means of — potassium per-

manganate.

It is usually sufficient to mix the solutions cold, but it is better to employ heat after mixing. A large excess of

* This flask should be provided with a stopper having two perforations through which glass tubes are passed, one of these, which serves to convey carbon dioxid gas, should reach close to the surface of the liquid, the other tube should end just below the stopper and serve as the outlet tube. See Fig. 44.

ANALYSIS BY OXIDATION AND REDUCTION 177

ferrous salt is unnecessary. It is imperative to dilute the solution highly before titration, as then only can the end color point be accurately determined in the green solution. The use of an excess of sulphuric acid before titration is likewise demanded. A violet-red color marks the end-point, and unless too great a quantity of chromate be taken, or the solution be insufficiently diluted, it can be easily recognized. This method is applicable not only to free chromic acid and soluble chro- mates, but also to chromates which are insoluble in water.* It can therefore be employed for the indirect estimation of such bases as are precipitable by chromic acid, out of neutral, ammoniacal or acetic acid solutions, as for instance lead, bismuth and barium.

Finally, the method may be employed for the estimation of chromic oxids. The solution of the latter is treated with an excess of sodium carbonate, bromine water added, and heat applied until a clear solution results. This solution, which contains all of the chromium in the form of sodium chromate, is evaporated, the residue dissolved in dilute acetic acid and the chromium completely precipitated by means of lead acetate. The precipitated lead chromate is then treated as above.

The calculation is made with reference to the equation, in which it is shown that one molecule, 100 of chromic oxid (CrO3), is equivalent to three molecules (167.46) of metallic

100 iron. The quantity of iron oxidized, multiplied by -^ — 7,

gives the weight of chromic oxid present, and from this its equivalent in potassium, sodium, lead, bismuth or barium chromate is calculated.

*In the case of insoluble chromates the salt is shaken directly with the ferrous solution, and the mixture more highly diluted, and more strongly heated, than in the .case of soluble salts.

178 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

In the case of potassium dichromate (K2Cr2O7) one mole- cule (294.2) is equivalent to six atoms (334.92) of metallic

iron. The quantity of iron oxidized is multiplied by — — — .

y 334.92

Example. To 1.5 gms. of ammonio-ferrous sulphate (con- taining 0.2142 gm. Fe) add 0.1241 gm. of K2Cr2O7 (molecular weight 294.2), and after complete oxidation, titrate the solu-

N tion with — KMnO4 to determine the quantity of unchanged

ferrous salt. Thirteen mils are required. Each mil represents 0.005582 gm. of Fe.

Thus, 13X0.005582 gm. =0.0725 gm., the quantity of iron which was not oxidized by the dichromate. This, deducted from the quantity of iron originally added (0.2142-0.0725 ^=0.1417 gm.), gives the quantity which was oxidized.

Then,

0.1417X100

i6746~= 475+ gm* 3

or

0.1417X294.2 __

^—=0.1247 gm. of K2Cr2O7. 334-92

Example. To 1.5 gms. of ammonio-ferrous sulphate (con- taining 0.2142 gm. of Fe) add the precipitate of barium chromate obtained from 0.2491 gm. of BaCl2 + 2H2O (molec- ular weight 244.29) and after complete oxidation, titrate with

N - permanganate. 7.8 mils are consumed, thus 7.8X0.005582

= 0.04353 gm. the quantity of unoxidized iron present. Then 0.2142 -0.04353=0.17057 gm. of iron oxidized by the barium chromate.

ANALYSIS BY OXIDATION AND REDUCTION 179

Volumetric Analysis by Means of Potassium Bichromate

In some respects the dichromate possesses advantages over permanganate :

1. It may be obtained in a pure state.

2. Its solution does not deteriorate upon standing as does that of permanganate.

3. It is not decomposed by contact with rubber as the permanganate is, and may therefore be used in Mohr's burette. Its great disadvantage, however, is that when used in the estimation of ferrous salts the end-reaction can only be found by using an external indicator. The indicator which must be used is freshly prepared potassium ferricyanid T.S., a drop of which is brought in contact with a drop of the solu- tion being tested, on a white slab, at intervals during the titration, the end of the reaction being the cessation of the production of the blue color, when the two liquids are brought together. Thus the estimation by potassium dichromate is cumbersome, and very .exact results are not as easily obtained as with permanganate.

Besides ferrous salts, a great many other substances may be estimated by oxidation analysis with dichromate. Among them nitrates, sulphates, arsenous acid, barium, lead, ferric salts after reduction by stannous chlorid or an alkaline sul- phite, but not after reduction by means of metallic zinc The presence of the dissolved zinc salt interferes with the reaction of the ferricyanid indicator. Ferrous salts may be estimated in the presence of hydrochloric acid, by means of dichromate, without the precautions that apply in the case of permanganate. Chromium as chromate may be indirectly estimated; an excess of a solution of a ferrous salt being added and then the excess determined by dichromate. lodids, thiosulphates and alkalies may also be estimated by means of potassium dichromate.

180 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

T\T

Preparation of Decinormal Potassium Bichromate

10

N (K2Cr2O7 = 294.2; — V.S. = 4.903 gms. in 1000 mils).

4.903 gms. of pure potassium dichromate * which has been pulverized and dried at 120° C. is dissolved in suffi- cient water to make 1000 mils of solution.

It will be noticed that ^- of the molecular weight of the dichromate (expressed in grams) is taken in the preparation of 1000 mils of this solution. The reason for this is that one molecule of potassium dichromate when treated with an acid yields three atoms of nascent oxygen which are available for oxidizing purposes, thus

and since each atom of oxygen is equivalent to two atoms of hydrogen, one molecule of the dichromate must be equiv- alent to six atoms of hydrogen. Hence a normal solution of potassium dichromate, when used as an oxidizing agent, should contain one-sixth of the molecular weight, expressed in grams, in 1000 mils (see definition for normal solution) and its deci normal solution ^V

If a standard solution of potassium dichromate is to be made for use as precipitant, as in the titration of barium, one-fourth of the molecular weight is to be taken for 1000 mils of the normal solution, as explained in Chapter III.

Standard solution of potassium dichromate is sometimes used as a neutralizing solution for estimating alkalies, phenol- phthalein being used as indicator.

When used for this purpose the normal solution con- tains 147.1 gms. in i liter (one-half the molecular weight

* Potassium dichromate for use in volumetric analysis should respond to all the tests for purity given in the U.S.P., or it should be recrystaUized several times and then dried.

ANALYSIS BY OXIDATION AND REDUCTION 181

in grams). It is then the exact equivalent of any normal acid V.S.

2KOH + K2Cr2O7 = 2K2CrO4 + H2O.

2)112.2 2)294.2

56.1 gms. 147.1 gms., or 1000 mils normal V.S.

Decinormal potassium dichromate may also be used in conjunction with potassium iodid and sulphuric acid for standardizing sodium thiosulphate. Iodin is liberated from potassium iodid in this reaction. The reaction is expressed by the equation

Thus one molecule of the dichromate will liberate six atoms of iodin, therefore a normal solution should contain one-sixth of the molecular weight, and a decinormal solution A in 1000 mils. The solution is hence of the same strength as that which is used for oxidizing purposes. If the deci- normal solution containing 14.71 gms. in i liter is used, it

3N

has the effect of a — solution. 10

The decinormal solution which is used as an oxidizing agent is chemically equivalent to decinormal potassium per- manganate. When used for the purpose of liberating iodin from potassium iodid, it is the equivalent of an equal volume of decinormal sodium thiosulphate.

Standard potassium dichromate may be checked in the same way as standard permanganate, with pure iron wire.

ESTIMATION OF FERROUS SALTS WITH POTASSIUM DICHROMATE

For titrating ferrous salts the decinormal solution of dichro- mate is used in the following manner:

Make an aqueous solution of the ferrous salt, introduce

182 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

it into a flask, and acidulate it with sulphuric or hydrochloric acid. Now add gradually from a burette the decinormal potassium dichromate until a drop taken out upon a white slab no longer shows a blue color with a drop of freshly prepared potassium ferricyanid T.S. Note the number of mils of the standard solution used, multiply this number by the factor, and thus obtain the quantity of pure salt in the sample taken.

Ferrous salts strike a blue color with potassium ferricyanid, but as the quantity of ferrous salt gradually diminishes during the titration, the blue becomes somewhat turbid, acquiring first a green, then a gray, and lastly, a brown shade. The process is finished when the greenish-blue tint has entirely disappeared.

The reaction of potassium dichromate with ferrous salts always takes place in the presence of free sulphuric or hydro- chloric acid at ordinary temperatures. Nitric acid should not be used.

If it is desired to estimate ferric salts by this standard solution it is necessary to first reduce them. This may be done by metallic magnesium, sulphurous acid, the alkali sul- phites, or by stannous chlorid.

One molecule of potassium dichromate yields, under favor- able circumstances, 3 atoms of oxygen. This is shown by the following equation:

K2Cr2O7 = Cr2O3 + K2O + O3.

Here it is seen that the three liberated atoms of oxygen combine at once with the ferrous oxid, converting it into ferric oxid:

6FeO-f-O3 = Fe6O9 or

In the oxidation of a ferrous salt, the reaction takes place only in the presence of an acid.

ANALYSIS BY OXIDATION AND REDUCTION 183

The dichromate then gives up its oxygen. Four of its oxygen atoms combine at once with the replaceable hydrogen of the accompanying acid, the other three being liberated. The three oxygen atoms thus set free are available either for direct oxidation or for combination with the hydrogen of more acid. In the latter case a corresponding quantity of acidulous radicals is set free.

K2Cr2O7 +4H2SO4 = K2SO4 + Cr2(SO4)3 + 4H2O + O3.

In this case four of the liberated atoms of oxygen combine with eight of the atoms of hydrogen of sulphuric acid and liberate four SO4 radicals, which at once combine with the K2 and Cr2 of the dichromate. The other three atoms are set free. If seven sulphuric acid molecules are used instead of four molecules, the three free atoms of oxygen will liberate 3(S04):

K2Cr2O7 + 7H2SO4 = K2SO4 + Cr2(SO4)3 + 7H2O + (SO4)3.

If this liberation of 3(SO4) takes place in the presence of a ferrous salt, the 3(SO4) will combine with six molecules of the ferrous salt, converting it into a ferric salt:

6FeSO4 + K2Cr2O7 + 7H2SO4

= K2SO4 + Cr2(SO4)3 + 7H20 + [3Fe2(SO4)3].

If in the above case hydrochloric acid is used instead o/ sulphuric, fourteen molecules of the former must be taken to supply the neccessary hydrogen.

The seven liberated atoms of oxygen must have fourteen atoms of hydrogen to combine with.

Three of these atoms of oxygen liberate six univalent or three bivalent acidulous radicals.

184 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

Therefore, since one molecule of K^Cr^Oy will give up for oxidizing purposes three atoms of oxygen, which are equiva- lent chemically to six atoms of hydrogen, one-sixth of the molecular weight in grams of the dichromate, dissolved in sufficient water to make one liter, constitutes a normal solution, and one-tenth of this quantity of K2Cr2O7 in a liter, a deci- normal solution.

Thus the estimation of ferrous salts is effected by oxidizing them to ferric with an oxidizing agent of known power, the strength of the ferrous salt being determined by the quantity of the oxidizing agent required to convert it to ferric.

Saccharated Ferrous Carbonate (FeCO3= 115.82). One gm. of saccharated ferrous carbonate is dissolved in 10 mils of diluted sulphuric acid and the solution diluted with water to about 100 mils. The decinormal potassium dichromate is carefully added, until a drop of the solution taken out and brought in contact with a drop of freshly prepared solution of potassium ferricyanid ceases to give a blue color.

The number of mils of the dichromate solution is read off and the following equations applied:

6FeCO3 + 6H2SO4 = 6FeSO4 + 6H2O + 6CO2 ; 694.9 911.34

then

6FeCO3 or

6)694.9 6)911.34 6)294.2

10)115.84 10)151.89 10)49.03 jg.

11.584 gms. 15.189 gms. 4.903 gms., or 1000 mils — K2Cr2O7 V.S.

N Thus each mil of — K2Cr2O7 represents 0.011584 gm. of

pure ferrous carbonate or 0.005582 gm. of metallic iron.

ANALYSIS BY OXIDATION AND REDUCTION 185

If strong sulphuric acid is added to saccharated ferrous carbonate it will char the sugar, and a black mass of burnt sugar is obtained. This may be prevented by adding water first, and then, slowly, the sulphuric acid.

Instead of sulphuric acid, hydrochloric acid may be used.

FIG. 45-

This will not char the sugar, but the ferrous chlorid which is then formed is too readily oxidized by the air.

It has also been suggested that as hydrochloric acid so rapidly converts ordinary sugar into invert sugar as to render it easily attacked by the dichromate, it should be cautiously used, if at all. Phosphoric acid has none of these disad- vantages, and may be employed with good results.

In making estimations of ferrous salts with potassium dichromate, care should be taken to avoid atmospheric oxida- tion. It is good practice to calculate approximately how much of the standard solution will probably be required to

186 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

complete the oxidation, and then add almost enough of the standard solution at once, instead of adding it slowly.

A white porcelain slab is then got ready, and placed alongside of the flask in which the titration is to be performed. Upon this slab are placed a number of drops of the freshly prepared solution of potassium ferricyanid, and at intervals during the titration a drop is taken from the flask on a glass rod and brought in contact with one of the drops on the slab. The glass rod should always be dipped in clean water after having been brought in contact with a drop of the indicator. See Fig. 45.

When a drop of the solution ceases to give a blue color on contact with the indicator, the reaction is complete.

Pills of ferrous carbonate and mass of ferrous carbonate are assayed by the same process.

Ferrous Sulphate (FeSO4 + 71120 = 277.89). Dissolve about one gram of crystallized ferrous sulphate in a little water, add a good excess of sulphuric or hydrochloric acid, titrate with the decinormal potassium dichromate, as directed under Ferrous Carbonate, and apply the following equation:

6(FeS04.7H20) + K2Cr2O7 + 7H2SO4

6)1667.34

10)277.89 N

27.789 gms., or 1000 mils — K2Cr2O7 V.S.

= 3Fe2(SO4)3 + K2SO4 + Cr2(SO4)3

N Thus each mil of the — K2Cr2O7 V.S. represents 0.027789

gm. of crystallized ferrous sulphate or 0.015189 anhydrous. If i gm. of the salt is taken and dissolved as above, it should require about 37 mils of the standard solution, equivalent to about 100 per cent.

ANALYSIS BY OXIDATION AND REDUCTION 187

TABLE OF SUBSTANCES WHICH MAY BE ESTIMATED BY MEANS OF POTASSIUM PERMANGANATE OR POTASSIUM DICHROMATE.

Name.

Formula.

Molecular Weight.

N — Factor. 10

Acid, chromic

CrO3

IOO

O.OO23

" hypophosphorous

HPH2O2

66 04

o oo 1 6 c. i

" nitric

HNO3

63 .01

O.OO2I

" nitrous

HNO2

47 OI

O OO23C.

" oxalic (crystallized)

H2C2O4+2H2O

126 .05

o 0063

Barium dioxid

BaO2

160. T.J

0.00841;

Calcium chlorid

CaCl2

III OI

o oo c c c

" hypophosphite

Ca(PH2O2)2

I 7O. 17

o 002127

Ferric chlorid

FeCl3

162 2O

o 01622

" hypophosphite . ....

Fe(PH,O,),

2C.O Q4

o 02^004

" sulphate

Fe2(SO4),

300.85;

O OIQQQ

Ferrous carbonate

FeCO3

115 82

o 011^82

" oxid . ...

FeO

71 82

o 007182

" sulphate (anhydrous)

FeSO4

i si .80

o oi^iSo

" " (crystallized)

FeSO4+7H2O

277 80

o 027780

Ferrum (metallic)

Fe,

re 82

o 00^82

Hydrogen dioxid

H2O,

34 °

o 0017

Manganese dioxid

MnO2

86. Q?

o 004346

Potassium hypophosphite

KPH2O2

104. 15

0.0026

Sodium hypophosphite

NaPH2O2

88 oc

O OO22

" nitrite .

NaNO2

60. 01

o 0034?

Analysis by Indirect Oxidation

This method of analysis is based upon the oxidizing power of iodin.

lodin acts upon the elements of water, forming hydri- odic acid with the hydrogen, and liberating oxygen in a nascent state.

Nascent oxygen is a very active agent, and readily com- bines with and oxidizes many substances, such as arsenous oxid, sulphurous acid, sulphites, thiosulphates, hydrosulphuric acid, the lower oxids of antimony and their salts.

188 THE ESSENTIALS OF VOLUMETRIC ANALYSIS As2O3 + 2H20 + 2l2 = As205+4Hl ;

Therefore iodin is said to be an indirect oxidizer, and may be used for the estimation of a great variety of substances with extreme accuracy.

When iodin is brought in contact with certain oxidizable substances it is decolorized. This decolorization occurs as long as some of the oxidizable substance is present, and ceases when oxidation is complete. Hence when the yellow color of iodin shows itself in the solution being analyzed the reaction is known to be at an end. In most cases a more delicate end-reaction is obtained by using starch solution as an indicator. This gives a distinct and unmistakable blue color with the slightest excess of iodin.

In making an analysis with standard iodin solution, the substance under examination is brought into dilute solution (usually alkaline), the starch solution added, and then the iodin, in the form of a standard solution, is delivered in from a burette, stirring or shaking constantly, until a final drop colors the solution blue— a sign that a slight excess of iodin has been added.

N Preparation of Decinormal Iodin (1=126.92; — V.S. =

12.692 *gms. per liter). Dissolve 12.692 gms. of pure* iodin

* If pure iodin be not at hand, it may be prepared from the commercial article as follows:

Powder the iodin and heat in it a porcelain dish placed over a water- bath, stirring constantly with a glass rod for twenty minutes. Any adhering moisture, together with any cyanogen iodid, and most of the iodin bromid and iodin chlorid, is thus vaporized.

Then triturate the iodin with about 5 per cent of its weight of pure, dry potas- sium iodid. The iodin bromid and chlorid are thereby decomposed, potassium bromid and chlorid being formed and iodin liberated from the potassium iodid.

The mixture is then returned to the porcelain dish, covered with a clean glass funnel, and heated on a sand-bath. A pure resublimed iodin is then obtained.

ANALYSIS BY OXIDATION AND REDUCTION 189

in 300 mils of distilled water containing 18 gms. of pure potas- sium iodid. Then add enough water to make the solution measure at 25° C. exactly 1000 mils.

The solution should be kept in small glass-stoppered vials, in a dark place.

The potassium iodid used in this solution acts merely as a solvent for the iodin.

If pure iodin is used in making this solution, there is no necessity for checking (standardizing) it.

But if desired the solution may be checked against pure arsenous acid or sodium thiosulphate. It there is any doubt as to the purity of the iodin, it is best to take a larger quantity, say 14 gms. instead of the 12.692 gms. directed above, and then dilute the resulting solution to the proper strength after standardizing.

Standardization of Iodin V.S. by Means of a Decinor- mal Sodium Thiosulphate Solution. 25 mils of the iodin solution are accurately measured off into a beaker, and then

N

from a burette the — thiosulphate is delivered until the solu- tion is of a pale yellow color,- two or three drops of starch solu- tion are then added, and the titration with the thiosulphate solution continued until the blue color of starch iodid is dis- charged.

If the iodin solution is exactly decinormal, the 25 mils will require 25 mils of decinormal sodium thiosulphate to exactly complete the reaction.

If on the other hand more than 25 mils of thiosulphate solution is required, it indicates that the iodin solution is too concentrated, and must be diluted so as to correspond with the thiosulphate solution, volume for volume.

Example. Assuming that in the above titration, 27 mils of the thiosulphate solution were used, then each 25 mils of the

190 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

iodin solution must be diluted with water to make 27 mils in order to convert the iodin solution into a strictly decinormal solution. If, however, the iodin solution is found to be weaker,

N as evidenced by its using up less than its own volume of -

thiosulphate, its relative strength should be noted on the label of the container.

Thus if only 24.8 mils of the thiosulphate solution are

used up, then i mil of the latter equals — - mils or 1.008

24*0

mils of the iodin solution.

One mil of this iodin solution is equivalent to 0.992 mil

N of — thiosulphate, which is the same as saying i mil =0.992

N

mil of — iodin, or expressed in another way, i mil of this 10

iodin solution contains 0.01259+ gm. of iodin.

Such an iodin solution may be used as an empirical solu- tion, and in any assay the quantity of it (in mils) which is

consumed is divided by 1.008 or multiplied by — - or by 0.992,

0

and then multiplied by the decinormal factor for the sub- stance analyzed. Another way is to multiply the mils of this iodin solution used by the weight of iodin contained in each mil, and then by a fraction in which the numerator represents the quantity of the substance analyzed equal to an atom of iodin, and the denominator is the atomic weight of iodin.

Example, o.i gm. of arsenous acid consumes 20 mils of this empirical solution. How much absolute As20s does it

N contain? The — factor for As2O3 is 0.004948 gm.

,, ., , , v 20X0.004948

Method (a) = 0.0981 gm.

ANALYSIS BY OXIDATION AND REDUCTION 191

24.8 Method (b) 20 X X 0.004948 =--0.0981 gm.

0

Method (c) 20X0.992X0.004948 = 0.0981 gm.

40.48 Method (d) 20X0.012 $gX— ^ — =0.0981 gm.

It is a good plan to have the factors marked on the labels. In the above case the label may be marked

or X 0.992 or i mil =0.01259 gm. iodin.

D

Standardization of Iodin V.S. by Means of Arsenous Oxid. 0.2 gm. of pure resublimed vitreous arsenous oxid is weighed off very carefully into a flask 50 mils of water are added, and then, after the addition of 2 gms. or more of sodium bicarbonate, the mixture is gently warmed and shaken until the arsenous oxid is completely dissolved.*

To this solution a few drops of starch indicator are added, and then the iodin solution delivered carefully from a burette until a blue color marks the end of the reaction.

As2O3 + I4 + 2H2O = As2O5 + 4HI. 197.92 4x126.92

49.48 gms. of As2O3= 126.92 gms. of iodin;

N 4.948 " •" As2O3 = 12.692 " " " or 1000 mils — V.S.

* Arsenous oxid is much more readily soluble in alkali hydroxid, than in carbonated alkalies, therefore the following method of making the solution is preferred: 0.2 gm. of arsenous oxid is dissolved in a small quantity of boiling water with the aid of potassium hydroxid (free from sulphur), the solution is then acidified with hydrochloric acid, and then again made alkaline by the addition of sodium bicarbonate. The latter must be added in con- siderable excess, being careful, however, to avoid loss of solution during effervescence.

192 THE ESSENTIALS OF VOLUMETRIC ANALYSIS 0.2 gm. of As2O3 will require

1000 X 0.2 N .

— = 40.44 mils of a true — loam V.S.

4.948 10

Assuming that in the above titration 37.4 mils of the iodin solution were used, then the iodin solution is too concentrated and must be diluted so that each 37.4 mils will be made up to 40.44 mils.

After diluting in this way a new trial should be made.

It is a good plan to make a decinormal solution of the arsenous oxid by dissolving 4.948 gms. of the pure oxid and 30 gms. of sodium bicarbonate in sufficient water to make 1000 mils at 15° C. and to titrate this with the iodin solution. 25 mils of this solution should require for complete oxidation exactly 25 mils of the iodin solution, if the latter is strictly of decinormal strength.

The Starch Solution. This solution, which is used as an indicator in iodometric determinations, is made as follows :

One gm. of starch (potato, arrowroot, or corn starch) is tri- turated with 10 mils of cold water, until a smooth mixture is obtained, then sufficient boiling water is added, with constant stirring, to make 200 mils of a thin, translucent fluid. If the solution is not translucent it should be boiled for about three minutes, then allowed to cool, and filtered. This solution does not keep very long, in fact it becomes useless after standing one day, therefore it should be freshly prepared when required.

This indicator is very sensitive to iodin — it will detect one part of iodin in 3,500,000. If the solution is not clear, or contains flocks of insoluble starch, the characteristic beautiful blue color is not obtained with iodin; instead, a greenish or brownish color is produced, and the insoluble

ANALYSIS BY OXIDATION AND REDUCTION 193

particles are even colored black and are decolorized with difficulty.

The blue color which starch gives with iodin constitutes a very delicate indication of the slightest excess of iodin. This color is usually regarded as being due to the formation of a compound of starch and iodin, called iodid of starch. It is a compound of very unstable character and of doubtful composition.

Sodium thiosulphate behaves towards iodid of starch exactly as it does toward free iodin — it takes up the iodin and thus discharges the blue color.

Iodid of starch dissociates upon heating, but reunites upon cooling, hence it is advisable to avoid heat in estimations where starch is used as an indicator.

In order, to prevent the deterioration of this solution a few drops of chloroform may be added; this will -preserve it for a long time. Oil of cassia is also recommended as a preservative. Moerk adds 2 mils of the oil to a liter of the cooled starch solution. Zinc chlorid or iodid added to the boiling starch solution will prevent its decomposition for a long time. A starch solution so made, however, should not be used in titrations of sulphids, because zinc reacts with sulphids.

In the case of solutions containing carbonates, the pre- cipitate of zinc carbonate is so small in amount that it does not interfere in the least with the recognition of the end- reaction tint. Mercuric iodid is also a very valuable preserva- tive.

o.o i gm. of mercuric iodid in a liter of the starch solution is quite sufficient. A very satisfactory indicator is the com- mercial soluble starch which is made by heating potato starch with glycerin and precipitating the starch by repeated treat- ment with alcohol. This starch dissolves readily in hot water,

194 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

forming a clear solution, which gives a very delicate reaction with iodin. It is best preserved under alcohol, the latter being removed by filtration and evaporation, when the starch is wanted for making a solution.

In making starch solution for use as an indicator, long continued boiling should be avoided, as this converts some of the starch into dextrin.

On the Use of Sodium Bicarbonate in Titrations with Iodin. In these titrations an excess of alkali is necessary in order to neutralize the hdriodic acid formed.

If the hydriodic acid is not removed by neutralization it will react with the arsenic oxid (As2O 5), reducing it to arsenous oxid (As2O3) and liberate iodine, as shown by the following equation :

4HI + As2O5 = As2O3 + 2H2O + 2l2.

Sodium bicarbonate is usually employed to neutralize the HI and should be used in slight excess.

Alkali hydroxids or carbonates cannot be used for this purpose, because they react with free iodin or even with starch iodid. Bicarbonates ordinarily have no such action, and therefore sodium bicarbonate is usually directed to be added in excess to the solution to be titrated with iodin.

It is well known that sodium hydroxid solution reacts with free iodin, with formation of hypoiodite and iodid.

2NaOH+I2 = NaIO+NaI + H2O, the hypoiodite quickly forming iodate. v K

3NaIO = 2NaI + NaIO3. It is also now a recognized fact that sodium carbonate

ANALYSIS BY OXIDATION AND REDUCTION 195

is partly hydrolyzed when in solution, with formation of some sodium hydroxid, as per equation,

Na2C03 + H2O = NaOH + NaHCO3.

It therefore reacts in much the same way with iodin as the hydroxid, though to a less extent.

On the other hand, it is generally supposed that bicar- bonate of soda is without effect on iodin, and when, in iodo- metric estimations, addition of sodium bicarbonate is indicated, little attention is given to amount added, as long as it be in excess.

The experiments of W. A. Puckner, Proc. A. Ph. A., 1904, 408, prove that we are entirely wrong in the supposition that sodium bicarbonate has no effect upon iodin. He showed that when using i to 2 gms. of the bicarbonate, an error of 1.5 to 4.5 mils of decinormal iodin may be introduced, even when the sodium bicarbonate used is of exceptional purity, and especially proven to be free from carbonate, sulphite or thiosulphate. He shows that when sodium bicarbonate is added to a decinormal iodin solution, residual titration with sodium thiosulphate will show a considerable loss of free iodin, which went into combination in some form or other (probably iodid) and that the quantity so lost is proportional to (i) the mass of sodium bicarbonate; (2) the time of the interaction (the reaction is slow); (3) the concentration of the solution; (4) the temperature, and (5) the size of the flask in which reaction occurs. These phenomena are due to the fact that sodium bicarbonate when dissolved in water under- goes hydrolysis, thus

2NaHCO3 = Na2CO3 + H2CO3 or (H2O + CO2).

This breaking up of the NaHCO3 into Na2CO3 and H2CO3, and the latter into H3O and CO2, continues until the

196 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

pressure of the CO 2 above is equal to the pressure of the gas in the solution, i.e., until equilibrium has been reached. In concentrated solutions of NaHCOs the amount hydrolyzed is much greater than in dilute solutions. An elevation of temper- ature materially increases the absorption of iodin.

Less iodin is lost when smaller flasks are used, provided the glass stopper completely shuts off communication with the atmosphere. The CO2 will escape from the solution until its pressure in the solution is equal to that of the gas above. Thus, since a larger volume of air is contained in a larger flask, more CO2 passes from the liquid before equi- librium is established, hence more NaHCO3 is decomposed, and more iodin in consequence absorbed.*

Reasoning from the above observations it may be said that: i, though sufficient sodium bicarbonate be used to more than neutralize the hydriodic acid formed, the solution titrated should be well diluted; 2, that the titration should be done cold; 3, that the titration should be done in small stoppered flasks, and 4, it should be done quickly.

Estimation of Arsenous Compounds

These compounds are estimated by means of iodin in a manner simitar to that described under standardization of iodin solution by means of arsenous oxid. The method is as follows:

Arsenous Oxid (Arsenous Acid, Arsenous Anhydrid, Arsenic Trioxid)(As2 03=197. 92). When arsenous acid is brought in contact with iodin in the presence of water and an alkali,

* For further study of equilibrium, see the work of Dr. H. N. McCoy, Am. Ch. J., vol. XXIV, 437

ANALYSIS BY OXIDATION AND REDUCTION 197

it is oxidized into arsenic acid and the iodin is decolorized. The reaction is:

The alkali should be in sufficient quantity to combine with the hydriodic acid formed, and must be in the form of potassium or sodium bicarbonate.

The hydroxids or carbonates should not be used. Starch solution is used as the indicator, a blue color being formed as soon as the arsenous acid is entirely oxidized into arsenic acid.

Dissolve about 0.2 gm. of arsenous acid accurately weighed, in 20 mils of boiling distilled water by the gradual addition of sodium hydroxid T.S. until complete solution results. Neutralize this solution with a diluted sulphuric acid V.S., using phenolphthalein T.S. as indicator, cool, dissolve in it 2 gms. of sodium bicarbonate and titrate the mixture with decinormal iodin V.S., using starch T.S. as indicator, shak- ing or stirring the mixture constantly until a permanent blue color is produced. The following equation illustrates the reaction :

As2O3 + 2H2O + 2l2 = 4HI -1- As2O5.

10)49.48 10)126.92 N

4.948 gms. 12.692 gms. or 1000 mils — I V.S.

N Thus each mil of — I V.S. represents 0.004948 gm. of

pure As2O3.

Solution of Arsenous Acid and Solution of Potassium Arsenite are assayed in the manner above described. Twenty mils are taken for the assay, i gm. of sodium bicarbonate

198 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

added, the solution diluted to 100 mils, and titrated with the decinormal iodin solution. No indicator is required though starch may be used.

In the case of solution of potassium arsenite, it is advisable to slightly acidify with hydrochloric acid, then to make the solution alkaline with sodium bicarbonate before titrating. The hydrochloric acid is employed here in order to neutralize any potassium hydroxid which may have been formed through hydrolysis of the potassium bicarbonate contained in the solution.

Arsenous lodid (Asia = 455. 72). This salt is estimated in the same way as described for arsenous oxid. The reaction is illustrated by the following equation:

2AsI3 + 5H2O + 2l2 = As2O5 + ioHI.

4)911.44

io)_227.86_ ^

22.786 gms.= 1000 mils — I V.S.

The Direct Percentage Assay of Arsenous Compounds. A. quantity of arsenous acid is taken, which is equal to the weight of pure As2O3, oxidized by 100 mils of decinormal

iodin, i.e., 0.4048 gm.

N If 0.4948 gm. of the sample be taken then each mil of —

I V.S. will represent T^ of this quantity or i per cent of pure As2O3. In the case of weak solutions of arsenic, as liquor acidi arsenosi, liquor potassii arsenitis, etc., which contain only one per cent of arsenous acid. A much larger quantity should be taken for analysis, otherwise the quantity of standard iodin solution used will be so small as to diminish the accuracy of the test.

Thus, if only 0.4948 gm. of either of the above solutions

ANALYSIS BY OXIDATION AND REDUCTION 199

be taken, no more than i mil of the standard solution would be required. It is better to take enough of the preparation to use up 30 to 50 mils of standard solution.

Estimation of Antimony Compounds

Antimonous oxid (Sb2C>3) or any of its compounds may be accurately estimated by means of iodin, in a manner similar to that described for the estimation of arsenous oxid, the antimonous oxid being oxidized to antimonic oxid, as per equation,

Sb2O3 + 2H2O + 2l2 = 4HI + Sb2O5.

