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TEXT -BOOK  OF  CHEMISTRY 


STUDENTS  OF  MEDICINE, 


W.  L.  GOODWIN,  D.Sc.  (EoiN.), 

QUEEN'S  UNIVERSITY,  KINGSTON. 


TORONTO : 
THE  COPP,  CLARK  CO.,  LIMITED,  PRINTERS,  COLBORNE  STREET. 

1887. 


Entered  according  to  Act  of  the  Parliament  of  Canada,  in  the  year  one 
thousand  eight  hundred  and  tighty-seven,  by  THB  COPF,  CLARK  Co., 
LIMITED,  in  the  Office  of  the  Minister  of  Agriculture. 


PREFACE. 


Tin's  book  was  written  to  meet  a  want  in  my  own  teaching, 
and  is  intended  to  give  an  outline  of  Chemistry,  from  the  medical 
point  of  view,  as  far  as  possible.  While  no  attempt  has  been 
made  to  treat  the  theory  of  chemistry  exhaustively,  I  have 
tried  to  explain  and  apply  the  fundamental  laws  and  principles 
of  the  subject  in  such  a  way  as  to  render  the  mastering  of  the 
facts  an  easier  task.  The  writers  of  text-books  of  chemistry 
have  perhaps,  in  many  cases,  erred  in  arranging  their  matter  too 
logically,  e.g.,  adopting  an  invariable  order  in  describing  the 
compounds  of  the  metals  ;  while,  in  other  cases,  the  error  has 
been  in  the  opposite  direction.  It  has  been  my  aim  to  steer  a 
course  between  these  two.  The  principle  adopted  has  been  to 
proceed  from  the  known  to  the  unknown,  as  far  as  is  consistent 
with  the  limits  of  time  and  space  imposed  upon  the  teacher. 
Thus,  water  is  studied  before  oxygen  and  hydrogen  are  introduced, 
the  study  of  air  precedes  that  of  nitrogen,  &c.,  &c.  I  have 
found  it  advantageous  in  teaching  chemistry  to  medical  students 
to  let  them  make  for  themselves  such  experiments  as  those  de- 
scribed in  the  text  of  this  book.  Nearly  all  the  experiments  de- 
scribed are  such  as  can  be  made  by  students  easily  and  with 
very  simple  apparatus.  Class  experiments  have  a  rather  limited 
value,  and  can  be  advantageously  replaced  in  most  cases  by 
simpler  experiments  made  by  the  students  themselves.  The 
compounds  of  carbon  have  been  described  along  with  that  ele- 
ment, as  there  seems  to  be  no  longer  any  necessity  for  relegating 
them  to  the  end  of  the  book,  as  is  generally  done.  In  the  Ap- 
pendix is  a  Table  of  solubilities,  which  I  have  found  very  useful 
for  reference.  Tests  are  given  at  the  end  of  the  description  of 
each  acid,  metal,  &c.,  and  are  collected  into  analytical  tables  in 
the  concluding  chapter. 

I  wish  to  express  my  gratitude  to  my  colleagues,  Professors 
Dupuis  and  Shortt,  and  to  my  friend  Dr.  John  Waddell,  for 
valuable  assistance  in  revising  proofs. 

W.  L.  GOODWIN. 

Queen's  University,  Kingston,  Ontario. 
March  1st,  1887 


TABLE   OF  CONTENTS. 


CHAPTER  I. 

Chemical  Action. — States  of  Matter — Pure  Substances — Elu- 
triation — Filtration — Solution—  Crystallisation — Distillation 
— Sublimation — Fusion — Chemical  Decomposition  and  Com- 
bination— Elements  and  Compounds. — pp.  1-6. 

CHAPTER    II. 

Weights  and  Measures. — Specific  Weight— Hydrometer  or 
Areometer. — pp.  6-12. 

CHAPTER   III. 

Water. — Heat— Conduction — Radiation — Relation  of  Light  and 
Heat — Convection  —  Expansion  by  Heat  —  Temperature — 
Thermometers — Maximum  Density  Point  of  Water — Freez- 
ing and  Melting — Change  of  Freezing  Point  by  Pressure — 
Change  of  Volume  on  Freezing — Latent  Heat  of  Water — 
Specific  Heat — Evaporation  and  Ebullition — Latent  Heat  of 
Steam — Solution  —Freezing  Mixtures — Crystallisation—  In- 
fusion, &c. — Decomposition  of  Water — Composition  of 
Water.— pp.  12-31. 

CHAPTER  IV. 

Oxygen- — Preparation,  &c. — Combustion  in  Oxygen — Tempera- 
ture of  Ignition — Slow  Combustion—  Chemisin — Metals  and 
Non-Metals  —Table  of  the  Elements.— pp.  32-40. 

CHAPTER  V. 

Conservation  of  Matter-  —  Definite  Proportions — Combining 
Weights  —Equivalents — Multiple  Proportions — The  Atomic 
Theory — Avogadro's  Law— Combination  by  Volumes — Mole- 
cules and  Atoms — Molecular  Weight  of  Gases — Chemical 
Notation — Atomic  Weights — Chemical  Equations — Chemical 
Calculations — Ozone.— pp.  41-54. 

CHAPTER  VI. 

Hydrogen.— Hydroxides — Valence — Diffusion —  Hydrogen  Di- 
oxide— pp.  55-66. 


TABLE  OF  CONTENTS. 


CHAPTER  VII. 

Air. — Boyle's  Law — Charles'  Law — Measurement  of  Volumes  of 
Gases — Composition  of  Air — Combustion  in  Air— Respira- 
tion.—pp.  67-78. 

CHAPTER  VIII. 

Nitrogen. — Ammonia — Nitric  Acid  —  N itrates — Basicity— Salts 
— Oxides  of  Nitrogen  —  Monoxide  —  Dioxide  —  Trioxide — 
Nitrous  Acid  — Tetroxide — Pentoxide — pp.  79-97. 

CHAPTER  IX. 

The  Halogens. — Sea  Water — Chlorine — Hydrochloric  Acid — 
Chlorides — Oxides  of  Chlorine — Monoxide — Trioxide — Te- 
troxide— Oxygen  Acids  of  Chlorine — Hypochlorous  Acid — 
Chlorous  Acid — Chloric  Acid — Chlorates — Perchloric  Acid 
— Bromine — Hydrobromic  Acid — Bromine  and  Oxygen — 
Iodine — Hydriodic  Acid— Iodides — Iodine  and  Chlorine  — 
Iodine  and  Oxygen — lodic  Acid — Fluorine* — Hydrofluoric 
Acid— Fluorides. —pp.  98- 1 1 9. 

CHAPTER  X. 

The  Sulphur  Group. — Sulphur — Sulphur  Dioxide — Sulphur 
Trioxide — Oxygen  Acids  of  Sulphur — Hyposulphurous  Acid 
— Sulphurous  Acid — Sulphites — Sulphuric  Acid — Sulphates 
— Normal  and  Acid  Salts  -Fuming  Sulphuric  Acid — Thio- 
sulphuric  Acid — Thiosulphates — Hydrogen  Sulphide — Chlor- 
ides, &c.,  of  Sulphur — Selenium — Tellurium. — pp.  120-141. 


CHAPTER  XI. 

Phosphorus.— Oxides  of  Phosphorus — Pentoxide — Trioxide — 
Phosphoric  Acid — Phosphates — Phosphorous  Acid — Hypo- 
phosphorous  Acid — Hypophosphites — Phosphoretted  Hydro- 
gen— Phosphorus  and  the  Halogens. — pp.  142-154. 


CHAPTER  XII. 

Arsenic. — Trioxide — Pentoxide — Arsenious  Acid — Arsenites — 
Arsenic  Acid — Arsenates — Sulphides  of  Arsenic— Arseniu- 
retted  Hydrogen — Marsh's  Test— Arsenic  Chloride,  &c. — 
Tests  for  Arsenic. — pp.  155-163. 


*  This  element  has  been  lately  prepared  by  the  electrolysis  of  dry  hydro- 
fluoric aeid.     It  is  a  gas  having1  powerful  chemism. 


VI  TABLE    OF    CONTENTS. 

CHAPTER   XIII. 

Carbon. — Carbon  Compounds— Sources  of  Carbon  Compounds — 
Dioxide — Carbonic  Acid  —  Carbonates  —  Monoxide  —  Bisul- 
phide— Hydrocarbons — Marsh  Gas — Chloroform —  Ethylene 
— Isomerism — Acetylene — Cyanogen  Compounds — Potassic 
Ferrocyanide — Hydrocyanic  Acid- -Cyanides — Cyanic  Acid 
— Sulphocyanates  — Urea — Uric  Acid — Urates — A  Icohols — 
Methyl  Alcohol — Ethyl  Alcohol — Fermentation — Amyl  Al- 
cohol— Isomeric  Alcohols — Amines — Ethers — Aldehydes — 
Ketones — Chloral — Fatty  Acids — Formic  Acid — Acetic  Acid 
— Acetates — Butyric  Acid — Valerianic  Acid — Acids  of  Fats 
and  Oils — Glycol — Oxalic  Acid — Oxalates — Succinic  Acid — 
Glycerine— Hydroxy-acids — Lactic  Acid — Tartaric  Acid — 
Tartrates  —  Citric  Acid  —  Citrates  —  Carbohydrates  —  Sac- 
charoses— Cane  Sugar — Milk  Sugar — Malt  Sugar — Glucoses 
— Dextrose — Levulose — Amyloses  —  Starch —  Dextrin — Gly- 
cogen— Gums — Cellulose. — pp.  164-231. 

CHAPTER  XIV. 

Aromatic  Compounds. — Coal  Tar — Benzene  Series— Benzene 
— Nitrobenzene — Aniline — Carbolic  Acid — Creosote — Picric 
Acid — Benzylic  Alcohol — Benzole  Aldehyde — Benzoic  Acid 
— Saccharine — Salicylic  Acid — Gallic  Acid — Tannic  Acid— 
Terpenes  —  Camphor  —  Cinnamic  Acid  —  Essential  Oils — 
Indigo— Naphthalene — Anthracene — Glucosides — Alkaloids 
— Conine  —  Nicotine  —  Morphine  —  Quinine — Cinchonine — 
Strychnine — Cocaine — Atropine — Kairine,  Antipyrine,  Thai- 
line — Albuminoids.  — pp.  232-255. 

CHAPTER  XV. 

Silicon. — Silica — Silicic  Acid  and  Silicates — Fluosilicic  Acid- 
Boron— Boric  Acid— Borax,  —pp.  25G  -262. 

CHAPTER  XVI. 

The  Metals. — General  Characters — Ores — Alloys — Compounds 
— Oxides — Sulphides — Chlorides,  &c. — Oxygen  Salts — Clas- 
sification— Analysis.  — pp.  263-271 . 

CHAPTER  XVII. 

Metals  of  Group  I. — Lead — Oxides — Salts — Acetate — Nitrate 
— White  Lead — Chloride— Iodide — Lead  Plaster — Sulphate 
— Commercial  Preparations— Lead  Poisoning. — Silver — Ox- 
ides— Salts  —  Nitrate  —  Mercury  —  Amalgams  —  Mercurous 
Compounds  —  Nitrate  —  Chloride  —  Iodide — Mercuric  Com- 
pounds —  Nitrate  —  Sulphate  —  Chloride — Oxide — Iodide  — 
White  Precipitate  —  Sulphide  —  Mercurial  Poisoning. — pp. 
271-292. 


TABLE    OF    CONTNETS.  Vll 

CHAPTER   XVIII. 

Metals  of  Group  II. — Copper — Compounds — Cupric  Sulphate 
— Oxide — Commercial  Preparations — Cadmium — Nitrate — 
Sulphate — Iodide — Bismuth — Nitrate — Subnitrate — Trioxide 
— Bismuthyl  —  Carbonate  —  Antimony  —  Trisulphide  —  Tri- 
chloride— Trioxide — Tartar  Emetic — Tin—  Stannic  Oxide — 
Stannous  Chloride — Stannic  Chloride—  Gold — Compounds — 
Phitinum — Compounds — Palladium,  SLC. — pp.  293-315. 

CHAPTER  XIX. 

Metals  of  Group  III. — Iron — Ferrous  Salts — Sulphate — Car- 
bonate— Arsenate  —  Phosphate  —  Ferric  Salts  —  Chloride — 
Sulphate  —  Nitrate  - —  Hydroxide  —  "  Scale  "  Compoimds — 
Chromium — Potassic  Bichromate — Chromic  Acid—  Chromates 
— Chrome  Alum — Chromic  Hydroxide — Aluminium — Alu- 
mina— Alums — Aluminic  Sulphate — Porcelain,  &c. — Zinc — 
Oxide — Chloride — Sulphate — Carbonate— Acetate — Manga- 
nese— Dioxide — Manganous  Salts — Manganic  Salts — Man- 
ganates — Permanganates  —  Cobalt — Oxides — Nitrate — Chlo- 
ride— Nickel —  Oxides — Sulphate — Cerium . — pp.  31 6  -350. 