The antimonous oxid is dissolved and kept in solution by the aid of tartaric acid, and then after the addition of an

N

excess of sodium-bicarbonate, the solution is titrated with -

10

iodin, using starch as an indicator. Accurate results can only be obtained if the solution is sufficiently alkaline to neutralize the hydriodic acid formed during the reaction. The titration should -be conducted without delay after the addition of the bicarbonate, otherwise a precipitate of antimonous hydrate will be formed, upon which iodin has little effect. The anti- mony must be in solution to be properly attacked by the iodin.

To o.i gm. of antimonous oxid 20 mils of water are added and the mixture heated to boiling; to this tartaric acid is added in small portions at a time until the oxid is completely dissolved. The solution is then neutralized by means of sodium carbonate, and sufficient of a saturated solution of sodium bicarbonate is added to make the solution distinctly alkaline (about 10 milsjis required for o. i gm. of the antimonous oxid). The mixture is now ready for titration with standard

200 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

iodin solution. This should be done immediately. The appearance of a permanent blue color marks the end-point, starch being used as indicator.

Sb2O3 + 2H2O + 2l2 = 4HI + Sb2O5.

4)288.4 4)507-68

io)72.i_ 10)126.92 N,

7.21 gms. 12.692 gms. or 1000 mils — ! V.S.

N One mil of — iodin represents 0.00721 gm. of Sb2O3.

The solution of the oxid may be made by means of hydro- chloric acid, and after adding a portion of tartaric and diluting with water, sodium bicarbonate is added and the titration conducted as above.

Other compounds of antimony may be estimated in the same way. Antimonic compounds are reduced to antimonous sulphid (Sb2Ss) by precipitating with hydrogen sulphid, and after thoroughly washing the precipitate, dissolving it in hydro- chloric acid; thus a solution of antimonous chlorid is obtained from which all traces of hydrogen sulphid are expelled by boiling. This solution is diluted with water, tartaric acid added, and finally, after making alkaline with sodium bicar- bonate, titrated with the standard iodin solution as above described.

Antimony and Potassium Tartrate (Tartar Emetic) [2(K[SbO]C4H4O6)+H2O = 664.7]. One gm. of the salt is dissolved in sufficient water to make 100 mils. 30 mils of this solution, representing 0.3 gm. of the salt, are taken for assay. 20 mils of a cold saturated solution of sodium bicarbonate are added, then a little starch solution, and the

N mixture titrated with — iodin until a permanent blue color

appears.

The calculation is as follows:

ANALYSIS BY OXIDATION AND REDUCTION 201 2K(SbO)C4H,O6 + H2O + 2l2

4)664.7 4)507-68

10)166.17 10)126.92

16.617 gms. 12.692 gms. = 1000 mils — I V.S.

= 4HI + 2KHC4H406 + 2HSb03.

N i mil of — iodin represents 0.016617 gm. of K(SbO)C4H4O6

+|H2O (crystallized tartar emetic).

K(SbO)C4H406 (anhydrous tartar emetic) =323.34.

2)323-34 10)161.67 N

16.167 gms. = 1000 mils — V.S.

N Thus i mil of — iodin represents 0.016167 gm. of anhydrous

tartar emetic.

Estimation of Sulphurous Acid and Sulphites

These substances may be accurately estimated by means of a standard solution of iodin. When sulphurous acid or one of its salts is brought in contact with iodin, a complete oxidation takes place. The sulphurous acid is oxidized to sulphuric acid and the sulphite to a sulphate, as the equations show:

H2S03 + H20 + 12 = 2HI + H2S04,

Na2SO3 + H2O + 12 = 2HI + Na2SO4, NaHSO3 + H2O + 12 - 2HI + NaHSO4.

There are two methods which may be employed. In one method the substance is brought into solution in water, an excess of sodium bicarbonate is added, and then the standard iodin solution is run in until a faint yellow color of free iodin

202 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

marks the end-reaction. If starch solution is used as indicator the end-point is the production of a blue color. The other method is that of Giles and Shearer, who, in a very voluable series of experiments detailed in the J. S. C. L, III, 197, and IV, 303, suggest the following modification:

The weighed sulphurous acid or the sulphite (in fine

N powder) is added to an accurately measured excess of :

iodin, without diluting with water. After the mixture has been allowed to stand for about one hour, with frequent shaking, the oxidation is complete, and the excess of iodin

N is ascertained by titrating back with — sodium thiosulphate.

The quantity of the latter deducted from the quantity

N of — iodin solution added, will give the quantity of the latter

which reacted with the sulphite.

The neutral and acid sulphites of the alkalies, alkali earths, and even zinc and aluminum, may be accurately estimated in this manner. The less soluble salts requiring, of course, more time and shaking, to insure their complete oxidation. The latter is the U. S. P. method.

Sulphurous Acid. This is an aqueous solution of sulphur dioxid (802 = 64.07).

Sulphurous acid when brought in contact with iodin is oxidized into sulphuric, the iodin being decolorized because of its union with the hydrogen of the accompanying water, forming hydriodic acid.

Two grams of sulphurous acid are taken and diluted with distilled water (recently boiled and cooled*) to about 25

*" Recently boiled" insures absence of air, the oxygen of which would partially oxidize the sulphurous acid, and "cooled" is directed to avoid loss of SO2, which would occur if hot water were used.

ANALYSIS BY OXIDATION AND REDUCTION 203

mils. Two grams of sodium bicarbonate are added, and then the decinormal iodin V.S. is delivered into the solution (to which a little starch solution had been previously added) until a permanent blue color is produced.

The following equations, etc., show the reactions that take place:

H2S03 + H20 + 12 = 2HI + H2S04.

Sulphurous acid being, however, looked upon as a solu- tion of SO2 in water, the quantity of this gas is generally estimated in analyses.

H20,S02 + H20 + I2 = 2HI + H2S04.

2)64.07 2)253.84

10)32.04 10)126.92

3.204 gms. 12.692 gms.

N

Thus each mil of — iodin consumed before the blue coior 10

appears, represents 0.003204 gm. of SO2.

The Residual Method. Because of the volatile nature of this acid the residual method described below is the most satisfactory in that loss by volatilization is avoided and com- plete oxidation of the acid assured. When the direct method described is used there is more or less loss of SO2 and incom- plete oxidation, with separation of sulphur.

Measure 2 mils of the sulphurous acid into a stoppered weighing flask and find its exact weight. Add this to the

N

so mils of — iodin contained in a titration flask and let the 10

N

solution stand for about five minutes. Then titrate with —

10

sodium thiosulphate until the mixture is decolorized. Sub- tract the number of mils of the thiosulphate used from the

204 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

N 50 mils of — iodin added, and multiply the difference by the

N

— factor for SO 2, which is 0.0032035 gm. This will give the

weight of SO 2 in the quantity of acid taken for analysis.

When a solution containing sulphur dioxid is to be measured

by means of a pipette, it is never advisable to fill the instrument by suction in the usual manner, as this would cause a loss of the gas. A better plan is to fill the pipette by pressure by the use of an arrangement similar to that shown in Fig. 46.

The solution containing sulphur dioxid or other volatile substance is poured into a flask which is provided with a stopper through which two glass tubes pass; one of these tubes reaches nearly to the bottom of the flask and the other projects about one-half an inch below the stopper and is bent outward above. To the upper end of the former the pipette is attached by means of a piece of rubber tubing. By blowing into the flask through the shorter tube the liquid is caused to rise and fill the pipette, which may then be easily pulled out of the rubber tube connection.

Sodium Sulphite (Na2SO3 + 7H2O = 252.2) . Take 0.5 gm. of the finely powdered crystals,

N FIG. 46. add to 50 mils of — iodin, contained in a

loo-mil glass-stoppered flask, and allow to stand for one

N hour (shaking frequently); then titrate with — sodium thio-

sulphate until the color is discharged. The reaction is expressed as follows:

ANALYSIS BY OXIDATION AND REDUCTION 205

Na2SO3 + 7H2O + I2 = 2HI + Na2SO4 + 6H2O.

2)252.2 2)253.84

10)126.1 10)126.92 T

1 2. 6 1 gms. 12.692 gms. or 1000 mils — iodin V.S.

Thus each mil of the standard solution represents 0.01261 gm. of crystallized sodium sulphite.

If i gm. of the salt is taken, to find the percentage multi- ply the factor by the number of mils of standard solution consumed, anpl the result by 100.

Potassium Sulphite (K2SO3 4- 2H2O = 194.37). Operate upon 0.5 gm. in the same manner as for sodium sulphite.

K2S03 + 2H20 + I2 = 2HI + K2S04 + H20.

2)194-37

10) 97.18

9.718 gms. or looo mils of standard V.S.

N Each mil of the — iodin represents 0.009718 gm. of crys-

tallized potassium sulphite.

Sodium Bisulphite (NaHSO4= 104.08). Operate upon about 0.25 gm. in the same manner as for sodium sulphite, and apply the following equation:

Sodium Thiosulphate (Sodium Hyposulphite) (Na2S2O3 = 248.24). This salt, when brought in contact with iodin, is converted into sodium iodid and sodium tetrathionate. The reaction is expressed by the equation

2Na2S2O3 + 12 = 2NaI + Na2S4O6. It is estimated as follows: i gm. of the salt is dissolved

206 'THE ESSENTIALS OF VOLUMETRIC ANALYSIS

in 20 mils of water, a few drops of starch solution are added,

N and then the — iodin is delivered in from a burette, until

the appearance of blue starch iodid indicates an excess of iodin.

Hydrogen Sulphid (H2S = 34.07) . When iodin and hydrogen sulphid are brought together in solution the following reaction occurs :

The reaction is not regular, however, when performed in an acid solution, but in the presence of alkali bicarbonates the results are constant. The method may be employed for the estimation of alkali sulphates.

The process may be conducted as follows:

Into 30 mils of a cold saturated solution of sodium bi- carbonate, contained in a 5oo-mil flask, measure a suitable quantity of the solution of hydrogen sulphid, stopper the flask and mix contents by shaking. Dilute the solution with about 300 mils of water, add starch solution and titrate

N with — iodin V.S. until a distinct and permanent blue color

appears.

N Each mil of — iodin represents 0.0017035 gm. of H2S.

The residual method may also be employed. A suitable

N volume of the sample is added to an excess of — iodin V.S.

mixed with some sodium bicarbonate solution, the solution is

N thoroughly shaken, and then titrated with — thiosulphate; the

N quantity of the latter, deducted from the quantity of — iodin

ANALYSIS BY OXIDATION AND REDUCTION 207

N .

- lodin 10

which reacted with

added, gives the quantity of

the H2S.

Sulphids. Soluble sulphids may be estimated by either of the above methods. The solu- tion of sulphid containing about 0.2 gm. being treated like an H2S solution.

Sulphids insoluble in wale? but decomposable by dilute acids may be estimated as fol- lows:

A weighed quantity of the sulphid is introduced into a flask, provided with a double perforated stopper; through one of the perforations the stem of a separatory funnel is passed, through the other a glass delivery tube (see Fig. 47). The funnel tube extends nearly to the bottom of the flask and is bent to form a hook, the opening of which is under water. The delivery tube begins at the lower end of the stopper and ends in another flask containing sodium bicarbonate solution. The funnel contains diluted sulphuric acid, which, upon open- ing the glass stop-cock, is allowed to flow into the flask, upon the contained sulphid; the H2S liberated is conducted into the solution of sodium bicarbonate which absorbs it com- pletely. A current of air aspirated through the apparatus insures absorption of the entire H2S developed. The sodium bicarbonate solution of H2S is then titrated with the standard iodin, in the presence of starch.

FIG. 47-

208 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

TABLE OF SUBSTANCES WHICH MAY BE ESTIMATED BY MEANS OF STANDARD IODIN SOLUTION

Name.

Formula.

Molecular Weight.

—Factor. 10

Acid sulphurous

H,SO4

82 oo

O OO4.IO3

Antimonous oxid

Sb2O3

288.4

o 00721

Antimony and potas'm tartrate Arsenous iodid

2 [K(SbO)C4H8O6] + H2O AsT3

664.7 4.rr 72

0.0l66l7

o 022786

" oxid

As/),

IO7 02

o 004048

Cyanogen

CN

26 01

o 001300^

Hydrogen sulphid

H,S

7/1 OO

O OOI7O3C

Iron (metallic)

Fe,

in 68

o 002701

ATercuric chlorid

HeCl,

o 027002

M^ercurous chlorid . .

HgCl

0-26 06

o 023^46

Potassium cyanid

KCN

6c ii

O OO T. 2 ^ ^

" sulphite (anhydrous) (crystallized) Sodium bisulphite

K2S03 K2SO3+ 2H2O NaHSO3

158-27 194.30

0.007913 0.009718

o 00^204

" sulphite (anhydrous).. " " (crystallized) . " thiosulphate

Na2SO3 Na2SO3+7H2O Na,S2O,+ sH,O

126.07 252.18

248 22

o . 006303 0.01261 o 024824

Sulphur dioxid

SO,

64 O7

O OO32O3C

Tin in stannous compounds .. . Zinc

Sn2 Zn2

238.0 1^0.74

0.00595 O.OO3268

Estimation of Substances Readily Reduced.

Any substance which readily yields oxygen in a definite quantity, or is susceptible of an equivalent action, which involves its reduction to a lower quantivalence, may be quan- titatively tested by ascertaining how much of a reducing agent of known power is required by a given quantity of the sub- stance for its complete reduction.

The principal reducing agents which may be employed in volumetric analysis are sodium thiosulphate, sulphurous acid, arsenous acid, oxalic acid, metallic zinc, and magnesium.

The sodium thiosulphate is the only one which is employed officially in the U. S. P. in the form of a volumetric solution.

ANALYSIS BY OXIDATION AND REDUCTION 209

It is used in the estimation of free iodin, and indirectly of other free halogens, or compounds in which the halogen is easily liberated, as in the hypochlorites, etc.

Estimations Involving the Use of Sodium Thiosulphate V.S.

(lodometry)

When sodium thiosulphate acts upon iodin, sodium tetra- thionate and sodium iodid are formed, and the solution is decolorized.

This reaction takes place in definite proportions: one molecular weight of the thiosulphate absorbs one atomic weight of iodin.

2Na2S2O3 + 12 = 2NaI + Na2S4O6.

Chlorin cannot be directly titrated with the thiosulphate, but by adding to the solution containing free chlorin an excess of potassium iodid, the iodin is liberated in exact proportion to the quantity of chlorin present, atom for atom.

C12 + 2KI = 2KC1+I2.

Then by estimating the iodin, the quantity of chlorin is ascertained. All bodies which contain available chlorin or bromin, or which when treated with an acid evolve chlori:i or bromin, may be estimated by this method.

Also, bodies which contain available oxygen, and which when boiled with hydrochloric acid evolve chlorin, such as manganates, chromates, peroxids, etc., may be estimated in this way.

Solutions of ferric salts, when acidulated and boiled with an excess of potassium iodid, liberate iodin in exact propor- tion to the quantity of ferric iron present.

210 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

Arsenates, copper salts, and others also liberate iodin from potassium iodid, quantitatively.

Thus sodium thiosulphate may be used in the estimation of a great variety of substances with extreme accuracy.

Preparation of Decinormal Sodium Thiosulphate (Hypo- sulphite) (Na2S2C>3 + 5^120 = 248. 22; contains 24.822 gms. in i liter). Sodium thiosulphate is a salt of thiosulphuric acid in which two atoms of hydrogen have been replaced by sodium; it therefore seems that a normal solution of this salt should contain one-half the molecular weight in grams in one liter.

But this salt is used chiefly for the estimation of iodin, and, as stated before, one full molecular weight reacts with and decolorizes one atomic weight of iodin, and since one atom of iodin is chemically equivalent to one atom of hydrogen, a full molecular weight of sodium thiosulphate must be con- tained in a liter of its normal solution.

Sodium thiosulphate is easily obtained in a pure state, and therefore the proper weight of the salt, reduced to powder and dried between sheets of blotting-paper, may be dissolved directly in water, and made up to one liter.

A stronger solution than decinormal is usually made, its titer found, and then the solution diluted to the proper measure.

Thirty gms. of selected crystals of the salt are dissolved in enough water to make, at or near 25° C., 1000 mils.

This concentrated solution is then standardized by one of the following methods :

a. Standardization by Means of JQ Iodin. Transfer 10

mils of this solution into a flask or beaker, add a few drops of starch T.S., and then gradually deliver into it from a burette decinormal iodin solution, in small portions at a time, shaking the flask after each addition, and regulating the flow to drops toward the end of the operation. As soon as a blue color is

ANALYSIS BY OXIDATION AND REDUCTION 211

produced which does not disappear upon shaking, but is not deeper than pale blue, the reaction is completed. Note the number of mils of iodin solution used, and then dilute the thiosulphate solution so that equal volumes of it and the deci- normal iodin will exactly correspond to each other, under the above-mentioned conditions.

Example. The 10 mils of sodium thiosulphate, we will assume, require 10.7 mils of decinormal iodin.

The sodium-thiosulphate solution must then be diluted in the proportion of 10 mils to 10.7 mils, or 1000 mils to 1070 mils.

After the solution is thus diluted a new trial should be made, in the manner above described, in which 50 mils of trie thiosulphate solution should require exactly 50 mils of the decinormal iodin to produce a faint blue color.

The solution should be kept in small dark amber-colored, glass-stoppered bottles, carefully protected from dust, air, and light.

One mil of this solution is the equivalent of:

Iodin 0.012692 gram.

Bromin 0.007992 "

Chlorin 0.003546 "

Iron in ferric salts 0.005584 "

b. Standardization by Means of Potassium Dichromate.

The potassium dichromate should be pure, if not it should be purified by triple recrystallization and then heated in a porce- lain crucible until the entire mass is just fused; then set aside to cool in a desiccator over calcium chlorid or sulphuric acid. The mass falls to a crystalline powder. Of this purified potassium dichromate, weigh off a definite quantity, say 0.2 gm., dissolve it in a small quantity of water, and pour the solution into a beaker containing 2 gms, of pure potassium

212 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

iodid (free from iodate) and 100 mils of water. Acidulate the solution with 5 mils of concentrated hydrochloric or sulphuric acid, cover the beaker, and let stand for about five minutes, then titrate with the thiosulphate solution to be standardized (using starch as an indicator) until the blue color is just discharged. The calculation is then made as follows :

K2Cr2O7 + 6KI + i4HCl = 2CrCl3 + 8KC1 + 7H2O + 3X2.

294.4 996.12 76l-52

Thus 294.4 gms. of potassium dichromate oxidizes 996.12 gms. of potassium iodid and liberates therefrom 761.52 gms. of iodin.

The potassium iodid must be in excess; in fact for each atom of iodin liberated one molecule of potassium iodid must be present at the completion of the reaction in order to keep the iodin in solution, and thus prevent loss by volatilization.

If 294.4 gms. of potassium dichromate liberate 761.52 gms. of iodin, 0.2 gm. will liberate

761.52X0.2

— = 0.5173 gm. of iodin. 294.4

N 0.012692 gm. of iodin = i mil of — thiosulphate.

N 0.5173 =40.172 mils of — thiosulphate.

Therefore if in the above assay 37.60 mils of the thiosul- phate V.S. were consumed, it must be diluted so that each 37.60 mils will measure 40.72 mils in order to convert the thio- sulphate solution into a true decinormal solution. A new trial should then be made with the diluted solution to see if its strength is correct.

It is usually more convenient to use a decinormal dichro-

ANALYSIS BY OXIDATION AND REDUCTION 213

mate V.S. for standardizing decinormal thiosulphate V.S. This may be done as follows :

Dissolve two grams of pure potassium iodid in a small quantity (10 to 15 mils) of sulphuric acid (i-io). Place this into a 5oo-mil flask and add to it very slowly 25 mils (accu- rately measured) of decinormal potassium dichromate V.S., mix well and let stand for five minutes, the flask being kept closed. The thiosulphate solution to be standardized is then run in from a burette in small portions, shaking after each addition until the solution is a pale yellow color. Two mils of starch solution are then added and the titration continued drop by drop until the blue color of starch iodid is just dis- charged.

Several trials are made, and from an average of three or four closely agreeing results the quantity of water required for dilution to decinormal strength is readily calculated.

Example. Assuming that in the above trial 23.8 mils of the thiosulphate solution were required to react with the iodin liberated by 25 mils of decinormal dichromate, then the thio- sulphate must be diluted with distilled water so that each 23.8 mils will measure 25 mils. After dilution a new trial should be made, in which 25 mils of decinormal dichromate should require when treated as above described exactly 25 mils of the thiosulphate.

c. Standardization, by Means of Potassium Bi-iodate. This method depends upon the fact that when potassium iodid and bi-iodate are brought together in the presence of a small quantity of an acid, an equivalent amount of iodin is set free. The reaction is illustrated by the equation:

KH(IO3)2 + ioKI + nHCl = 12! + nKCl + 6H2O.

12)389.94 12)1523.04

10)32.495 10)126.92 N

3 . 2495 gms- 12.692 gms. or 1000 mils — V.S.

214 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

Thus it is seen that one molecule of the bi-iodate causes the liberation of 12 atoms of iodin, against which the thio- sulphate solution is standardized.

In order not to use up too large a quantity of the thio- sulphate solution in the titration, a very small quantity of the bi-iodate must be taken, and since a small error in weighing this would entail a relatively large error in the results it is best to use the bi-iodate in the form of a solution of known strength and thus obviate the difficulty.

A decinormal solution of the bi-iodate, i.e., one containing in 1000 mils 3.2245 gms. of the salt may be used to advantage. Such a solution will keep unchanged for years, and may be employed as follows for the standardization of sodium thio- sulphate solution:

Into a glass-stoppered flask of 250 mils capacity introduce 10 mils of a 5 per cent solution of potassium iodid,* i mil of diluted hydrochloric acid and exactly 25 mils of the above potassium bi-iodate solution.

This brownish-yellow solution is now titrated with the sodium thiosulphate solution which is slowly delivered from a burette until the solution becomes pale yellow in color; a few drops of starch solution are now added, and the titration continued (the flow being reduced to drops) until the blue color is just discharged. During the titra- tion, the flask should be frequently stoppered and vigorously shaken.

When the blue color is discharged, note the number of mils used. If an exactly decinormal thiosulphate solution is taken, 25 mils will be required to react with the 25 mils of potassium bi-iodate solution, under the above conditions. If,

* The potassium iodid must be in sufficient quantity not only to react with the bi-iodate quantitatively, as shown in the equation, but also to dissolve the iodin which is liberated.

ANALYSIS BY OXIDATION AND REDUCTION 215

on the other hand, only 22 mils of the thiosulphate solution are used in the titration, then the latter is too strong, and must be diluted so that each 22 mils will measure 25 mils in order to make it strictly decinormal.

d. Standardization by Means of Potassium Permanga- nate. Sodium thiosulphate solution may be accurately standardized by means of potassium permanganate, if the other substances used for this purpose are not at hand. The method is easily understood by referring to the iodometric standardization of potassium permanganate solution. See page 145.

Estimation of Free lodin (1=126.92). lodin in the dry state or in the form of a tincture is brought into aqueous solution by means of pure potassium iodid and then titrated with standard sodium thiosulphate. The potassium iodid is used here to dissolve the iodin; it must be free from iodate (KIO3) because the presence of this salt would cause a libera- tion of iodin from the potassium iodid.

Dry iodin is assayed as follows:

About 0.5 gm. of iodin are placed in a tightly stoppered weighing bottle and accurately weighed. One gram of potas- sium iodid and 50 mils of water are added, and when the iodin

N is dissolved, — sodium thiosulphate is delivered from a burette

in small portions at a time, shaking after each addition until the reaction is nearly completed, and the solution is of a faint yellow color. A few drops of starch indicator are now added, and the titration with the thiosulphate continued drop by drop until a final drop just discharges the blue color. The number of mils of the thiosulphate solution used is noted. This number, multiplied by the decinormal factor for iodin, gives the weight of the latter present in the sample assayed.

216 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

I2 + 2(Na2S2O3 + 5H2O) = Na2S4O6 + 2NaI + ioH2O.

2)253.84 2)406.96

10)126.92 10)248.48 N

12.692 gms. 24.848 gms. or 1000 mils — V.S.

N Each mil of — thiosulphate represents 0.012692 gm. of

iodin.

N If in the above assay 39 mils of — sodium thiosulphate

were consumed, then

0.012692X39 = 0.495 gm.

0.495 + 10°

—=oo per cent. 0.5

LugoPs Solution. This is an aqueous solution of iodin and potassium iodid.

It is estimated for iodin in the same way as the foregoing. The potassium iodid acts merely as a solvent for free iodin, and does not enter into the reaction.

Ten or twelve grams of the solution is a convenient quantity to operate upon. Starch solution is the indicator.

Tincture of Iodin. An alcoholic solution of free iodin must be diluted with a solution of potassium iodid, before titration, in order to prevent the precipitation of iodin, which would result upon the addition of the aqueous standard solu- tion. The U.S. P. tincture of iodin contains potassium iodid and therefore may be titrated as it is.

Indirect lodometric Estimations. The titration methods previously described, in which iodin is used as a standard solution and the estimation of free iodin by means of sodium thiosulphate, are classed as direct iodometric methods. The following methods, in which the strength of the substance under analysis is determined by the quantity of iodin which

ANALYSIS BY OXIDATION AND REDUCTION 217

it liberates from an iodid, are known as indirect iodometric methods. Potassium iodid is added in excess * to an acidulated solution of the substance, and the liberated iodin estimated by means of standard thiosulphate. These methods are among the most accurate of all volumetric analyses, and take in a very large class of substances. Among the sub- stances which may be analyzed by this method are chlorin and bromin and all substances which readily liberate these elements: ferric salts, manganates, chromates, metallic perox- ids, and other substances from which oxygen can. be easily liberated.

Free Chlorin or Bromin. Free chlorin acts upon potas- sium iodid, liberating iodin, as per the equation

Thus it is seen that each atom of chlorin will liberate one atom of iodin, hence by determining the quantity of iodin by means of a standard thiosulphate solution the quantity of chlorin present is easily ascertained (126.92 gms. of 1 = 35.46 gms. of Cl). The same applies to free bromin, one atom of bromin (79.92) will liberate one atom of iodin

N (126.92) Br2 + 2KI = 2KBr + l2. 1000 mils of — sodium thio-

sulphate is equivalent to 12.692 gms. of iodin, and hence to 3.546 gms. of chlorin or 7.992 gms. of bromin.

N Thus i mil of — sodium thiosulphate is equivalent to

0.012692 gm. of iodin; 0.003546 gm. of chlorin; 0.007992 gm. of bromin.

Chlorin cannot be directly titrated with sodium thiosul-

*The iodid should be in sufficient excess to keep the liberated iodin in solution as KI.L

218 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

phate because, instead of the tetrathionate being formed as with iodin, sulphuric acid is produced; furthermore there is no readily observable end-point as there is with iodin.

Chlorin Water. This is an aqueous solution of chlorin, 01 = 35.46, containing at least 0.4 per cent of the gas.

The estimation of chlorin is effected in an indirect way, namely, by determining the quantity of iodin which it liberates from potassium iodid.

A definite quantity of chlorin will liberate a definite quan- tity of iodin from an iodid; these quantities are in exact proportion to their atomic weights, as the equation shows:

C12 + 2KI - 2KC1 + I2.

2)70.92 2)253.84

10)35.46 10)126.92

3.546 12.692 gms.

Thus it is seen that by estimating the liberated iodin the quantity of chlorin may be determined with accuracy.

Ten grams is a convenient quantity to operate upon. To this about half a gram of potassium iodid is added. A little starch solution is then introduced, and the titration is begun with decinormal sodium thiosulphate.

When the blue color of starch iodid has entirely disappeared the reaction is finished.

The reaction between iodin and sodium thiosulphate is illustrated by the following equation:

I2 + 2(Na2S2O3 + 5H2O) = 2NaI+ Na2S4O6 + ioH2O.

2)253.84 2)496.96

10)126.92 10)248.48 j_

12.692 gms. 24.848 gms. or 1000 mils — V.S.

N Thus we see that 1000 mils of — Na2S2O3 . 5H2O represent

ANALYSIS BY OXIDATION AND REDUCTION 219

12.692 gms. of iodin, which are equivalent to 3.546 gms. of chlorin.

Each mil therefore is equivalent to 0.003518 gm. of chlorin. This number is the factor which, when multiplied by the

N number of mils of — thiosutohate used, gives the weight in

grams of chlorin contained in the quantity of chlorin water acted upon.

Chlorinated Lime (Calx Chlorinata, Chlorid of Lime, Bleaching-powder). This substance was formerly supposed to be a compound of lime and chlorin, CaOCb, and hence the name chlorid of lime. It is now generally considered to be a mixture principally of calcium chlorid and calcium hypo- chlorite, CaCl2 + Ca(ClO)2 or Ca(OCl)Cl. The hypochlorite is the active constituent. This is a very unstable salt, and is readily decomposed even by carbonic acid. When treated with hydrochloric acid it gives off chlorin.

The value of chlorinated lime as a bleaching or disinfecting agent depends upon its available chlorin, that is, the chlorin which the hypochlorite yields when treated with an acid.

In estimating the available chlorin, the latter is liberated with acetic acid. This liberated gas, then, acting upon potas- sium iodid, sets free an equivalent amount of iodin. The quantity of iodin is then determined, and thus the amount of available chlorin found. 0.2 to 0.4 gm. are convenient quantities to operate upon.

Introduce into a stoppered weighing bottle between 3 and 4 gms. of chlorinated lime and weigh accurately. (In order to make the descriptions simpler we will assume that 3.5 gms. is the weight taken.) This is triturated thoroughly with 50 mils of water, and the mixture transferred to a graduated vessel, together with the rinsings, and made up to 1000 mils with water. This is thoroughly shaken, 100 mils of it (repre-

220 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

sen ting 0.35 gm. of the sample) are removed by means of a pipette and treated with i gm. of potassium iodid * and 5 mils of acetic acid, and into the resulting reddish-brown

N liquid the — sodium thiosulphate is delivered from a burette.

Towards the end of the titration, when the brownish color of the liquid is very faint, a few drops of starch solution are added and the titration continued until the bluish or greenish color produced by the starch has entirely disappeared. Not less than 30 mils of the volumetric solution sould be required to produce this result.

The reactions which take place in this process are illus- trated by the following equations :

Ca(OCl) Cl + 2HC1 = CaCl2 + H2O + C12 or

2)70.92 2)253.84

10)35.46 10)126.92

3.546 gms. 12.692 gms.

I2 + 2(Na2S2O3 + sH2O) = 2NaI + Na2S4O6 + ioH2O.

2)253-84 10)126.92 N|

12.692 gms. =1000 mils — thiosulphate V.S.

It is thus seen that i mil of the decinormal sodium thio- sulphate represents 0.012692 gm. of iodin, which in turn is equivalent to 0.003546 gm. of chlorin. Then

0.003546X30 = 0.1063 gm. 0.1063X100

0-35

= 30.37 per cent of available chlorin.

* In order to assume a sufficient excess of potassium iodid, take twice as much of it as of the bleaching-powder.

ANALYSIS BY OXIDATION AND REDUCTION 221

This is a very rapid method for estimating chlorin; but when calcium chlorate is present in the bleaching-powder (and it often is, through imperfect manufacture) the chlorin from it is recorded, as well as that from the hypochlorite, the chlorate being decomposed into chlorin, etc., by hydro- chloric acid (which is sometimes used). The chlorate, however, is of no value in bleaching; its chlorin is not available. Hence, unless the powder is known to be free from chlorate, the anal- ysis should be made by means of arsenous-acid solution, or by using acetic acid instead of hydrochloric, and thus avoid liberating chlorin from the chlorate which may be present.

The various bleaching preparations of the market which depend upon their available chlorin are all salts of hypochlorous acid (HCIO) or solutions of such salts.

Eau de Javelle (Javelle's Water) is a solution of potas- sium hypochlorite and potassium chlorid. A solution of mag- nesium hypochlorite is known in commerce as Ramsay's or Grouvelle's Bleaching Fluid. The solution known as Wilson's Bleaching Fluid contains aluminum hypochlorite.

Solution of Chlorinated Soda (Labarraque's Solution). This is an aqueous solution of several chlorin compounds of sodium, principally sodium chlorid and hypochlorite, con- taining at least 2.4 per cent by weight of available chlorin.

In this solution, as in chlorinated lime, it is the available chlorin which is estimated. The chlorin is first liberated with hydrochloric or acetic acid; ' this then liberates iodin from potassium iodid, and the free iodin is then determined by standard sodium thiosulphate.

Seven grams of chlorinated soda solution are mixed with 50 mils of water, 2 gms. of potassium iodid, and 10 mils of acetic acid are then added, together with a few drops of starch solution. Into this mixture the decinormal sodium thiosul- phate is delivered from a burette until the blue or greenish

222 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

N tint of the liquid is just discharged. Each mil of — thiosul-

phate used up represents 0.003546 gm. of available chlorin. The potassium iodid should always be added before the acetic acid, so that the chlorin has potassium iodid to act upon as it is liberated, and thus loss of chlorin is obviated.

Bromin Water, or any substance containing free bromin may be assayed in exactly the same manner as that described for chlorin water. Free chlorin must, however, be absent.

N Each mil of — thiosulphate solution represents 0.007992 gm.

of bromin.