CHAPTER  XX. 

Metals  of  Groups  IV.  and  V. — Calcium — Oxide — Hydroxide 
—Carbonate  —  Chloride  —  Sulphate  —  Bleaching  Powder — 
Phosphates — Mortars  and  Cements — Strontium—  Barium — 
Oxides — Chlorides  —  Nitrate  —  Magnesium — Sulphate — Car- 
bonate— Magnesia — pp.  350-367. 

CHAPTER   XXI. 

Metals  of  Group  VI.—  Sodium — Chloride — Sulphate— Carbon- 
ate— Bicarbonate —  Hydroxide  —  Nitrate  —  Sulphite — Phos- 
phate— Bromide  —  Sulphide— Glass — Potassium — Carbonate 
— Bicarbonate — Hydroxide — Chlorate— Nitrate— Bromide — 
Iodide  —  Ammonium  - —  Sulphate  —  Chloride  —  Carbonate  — 
Phosphate-  -Microcosmic  Salt — Sulphide — Lithium — Carbon- 
ate— Rubidium — Ccesium — Spectrum  Analysis. — pp.  368-390. 

CHAPTER  XXII. 

Electricity.  —  Electrolysis  —  Electro-chemical  Series.  —  pp. 
390-392. 

CHAPTER  XXIII. 

Analysis.  —  Chemical  Toxicology  —  Analytical  Tables.  —  pp. 
393-404. 

Appendix.— Table  of  Solubilities— pp.  405-408. 
Index.— pp.  409-416. 


ADDENDA  ET  CORRIGENDA. 


Page    7,  table,  for  0.064  read  0,0648. 

"       7,  table,  for  28.3549  read  28.3495. 

"     11,  1.  8,  for  96  read  88, 

"     26,  1.  18,  for  no  read  not  much. 

"     35,  1.  3,  from  bottom,  for  is  read  was  called. 

"     48j  1.  15,  for  always  read  generally. 

"    48,  1.  15,  for  6. 6  read  6. 3. 

"     48,  Is.  19  and  20,  for  specific  read  atomic. 

11     53,  last  line,  for  HgO  read  2HgO, 

"    56,  for  Experiment  28  read  Experiment  28  A. 

"    57,  1.  3,  for  atomic  read  atom. 

"    83,  1.  16,  for  red  read  blue. 

"  122,  1.  4  from  bottom,  for  that  read  the  atom. 

"  141,  1.  8,  for  KOH  read  NaOH 

"  145,  1.  7,  for  amphorous  read  amorphous. 

"  154,  last  line,  for  Cas(POt)2  read  3(Caa(POt).2). 

"  163,  first  equation,  for  Na,2CO  read  Na2COs. 

"  179,  1.  3  from  bottom,  for  151  read  159. 

"  193,  1.  6,  after  absent,  insert,  Fermentations  are  also  caused 
by  certain  nitrogenous  organic  compounds  called  tin- 
organised  ferments,  e.g.  diastase,  synaptase,  &c. 

"  201,  1.  4  from  bottom,  after  chlorine  read  by  a  series*  of 
reactions. 

"  203,  1.  17,  for  C3H60S  read  CsffsOa. 

"  209,  1.  2,  for  a  solution  of  read  sparingly  soluble. 

"  219,  1.  15,  omit  sparingly. 


CHEMISTRY 


STUDENTS   OF   MEDICINE. 


CHAPTER     I. 


INTRODUCTORY. 

1.  Chemical  Action. — Iron  rusts  in  air,  and  the 
rust  differs  in  properties  from  the  iron.  Iron  heated  in 
air  becomes  changed  to  a  black  substance  (as  in  a  black- 
smith's shop)  unlike  the  iron  in  many  respects.  When 
placed  in  vinegar  iron  gradually  disappears  into  the 
vinegar,  i.  e.  dissolves,  and  a  red  liquid  is  formed,  having 
the  properties  of  neither  vinegar  nor  iron.  Wood  burns, 
leaving  only  a  small  portion  of  ash,  unlike  the  wood  in 
colour,  and  other  properties.  The  greater  part  of  the 
wood  has  been  changed  into  substances  like  air.  A  plant 
takes  food  from  the  soil  and  from  the  air  in  the  form  of 
water,  carbonic  acid,  ammonia,  and  various  mineral  sub- 
stances. From  these  it  elaborates  sugars,  starch,  gums, 
wood,  &c.,  substances  totally  different  from  the  original 
articles  of  food,  and  not  to  be  found  in  the  sources  from 
which  the  plant  gets  its  food.  An  animal  brings  about 
similar  changes,  forming  out  of  its  food  substances  quite 
2 


2  PURE   SUBSTANCES. 

unlike  that  food,  converting  starch  into  sugar,  albumin 
into  fat,  &c.  In  all  these  cases,  substances  undergo  such 
changes  that  they  become  converted  into  other  and  differ- 
ent substances.  Such  processes  are  called  chemical  actions, 
and  Chemistry  is,  for  the  most  part,  the  study  of  chemical 
actions. 

2.  Three  States  of  Matter. — All  substances  can 

be  grouped  into  three  classes,  viz. :  solids,  liquids,  and 
gases.  Solids  have  a  definite  form  and  volume.  Liquids 
have  no  definite  form,  but  their  volume  does  not  tend  to 
change.  Gases  are  indefinite  both  in  form  and  volume. 
They  take  the  shape  of  the  vessel  in  which  they  are  con- 
fined, and  readily  undergo  compression  and  expansion. 

3.  Pure  Substances. — Simple  inspection  of  granite 
shows  it  to  contain  more  than  one  substance.     It  is  a 
mixture  or  mixed  substance.     Soil,  pudding  stone,  milk, 
and  air,  are  other  examples  of  mixtures.    There  are  many 
methods  of  separating  mixtures  into   their  ingredients, 
and  thus  obtaining  pure  substances,  or  chemical  indivi- 
duals.   The  commoner  methods  are  given  in  the  following 
sections. 

4.  Elutriation,  or  "  washing  out,"  is  a  process  used 
by  the  gold  miner  who  washes  away  the  light  earth,  sand, 
&c.,   from  the  heavier  gold.      Winnowing  is  a  similar 
process. 

5.  Filtration  is  separating  a  liquid  from  a  solid  by 
allowing  the  former  to  flow  through  some  porous  sub- 
stance which  retains  the  latter.     Thus,  muddy  water  can 
be  separated  into  mud  and  water.     Unsized  paper,  called 
Jilter  paper,  is  very  commonly  used. 


SEPARATION    OF    MIXTURES.  6 

C.  Solution. — In  this  process,  a  liquid  called  a  solvent 
is  used  to  separate  a  soluble  substance,  sugar  for  instance, 
from  an  insoluble  substance  such  as  sand. 

7.  Crystallisation. — When  all  the  ingredients  of 
a  mixture  are  soluble  in  some  one  solvent,  e.  g.  in  water, 
a  separation  can  often  be  brought  about  by  evaporating 
the  solvent.     When  there  is  not  enough  of  the  solvent 
left  to  dissolve  all  the  substances,  that  one  which  is  most 
difficult  to  dissolve  (or  the  least  soluble),  separates  from 
the  liquid,  generally  in  regularly  formed  crystals.     At 
another  stage  of  the  evaporation,  the  substance  next  in 
solubility  crystallises,  and  thus  a  separation  more  or  less 
complete  is  effected.     It  is  by  this  process  that  common 
salt  is  separated  from  the  other  substances  dissolved  in 
sea  water. 

8.  Distillation. — If  a   dish   of  alcohol  and   one  of 
water  be  set  side  by  side  on  a  hot  stove  the  alcohol  begins 
to  boil  much  before  the  water.     Alcohol  boils  at  a  lower 
temperature  than   water,   and  a  mixture   of    these  two 
liquids  can  be  separated  by  distilling  them.     For,  when 
heat  is  applied  to  the  vessel  (the  boiler,  or  retort,)  contain- 
ing the  mixture,  the  alcohol  is  first  changed  to  vapour, 
passes  as  vapour  into  the  cold  tube  (or  condenser),  and  is 
there  cooled  and  condensed  (or  made  liquid).     Thence  it 
runs  into  the  receiver.     Later,  water  begins  to  distil  over 
and  may  be  collected  in  a  different  vessel.     In  a  similar 
way  fresh  water  can  be  prepared  from  sea  water,  the  pure 
water  distilling,  and  the  salt  remaining  in  the  boiler. 

9.  Sublimation. — This    is    distillation    applied    to 
solids  which  can  be  changed  to  gases  by  beat.     Sulphur 


4  CHEMICAL   DECOMPOSITION. 

is  purified  by  lieating  it  until  it  is  changed  to  a  gas 
which  is  then  condensed  in  a  clean  vessel.  The  impuri- 
ties remain,  since  they  are  non-volatile. 

10.  Fusion. — Some  substances  melt  or  fuse  at  lower 
temperatures    than   others.     Thus,   butter  fuses  at  the 
temperature  of  the  hand ;  while  salt  can  be  fused  only 
by  a  strong,  red  heat.     If  a  mixture  of  butter  and  salt 
be  gently  heated,  the  butter  melts,  the  salt  sinks  to  the 
bottom,  and  the  liquid  butter  can  be  poured  off.     This  is 
the    pharmaceutical    process    of   clarification.     In    some 
cases  the  impurities  are    lighter  than    the    melted  sub- 
stance to  be  clarified.     They,  then,  rise  as  a  scum  and 
are  removed  by  skimming. 

11.  Chemical  Decomposition  and  Combina- 
tion.— In  the  processes  described  above,  no  permanent 
change  is  brought  about  in  the  pure  substances  which 
are  separated  out  of  the  mixtures.     They  preserve  their 
identity.       In   chemical    actions    substances    lose    their 
identity. 

Experiment  1. — Heat  a  little  red  oxide  of  mercury  in  a 
tube  of  glass.  The  oxide  disappears  gradually,  and  a  silvery 
liquid,  quicksilver,  gathers  on  the  inside  of  the  tube.  Thrust  a 
splinter  of  wood  with  its  end  still  glowing  into  the  tube.  The 
splinter  begins  to  burn  very  brightly. 

Two  substances  have  been  formed  from  the  oxide  of 
mercury,  viz.,  the  liquid  metal  quicksilver  or  mercury, 
and  the  gas  oxygen,  which  causes  a  live  coal  to  burst  into 
bright  flame  and  therefore  differs  from  ordinary  air. 
This  is  an  example  of  chemical  decomposition.  Red  oxide 
of  mercury  has  been  decomposed  by  heat,  and  the  pro- 


ELEMENTS    AND    COMPOUNDS.  5 

ducts  of  its  decomposition  are  mercury  and  oxygen.  This 
process  differs  from  the  separation  of  a  mixture  into  its 
ingredients,  because  the  products  of  decomposition  have 
not  the  properties  of  the  substance  decomposed.  The 
chief  agents  which  bring  about  chemical  decomposition 
are  heat,  light,  electricity,  mechanical  force  (as  in  the 
explosion  of  dynamite),  and  contact  with  certain  sub- 
stances. Chemical  decompositions  are  also  brought 
about  in  some  unknown  way  by  living  beings,  as  in 
fermentation. 

Experiment  2- — Mix  well  four  parts  of  flowers  of  sulphur 
with  seven  parts  of  very  fine  iron  filings.  A  powder  is  obtained 
in  which  the  presence  of  both  iron  and  sulphur  can  be  easily 
recognized.  Shake  up  a  little  of  the  mixture  with  water  in  a 
test-tube.  The  iron  sinks  to  the  bottom  more  quickly  than  the 
sulphur  and  the  two  are  separated.  Move  a  magnet  over 
another  small  portion  of  the  mixture.  The  iron  sticks  to  the 
magnet  and  the  sulphur  remains.  The  two  substances  were 
merely  mixed.  Now,  heat  the  remainder  of  the  mixture  in  a 
a  small  porcelain  dish.  It  gets  red  hot,  blackens,  aud  becomes 
quite  uniform  in  appearance.  The  closest  examination  does  not 
show  the  presence  of  either  sulphur  or  iron.  Powder  a  little  of 
the  black  substance  and  shake  it  up  with  water  as  before.  No 
separation  takes  place.  Move  a  magnet  over  another  part.  No 
iron  sticks  to  it.  The  iron  and  sulphur  have  disappeared  and 
this  single  black  substance  has  taken  their  place. 

This  is  an  example  of  chemical  combination.  Iron  and 
sulphur  have  combined  to  form  sulphide  of  iron  (ferrous 
sulphide),  a  substance  differing  altogether  in  properties 
from  iron  or  sulphur.  Chemical  combination  must  be 
distinguished  from  mixture. 