Assay of Hydrogen Dioxid (H2O2 = 34). The iodometric method, which originated with Kingzett,* is based upon the fact that iodin is liberated from potassium iodid by hydrogen dioxid, in the presence of sulphuric acid, and that this libera- tion of iodin is in direct proportion to the available oxygen contained in the dioxid.

Then by' determining the amount of iodin liberated, the available oxygen is readily found.

H202 + H2S04 + 2KI - K2S04 + 2H2O + I2.

2)34_ 2)j^. 2)253.84

17 = 1 available O= 8 126.92

This shows that 126.92 gms. of iodin are liberated by 17 gms. of absolute dioxid, which are equivalent to 8 gms. of available oxygen.

N Thus 1000 mils of — sodium thiosulphate V.S., which

absorb and consequently represent 12.692 gms. of iodin, are equivalent to 1.7 gms. of H2O2 or 0.8 gm. of available oxygen.

* J. Chem. Soc., 1880, Vol. 37, p. 792.

ANALYSIS BY OXIDATION AND REDUCTION 223

N Each mil of this — V.S., then, represents 0.0017 gm- of

H2O2, and 0.0008 gm. of available oxygen.

The coefficients for weight of H2O2 and of oxygen, it is seen, are identical with those used in the permanganate process. Therefore the coefficient for volume is also the same in this method as in the other if i mil be taken for assay.

The process is carried out as follows: Take 2 or 3 mils of sulphuric acid, dilute it with about 30 mils of water, add an excess of potassium iodid (about i gm.), and then i mil of hydrogen dioxid. After the mixture has been allowed to stand five minutes, starch solution* is added, and the titra-

N tion with — sodium thiosulphate begun.

Note the number of mils required to discharge the blue color, and multiply this number: by 0.0017 gm. to find the quantity, by weight, of H2O2; by 0.0008 gm. to find the weight of available oxygen; by 0.57 mil to find the volume of avail- able oxygen.

If 18 mils are required, the solution is of 0.57X18=10.26 volume strength.

0.0017X18=0.0306 or 3.06 per cent H2O2. 0.0008X18=0.0144 or 1.44 per cent of oxygen.

With this method the author has always obtained satisfactory results. The lack of uniformity in the reaction, which is frequently reported, is doubtless due to the use of insufficient acid or to taking a too concentrated solution of the dioxid.

The best results are obtained if the solution is not more than two volumes strength.

* Starch solution may be omitted, as the decolorization of the iodin is distinctly seen if the beaker is placed upon a white surface.

224 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

The sulphuric acid used in this assay must be free from sulphurous acid, arsenous acid and nitric acid, and the potas- sium iodid must contain no iodate.

Distillation Methods. Manganates, chromates, metallic peroxids, and a great variety of substances containing oxygen, including antimonic oxid and arsenic pentoxid, will, when heated with concentrated hydrochloric, liberate an equivalent amount of chlorin. This is illustrated by the following equation :

The chlorin which is evolved, is passed into a solution of potassium iodid and liberates an equivalent of iodin, which latter substance is then estimated by titration with sodium thiosulphate solution. The quantity so found is therefore a measure of the original substance and of its oxygen content. The process may be carried out by means of the apparatus

FIG. 48.

devised by Bunsen, Fig. 48, or by that of Fresenius, Fig. 49, or Mohr, Fig. 50.

An accurately weighed quantity of the substance to be analyzed is introduced into the round-bottomed flask a, Fig. 48. The flask is then filled to about two-thirds its capacity with concentrated hydrochloric acid, and quickly connected by means of a short rubber tube with a long-bulbed delivery

ANALYSIS BY OXIDATION AND REDUCTION 225

tube, b, which is introduced into and extends to the bottom of an inverted bulbed retort, c. The larger bulb of the retort is filled to two- thirds of its capacity with a 10 per cent solution of potassium iodid. Heat is applied to the flask, and the chlorin distils over into the potassium iodid solution, which becomes brownish-red through liberation of iodin. The distillation is continued until about one-third of the acid fluid has passed over or until a peculiar cracking sound indicates the absorption of hot hydrochloric acid vapor.

The flask, together with its delivery tube, is then slowly removed, the heating, however, is continued until the tube

FIG. 49.

is entirely withdrawn, in order to prevent the iodid solution being drawn over into the flask. The retort is then shaken so that any traces of chlorin which may have escaped absorp- tion, are taken up and the contents of the retort poured into

226 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

a beaker; the retort and the delivery tube are then rinsed with water and the rinsings added to the fluid in the beaker, and titration with standard sodium thiosulphate begun immediately.

It is important that the quantity of potassium iodid is suffi- cient to keep the liberated iodin in solution, and the potas- sium iodid free from iodate, and the titration started with- out delay to avoid liberation of iodin through action of the air upon the strongly acid potassium iodid solution. When all the chlorin has passed over and hydrochloric acid gas begins to distil, the liquid in the retort is apt to be drawn back into the flask because of the great affinity which hydrochloric acid gas has for water, and the resultant condensation in the flask. This regurgitation may be avoided by introducing into the generating flask a small piece of magnesite, which slowly dissolves in the acid solution and so keeps up a constant flow of carbon dioxid, which by its pressure prevents back-flow of the fluid. The bulbs in the retort and delivery tube are also calculated to prevent this regurgitation.

The Fresenius apparatus is illustrated in Fig. 49. In this the potassium iodid solution is contained in two joined U-shaped tubes. The delivery tube from the distilling flask enters one of the U-tubes through a paraffin-soaked cork (which fits tightly), and terminates just above the potassium iodid solu- tion. In operation the U-tubes should be kept in ice water, and all the fittings should be air-tight. Paraffin-covered cork stoppers only should be used.

After all the chlorin has passed over or when about one-third of the acid has distilled over, the apparatus is allowed to stand for a few minutes, to permit all traces of chlorin to become absorbed; the application of a suction pump to the rear outlet tube will help to bring about this result.

Mohr's apparatus, shown in Fig. 50, is of very simple con- struction and easy to use.

ANALYSIS BY OXIDATION AND REDUCTION 227

The distilling flask is fitted with a paraffin-soaked cork, through which a delivery tube containing one bulb passes; this delivery tube again passes through a common cork which loosely fits a stout, large test tube, containing the potassium iodid solution. The delivery tube is drawn out to a fine point and reaches to near the bottom of the test tube. The latter is placed, when in operation, in a hydrometer jar containing cold water.

Estimation of Manganese Dioxid (MnO2 = 86.93). °-4 gm. of pulverized manganese dioxid is placed into the distil-

FIG. 50.

ling flask (a, Fig. 48 or Fig. 49) and the latter filled to two- thirds of its capacity with concentrated hydrochloric acid and

228 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

connected without delay with the vessel containing the potas- sium iodid solution. The flask is then gradually heated so that a steady current of chlorin passes over into the potassium iodid solution. When the evolution of chlorin gas begins to diminish, the heat is slowly raised to boiling, and continued at this point until about one-third of the acid liquid has dis- tilled over. The delivery tube is then removed and rinsed, as previously described, and the liberated iodin titrated by means

N of - - sodium thiosulphate solution, of which we will assume

60 mils were consumed.

The reactions are as follows:

(a) MnO2 +4HC1 = MnCl2 + 2H2O + C12.

86.93 70.92

(b) C

70.92 253.84

(c) I2 + 2Na2S2O3-5H2O = 2NaI + Na2S4O6 + ioH2O.

2)253.84 2)496.96 10)126.92 10)248.48

12.692 gms. 24.848 gms. = iooo mils — V.S.

N 0.012692 gm. =i mil — V.S.

These equations show that 494.96 gms. of sodium thiosul- phate will decolorize 253.84 gms. of iodin, which quantity is liberated by 70.92 gms. of chlorin, which is itself liberated from hydrochloric acid by 86.93 §ms- °f manganese dioxid.

Therefore i mil of a decinormal solution of sodium thio- sulphate (containing 24.848 gms. in 1000 mils) is equivalent to 0.012692 gm. of I; 0.003546 gm. of Cl; 0.0043465 gm. of MnO2; 0.0008 gm. of O (available).

The 60 mils of the thiosulphate solution used in this assay will therefore represent 0.0043465X60 = 0.2608 gm. of pure MnO2, or 65.2 per cent.

ANALYSIS BY OXIDATION AND REDUCTION 229

0.2608X100

= 65.2 per cent.

0.4

This is the method which should be used for the assay of native manganese dioxid. The freshly precipitated manganese dioxid of the Pharmacopoeia may be assayed by the more easily performed digestion method described on page 233.

Estimation of Chromic Acid and Chromates. Chromic anhydrid, chromium trioxid (CrO3 = ioo), when heated with concentrated hydrochloric acid, liberates chlorin as per the equation,

CrO3 + 6HC1 = CrCl3 + 3H2O + C13.

ioo 3X35-46

One hundred parts of CrO3 liberates 3X35.46 parts of Cl, hence one atomic weight of chlorin, 35.46 parts, represents 33.3

N

+ parts of CrO3. Or i mil of — sodium thiosulphate repre- sents 0.00333 gm. of CrO3.

Estimation of Potassium Bichromate (K2 0*207 = 294. 20). This salt, as explained in a previous chapter, has three atoms of oxygen available for oxidation. A molecule of this salt is therefore equivalent to six atoms of chlorin, and when boiled with hydrochloric acid will liberate six atoms of chlorin, as the equation shows.

K2Cr207 + i4HCl = 2KC1 + 2CrCl3 + 7H2O + 3C12. 294.2 6X35.46

Thus one atom of liberated chlorin will represent one-sixth of 294.2, which is 49.03 + parts of potassium dichromate. Then

N i mil of — sodium thiosulphate will represent 0.004903 gm.

of K2Cr2O7. In the same way all other chromates may be treated, but these compounds will, when treated with hydro- chloric acid, liberate chlorin at once and without the applica-

230 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

tion of much heat, hence some chlorin is apt to be lost before the distillation flask can be connected with the apparatus, and therefore it is more convenient 'to employ the digestion method later described.

The reaction in the case of neutral potassium chromate is as follows:

K2CrO4 +8HC1 = 2KC1 +4H2O + CrCl3 + C13. 194.2

Lead peroxid, PbO2; cobaltic oxid, Co2O3; nickel oxid, Ni2O3, as well as many other substances, may be assayed by this distillation method.

Estimation of Alkali lodids by the Distillation Method. This method is based upon the fact that metallic iodids, when

FIG. 51.

treated with ferric salts in acidulated solution, yield up all of their iodin. As shown in the equation

The iodin thus set free is distilled into a solution of potassium iodid and its quantity determined by titration with sodium thiosulphate in the usual manner.

ANALYSIS BY OXIDATION AND REDUCTION 231

For the reaction, ferric sulphate or ammonio-ferric alum may be used; the latter is, however, preferred because of its paler color. Ferric sulphate and ferric chlorid are so dark in color that the determination of the end-reaction is quite a difficult matter; furthermore, these salts frequently contain traces of nitrates which, if present, liberate chlorin from the chlorids or distil over and liberate iodin from the potassium iodid solution in the receiving vessel. Ferric chlorid is par- ticularly objectionable because of the tenacity with which it holds the last portions of iodin.

The distillation may be done in the Fresenius apparatus, Fig. 49, or better in that shown in Fig. 51. The latter consists of a ico-mils distilling flask (o), connected by means of a glass tube with a nitrogen flask (b), which contains a 10 per cent potassium iodid solution, and which is kept in a vessel of ice water when in use. The stoppers used are cork, well soaked in paraffin. The construction of the flask is particularly suit- able, because it presents a large surface to the vapor of iodin which distils over, and because the titration can be done directly in it, thus avoiding the necessity of transferring its contents to a beaker or other vessel.

The glass tube which conveys the iodin vapor must not be carried into the solution of potassium iodid and must not be drawn to a fine point. The reason for this is that the iodin condensing at the point would soon choke up the tube and prevent the further passage of iodin vapor. Any iodin which condenses in the tube is washed down into the potassium iodid solution by the steam during the distillation.

The Process. Into the flask (a) is introduced about 5 gms. of ammonio-ferric alum, 50 mils of water, 20 mils of diluted sulphuric acid (i no), and the iodid to be examined, accurately weighed. Take about 0.5 gm. The flask is then connected with the receiving vessel (b), which is about half filled with a

232 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

10 per cent potassium iodid solution, and after connections are made tight, heat is gradually applied to the distilling flask. After most of the iodin has passed over, the heat is raised to boiling, and continued at this temperature until about one- fourth of the liquid has passed over, and the solution in the distilling flask is no longer of a brown color.

When the receiving vessel has sufficiently cooled, it is dis- connected and its contents titrated with decinormal sodium thiosulphate, using starch as indicator. Before beginning titration, however, it is necessary to rinse the lower extremity of the tube and the stopper into the solution in the receiving vessel, in order that every trace of iodin be collected.

The calculation is then made as follows-

(a) Fe2(S04)3.(NH4)2S04

332.04

- K2SO4 + (NH4)2SO4 + 2FeSO4 +I2.

253-84

(b) I2 + 2Na2S2O3 + 5H2O = 2NaI + Na2S4O6 + ioH2O. 2)253-84

10)126.92 N

12.692 gms. = looo mils — V.S.

N 0.012692 gm.= i mil — V.S

Referring to the above equations we see that 253.84 gms. of iodin are liberated from 332.04 gms. of potassium iodid; thus 126.92 gms. of iodin represent 166.02 gms. of potassium iodid, and therefore i mil of decinormal sodium thiosulphate solution, representing 0.012692 gm. of iodin, will at the same time represent 0.016602 gm. of KI.

N If in the above assay 29 mils of — sodium thiosulphate

are consumed, we multiply the factor for KI =6.016602 gm. by 29, this gives 0.4814+ gm., the quantity of pure KI in the 0.5 gm. taken, which is about 96 per cent.

ANALYSIS BY OXIDATION AND REDUCTION 233

Digestion Methods. The distillation methods above de- scribed may be avoided in many cases and the more easily performed digestion process used. For instance, freshly pre- cipitated manganese dioxid, lead peroxid, chromic acid, chlor- ates, bromates, iodates, ferric salts, and a great many other substances, may be assayed by mere digestion with hydro- chloric acid at a slightly elevated temperature.

The digestion is performed in a strong glass bottle, provided with an accurately fitting ground-glass stopper which is tied down by means of wire or secured by a clamp. See Fig. 52.

Before using the bottle for this opera- tion it should be tested by securely tying down the stopper and immersing the bottle entirely in hot water to see if the stopper fits sufficiently tight. If it does not, bubbles of air will escape from inside, and the bottle is useless for the purpose intended. In that event the stopper must be reground into the neck of the bottle with a little very fine emery and water. The capacity of the bottle may vary from 50 to 150 mils.

The Process. The substance is accurately weighed and introduced into the bottle together with a small quantity of coarsely powdered glass or small pure flint pebbles (to prevent caking, especially in the case of insoluble powders). A suffi- cient excess of potassium iodid solution is then added, followed by some pure concentrated hydrochloric acid. The stopper is then quickly inserted, firmly secured by wire or a clamp, and the bottle placed in a water bath, and the water gradually heated to boiling; this temperature being continued until decomposition is complete, which is usually in about half an

FIG. 52.

234 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

hour. The bottle is then allowed to cool slowly and its con- tents emptied into a beaker. Then, after washing the bottle and adding the washings to the contents of the beaker, the liberated iodin is estimated by titration with sodium thio- sulphate.

The potassium iodid used in this process must be absolutely free from iodate.

N One mil of — sodium thiosulphate is equivalent to

KClOs 0.002042 gm.

NaClO3 . 0.001784 "

KBrO3 0.0027836 "

KIO3 0.003567 "

MnO2 0.0043465 "

PbO-2 0.011955 "

Cr03 0.00333 "

Estimation of Chlorates, Bromates and lodates. The

estimation of these salts is based upon the fact that in each case one equivalent of the acid or its monobasic salt liberates six equivalents of chlorin and consequently six equivalents of iodin when decomposed by the digestion method. This is illustrated by the equations:

(a) KC103 + 6HC1 = 3H2O + KC1 + C16 ;

(6) KBrO3 + 6HC1 = 3H2O + KBr + C16 ;

(c) KI03 + 6HC1 = 3H20 + KI + C16. and

(d) C16 + 6KI = 6KC1 + I6.

In the distillation process, however, bromates and iodates liberate only four equivalents of iodin, while bromous chlorid

^ ANALYSIS BY OXIDATION AND REDUCTION 235

and iodous chlorid remain in the retort, therefore in these cases the digestion is preferable to the distillation method.

If the bromate or iodate to be assayed contains any bromid or iodid, bromin or iodin respectively will be liberated upon the addition of the acid, according to the equations

(a) 5KBr + KBrO3 + 6HCl=6KCl+3H2O+Br6;

(b) 5 KI + KIO3 + 6HC1 = 6KC1 + 3H2O + IG ;

therefore the method is not applicable for the assay of such mixtures.

The presence of either of these salts may be ascertained by moistening a small quantity of the salt with dilute sulphuric acid, when if a yellow or brown coloration results, either a bromid or an iodid respectively is present.

Example. Estimation of Potassium Chlorate. 0.2 gm. of the salt is introduced into the digestion bottle of about 100 mils capacity, 10 mils of water added, and about 4 gms. of potas- sium iodid (or sufficient of its saturated solution). This is followed by 10 mils of concentrated hydrochloric acid, the stopper quickly inserted, firmly secured by wiring or a clamp, and the flask placed, stopper downward, in a water bath. The water is gradually raised to boiling and kept at this temperature for about half an hour. It is then allowed to cool slowly, and the contents of the bottle washed into a beaker and titrated with decinormal sodium thiosulphate,

N

using starch as indicator. The number of mils of — thio-

10

sulphate solution used, multiplied by 0.002042 gm^ gives the weight of pure KC1O3 present in the sample.

In the assay of bromates and iodates a smaller quantity of hydrochloric acid may be used, and a lower temperature, say 50° €„ is sufficient for decomposition.

236 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

Example. Estimation of Potassium Br ornate, 0.2 gm. of the salt is dissolved in 15 mils of water, 4 gms. of potassium iodid are added, followed by 4 mils of concentrated hydro- chloric acid. The bottle is securely closed, as in the fore- going assay, and heated for half an hour at 50° C. Then,

N after decomposition, titration with — sodium thiosulphate

solution is begun, each mil of which represents 0.2836 gm. of pure KBrOs.

If r^nr of one-sixth of the molecular weight of either salt

N be taken for assay, each mil of the — thiosulphate solution

used will indicate i per cent purity.

Estimation of Ferric Salts. When a ferric salt in an acidulated solution is digested with an excess of potassium iodid^the salt is reduced to the ferrous state, and iodin is set free.

One atoin of iodin is liberated for each atom of iron in the ferric state. The liberated iodin is then determined by sodium thiosulphate in the usual way. 126.92 gms. of iodin = 55.82 gms. of metallic iron.

The method is exemplified in the following assay: Ferric Chlorid (Fe2Cl6 or FeCl3). 0.5 of the dry salt accurately weighed is put into a zoo-mil glass-stoppered bottle, 15 mils of water are added to dissolve the salt, and then 3 mils of concentrated hydrochloric acid and 2 gms. of pure potassium iodid are introduced. The bottle is securely closed and its contents heated to 40° C. (104° F.) for half an hour. A higher temperature must be avoided, otherwise iodin will be volatilized. The liquid is then cooled and the liberated

N iodin titrated with — sodium thiosulphate until the color is

just discharged. Starch may used as indicator-

ANALYSIS BY OXIDATION AND REDUCTION 237

N Each mil of — thiosulphate solution used represents 0.005582

gm. of metallic iron or 0.01622 gm. of pure ferric chlorid. The following equations illustrate the reactions :

Fe2Cl6 + 2KI = 2FeCl2 4- 2KC1 + I2. 2)324.4 2)253.84

162.2 126.92

Then

I2 + 2Na2S2O3.5H2O =2NaI + Na2S4O6 + ioH2O.

2)253-84 10)126.92 N

12.692 gms. = 1000 mils — V.S.

The assay of other ferric salts is practically the same as that described above.

N Each mil. of — sodium thiosulphate is equivalent to

Ferrum, Fe 0.005582 gm.

Ferric ammonium sulphate, Fe2 (SO4) 3. (NH4) 2SO4 o . 0266 "

Ferric chlorid, Fe2Cl6 0.01622 "

" (cryst), Fe2Cl6 + i2H20 0.02702 "

" nitrate, Fe2(NO3)6 0.024185 "

" oxid, Fe2O3 0.007982 "

" sulphate, Fe2(SO4) 3 0.019992 "

In the case of the scale salts of iron, which are mostly of indefinite and variable composition, it is the quantity of metallic iron present which is determined.

The solutions of ferric salts are estimated in the same man- ner. It is the rule to take i to 2 gms. of the solution for assay.

Assay of Chromic Acid, Chromates and Bichromates. Dissolve about i gm. of the substances accurately weighed in a stoppered weighing bottle, in sufficient distilled water to measure 100 mils.

To 15 mils of this solution add 3 mils of hydrochloric acid

238 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

and 2 gms. of potassium iodid in a glass-stoppered container, shake the mixture, allow it to stand for five minutes, th°n dilute with distilled water to measure 100 mils, and titrate with tenth-normal sodium thiosulphate V.S., using starch T.S. as indicator.

Each mil of the thiosulphate V.S. used corresponds to °-°°333 g111- of CrOs.

See equations on page 229.

Assay of Arsenic Oxid (As2Os) and Arsenates. Also anti- monic compounds. Dissolve about 0.5 gm. of the substance accurately weighed in 25 mils of distilled water, heat the solu- tion to 80° C. and add 10 mils of hydrochloric and 3 gms. of potassium iodid. Allow the mixture to stand for fifteen minutes at 80° C., then cool and titrate with tenth-normal sodium thiosulphate V.S., using starch as indicator. The solution must be decidedly acid in reaction. The function of the acid is expressed by the equation

KI+HC1 =

The HI then reacts with the As2C>5, reducing it to As2Oa and liberating iodin, as the following equation shows:

As2O5 + 4HI = As2O3 + 2H2O + 2l2.

Each mil of the tenth-normal thiosulphate V.S. corre- sponds to 0.005748 gm. of As2O5 or 0.0092985 gm. of Na2HAsO4 (sodium arsenate.)

Assay of Copper. This method depends upon the reaction between copper salt and potassium iodid, in which iodin is quantitatively liberated.

2CuSO4 +4 KI = Cu2I2 + 2K2SO4 + I2. 2X159.64 2X126.92

The solution of the copper sulphate should be weakly acid. Acetic acid is usually used. All iodin liberating substances

ANALYSIS BY OXIDATION AND REDUCTION 239

must be absent, as well as all metals which react with iodin, as Pb, As, Sb, Bi, etc.

The Process. About i gm. of the copper salt, accurately weighed, is dissolved in 50 mils of distilled water; to this is added 4 mils of acetic acid and 3 gms. of potassium iodid. The liberated iodin is then titrated with tenth-normal sodium thiosulphate V.S., using starch as indicator. Each mil of the thiosulphate V.S. corresponds to 0.006357 gm. metallic copper or 0.015964 gm. of CuSO-i.

If copper nitrate is to be assayed, the NO3 ion must be taken up with ammonia-water, which is added until a clear blue solution results, this is boiled for about one minute, and then acidified with acetic acid, or the solution is evaporated with sulphuric acid until the nitric acid is expelled.

Assay of Mercuric Salt. The mercury is precipitated out of an alkaline solution by formaldehyde, and then dissolved in an excess of standard iodin solution in the presence of potas- sium iodid. The excess of iodin is then titrated with standard 'sodium thiosulphate V.S.

Hg + I2 = HgI2.

Mercuric Chlorid may be reduced to mercurous chlorid by means of hydrogen dioxid in the presence of tartaric and hydro- chloric acids. Frequent additions of hydrogen dioxid are made and the mixture warmed and filtered. The precipitate is washed and treated as described under assay of mercurous chlorid.

Assay of Mercurous Chlorid or Iodid. The mercurous

N chlorid or iodid is treated with — iodin V.S. and potassium

iodid until solution is complete. The reaction is as follows: 2HgCl + 6KI + 12 = 2HgK2I4 + 2KC1.

240 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

The excess of iodin is then titrated with sodium thiosul- phate, starch being used as indicator.

The Process. One gm. of the salt, accurately weighed, is heated with 10 mils of distilled water, 50 mils of tenth-normal iodin V.S. and 5 gms. of potassium iodid. The flask is stoppered, and the mixture allowed to stand, with occasional agitation, until complete solution has taken place. The excess of iodin is then titrated as above described.

N Each mil of — iodin V.S. used corresponds to 0.023606

gm. of HgCl or 0.032752 gm. of Hgl.

Reduction Methods Involving the Use of Standard Arsenous Acid Solution (Chlorometry)

As previously described, arsenous oxid when brought in contact with iodin in an alkaline solution, results in an oxida- tion of the former to arsenic oxid, and a conversion of the iodin to hydriodic acid, as shown in the equation

As2O3 + 2H2O + 14 = As2O 5 + 4HI.

Advantage is taken of this reaction for the estimation, not only of arsenous and antimonous compounds, but also of iodin and the other halogens, chlorin and bromin, as well as of all those bodies which, when heated with hydrochloric acid, evolve chlorin, as for instance the peroxids.

The reaction with chlorin is as follows:

This reaction is really an oxidation, so far as the formation of arsenic oxid (As2O->) is concerned, but there is no accom- panying reduction. The conversion of the halogen to an haloid acid is not strictly a reduction in the accepted sense

ANALYSIS BY OXIDATION AND REDUCTION 241

of the word. Nevertheless, for obvious reasons, we speak of analyses done by means of arsenous acid as reduction methods.

The chief value of this method is found in the estimation of free chlorin, as in chlorin water, and the available chlorin existing in hypochlorites or that evolved from hydrochloric acid by heating with peroxids. Hence the designation " chlor- ometry."

In carrying out this method free alkali must be present to combine with the haloid acid which is formed. The alkali must be in the form of bicarbonate. Normal carbonates or hydroxids are not suitable (see page 194).

The solutions required are:

Decinormal iodin (see page 188);

Decinormal arsenous acid;

Starch solution (see page 192); or iodized starch test paper.

Preparation of Decinormal —Q Arsenous Acid (As2O3 =

N 197.92; — V.S. = 4.948 gms. in i liter). 4.948 gms. of the

purest sublimed arsenous anhydrid (As2Os) are dissolved in a minimum of concentrated sodium hydroxid solution. When complete solution is effected add 100 mils of distilled water and neutralize the excess of alkali with diluted sulphuric acid, using phenolphthalein as indicator. Add to this 20 gms. of pure sodium bicarbonate dissolved in 500 mils of distilled water, and finally dilute to 1000 mils at standard temperature. It is then standardized with decinormal iodin, using starch as indi- cator. Decinormal arsenous acid solution should correspond, volume for volume, with decinormal iodin solution.

If this solution is made from pure arsenous acid, it will hold its titer for years, but if any sulphur is present there will be an absorption of oxygen from the air and a consequent

242 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

oxidation to arsenic oxid. If the presence of sulphur is sus- pected, the solution should be tested with silver nitrate, when its presence will be indicated by the formation of a reddish precipitate.

Iodized Starch Test Paper. A portion of starch solution is mixed with a few drops of potassium iodid solution and in this are soaked strips of pure white filtering paper. This test paper is used in the damp state; it is then far more sensitive.

Estimation of Free Halogens. The estimation of chlorin, bromin or iodin by the chlorometric method depends, as before stated, upon their power of oxidizing arsenous acid. When a free halogen is brought in contact in alkaline solution with arsenous acid, the latter is oxidized to arsenic acid, while the halogen is transformed into a haloid acid, as per equations

CU] [4HC1

Br4 +2H2O+As2O3-As2O5+ 4HBr. • I4 J I 4HI

The estimation may be carried out in two ways: ist, by direct titration with a standard arsenous oxid solution, using iodized starch test paper as indicator; 2d, by residual titra- tion, an excess of the standard arsenous oxid being taken, and then retitrating with standard iodin solution, using starch as indicator. The residual titration method need not be employed for free iodin, as this can be titrated direct with the arsenous oxid solution, using starch as indicator. Furthermore, iodin need not be brought into solution to be titrated by this method.

The estimation of free halogens by the direct chlorometric method is as follows:

An accurately weighed quantity of substance made alka- line by the addition of sodium bicarbonate is titrated with decinormal arsenous acid solution, and from time to time

ANALYSIS BY OXIDATION AND REDUCTION 243

during the titration a drop of the solution is removed on the end of a pointed glass rod and brought in contact with a piece of iodized starch test paper. So long as free chlorin or bromin is present the liquid will cause a blue stain on the test paper, but when the halogen is all taken up no blue color is produced.

If the exact point is overstepped the residual method must be used. A little additional excess of the arsenous acid solu- tion may be added, together with a few drops of starch solu- tion, and the excess then titrated by means of decinormal iodin solution until the blue color is produced. The volume of decinormal iodin solution so used, deducted from the total volume of arsenous acid solution taken, gives the exact quan- tity which was oxidized by the halogen, and from this the percentage of chlorin or bromin may be calculated.

Example i. Estimation of Chlorin in Chlorin Water. Twenty mils of chlorin water (sp.gr. i.o) titrated by the direct method required 22 mils of decinormal arsenous acid solution before the iodized starch test paper indicated the completion of the reaction.

By referring to the equation we see that each mil of the arsenous acid solution represents 0.003546 gm. of chlorin. Therefore, if 22 mils were used, the 20 mils of chlorin water must have contained 22X0.003546 gm. of chlorin, which is 0.078012.

The 20 mils of chlorin water (sp.gr. i.o) weigh 20 gms. Hence

0.078012X100

— ^— =39.0+ percent.

Example 2. The 20 mils of chlorin water weighing 20 gms. were treated with 26 mils of the arsenous acid solution, starch solution was then added, and the excess of arsenous acid solution titrated by means of decinormal iodin. Four

244 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

mils of the latter were required, then 4 from 26 mils leaves

N 22 mils, the quantity of the - - arsenous acid solution which

reacted with the chlorin. The calculation is the same as in Example i.

The reaction is as follows:

C14 + 2H2O + As2O3 = 4HC1 + As2O5. . 4)141.84 4)197.92

10)35.46 10)49.48 N

3.546 gms.= 4-948 gms. = 1000 mils — V.S.

10

0.003546 gms. = i mil — V.S.

10

Estimation of Available Chlorin in Bleaching Powder.

Three and one-half gms. of the bleaching powder (chlorin- ated lime) are triturated thoroughly with 50 mils of water, and the mixture transferred to a graduated vessel, together with the rinsings, and made up to 1000 mils with water. This is thoroughly shaken. 100 mils of it (representing 0.35 gm. of the sample) is removed by means of a pipette and titrated with decinormal arsenous acid solution, as described in the foregoing assay, using either the iodized starch test

N paper as indicator or retitrating the excess of — arsenous

N . acid solution added by means of — iodin solution.

N Fach mil of — As2O3 V.S. represents 0.003546 gm. of

available chlorin.

Ca(OCl)Cl + As2O3 = As2O5 + 2CaCl2.

4)141-84 4)197-92

10)35.46 10)49.48 N

3.546 gms.= 4.948 gms. or looomils — V.S.

10

As seen by referring to the above equation this process determines the value of the chlorinated lime by measuring

ANALYSIS BY OXIDATION AND REDUCTION 245

the amount of arsenous acid which the oxygen present in the active constituent (Ca(OCl)Cl) is capable of oxidizing. In the formula of this compound there are two atoms of chlorin and one atom of oxygen. Therefore the quantity of bleaching powder which yields 35.46 parts of available chlorin will also supply 8 parts of oxygen; this may therefore be taken as the measure of the chlorin. The same method may be employed for the assay of all other solutions containing available chlorin.

Assay of Manganese Dioxid (Chlorometric) . The chloro- metric assay of manganese dioxid, as well as that of all other bodies which liberate chlorin when heated with hydrochloric acid, may be made in similar manner to that described for the iodometric assays of these substances, the same apparatus, etc., being used.

N

The liberated chlorin is, however, titrated with — arsenous

10

acid solution.

The chlorin may be distilled into a solution of sodium

N

carbonate, and this solution is then titrated with — arsenous

10

acid or the chlorin may be distilled directly into a measured

N

volume of — arsenous acid solution and the latter then titrated 10

N with — iodin solution, using starch as indicator, the difference

between the volume of iodin solution used and that of the arsenous acid solution taken is the measure of the latter which reacted with the chlorin.

It is a good plan in each case to divide the solution into two or three equal parts and to titrate each separately.

246 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

Reduction Methods Involving the Use of Stannous Chlorid

Stannous chlorid (SnCl2) is a very powerful reducing agent. Its action in this respect depends npon its affinity for chlorin which it readily abstracts from most other chlorids. In its action upon mercuric chlorid a portion of the latter is always reduced to the metallic state.

This reducing action of stannous chlorid is utilized in certain volumetric processes, especially in the estimation of iron. In this case it possesses an advantage over perman- ganate, in that the iron must be in the ferric state, in which condition it is most usually found, while if permanganate is used, a preliminary reduction to the ferrous state is necessary before titrating. The great disadvantage, however, is in the fact that even short contact with air will quickly oxidize it, and thus spoil its titer. In consequence of this it must be frequently tested, and can be used only in the form of empirical solutions.