12.  Elements  and  Compounds. — The  majority 

of  substances  can  be  decomposed,  but  there  are  certain 


b  WEIGHTS    AND    MEASURES. 

which  cannot,  so  far  as  known.  Thus,  it  has  been  found 
impossible  to  obtain  from  a  portion,  of  pure  mercury  any- 
thing else  ;  and  the  same  is  true  of  iron  and  sulphur. 
These  substances  will  combine  with  others  to  form  new 
substances,  but  from  themselves,  taken  alone,  nothing  dif- 
ferent has  been  obtained  by  any  process  yet  tried.  Such 
substances  are  called  Elements,  or  Simple  Substances. 
They  unite  to  form  compounds. 

Definitions. — A  chemical  action  or  reaction  is  a  change  in 
which  from  one  or  more  substances  there  are  formed  other  sub- 
stances differing  from  the  original  in  essential  properties,  e.y.  in 
colour,  taste,  smell,  &c.  In  chemical  combination  simpler  sub- 
stances unite  to  form  more  complicated.  In  chemical  decomposition 
complex  substances  are  broken  up  into  simpler.  An  element  is  a 
pure  substance  which  has  never  been  decomposed  into  substances 
differing  from  it  in  properties.  A  compound  is  a  substance 
formed  of  two  or  more  elements  chemically  combined. 

There  are  at  present  sixty-seven  elements  known,  and 
all  the  known  substances  on  and  in  the  earth  consist  of 
these  elements  and  their  compounds  with  each  other. 
A  list  of  the  elements  is  given  at  page  39. 


CHAPTER    II. 


WEIGHTS  AND  MEASURES— SPECIFIC  WEIGHT. 

13.  Weights  and  Measures. — In  order  to  study 

chemical  actions  completely  the  quantities  of  substances 
taking  part  in  them  must  be  known.  The  metrical  sys- 
tem of  weights  and  measures  has  been  adopted  by  chemists 
everywhere.  It  involves  very  little  calculation  since  its 


WEIGHTS    AND    MEASURES.  7 

units  all  increase  in  magnitude  by  tens,  so  that  the  opera- 
tions  of  reduction  can  be  performed  by  merely  moving 
the  decimal  point.  It  is  hence  called  the  decimal  system 
of  weights  and  measures.  The  unit  of  length  in  this 
system  is  the  metre  intended  to  be  equal  to  one  ten- 
millionth  part  of  a  quarter  of  the  earth's  circumference 
through  the  meridian  of  Paris. 

Linear  Measures. — 1000  millimetres  (mm.)  =  100 
centimetres  (cm.)  =  10  decimetres  (dcm.)  =  1  metre  (m.) 
=  ra  decametre  =  TJO  hectometre  =  r^iys  kilometre  = 
39.371  English  inches. 

Measures  Of  Capacity. — 1000  cubic  centimetres 
(c.c.)  =  1  cubic  decimetre  =  1  litre  (1.)  =  61.027  cubic 
inches  =  1.761  pint. 

Measures  Of  Weight. — 1000  milligrams  (mgms.) 
=  100  centigrams  (cgms.)  =10  decigrams  (dgms.)  =  1 
gram  (gm.)  =  •£$  decagram  =  i^  hectogram  =  T^^  kilo- 
gram =  15.43  grains. 

1  gram  =  the  weight  of  1  cubic  centimetre  of  pure 
water  measured  at  4-°  centigrade  (its  point  of  greatest 
density). 

Relations  of  Pharmaceutical  to  Metrical  Weights  and  Measures. 

CUBIC  CENTIMETRE. 

1  minim  =        0.059 

1  grain    =      0.064  gram.  1  fluid  drachm  =         3"549 

1  ounce  =    28.3549    »  ounce     =       28'396 


1  pound  =  453.5925 


1  pint  =     567.936 

1  gallon  =  4543.487 

or  about  4J  litres. 


8  SPECIFIC    WEIGHT. 

Exercises.— 1.   In  1764  millimetres  how  many  metres  ? 

2.  In  37  metres  how  many  cms.  ? 

3.  How  many  litres  capacity  has  a  tank  2  metres  long,   1.5 
metres  broad  and  1.5  metres  deep  ? 

4.  How  many  grams  in  a  pound  avoirdupois  (—  7000  grains)  ? 

5.  What  weight  of  water  will  fill  the  tank  in  (3)  ? 

6.  How  many  c.  c.  in  a  fluid  ounce  ?     (Note.  — The  fluid  ounce 
is  the  volume  of  1  ounce  weight  of  water  =  437.5  grains). 

7.  Reduce  38,674  cubic  centimetres  to  litres. 

8.  In  1  kilogram  how  many  pounds  ? 

9.  Find  the  number  of  kilometres  in  a  mile. 

10.  How  many  inches  in  a  kilometre  ? 

11.  Find  the  capacity  in  litres  of  a  tank  4  feet  long,  1\  feet 
wide  and  3  feet  deep. 

12.  What  is  the  weight  in  grams  of  8  fluid  ounces  of  mercury  ? 
(mercury  weighs  13.596  times  as  much  as  the  same  volume  of 
water). 

14.  Specific  Weight. — When  it  is  said  that  lead 
is  heavier  than  water  it  is  meant  that  if  equal  volumes 
of  lead  and  of  water  be  weighed,  the  lead  will  be  found 
to  be  heavier.  Thus  Specific  Weights,  or  Specific  Gravi- 
ties, of  substances  are  found  by  comparing  the  weights 
of  equal  volumes  of  the  substances  with  that  of  the  same 
volume  of  some  substance  chosen  as  a  standard.  For 
liquids  and  solids  water  has  been  chosen  as  the  standard ; 
and  for  gases,  air,  or  hydrogen.  For  example,  1  cubic 
centimetre  of  water  weighs  one  gram ;  and  the  same 
volume  of  mercury  weighs  13.596  grams.  Then,  the 
specific  weight  of  mercury  is  13.596. 


SPECIFIC    WEIGHT. 

Definition. — The  Specific  Weight  of  a  substance  is  a  number 
expressing  how  many  times  heavier  the  substance  is  than  an 
equal  volume  of  some  substance  chosen  as  a  standard. 

15.  Specific  Weights  of  Solids. — If  the  solid  is 

heavier  than  water  it  is  first  weighed  in  the  air,  and 
then  hanging  in  water.  Its  weight  in  water  (w]  is 
less  than  its  weight  in  air  (W)  by  the  weight  of  the 
water  which  it  displaces,  i.e.,  an  equal  volume  of  water. 

W 

Or,  if  S  represents  the  specific  weight,  then  S  =  ~  — 

16.  Specific  Weights  of  Liquids.— (l)  Fill  a 

weighed  narrow-necked  flask  with  water  up  to  a  mark 
on  its  neck,  and  weigh  the  full  flask  to  determine  the 
weight  (W)  of  water.  Fill  the  same  flask  with  the 
liquid  of  which  the  specific  weight  is  to  be  found  and 

W 

find    the    weight    ( W)   of    the   liquid.     Then    S  =  ~  ' 

(2)  The  Hydrometer  or  Areometer  is  a  glass  tube  with  a 
bulb  blown  on  one  end.  In  the  bulb  is  a  small  quantity 
of  mercury  which  causes  the  tube  to  swim  upright  in 
any  liquid  in  which  it  is  placed.  The  stem  of  the  hydro- 
meter is  graduated  and  numbered.  When  placed  in 
water  it  sinks  until  the  water  reaches  a  certain  mark  on 
the  stem.  When  placed  in  a  liquid  heavier  than  water 
it  does  not  sink  so  far.  In  a  liquid  lighter  than  water 
the  instrument  sinks  farther  than  in  water.  Numbers 
marked  on  the  stem  indicate  the  specific  weights  of  the 
liquids.  This  instrument  in  various  forms  is  constantly 
used  in  medical  practice. 

1 7.  Specific    Weights    Of    Gases.  —  These    are 
found   by   methods  the  same  in  principle  as    the   first 


10 


SPECIFIC    WEIGHT. 


method  for  liquids  (§16).  A  large  glass  globe,  whose 
capacity  is  known,  is  made  as  nearly  as  possible  empty 
of  air  by  means  of  the  air  pump,  and  is  then  weighed. 
The  gas  of  which  the  specific  weight  is  to  be  determined 
is  then  allowed  to  flow  in,  and  the  flask  is  reweighed. 
The  increase  in  weight  is  the  weight  of  the  flask  full  of 
gas.  This  weight  divided  by  that  of  an  equal  volume  of 
the  standard  gives  the  specific  weight  required.  There 
are  many  minute  precautions  and  corrections  which  can- 
not be  described  here. 


18.  In  the  following  table  are  given  the  specific 
weights  of  some  of  the  commoner  solids  and  liquids, 
water  at  4°  centigrade  being  the  standard  : — 


Platinum 22.069 

Gold 19.362 

Lead  11.352 

Silver 10.474 

Copper 8.788 

Brass 8.383 

Steel 7.820 

Iron  (wrought) 7.788 

Iron  (cast) 7.207 

Tin 7.291 

Zinc 6.861 

Diamond 3.531 

Flint  Glass 3.329 

Marble 2.840 

Bottle  Glass 2.600 

Plate  Glass 2.370 

Porcelain 2.300 

Sulphur 2.030 

Ivory 1.917 

Graphite 1.8  to     2.400 

Anthracite..    1.800 

Phosphorus 1.770 

Magnesium 1.740 

Amber 1.080 

Water  at  4°  0 1.000 

White  Wax 0.970 


Sodium 0.970 

IceatO°C 0.918 

Potassium 0.860 

Mahogany 1.060 

English  Oak 0.970 

Beech 0.852 

Ash 0.840 

Yellow  Pine 0.657 

Cork 0.240 

Mercury 13.596 

Oil  of  Vitriol 1.840 

Chloroform 1 .525 

Nitric  Acid    1.500 

Hydrochloric  Acid 1.220 

Blood  (human) 1.045 

Milk    1.030 

Sea  Water 1.028 

Port  Wine 0.990 

Castor  Oil 0.970 

Linseed 0.940 

Proof  Spirit     0.930 

Oil  of  Turpentine 0.870 

Brandy 0.837 

Absolute  Alcohol 0. 780 

Ether...  0.720 


QUESTIONS    AND    EXERCISES.  11 

QUESTIONS  AND  EXERCISES. 

1.  What  is  a  unit  of  measurement  ?     Why  are  there  many  units 
of  length  in  use  instead  of  only  one  ? 

2.  A  flask  filled  with  water  was  found  to  weigh  72  grams,  the 
flask  alone  weighing  22  grams.     The  same  flask  filled  with  oil 
of  vitriol  weighed  114  grams.     Calculate  the  sp.  wt.  of  oil  of 
vitriol. 

3.  A  piece  of  iron  weighing   96  grams  in  air  was  found  to 
weigh  76.5  grams  in  water.     Calculate  the  sp.  wt.  of  iron. 

4.  A  flask  filled  with  water  weighed  153  grams  ;  25  (j.  of  cop- 
per are  dropped  in.     The  flask  and  contents  then  weighed  175.19 
g.     What  is  the  sp.  wt.  of  copper  ? 

5.  '"'When  a  body  is  weighed  in  air  its  true  weight   is   not 
found."     Why  not  ?     How  must  it  be  weighed  in  order  to  find 
its  real  weight  ? 

6.  A  man  of  160  Ibs.  weight  immersed  his  body  completely  in 
a  bath  7  ft.  long  and  3  ft.  wide.     The  water  rose  1J  in.     What 
was  the  sp.  wt.  of  his  body  ? 

7.  A  piece  of  cork  weighing  in  air  15  grams  is  immersed  in  a 
vessel  of  water  12  cm.  long  and  5  cm.  wide.     The  water  rises  1 
cm.     Calculate  the  sp.  wt.  of  the  cork. 

8.  What  is  the  volume  in  cu.  inches  of  a  piece  of  lead  weigh- 
ing 10  Ibs.  ? 

9.  Calculate  the  volume  in  cubic  centimetres  of  100  grams  of 
copper. 

10.  What  is  the  weight  in  grams  of  a  gallon  of  pure   water 
at  4°  C. 

11.  What  is  the  sp.  wt.  of  a  body  which  floats  with  one  third 
of  its  bulk  out  of  water  ? 

12.  When  the  lungs  are  inflated  the  specific  weight  of  the 
human  body  is  less  than  1  ;  but  when  the  lungs  are  tilled  with 
water,  it  is  greater  than  one.     Explain  this. 

13.  Write  a  short  essay  on  the  convenience  of  the  metrical 

system. 


12  CONDUCTION. 


CHAPTER    III. 


WATER— HEAT—  SOLUTION. 

19.  Water. — About    two-thirds    of    the    weight    of 
animals  and  a  large  fraction  of  the  weight  of  plants  con- 
sist of  water.     As  it  occurs  in  nature  water  is  not  a  pure 
substance.     This  can  be  shown  by  distilling  any  sample 
of  sea,  river,  lake,  spring  or  rain  water.     A  solid  residue 
is  left  in  each  case.     But  if  this  distilled  water  (aqua 
distillata)  be  redistilled    again  and  again  no  residue  is 
left  in  the  retort.     Distilled  water  is  a  pure  substance,  a 
chemical   individual.     It   is  generally  a   liquid,   but   if 
sufficient  heat  be  removed  from  it,  it  becomes  solid  (ice); 
and  when  heat  is  added  liquid  water  is  changed  to  a  gas 
(water  vapour,  steam). 

20.  Heat  -  Conduction. — Heat    and    light    were 
once  thought  to  be  substances.     They  are  now  known  to 
be  motion  of  some  sort.     When  a  body  becomes  hotter  its 
particles  move  (vibrate)  faster,  and  the  motion  is  com- 
municated to  any  other  body  in  contact  with  it.     Simi- 
larly, heat  passes  from  a  hot  to  a  cooler  part  of  a  body 
without  any  movement  of  the  body  as  a  whole.     This 
process  is  called  conduction  of  heat. 

Experiment  3. — Choose  a  piece  of  copper  wire  and  another 
of  iron  wire  of  about  the  same  size  and  length.  Hold  the  end  of 
each  in  the  flame  of  a  burner,  and  observe  the  time  required  for 
the  heat  to  become  unpleasant  at  the  other  end  held  between 
the  fingers.  The  time  is  much  longer  for  the  iron  than  for  the 
copper.  Copper  is  a  better  conductor  of  heat  than  iron. 


RADIATION.  13 

A  glass  rod  may  be  heated  at  one  end  until  it  melts 
before  the  other  end  becomes  warm.  Glass  is  a  poor 
conductor  of  heat.  The  metals  generally  are  the  best 
conductors  of  heat.  Wool,  feathers,  asbestos,  fur,  air 
and  gases  generally,  and  most  liquids  are  bad  conductors. 
We  use  bad  conductors,  as  clothing,  to  keep  the  heat 
from  leaving  our  bodies  ;  but  we  also  use  them  as  pack- 
ing for  ice,  &c.,  to  keep  the  heat  out. 

Experiment  4. — Wrap  a  piece  of  copper  wire  round  a 
small  lump  of  ice  and  allow  the  ice  to  fall  to  the  bottom,  of  a 
test-tube  full  of  cold  water.  Holding  the  test-tube  aslant,  heat 
the  water  near  the  top  until  it  begins  to  boil.  The  ice  at  the 
bottom  remains  unmelted.  Water  is  a  bad  conductor  of  heat. 

21.  Radiation. — Heat  reaches  us  from  the  sun  by 
a  process  different  from  conduction.     It  may  warm  the 
earth  to  90°  or  100°  F.,  and  still  leave  the  air  through 
which  it  passes  cool.     In  this  way,  too,  heat  and  light 
pass  through  space  empty  of  everything  as  far  as  known. 
This  movement  of  heat  and  light  is  called  radiation.     It 
is  extremely  rapid — about  186,000  miles  per  second. 

22.  Relation  of  Light  and  Heat. — Light  and 

heat  are  related  to  each  other  just  as  the  high  notes  of 
music  are  related  to  the  low.  Light  consists  of  short 
waves,  or  rapid  vibrations,  which  produce  in  the  eyes  the 
sensation  ordinarily  called  light.  Heat  consists  of  the 
longer,  slower  waves  which  are  not  capable  of  exciting 
this  sensation, — in  our  eyes  at  least.  When  light  falls 
upon  a  body  which  is  not  a  good  reflector,  it  is  absorbed 
and  may  be  thus  transformed  into  heat.  It  is  in  this 
way,  principally,  that  the  earth  is  heated  by  the  sun. 


14  CONVECTION — EXPANSION. 

23.  Convection. 

Experiment  5. — Fill  a  tall  glass  vessel  (beaker)  with  water 
containing  tine  sawdust,  and  heat  at  the  bottom.  The  move- 
ments of  the  sawdust  show  that  the  water  as  it  is  heated  rises  to 
the  top,  while  the  colder  water  sinks. 

Thus,  heat  is  conveyed  from  place  to  place  by  motion 
en  masse  of  the  substance  which  is  being  heated,  i.  e.,  by 
convection  currents.  Winds  are  convection  currents  on 
a  grand  scale ;  drafts  are  the  same  on  a  smaller  scale. 
From  experiments  4  and  5  it  is  easily  seen  that  a  mass 
of  water  can  be  much  more  quickly  heated  by  applying 
the  heat  at  the  bottom,  than  by  applying  it  at  the  top. 

24.  Expansion  by  Heat. — In  Experiment  5  the 
water  rises  as  it  becomes  warm,  because  it  is  lighter  than 
the  cold  water.     Its  volume  has  been  increased.     Heat 
expands  water. 

Experiment  6. — Fill  a  glass  flask  full  with  cold  water,  and 
heat  the  water  to  boiling.  A  small  quantity  runs  over  as  the 
water  becomes  hot.  Allow  the  water  to  cool.  It  sinks  down 
the  neck  of  the  flask.  Water  expands  when  heated  and  con- 
tracts when  cooled. 

Other  liquids  expand  when  heated  and  contract  on 
cooling,  but  different  liquids  expand  differently.  For 
example,  mercury  expands  less  than  water  for  the  same 
increase  in  degree  of  heat.  Solids  also  expand  when 
heated,  but  not  so  rapidly  as  liquids.  It  will  be  shown 
later  on  that  gases  expand  when  heated  much  more  than 
either  liquids  or  solids,  and  that  all  gases  expand  almost 
exactly  at  the  same  rate. 


TEMPERATURE.  15 

Experiment  7- — Fit  a  cork  to  a  glass  flask,  bore  a  hole 
through  the  cork,  and  through  the  hole  put  a  glass  tube  fitting 
tightly.  Fill  the  flask  with  water,  push  in  the  cork  tightly  so 
as  to  cause  the  water  to  rise  an  inch  or  two  in  the  tube.  Pour 
hot  water  on  the  flask.  The  water  in  the  tube  at  first  sinks,  but 
immediately  afterwards  begins  to  rise  steadily.  When  it  has 
stopped  rising  it  begins  to  fall,  and  finally  remains  at  about  the 
same  level  as  before  the  hot  water  was  poured  on  the  flask. 

The  flask  expands  first  when  the  hot  water  is  poured 
on  and  the  contained  water  sinks  down  the  tube.  Then 
the  heat  penetrates  to  the  water  which  expands  faster 
than  the  glass,  and  therefore  rises  in  the  tube. — Builders 
of  bridges  allow  for  the  expansion  and  contraction  of  iron 
with  change  of  season,  &c.  In  laying  rails  in  cold 
weather  the  ends  are  not  put  close  together,  but  space  is 
allowed  for  the  increased  length  in  warm  weather. 

25.  Temperature  — If  an  instrument  such  as  that 
described  in  Experiment  7  were  put  into  a  body  of  water 
which  caused  the  water  in  the  tube  to  rise,  the  conclu- 
sion drawn  from  such  an  experiment  would  be  that  the 
body  of  water  was  hotter  than  that  in  the  flask.  If,  on 
the  other  hand,  the  liquid  should  fall  in  the  tube,  it  would 
be  concluded  that  the  water  under  examination  was  cooler 
than  that  in  the  flask  ;  if  no  change  were  produced,  the 
conclusion  would  certainly  be  that  it  was  neither  hotter 
nor  colder.  The  sense  of  touch  would  confirm  these  con- 
clusions in  each  case,  if  the  hand  were  immersed  succes- 
sively in  the  two  quantities  of  water. 

Experiment  8- — Lay  a  piece  of  iron  and  a  block  of  wood 
side  by  side  on  a  table  and  after  half  an  hour  touch  each  of  them 
with  the  hand.  The  iron  feels  colder  than  the  wood  ;  and  yet, 


16  THERMOMETERS. 

if  an  instrument  such  as  that  used  in  Experiment  7  be  touched 
first  to  one  and  then  to  the  other  it  shows  them  to  be  exactly  in 
the  same  state. 

The  sense  of  touch  does  not  always  give  the  same  ver- 
dict regarding  heat  as  the  effect  on  volume  does.  The 
iron  feels  colder  because  it  conducts  heat  away  from  the 
hand  more  rapidly,  being  a  better  conductor  than  wood. 
For  the  same  reason  blankets  feel  warmer  than  sheets. 
The  words  "  hot "  and  "  cold  "  do  not  convey  a  correct 
idea  of  the  state  of  a  body  with  regard  to  its  sensible  heat. 
Generally,  bodies  which  have  been  in  contact  for  some 
time  are  in  the  samo  state  with  regard  to  sensible  heat, 
for,  if  they  are  at  first  in  different  states,  heat  flows  from 
the  hotter  to  the  colder  until  they  are,  as  it  were,  at  the 
same  heat-level,  or  at  the  same  temperature. 

Definition. — Temperature  is  the  state  of  a  body  with  regard 
to  sensible  heat ;  or  temperature  is  heat-level. 

26.  Thermometers. — Temperatures  are  measured 
by  thermometers,  or  heat-measures.  In  most  thermo- 
meters the  temperature  is  indicated  by  the  amount  of 
expansion  of  a  liquid  enclosed  in  a  graduated  glass  tube 
with  a  bulb  or  resorvoir  at  the  lower  end.  The  most 
convenient  liquid  for  ordinary  ranges  of  temperature  is 
mercury.  It  expands  regularly,  freezes  only  at  a  low, 
and  boils  only  at  a  high  temperature.  Alcohol  is  used 
for  very  low  temperatures  as  it  does  not  freeze  until  the 
temperature  sinks  much  below  any  that  occurs  naturally. 
The  thermometer  is  graduated  by  placing  it  first  in  a 
mixture  of  ice  and  water,  then  in  steam  at  ordinary 
pressure,  and  marking  on  the  stem  or  scale  the  level  at 
which  the  mercury  stands  in  each  case.  These  are  the 


THERMOMETERS.  17 

fixed  points,  and  always  when  the  thermometer  is  placed 
in  melting  ice  the  mercury  stands  at  the  lower  fixed 
point,  and  when  it  is  placed  in  water  boiling  under 
ordinary  circumstances,  the  mercuiy  remains  stationary 
at  the  higher  fixed  point.  Under  the  same  circumstances 
water  freezes  always  at  the  same  temperature,  and  also 
boils  always  at  the  same  temperature.  It  only  remains 
to  divide  the  interval  between  the  fixed  points  into  equal 
parts  called  decrees,  and  to  number  these  in  regular 
order.  The  numbers  chosen  are  different  in  thermo- 
meters devised  by  different  men.  Any  numbers  may  be 
chosen,  according  to  taste  or  convenience.  On  the  Cen- 
tigrade or  Celsius  thermometer,  the  one  always  used  for 
scientific  purposes,  the  lower  fixed  point  is  marked  0°, 
the  upper  100°,  and  the  space  between  is  divided  into 
100  degrees.  On  the  Fahrenheit  thermometer,  generally 
used  in  this  country,  the  two  points  are  marked  32°  and 
212°  respectively,  so  that  the  zero  Fahrenheit  is  32  de- 
grees below  the  freezing  point  of  water.  On  the  Reau- 
mur thermometer  the  space  between  the  fixed  points  is 
divided  into  80  degrees  numbered  from  0°  to  80°.  Thus, 
the  degrees  on  the  Reaumur  scale  are  longer  than  on  the 
Centigrade,  which  in  their  turn  are  longer  than  those  on 
the  Fahrenheit.  The  relations  between  the  degrees  are 
expressed  as  follows  : — 1  degree  F.  =  |-f§  --  f  degree 
C.  =  ^j  =  ^  degree  R.  To  change  any  temperature 
Fahrenheit  to  Centigrade  (or  Reaumur)  subtract  32  and 
multiply  by  f  (or  f).  The  32  degrees  between  the 
Fahrenheit  zero  and  the  freezing  point  must  be  sub- 
tracted first,  because  the  other  thermometers  reckon  from 
the  freezing  point.  To  change  any  temperature  Centi- 
grade (or  Reaumur)  to  Fahrenheit,  multiply  by  f  (or  f) 
and  add  32.  Or,  if  t  represent  any  temperature  Fahreii- 
3 


18  MAXIMUM    DENSITY    POINT. 

heit  (t  —  32)  %  represents  the  same  temperature  Centi- 
grade, and  (t  —  32)  £  the  same  temperature  Reaumur. 
Also  let  t  be  any  temperature  Centigrade,  then  f  t  +  32 
represents  that  temperature  Fahrenheit.  Similarly  for 
Reaumur.  Temperatures  below  zero  are  indicated  by 
the  minus  sign  (  —  ). 

27.  Maximum  Density  Point  of  Water. 

Experiment  9- — Bore  a  second  hole  in  the  cork  in  the  appa- 
ratus of  Experiment  8,  fit  a  centigrade  thermometer  in  it,  and 
arrange  the  apparatus  as  in  Experiment  8.  Surround  the  flask 
with  snow  or  ice.  The  water  in  the  tube  falls  until  the  ther- 
mometer marks  4°  C.  It  then  begins  to  rise,  and  continues  to 
do  so  until  the  temperature  sinks  to  0°. 