It is particularly useful in the titration of ferric salts, which salts can be accurately estimated by direct titration with it, the end-point being recognized by the disappearance of the yellow color of the ferric solution. These salts may also be estimated residually by adding an excess of stannous chlorid solution of known strength and retitrating the excess by means of standard iodin, using starch as indicator.

The reactions are:

Fe2Cl6 + SnClo = 2FeCl2 + SnCl4 and

SnCl2 + 2HC1 + 12 = SnCl4 + 2ffl.

The Estimation of Iron by Means of Stannous Chlorid Solutions may be accurately affected by the following pro- cedure, as suggested by Fresenius. The solutions necessary are:

ANALYSIS BY OXIDATION AND REDUCTION 247

(a) A solution of ferric chlorid containing 10 gms. of pure iron in a liter.

This is made by dissolving 10.04 gms. of thin annealed binding wire (which contains 99.6 per cent of pure iron) in a sufficient quantity of pure hydrochloric acid. A small quantity of potassium chlorate is then added to effect complete oxidation of the iron, and the excess of chlorin expelled by boiling. This solution is then cooled and diluted to i liter.

(b) A solution of stannous chlo- rid made by dissolving about 10 gms. of pure tin in 200 mils of strong, pure hydrochloric acid. This may be done by heating the tin in small pieces with the acid in a flask, and introducing a few pieces of platinum foil to excite galvanic action. The solution so

obtained is diluted to about one liter with distilled water and should be preserved in a bottle, such as shown in Fig. 53, to which air can only gain access through a strongly alkaline solution of pyrogallic acid. When so kept the strength of the solution can be preserved for several weeks.

(c) A solution of iodin in potassium iodid. This may be approximately or exactly decinormal.

The procedure is as follows:

i st. The relation between the tin solution and the iodin solution is found.

2d. The relation between the tin solution and the iron solution is determined.

FIG. 53.

248 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

3d. The assay.

The relation between the tin solution and the iodin solution is found as follows :

Two mils of the tin solution are put into a beaker, a little starch solution added, and the iodin solution then delivered in from a burette until the blue color occurs. If 4 mils are used, then each 2 mils of iodin solution represents i mil of tin solution.

The relation between the tin solution and the iron solution is found as follows:

Fifty mils of the iron solution (representing 0.5 gm. of iron, are put into a small flask together with a little hydrochloric acid and heated to gentle boiling. The tin solution is then delivered from a burette until the yellow color of the iron solution is nearly discharged. It is then added continuously, drop by drop, until the color is entirely gone. Assuming that 35 mils were re- quired, then each 35 mils of tin solution are equivalent to 0.5 gm. of pure iron. If the end-point is not clearly recognized and an excess of the tin solution was added, the solution should be quickly cooled, a few drops of starch solution added, and the excess estimated by titrating with the iodin solution, each mil of which represents 0.5 mil of the tin solution. The excess so found, de- ducted from the total quantity of tin solution added, gives the quantity of the latter, which corresponds to 0.5 gm. of iron.

Having determined these data, the analyst can readily estimate any unknown quantity of iron in solution in the ferric state.

If the iron is partly or wholly in the ferrous state it may be oxidized by adding some potassium chlorate and boiling to expel excess of chlorin.

The Assay. A solution of iron taken for analysis, required 24 mils of the tin solution. The quantity of iron present is calculated by proportion as follows :

35 mils : 0.5 gm. :: 24 mil : x\ # = 0.34 gm.

ANALYSIS BY OXIDATION AND REDUCTION 249

To secure accurate results the iron solution assayed must be fairly concentrated, because then the end-reaction is more readily seen, and also because the greater the dilution the larger the amount of tin solution will be required. It is good policy to use very little excess of the tin solution, so that only a very small quantity of iodin solution is required.

Estimation of Mercuric Salts (Laborde). This depends upon the fact that stannous chlorid solution added to a solution of a mercuric salt reduces the latter first to mercurous chlorid (calomel) and finally the calomel to metallic mercury. The reduction to calomel results in the formation of a white pre- cipitate, and when the mercuric salt is completely reduced the stannous chlorid acts upon and reduces the calomel to metallic mercury, which results in the production of a char- acteristic brownish color.

The reactions are as follows:

SnCl2 + 2HgCl2 = SnCl 4 + 2HgCl ; SnCl2 + 2HgCl = SnCl4 + Hg2.

According to Laborde the tin solution is made by dissolving 8 gms. of pure tinfoil by means of heat in 100 mils of pure hydrochloric acid, and diluting to 2 liters.

This tin solution is checked against a solution of mercuric chlorid containing 10 gms. per liter. To counteract the hin- dering effect of the hydrochloric acid the solution under anal- ysis, containing o.i gm. of mercuric chlorid, is mixed with 5 mils of a solution containing 100 gms. of ammonium acetate and zoo mils of acetic acid to the liter. The acetic acid pro- motes the disappearance of the brown color which occurs at the point where the tin solution is in excess, but before reduc- tion is complete. The titration with the tin solution is con- tinued until a permanent brown color occurs.

250 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

If the brown color is too dark from overstepping of the end-point, the addition of i mil of the mercuric chlorid solution will render the solution white again, and the titration can then be carried further.

This method, which is convenient, rapid, and very accurate, can be employed in many cases. If the mercuric solution contains any free mineral acids, the latter must be neutralized with ammonia (in the presence of ammonium acetate, to prevent formation of ammoniated mercury).

The presence of alkali and alkali earth salts or most salts of other metals (except iron, gold, and platinum) do not in the least interfere with the accuracy of the results. The same is true of organic acids, either free or in combination with alkali.

PART II

CHAPTER XI ESTIMATION OF ALKALOIDS

IN making alkaloidal assays of drugs it has long been the custom to evaporate the final ethereal or chloroformic extract, and to weigh the residue as alkaloid. This residue seldom, if ever, consists of the pure alkaloid, and the impurities — i.e., non-alkaloidal matter — is variable in amount and difficult to entirely remove, consequently gravimetric results are in many cases very wide of the truth, and hence unreliable.

The volumetric methods are in most cases much more satisfactory.

While the results of the titration of the total alkaloids of drugs cannot be called absolutely accurate, nevertheless experi- ence has shown that they are nearer the truth than those obtained by the gravimetric method.

In estimating an alkaloid by titration it is essential to know the formula and molecular weight of the alkaloid, as well as the equivalent of acid with which it will combine.

The quantity of alkaloid present is easily calculated when we know that a molecular weight of a monobasic acid or half a molecular weight of a dibasic acid will combine with and neutralize a molecular weight of an alkaloid, provided the alkaloid is a monacid base. If the alkaloid is a diacid base, one molecular weight will combine with two molecules of a monobasic acid or one molecular weight of a dibasic acid.

251

252 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

Sparteine and emetine (?) are diacid alkaloids; most of the others are monacid bases. Examples. Monacid alkaloids:

C2oH24N202 + HC1 = C20H24N202 . HC1 ;

Quinine

(C17H19N03)2+H2S04=(C17H19N03)2.H2S04;

Morphine

(C2iH22N2O2) 2 + H2SO4 = (C2iH22N2O2) 2 . H2SO4;

Strychnine

Diacid alkaloids:

Sparteine

C3oH4oN205 + H2S04 = C3oH4oN205 - H2SO4.

Emetine (Kunz)

The quantity of alkaloid present in the substance is easily calculated, as illustrated by this equation:

C20H24N2O2 + HC1 = C20H24N2O2 . HC1.

Quinine

N 324.2 gms. 36.46 gms. = 1000 mils — V.S.;

N 32.42 " 3.646 " =1000 mils — V.S.

N Thus i mil of — V.S. = 0.03242 gm. of quinine.

C15H26N2 + 2HC1 = C15H26N2(HC1)2.

Sparteine

2)234.2 2)72.92 N

117.1 gms. 36.46 gms. = 1000 mils — V.S.;

11.71 " 3.646 " = 1000 mils — V.S.

10

N One mil of — V.S. hence = o.oii7i gm. of sparteine.

N Thus icoo mils of — hydrochloric acid will combine with

ESTIMATION OF ALKALOIDS 253

•fa of the molecular weight in grams of a monacid alkaloid, or TjV of the molecular weight of a diacid alkaloid.

In the case of drugs where two or more alkaloids are present, accurate results can only be obtained by determining how much of each alkaloid is present by a separate assay. But often it is assumed that the alkaloids are present in equil quantities, and the mean of their molecular weights is taken as the basis for the calculation.

It must be borne in mind, however, that in titrating alka- loids the greatest care must be exercised and all precautions closely observed in order to attain any degree of accuracy. The volumetric solutions must be prepared with the greatest care and must be absolutely accurate. The eye must be trained (as it can only be through practice), to distinguish the end-colors of the indicators employed, a matter of some difficulty. Furthermore, all measuring instruments used must be accurate, or they should be carefully calibrated in order to find the necessary factor for correction.

The Volumetric Solutions usually employed in titrating

N N N N N

alkaloids are — , — , — , — , and — . The weaker solutions 10 20 25 50 loo

give more accurate results, as will be understood if we remem-

N N

ber that a — is loX stronger than a — and looX stronger

N N

than — V.S. Then if — solution be used, one drop may

100 I

overstep the neutral point, while if the same solution were

N treated with — solution 5 drops would be required to neutralize,

N which is equivalent to using one half a drop of — solution. A

N — solution will of course be capable of even more delicate

254 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

N work. In the case just mentioned 45 drops of the - - may

exactly neutralize the solution, hence less than half a drop

N N

of : - and between 4 and 5 drops of \ - are represented.

Therefore in all delicate alkaloidal titrations weak standard solutions should be used — in fact, in all titrations where great accuracy is required.

// the alkaloid be from a recent extraction and is in the form of a free alkaloid, it is dissolved in a measured volume

N

of — acid solution and the excess of acid solution then deter- 10

N

mined by residual titration with -- alkali solution.

100

N ; In this the — sulphuric acid solution is preferred, except

N in the case of quinine or cinchonine, in which — hydrochloric

acid gives better results.

The process in detail is as follows: Place 2 gms. of the

N alkaloid into a beaker, add 75 mils of — sulphuric acid solu-

tion, and warm1 on a water-bath until the alkaloid is com- pletely dissolved. The solution is then allowed to cool and diluted to 100 mils.

Ten mils of the solution (containing 0.2 gm. of the alkaloid

N and 7.5 mils of — sulphuric acid solution) are removed by

N means of a pipette and retitrated with — potassium hydroxid

fabJI

N

solution. One-tenth of the quantity of the — alkali used

100

N is deducted from the 7.5 mils of the — acid solution, and the

ESTIMATION OF ALKALOIDS 255

remainder is the quantity of the latter which combined with and hence represents the alkaloid present. This, if multiplied by the factor, gives the weight of the alkaloid.

Either haematoxylin solution or Brazil-wood T.S. may be employed as the indicator.

// the alkaloid is soluble in alcohol, as are quinine and ( odeine, it may be treated as follows :

Place 2 gms. of the alkaloid in a graduated cylinder, dis- solve in alcohol, and dilute the solution up to 100 mils with alcohol. Remove 10 mils of this solution (containing 0.2 gm. of the alkaloid) and place in a beaker, add the indicator and run in the decinormal acid solution to slight excess. Rotate the beaker several times, let stand for a few minutes, wash down the sides of the beaker with distilled water, using about

N 40 mils, and retitrate the excess of acid with — potassium

hydroxid solution, until end-color is given by the indicator.

N Deduct one-tenth of the quantity of the — solution used

N from the quantity of — acid solution added, and the remainder

is the quantity of the latter, which combined with and hence represents the alkaloid present.

The indicators best suited for most alkaloids are Hcema- toxylin, Brazil-wood, Cochineal, lodeosin and Litmus.

The color changes produced by the principal indicators used in alkaloidal titrations are as follows:

256 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

Acid.

Alkali.

Haematoxylin

Yellow

Blue

Brazil- wood

«

Purplish-red

Cochineal

Ye llo wish-red

Purplish

Litmus

Red

Blue

lodeosin

Yellow

Rose-red

Lacmoid

Red

Blue

Fluorescein

Green fluorescence

No fluoresc., yellowish

Congo red

Blue

Red

Methyl-orange

Red

Straw-yellow

Phenolphthalein

Colorless

Red

TABLE SHOWING THE FACTOR FOR VARIOUS ALKALOIDS

WHEN TITRATING WITH — ACID OR ALKALI

10

Name.

Formula.

VIolecular Weight.

Factor.

Aconitinc

CwHi-NO,,

64^.48

0.064548

Atropine

C,,H,,NO,

280 24

o 028024

Brucine

C,,H2flN2O4

^04.28

0.039428

Cephaeline

C14H19NQ2

2T.T..2O

0.023320

Cinchona alkaloids (combined)

0.03060

Cinchonine

C19H22N,O

204-24

0.029424

Cinchonidine

CiaH,,N,O

2Q4 24

o 020424

Cocaine

C17H21NO4

"?O?.22

0.030322

Codeine

C,«H,,NO.

2QO 22

O O2QQ22

Coniine

C8H17N

127 18

O.OI27I8

Emetine -1

C30H44N2O4 (Glenard) C.nlLnNiOc (Kunz)

496.32

co8 %4.

O.0248l8

o 02^417

Hydrastine

C15H21N02 (U. S. P.) C,,H,,NO.

247.22

•2g-2 22

O.O24722

o 038322

Hvoscine

C,7H21NO4

•JQ7 . l8

0.030318

Hvoscy am ine

Ci,H,,NOo

280 24

o 028924

Ipecac alkaloids (combined)

o 024021

Morphine

C,,H,0NO»

28? 2O

o 028^2

Nicotine

C5H7N

81 06

o 008106

Physostigmine

CnH,,N,O,

27C 24

1.027^24

Pilocarpine

C,,Hi«N,O,

208 18

o 020818

Quinine

C2ftH,,N2O,

^24.26

0.032426

Sparteine

C,5H26N2

234.28

0.011714

Strychnine

C,,H,2N,O,

334.24

0.033424

ESTIMATION OF ALKALOIDS 257

GORDIN'S MODIFIED ALKALIMETRIC METHOD, USING PHENOL-

PHTHALEIN AS INDICATOR

The alkaloidal residue obtained by any of the extraction

N

methods in use is dissolved in a measured excess of — hydro- chloric acid solution. Wagner's or Mayer's reagent is then added, a little at a time, with frequent shaking, until the alkaloids are completely precipitated. The mixture is then diluted with water to 100 mils and shaken until the double salt of the alkaloid and reagent completely separate. When allowed to stand a few minutes the supernatant liquid should be clear, and if Wagner's reagent has been used, this will be of a dark-red color. The liquid is now filtered, and 50 mils of the filtrate (representing one-half of the alkaloid) is treated with a 10 per cent solution of sodium thiosulphate added, drop by drop, until the color of the free iodin dis- appears. This discolorization is not needed if Mayer's reagent has been used. A few drops of the indicator phenolphthalein are now introduced, and the excess of acid estimated by reti-

N tration with — potassium hydroxid solution. This, deducted

from one-half of the volume of the standard acid solution

N employed, indicates the number of mils of — hydrochloric

acid solution, which combined with the alkaloid in 50 mils of the solution. This number, multiplied by two and then by the factor for the alkaloid present, gives the total quantity of alkaloid.

Example. An alkaloidal residue consisting of morphine

N was dissolved in 30 mils of — HC1 V.S. Then Wagner's reagent

was added in excess as described, and the mixture made up

258 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

to 100 mils with water. Fifty mils of this were filtered off,

N

decolorized as directed, and titrated with — KOH V.S. Ten

20

mils were required, which corresponds to 20 mils for the entire

N quantity. Then 20 mils deducted from the 30 mils of -

HC1 V.S. added, leaves 10 mils, the quantity required to neu-

N tralize the morphine. The — factor for morphine (0.0137

gm., Gordin), multiplied by 10=0.137 gm., the quantity oi alkaloid in the residue examined.

N The — factor for morphine here given is somewhat lower

than the theoretical equivalent. It was ascertained by exper- ment with a sample of the anhydrous alkaloid.

The following factors (by Gordin) were obtained by com- paring their molecular weights with that of morphine, which factor was determined by experiment:

N Morphine 0.0137 £m- = i mil — HC1 V.S.

Hydrastine 0.0184 " =

Strychnine 0.0160 " = "

Caffeine, cryst 0.0102 " =

Cocaine 0.0146 " = "

Atropine 0.0139 " = "

CHAPTER XII

VOLUMETRIC ASSAYING OF VEGETABLE DRUGS AND THEIR PREPARATIONS

EXTRACTION OF THE ALKALOIDS

Selection of the Sample. Care must be taken to secure a fairly representative sample.

If the drug is in small pieces or consists of seeds or leaves, mix it well, take a portion, pulverize it, and of the powder take a sufficient quantity for the assay. If the drug be in large lumps, which vary in quality, select a few representative lumps and cut from each a fairly representative section; pul- verize these, mix well, and weigh off a sufficient quantity for the assay.

If drying is necessary, the loss of weight in drying must be made note of.

The Exhaustion of the Drug is usually effected by macer- ation in a suitable menstruum, although percolation, boiling, and hot repercolation must be employed in some cases.

The Choice of Solvent depends upon the nature of the drug. Water dissolves, besides the alkaloids, so much inert matter that the subsequent steps in the assay are liable to be interfered with. Alcohol dissolves too much of the resinous matter, and besides does not penetrate the drug very well. Acidulated water has been much used, but chloroform and ether, separately and in various combinations, are now most generally employed for exhausting drugs in conjunction with

259

260 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

alcohol and ammonia. Petroleum benzin has of late been recommended.

Prollius' fluid, or some modification of it, is very satis- factory.

Prollius' Fluid consists of ether 325 mils, alcohol 25 mils, and concentrated ammonia-water 10 mils.

Modified Prollius' Fluid consists of ether 250 mils, chloro- form 80 to 100 mils, alcohol 25 mils, concentrated ammonia- water 10 mils.

The Separation of the Alkaloids and the Use of Immiscible Solvents. " The Shaking-Out " Process. Most alkaloids are insoluble in water but are readily soluble in certain " volatile solvents " such as alcohol, ether, chloroform, amyl alcohol, benzene, petroleum benzin, acetic ether, etc., but not in carbon tetrachlorid. Those alkaloidal solvents which do not mix readily with water are known as " the immiscible sol- vents." Salts of alkaloids on the other hand are as a rule soluble in water, but practically insoluble in the immiscible solvents. This property of the alkaloids makes it possible to separate them from their natural sources.

The alkaloid in solution in an immiscible solvent may be separated by " shaking out " with a dilute acid, the alkaloid is thereby converted into an alkaloidal salt which readily dissolves in the aqueous acid solution. Conversely, an aqueous solution of an alkaloidal salt when treated with an alkali to set free the alkaloid and shaken with an immiscible solvent gives up its alkaloid in a pure state to the latter. The process of assay by this " shaking-out " process is carried out by treating the liquid extracts (after having been freed from alcohol) with an immiscible solvent and a slight excess of an alkali. The alkaloid is thus dissolved in the immiscible solvent, which is then separated from the aqueous liquid, and transferred to another vessel in which it is shaken with an

VOLUMETRIC ASSAYING OF VEGETABLE DRUGS 261

excess of acid largely diluted with water. The acid combines with the alkaloid and forms an alkaloidal salt, which leaves the immiscible solvent, and enters the aqueous solution. If the alkaloidal solution so obtained is still colored, the shaking out is repeated until a pure alkaloidal solution in the immis- cible solvent is obtained. The ap- paratus used in this shaking-out operation is known as a separator. See Fig. 54.

Separators are conical or pear- shaped; the neck is provided with a well-ground glass stopper and the outlet tube or stem at the bottom with an accurately-fitting glass stop- cock. The solvents usually used in alkaloidal drug assaying are alcohol, chloroform, ether and vari- ous mixtures of these containing at least 75 per cent of ether. When chloroform is used it will collect at the bottom of the separator, and can be easily drawn off; but if ether, or ether-chloroform mixture is used, it will form the upper layer in the separator, and may be syphoned off, or the lower aqueous layer drawn off, and the ethereal layer then transferred to another separator. Violent shaking of the contents of the separator is to be avoided, gentle shaking or rotating is quite sufficient to bring about the desired transfer of alkaloid. Excessive shaking will cause the formation of an emulsion of the water and solvent, which is often hard to break up.

The final operation should always be the collection of the free alkaloid by means of a portion of the immiscible solvent. This solution is drawn off into a beaker, the solvent evap-

Squibb's Pattern

FIG. 54.

262 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

orated over a water-bath, using gentle heat, and the dry residue in the beaker dissolved in a measured quantity of standard acid V.S. The excess of acid V.S. is then titrated with a standard alkali V.S. in the presence of a suitable indicator.

In connection with this, attention is called to the fact that some of the most important methods of isolating alkaloids in drug assays and in toxicological investigations depend upon the solubility of the free alkaloids in ether, chloroform, benzene, etc., and the relative insolubility of alkaloidal salts in the same solvents. There are, however, a number of exceptions to this rule, and special attention is called to the fact that in many cases alkaloids pass from decidedly acid aqueous solu- tions (in which they certainly occur as alkaloidal salts) into chloroform and ether, in the shaking-out methods. While the amount of alkaloid so dissolved by these solvents is never very great, still the quantity is an appreciable one. This behavior occurs in the case of caffeine, colchicine, and narco- tine; and under certain conditions also in the case of strychnine, atropine, veratrine and other bases. This transfer of alkaloid to chloroform occurs more particularly in the case of alka- loids of weak basic character, and when the solution is neutral or only feebly acid, or when the alkaloid is in combination with a comparatively weak acid, as citric, tartaric, etc. Fur- thermore, some alkaloidal salts, notably those of weak bases, are transferred as such to chloroform, especially the salts of hydrochloric acid, hydrobromic acid, and nitric acid. In the case, however, of the sulphates, phosphates, tartrates, and citrates of strongly basic alkaloids, no transfer occurs, or at most, only minute quantities of the alkaloids pass over. In order to prevent a transfer of alkaloid or alkaloidal salt out of an aqueous solution to an immiscible solvent, the use of sulphuric acid is to be recommended, and in all toxicological

VOLUMETRIC ASSAYING OF VEGETABLE DRUGS 263

investigations due regard should be paid to the above named conditions. For more details, see the paper by Edward Schaer, Proc. A. Ph. A., 1906, 425.

GENERAL METHODS OF ASSAYING DRUGS

No rule can be formulated as to the method of extraction, or the solvent to be employed, which can be applied to all drugs. Each must be dealt with in accordance with the properties of the contained alkaloids and their state of com- bination. Several methods, however, are in use which may be applied to a large number of different drugs, and with slight special modifications to many more.

Kebler's Modification of the Keller Method.* Treat 10 gms. of the dry powdered drug in a 25o-mil flask with 25 gms. of chloroform and 75 gms. of ether: stopper the flask securely, agitate well for a few minutes and add 10 gms. of ammonia- water and shake frequently during one hour. Then on adding 5 gms. more of the ammonia-water and shaking, the suspended powder agglutinates into a lump, leaving the solution clear after a few minutes' standing. Then proceed by A or B.

A. When the mixture has completely separated, 50 gms. (representing 5 gms. of the drug) are poured off into a beaker and heated on a water-bath until the solvent is evaporated. Ten mils of ether are then added and again evaporated. The varnish- like residue is then dissolved in 15 mils of warm alcohol and water added to slight permanent turbidity, then the indicator is added; followed by an excess of standard acid solution, and the mixture retitrated with standard alkali.

B. When the mixture has completely separated, pour off 50 gms. into a separatory funnel, and add 20 mils of acidulated water, agitate, and when the liquids have separated draw

* J. A. C. S., XVII, 828,

264 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

off the aqueous solution into a second separatory funnel. Repeat this operation with two more portions of 15 mils of acidulated water. Now render the contents of the separatory funnel alkaline by adding ammonia- water. This liberates the alkaloids, which are then separated by treatment with a mix- ture of chloroform, three parts (by volume), and ether, one part, using three successive portions, first 20 mils, then twice 15 mils.

The chloroform-ether solution is heated on a water-bath until the. solvent is evaporated, and then the varnish-like residue treated twice with 8 mils of ether and again evaporated.

The residue is then dissolved in 15 mils of alcohol, water added to slight permanent turbidity, and then the indicator. Titrate in usual way with decinormal acid and centinormal alkali.

A serious objection to this method lies in the taking of a so-called aliquot part: First, because of the well-known solu- bility of ether in water, and conversely of water in ether, as a result of which the volume of the ethereal stratum is materially changed. Furthermore, commercial ether contains variable quantities of alcohol, hence the change in volume will not always be the same.

Another source of error in the aliquot part is found in the volatile nature of the solvents used. In warm weather it is impossible to avoid loss by volatilization, hence the aliquot part taken is too large.

W. A. Puckner * has described a modification of the Keller method which avoids the use of the aliquot part. He uses only one-half of the ethereal solvent for the maceration, and after the usual maceration transfers the drug to a small percolator in which, after the ethereal solution has been well

* Ph. Rev, XVI, 180, and XX, 457.

VOLUMETRIC ASSAYING OF VEGETABLE DRUGS 265

drained off, the marc is percolated with the same menstruum to complete exhaustion. The quantity of ethereal solvent required is not materially greater than in the Keller method, while the quantity of alkaloid obtained for weighing or titrating is larger because it represents the whole of the sample taken for the assay. In the case of drugs containing a very small proportion of alkaloid this is an important advantage.

The objection to this plan is that the transfer of the mass from the flask in which the maceration has been conducted to a suitable percolator, which should not be more than 3 cm. in diameter, requires very dextrous manipulation, or it will be attended with loss of alkaloid.

Assay of Galenical Preparations.

J. U. Lloyd's Method. One gm. of a solid extract which has been dissolved in 5 to 8 mils of an alcoholic menstruum or a corresponding volume of the tincture evaporated to this bulk, or 5 mils of the fluid extract in a flat-bottomed porcelain mortar, are mixed with 2 mils of a solution of perchloride of iron. To this is added sodium bicarbonate with constant trituration until a stiff magma * results. Extract this magma by repeated trituration with chloroform, using first 20 mils and then three portions of 10 mils each, decanting them severally by means of a guiding-rod, being careful that no suspended portions of the magma are drawn off. In order to make sure that all of the alkaloid has been extracted, add 5 mils more of chloroform, draw it off, evaporate on a watch-glass, dissolve

* The ferric hydroxid which is produced in the above process serves to attract most of the tannates, gums, vegetable acids, and coloring matters, while the excess of sodium bicarbonate liberates the alkaloids, which are then dissolved by the chloroform. If the fluid extract is strongly alcoholic the chloroform will not separate easily, in which case the addition of a few mils of water containing a very little glucose will cause a sharp separation,

266 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

residue in dilute sulphuric acid, and test for alkaloids by Wagner's or Mayer's reagent.

The mixed chloroformic extracts are collected and may be estimated volumetrically as follows:

METHOD A. To be used if the chloroformic extract is not colored.

The chloroformic solution is evaporated to dryness in a flask placed on a water-bath. To this residue is added an

N accurately measured excess of — sulphuric V.S. and the solution

diluted with a little water, the indicator added and the excess

N of standard acid solution estimated by titrating with — potas-

0

N sium hydroxid V.S. The number of mils of the — alkali V.S.

N used, divided by 2 and subtracted from the volume of — acid

25 V.S. originally added, will give the number of mils of the

latter required for the alkaloid. This number, multiplied by the proper factor, will give the total alkaloid present in the fluid extract.

Example. The chloroformic residue obtained from 5 mils of a fluid extract of hyoscyamus was dissolved in 12 mils of

N N

— acid V.S., the solution titrated with — alkali V.S. 20.6

25 5°

mils of the latter were used.

N N

20.6 mils of — V.S. is the equivalent of 10.3 mils of -

O j

N V.S. 10.3 mils subtracted from 12 mils, the amount of — acid

^7 N

originally added, leaves 1.7 mils, the quantity of — acid V.S.

0

which was required for the alkaloid. This, multiplied by the

VOLUMETRIC ASSAYING OF VEGETABLE DRUGS 267

N - factor for total alkaloids of hyoscyamus, 0.01157 gm., gives

the quantity of alkaloids present in the 5 mils, which quantity multiplied by 20 gives the per cent :

1.7X0.01157=0.019669 gm.X2o = o.39338 per cent.

METHOD B. This method may be employed if the chloro- formic extract is highly colored, the indicator fluorescein being used.

The residue from the evaporation of chloroform is dis- solved in 10 mils of acid-free alcohol ; then water is added to

N slight turbidity, followed by a measured excess of — acid V.S.;

0

N then the titration is completed with — alkali V.S. The first

50 appearance of fluorescence marks the completion of the reaction.

This is best observed by holding the flask over a dark surface and viewing by reflected light.

METHOD C. This is to be used in the case of highly colored extracts. It consists in removing the alkaloid in a pure state by shaking out in a separating funnel with immiscible sol- vents. The chloroformic extract is shaken out with several portions of acidulated water; this removes the alkaloid, leaving resins, fats, coloring matters, etc., in the chloroform. The acid alkaloidal solution is then treated with ether in a second separator, ammonia added to alkaline reaction, and the alka- loid thus liberated by ammonia dissolves in the ether, from which it is obtained by evaporation and estimated acidimetrically.

ANOTHER METHOD. Add ammonia and shake out directly with an immiscible solvent. Wash the alkaloid out of this with acidulated water and again shake out with a suitable immiscible solvent.

J. Katz's Method (Arch. d. Pharm., 1898, i; Aw. Dr., 1898, 281) . This method has the advantage of all other methods

268 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

in that it enables one to rapidly and accurately estimate the alkaloids in a preparation without the necessity of applying heat for any purpose whatever during the process. The assay may be made in from one to three hours. Twenty-five mils of the tincture of an alcoholic strength of about 45 per cent are placed in a separatory funnel, i mil of a 33 per cent solution of soda added, and the mixture shaken for five minutes with 50 mils of ether, and set aside until the liquids have com- pletely separated into two layers. The lower dark-colored aqueous layer is drawn off into a beaker. The ethereal layer which, besides the alkaloid, has taken up most of the alcohol and some coloring matter, is shaken up with 3 mils of water which, after separating, is drawn off and added to the first aqueous liquid. The ethereal liquid is then poured into a suitable flask, while the combined aqueous liquid is further treated with two portions of ether (25 mils each) the ether to contain about 10 per cent of alcohol. These ethereal extractions are washed each with 1.5 mils of water; the first extraction after washing may be added to the original ethereal solution, while the second is reserved and later on employed for washing the flask. The ethereal solution, which contains still some traces of aqueous fluid containing alkali, is dehydrated by treatment with 2 or 3 gms. of exsiccated calcium sulphate, and finally filtered into a glass-stoppered flask containing 50

mils of water.

N Titration by means of acid V.S. is employed, using

alcholic solution of iodeosin (1-250) as indicator.

The Method as above described is obviously applicable only to such alkaloids as are readily soluble in ether. If an estimation of alkaloids insoluble or only slightly soluble in ether, but soluble in chloroform, is to be made, the method is modified as follows:

VOLUMETRIC ASSAYING OF VEGETABLE DRUGS 269

Twenty-five mils of the tincture of 45 per cent alcoholic strength are vigorously shaken for five.minutes with 30 mils of a mixture of i part of chloroform and 2 parts of ether. The solution so obtained is washed with 3 mils of a 30 per cent solution of sodium chlorid. This operation is repeated twice, using each time 15 mils of the ether-chloroform mixture and washing with 1.5 mils of sodium chlorid solution.

If the Separation of the aqueous and ethereal liquids is not distinct an additional 2 or 3 gms. of sodium chlorid may be used. This prevents the emulsification, which, if pure water were employed, would occur.

If the Tincture to be assayed contains more than 45 per cent of alcohol it is necessary to add water to reduce it to 40 or 50 per cent.

Tinctures containing Chlorophyll or fat or fatty acids must first be deprived of these constituents, as these substances, possessing acid properties inferior to that of iodeosin, will act the part of an alkali toward it and thus be recorded as alkaloid.

To remove the Chlorophyll and fatty acids, acidulate a mixture of equal parts of the tincture and water with a few drops of sulphuric acid, shake with talcum during several hours, and, after subsidence of the latter, filter. Of this filtrate 25 gms. (not mils, on account of the admixture of alco- hol and water causing change of volume) are taken and the alkaloid estimated in the manner already described, after first removing, if necessary, the last traces of fat by a single shaking of the acid solution with petroleum ether.

For the Assay of extracts i to 1.5 gms. are dissolved in from 40 to 50 mils of 45 per cent alcohol to make a solution containing less than 3 per cent extractive. For the assay of fluid extracts 10 mils are taken.

For details of special assays, see the author's Manual of Volumetric Analysis.