Water  has  the  smallest  volume  at  4°  C.,  and  has, 
therefore,  the  greatest  specific  weight  at  that  tempera- 
ture, which  is  for  this  reason  called  the  maximum  density 
point  of  water.  Most  liquids  are  densest  at  their  freezing 
points.  Water  is  exceptional  in  this  respect.  When  a 
body  of  water  cools  at  the  surface,  as  a  lake  in  autumn, 
the  cooled  water  sinks  until  the  whole  lake  is  at  the  tem- 
perature 4°  C.  When  the  water  at  the  surface  is  cooled 
below  this  it  remains  at  the  surface,  because  it  expands 
and  thus  becomes  specifically  lighter  than  the  warmer 
water  below.  As  water  is  a  very  poor  conductor  of  heat 
(see  Experiment  4),  the  water  below  cools  very  slowly, 
even  if  the  surface  becomes  changed  into  ice.  Were  it 
not  for  this  exceptional  feature  in  the  effect  of  cooling  on 
water,  our  lakes  would  become  masses  of  ice  every  winter. 

28.  Freezing  and  Melting. 

Experiment  10. — Heat  slowly  a  vessel  containing  snow  or 
ice,  stirring  constantly  with  a  thermometer.  The  temperature 


FREEZING    AND    MELTING.  19 

of  the  whole  mass  remains  at  0°  C.  until  the  last  portion  of  snow 
or  ice  is  melted.     It  then  begins  to  rise. 

Ice  melts  at  0°  C.  always,  under  ordinary  circum- 
stances. Other  solids  have  different  melting  points,  e.  g. 
Lead  melts  at  300°  C.  The  word  fuse  is  used  instead  of 
melt,  especially  of  solids  which  become  liquid  only  at 
high  temperatures.  If  a  portion  of  water  be  cooled  it 
begins  to  freeze  at  0°  C.,  and  the  temperature  does  not 
sink  below  that  until  the  whole  of  the  water  is  changed 
to  ice.  The  freezing  point  of  water  is  the  same  as  the 
melting  point  of  ice.  This  is  true  of  most  liquids,  but 
not  of  all. 

29.  Change  of  Freezing  Point  by  Pressure. 

Experiment  11- — Support  a  slab  of  clear  ice  at  either  end 
and  lay  over  it  a  thin  clean  wire  having  a  weight  attached  to 
each  end  so  as  to  cause  the  wire  to  press  upon  the  ice.  The  wire 
gradually  passes  through  the  ice,  but  does  not  cut  it  in  two. 

The  ice  is  melted  by  the  pressure  of  the  wire  upon  it, 
but  solidifies  again  as  soon  as  the  pressure  is  removed  by 
the  onward  movement  of  the  wire.  Ice  at  0°  C.  can  be 
melted  by  pressure,  and  cannot  be  frozen  as  long  as  the 
pressure  continues,  without  reducing  the  temperature 
below  0°  C.  In  other  words  the  freezing  point  of  water 
is  lowered  by  pressure.  Water  is  exceptional  in  this 
respect  also ;  most  liquids  have  their  freezing  points 
raised  by  pressure. 

30.  Change  of  Volume  on  Freezing. — Since 

ice  floats  on  water  it   must  be  specifically  lighter  than 
water.      Water  expands  on  freezing.     Most  liquids  con- 


20  LATENT    HEAT    OF    WATER. 

tract  on  solidifying,  e.g.,  solid  lead  sinks  in  liquid  lead. 
Water  expands  about  one-eleventh  of  its  volume  on 
freezing.  Enormous  force  is  exerted  during. this  expan- 
sion ;  hence  the  bursting  of  water-pipes,  the  splitting 
of  rocks,  the  "heaving"  of  the  foundations  of  build- 
ings, &c. 

31.  Latent  Heat  Of  Water. — From  Experiment 
10  we  learn  that  when  ice  is  melting  heat  goes  into  it 
without  raising  its  temperature.  This  heat  becomes 
latent  or  hidden.  In  order  to  freeze  the  water  again  this 
latent  heat  must  be  removed  from  it. 

Experiment  12. — Mix  1,000  grams  ice-cold  water  with  1,000 
grams  boiling  water.  The  temperature  of  the  mixture  is  50°  C. 
Now,  mix  1,000  grams  ice  with  1,000  grams  boiling  water,  stir 
well  until  the  ice  is  melted.  The  temperature  of  the  mixture  is 
only  10.5°  C. 

From  these  experiments  it  is  plain  that  the  heat 
rendered  latent  by  melting  1,000  grams  ice  would  warm 
2,000  grams  water  from  10.5°  C  to  50°  C.  If  we  take 
as  the  unit  of  heat  (or  thermal  unit)  the  quantity  of  heat 
required  to  raise  the  temperature  of  1  gram  of  water  1 
degree  Centigrade,  then  2,000  x  39.5  =  79,000  units 
of  heat  are  rendered  latent  in  melting  1,OOD  grams  ice, 
or  79  for  1  gram.  The  Latent  Heat  of  Water  is  79 
thermal  units  ;  or,  to  change  ice  at  0°  into  water  at  Oc 
as  much  heat  is  rendered  latent  as  would  warm  the  same 
weight  of  water  from  0°  to  79°.  The  latent  heat  of 
water  is  much  greater  than  that  of  most  substances. 
That  of  mercury  is  only  2.28;  that  of  molten  lead  5.4. 
Ice  is  thus  an  excellent  cooling  agent. 


SPECIFIC    HEAT.  21 

32.  Specific  Heat. 

Experiment  13- — Take  equal  weights  of  iron  and  lead,  about 
1,000  grams  ;  cool  them  to  0°  and  immerse  them  in  equal  quan- 
tities of  water  (500  c.  c. )  at  50°  C.  After  stirring  for  a  moment 
note  the  temperatures.  That  of  the  water  in  which  the  iron  is 
will  be  about  40°  C. ;  of  that  in  which  the  lead  is,  about  47°  C. 

The  iron  has  taken  more  heat  from  the  water  than  the 
lead  has,  and  yet  its  temperature  has  not  been  raised  so 
high.  Iron  is  harder  to  heat  than  an  equal  weight  of 
lead,  i.e.,  requires  more  heat  to  raise  its  temperature,  say, 
one  degree.  Different  substances  require  different  quan- 
tities of  heat  to  raise  the  temperature  of  equal  weights 
of  them  one  degree. 

Definition. — The  Specific  Heat  of  a  substance  is  the  quantity 
of  heat  required  to  raise  the  temperature  of  unit  weight  (1  gram 
or  1  Ib.)  of  the  substance  one,  degree  (from  0°  C.  to  1°  C.) 

The  specific  heat  of  water  is  of  course  the  thermal 
unit.  That  of  most  other  substances  is  much  less.  That 
of  lead  is  only  0.031,  of  iron,  0.1137,  and  of  copper, 
0.095. 

33.  Evaporation  and   Ebullition. — Water  in 

the  state  of  gas  (water  vapour)  is  constantly  escaping 
into  the  air  from  every  exposed  surface  of  water  or  ice ; 
and,  other  things  being  equal,  the  hotter  the  water  the 
more  rapidly  does  it  become  changed  to  vapour.  This 
process  is  called  evaporation.  Thus  the  air  always  con- 
tains water-vapour.  At  any  particular  temperature  a 
portion  of  air  is  capable  of  containing  only  a  certain 
quantity  of  water-vapour.  When  air  contains  as  much 
water- vapor  as  it  is  capable  of  holding,  it  is  saturated  ; 


22  EVAPORATION    AND    EBULLITION. 

and  if  it  bo  then  cooled,  some  of  the  vapour  is  condensed  ; 
if  it  be  heated,  it  becomes  capable  of  containing  more 
vapour.  The  air  corning  from  our  lungs  is  warm  and 
saturated  with  water-vapour.  When  it  comes  in  contact 
with  cold  air  in  winter  some  of  the  vapour  is  condensed 
and  forms  a  "  cloud."  Dew,  fogs,  clouds,  mists.  &c.,  are 
explained  similarly.  Air  is  moist  or  dry  according  as  it 
is  near  to,  or  far  away  from,  its  point  of  saturation. 
Thus,  if  moist  air  be  heated  it  becomes  dry,  i.e.,  capable 
of  receiving  more  water-vapour.  Evaporation  is  hastened 
by  heat,  or  by  a  current  of  dry  air. 

Experiment  14. — Heat  some  water  in  a  glass  flask  until  it 
begins  to  boil,  observing  the  temperature  by  a  thermometer  in 
the  water.  The  water  begins  to  boil  when  the  temperature  rises 
to  160°  0.  Observe  the  temperature  of  the  steam.  It  is  also 
100°.  Use  a  larger  flame  so  as  to  make  the  water  boil  more 
vigorously.  The  temperature  does  not  rise.  Now,  close  the 
flask  with  a  cork  through  which  the  thermometer  passes.  Heat 
very  carefully,  and  observe  that  the  temperature  rises  above  100° 
and  yet  the  water  scarcely  boils.  Remove  the  lamp  when  the 
mercury  has  risen  one  or  two  degrees,  otherwise  the  cork  will  be 
driven  out  by  the  pressure  of  the  steam.  Take  out  the  cork, 
boil  the  water  vigorously  for  a  minute  or  two  and  while  steam  is 
still  coming  out  put  in  a  cork  firmly  at  the  same  time  taking  the 
flask  away  from  the  lamp.  Allow  the  water  to  cool  a  little,  and 
then  pour  cold  water  on  the  flask.  The  water  in  the  flask  begins 
to  boil. 

Water  in  an  open  vessel  boils  at  100°  C.,  and  the  tem- 
perature cannot  be  raised  above  that,  unless  the  vessel  be 
closed  so  as  to  increase  the  pressure  on  the  surface  of  the 
water.  When  the  pressure  on  the  surface  is  lowered,  as 
by  condensing  the  steam — filling  the  closed  flask  in  Ex- 
periment 14,  water  boils  at  temperatures  below  100°  C. 
The  boiling  point  of  water  depends  on  the  pressure  on  its 


LATENT    HEAT    OF    STEAM.  23 

surface,  being  higher  the  greater  the  pressure.  On  the 
top  of  a  high  mountain  where  the  weight  of  air  pressing 
on  surfaces  is  less  than  it  is  lower  down,  water  boils  at 
so  low  a  temperature  that  it  can  not  cook. 

Experiment  15- — Boil  some  alcohol  in  a  flask,  observing  the 
temperature  of  the  boiling  liquid  and  of  the  vapour.  If  the 
alcohol  is  pure,  the  temperature  in  each  case  is  78.4°  C.  Try 
the  boiling  point  of  ether,  heating  it  by  means  of  hot  water. 
Ether  boils  at  34.2°. 

Every  liquid  boils  at  a  particular  temperature,  just  as 
every  solid  fuses  at  a  particular  temperature. 


BOILING    POINTS. 


Nitrous  oxide — 87.9 

Carbon  dioxide —78.2 

Ammonia — 33.7 

Sulphur  dioxide — 10.5 

Aldehyde +19.8 

Ether 34.2 

Carbon  Bisulphide 47.9 

Methylic  alcohol 55.1 

Chloroform 61.0 

Bromine...  63.0 


Ethylic  Alcohol 78.4 

Benzene 80.4 

Water 100.0 

Acetic  Acid 116.9 

Naphthalene 216.8 

Phosphorus 290.0 

Mercury 350.0 

Sulphur 440.0 

Cadmium 860.0 

Zinc..                              .  1040.0 


34.  Latent  Heat  Of  Steam. — From  Experiment 
14  we  learn  that  the  heat  which  goes  into  boiling  water 
and  changes  it  to  steam  does  not  cause  any  rise  of  tem- 
perature. This  heat  becomes  latent,  just  as  in  the  case 
of  melting  ice.  The  latent  heat  of  steam  is  536  thermal 
units ;  i.e.,  to  change  any  weight  of  water  at  100°  C.  to 
steam  at  the  same  temperature  heat  is  required  sufficient 
to  raise  the  temperature  of  536  times  the  weight  of 
water  one  degree.  The  latent  heat  of  steam  is  very 
great  compared  with  that  of  other  vapours.  Water  in 
evaporating  from  the  earth's  surface  carries  away  a  great 


24  SOLUTION. 

deal  of  heat.  Wet  clothes  feel  cold  because  of  the  heat 
consumed  by  the  evaporating  water.  Much  heat  is  lost 
to  the  body  by  the  evaporation  of  water  through  the 
skin  and  lungs.  The  cooling  effect  of  ether  and  other 
rapidly  evaporating  liquids  is  explained  in  the  same  way. 
Water  can  be  frozen  by  rapidly  evaporating  ether.  That 
such  an  enormous  amount  of  heat  is  required  to  evapo- 
rate a  liquid  is  accounted  for  by  the  great  change  of 
volume.  For  example,  1  cubic  inch  of  water  forms  1,700 
cubic  inches  of  steam. 

35.  Solution. 

Experiment  16. — Mix  100  grams  finely-powdered  saltpetre 
with  100  grams  distilled  water  in  a  beaker,  stirring  for  some 
time  with  a  glass  rod.  Part  of  the  saltpetre  disappears  ;  the 
water  has  now  the  taste  of  saltpetre,  and  is  colder.  Apply  heat, 
carefully  stirring  all  the  time.  More  of  the  saltpetre,  and  at 
length  all  of  it,  disappears.  Allow  the  liquid  to  cool ;  some  of 
the  saltpetre  reappears  in  long  crystals  which  grow  as  the  liquid 
cools.  When  quite  cool,  pour  off  some  of  the  liquid  into  a  small 
porcelain  basin,  boil  away  some  of  the  water,  and  allow  the  re- 
maining liquid  to  cool  again.  More  saltpetre  crystallises  out. 
This  process  is  called  concentration.  Repeat  this  experiment, 
using  washing  soda  instead  of  saltpetre. 

Experiment  17- — Carefully  heat  300  grams  saltpetre  with 
100  grams  water.  (Do  not  apply  the  flame  to  the  beaker.) 
Only  part  of  the  saltpetre  disappears  even  when  the  water 
begins  to  boil.  Mix  10  grams  of  gypsum  with  100  of  water,  and 
heat  to  boiling.  The  gypsum  does  not  disappear.  Filter  some 
of  the  liquid  into  a  porcelain  basin  and  evaporate  to  dryness . 
A  small  quantity  of  white  substance,  gypsum,  is  left.  Repeat 
this  experiment,  using  clean  white  sand,  and  again  with  chalk. 
In  each  case  little  or  nothing  remains  when  the  water  is  evapo- 
rated. 


SOLUTION.  25 

Saltpetre  and  washing  soda  dissolve  readily  in  water, 
imparting  their  taste  and  other  properties  to  the  water. 
They  are  soluble  in  water,  and  form  solutions  from  which 
they  can  be  got  again  by  evaporating  the  water. 

Experiment  18- — Mix  10  c.  c.  hydrochloric  acid  with  about 
four  times  as  much  water,  pour  it  over  three  or  four  iron  tacks 
in  a  porcelain  basin,  and  warm  gently.  Bubbles  rise  from  the 
tacks,  which  gradually  disappear,  giving  a  greenish  color  to  the 
water.  Evaporate  the  water.  A  solid  remains,  but  it  is  not 
iron ;  it  is  ferrous  chloride.  Repeat  the  experiment,  using 
marble  instead  of  iron. 

In  Experiment   18  the  process  of  solution  is  accom- 
panied  by  chemical  action.      This  kind   may  be  called 
chemical  solution. — Substances  are  generally  more  soluble 
in  hot  than  in  cold  water. — Gypsum  is  much  less  soluble 
than  saltpetre  or    washing  soda.     Sand  and  chalk   are 
insoluble.     A  given  quantity  of  water  can  dissolve  only 
a  certain  maximum  weight  of  each  soluble  substance  at 
a  given  temperature,  forming  in  each    case  a  saturated 
solution.     For  example,  100  grams  of  water  will  always 
dissolve  13J  grams  of  saltpetre  at  0°  C.,  and  always  250 
grams  at  100°  C.     These  weights  may  be  taken  to  repre- 
sent the  solubility  of  saltpetre  at  the  two  temperatures. 
(For  Table  of  Solubilities,  see  Appendix.)     To  make  a 
saturated  solution  of  a  substance  at  any  temperature, 
either  stir  it  up  with  water  at  that  temperature  until  no 
more  will  dissolve  (this  requires  much  time  and  stirring), 
or  heat  the  water  to  a  higher  temperature,  dissolve  in  it 
as  much  as  possible  of  the  substance,  cool,  and  decant  or 
filter.     Water  saturated   with  one   substance   will   still 
dissolve  others.     Solutions  of  solids  in  water  are  heavier 
than  pure  water  and  boil  at  higher  temperatures.     Heat 


26    SUPERSATURATED  SOLUTIONS  —  FREEZING  MIXTURES. 

is  rendered  latent  when  solids  dissolve  in  liquids,  ;md 
reappears  when  they  become  solids  again  (compare 
Art.  31). 

Experiment  19. — Heat  some  sodic  acetate  in  a  flask  with 
very  little  water  until  the  salt  is  dissolved  and  steam  issues 
from  the  flask.  Put  a  plug  of  cotton  wool,  or  a  cork,  in  the 
mouth  of  the  flask,  and  let  cool.  The  solution  remains  liquid. 
Open  the  flask  and  drop  in  a  small  grain  of  sodic  acetate.  Crys- 
tals shoot  out  from  the  point  where  it  falls,  and  soon  the  whole 
mass  becomes  solid.  At  the  same  time  the  flask  becomes  warm. 
(Why?)  A  solution  of  this  sort  is  supersaturated.  It  contains 
more  of  the  salt  than  can  be  dissolved  at  the  lower  temperature. 

Water  dissolves  a  great  many  substances,  solids,  liquids, 
and  gases.  It  is  a  good  menstruum,  or  solvent.  Many 
substances  insoluble  in  water  will  dissolve  in  other 
liquids.  Thus,  mercury  dissolves  most  of  the  metals ; 
alcohol  dissolves  resins  ;  and  ether  dissolves  fats  ;  all  in- 
soluble in  water.  No  heat  becomes  latent  when  one  liquid 
dissolves  another ;  there  is  no  change  of  state.*  When  a 
liquid  dissolves  a  gas  the  latent  heat  of  the  gas  becomes 
sensible,  but  becomes  latent  again  when  the  gas  resumes 
its  former  state.  This  accounts  for  the  coolness  of  effer- 
vescing drinks. 

36.  Freezing  Mixtures. 

Experiment  20- — Mix  quickly  and  thoroughly  5  parts  by 
weight  of  snow  with  2  parts  of  salt.  Observe  the  intense  cold. 
Freeze  some  water  in  a  test-tube.  Note  the  melting  of  both 
snow  and  salt.  The  temperature  can  be  reduced  in  this  way  to 
—20°  C.  What  becomes  of  the  heat  ?  (See  Art.  31  and  Art.  35). 
Salt  water  will  not  freeze  until  the  temperature  sinks  much 
below  0°  C. 

*  In  some  cases  heat  becomes  sensible,  as  when  alcohol  and  water  are  mixed. 


CRYSTALLISATION.  27 

Another  freezing  mixture  is   16  parts  water,  5  parts 
saltpetre,  and  5  of  sal-ammoniac. 

37.  Crystallisation. — Crystals  are    beautiful    geo- 
metric forms,  e.g.,  cubes,  pyramids,  in  which  solids  can  be 
obtained.     They  may  be  produced  in  several  ways,  par- 
ticularly by  cooling  or  evaporating  solutions,  as  in  the 
experiments  with  saltpetre  (Art.   35).     A  substance  is 
known  to  be  at  least  nearly  pure,  if  it  consists  of  distinct 
and  well-formed  crystals.     This  is  the  reason  for  the  care 
taken  by  manufacturers  to  prepare  well-crystallised  chemi- 
cals in  particular  cases  where  adulterations  are  common. 
Crystals  are  generally  obtained  by  gradually  cooling  hot 
saturated  solutions,  or  by  slowly  evaporating  cold  solu- 
tions.    The  impurities  remain  dissolved  in  the  mother- 
liquor.     Substances  are  often  purified  by  crystallisation. 
Many  substances  combine  with  water  as  they  crystallise. 
Washing-soda  crystals  contain  63%  of  water  of  crystallisa- 
tion.   The  water  is  necessary  to  the  form  of  the  crystals. 
It  is  chemically  combined,  and  is  solid.     If  it  be  driven 
off  by  heat  the  cr}*stals  fall  to  powder.     Many  crystalline 
substances  containing  water  of  crystallisation  lose  this  on 
exposure  to  air,  and  become  powdery.     This  phenomenon 
is   called    efflorescence.     Other    solid    substances   attract 
moisture  from  the  air  and  at  last  become  liquid.     This  is 
called  deliquescence. 

38.  Infusion,  &C. — Water  and  other  solvents  are 
used  to  separate  the  soluble  from  the  insoluble  parts  of 
medicinal  plants,  &c.     Maceration,  or  cold  infusion,  is 
carried  on  at  the   ordinary   temperature  of  the  atmos- 
phere.    The  substance  to  be  macerated  is  ground  up,  or 
in  some  other  way  intimately  mixed  with  the  solvent,  is 


28  INFUSION,    ETC. 

allowed  to  stand  for  some  time,  and  the  liquid  is  then 
strained  or  filtered  off.  An  infusion  (hot  infusion)  is 
made  by  pouring  the  boiling  solvent  on  the  substance, 
and  allowing  it  to  cool  gradually.  Digestion  (or  simmer- 
ing) is  the  process  of  making  a  solution  with  the  solvent 
kept  just  below  its  boiling  point.  The  differences  in 
these  methods  are  mainly  differences  of  temperature. 
When  it  is  required  to  extract  a  substance  which  is  de- 
composed at  or  below  the  boiling  point  of  the  solvent, 
or  a  substance  which  is  very  volatile  (easily  evaporated), 
the  method  of  cold  infusion  or  digestion  is  used.  Tea 
should  be  infused,  or  at  most  digested,  and  never  boiled, 
since  boiling  extracts  a  bitter  and  unwholesome  principle, 
which  is  not  dissolved  in  so  great  quantity  by  infusion  or 
digestion.  Tinctures  are  mostly  macerations  with  alcohol 
as  the  solvent.  Percolation  has  now  largely  superseded 
maceration  in  the  preparation  of  tinctures.  The  sub- 
stance is  placed  in  a  filter  (percolator),  and  the  solvent  is 
poured  upon  it  and  allowed  to  percolate  (run  through) 
slowly.  This  method  has  the  advantage  of  allowing  a 
fresh  portion  of  the  solvent  to  be  poured  on  the  partially 
exhausted  material ;  so  that  the  extraction  can  be  com- 
pleted more  quickly  and  thoroughly.  (Why  ?)  The  same 
principle  is  employed  in  filtration  when  it  is  wished  to 
wash  away  a  soluble  from  an  insoluble  substance  with  as 
little  as  possible  of  the  solvent.  Very  little  is  poured  on 
at  a  time,  and  one  portion  is  allowed  to  run  through  before 
another  is  added. 

39.  Decomposition  Of  Water.— In  Experiment 
1  heat  was  the  kind  of  energy  used  to  decompose  a  com- 
pound. Very  many  chemical  substances  can  be  decom- 
posed by  heat,  and  under  certain  conditions  water  under 


DECOMPOSITION    OF    WATKR.  29 

the  action  of  heat  breaks  up  into  two  different  substances  ; 
but  another  kind  of  energy,  electricity,  is  more  convenient 
for  showing  the  compound  nature  of  water.  It  must  be 
remembered  that  electricity  is  no  more  a  substance  than 
heat  is.  It  is  classed  along  with  light,  heat,  &c.,  as  a 
kind  of  energy,  power  of  doiny  work  of  some  kind.  One 
kind,  of  work  for  which  electricity  is  specially  suited  is 
to  separate  compound  into  simpler  substances.  Chemical 
action  is  generally  accompanied  by  the  development  of 
heat,  and  in  the  galvanic  battery,  a  current  of  electricity 
takes  the  place  of  some  of  the  heat  which  would  be  set 
free  if  the  metal  used  were  allowed  to  dissolve  freely  in 
the  acid  of  the  battery.  Good  conductors  of  heat  are 
generally  good  conductors  of  electricity.  Compounds 
which  will  not  conduct  electricity,  i.  e.,  through  which  it 
will  not  flow,  are  not  decomposable  by  it.  They  may, 
however,  become  conductors  by  mixing  with  other  sub- 
stances, by  melting,  &c.,  and  then,  they  may  be  decom- 
posed. Water  is  a  very  bad  conductor  of  electricity 
when  it  is  pure,  but  when  slightly  acidified  with  sulphuric 
acid  it  conducts  well. 

Experiment  21. — Fasten  two  pieces  of  platinum  foil  or  wire 
to  the  ends  of  the  wires  from  a  galvanic  battery.  Dip  the  foils 
in  a  vessel  of  distilled  water,  keeping  them  about  an  inch  apart. 
Unless  the  battery  is  very  powerful,  there  is  no  result.  Add  a 
few  drops  of  sulphuric  acid  to  the  water.  Bubbles  of  gas  gather 
on  the  foil  and  rise.  Collect  the  gases  in  two  graduated  tubes 
filled  with  water  and  inverted  over  the  pieces  of  platinum  foil 
(electrodes).  Observe  that  one  tube  fills  faster  than  the  other, 
and  after  some  time  the  quantity  in  the  one  tube  is  seen  to  be 
double  that  in  the  other.  Remove  the  tubes  one  at  a  time, 
closing  them  with  the  thumb,  and  put  a  lighted  match  into  each. 
The  larger  quantity  of  gases  catches  tire  and  burns,  putting  out 
the  match,  but  the  smaller  causes  the  match  to  burn  more 


30  COMPOSITION    OF    WATER. 

brightly.  With  a  very  powerful  battery  a  small  portion  of  pure 
water  could  be  completely  decomposed  into  these  two  gases 
differing  in  this  and  in  other  respects. 