270 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

The Influence of the Presence of Certain Volatile Solvents upon the accuracy of alkaloidal titrations must be borne in mind. Alcohol in may instances influences the color changes of the indicators to the extent of rendering them indistinct or entirely unreliable; while ether or chloroform materially diminishes the sensitiveness of many of our most valued indicators, among them phenolphthalein, rosalic acid, congo- red, and luteol. On the other hand, fluorescein and gallein may be mentioned as being more sensitive in the presence^ of alcohol or even of ether. It is advisable, therefore, in most cases, to perform the titration without the presence of the above-named solvents.

Indicators Vary in their Degree of Sensitivenness to the same Alkaloid, hence the choice of an indicator in a particular case is a matter of importance. The following table, based upon that of Kippenberger, will serve as an aid in the selection. The smallest quantity which will procure a distinct tint should be taken.

Atropine. Lacmoid, Fluorescein, lodeosin, Haematoxylin.

Brucine. Cochineal, lodeosin, Haematoxylin.

Cocaine. Lacmoid, Fluorescein, Cochineal, Haematoxylin.

Coniine. Cochineal, Lacmoid, lodeosin, Haematoxylin.

Codeine. lodeosin, Lacmoid, Haematoxylin.

Emetine. lodeosin, Cochineal.

Morphine. H&matoxylin, Cochineal, Lacmoid.

Quinine. Hamatoxylin, Azolitmin, Fluorescein.

Strychnine. Azolitmin, lodeosin, Haematoxylin.

Sparteine. Hamatoxylin, Azolitmin.

CHAPTER XIII

ESTIMATIONS INVOLVING USE OF DECINORMAL BROMIN V.S.

Preparation of Decinormal Bromin V.S. (Koppeschaar's Solution), 61 = 79.92; 7.992 gm. in i liter.

KBr =119.02 NaBr =102.92 KBrO3 = 167.02

This solution does not contain free bromin, but it contains two salts, a bromid and a bromate, which, when treated with hydrochloric acid, liberate a definite quantity of bromin.

It is made as follows: Dissolve 3 gms. of sodium bromate and 50 gms. of sodium bromid (or 3.2 gms. of potassium bromate and 50 gms. of potassium bromid) in sufficient water to make 900 mils.

Transfer 20 mils of this solution by means of a pipette into a bottle having a capacity of about 250 mils, provided with a glass stopper; add 75 mils of water, then 5 mils of pure hydro- chloric acid, and immediately insert the stopper.

Shake the bottle a few times, then remove the stopper just sufficiently to quickly introduce 5 mils of potassium iodid solution (1-5), taking care that no bromin vapor escape, and immediately stopper the bottle.

Agitate the bottle thoroughly, remove the stopper and rinse it and the neck of the bottle with a little water so that the washings flow into the bottle, then add from a burette decinormal sodium thiosulphate until the color of the free iodin is nearly all discharged, then add a few drops of starch

271

272 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

N solution, and continue the titration with — thiosulphate V.S.

until the blue color disappears.

N

Note the number of mils of the - - sodium thiosulphate

10

thus used, and dilute the bromin solution so that equal volumes of it and the -- sodium thiosulphate will exactly correspond

to each other under the above-mentioned conditions.

Example. Assuming that the 20 mils of bromin solution

N required 25.2 mils of the — thiosulphate to completely absorb

the iodin, the bromin solution must be diluted in the pro- portion of 20 to 25.2; that is, each 20 mils must be diluted to make 25.2 mils.

Thus if 850 mils are left, they must be diluted to make 1071 mils, and the solution is decinormal.

A new trial should always be made after diluting, and the bromin solution should correspond, volume for volume, with the decinormal sodium thiosulphate.

The first step in the preparation of this solution is to dissolve the salts; then hydrochloric acid is added, which liberates a definite quantity of bromin, as the equation illus- trates:

5NaBr + NaBrO3 + 6HC1 = 6NaCl + 3Br2 + 3H2O.

The stopper should be inserted into the bottle as soon as the hydrochloric acid has been added, in order that no bromin vapor escape, and the bottle rotated so as to mix the acid thoroughly with the liquid.

The next step is to determine the quantity of bromin which a definite volume of solution will liberate. The bromin solution should be of such strength that 1000 mils of it will

ESTIMATION OF PHENOL 273

contain 7.976 gms. of available bromin. Bromin, like chlorin, liberates iodin from potassium iodid, and is estimated in the same manner.

One atomic weight of iodin is liberated by one atomic weight of bromin:

Thus by determining the quantity of iodin liberated the quantity of bromin is found.

N The iodin is determined by the - - sodium thiosulphate,

one liter of which represents 12.692 gms. of iodin, which is equivalent to 7.992 gms. of bromin, as shown by the following equation :

(Br2) = I2 + 2(Na2S2O3

20)159.84 20)253.84 20)496.44

7.992 gms. 12.692 gms. 24.822 gms. or 1000 mils — V.S.

= 2NaI + Na2S4O6 + ioH2O. i

The Assay of Phenol. Dissolve i gm. of the sample in sufficient water to make 500 mils of solution at the standard temperature. Twenty mils of this solution containing 0.04 gm. of the sample are transferred to a glass-stoppered bottle, having a capacity of about 200 mils.

To this, 30 mils of decinormal bromin, followed by 5 mils of hydrochloric acid, are added, and the bottle immediately stoppered, and shaken repeatedly during half an hour.

Then the stopper is removed just sufficiently to introduce 5 mils of a 20 per cent aqueous solution of potassium iodid, being careful that no bromin escape.

The bottle is then thoroughly shaken and the neck rinsed with a little water, the washings being allowed to flow into the bottle, then i mil of chloroform is added, the mixture thoroughly

274 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

shaken, and titration with decinormal sodium thiosulphate V.S. begun, starch being used as indicator, until the blue color is just discharged.

The precipitated tribromphenol interferes somewhat with the end-reaction when starch is used, and frequently with old phenol solutions * the precipitate possesses a bluish color which is not removed by an excess of sodium thiosulphate and which makes the end-reaction difficult. This difficulty is over- come by the use of a small quantity of chloroform which dissolves the tribromphenol and admits of a very sharp end-reaction.

When chloroform is alone used, the end-reaction is very clearly defined, and is known by a colorless aqueous solution and the chloroform being free from any tinge of pink, due to traces of iodin.

N Note the number of mils of — thiosulphate used; deduct

N this number from 30 mils (the quantity of — bromin originally

N added), and the quantity of — bromin which went into com-

bination with the phenol is obtained.

N Each mil of — bromin represents 0.001556 gm. of pure

phenol.

N Example. Assuming that 6 mils of — sodium thiosulphate

were required to discharge the color of the starch iodid, this deducted from 30 mils leaves 24 mils, the quantity which combined with the phenol.

0.001566X24 = 0.037584 gm.

= 93.96 per cent of pure phenol.

* F.X.Moerk, A. J. Ph., 1904,475-

ESTIMATIONS OF DECINORMAL BROMIN V.S. 275

The above method originated with Koppeschaar, and is the only volumetric method by which accurate results may be obtained.

It is based upon the fact that bromin reacts with phenol, producing an insoluble precipitate of tribromphenol.

The titration is not made directly; but the phenol solution is treated with an excess of standard bromin solution in the presence of some hydrochloric acid. The hydrochloric acid liberates the bromin, and the freed bromin reacts with the phenol, as shown by the equations:

(a) 5NaBr + NaBrO3 + 6HC1 = 6NaCl + 3H

(b) C6H5OH + 3Br2 = C6H2Br3OH + 3HBr.

6)94 6)479.92

10)15.66 10)79.92 N

1.566 gms. 7.992 gms. or 1000 mils — bromin V.S.

N Thus each mil of the — bromin represents 0.001566 gm.

of pure phenol.

The bromin solution which was added in excess, and the liberated bromin of which is not fixed by phenol, is then

N found by residual titration with — sodium thiosulphate after

the addition of some potassium iodid.

The decinormal bromin solution and the decinormal sodium thiosulphate solution being equivalent, each mil of the latter consumed represents one mil of the former. Then by sub- tracting the number of mils of the sodium thiosulphate solu- tion from the number of mils of bromin solution originally added, the quantity of the latter which was actually consumed by the phenol present is found. This number, when mul- tiplied by the factor for phenol, then gives the quantity of pure phenol present.

276 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

The hydrochloric acid used in the above estimation must contain no free chlorin. The potassium iodid must be free from iodate. The starch T.S. should not be added until most of the free iodin has been taken up, and the color of the solution has diminished to light yellow.

The carbolic acid should be diluted with water before titration, and should never be stronger than o.i gm. in 25 mils.

Resorcinol, C6H4(OH)2 i : 3 (110.05). Dissolve about 1.5 gm., accurately weighed, in sufficient distilled water to measure 500 mils. Transfer 25 mils of this solution (representing 0.075 gm.) to a 5oo-mil 'glass-stoppered flask, add 50 mils of deci- normal bromin V.S., and dilute with 50 mils of distilled water. Then add 5 mils of hydrochloric acid and at once stopper the flask, shake well, dilute it with 20 mils of distilled water, add 5 mils of potassium iodid T.S., let stand five minutes, and then titrate the liberated iodin with decinormal sodium thio- sulphate V.S., using starch as indicator.

Calculate as in preceding assay. Each mil of decinormal bromin corresponds to 0.001834 gm. of C6H4(OH)2.

Phenylsulphonates, Sulphocarbolates. Dissolve about 0.25 gm. of the salt, accurately weighed, in 50 mils of distilled water, add 50 mils of decinormal bromin V.S. and 5 mils of hydrochloric acid. Allow the mixture to stand for fifteen minutes, then add 2 gms. of potassium iodid dissolved in 5 mils of water, and titrate the liberated iodin with decinormal sodium thiosulphate V.S., using starch as indicator.

Each mil of decinormal bromin V.S. used corresponds to 0.04903 gm. of NaC6H5O .

CHAPTER XIV

SOME TECHNICAL EXAMINATION METHODS FOR FATS, OILS AND WAXES

The Acid Value or Proportion of Free Fatty Acids. This indicates the number of milligrams of KOH required to neutral- ize the free fatty acids in i gm. of oil, fat or wax. This standard

N N alkali used is in alcoholic solution and may be — , — , or weaker,

depending upon the nature of the fat. Phenolphthalein is the indicator. The fat is dissolved, according to Geissler, in ether, but alcohol or purified methylated spirit, chloroform or a mix- ture of alcohol and ether may be used. The solvent, whichever is used, must be free from acidity, and should be neutralized

N with — alkali if necessary.

The Process. Ten gms. of the oil are accurately weighed into a flask and about 50 mils of solvent added. A few drops of phenolphthalein are then added and the titration with

N

alcoholic -- potassium hydroxid solution begun, shaking con- stantly until the first appearance of a pink color. Care must be taken not to add too great an excess of the alkali, other- wise saponification will occur. A small excess may, however, be added and retitrated with standard acid; a more distinct end-point is then obtained. In the case of waxes or solid fats, the solvent is added, heat applied until it boils, and the titration at once started,

277

278 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

N One mil of — KOH = 0.02805 gm. of KOH or 28.05 mgms.

The number of mils used, multiplied by 28.05 and then divided by the weight of oil taken, gives the milligrams of KOH neutralized by the free fatty acids of the oil, i.e., the acid value.

The Saponification Value (Kottstorfer Number).* This indicates the number of milligrams of KOH required for the complete saponification of one gram of fat or oil. Reagents required are:

Alcoholic Potassium Hydroxid Solution, made by dissolving 40 gms. of potassium hydroxid in i liter of 95 per cent alcohol.

Half-Normal Hydrochloric Acid Solution. Each mil of which = 28.05 mgms- of KOH.

Indicator. Phenolphthalein i gm. in 100 mils of 95 per cent alcohol.

The Process. Into an Erlenmeyer flask capable of holding 200 mils, accurately weigh about 1.5 gms. of the fat (previously purified and filtered). Run into this from a burette 25 mils of the alcoholic potassium hydroxid solution. Then insert into the neck of the flask a perforated stopper provided with a glass tube 70 to 80 cm. in length and from 5 to 8 mm. in diameter, and set it in a water-bath for half an hour or until the fat is entirely saponified. The operation is facilitated by occasional agitation. The flask is then removed, its con^ tents cooled and titrated with the half-normal hydrochloric acid, using phenolphthalein as indicator. The alcoholic potas, sium hydroxid solution is standardized by conducting a blank experiment similar in every detail to the above with the exception that the fat is omitted.

* J. Kottstorfer, 1879, Zeitschr. anal. Chem., XVIII, 199.

TECHNICAL EXAMINATION METHODS FOR FATS 279

The Kottstorfer number is then ascertained by subtracting the number of mils of the standard hydrochloric acid used in the analysis from the number necessary to neutralize 25 mils of the alcoholic potassium hydroxid solution in the blank experiment, multiplying the difference by 28.05 and dividing by the weight of fat taken.

Example. 1.5 gms. of the fat was treated with 25 mils of alcoholic potassium hydroxid solution, under the above described

N

conditions, and in titrating the excess, 21.5 mils of -- hydro- chloric acid was required. In the blank experiment 25 mils of the alcoholic potassium hydroxid required 32 mils of the

N

— acid. Then 32 — 21.5 = 10.5.

10.5X28.05

196.

196 is the milligrams of KOH neutralized by i gm. of the oil, or the Kottstorfer saponification number.

TABLE SHOWING REQUIREMENT AS TO SAPONIFICATION

NUMBER

Lard oil 195 to 197

Almond oil (expressed) . . 191 to 200

Cottonseed oil 191 to 196

Linseed oil 187 to 195

Cod-liver oil 175 to 185

Olive oil 191 to 195

Castor oil 179 to 180

Theobroma oil 188 to 195

Croton oil 212 to 218

Yellow wax 90 to 96

The Reichert Number or Volatile Fatty Acid Value. This indicates the number of mils of decinormal KOH required to neutralize the volatile fatty acids distilled from 2.5 gms. of fat. This method is conducted as follows: 2.5 gms. of the clear filtered fat are taken in an Erlenmeyer flask of about 150 mils capacity with i gm. of potassium hydroxid and 20 mils

280 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

of 80 per cent alcohol, and the whole digested on a water-bath (rotating the flask frequently) until saponification is complete and the alcohol all removed. 50 mils of water are then added, and 20 mils of dilute sulphuric acid (1:10) and the mixture distilled. The distillate is collected in a 5o-mil flask into which is set a funnel with a wetted filter to collect any insoluble fat acid. The first 10 or 20 mils of distillate are returned to the flask, and the 50 mils distilled. This is treated with a few drops of phenolphthalein and then titrated with decinormal potassium hydroxid. The number o'f mils consumed is the Reichert Number.

In the Meissl modification (the Reichert-Meissl method) 5 gms. of the fat are taken and a more complete distillation of the volatile acids effected. The method is particularly useful in determining the genuineness and purity of butter.

The Reichert-Meissl Number. This is undoubtedly the best process for detecting the admixture of foreign fats with butter. This process depends upon the fact that butter con- tains certain constituents which, when appropriately treated, yield volatile acids in much larger quantity than is obtained from any of the practicable substitutes for butter. These acids are principally butyric and caproic. The process con- sists in saponifying the fat with an alkali, then separating the fatty acids by neutralization, and distilling off the volatile acids for titration with standard alkali.

The operations involved in this process do not admit of any arbitrary variation, and reliable and comparable results can only be obtained by strictly adhering to the prescribed details.

The following process is adopted by the Association of Official Agricultural Chemists. The solutions required are:

Sodium Hydroxid Solution. One hundred gms. of sodium hydroxid are dissolved in 100 mils of distilled water. The

TECHNICAL EXAMINATION METHODS FOR FATS 281

alkali should be as free as possible from the carbonates, and be preserved out of contact with the air.

Potassium Hydroxid Solution. One hundred gms. of the purest potassium hydroxid are dissolved in 58 mils of hot dis- tilled water, cooled in a stoppered vessel, and the clear liquid decanted and preserved out of contact with the air.

Sulphuric Acid. Mix 200 mils of the strongest acid with 1000 mils of water.

Alcohol of about 95 per cent, redistilled from caustic soda.

Standard Barium Hydroxid Solution, accurately standard- ized, approximately decinormal.

Indicator. Dissolve i gm. of phenolphthalein in 100 mils of 95 per cent alcohol.

The process:

Weighing the Butter. The butter to be examined should be melted and kept in a dry, warm place at about 60° C. for two or three hours, until the water and curd have entirely settled out. The clear supernatant fat is poured off and filtered through a dry filter-paper in a jacketed funnel con- taining boiling water. Should the filtered fat in a fused state not be perfectly clear, it must be filtered a second time. This is to remove all foreign matter and any trace of moisture. The saponification flasks are prepared by thoroughly washing with water, alcohol and ether, wiping perfectly dry on the outside, and heating for one hour at the temperature of boiling water. The flasks should then be placed in a tray by the side of the balance and covered with a silk handkerchief until they are perfectly cool. They must not be wiped with a silk handkerchief within fifteen or twenty minutes of the time they are weighed. The weight of the flasks having been accurately determined, they are charged with the melted fat in the following way:

A pipette with a long stem, marked to deliver 5.75 mils,

282 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

is warmed to a temperature of about 50° C. The fat, having been poured back and forth once or twice into a dry beaker in order to thoroughly mix it, is then taken up in the pipette and the nozzle of the pipette carried nearly to the bottom of the flask, having been previously wiped to remove any adhering fat, and 5.75 mils of fat are allowed to flow into the flask. After the flasks have been charged in this way they should be recovered with the silk handkerchief, allowed to stand fifteen or twenty minutes, and again weighed.

The Saponification. Three methods may be employed: i. Under Pressure with Alcohol. Ten mils of 95 per cent

alcohol are added to the fat in the flask, and then 2 mils of the caustic soda solution. A soft cork stopper is now inserted in the flask and tied down with a piece of twine. The saponi- fication is then completed by placing the flask upon the water or steam bath (see Fig.

55). The flask during the saponification, which should last one hour, should be gently rotated from time to time, being careful not to pro- ject the soap for any distance up its sides. At the end of an hour the flask, after having been cooled nearly to room temperature, is opened.

2. Under Pressure without the Use of Alcohol. Place 2 mils of the potassium hydroxid in the flask containing the fat, which must be round-bottomed and made of well-annealed

TECHNICAL EXAMINATION METHODS FOR FATS 283

glass to resist the pressure; cork, and heat as in the previous method. Rotate the flask very gently during the saponifica- tion, taking great care that none of the fat rises on the sides of the flask out of reach of the alkali. Potash makes a softer soap than soda and thus allows a complete saponification without the use of alcohol. This method avoids the danger of formation of esters and the trouble of removing the alcohol after saponification.

3. With a Reflux Condenser and the Use of Alcohol. Place 10 mils of the 95 per cent alcohol in the flask containing the fat, add 2 mils of the sodium hydroxid solution with a reflux condenser (a glass tube not less than i meter in length is allowable) and heat on the steam bath until the saponification is complete.

Removal of the Alcohol. The stoppers having been laid loosely in the mouth of the flask, the alcohol is removed by dipping the flask into a steam bath. The steam should cover the whole of the flask except- the neck. After the alcohol is nearly removed, frothing may be noticed in the soap, and to avoid any loss from this cause or any creeping of the soap up the sides of the flask, it should be removed from the bath and shaken to and fro until the frothing disappears. The last traces of alcohol vapor may be removed from the flask by waving it briskly, mouth down, to and fro, or better, by a current of carbon dioxid free air.

Dissolving the Soap. After the removal of the alcohol the soap should be dissolved by adding 135 mils of recently boiled distilled water (or 132 mils if potassium hydroxid was used in the saponification), warming on the steam bath, with occasional shaking until solution of the soap is complete.

Setting Free the Fatty Acids and Distilling. Cool to from 60° to 70° C., throw in a few pieces of pumice stone, add 5 mils of the dilute sulphuric acid (or 8 mils if potassium hydroxid

284 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

was used in the saponification), stopper as in the method of saponification, and heat on the water-bath until the fatty acids form a clear, transparent layer on top of the water. This may take several hours. Cool to room temperature, add a few pieces of pumice stone, and connect with a glass condenser, as in Fig. 56.

Heat slowly with a naked flame until ebullition begins, and distil, regulating the flame hi such a way as to collect no mils of distillate in as nearly thirty minutes as possible.

Mix this distillate, filter through a dry filter, and titrate 100 mils with the standard barium hydroxid solution, using

FIG. 56.

0.5 mil of phenolphthalein as indicator. The red color should remain unchanged for two or three minutes.

Increase the number of cubic centimeters of tenth-normal alkali used by one-tenth, divide by the weight of fat taken, and multiply by five to obtain the Reichert-Meissl number. Correct the result by the figure obtained in a blank exper- iment.

When treated as above described, 5 gms. of genuine butter

N never yields less acidity than is represented by 24 mils of —

TECHNICAL EXAMINATION METHODS FOR FATS 285

alkali. It is true, however, that the butter made from the milk of a single cow, especially towards the end of her period of lactation, has been known to fall slightly below this figure, but the average butter, as produced from the mixed milk of a herd, usually requires 27 mils or more. Oleomargarin re- quires about i mil of beef-fat, and lard about the same. Cacao butter requires about 7 mils.

The percentage of butter-fat in a mixture of fats, 5 gms. being taken: (n - 0.6) X 3.65 = percentage of true butter-fat.

lodin Absorption Number of Fats and Oils (HiibPs Num- ber.)* This is the percentage of iodin absorbed by a fat or an oil under certain conditions. In other words it is the number of parts of iodin absorbed by 100 parts of an oil.

Reagents required:

(a) HiibFs Iodin Solution. Dissolve 25 gms. of pure iodin in 500 mils of alcohol, and mix this solution with 500 mils of alcohol containing 30 gms. of pure mercuric chlorid, and set aside for twenty-four hours. The mercuric chlorid solution should be filtered if necessary before it is mixed with the alcoholic iodin solution. This solution loses its strength rapidly, and should therefore be tested before using.

(b) Decinormal sodium thiosulphate.

(c) Potassium iodid solution, 20 gms. in 100 mils.

(d) Starch paste indicator.

The Process. To about 0.3 gm.| of the fat or oil, accu- rately weighed and dissolved in 10 mils of chloroform, con-

* Dingler's Polyt. Jour., 1884, 281; Am. Ch. Jour., VI, 285.

t In the case of drying oils which have a very high absorbent power, as linseed oil, use from 0.15 to 0,20 gm.; for oil of theobroma and similar fats use 0.80 gm,

286 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

tained in a glass-stoppered bottle of 250 mils capacity, add 25 mils of the iodin solution. Stopper the bottle securely ana place in a cool, dark place for four hours.* At the expiration of this time, the mixture must still possess a brown color; if it does not, a further measured quantity of the iodin solution must be added and the mixture again set aside. Finally 20 mils of the potassium iodid solution are added together with 50 mils of water, and the mixture titrated with the decinormal sodium thiosulphate until the color is almost discharged, when a few drops of starch indicator are added and the titration continued until the solution is colorless.

The Standardization of the Iodin Solution is effected by subjecting it to the same treatment as in the assay, but with the oil omitted, and at the end of four hours titrating with the decinormal thiosulphate.

The difference between the number of mils of thiosulphate solution used in the blank test and the number used in the actual assay, is multiplied by 12.59, and this divided by 3, gives the iodin value of the oil under analysis.

When the quantity of the oil or fat taken is not 0.3 gm., then the product is not divided by 3, but by the figure corre- sponding to the quantity taken; thus if 0.15 gm. are taken the product is divided by 1.5.

Another way of making the calculation is as follows:

The difference between the mils of thiosulphate used in the blank test and the mils used in the analysis, is multiplied by 0.01259, then by 100, and the product divided by the weight in grams of the oil taken for analysis.

* The time allowed does not give the complete iodin absorption power of an oil or fat and cannot be compared with determinations in which six to twelve hours have been used. It gives, however, very satisfactory comparative results, but the time factor must be very closely observed.

TECHNICAL EXAMINATION METHODS FOR FATS 287

Example:

Number of mils of thiosulphate used in blank test ....... 45-4

" " " " " " " analysis ......... 26.1

" " " " representing iodin absorbed .......... 19.3

10.3X0.01259X100 Iodin number is - J = 80.9.

This method, as is well known, is based upon the fact that the unsaturated glycerids in the oils form addition products with the iodin. The mercuric chlorid and iodin contained in the alcoholic solution interact with a formation of mercuric chloriodid and iodin chlorid; the latter is supposed to be the active agent.

2=HgClI+ICl.

The mercuric chlorid also acts the part of a carrier of halogen similar to that played by mercury in the Kjeldahl process when dissolving the substance in sulphuric acid.

Gill and Adams, J. A. C. S., xxn, 12, call attention to the fact that not only addition products, but also substitution products, are formed in this process, which vary in amount with the time of action and the strength of the solution. This is a feature which interferes with the accurate determination of the iodin number, giving as it does a higher figure. In order to prevent this formation of substitution products, and thus overcome the discrepancy in results, these authors suggest the use of mercuric iodid instead of mercuric chlorid, making the solution with methyl alcohol (free from acetone and anhydrous). They claim that by the use of a solution so made a truer iodin absorption number is obtained. A great disadvantage of the Hiibl method is that the solution quickly loses its strength; another is, the length of time required

288 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

for the absorption. As above described, four hours are required, but this is not sufficient time to allow of complete absorption of the iodin. It is, however, a good practice to have a fixed time limit; the process then gives very satisfactory comparative results.

Iodin Absorption Number. The Hanus Method.* This method is conducted like the foregoing. It differs in the com- position of the iodin solution used. This is prepared as fol- lows: Dissolve 13.2 gms. of powdered iodin in 1000 mils of pure glacial acetic acid with the aid of gentle heat if necessary, to facilitate solution. Cool to 25° C. and determine the iodin content in 20 mils of the solution by means of decinormal sodium thiosulphate V.S. ; then add to the solution a quantity of bromin, molecularly equivalent to that of the iodin present (3 mils is the usual approximate amount). Keep the solution in glass-stoppered bottles protected from light.

The Process. Introduce about 0.8 gm. of a solid fat or 0.3 gm. of an oil, accurately weighed, into a glass-stoppered bottle of 250 mils capacity. Dissolve it in 10 mils of chloro- form, add 25 mils of the iodin solution, stopper the bottle securely, and allow the mixture to stand for half an hour in a cool, dark place. After this time it must still retain a brown color; if it is not brown, a new test should be started using a smaller quantity of the fat or oil. Then add in the order named, 30 mils of potassium iodid T.S., 100 mils of distilled water, and decinoimal sodium thiosulphate V.S. in small successive portions, shaking thoroughly after each addi- tion, until the color of the mixture becomes quite pale. Then add a few drops of starch solution and continue the addition of the thiosulphate V.S. until the blue color is discharged. Then make a blank test by mixing exactly the same quantities

* Zeitschr. Nahr. u. Genus. (1901), 913.

TECHNICAL EXAMINATION METHODS FOR FATS 289

of the iodin solution and chloroform, and titrating the free iodin with thiosulphate V.S. as directed above. The dif- ference in the number of mils of the thiosulphate V.S. con- sumed in the blank test and the actual test multiplied by 1.269 and divided by the weight of the fat or oil taken gives the iodin number.

In the case of oils which have a high iodin-absorbing power, smaller quantities of the oils should be used. 0.15 to o.i 8 gm. for linseed oil and about 0.2 gm. for cod liver oil. The lime allowed for absorption in the case of castor oil and linseed oil should be one hour.

The consensus of opinion among chemists is that the Hanus method is the most satisfactory. Its principal advantage over the Hiibl method lies in the facts (a) that the Hanus iodin solution is much more stable, (b) that the time required for the reaction is comparatively much shorter.

TABLE SHOWING IODIN ABSORPTION NUMBER FOR SOME COMMON OILS

Hubl's No. Hanus' No.

Four Hours. Half an_Hour.

Almond oil (expressed) 98. 2 98. 7

Butter 35.4 35.4

Castor oil 87.3 87.5

Cocoanut oil 8.4 8.6

Cod liver oil 144 .3 143 . 9

Cottonseed oil 104 . 3 106 . 6

Lard oil 69.5 69.7

Linseed oil 179. 6 186. 2

Mustard oil 113.1 i*5-3

Olive oil 86.1 86.8

Oleomargarin 52.1 52.3

Peanut oil 96 . 3 97-3

Rape oil 102 . 4 105 . 2

Sesame oil 106 . 4 106 . 5

Theobroma oil 35.4 35 . 5

290 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

The Acid Number for Resins. This is the number of milligrams of potassium hydroxid required to neutralize i gm. of a resinous substance.

The Process. About 2 gms. of the resinous substance, accurately weighed, is dissolved in alcohol, and the solution titrated with half-normal potassium hydroxid V.S., using phenol- phthalein as indicator. The number of mils of the half- normal alkali V.S. required for neutralization when multiplied by 0.028055 and the product divided by the weight, gives the acid number of the resin.

CHAPTER XV ESTIMATION OF SUGARS

Preparation of Fehling's Solution, (a) The Copper Solu- tion. 34.66 gms. of carefully selected small crystals of pure cupric sulphate are dissolved in sufficient water to make, at or near 25° C., exactly 500 mils. Keep in small well-stoppered bottles.

(b) The Alkaline Tar irate Solution. 173 gms. of potassium and sodium tartrate (Rochelle salt) and 75 gms. of potassium hydroxid are dissolved in sufficient water to make, at or near 25° C., exactly 500 mils. Keep in small rubber-stoppered bottles.

For use, equal quantities of the two solutions should be mixed at the time required.

One molecular weight of water-free glucose will reduce five molecular weights of cupric oxid, i.e., 180.12 gms. of glucose will reduce 1248.7 gms. of crystallized copper sul- phate (CuSO4 + sH2O). If pure chemicals are used in the preparation of this solution there will be no need of standard- izing it, but if the solution is old and its titer doubtful, the following method of standardization may be employed.

The Standardization. Dissolve 0.95 gm. of pure cane sugar in 50 mils of water, add 2 mils of hydrochloric acid, and heat to 70° C. for ten minutes. Then neutralize with sodium carbonate and dilute to i liter.

Fifty mils of this solution should exactly reduce 10 mils of Fehling's solution, when treated as described below.

291

292 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

Ten mils of the mixed Fehling's solution is equivalent to

Glucose 0.050 gm.

Maltose ... 0.0806 "

Inverted cane-sugar 0.0475 "

Inverted starch 0.045 "

Lactose 0.0678 "

Estimation of Reducing Sugars. (Glucose, Maltose, In- verted Cane-sugar, Lactose and Inverted Starch. 0.5 gm. or less of the sugar is dissolved in 100 mils of distilled water, and put into a burette. Ten mils of the mixed Fehling's solu- tion, accurately measured, is put into a loo-mil Erlenmeyer flask, or in a white porcelain dish, diluted with 40 mils of distilled water, and rapidly heated to boiling. The solution should remain perfectly clear and retain its blue color. If it does not it should be rejected.

The hot, clear blue solution is immediately titrated with the sugar solution, which should be added in small portions at a time, boiling after each addition until the copper is com- pletely precipitated and the blue color of the solution is entirely destroyed. The solution is then boiled about two minutes longer, and the amount of the sugar solution used is read off.

The Calculation. Ten mils of Fehling's solution are always taken; and whatever the quantity of glucose or other sugar solution is required to effect its complete reduction, that quantity contains the equivalent of 10 mils of Fehling's solu- tion, as shown in the table above.

Thus if 12 mils of the sugar solution were required to reduce 10 mils of Fehling's solution, the 12 mils contain 0.05 gm. of glucose or 0.0806 gm. of maltose, etc. One hundred mils of the solution therefore contain x gm. of glucose.

0.05 X 100

12

= 0.416 gm. glucose.

ESTIMATION OF SUGARS 293

fci order to obtain reliable results it is important that the process be carried out exactly as laid down in the above directions, and that the quantity of sugar present in solution be no greater than one per cent. The degree of heat and the time occupied in the process, as welt as the concentration of the Fehling's solution, have a very important bearing upon the accuracy of the results.

It is advisable to complete the titration in as short a time as possible. A preliminary test should always be made, in which the approximate quantity of the solution required is found; then a second and more accurate titration can be done in which the sugar solution may be added more boldly, and the time of boiling and exposure of the copper solution to the air much lessened. The complete reduction of the copper (using undiluted Fehling's), after the addition of the requisite quantity of sugar, does not take place instantly. The time required varies somewhat with the different sugars. For instance, with glucose, invert sugar and levulose, the reduction is not complete until after heating two minutes; with maltose four minutes, and with lactose six minutes are required.

If the sugar to be examined be either glucose, maltose, or lactose, it may be titrated directly; but if it be sugar-cane, it must first be inverted. This is done by dissolving the sugar (0.5 gm.) in about 100 mils of water, adding 3 or 4 drops of strong hydrochloric acid, and boiling briskly for ten or fifteen minutes. This is then allowed to cool, neutralized with potas- sium hydroxid, and made up to 100 mils with water.

The sugar in urine may be estimated by this process. The urine is placed in the burette and run into the boiling Fehling's solution in the usual manner. If it contain a large quantity of sugar, it must be diluted two or three times.

In estimating with Fehling's solution it is well to attach a

294 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

rubber tube eight to twelve inches in length to the lower end of the burette, so that the boiling need not be done directly under the burette, and thus cause incorrect readings through the expansion of the liquid therein.