40.  Composition  of   Water. — We  thus  prove 

water  to  be  a  chemical  compound  of  two  gases,  Hydrogen 
and  Oxygen.  These  gases  have  never  been  further  de- 
composed and  are  therefore  held  to  be  elements.  Water 
has  been  decomposed  in  a  variety  of  ways,  and  always 
into  these  two  gases ;  further,  the  volumes  of  the  gases 
have  always  been  in  the  ratio  of  two  of  hydrogen  and 
one  of  oxygen.  The  next  thing  to  discover  is  the  weights 
of  oxygen  and  hydrogen  which  form  a  given  weight  of 
water.  In  Experiment  21  water  is  analysed — its  com- 
position by  volume  has  been  found  by  analysis,  i.e.,  by 
decomposing  it  into  elements.  It  will  be  seen  later  that 
the  word  analysis  is  used  also  in  a  broader  sense.  In 
order  to  discover  the  composition  by  weight,  the  method 
of  building  up,  or  synthesis,  is  used.  A  current  of  hydro- 
gen gas  (How  dried  ?  See  Sulphuric  Acid.)  is  allowed 
to  flow  through  a  weighed  tube  containing  black  oxide 
of  copper  (a  compound  of  oxygen  and  copper)  heated  to 
redness.  The  hydrogen  takes  away  oxygen  from  this 
substance,  forming  water.  This  water  is  collected  and 
weighed,  and  the  tube  of  cupric  oxide  is  re-weighed. 
The  loss  of  weight  in  the  latter  case  is  the  weight  of 
oxygen  consumed,  and  the  difference  between  this  and 
the  weight  of  water  collected  gives  the  weight  of  hydro- 
gen. (What  law  is  assumed  here  1)  This  experiment 
has  been  made  thousands  of  times  and  with  the  utmost 
nicety,  and  always  with  the  same  result — for  every  8  parts 
by  weight  of  oxygen  taken  away  from  the  oxide  of  copper 
9  parts  of  water  are  formed.  Water  is  then  composed 


QUESTIONS    AND    EXERCISES.  31 

of  the  two  gaseous  elements,  hydrogen  and  oxygen, 
united  in  the  proportion  of  1  of  hydrogen  to  8  of  oxygen. 
(Deduce  the  relative  weights  of  hydrogen  and  oxygen.) 


QUESTIONS  AND  EXERCISES. 

1.  Is  steam  visible? 

2.  It  is  observed  that  arctic  animals  are  generally  white.     Of 
what  advantage  is  this  to  the  animals  ? 

3.  Why  does  hot  glass  crack  when  a  drop  of  water  falls  on  it  ? 

4.  Change  10°,  32°,  and  —  40°  Fahrenheit  to  centigrade. 

5.  At   what   temperature   is   the  number  on  the  Fahrenheit 
scale  double  that  on  the  centigrade  ? 

6.  A  Fahrenheit  and  a  centigrade  thermometer  are  placed  side 
by  side  in  a  vessel  of  water.     The  centigrade  thermometer  reads 
80°.     What  does  the  Fahrenheit  read  ? 

7.  What  is  the  latent  heat  of  water  if  the  Fahrenheit  ther- 
mometer is  the  standard  ? 

8.  Twenty  grams  of  ice  at  0°  C.  is  melted  with  1,000  grams 
of  water  at  50°.     What  is  the  temperature  afterwards  ? 

9.  What  weight  of  water  at  100°  C.  is  required  to  melt  10  Ibs. 
of  ice  at  0°  C.  ? 

10.  Why  does  snow  fall  powdery  in  very  cold  weather,  and  in 
large  flakes  in  mild  weather  ? 

11.  If  100  grains  of  steam  at  100°  be  condensed  in  1  Ib.  of  water 
at  0°  C.,  what  is  then  the  temperature  of  the  water  ? 

12.  A  piece  of  silver  weighing  120  grams  is  heated  in  boiling 
water  and  then  put  into  a  vessel  containing  500  grams  of  water 
at  10°  C.     After  stirring,  the  temperature  of  the  water  is  found 
to  be  11.2°.     Calculate  the  specific  heat  of  silver. 

13.  Why  does  stirring  facilitate  the  solution  of  a  substance  ? 

14.  How  much  steam  must  be  condensed  to  heat  10  kilograms 
of  water  from  15°  C.  to  100°  C.  ? 


DISCOVERY    OF    OXYGEN. 


CHAPTER     IV. 


OXYGEN—  COMBUSTION—  CHEMISM—  METALS    AND 
NON-METALS. 

41.  Discovery  Of  Oxygen.  —  Dr.  Joseph  Priestly 
was  the  father  of  pneumatic  chemistry,  or  the  chemistry 
of  gases.  He  invented  the  pneumatic  trough  and  devised 
the  method  of  collecting  gases  over  liquids.  Simple  as 
this  method  seems  to  us  at  this  day,  it  required  to  be 
invented.  The  pneumatic  trough  is  a  vessel  provided 
with  a  perforated  shelf.  The  vessel  is  filled  with  water 
or  other  liquid  until  the  shelf  is  covered  ;  jars  or  bottles 
are  filled  with  the  liquid  in  the  trough,  inverted,  and  set 
upon  the  shelf.  These  jars  or  bottles  can  be  filled  with  any 
gas  by  allowing  it  to  rise  in  bubbles  through  the  liquid, 
which  it  displaces  downward.  In  1744  (August  1st), 
Priestly  heated  a  compound  of  mercury  known  as  red 
precipitate,  or  precipitate  per  se,  and  collected  a  colourless 
gas  which  was  drivsn  off.  (He  called  it  an  air,  the  name 
gas  not  being  used  until  later).  He  found  that  this  gas 
had  the  remarkable  property  of  causing  a  glowing  coal 
to  burst  into  flame.  For  reasons  to  be  presently  explained 
it  was  afterwards  called  oxygen. 

12.  Preparation  of  Oxygen.  —  The  most  conven- 
ient way  of  preparing  oxygen  is  by  heating  potassic 
chlorate  mixed  with  a  little  black  oxide  of  manganese. 


22-  —  -Grind  about  10  grains  of  potassic  chlorate 
to  powder  in  a  mortar.  Dry  it  by  heating  in  a  porcelain  basin, 
taking  care  not  to  fuse  it.  It  is  best  to  stir  constantly  with  a 


PREPARATION  AND  PROPERTIES  OF  OXYGEN.      33 

glass  rod  while  heating  it.  Dry  in  the  same  way  about  1  grain 
manganese  dioxide.  Heat  a  small  part  of  the  potassic  chlorate 
in  a  small  glass  tube  closed  at  one  end,  until  the  salt-like  sub- 
stance fuses  (melts),  and  bubbles  of  gas  rise  through  it.  Insert 
a  small  glowing  splinter,  and  observe  its  bright  flame.  Mix  the 
remainder  of  the  chlorate  thoroughly  with  the  manganese  dioxide, 
and  put  it  in  a  larger  hard-glass  tube  (test-tiibe  size)  provided 
with  a  gas  delivery  tube  leading  from  the  test-tube  and  dipping 
under  the  surface  of  water  in  the  pneumatic  trough.  The  tube 
passes  tightly  through  a  cork  in  the  mouth  of  the  test-tube. 
Arrange  the  test-tube,  &c.,  conveniently,  and  heat  at  the  bottom, 
applying  the  heat  gradually,  but  not  removing  the  lamp  when 
once  applied,  unless  the  open  end  of  the  delivery  tube  is  first 
lifted  out  of  the  water.  Otherwise,  as  the  heated  gas  in  the 
retort  (test-tube)  cools,  it  will  contract,  and  the  cold  water  will 
run  back  into  the  retort,  breaking  it.  Collect  the  gas  which 
bubbles  off,  until  about  as  much  as  would  fill  the  retort  and 
delivery  tube  is  collected,  then  change  the  receiver  (collecting 
vessel),  and  fill  5  small  bottles  with  the  gas.  (Why  reject  the 
first  portion  ?) 

Oxygen  can  also  be  prepared  by  heating  black  oxide  oj 
manganese  or  potassic  bichromate,  either  alone  or  with 
oil  of  vitriol.  The  black  oxide  of  manganese  used  in 
Experiment  22  undergoes  no  change.  It  enables  the 
oxygen  to  escape  at  a  lower  temperature.  This  is  a  case 
of  so-called  contact-action.  When  oxygen  is  prepared 
from  potassic  chlorate  it  is  generally  contaminated  by 
the  poisonous  gas,  chlorine,  from  which  it  should  be 
freed  by  bubbling  it  through  water,  in  which  chlorine 
is  very  soluble.  It  is  very  necessary  to  do  this  when 
the  oxygen  is  to  be  inhaled. 

43.   Properties  Of  Oxygen. — Oxygen    is   an    in- 
visible gas,  odourless,  tasteless,  and  a  little  heavier  than 
air.     Below  — 130°  C.  it  can  be  condensed  to  a  liquid. 
There  is  a  certain  temperature  for  each  gas,  above  which  it 
4 


34  COMBUSTION    IN    OXYGEN. 

cannot  be  condensed  to  a  liquid  by  pressure,  below  which 
it  can.  This  is  the  critical  temperature,  about  — 130°  C. 
for  oxygen.  Oxygen  is  a  "supporter  of  combustion" — 
many  substances  burn  in  it.  It  is  necessary  to  animal 
and,  generally,  to  vegetable  life.  It  is  obtained  from  the 
air  by  animals  when  breathing,  diffusing  through  the 
thin  walls  of  the  lungs  into  the  blood.  In  the  air, 
oxygen  is  diluted  with  four  times  its  volume  of  another 
gas,  nitroyen,  so  that  its  action  is  moderated.  An  at- 
mosphere of  pure  oxygen  would  throw  us  into  a  state  of 
excitement  and  fever,  and  would  cause  a  vast  conflagra- 
tion over  the  surface  of  the  earth.  It  was  once  often 
given  as  a  stimulant.  In  cases  of  asphyxia  and  dyspnoea 
it  is  valuable.  The  patient  obtains  tho  necessary  supply 
of  oxygen  in  a  much  smaller  space,  therefore  with  less 
breathing.  Oxygen  is  somewhat  soluble  in  water,  about 
4  c.  c.  in  100  c.  c.  of  pure  water.  When  a  glass  vessel 
filled  with  water  is  heated  slowly,  bubbles  of  gas  collect 
on  the  sides,  and  at  length  rise  to  the  surface.  These 
have  been  collected  and  proved  to  be  largely  oxygen. 
Fish  obtain  their  oxygen  by  diffusion  from  the  water. 
Hence,  the  necessity  for  a  current  of  water,  through  the 
mouth  and  out  over  the  gills.  When  fish  are  drowned 
they  are  in  reality  smothered.  Water  dissolves  gases 
according  to  a  certain  law.  The  quantity  of  any  gas 
which  is  dissolved  by  a  given  quantity  of  water  is  pro- 
port 'onal  to  the  pressure  and  inversely  as  the  temperature. 
This  law  does  not  hold  for  gases  which  unite  chemically 
with  water,  e.g.,  ammonia,  carbon  dioxide. 

44.  Combustion  in  Oxygen. — When  chemical 
action  is  accompanied  by  light  and  much  heat  it  is  called 
combustion.  Flame  also  usually  accompanies  combus- 


ACIDS.  35 

tion.  Ordinary  combustion  takes  place  in  an  atmos- 
phere which  unites  partially  at  least  with  the  substance 
burned.  It  is  a  supporter  of  the  combustion.  Oxygen 
is  the  supporter  of  combustion  in  the  case  of  wood  and 
coal  fires,  lamps,  &c. 

Experiment  23-— Fasten  a  small  piece  of  charcoal  to  the  end 
of  a  copper  wire  ;  heat  the  charcoal  until  it  glows,  and  then 
thnist  it  into  a  jar  of  oxygen.  It  glows  more  brightly  and 
bursts  into  flame.  Pour  a  few  drops  of  litmus  solution  into  the  jar, 
close  it  with  the  hand  and  shake  it.  The  blue  litmus  turns  red. 

Litmus  is  colouring  matter  extracted  from  a  lichen. 
It  turns  red  when  acted  on  by  sour  substances  (acids), 
e.g.,  citric  acid,  hydrochloric  acid,  and  vinegar.  We  can 
conclude  then  that  an  acid  is  formed  in  this  experiment. 
It  is  called  Carbonic  Acid. 

Experiment  24. — Set  fire  to  a  piece  of  sulphur  in  a  small 
long-handled  iron  cup,  and  then  plunge  it  into  a  jar  of  oxygen. 
The  sulphur  burns  with  a  dazzling  purplish  light,  and  suffocat- 
ing fumes  are  formed.  Shake  up  with  a  little  water,  and  note 
the  sourish  taste.  Add  a  few  drops  of  litmus  solution.  It  is 
reddened.  Sulphurous  Acid  has  been  formed. 