Determination of the End-point. It is always somewhat difficult to determine the exact point at which the blue color disappears, owing to the presence of the precipitated suboxid of copper. This difficulty may be overcome by the addition of some substance which will prevent the precipitation of the cuprous oxid, such as ammonium hydroxid or potassium ferrocyanid. When the latter is used the disappearance of the blue color can then be readily seen, as the solution remains clear to the end, turning from blue to green, and finally brown, which indicates the end of the reaction.

Professor Bartley reports this method as accurate, reliable, and rapid, provided the solution be not boiled during the reduction. He recommends adding to the Fehling's solution in the porcelain basin 10 mils of a 10 per cent freshly prepared solution of potassium ferrocyanid and 30 mils of water. The ferrocyanid does not precipitate the copper in alkaline solution.

L. Beulaygue (Compt. rend., 138, 51) suggests the following method for determining the end-reaction when titrating sugaf solution in the usual manner with Fehling's reagent, using solution of sodium monosulphid as the indicator. When the end of the reaction is near, a little of the hot solution is applied by means of a glass rod, to two superimposed white filter papers. The upper one acts as a filter, retaining the par- ticles of cuprous oxid. The lower paper is withdrawn, and the moist spot touched with a drop of the sodium monosulphid reagent, when an immediate black stain of cupric sulphid is formed if the reaction is not complete. By successive spotting out and testing in this manner, a very accurate reading of the end-reaction may be obtained. It is important, when

ESTIMATION OF SUGARS 295

standardizing the Fehling's solution, that the same indicator should be employed. Potassium ferrocyanid in solution, acidified by either hydrochloric or acetic acid, may be employed in a similar manner. The end-reaction in then indicated by the disappearance of the red color from the last spot.

E. F. Harrison,* who has employed Fehling's solution somewhat extensively in quantitative sugar estimations, has found the indicators usually recommended to determine the end-point of reaction to be unsatisfactory. It was suggested by him that the action of cupric salts in liberating iodin from iodid might be utilized for this purpose with advantage, and his experiments determine the superiority of this over the other indicators heretofore proposed. The indicator is pre- pared by boiling 0.05 gm. of starch with a few mils of water, adding 10 gms. of potassium iodid and diluting to 100 mils. This indicator should be prepared as required. In use 0.5 to i.o mil of this solution is acidified with about 5 or 10 drops of acetic acid, and one drop or more of the liquid in process of titration added. As long as unreduced copper is present, a color is produced, varying from red to blue, and of greater or less intensity, according to the nearness of the end-point. The production of no color marks the end of the reaction. The indicator is available with one drop of a solution containing one part of cupric sulphate in twenty thousand parts.

Estimation of Maltose in Malt Extract. A half per cent solution of the sample is prepared, and titrated into 10 mils of Fehling's solution in the manner described above.

Estimation of Starch After Inversion. This method con- sists in converting the starch into glucose and then estimating the glucose with Fehling's solution.

Two gms. of the starch are mixed with water, and boiled

* Trans. Brit. Ph. Conf., 1903, 568-9.

296 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

for fifteen minutes; the solution is then cooled, and diluted to about 200 mils. Eighteen mils of strong hydrochloric acid are then added, and the solution heated under a return con- denser for two or three hours. The solution is then cooled, neutralized with potassium hydroxid and diluted to 250 mils. This is put into a burette and titrated into 10 mils of Fehling's solution, as described above, under reducing sugars.

In estimating the starch in baking powder, 2 to 5 gms. of the powder are introduced into an Erlenmeyer flask, 150 to 200 mils of a 4 per cent solution of hydrochloric acid are added and the solution gently boiled for four hours, after which the flask and contents are cooled, neutralized by adding sodium hydroxid, and made up to a definite volume. It is then ready for testing with Fehling's solution.

Ten mils of Fehling's solution is the equivalent of 0.0475 gm. of invert starch.

Estimation of Starch after Inversion by Means in Diastase. The treatment of starch with malt infusion or pure diastase at a temperature not above 70° C., readily converts the starch into maltose, but the solution also contains, besides maltose, various dextrins in proportions varying with the temperature at which the diastase acts. The digestion may vary from fifteen minutes to fifteen hours. Complete conversion of the starch may be determined by testing occasionally with iodin. A blank experiment should be made, especially if the digestion is carried on beyond half an hour. A like quantity of the same diastase solution should be digested at the same tem- perature and for the same time, and the amount of sugar found deducted from the total quantity found in the analysis. Faulenbach* makes use of the following solution of diastase: Crush 3.5 kilos of fresh green malt, treat with a mixture of two liters of water and four liters of glycerin and let stand

* Zeitschr. f. physiol. Chem., VII, 510; and Chem. Centralh., 1883, 632,

ESTIMATION OF SUGARS 297

for one week, stirring occasionally; then express and filter.

This solution is very stable. Five drops of it will dissolve

1 gm. of starch; 15 drops of it contain a quantity of carbo- hydrate = 0.00 1 gm. of glucose.

A quantity of the substance to be tested (containing about

2 gms. of starch) is boiled to gelatinize the starch. Fifteen drops of the diastase solution are then added, and the mix- ture digested at 63° C. It is then filtered to separate the undissolved cellulose, etc., and heated with 20 mils of hydro- chloric acid in a water-bath for three hours. The acidity is then just destroyed by means of caustic soda, the glucose determined, o.ooi gm. deducted, and the starch calculated from the glucose.

O'Sullivan* employs pure diastase, prepared as follows: Pour sufficient water over 2 or 3 kilos of finely crushed pale malt to just cover it. Let stand for three or four hours, then express and filter the solution. Add alcohol (sp.gr. 0.83) until the liquid above the flocculent precipitate becomes opal- escent. Collect the precipitate, wash it with alcohol (sp.gr. 0.86 to 0.88) then with absolute alcohol and press it between linen. Finally dry it in a vacuum over sulphuric acid.

Estimation of the Diastasic Value of Malt Extract. The diastasic value of malt extracts may also be determined by estimating the amount of maltose produced by a given amount of the extract in a given time, when brought in contact with an excess of gelatinized starch solution. It is always necessary to estimate the copper-reducing power of the extracts with Fehling's solution upon a separate sample, and to deduct this from the total reducing power found after treatment with the starch.

The process is briefly as follows: A definite quantity (say 30 mils) of gelatinized starch, made from the best Bermuda

* Jour. Chem. Soc.? XLV (1884), i.

298 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

arrowroot, of 3 per cent strength, is placed in a flask and heated to 40° C. A weighed amount of the malt extract (say 20 mils of a i per cent solution) is then added, and the temper- ature kept at 40° C. for exactly half an hour. At the end of this time some sodium hydroxid solution is added in order to check the further action of the diastase upon the starch (heating to 100° C. accomplishes the same result) . The solution is then diluted to a definite volume with water and the maltose produced, estimated by Fehling's solution in the usual manner. The quantity of reducing sugar originally present in the sample must be previously determined, and this amount deducted from the total amount found after treatment with starch, and the remainder calculated as maltose. Ten mils of Fehling's solution is the equivalent of 0.0806 gm. of maltose.

CHAPTER XVI ESTIMATION OF FORMALDEHYDE *

The Ammonia Method (Legler). This method is based upon the reaction between free ammonia and formaldehyde in which hexamethylene-tetramin is formed. It is for ordinary purposes sufficiently accurate.

It is this method which is recommended by Lederle f for use in the laboratory of the New York City Health Depart- ment, and by Prescott in the laboratory of the Universtiy of Michigan.

The assay is conducted as follows: 2 mils of the solution are placed in a glass-stoppered bottle, the stopper of which

N is thickly coated with petrolatum, and 50 mils of — ammonia

solution added; let stand twelve hours, shaking occasionally.

N Then determine the excess of ammonia by titrating with —

sulphuric acid solution, using rosolic acid or litmus as indicator. The excess of ammonia subtracted from the quantity added gives the quantity which combined with the formaldehyde, and thus the amount of the latter is ascer- tained.

* Berichte d. Chem. Ges., XVI., 1335, 1883. f Am. Drug., 1897, 246.

299

300 THE ESSENTIALS OF VOLUMETRIC ANALYSIS The reaction is represented as follows:

6CH20 + 4NH3 = (CH2)6N4 + 6H2O.

4)180 4)68.04 Hexamethylenetetramin

2)45_ 2)17.01 N

22.5 gms. 8.5 gms. or 1000 mils — V.S.

0.0225 gin. 0.0085 gm. or i mils " "

N Assuming that 22 mils of — sulphuric acid were employed

N in the titration, 22 mils of the — ammonia solution must have

been in excess, hence 28 mils of the latter went into combina- tion with the formaldehyde. Thus the 2 mils of formaldehyde solution contained 28X0.0225=0.63 gm.

A. G. Craig* says that the chief difficulty in using the Legler method is the volatility of the ammonia. The difficulty is not so much the loss of strength in the standard solution, but the loss during the determination. He proposes the follow- ing scheme by which this error is removed.

Prepare a normal solution of sulphuric acid. Make up an approximately normal solution of ammonia, the exact strength being immaterial. Procure several three-ounce prescription bottles with smooth sides and close-fitting soft rubber stoppers. Prepare a methyl orange solution. Procure a boiler in which the bottles may be immersed to the neck without upsetting (a large beaker will do). Take as much of the sample as will contain 0.5 gm. of formaldehyde. Measure with the pipette, 25 mils of the ammonia solution into each of the bottles, and to half of them add a sample of formaldehyde; stopper tightly. If the necks of the bottles are small, the stoppers need not be tied down. Place the bottles in the boiler, add cold water to the necks, and heat to boiling. Boil for one

* J. A. C. S., XXIII, 642 (1901).

ESTIMATION OF FORMALDEHYDE 301

hour, and cool by running in cold water slowly, being careful not to allow the cold water to touch the hot bottles. Titrate with sulphuric acid and methyl orange to the first indication of a color change. Take the difference between the readings for the blanks and those for the samples, as the ammonia consumed in normal mils. Of this difference, i mil =0.0601 gm. of formaldehyde.

The Legler method is also liable to error because the compound formed is a weak base, and as such combines with acid, while at the same time it is liable to decompose into ammonia and formaldehyde and thus give an indefinite end-point when the residual ammonia is titrated with acid. Error is also liable to be introduced, through the presence of carbonic acid in the ammonia water, which, with the indicator rosolic acid, gives no sharp end-reaction.

The Ammonium Chlorid Method. In this method a solu- tion of ammonium chlorid is used, from which ammonium is evolved by treatment with sodium hydroxid. The excess of alkali is then determined by titration with standard solution of sulphuric acid. This method, as devised by H. Schiff,* and Modified by C. A. Male,f is as follows:

Introduce 2 gms. neutral ammonium chlorid, dissolved in 20 mils of water, into a flask or bottle of about 200 mils capacity, having a well-fitting stopper. Dilute 10 mils of the formaldehyde solution to 100 mils with water, and neutralize with sodium hydroxid solution, as the formaldehyde solution generally con- tains varying quantities of formic acid. Add 20 mils of this neutralized solution to the ammonium chlorid solution, then

N 25 mils of — NaOH, and immediately stopper the flask, and

* Chem. Ztg., XXVII, 14 (1903).

f Pharm. Jour., June, 1905, 844. See also Carl E. Smith, A. J. Ph., LXX, 86, (1898).

302 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

leave for one hour. Afterwards determine the excess of alkali

N

with — H2SO4, using rosolic acid or litmus solution as indi- cator, both of which give up a sharp end-reaction. The reaction and calculation is based upon the following equation:

2NH4C1 +3CH20 + 2NaOH= N2(CH2)3 + 2NaCl + sH2O.

N i mil — NaOH is equivalent to 0.045 g111- °f formic aldehyde.

This modified method is quite as simple, and gives results practically identical with that of Romijn.

Oxidation by Hydrogen Dioxid (Blank and Finkenbeiner) * This method depends upon the use of hydrogen dioxid for oxidizing formaldehyde into formic acid, in alkaline solution, The formic acid so produced neutralizes a portion of the alkali, and the excess of the latter is then determined by titratioi with standard acid. The method gives good results and can be very rapidly carried out.

Three mils of the formaldehyde solution are placed into a 100- mil Erlenmeyer flask, the latter closed with a well-fitting stopper, and the weight of the solution carefully taken. 50 mils of normal NaOH V.S. are then added, and followed immediately, but slowly, through a small funnel, by 50 mils of a 3 per cent hydrogen • dioxid solution, previously neutralized with normal NaOH V.S., using a drop of litmus solution. The solution is allowed to stand about 15 minutes, or until the reaction has ceased. The funnel and the sides of the vessel are then rinsed with distilled water, a few drops of litmus solution added and the unconsumed alkali titrated with normal H2SO4 V.S. The mils of the latter deducted from the 50 mils of nor- mal NaOH V.S. originally taken gives the quantity of the * Berichte d. Chem. Ges., XXXI, 2979 (1898); and A. J. Ph., 1899, 486.

ESTIMATION OF FORMALDEHYDE 303

alkali V.S. which was consumed in the reaction. The dif- ference multiplied by 0.03 gm. gives the weight of CH2O in the sample. To find per cent multiply the result by 100 and divide by the weight of sample taken for analysis.

In the original process, double normal sodium hydroxid is used.*

The reaction is as follows:

CH2O + H2O2 + NaOH = NaCOOH + 2H2O

HsO2 oxidizes formaldehyde to formic acid, HCOOH, which reacts with and neutralizes NaOH.

The lodometric Method (Romijn).| This method, which is considered the most rapid, most accurate, and most readily applied, depends upon the fact that iodin in the presence of an alkali acts as an indirect oxidizing agent, giving, when formaldehyde is present, the iodid of the base and formic acid:

As modified by L. Reuter { it is as follows: Twenty mils of 35 to 40 per cent forma dlehyde are intro- duced into a graduated flask and distilled water added to bring up the volume to 500 mils. Of this thoroughly mixed fluid, 5 mils are introduced into a bottle capable of being perfectly closed. 30 mils of normal sodium or potassium hydroxid are

N added, and then — iodin solution allowed to flow in from

a burette with constant agitation, until the fluid remains a bright yellow color (36-60 mils). The shaking is then vig- orously continued for one minute, when 40 mils of normal

*See C. Allen Lyford, J. A. C. S., XXIX, 1227 (1907) "The Action of Barium Peroxid and Hydrogen Peroxid upon Formaldehyde."

fZeitschr. anal. Chem., XXXVI, 18-24 (1897); ibid, XXXIX, 60-63 (1900).

J Ph. Rev., 1903, 207.

304 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

sulphuric acid are added, and then the excess of iodin titrated

N with decinormal thiosulphate solution. Each mil of - - iodin

solution consumed corresponds to 0.0015 gm. formaldehyde. A blank titration should always be made.

A solution containing as much as 5 per cent may be accurately estimated by this method, provided not more than 2 gms. be taken. The method is, however, not accurate in the presence of other aldehydes. F. O. Taylor,* commenting upon this method, says that it is quite satisfactory, but that the quantities of formaldehyde and reagents used are unnec- essarily large and cumbersome, and the method is hence modi- fied by him as follows:

From a weighing bottle, consisting of a small Erlenmeyer flask, fitted with a perforated rubber stopper through which passes a dropper, and containing about 25 or 30 mils of the formaldehyde solution, weigh out accurately about 10 gms. of the solution into a stoppered 5oo-mil flask and fill this to the mark with distilled water. For titration remove 5 mils of this solution, corresponding to o.i gm. of the weighed quan- tity of formaldehyde, and put into a 2oo-mil Erlenmeyer flask. Into another flask put 5 mils of water for a blank titration. To both now add 20 mils of normal NaOH and then 20 mils

N of an approximately — iodin solution, whose exact strength

need not be known, and let stand for five or ten minutes for the entire completion of the reaction. Now add 25 mils of

N

normal H2SO4 and titrate the excess of iodin with — Na2S2Os.

10

The difference between the mils of thiosulphate used on the

N assay and the mils of the blank is the number of mils of — iodin

J 10

* Bull. Ph., Aug., 1903, 323.

ESTIMATION OF FORMALDEHYDE 305

N consumed by the formaldehyde. Each mil of - - iodin so

used equals 0.0015 gm. of CH2O.

The Potassium Cyanid Method.* This method is especially applicable to solutions containing small quantities of formal- dehyde. It depends upon the fact that potassium cyanid and formaldehyde combine to form an addition product, in which one molecule of potassium cyanid combines with one molecule of formaldehyde, as shown by the following equation:

H

I CH2O + KCN = NEEEC— C— O— K.

H

In the estimation, the formaldehyde is mixed with a known quantity of potassium cyanid (in excess), the excess of the latter being determined by the use of standard silver nitrate solution, and thus the quantity of potassium cyanid, which combined with formaldehyde is found, and from this the quantity of formaldehyde is calculated.

N The process is carried out as follows:^ Ten mils of —

silver nitrate are treated with 6 drops of 50 per cent nitric acid in a 5o-mil flask. Ten mils of a solution of potassium cyanid (containing i gm. of KCN in 500 mils of water) are then added and well shaken. An aliquot part of the nitrate,

N say 25 mils, is then titrated by Volhard's method with —

ammonium sulphocyanate for excess of silver.

Another 10 mils of silver nitrate solution is then acidified

* Romijn, Zeitschr. anal. Chem., XXXVI, 18-24 (1897). t Bernard H. Smith, J. A. C. S., XXV, 1032 (1903).

306 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

with nitric acid and treated with 10 mils of the potassium cyanid solution to which has been added a measured quantity of dilute formaldehyde solution. The whole is made up to 50 mils and then filtered. 25 mils of the filtrate are titrated

N with — ammonium sulphocyanate for excess of silver as

before.

The difference between these two results multiplied by 2 gives the amount of potassium cyanid which was used by

N the formaldehyde, in terms of — sulphocyanate.

N Each mil of — V.S. represents 0.003 gm- °f formaldehyde.

The best results by this method are obtained, if the solu- tion of formaldehyde is diluted to below i per cent. With i per cent solutions it is necessary to use 15 mils of the silver solution. In estimating very dilute solutions, it is advisable to use a 200-mil flask and to take 100 mils of the filtrate for the titration. It is possible by this method to determine with accuracy one part in 100,000.

Paraformaldehyde (CH2O)3. This is a polymeric form of formaldehyde and may be assayed by the same methods. The U.S.P. IX recommends the process in which hydrogen dioxid is used, the procedure being practically the same.

Each mil of normal alkali hydroxid V.S. corresponds to 0.03 gm. of (CH2O)s.

Acetone (Dimethyl-ketone) (CH3. CO. CH3 = 58.05). This may be assayed by the iodometric method, used for formalde- hyde. One mil of acetone is accurately weighed in a stoppered weighing-flask tared together with some distilled water. This is transferred to a looo-mil flask and diluted to the mark with distilled water. Twenty-five mils of normal potassium hydroxid V.S. are placed in a 25o-mil glass-stoppered flask and to this

ESTIMATION OF FORMALDEHYDE 307

is added exactly 25 mils of the acetone solution (representing one-fortieth of the weight of acetone taken), then with constant agitation of the flask, 35 mils of decinormal iodin V.S. is added, and the mixture allowed to stand for fifteen minutes. At the end of this time 26 mils of normal hydrochloric acid V.S. is added and the titration begun at once, with decinormal sodium thiosulphate V.S., using starch as indicator.

A blank test is conducted with the same quantities of re- agents. The difference in the quantity of decinormal iodin V.S. consumed in the blank test and that in the assay, repre- sents the quantity taken up by the acetone.

Each mil of this difference corresponds to 0.0009675 gin. of acetone. The reaction is

CHAPTER XVII

ESTIMATION OF ALCOHOL IN TINCTURES AND BEVERAGES

THE quantity of alcohol contained in dilute spirit, which leaves no residue upon evaporation, may be ascertained by taking the sp.gr. and referring to the alcohol table. When taking the specific gravity, the temperature of the liquid should be i5|° C. (60° F.).

In Wines, Beer, Tinctures and other alcoholic liquids containing vegetable matter, the sp.gr. of the sample is taken at 15!° C. (60° F.) and noted. A certain quantity (say 100 mils) is measured off and evaporated to one half, or until all odor of alcohol has passed off, the evaporation being con- ducted without ebullition, in order that particles of the material may not be carried off by the steam. The liquid left is then diluted with distilled water, cooled to 60° F. and made up to the original volume (100 mils), and the sp.gr. taken. Lastly, we calculate: the sp.gr. before evaporating is divided by the sp.gr. after evaporating, and the quotient will be the sp.gr. of the water and alcohol only of the liquor. Then by referring to the alcohol table the percentage of alcohol contained in the liquor is obtained.

Example. The liquor before evaporating had a sp.gr. of 0.9951; after evaporation and dilution to 100 mils the sp.gr. was found to be 1.0081.

•—- = 0.987, the sp.gr. of the contained spirit.

308

ALCOHOL IN TINCTURES AND BEVERAGES 309

TABLE FOR ASCERTAINING THE PERCENTAGES RESPECTIVELY OF ALCOHOL BY WEIGHT, BY VOLUME, AND AS PROOF SPIRIT, FROM THE SPECIFIC GRAVITY.

Condensed from the excellent Alcohol Tables of Mr. Hehner in the " Analyst," vol. v. pp. 43-63.

Specific Gravity 15-5°.

Absolute Alcohol by w'ght. Per cent.

Absolute Alcohol by vol'me. Per cent.

Proof Spirit. Per cent.

Specific Gravity 15-5°.

Absolute Alcohol by w'ght. Per cent.

Absolute Alcohol DV vol'me. Per cent.

Proof Spirit. Per cent.

I .0000

o.oo

0.00

o.oo

-9489

35-05

41-90

73-43

• 9999 .9989

0.05 0.58

0.07 0.73

0. 12 1.28

.9479 .9469

35-55 36.06

42.45 43-Qi

74-39 75-37

•9979

I . 12

1.42

2.48

• 9459

36.61

43 63

76.45

.9969

i-75

2 . 2O

3-85

• 9449

37-17

44.24

77-53

• 9959

2.33

2-93

5-13

• 9439

37-72

44.86

78.61

• 9949

2.89

3-62

6-34

.9429

38.28

45-47

79-68

• 9939

3-47

4-34

7-61

.9419

38.83

46.08

80.75

.9929

4.06

5-o8

8.90

.9409

39-35

46.64

81.74

.9919

4.69

<.86

10.26

• 9399

39-85

47.18

82.69

• 9909

5-31

6.63

ii .62

.9389

40.35

47-72

83.64

.9899

5-94

7.40

12.97

• 9379

40.85

48.26

84.58

.9889

6.64

8.27

14.50

.9369

41-35

48.80

85.53

.9879

7-33

9-13

15-99

•9359

41.85

49-34

86.47

.9869

8.00

9-95

17-43

• 9349

42.33

49-86

87-37

.9859

8.71

10.82

18.96

• 9339

42.8!

50.37

88.26

.9849

9-43

ii . 70

20.50

•9329

43.29

50.87

89.15

• 9839

10. 15

12.58

22.06

.9319

43.76

51.38

90.03

.9829

10.92

13-52

23-70

• 9309

44.23

51-87

90.89

.9819

ii .69

14.46

25-34

.9299

44-68

52.34

91-73

.9809

12 .46

15.40

26.99

.9289

45-14

52.82

92.56

• 9799

13-23

16.33

28.62

.9279

45-59

53-29

93-39

.9789

14.00

17 .26

30.26

.9269

46.05

53-77

94-22

•9779

14.91

18.36

32.19

.9259

46.50

54-24

95-05

.9769

15-75

19-39

33-96

.9249

46.96

54.71

95-88

•9759

16.54

20.33

35-63

•9239

47-41

55-18

96.70

• 9749

17-33

21 .29

37-30

.9229

47-86

55-65

97.52

• 9739

18.15

22.27

39-03

-9219

48.32

56. ii

98.34

• 9729

18.92

23.19

40.64

.9209

48.77

56.58-

99.16

.9719

19-75

24.18

42.38

.9199

49-20

57.02

99-93

• 97O9 .9699 .9689

20 . 58 21.38 22.15

25- X7 26. 13 27.04

44. 12

45-79 47-39

.9198

49-24

57-06

loo.ooPs

9670

2 2 O2

2 7 • 9 5

48. 98

.9669

44 .y*

23 69

28.86

50-57

.9189

49-68

57-49

100. 76

• 9659

24.46

29.76

52.16

.9179

50-13

57-97

101 .59

.9649

25-21

30.65

53-71

.9169

50-57

58.41

102.35

• 9639

25-93

31.48

55-18

•9159

51 .00

58-85

103.12

.9629 .9619

26.60 27.29

32.27 33-06

56.55 57-94

.9149 • 9139

51.42 51-83

59-26 59-68

103-85 104.58

.9609

28.00

33.89

59-40

.9129

52.27

60. 12

I05-35

• 9599

28.62

34.61

60.66

.9119

52.73

60.56

106.15

.9589

29.27

35-35

6i.95

.9109

53-17

61 .02

106.93

• 9579

29-93

36. 12

63.30

• 9099

53-6i

61.45

107.69

• 9569

30.50

36.76

64-43

.9089

54-05

61.88

108.45

• 9559

3L06

37-41

65-55

.9079

54.52

62.36

109.28

• 9549

31.69

38.11

66.80

.9069

55-00

62.84

IIO. 12

•9539

32.31

38.82

68.04

• 9059

55-45

63.28

110.92

• 9529

32.94

39-54

69.29

• 9049

55-91

63-73

III.?!

• 9519

33-53

40.20

70.46

•9039

56.36

64-18

112.49

• 9509

34-10

40.84

7I-58

.9029

56.82

64-63

113 .26

• 9499

34-57

41-37

72.50

.9019

57-25

65-05

113-99

310 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

Specific Gravity 15-5°.

Absolute Alcohol by w'ght. Per cent.

Absolute Alcohol by vol'me. Per cent.

Proof Spirit. Per cent.

Specific Gravity 15-5°-

Absolute Alcohol by w'ght. Per cent.

Absolute Alcohol by vol'me. Per cent.

Proof Spirit. Per cent.

.9009

57-67

65-45

114.69

.8429

82. 19

87.27

!52.95

.8999 .8989

58.09 58.55

65.85 66.29

115.41 116.18

.8419 .8409

82.58 82.96

87.58 87.88

I53-48 154.01

59-00

66.74

116.96

.8399

83.35

88.19

154.54

.8969

59-43

67-15

117.68

-8389

83-73

88.49

155-07

• 8959

59.87

67.57

118.41

-8379

84.12

88.79

155-61

• 8949

60.29

119.12

.8369

84.52

89.11

156. 16

• 8939

60. 71

68 36

I I 9 . 80

• 8359

84.92

89.42

156.71

.8929

61.13

68.76

120.49

.8349

85.31

89.72

157-24

.8919

6i.54

69-15

121.18

.8339

85.69

90.02

157-76

.8909

61.96

69-54

121.86

.8329

86.08

90.32

158.28

.8899

62.41

69.96

122 .6l

.8319

86.46

90.61

158.79

.8889

62.86

70.40

123.36

.8309

86.85

90.90

159.31

,8879

63-30

70.81

124.09

.8299

87.23

91 .20

159.82

.8869

63-74

71.22

124.80

.8289

87.62

91.49

160.33

.8859

64.17

71.62

125.51

.8279

88.00

91-78

160.84

.8849

64.61

72.02

126.22

.8269

88.40

92.08

161.37

.8839

65.04

72.42

126.92

.8259

88.80

92-39

161.91

.8829

65.46

72.80

127-59

.8249

89-19

92.68

162.43

.8819 .8609

65-88 66.30

73-19 73-57

128.25 128.94

.8239 .8229

89.58 89.96

92.97 93-26

162.93

163.43

.8799

66.74

73-97

129.64

.8219

90.32

93-52

163.88

.8789 .8779

67-17 67.58

74-37 74-74

130.33 130.90

.8209 .8199

90.68 91.04

93-77 94 • 03

164.33 164.78

.8769

68.00

75-12

131.64

.8189

91-39

94.28

165.23

• 8759

68.42

75-49

132.30

.8179

91-75

94-53

165.67

-8749

68.83

75.87

132.95

.8169

92 . ii

94-79

166. 12

-8739

69.25

76.24

133.60

.8159

92.48

95.06

166.58

.8729 .8719

69.67 70.08

76.61 76.98

134.25 134.90

.8149 .8139

92.85

93-22

95-32 95.58

167.04 167.50

.8709

70.48

77-32

135-51

.8129

93-59

98.54

167.96

.8699

70.88

77.67

!36.i3

.8119

93.96

96. ii

168.24

.8689

71.29

78.04

136.76

.8109

94.31

96.34

168.84

.8679

71.71

78.40

137-40

.8099

94 66

96.57

169.24

.8669

72.13

78.77

138.05

.8089

95-00

96.80

169.65

.8659

72.57

79.16

138.72

.8079

95-36

97-05

170.07

.8649

73.00

79-54

139-39

.8069

95-71

97-29

170.50

.8639

73-42

79.90

140.02

.8059

96.07

97-53

170.99

.8629

73.83

80.26

140.65

.8049

96.40

97-75

171.30

.8619

74-27

80.64

141.33

.8039

96.73

97-96

171.68

.8609

74-73

81.04

142.03

.8029

97-07

98.18

172.05

• 8599

75-iS

81.44

I 42 . 77

.8019

97-40

98.39

172.43

.858?

75.64

81.84

143.42

.8009

9? • 73

98.61

172.80

-8579

76.08

82.23

144.10

•7999

98.06

98.82

173-17

.8569 .8559

76.50 76.92

82.58 82.93

144.72 145-34

.7989 •7979

98.37 98.69

99-00 99-i8

173-50 173-84

-8549

77-33

83.28

145-96

.7969

99-00

99 • 3 7

174-17

.8535

77-75

83.64

146.57

.7959

99-32

99-57

174-52

.8529

78.16

83.98

147.17

• 7949

99.65

99-77

174-87

.8519

78.56

84-31

147-75

•7939

99-97

99.98

175-22

.8509

78.96

84-64

148.32

.8499

79.36

84.97

148.90

.8489 .8479

79-76 SO. I 7

85.29 85-63

149-44 150.06

Absolute Alcohol.

.8469

80.58

85-97

150.67

.8459

Sl.OO

86.32

151.27

.8449

81.40

86.64

151.83

• 8439

81.80

86.96

152.40

.7938

IOO.OO

IOO.OO

175-25

ALCOHOL IN TINCTURES AND BEVERAGES 311

Then by referring to the table we find that this sp.gr. cor- responds to 7.33 per cent, by weight, of absolute alcohol.

Another Way is to boil the liquid in a retort, condense the vapor, and when all the alcohol has passed over add sufficient water to the distillate to make up the original vol- ume, at the temperature of 15$-° C. (60° F.). Then, by taking the sp.gr. of this diluted distillate, the quantity of absolute alcohol is found by reference to the table. This latter method requires the taking of the sp.gr. but once and gives more accurate results.

The details of this method are as follows: Place 25 mils of the liquid measured at any definite temperature between 15° and 30° C. in a distilling flask of about 200 mils capacity and connected with a suitable condenser; add 50 mils of dis- tilled water, distil into a 5o-mil graduated flask at such a rate that about 48 mils of distillate will be received in a half hour. Dilute with water of the same temperature to 50 mils. Take the specific gravity at any definite temperature between 15° and 30° C. by means of a pycnometer, and refer to the alcohol- ometric table for per cent of absolute alcohol and multiply by the dilution factor to find the per cent of alcohol in the original liquid.

If the preparation is suspected to contain less than 25 per cent of alcohol, take 50 mils for analysis; if it is suspected to contain over 50 per cent, take 12.5 mils for analysis. In either case, the volume of the distillate should be made up to 50 mils.

If the preparation contains iodin, conbine it with zinc dust, or add sufficient sodium thiosulphate to take up the iodin. In the latter case add a few drops of sodium hydroxid T.S. to prevent sulphur from distilling over. If volatile acids are present, neutralize with sodium hydroxid T.S. If volatile bases are present neutralize with diluted sulphuric. If both

312 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

volatile acids and volatile bases are present, neutralize first with dilute sulphuric acid and distil about 50 mils, then neu- tralize the distillate with sodium hydroxid T.S. and again distil to obtain 50 mils of distillate.

If acetone, camphor, chloroform, ether, glycerin (30 per cent or over) volatile oils or other volatile products are sus- pected to be present, transfer the first distillate to a separator, saturate it with sodium chlorid, add 15 mils of petroleum benzin (boiling point 40-50° C.) and shake for one or two minutes. After the liquids have completely separated, draw off the lower alcoholic salt solution into a second separator and repeat the extraction with 15 mils of petroleum benzin. Again draw of the lower alcoholic salt solution. Introduce this into a 2OO-mil distilling flask. Wash the combined benzin solutions with about 25 mils of saturated sodium chlorid solu- tion and add the washings to the distilling flask, and distil about 45 or 48 mils, bring to the original temperature, dilute with water of the same temperature to 50 mils, determine specific gravity and alcohol per cent as directed above.

PART III

A FEW GASOMETRIC ANALYSES

CHAPTER XVIII THE NITROMETER

FOR general gas analysis, and for the rapid estimation of such substances as* ethyl nitrite, hydrogen dioxid, urea, bleaching powder, manganese dioxid, etc., an instrument called the nitrometer is used.

The apparatus in its simplest form is shown in Fig. 57. It consists of a measuring tube, a, of 50- or loo-mils capacity, and graduated in tenths of a mil. This is connected by means of a stout rubber tube with an open equilibrium tube, b, also called " control-tube," " pres- sure-tube," or " level- tube." Both tubes are preferably provided with a globular expan- sion near the lower end, and are held by suitable clamps upon a stand, in such a manner that either tube may be readily and quickly clamped at a higher or lower level. The measuring tube is fitted at the top with a stop-cock, c, and a graduated glass tube or cup, d. Some nitrometers are pro- vided with a three-way stop-cock, so arranged that according to the way it is' turned, it will discharge the contents of the cup either into the measuring tube below, or out into the

313

FIG. 57.

314 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

waste opening which is usually placed at e, or it will discharge the contents of the measuring tube into the waste opening.

With this apparatus gases can be rapidly and accurately measured at definite temperature and pressure.

In measuring the gas the instrument is filled with some liquid in which the gas is insoluble — generally mercury. In many cases a saturated solution or salt may be used.

Suppose we fill the instrument with mercury in such quan- tity that when the stop-cock is opened and the control-tube raised, the mercury will rise as far as the top, and about two inches in the control-tube.

The top is now closed, the control-tube lowered, and a little carbonic acid gas admitted through e. The top is then again closed, and the instrument allowed to stand until its contents have acquired the temperature of the room. A centigrade thermometer suspended to the stand will then give the temperature of the gas.

The control-tube is now raised or lowered so as to make the level of the liquid in both tubes the same. This makes the pressure in the tube the same as the atmospheric pressure outside, and, by referring to a barometer standing near, this pressure is ascertained.

We now have a definite volume of the gas at a known temperature and pressure.

It now only remains to read off the volume of the gas, and correct it to the normal temperature and pressure by Charles' and Boyle's laws, respectively.

The normal temperature and pressure is o° C. and 760 mm. pressure. Although the U.S.P. adopts a standard tem- perature of 25° C.

The weight of the gas in grams may then be calculated from its volume by multiplying the number of mils at the normal tempera- ture and pressure, by the weight of i mil of the gas in grams

THE NITROMETER 315

This weight may be found as follows:

1000 mils of hydrogen at normal temperature and pres- sure weigh 0.0896 gm. One mil of H then weighs 0.0000896 gm.

One mil of oxygen weighs sixteen times as much, and i mil of nitrogen weighs fourteen times as much. There- fore, by multiplying the weight of i mil of H by the atomic weight of an elementary gas, or half the molecular weight of a compound gas, the weight of i mil of that gas is obtained.

According to the Law of Charles, the volume of a gas under constant pressure varies directly with the absolute tem- perature.

All gases expand or contract by -ji^- of their volume for each centigrade degree of temperature increased or decreased.

We may regard a gas at o° C. as having passed through 273° C. In other words, 273° below zero must be regarded as the absolute zero, and o° C. as 273° absolute temperature.

Thus the absolute temperature centigrade is the observed temperature + 2 73°.

Example. A given volume of oxygen gas at 15° C. measures 20 mils. What will it measure at o° C. ?

273° X 20

15° 4-273°

Boyle's Law. The volume of a confined gas is inversely proportional to the pressure brought to bear upon it. That is, the less the pressure, the greater the volume, and vice versa.

Rule. Multiply the observed volume by the observed pres- sure, and divide by the normal pressure.

Example. A given volume of gas at 750 mm. pressure

316 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

measures 20 mils. What will it measure at 760 mm. (the nor- mal pressure) ?

7 50X20 mils

76o =19.73 mils. Ans.

Now let us take an example in which both laws are involved.

A given volume of oxygen at 15° C. subjected to a pressure of 750 mm. measures 20 mils. What will it measure at the normal temperature and pressure? — i.e., o° C. and 760 mm.

In the first example we find that 20 mils of oxygen at 15° C. will measure at o° C. 18.95 mu<s- Then

750X18.95 mils

-=18.70 mils. Ans. 760

Now to find the weight of this volume of oxygen, we pro- ceed as follows :

i mil of H weighs 0.0000896 gm.;

i mil of O weighs 16X0.0000896 = 0.0014336 gm.;

18.70 mils of 0 = 18.70X0.0014336 gm., or 0.02680832 gm.

In the U.S. P. the standard for temperature and pressure is 25° C. and 760 mm. and hence all pharmaceutical assays should be made in accordance with this standard, and all gases measured at this temperature and pressure, or corrections made by calculation.

Example. A volume of gas measures 20 mils at 15° C. What will it measure at 25° C. ?

25° + 273° X 20 298X20

15° + 273° ^8^=20-7mlls-

The following tables from the United States Pharma- copoeia IX will be found very useful for making temperature and pressure corrections.

THE NITROMETER

317

FACTORS FOR TEMPERATURE CORRECTIONS

(Normal Temperature, 15° C.)

Temperature.

Factor.

Temperature.

Factor.

Temperature.

Factor.

15° C.

•035

22° C.

I.OIO

29° C.

0.987

i6°C.

.031

23° C.

1.007

30° C.

0.983

17° C.

.028

24° C.

1.003

31° C.

0.980

i8°C.

.024

25° C.

I.OOO

32° C.

0.977

19° C.

.021

26° C.

0.997

33° C.

0.974

20° C.

.017

27° C.

0-993

34° C.

0.971

21° C.

.014

28° C.

0.990

35° C.

0.968

Example. Assuming that the volume of a gas read off was 41 mils at 30° C. and it is desired to ascertain the corre- sponding volume at 25° C., then the 41 mils must be multiplied by 0.983. The result will be 40.30 mils as the equivalent volume of gas at 25° C.

FACTORS FOR CORRECTION FOR BAROMETRIC PRESSURE

(Normal Barometer, 760 mm.)

BAROMETEB Millimeters.

READING. Inches.

Factor.

BAROMETEF

Millimeters.

,. READING. Inches.

Factor.

790

31.10

1.039

660

25-98

0.868

78o

30.71

1.026

650

25.59

0-855

770

30.31

1.013

640

25.20

0.842

76o

29.92

I.OOO

630

24.80

0.829

750

29-53

0.987

620

24.41

0.816

740

29.13

0.974

610

24.02

0.803

730

28.74

0.961

600

23.62

0.789

720

28.35

0.947

590

23-23

0.776

710

27-95

0-934

580

22.83

0.763

700

27.56

0.921

570

22.44

0.750

690

27.17

0.908

560

22.05

0-737

680

26.77

0.895

550

21.65

0.724

670

26.38

0.882

540

21. 26

0.711

Example. Assuming that the volume of gas read off was 41 mils at 590 mm. barometric pressure, and it is desired to ascertain the corresponding volume at normal pressure (760 mm.), then the 41 mils must be multiplied by 0.776. The result will be 31.81 mils.

CHAPTER XIX ASSAY OF NITRITES

Spirit of Nitrous Ether. This is an alcoholic solution of ethyl nitrite ^HsNC^^ 75.05), yielding, when freshly pre- pared and tested in the nitrometer, not less than 1 1 times its own volume of nitrogen dioxid (NO = 30.01).

When nitrites are mixed with an excess of KI and acid- ulated with H2SO4, iodin is liberated, and all the nitrogen of the nitrite is evolved in the form of NO, as shown in the equation

2C2H5NO2+2KI+ 2H2SO4= 2C2H5OH + 2KHSO4 + 12 + 2NO. 150.1 60.02

The process is conducted as follows:

Open the stop-cock of the measuring tube, raise the control- tube, and pour into the latter a saturated solution of NaCl until the measuring tube, including the bore of the stop-cock, is completely filled. Then close the stop-cock and fix the control-tube at a lower level. Now introduce into the funnel at the top of the measuring tube a weighed quantity (about 4 gms.) * of spirit of nitrous ether; open the stop-cock and allow the spirit to run into the nitrometer, being careful that

* It is conveninent to take 5 mils accurately measured, and calculate its weight by multiplying by the specific gravity, but better to take about 40 mils of the spirit, weigh it accurately, and then add sufficient alcohol to make exactly 100 mils. This is mixed thoroughly and 10 mils of the solution taken for analysis.

318

ASSAY OF NITRITES 319

no air enters at the same time. Ten mils of potassium iodid T.S. are now added in the same manner, and followed by 10 mils of normal sulphuric acid V.S. Effervescence takes place immediately, and after thirty to sixty minutes, when the volume of gas has become constant, the control-tube is lowered so as to make the level of the liquid in both tubes the same, and the volume of the gas in the graduated tube read off.

This volume, multiplied by 0.00307 gm., gives the weight of ethyl nitrite in the spirit taken for analysis. The product multiplied by 100, and then divided by the weight of the spirit taken, gives the per cent of pure ethyl nitrite present.

The temperature correction is one-third of one per cent of the total percentage found, for each degree, additive if fhe temperature is below, sub tractive if above 25° C. The barometric correction is ^ of one per cent for each millimeter, additive if above, subtractive if below, 760.

The volume of NO generated at the ordinary indoor tem- perature (assumed to be at or near 25° C., 77° F.) should not be less than 55 mils if 5 mils of the spirit are taken, corre- sponding to about 4 per cent of pure ethyl nitrite.

Sodium chlorid solution is used in the above assay, because, owing to its density, the spirit will float on top, and the gas evolved will not dissolve in it. At the same time the expense of using mercury is saved. It is important that no air be allowed to get into the measuring tube, because this would convert the NO into a higher oxid of nitrogen, which would dissolve in the salt solution, and thus vitiate the result.

Example. Four gms. of spirit of nitrous ether (sp.gr. 0.823) are treated in a nitrometer, and the NO evolved measures 55 mils.

The temperature at which the operation is conducted is 25° C., and the atmospheric pressure normal.

320 THE ESSENTIALS OF VOLUMETRIC ANALYSIS What per cent of ethyl nitrite is present in the sample ?

0.00307X55X100

- : - =3.^2 per cent.

Amyl Nitrite is a liquid containing about 80 per cent of amyl nitrite (principally iso-amyl nitrite), C5HuNO2= 117.09, together with variable quantities of undetermined compounds.

The assay is as follows: Take about 3 mils of the amyl nitrite, which has been previously shaken with 0.5 gm. of potassium bicarbonate and carefully decanted. Put into a tared i co-mil measuring flask, and weigh it accurately; add alcohol to bring the volume to exactly 100 mils. Ten mils of this alcoholic solution are introduced into the nitrometer as directed for spirit of nitrous ether; 10 mils of potassium

N iodid solution (20 per cent) and 10 mils of — H2SO4 V.S.

are then added, and the volume of NO generated, measured at the ordinary indoor temperature (25° C. or 77° F.), should be about 40 mils. Each mil at this temperature represents 0.0048 gm. of pure amyl nitrite.

Sodium Nitrite (NaNO2 = 69.01). This, like the other nitrites mentioned, when treated with potassium iodid and sulphuric acid, is decomposed, and NO is given off. The reaction is here illustrated:

= K2SO4 + Na2SO4 + 2H2O + 2NO + 12.

A molecule of NaNO2 (69.01) evolves, when properly treated, one molecule of NO (30.01).

The assay process is as follows: Weigh out 0.15 gm. of NaNO2, dissolve it in about 5 mils of water, and introduce the solution into a nitrometer. This is followed by a solution

ASSAY OF NITRITES 321

N of i gm. of KI in 6 mils of water and 15 mils of — H^SCU.

The gas which is liberated should measure not less than 50 mils at 15° C. (59° F.) or 51.7 mils at 25° C. (77° F.), corre- sponding to not less than 97.6 per cent of the pure salt. Each mil at 25° C. represents 0.002837 gm. and at o° C. 0.0030916 gm. of pure NaNC>2.

Nitric Acid in Nitrates. This may also be effected by the use of the nitrometer.

When a nitrate is shaken up with an excess of sulphuric acid and mercury, the nitrate is decomposed and NO is evolved; as seen in the following equation:

2)202.22 2)60.02

IOI.II 30.01

Thus each molecule of the nitrate radical NOs gives off a molecule of NO.

Not more than 0.2 gm. of nitrate should be taken for analysis, since, if this quantity is exceeded, the volume of gas evolved willl be greater than the instrument can conve- niently hold. In this estimation the nitrometer is filled with mercury instead of brine; the nitrate is dissolved in 5 cc. of water, introduced into the nitrometer, and followed by excess of strong sulphuric acid. The instrument is well shaken for some time, and when action has ceased and the contents have cooled down to the temperature of the room, the level is adjusted and the volume of NO read off and cal- culated in the usual way.

322 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

ESTIMATION OF NITROGEN DIOXID

NO = 30.01 ; i liter at o° C. and 760 mm. = 1.3406 gms. at 25° C. and 760 mm. = 1.2281 gms.

ONE MILLILITER OF NITROGEN DIOXID IS THE EQUIVALENT OF:

At o° C. and

760 mm. Gram.

At 25° C. and

760 mm. Gram.

Nitrogen dioxid, NO = 30.01

o 0013406

O 0012281

Amyl nitrite, C6HnNO2 = 117.10

O. OO523I

0.0048

Ethyl nitrite, C2H6NO2= 75.05

O.OOHZ2&

0.003071

Sodium nitrite, NaNO2 = 6Q 01

O 0030828

O 002824

CHAPTER XX HYDROGEN DIOXID

As stated in a previous chapter, hydrogen dioxid when acted upon by an acidulated solution of potassium perman- ganate, is decomposed and oxygen is evolved. One-half of this oxygen comes from the dioxid and the other half from the permanganate.

Therefore if i mil of the dioxid be treated in this way and 20 mils of oxygen are evolved, the strength of the solu- tion is 10 volumes.

The nitrometer may be used for this estimation.

This instrument is charged with a concentrated solution of sodium sulphate (which in this case is better than brine), and i mil of the dioxid introduced from the funnel, followed by excess of solution of permanganate acidulated with sul- phuric acid.

This latter solution should be of such strength that when the reaction is completed, the solution should still have a purple color.

The reaction is thus illustrated:

5H2O2 + 3H2SO4 + 2KMnO4 = K2SO4+ 2MnSO4+8H2O + 5O2.

By the use of Squibb' s Urea Apparatus (Fig. 61) the estimation may be easily and rapidly made.

Into the generating bottle is put about 30 mils of a strong, acidulated solution of potassium permanganate, and a small test-tube containing i mil of H2O2 is carefully introduced. The two liquids must not be allowed to come in contact.

323

324 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

The larger flask is filled with water, or better, a solution of sodium sulphate; the connection is then made by means of the rubber tube, and the generating-bottle tipped over and agitated so that the liquids will mix and the reaction take place.

The liberated oxygen then passes into the larger bottle, displacing an equal volume of water, which is collected and measured. Half of this volume represents the volume strength of the H2O2.

An Improvised Nitrometer may be used. The author has found the following instrument convenient:

To the bottom of an ordinary 5o-mil burette is attached a suitable length of rubber tubing, to the other end of which is attached another burette or ungraduated tube, which serves as a control-tube.

Into the top of the burette is fitted a rubber stopper, through which passes a short glass tube, which is connected by means of a rubber tube to a generating-bottle similar to that used with Squibb1 s Urea Apparatus. Into the control- tube is poured the solution of sodium sulphate, sufficient to fill the burette to the zero mark and have the surface of the liquid in both tubes on a level.

About 30 mils of strong permanganate solution acidulated with sulphuric acid are now placed in the generating-bottle, and then the small test-tube or homeopathic vial, containing exactly i mil of hydrogen dioxid, is placed in. The gener- ating bottle is then stoppered and agitated, the evolved gas passes over, and forces down the liquid in the burette; the control-tube is then lowered so as to bring the surfaces of the liquid in both tubes* on a level.

The reading is then taken.

Each mil of gas represents one-half volume of oxygen evolved from the dioxid if i mil of the latter is used. Each mil

HYDROGEN DIOXID 325

of oxygen evolved from i mil of the dioxid represents also 0.0017 gm. of absolute H2O2, or 0.0008 gm. of available oxygen.

Thus if from i mil of the solution of hydrogen dioxid 20 mils of gas are evolved, it is a so-called ten-volume solution and contains 0.0017X20=0.034 gm. of absolute H2O2, or 0.0008X20 = 0.016 gm. of available oxygen.

According to Naylor and Dyer (Trans. Brit. Ph. Conf., 1901, 339) the gasometric permanganate method is unreliable, because under the conditions of the test sulphuric acid added to the brine solution in the nitrometer naturally liberates a little hydrochloric acid, and this in the presence of perman- ganate becomes to some extent decomposed into chlorin. It is the uncertainty as to the extent to which the chlorin is absorbed by the water, which renders the accuracy of the method doubtful. The results of this method are uniformly too high, whether the gas be collected over mercury, over saturated magnesium sulphate, or over brine, and in the latter case considerably higher. But when the dichromate V.S. is used (without acid), closely concordant results are obtained, whether the gas be collected over mercury, or the other solutions. The evolution of oxygen by the latter method is slower than when permanganate is used, but the oxygen obtained represents the volume available in the sample.

In the Hypochlorite Method, the nitrometer is filled with a saturated solution of sodium chlorid. Two mils of the hydro- gen dioxid are admitted into the measuring tube, the funnel tube filled with a little water, and this also let in, then 20 mils of the chlorinated lime solution introduced. From this point the procedure is the same as in the gasometric permanganate method.

Ca(ClO)2 + 2H2O2 = O4 + CaCl2 + 2H2O.

The presence of preservatives, except inorganic ones, gives low results.

326 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

The Hypobromite Method. W. M. Dehn (J. A. C. S., xxix (9), 1315) describes an accurate and rapid determination of hydrogen dioxid by means of sodium hypobromite, using a ureometer. The reaction involved is

H2O2 + NaBrO = NaBr + H2O + O2.

The apparatus is shown in Fig. 58. The following descrip- tion of the method is by Dehn from the Journal of the American Chemical Society.. " The stop-cock E is opened and the stop- cocks D and F are closed; then the solution of sodium hypobromite * is poured in at the top of C until it fills the tubes A and C to some point above the stop-cock E.

" The stop-cock E is then closed and the stop-cock F is opened, so that the hypobro- mite in C may run down to the constricted portion; the solution in A is then sustained by atmospheric pressure. The stop-cock D (arranged so as to deliver only in the two directions of a right-angle triangle) is turned from the position shown in the figure and is so controlled that B may first be washed with a little of the hydrogen peroxid and then be filled with the same to a readable height on the scale. Upon turning D so as to admit a regulated volume of the hydrogen peroxid solution an immediate evaluation of oxygen results. After admitting most of the hypo- bromite held above E, and letting stand for a minute or two so as to drain properly, the columns of hypobromite in A

FIG. 58.

* This solution is prepared as directed under Estimation of Urea, except that it is finally diluted with an equal volume of water.

HYDROGEN DIOXID

327

and C are brought to the same level, the volume of oxygen is read off and its weight and that of the corresponding hydrogen dioxid are calculated by the usual formulas."

The author of this method also gives the following table, by means of which the mils of oxygen, under various conditions of temperature and pressure, may be calculated into milli- grams of hydrogen dioxid, and claims that by the use of this instrument, this hypobromite method and the table for calcu- lating the assay of hydrogen dioxid, is not only rapid and accurate, but the necessity of preparing and correcting standard solutions is avoided and the presence of the usual preserva- tives used in the dioxid solution may be ignored.

WEIGHT IN MILLIGRAMS OF H2O2 CORRESPONDING TO ONE CUBIC CENTIMETER OF MOIST OXYGEN

t/mm.

728

732

736

740

744

748

4°

1 . 2664

I-2734

.2802

1.2872

.2942

.3011

8

.2463

I-253I

.2600

.2669

.2736

.2805

12

.2251

1.2317

-2387

-2454

.2522

.2589

16

.2044

I.2III

.2178

.2245

.2311

.2378

20

.1817

I.I884

.1948

.2015

.2080

•2145

24

-1583

1.1649

.1719

.1777

.1843

.1907

28

•J345

I.I4II

.1476

-1538

.1603

.1665

32

.1085

I. 1149

.1213

-1275

-1338

.1401

36

.0843

1.0905

.0967

.1030

. 1093

-1155

40

.0605

I. 0666

1.0725

.0786

.0849

.0909

t/mm.

752

756

760

764

768

4°

.3081

-3I51

1.3222

1.3290

-3359

8

.2876

,2944

.3014

1.3081

•3*5°

12

.2657

.2726

.2823

.2860

.'2928

16

.2444

.2512

-2578

.2946

.2713

20

.2213

.2279

-2345

.2410

-2475

24

.1972

.2036

.2100

.2169

.2230

28

•T73I

.1796

-1857

.1922

.1986

32

.1465

.1528

.1589

.1562

-1715

36

.1214

.1279

1.1341

.1402

.1465

40

.0971

-I033

1.1094

-"55

.1216

CHAPTER XXI

ESTIMATION OF SOLUBLE CARBONATES BY THE USE OF THE NITROMETER

THE nitrometer may be used for estimating ammonium carbonate in aromatic spirit of ammonia.

The nitrometer in this case must be charged with mercury, as the liberated CO2 is soluble in aqueous liquids.

A given volume of the spirit is introduced into the nitrometer followed by an excess of dilute HC1, and the evolved gas then read off; and from its quantity the proportion of ammonium carbonate may be calculated by applying the equation

(NH4) 2CO3 + 2HC1 = 2NH4C1 + H2O -f CO2.

*96 *44

The volume of gas liberated must first be reduced to its corresponding volume at o° C.

Each mil of CO2 at o° C. weighs 0.001966 gm. Now if 44 gms. of CO2 represent 96 gms. of normal ammonium car- bonate, how much ammonium carbonate does 0.001966 gm. of CO2 represent?

44:96: '.0.001966 :oc. #=0.004289 gm.

Thus each mil of CO2 at normal pressure and o represents 0.004289 gm. of (NH4)2CO3, approximately.

* The atomic weights are approximate.

328

CHAPTER XXII ESTIMATION OF UREA IN URINE

THE determination of urea is based upon the fact that when urea is decomposed by an alkaline hypochlorite or hypobromite, carbon dioxid and nitrogen are given off, as the equation shows:

The liberated N may be measured, and from its quantity the amount of urea calculated. The other products of the decomposition go into solution.

The hypobromite solution is prepared as follows: 100 gms. NaOH are dissolved in 250 mils of water, and when this solution has become cold 25 mils of bromin are added, and the solution kept cold. This solution contains sodium hypo- bromite, bromate and hydroxid. It readily undergoes decom- position, and should therefore always be freshly prepared when wanted for use. To 15 mils of the NaOH solution add i mil of bromin.

The solution of sodium hypochlorite is generally preferred to the hypobromite, because it is more stable, just as effica- cious, and the disagreeable handling of bromin is obviated.

Various forms of apparatus have been devised for the quantitative estimation of urea.

The Doremus Ureometer (Fig. 59) is the simplest of these. The long arm of the ureometer is filled with the hypobromite solution, and then i mil of the urine is introduced by the aid of the pipette. The pipette is introduced through the bulb

329

330 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

as far as it will go in the bend, and the nipple is then gently but steadily compressed, being careful that no air is admitted.

The volume of the liberated gas is read off after the froth has subsided.

The ureometer indicates, according to its graduation, either milligrams of urea in i mil or grains of urea per fluid ounce of urine.

It also indicates by > the signs +, N, and -- whether the

FIG. 59.

FIG. 60.

urea is present in an increased, normal or decreased quantity. Either Knop's or Squibb's solution, or Liquor Sodae Chloratae U. S. P. may be used in this instrument. Knop's solution is that described above. Squibb's solution contains potassium bromid as well as bromin. It is prepared by taking an equal weight of bromin and of potassium bromid, and adding eight times as many cc. of water as there were grams of bromin taken. For use mix equal volumes of this solution with the sodium hydroxid solution above described.

ESTIMATION OF UREA IN URINE

331

The Hinds-Doremus Ureometer. This apparatus, which is shown in Fig. 60, is capable of giving more exact results than the original apparatus, because the i mil of urine can be delivered more accurately. It consists of a bulb with an upright tube a, graduated like the original, so that each of the smallest divisions represents o.ooi gm. of urea in the urine used. The lower portion of this tube is, in connection with a smaller tube c, graduated with a capacity of 2 mils; between these tubes a glass stop-cock is situated. Closing the stop- cock b, Knop's or Squibb's fluid (diluted one half) or liquor sodae chloratae U. S. P. is introduced into tube a so as to

FIG. 61.

completely fill it. The apparatus is then placed in an upright position and the smaller tube c is filled to the zero mark with urine. The stop-cock is then turned slowly so as to admit gradually i mil of the urine to tube a. After fifteen minutes the reading is taken. If the reading be 0.015 and the amount of urine taken was i mil, then multiplying by 100 gives 1.5 per cent of urea.

Squibb's Urea Apparatus (Fig. 61) is a very simple appa- ratus, and can be easily improvised in a drug store. It con- sists of two wide-mouthed bottles, the larger of which, C, capable of holding about 250 mils, is fitted with a rubber stopper, through which is passed a curved delivery tube and a short straight tube, the latter connected by a piece of rubber tubing

332 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

to the short glass tube in the rubber stopper of the smaller or generating-bottle B. In the generating-bottle is a small test-tube A .

Into the test-tube A is placed 5 mils of urine, and into the smaller bottle B is put 20 mils of the hypobromite solution, or strong liquor sodae chloratae. The test-tube is then placed in the generating-bottle B, being careful that the urine and the reagent do not come in contact. The larger bottle C is now filled with water and the two bottles connected by the rubber tube, the larger bottle being placed on its side upon a block, and when all connections are tight, the generating- bottle is shaken so that the urine will mix with the reagent.

Decomposition takes place, and the generated gas passes into the bottle C, displacing water, which is caught in a grad- uated cylinder or other measuring vessel. The volume of water displaced is equivalent to the volume of gas evolved.

Each mil of nitrogen gas evolved at o° C. and normal pressure represents 0.00268 gm. of urea. Then by multi- plying the number of mils evolved by this number the quantity of urea in the 5 mils of urine taken is ascertained.

The volume of gas obtained when the operation is con- ducted at ordinary temperatures should always be reduced to its corresponding volume at o° C. and 760 mm.

The factor 0.00268 is found in the following manner:

loco mils of H at o° C. weigh 0.0896 gm. ; 1000 mils of N at o° C. weigh 1.255 gms.

By the equation it is seen that 60.02 gms. of urea evolve when decomposed 28.02 gms. of N.

60.02 28.02

ESTIMATION OF UREA IN URINE 333

Now we will find the volume occupied by 28.02 gms. of N at o° C.

1.255 gms. of N = iooo mils.

gms. mils gms. mils

1.255 : 10001:28.02 : x. x = 22326.

Thus 60.02 gms. of urea evolve 22326 mils of N. One mil of N thus represents 0.00268 gm. of urea.

APPENDIX DESCRIPTION OF INDICATORS

THE following list includes the more reliable indicators in common use, arranged alphabetically.

AliVflrirr alkalies = £«/

Alizarin. adds =Yellow

This dye was first found in the roots of madder (Rubia tinctorium) but is now also obtained synthetically. A half per cent solution in alcohol is employed as an indicator.

Azolitmln:

This is the color principle to which litmus owes its value as an indicator. Its extraction is explained under litmus. It is a high-priced article and is in consequence seldom used, the purified litmus tincture being preferred.

Brazil Wood Solution:

Boil 50 gms. of finely-cut Brazil-wood (the heart-wood of Peltophorum dubium (Sprengel) Briton, nat. ord. Leguminosce) with 250 mils of water during half an hour, replacing from time to time. Allow the mixture to cool; strain; wash the contents of the strainer with water until 100 mils of strained liquid are obtained; add 25 mils of alcohol and filter.

This indicator is especially serviceable for the titration of

334

DESCRIPTION OF INDICATORS 335

alkaloids, but it is useless in the presence of sulphurous acid, sulphites, or sulphids, as these substances decolorize the solu- tion.

alkalies = Violet

^yellowish-red

Cochineal is the dried female insect, Pseudococcus cacti Linne.

The test solution is made by macerating 3 gms. of unbroken cochineal for four days in 250 mils of a mixture of i part of alcohol and 3 parts of water, by volume. It should be neu- tralized with ammonia water before using.

It is a very valuable indicator, especially for carbonates of the alkalies and alkaline earths, because it is not affected by CO2. It is also useful for titrating alkaloids, alkalies, alkali earths, ammonia, and inorganic acids, but is useless for most organic acids.

Congo Red:

Congo red is a sodium tetrazo-diphenyl-naphtionate. It occurs in commerce in the form of lumps of a reddish-brown color. It is readily soluble in both water and alcohol. Its solution is exceedingly sensitive to free acids even in the presence of acid salts, and is likewise very sensitive to free carbonic and acetic acids. It may be employed for estimating free mineral acids, in the presence of most organic acids.

The test solution contains i per cent of the dye and 10 per cent of alcohol.

rs o llpin • alkali68 = Bright red acids = Pale brown

Anthracene violet or pyrogallol-phthalein was proposed by M. Dechan for use as an indicator.

It is prepared by heating a mixture of one part of phthalic

336 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

anhydrid and two parts of pyrogallol, and finally recrys- tallizing in a similar way to phenolphthalein.

It is described as a dark reddish crystalline solid, possessing a greenish luster. It is nearly insoluble in water, but readily soluble in alcohol. In commerce it is frequently found as a paste, mixed with water.

It forms a red coloration with alkalies, which is changed to yellowish-brown on addition to an acid in excess.

It is said to be more delicate toward alkalies than phenol- phthalein, and may be used in its stead for titrating many of the alkaloids. It may be used in the presence of ammonia or its salts. It indicates sharply with the organic acids. A solution in rectified spirit i-iooo is generally employed.

H»MTmtnYv1in • alkalies= Violet

naematoxyim . adds = Yellow to orangg

A peculiar principle obtained from logwood, having the composition CieH^Oe, and crystallizing with one or three molecules of water. It is an efflorescent yellowish-rose colored substance, but when pure is said to be colorless, reddening on exposure to light.

It is soluble in hot water or alcohol. Its alcoholic solution is largely used as an indicator in the titration of alkaloids, for which it is considered the indicator par excellence.

The solution is prepared by dissolving 0.3 gm. of the well- crystallized material in 100 mils of alcohol. In titrating use about three drops of this solution.

Its color reaction with carbonates and bicarbonates is interesting. When added to a solution of alkali-bicarbonate the reaction requires many seconds, and results in a gradually deepening carmine-red which is permanent, while in the case of soluble carbonates the reaction is instantaneous, a purple- red which changes rapidly through cherry, eosin-red to orange.

DESCRIPTION OF INDICATORS 337

The reaction with ammonium carbonate is similar to that with bicarbonate.

lodeosin:

Tetra-iodo-fluorescein Erythrosin B. This indicator is use- ful for minute quantities of alkali, as for instance, such as may be dissolved out from glass on contact with water. It is used in connection with highly dilute standard solutions only. The iodeosin solution is made by dissolving 0.002 gm. of the indicator in 1000 mils of pure ether. Titration with this indicator is carried out by introducing 50 to 100 mils of the liquid to be titrated into a stoppered bottle and adding 10 to 20 mils of the ethereal indicator solution and setting aside after shaking. The ethereal layer as well as the fluid will be colorless if the latter is neutral, but if traces of alkali are present the rose-red tint passes into the aqueous liquid leaving the ether colorless. If the fluid is acid, the ethereal layer is yellow. If preferred, 4 or 5 drops of a 1-10,000 alcoholic solution of the indicator may be added to the liquid, and ether then added. lodeosin is particularly useful in titration of alkaloids, especially those of weak basicity, as emetine, atropine, morphine, etc.

alkalies

adds =Red

Lacmoid is somewhat allied to litmus, but differs from it in many respects. It is a product of resorcin, and may be prepared by heating gradually to 110° C. a mixture of 100 parts of resorcin, 5 parts of sodium nitrite, and 5 parts of water. After the violent reaction moderates it is heated to 120° C. until ammonia ceases to be evolved. The residue is then dissolved in warm water and the lacmoid precipitated therefrom by HC1; the free acid is then removed by washing and the residue dried.

338 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

This constitutes the commercial lacmoid, which is not in a sufficiently pure state to be used as an indicator; its puri- fication is effected according to Forster by treating the powder with boiling water and acidulating the resulting blue solution with hydrochloric acid. After a few hours the precipitate is collected and washed with a little cold water and carefully dried, or it is dissolved in alcohol and the solution evaporated. Even after careful purification, lacmoid solution may still exhibit a violet tinge, which is a disturbing factor in accurate work. To remedy this defect, Forster suggests the addition of 5 gms. of beta-naphthol green to 3 gms. of purified lacmoid dissolved in 700 mils of water and 300 mils of alcohol. A sample of lacmoid which is only sparingly soluble in water should be rejected. The purer the article the more readily does it dissolve in water. A good T.S. may be made by dissolving i part in 12 parts of 20 per cent alcohol, filtering and evaporating in vacuo over sulphuric acid. Of the residue obtained in this way 0.2 gm. is dissolved in 100 mils of alcohol.

Lacmoid paper is prepared by dipping slips of calendered unsized paper into the blue or red solution and drying them.

Lacmoid is slightly affected by carbonic-acid gas. It may be used cold for the alkaline and earthy hydroxids, arsenites, and borates and the mineral acids. The carbonates and bicarbonates of the alkalies and alkali earths are titrated hot with this indicator.

Many of the metallic salts, such as the sulphates and chlorids of iron, copper and zinc, which are more or less acid to litmus, are neutral to lacmoid; therefore free acids in such solutions may be estimated by its aid.

Lacmoid paper reacts alkaline with the chromates of potas- sium or sodium, but neutral with the dichromates, so that a mixture of the two or of chromic acid and dichromate may be titrated by its aid.

DESCRIPTION OF INDICATORS 339

Litmus (Lacmus):

A pigment obtained by the fermentation of certain lichens, principally from Roccdla tinctoria and R. fuciformis, but also from other species of lichen.

It occurs in commerce in small, friable, light cakes or cubes, of a violet color.

The coloring principles of litmus are azolitmin, erythro- litmin, and erythrolein. The first, which is the most important, is soluble in water, but insoluble in alcohol. The other two are readily soluble in alcohol, but only sparingly soluble in water.

The process for making litmus test solution consists in exhausting coarsely powdered litmus with boiling alcohol.

The residue is then digested with about an equal weight of cold water so as to dissolve the excess of alkali present.

The blue solution thus obtained, after being acidulated, may be used to make red litmus-paper. Finally, the residue is extracted with about five times its weight of boiling water and the solution filtered.

The filtrate is preserved as test solution in wide-mouthed bottles, stoppered with loose plugs of cotton to exclude dust, but to admit air.

When kept in closed vessels litmus solution gradually loses color, but this returns upon exposure to air and conse- quent absorption of oxygen.

The fermentation to which the loss of color is due may be prevented by saturating the solution with NaCl or by the addition of thymol or phenol.

The British Pharmacopoeia recommends the boiling of litmus in powder with three succcessive portions of rectified spirit, and then to digest the residue in distilled water and filter, the object of these steps in the process being to get rid of

340 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

the greater portion of erythrolitmin and erythrolein, which are soluble in alcohol. Then by treating the residue with water a larger proportion of azolitmin is dissolved, and the solution is contaminated with very little of the other two principles.

Litmus test solution should be .of such strength that 3 drops added to 50 mils of water will impart to the latter a

N

distinct color. If one drop of — acid or alkali solution be

10

added to this the color should change to red or blue.

Litmus may be used in a very large number of titrations. It is of value in the titration of most mineral acids and of a few organic acids, e.g., benzoic and oxalic. It is also useful in the titration of alkaline hydroxids when the latter are free from carbonates.

But for carbonates, bicarbonates, etc., a reliable end-reaction can only be obtained by boiling the solution during the titra- tion, in order to dispel the liberated CC>2.

Free CC>2 has an acid reaction with litmus, and interferes very much with the finding of the end-reaction.

Litmus may be used for ammonia and for borax. It is of no use for phosphoric or arsenic acid, nor for sulphurous acid, phosphates, or arsenates, because the change of tint is too gradual.

It is unsatisfactory in titrating many organic acids, e.g., tartaric and citric, but may be used for oxalic or benzoic, as before stated.

Sometimes it is required to perform a titration with litmus at night. Gas- or lamp-light is not adapted for showing the reaction satisfactorily, but by using a monochromatic light, such as the sodium flame, a very sharp line of demarcation may be found.

The operation should be conducted in a dark room, using

DESCRIPTION OF INDICATORS 341

a piece of platinum foil sprinkled with salt or a piece of pumice- stone saturated with a solution of salt, heated in a Bunsen flame.

The red color then appears perfectly colorless, while the blue appears like a mixture of ink and water.

alkalies

adds =Colorless

Chemically it is an oxy-chlor-diphenyl-quinoxalin. It was suggested as an indicator by Autenrieth.

The solution for the purpose of an indicator is prepared by dissolving i part in 100 parts of alcohol. Of this, four drops are sufficient for 50 mils of fluid to be titrated.

In sensitiveness, luteol exceeds both litmus and phenol- phthalein. It is more sensitive toward ammonia than Nessler's solution. Ten mils of a solution containing one drop of am- monia water per liter, is colored yellow immediately upon add- ing luteol, whereas with Nessler's solution it takes quite some time before a reaction is obtained.

Methyl Orange:

Former's Orange III, Tropaeolin D, Helianthin, Mandarin- orange, Para-sulpho-benzeneazo-dimethylanilin.

This is prepared by the action of diazo-sulphanilic acid upon dimethylanilin; the acid so formed is converted into a sodium or ammonium salt, purified by reprecipitation with HC1, and again converted into a sodium or ammonium salt. If prepared carefully and from the purest materials, it is a bright orange-red powder, perfectly soluble in water and slightly in alcohol; but it is often found in commerce as a dull orange-brown powder, often not completely soluble in water. Many conflicting statements have been made by opera-

342 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

tors as to the value of methyl orange as an indicator, which have tended to bring this indicator into disrepute.

Sutton has examined many specimens, but has not found any in which the impurities sensibly affected its delicate action. He claims that the common error is the use of too much indicator, and that some eyes are more sensitive to a change of tint than others.

Methyl orange is no doubt a very good indicator, but practice with it must be had in order to obtain good results. The author has found one sample which had a beautiful orange color, but which was absolutely useless as an indicator.

A. H. Allen describes the characters and tests of a good article as follows:

1. Aqueous solution, not precipitated by alkalies. (Orange I becomes red-brown; orange II brownish-red.)

2. Hot concentrated aqueous solution yields with HC1 microscopic acicular crystals of the free sulphonic acid, soon changing to small lustrous plates or prisms having a violet reflection. (Orange I gives yellow-brown color or flocculent precipitate; orange II brown-yellow precipitate.)

3. Dissolves in concentrated H2SO4 with a reddish or yellowish-brown color, which on dilution becomes fine red.

4. BaCb yields a precipitate.

5. CaCb yields no precipitate. (Orange I gives a red precipitate.)

6. Pb(C2H3O2)2 yields an orange-yellow precipitate.

7. MgSO4 in dilute solutions precipitates the coloring matter in microscopic crystals.

Methyl-orange T.S. is made by dissolving i gm. of methyl orange in icoo mils of water. Add to it carefully diluted sul- phuric acid in drops until the liquid turns red and just ceases to be transparent. Then filter.

The great value of this indicator consists in the fact that

DESCRIPTION OF INDICATORS 343

it is not affected by carbonic-acid gas, sulphureted hydrogen, or silicic, oleic, stearic and many other acids.

It answers well for ammonia, but it is useless for most of the organic acids. Phosphoric and arsenic acids are rendered neutral to methyl orange when only one third of the acid has combined with the base, the end-reaction being well denned. (Phenolphthalein indicates neutrality when two-thirds of acid are combined.)

This indicator should not be employed when titrating with standard solutions which are weaker than decinormal, nor should it be used in any hot titrations, nor in excessive quan- tities. Two drops are sufficient for 50 mils of the fluid to be titrated, or just enough to give it a faint tint.

Methyl Red:

acid =Rose red

(Paradimethylaminoazobenzene - orthocarboxylic acid) — • (CH3)2 -N -C6H4 -N = NC6H4 • COOH. Violet crystals, very slightly soluble in cold water, yielding a red solution, readily soluble in alcohol and glacial acetic acid. When dissolved in sulphuric acid it shows the same color change as methyl orange. The applications of methyl red as indicator are the same as those of methyl orange, except in the case of titration of car- bonates.

Methyl Red T.S. Dissolve 0.2 gm. of methyl red in 100 mils of alcohol. Its end-points are sharper than those of methyl orange, and it is far more sensitive for dilute solutions of weak bases, surpassing hematoxylin or iodeosin in the titration of alkaloids.

344 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

,, ,• f carbonates = Red

Phenacetolin : LU \ hydroxids = Yellow

acids = Yellow

This indicator is prepared by heating together for several hours equal molecular weights of phenol, glacial acetic acid, and sulphuric acid in a vessel provided with a reflux con- denser. The product is then thoroughly washed with water to remove excess of acid and dried for use. It is only very slightly soluble in water, but dissolves readily in alcohol, forming a greenish-brown solution.

The solution yields with alkali hydroxids a scarcely per- ceptible pale yellow, but with normal carbonates of the alkalies, sulphids, and with ammonia it gives a decided pink color; with bi carbonates a more intense pink, while with acids a golden yellow.

This indicator is useful for estimating the amount of alkali or alkali earth hydroxids in the presence of carbonate, unless the hydroxid is present in too small a quantity. Ammonia must not be present. The titration is carried out by adding the acid until a faint red color appears; this indicates that the alkali hydroxid or the lime has been neutralized. The further addition of the acid intensifies the red until the car- bonate present in the mixture is neutralized, when a golden- yellow color appears. The proportion of alkali hydroxid must be far in excess of the carbonate in order to obtain reliable results; furthermore, considerable practice is required in the use of this indicator in order to accustom the eye to the color changes.

A convenient strength of solution is i : 100 in alcohol.

DESCRIPTION OF INDICATORS 345

Phenolphthalein (C20H14O4):

Preparation. Five parte of phthalic anhydrid (C8H4O3), 10 parts of phenol (C6H6OH), and 4 parts of H2SO4 are heated together at 120° to 130° C. for several hours. The product is then boiled with water, and the residue, which consists of impure phenolphthalein, is dissolved in dilute soda solution and filtered. By neutralizing this solution the phenolphthalein is precipitated and may be purified by crys- tallization from alcohol; or the alcoholic solution may be boiled with animal charcoal, filtered, and the phenolphthalein reprecipitated by boiling water.

Uses. Phenolphthalein is a very valuable indicator; is extremely sensitive, and exhibits a well-marked and prompt change from colorless to pink, and vice versa. A few drops of the solution of the indicator show no color in neutral or acid liquids, but the faintest excess of alkali produces a sudden change to red.

It may be employed in the titration of mineral and organic acids and most alkalies, but it is not suited for the titration of ammonia or its salts. It is very sensitive to CO2, and therefore in estimating carbonates the liquid must be boiled, as with litmus. It is inapplicable for borax, except in the presence of glycerin, because the color gradually fades away as the acid is added. One great advantage which phenol- phthalein possesses is that its indications may be clearly read in many colored liquids; another is that it may be used in alcoholic liquids or in mixtures of alcohol and ether, and therefore many organic acids which are insoluble in water may be accurately titrated by its help.

Phenolphthalein T.S. is a one per cent solution in alcohol.

346 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

n ,. f carbonates = Blue

Poirrier Blue (C4B): alkah \ hydroxids -Red

acids =Blue

This indicator, which is closely allied to Gentian Blue in properties, is obtained by the action of sulphuric acid on triphenylrosanilin. It is a blue powder with a coppery luster. It dissolves in water and in alcohol, yielding blue solutions. KOH and NaOH change the color to red, but ammonia decolorizes it. It is employed as an indicator in aqueous solution 1:500. This indicator is exceedingly sensitive to acids. Borax and boric acid give a blue color; in the titra- tion of boric acid the red color does not appear until the acid is completely neutralized. This indicator is recommended for the titration of hydrocyanic acid, toward which it is especially sensitive, the alkaline cyanids are alkaline in reaction to most indicators, but C4B does not show an alkaline reaction until the HCN is completely neutralized, and a minute excess of the alkali hydroxid has been added. C4B is of the character of a weak acid and its salts are very unstable; they are decom> posed by water alone when in very great dilution, therefore the indicator must be used in sufficient quantity. The addition of a few drops of alcohol facilitates the color change, which is indeed a very sharp one.

P e«wi 7 n r i n • alkalies Kesazuntn. adds =Red

This is a new indicator for alkalimetry, proposed by Crismer. It is prepared as follows: Dissolve 4 gms. of resorcin in 300 mils of anhydrous ether and add 40 to 45 drops of nitric acid (sp.gr. 1.25) saturated with nitrous anhydrid. Allow the mixture to stand in a cold place for two days, whereupon a deposit of blackish crystals, having a reddish-brown reflec- tion, will be formed in the bottom of the vessel. The super-

DESCRIPTION OF INDICATORS 347

natant clear red liquid is decanted and the crystals washed with ether until the washings show a blue color with ammonia water.

Resazurin (Ci2H7NO4) is slightly soluble in water, more so in alcohol and freely soluble in acetic ether. It produces a blue solution with water, alkalies, and alkali carbonates, which are turned red upon the addition of a slight excess of acid. To use this indicator in alkalimetry, Crismer recom- mends the following solution: Resazurin 0.2 gm. dissolved

N in 40 mils of — ammonia solution, and made up to 1000 mils

with distilled water.

This is deep blue in color and keeps well. Two or three drops are sufficient to color 200 mils of liquid.

This indicator is not suited for the titration of nitric acid or monobasic organic acids, and it is not very sensitve to carbonic acid. It is, however, extremely sensitive to alkalies. If the solution is acidulated to a rose-red color and heated in a white glass flask, the solution will turn blue through the alkaline reaction of the dissolved glass before the boiling- point is reached.

This indicator is especially useful for borax.

Rosolic Acid (C20H1403):

This compound is also called Aurin and Coralline, and is prepared as follows:

A mixture of phenol and sulphuric acid is placed upon a water-bath, and oxalic gradually added, waiting each time till the evolution of gas ceases, and using less oxalic acid than is required to attack all the phenol.

In this process the oxalic acid is decomposed into CO,

348 THE ESSENTIALS OF VOLUMETRIC ANALYSIS

CO2, and H2O. The CO immediately reacts with the phenol and forms rosolic acid, as the following equation shows :

3C6H5OH + 2CO = C2oH14O3 + 2H2O.

Commercial rosolic acid is a mixture of several derivatives^ among them the above, methylaurin C20Hi6O3 and others. Commercial pceonin (also known as Aurin R.) [chiefly Ci9Hi4O3] may be used in place of rosolic acid.

Rosolic acid is soluble in diluted alcohol. Its color is pale yellow, not changed by acid, but turns violet-red with alkalies.

It is an excellent indicator for the mineral acids and strong bases, weak ammoniacal solutions, oxalic acid and other or- ganic acids, except acetic.

The test solution is made. by dissolving i gm. of the com- mercial rosolic acid in 10 mils of diluted alcohol and then adding enough water to make 100 mils.

Trop^olin (OO):

This is used in the form of a solution containing 0.5 gm. to 1000 mils of alcohol.

Turmeric Tincture. Digest any convenient quantity of ground curcuma-root (from Curcuma longa Linne, nat. ord. Scitaminece) repeatedly with small quantities of water, and throw this liquid away. Then digest the dried residue for several days with six times its weight of alcohol and filter.

Turmeric Paper. Impregnate white, unsized paper with the tincture and dry it.

The color principle of turmeric is curcumin. It is seldom used in volumetric analysis, except in the form of turmeric paper. For high-colored solutions curcumin gives no reaction

DESCRIPTION OF INDICATORS 349

with acids, but becomes brown with alkalies. There is another color principle in turmeric besides curcumin, which is, however, useless in that it is indifferent to alkalies; it is soluble in water, and extracted by digestion with water, after which the cur- cumin is dissolved out with alcohol.

Turmeric paper is especially useful, because of its peculiar reaction with boric acid, with which it develops a brown color after drying, and which color, when touched with caustic soda solution is changed to dark green.

INDEX

PAGE

Acetate, lead 169

— potassium 85

— sodium 86

Acetic acid 108

Acetone 306

Acid, acetic 108

— arsenous 197

decinormal V.S 241

solution of 198

V.S., use of in reduction 240

— defined 58

^- benzoic 109

— boric 108

— chromic 176, 229, 237

— citric 109

— hydriodic 1 23

— hydriodic by sulphocyanate method 125

— hydriodic, syrup of 126

— hydrobromic 122

by sulphocyanate method 122

by Volhard's Method 122

using chromate as indicator 125

— hydrochloric 104

normal 63

action of , on permanganate 148

standardization of 67

standardization by sodium carbonate 65

— hydrocyanic 133

using chromate indicator 132

potassium iodid indicator 133

— hypophosphorous .» 107, 164

— lactic no

— nitric 107, 321

351

352 INDEX

PAGE

Acid, nitrous 163

— number of resins 290

— oxalic 159

and oxalates 159

decinormal 63

— phosphoric 105, 1 28

— rosolic 1 8, 339

— salicylic. 109

— sulphuric 104

normal 67

— sulphurous 201

— tartaric 109

— trichloracetic 109

— value 277

-of fats and oils 279

Acidimetry 95

— and alkalimetry 59

Acids, estimation of, by neutralization 95

— haloid 1 24

— inorganic 102

— organic 108

— quantity to be taken for assay 104

— weighing of, for assay 101

volatile 103

Alcohol, in tinctures and beverages 308

Alcoholometric table , 309

Alizarin 334

Alkali bicarbonates and carbonates mixed 78

— carbonates 71

— hydroxid and carbonate mixed 77

— hydroxid, standard solution, preservation of 97

— hydroxids, estimation 67

— iodids 230

— standard solutions, preparation of 97

Alkali earth hydroxids 90

and carbonates mixed 93

salts 90, 94

— earths, organic salts of 81

— organic salts of 80

Alkalimetry 61

— and acidimetry 59

Alkaloids, estimation of , 251

INDEX 353

PAGE

Alkaloids, extraction of 259

— separation of 260

Ammonia, aromatic spirits 328

— water 7°

stronger 71

Ammonium, benzoate 89

— carbonate 74

— salicylate 89

Amyl nitrite 320, 322

Anions 19

Anthracene violet 335

Antimonic compounds 200, 238

Antimonous oxid 199

Antimony compounds 199

— and potassium tartrate 200

Apparatus, cleaning of 4°

— used in volumetric analysis 29

— use of 40

Arsenates 238

Arsenic oxid 238

Arsenic trioxid 197

Arsenite of potassium solution 198

Arsenous acid 197

solution 198

decinormal V.S 240

V.S., use of in reduction 240

— anhydrid 197

— compounds 196

direct percentage assay of 198

— iodid 198

— oxid 196

standardization of iodin with 191

Atomic weights xii

multiples of xiii

Azolitmin 334

Barium chlorid 112

— dioxid 158, 162

— hydroxid V.S 101

— peroxid 162

— soluble salts of 90

Benzoate, ammonium 89

354 INDEX

PAGE

Benzoate, sodium 86

Benzoic acid 109

Berzelius' system of oxids 56

Bicarbonate of sodium 74

— of potassium 73

Bisulphite, sodium 205

Bitartrate, potassium 85

Bleaching powder 219. 244

Borax 79

Boric acid 108

Boyle's Law 315

Brazil wood 334

test solution 18, 334

Bromates 234

Bromate, potassium 235

Bromids 119, 122

Bromin free 217

— V.S 271

— water 222

Burette, automatic 31

— connected with reservoir 32

— clamps 34

— glass stop-cock 30

— holder 34

— Mohr's 29

Burettes, special forms of 45

Burette supports 34

Butter, examination of 281

Cacodylate, sodium 79

Calcium carbonate 91, 166

— chlorid 93

— lactate 87

— salts 166

— soluble salts of 166

Calculating results 50

Calibration of instruments , 46

Calx chlorinata 219

Cane sugar, inverted 292

Carbonate, ammonium 74

— and hydroxid of alkali mixed 77

— calcium 163

INDEX 355

PAGE

Carbonate, of lithium 74

— potassium 72

— of sodium (anhydrous) 74

(crystallized) 73

— sodium, normal V.S o,2

Carbonates and bicarbonates of alkalies, mixed 78

Carbonates and hydroxids of alkali earths, mixed 93

— of alkalies 71

— soluble, assay of by the use of the nitrometer 328

Cathions 19

C4B.; 345

Centinormal solutions 10

Charles' Law 315

Chloral hydrate 94

Chlorate potassium 175, 235

Chlorates ; 175, 234

Chlorid, calcium 93

— ferric 237

— of lime 219

— sodium 116, 120, 123

V.S 116

Chlorids 120, 123

Chlorin, in bleaching powder 219, 244

— in chlorin water 21 7, 237, 243

~ free 217, 237

— water 217, 237, 243

Chlorinated lime 219, 238, 244

— soda solution 221

Chlorometry , 240

Chromates 176, 229, 237

Chromic acid 176, 229, 237

— anhydrid 176, 229

— oxids 176, 229

Chromium trioxid 229

Chromophoric theory 21

Citrate lithium 86

— potassium 85

Citric acid 109

Cochineal , 18, 335

Coefficients for caclulating analyses 53

Cold way, titration 72

Congo red 335

356 INDEX

PAGE

Copper assay 238

Cream of tartar 83

Cyanid potassium 134

Cyanogen 131

Cylinder, graduated 39

Deciriormal solutions 10

Diastasic value of malt extract 297

Bichromate, potassium 229, 237

analysis by means of 179

preparation of V.S 180

Digestion methods 232

Dioxid, hydrogen 158, 222, 323

— manganese 170, 227, 245

Direct percentage estimations 50

table of quantities for in

Dissociation theory 19

Distillation methods 223

Double normal solutions u

Eau de Javelle 221

Elements, list of xii

Empirical permanganate solutions, use of 150

— solutions ii

End-reaction 17, 113

Eosin 18

Erdmann's float 44

Erythrosin B 337

Factors 53

Fats, waxes and oils 277

Fehling's solution 291

end-point 294

Ferric alum solution 122

— chlorid 237

— salts 209, 233, 237

Ferrous ammonium sulphate 146

— carbonate saccharated 184

— iodid, syrup of 127

— salts, estimation by dichromate 182

— sulphate : 154, 186

Ferrum reductum 155

INDEX 357

PAGE

Flask, liter 38

— measuring 38

Flasks, titration 39

Fluorescein 19

Formaldehyde 299

Free fatty acids 277

Galenical preparations 265

Gallein 335

Gasometric analyses 313

Gay-Lussac's method for haloid salts 120

General methods of assaying drugs 263

General principles 4

Glucose 292

Glycerophosphate, sodium 79

Gordin's modified alkalimetric assay . 257

Gravimetric method, tne. . .,^<\. i

Grethan's pipette \t\. 103

Grouvelle's bleaching fluid 1 221

Haematoxylin 336

Halogens, free 242

Haloid acids 1 24

— salts 118

estimation of with chromate as an indicator 118

— Mohr's method 118

Hanus' number 288

Helianthin 341

Hot way titrations 71

HiibPs number 285

Hydriodic acid 125

— syrup 126

Hydrobromic acid 1 22, 125

Hydrochloric acid 104

action of, on permanganate 148

normal 63

standardization of 65

Hydrocyanic acid 133

Hydrogen dioxid 158, 222, 323

Hydrogen sulphid ; 206

Hydroxid and carbonate of alkali mixed 77

of alkali earths, mixed 93

358 INDEX

.

PAGE

Hydroxid, potassium 69

— potassium normal V.S 97

— sodium 70

normal V.S 100

Hydroxids, alkali, estimation 67

— of alkali earths 90

— • sodium and potassium mixed 78

Hypobromite solution for urea estimation 329

Hypochlorite 221

Hypophosphite, calcium 130

Hypophosphites 130, 164

Hypophosphorous acid 107, 126, 164

Hyposulphite, sodium 206

V.S. preparation of 210

Immiscible solvents 260

Indicator 17, 61

Indicators, classification of 25

• — description of individual 334

• — guide for the selection of 27

Indicator, requirements of a good 27

— sensitiveness to alkaloids 270

— theories of ' 21

Indirect oxidation, analysis by 187

Inorganic acids 102

Instruments, calibration of . , 46

Introduction i

Inverted cane sugar 292

lodates 233. 234

lodeosin 337

lodid, arsenous 198

— ferrous syrup of 127

— potassium 123

lodids 117, 119

— alkali 230

lodin, free 215

— purification of 189

— tincture 213

— titrations, use of sodium bicarbonate in . 194

Iodine absorption number 285, 288

— V.S. preparation of 189

Iodized starch test paper 242

INDEX 359

PAGE

lodometry 209

lodometric estimations, indirect 216

lonization theory 19

Ions 19

Iron, estimation of by stannous chlorid 247

— reduced 152

Javelle's water 221

Katz's method 267

Kebler's Keller method 263

Kingzett's method 222

Knop's hypobromite solution 330

Kappeschaar's solution 271

Kottstorfer number 278

Labarraque's solution 221

Lacmoid • 337

Lacmus 339

Lactate, calcium 87

Lactic acid no

Lactose 292

Law of Boyle 3*5

— of Charles 315

Lead acetate 169

— peroxid 233

— salts . 168

— subacetate 168

Lime, chlorid of 219

— chlorinated 219

Liter flasks 38

Lithium carbonate 74

— citrate 86

— organic salts of 81

Litmus 339

— tincture of 18, 339

Lloyd's method 265

Lugol's solution 216

Lunge's pipette 103

Luteol 34i

Magnesium carbonate and hydroxid 93

— sulphate, use of in permanganate titrations 149

360 INDEX

PAGE

Malt extract, diastasic value of 297

maltose in 295

Maltose 292, 295

Mandarin-orange 341

Manganates 209, 223

Manganese dioxid 1 70, 227, 245

Manganous sulphate, use of in permanganate titrations 149

Measuring flask 38

Meniscus 43

Mercuric salts 136, 239, 249

— sulphate, use of, in permanganate titrations 149

Mercurous salts 239

Methyl orange 18, 341

— red 343

Mil 14

Milliliter 14

Mohr's burette 29

— salt 149

Multiples of atomic weights xiii

Neutralization analysis 58

— of acids 95

Nitrate, silver V.S 1 14

Nitrates, nitric acid in 321

— Pelouze method 171

Nitric acid 107. 173, 321

Nitrite, amyl 320, 322

— ethyl 318, 322

— sodium 320

Nitrites 163, 317

Nitrogen dioxid 322

Nitrometer, the 313

Nitrous acid 163

— ether 318

Normal oxalic acid 62

— solutions 7

Orange, methyl 341

Oils, fats and waxes 277

— iodin absorption number of 285, 288

— table showing iodin absorption number of , 290

INDEX 361

PAGE

Organic acids 108

— salts of alkalies 80

of the alkalies, table of factors for 87

of alkali earths 83

of lithium 83

Oxalic acid 159

and oxalates 159

— decinormal 63

normal solution 62

— standardization of permanganate by 144

Oxid, antimonous *. 198

— arsenous 197

Oxidation and reduction 139

— indirect analysis by 187

Oxy-chlor-diphenyl-quinoxalin 341

Paraformaldehyde 306

Paranitrophenol 23

Perborate, sodium 162

Percarbonates 163

Percentage rules for direct estimations 50

Permanganate, action of hydrochloric acid on 148

— potassium V.S 141

— solutions, empirical 150

— titration with in presence of hydrochloric acid 149

— typical analysis with 154

— volumetric analysis by means of 146

Peroxids 209, 223

Phenacetolin 343

— in estimating mixed alkali hydroxid and carbonate 77

Phenol 273-275

Phenolphthalein 18, 345

Phosphate, sodium. 129

Phenylsulphonates 276

Phosphoric acid 105, 128

Pinch cocks 33

Pipettes 34

Poirrier blue 345

Former's orange m 341

Potassium acetate 85

— arsenite solution 198

— bicarbonate , , , , 73

362 INDEX

Potassium bi-iodate, advantages of for standardizing V.S 100

preparation of 100

— bitartrate 85

purification of 99

— bromate 235

— bromid 122

— carbonate 72

— chlorate 175, 235

— chromate T.S 19

— citrate 85

— cyanid 134

— dichromate 229

— ferricyanid T.S 19

— hydroxid 69

normal V.S 97

alcoholic 270

— iodid 119

— permanganate V.S 141

— sodium tartrate 83

— sulphite 205

— sulphocyanate V.S 117

— tartrate 81

Precipitation analysis 113

Preparation of normal oxalic acid solution 62

— of standard acid solutions 61

Prollius' fluid 260

Puckner's method 264

Pyrogallol-phthalein 335

Ramsay's bleaching fluid 221

Reading of instruments 42

Reduced, estimation of substances readily 208

— iron 156

Reducing agents 208

— sugars 292

Reduction and oxidation analysis 139

— methods, involving the use of arsenous acid V.S 240

stannous chlorid 246

Reichert number 279

Reichert-Meissl number 280

Resazurin 346

Residual titration 13

INDEX 363

PAGE

Resins, acid number of .' 290

Resorcinol 276

Re-titration 15

Rochelle salt 83

Rosolic acid 18, 347

Rules for direct percentage estimations 50

— for finding percentage 54

Salicylate, ammonium 89

— sodium 86

— strontium 87

Salicylic acid 109

Saponification number 278

Seidlitz powder 87

Seminormal solutions. 10

Separators 261

Shaking-out process for alkaloids 260

Silver alloys , 135

— metallic 135

Silver nitrate, assay of by means of sodium chlorid V.S 134

— sulphocyanate 135

- V.S 114

— salts 134

Soda, chlorinated 221

Sodium acetate 86

— tartrate 83

— benzoate 86

— bicarbonate 74

' use of in titrations with iodin 194

— bisulphite 206

— borate 79

— cacodylate 79

— carbonate (anhydrous) 74

(crystallized) 73

normal V.S 92

pure, how to make 92

— chlorid 120

preparation of pure 116

V,S 1 16

— • glycerophosphate 79

— hydroxid 7°

— — V.S. . ioo

364 INDEX

PAGE

Sodium hyposulphite 206

V.S., preparation of 210

— nitrite 320

— perborate 162

— phosphate 1 29

— salicylate 86

— sulphite 205

— tetrazo-diphenyl-naphthonate 335

— thiosulphate 206

V.S., preparation of 210

Spirit of ammonia 71

aromatic 328

— of nitrous ether 318

Squibb's hypobromite solution 330

Standard solutions 7

— temperature 13

Stannous chlorid, estimation of iron by means of 247

solution, preparation of 247

use of in reduction methods 246

Starch 292

— as an indicator 192

— after inversion 295

— inversion by diastase 296

— iodized, test paper 242

— solution 192

Stating results 55

Strontium salicylate 87

Subacetate, lead , 169

Sugars 291

Sugar in urine 293

Sulphid, hydrogen 206

Sulphids, insoluble 207

Sulphuric acid 104

normal 67

Sulphite, potassium 205

— sodium 205

Sulphites 201

Sulphocarbolates 276

Sulphocyanate method 1 20

— potassium V.S 117

Sulphocyanate V.S., assay of silver nitrate by 134

Sulphurous acid 201, 203

INDEX 365

PAGE

Syrup of hydriodic acid 126

— of ferrous iodid 127

Table, alcoholometric 309

— for correction of volume for the temperature 42, 317

— for correction of pressure 317

— of elements xii

— of factors of organic salts of the alkalies 89

— of normal factors for acids, alkalies and alkali earths 55

— of normal factors for oxids, etc 57

— of multiples xiii

— of quantities for direct percentage estimations no

— of substances estimated by standard iodin solution 208

precipitation 137

permanganate and dichromate 187

— showing color changes of indicators 256

factors for alkaloids 256

iodin absorption number of oils 288

saponification number 273

Tartar, cream of 83

— emetic 200

Tartaric acid 109

Tartrate, antimony and potassium 200

— potassium 81

— sodium and potassium 83

Temperature, standards of 13

Test mixer 39

Tetra-iodo-fluorescein >, 337

Theories of indicators 21

Theory, chromophoric 21

— ionization of indicators 21

— of Ostwald 21

Thiosulphate sodium 206

V.S., estimations involving use of 209

preparation of 210

— standardization of iodin with 190

— V.S., standardization of by dichromate 211

of by potassium bi-iodate 213

of by iodin • 210

— of by permanganate 215

Titrate, to 14

Titrated solution ,,,,,, 7

366 INDEX

PAGE

Titration, influence of concentration of V.S 41

of rate of speed 41

of temperature 41

— residual. 15

Titer 7

Trichloracetic acid 109

Tropaeolin D 341

-(00) 348

Turmeric 18

— paper 348

— tincture 348

Urea apparatus, Squibb's 331

— eLtimation of by Doremus' ureometer 329

Ureometer, Doremus' 329

— Hinds-Doremus' , 331

Urine, sugar in 293

Valence : 12

Vegetable drugs, assaying of 259

Violet, anthracene 335

Volatile acids, weighing of 103

— fatty acid, value 279

— solvents, influence of in alkaloidal assays 270

Volhard's method for haloids 1 20

— solution 117

Volume strength of hydrogen dioxid 162

Volumetric method, the 2

— or standard solutions 7

Weighing bottle 102

Wilson's bleaching fluid 221

Zinc salts, ,,,,,,,,,,,,. 93

CD O

CV2

CQ

1

ctf

o

•H

JH -P 0

-H O H >

H

3*8

0

cti

^ 0

CH W

O TJ CO bO

CQ <D LQ LQ

to to

O CO

University of Toronto

Library

DO NOT

REMOVE

THE

CARD

FROM

THIS

POCKET

Acme Library Card Pocket

Uader Pat. "RJ. Index File" Made by LIBRARY BUREAU

.

.

* *

.

•*' 4i **p ".f ~t •*&