Experiment  25. — Cut  off  a  very  small  bit  of  phosphorus 
under  water,  dry  it  carefully  with  blotting  paper,  put  it  in  the 
small  cup  used  before  (deflagrating  spoon),  hold  it  in  a  jar  of 
oxygen,  and  set  fire  to  it  with  a  hot  wire.  It  burns  with  a 
bright  white  light,  and  white  fumes  are  formed.  Shake  up 
with  a  little  water,  and  note  the  sour  taste.  Test  with  litmus 
as  before.  Same  result.  Phosphoric  Acid  has  been  formed. 

Certain  substances  burn  in  oxygen  forming  compounds 
which  give  a  sour  taste  to  water  and  turn  blue  litmus 
red.  Oxygen  is  the  acid  generator  •  hence,  its  name, 
from  the  Greek.  To  oxidise  is  to  combine  with  oxygen, 
as  in  the  above  three  experiments.  Oxides  are  com- 


36  BASKS. 

pounds  of  oxygen  with  other  elements.  Thus,  in  the 
above  experiments,  oxides  of  carbon,  sxilphur,  and  phos- 
phorus are  formed.  The  greater  part  of  the  earth's 
crust  is  made  up  of  oxides  combined  in  various  ways. 
Oxygen  forms  compounds  with  all  the  elements  except 
fluorine.  To  reduce  an  oxide  is  to  deprive  it  of  a  part 
or  the  whole  of  its  oxygen,  i.e.,  to  lead  it  back  to  the 
original  condition  before  oxidation. 

Experiment  26. — Burn  a  small  piece  of  the  metal  sodium  in 
a  jar  of  oxygen  containing  a  little  water,  using  the  deflagrating 
spoon.  It  burns  with  a  bright  yellow  flame  forming  white  fumes 
and  a  grey  substance  which  remains  in  the  spoon.  Wash  the 
spoon  in  the  water,  and  note  the  soapy  feel  and  taste  of  the 
water.  Add  some  of  the  litmus  turned  red  in  Experiments  25 
and  26.  It  is  turned  blue  again.  The  same  results  can  be 
obtained  by  using  the  metal  potassium. 

The  substances  formed  by  burning  these  metals  in 
oxygen  in  the  presence  of  water  have  a  character  opposite 
to  that  of  the  acids.  They  are  called  bases.  That  formed 
from  sodium  is  called  caustic  soda.  Some  other  metals 
form  soluble  bases,  but  most  bases  are  insoluble  and  can 
only  be  pi'epared  indirectly  from  the  metals. 

Experiment  27. — Pour  a  few  drops  of  litmus  into  a  solution 
of  hydrochloric  acid  so  as  to  produce  a  distinct  red  colour. 
Then,  add  some  solution  of  caustic  soda,  dropping  it  from  a 
pipette  gradually,  and  stirring,  until  the  red  just  begins  to  turn 
blue,  i.e.,  becomes  purplish.  Observe  that  the  solution  now 
tastes  neither  sour  nor  soapy,  but  salt.  Both  acid  and  base  have 
disappeared ;  the  colour  is  neither  blue  nor  red  ;  the  solution  is 
neutral  in  its  action  on  litmus.  Evaporate  it  in  a  porcelain  dish. 
Crystals  of  common  salt  are  obtained. 

A  similar  result  can  be  produced  with  sulphuric  or 
nitric  acid,  and  caustic  potash.  In  each  case,  by  mixing 


TEMPERATURE    OF    IGNITION.  37 

solutions  of  the  acid  and  base  in  the  right  proportions  a 
point  is  reached  where  the  solution  turns  litmus  neither 
blue  nor  red,  both  sour  and  soapy  taste  disappear,  and  a 
salt-like  substance  can  be  separated  from  the  solution  by 
evaporation.  Substances  formed  by  the  action  of  acids  on 
bases  are  called  salts. 

Experiment  28. — Make  a  bundle  of  thin  strips  of  zinc  foil, 
dip  the  ends  in  a  little  melted  sulphur  ( Why  ?),  set  on  fire  and 
plunge  into  a  jar  of  oxygen.  The  zinc  burns  with  a  white  light, 
forming  zinc  oxide,  or  "  philosopher's  wool."  Pour  a  little  water 
into  the  jar,  and  observe  that  the  oxide  does  not  dissolve.  Add 
a  few  drops  of  sulphuric  acid  ;  the  oxide  dissolves  ;  evaporate  in 
a  porcelain  dish  ;  a  salt,  zinc  sulphate,  or  white  vitriol,  remains. 
Burn  some  fine  iron  wire  in  the  same  way. 

The  oxide  of  zinc  is  a  base-forming  oxide,  because  it 
acts  on  an  acid  to  form  a  salt ;  but  it  does  not  dissolve  in 
water  forming  a  base,  as  the  oxides  of  sodium  and  of 
potassium  do.  But  the  corresponding  base,  known  as 
zinc  hydroxide,  can  be  prepared  in  another  way.  It  is 
insoluble  in  water,  and  does  not  turn  red  litmus  blue,  nor 
feel  soapy  when  rubbed  between  the  fingers,  as  the  char- 
acteristic bases  do.  It  is  nevertheless  called  a  base, 
because  it  forms  salts  with  acids.  Iron  and  other  metals 
can  be  burned  in  oxygen.  They  all  produce  base-forming 
oxides  ;  but  it  is  to  be  noted  that  most  bases  have  not  the 
strongly  marked  characteristics  of  caustic  soda.  They  are 
mostly  insoluble  in  water,  and  do  not  turn  red  litmus  blue, 
but  they  are  all  acted  on  by  acids,  and  thereby  form  salts. 

45.  Temperature  of  Ignition. — It  is  necessary 

to  heat  combustible  substances  to  a  certain  tempei-ature 
before  they  will  catch  fire.  This  is  called  the  tempera- 
ture of  ignition.  It  is  very  different  for  different  sub- 


38  SLOW    COMBUSTION CHEMISM. 

stances.  Thus,  phosphorus  catches  fire  in  air  below 
100°  C.,  carbon  bisulphide  at  150°,  sulphur  at  400°, 
while  coal  and  wood  begin  to  burn  only  at  a  red  heat. 

46.  SlOW  Combustion. — Combustible    substances 
often  combine  slowly  with  oxygen  at  temperatures  below 
that  of  ignition.     In  such  cases  the  more  striking  ap- 
pearances of  combustion  are  absent.     Rusting  of  metals, 
respiration,  decay  of   animal  and  vegetable  substances, 
are  processes  of  this  kind,  modified  more  or  less  by  the 
action  of  water.     Heat  is  produced  in  slow  combustions 
(as  in  the  heating  of  grain,  of  damp  hay,  &c.),  but  more 
slowly  than  in  ordinary  combustions,  so  that  it  is  con- 
ducted away  by  surrounding  objects  and  does  not  raise 
the    temperature    so    much.     In   some    cases,    however, 
where  the  substance  is  finely  divided,  exposing  a  large 
surface  to   the  action  of  the  air,  oxidation  goes  on  so 
rapidly  as  to  heat  the  mass  up  to  the  temperature  of 
ignition.     Thus  occur  cases  of  spontaneous  ignition  and 
combustion.     Rags  soaked  with  oil,  and  coal  containing 
much  sulphur,  have  been  known  to  catch  fire  in  this  way. 

47.  Chemism. — That  which  causes  oxygen  to  unite 
with   carbon,    sulphur,   &c.,   is    called    chemical  affinity, 
chemical  attraction,  or  chemism.     Oxygen    may  be  com- 
pressed with  a  pressure   of  several  tons  to  the  square 
inch  without  liquefying  it,  and  yet,  when  it  combines 
with  iron  it  becomes  solid.     Chemical  attraction  must  be 
very  powerful  between  oxygen  and  iron.     Between  some 
of  the  elements  the  attraction  seems  to  be  very  slight  or 
even  wanting.     Thus,  oxygen  forms  no  compound  with 
fluorine,    and    cannot    be    got    to    unite    directly    with 
chlorine,  bromine,  iodine,  or  gold.     As  a  rule  elements 


THE    ELEMENTS. 


39 


which   are  unlike  have  the  greatest  attraction  for  each 
other ;  for  example,  the  metals  and  the  non-metals. 

48.  Metals  and  Non-Metals.  —  All  the  sub- 
stances described  as  burning  in  oxygen  to  form  oxides 
are  elements ;  observe  that  some  of  them,  as  sulphur, 
carbon,  and  phosphorus,  produce  acid-forming  oxides, 
while  others,  sodium,  potassium,  zinc,  iron,  produce  base- 
forming  oxides.  Elements  which  have  base-forming 
oxides  are  called  metals  ;  those  whose  characteristic  oxides 
are  acid-forming  are  called  non-metals.  Of  the  07  elements 
at  present  known,*  52  are  metals  and  15  non-metals  ;  but 
some  of  the  elements  have  the  characters  of  both  metals  and 
non-metals,  so  that  it  is  difficult  to  decide  to  which  class 
they  belong.  Thus,  antimony  is  classed  by  some  chemists 
among  the  metals,  by  others  with  the  non  metals. 

TABLE  OF  THE  ELEMENTS. 
(Non-metals  in  heavy  type,  imperfect  metals  in  italics.) 


English 
Names. 

Latin  Names. 

Symbols. 

Atomic 
Weights. 

t  Specific 
Weights. 

Remarks. 

Aluminium.  .  .  Aluminium  .  .  . 
Antiiiiuiiij  ..  .  .  Stibium  
Arsenic  Arsenicum.  ... 
Barium  Barium  

Al  v 
Sb  «» 
As  »  » 
Ba  i 
Be  i 
Bii  i  v 
B" 
Br    v 
Co" 
Csi 

Caii 

CiT 

Ceiii 

Cl  i  iii  v 
Crii  iv  vi 
Co  ii  iv 

27.3 
122. 
75  (74.9) 
136.8 
9. 
210. 
11. 
80  (79.75) 
111.6 
132.56 

40. 
12  (11.97) 
141.2 
35.37 
52.4 
58.6 

2.6 
6.71 
5.73 
3.75 
2.07 
9.8 
2.5 
3.19 
8.6 
1.88 

1.57 
3.52  (diam'd) 
6.7 
2.45 
6.5 
8.5 

Solid. 

Gas. 
Liquid. 
Solid. 
"     (melts 
at  32") 

Gas. 
Solid. 

Beryllium  .... 
Bismuth  
Boron  
Bromine  — 
Cadmium  
Caesium  

Calcium  
Carbon  
Cerium  

Beryllium  
Bismuthum  .  .  . 
Borium  
Bromum  
Cadminra  
Caesium  

Calcium  
Carbo  
Cerium  

Chlorine  — 

Chromium.  .  .  . 
Cobalt  

Chlorinum  .  .  . 
Chromium.  .  .  . 
Cobaltum  

*  New  elements  are  discovered  from  time  to  time. 

t  For  the  specific  weights  of  solids  and  liquids  water  at  4°  C.  is  the  standard  ; 
for  those  of  gases,  air. 


40  THE    ELEMENTS. 

TABLE  OF  THE  ELEMENTS— Continued. 


English 
Names. 

Latin  Names. 

Symbols. 

Atomic 
Weights. 

Specific 
Weights. 

Remarks. 

Copper  

Cuprum  

Cu  ii 

63 

8  88 

Solid 

Didymium.  .  .  . 

Jidyniium.  .  .  . 

Diiii 
Eriii 

142.3 
170  5 

6.54 

Fluorine  .  . 
Gallium  

Quorum  
Gallium  

Fi 
G'T 

19.1 
69.86 

1.31 
5.95 

3  as  prob'ly. 
Solid  (melts 

Gold  

Aurum  

All  1  "i 

196  5 

19  32 

at  30°) 

Hydrogen  .. 
Indium  
Iodine  
Iridium  
Iron  

rlydrogenum  . 
ndium  
odum  
ridium  
Ferrum  

Hi 
IniT 

liiiir 
Irii  iv  vi 
Fe  ii  iv 

1. 
113.6 
126.5 
193. 
56. 

0.0692 
7.42 
4.948 
22.42 
7  86 

Gas. 
Solid. 

Lanthanum.  .  . 
Lead  
Lithium  
Magnesium.  .  . 
Manganese  .  .  . 

lanthanum  .  .  . 
Plumbum  
Lithium  
Magnesium  ... 
Manganesium. 

La»i 

Pb  ii  iv 

Li' 
Mg« 

Mn  i>  iv  vi 

HIT  ii 

139. 

206.4 
7. 
24. 
54.8 
200  (199  8) 

6.1 
11.35 
0.59 
1.74 
8.03 
13  596 

Molybdenum.  . 
Nickel  

Molybdenum  . 

Moii  iv  iv 
Ni  ii  iv 

96. 

58  6 

8.6 
S  9 

Solid. 

Niobium  
Nitrogen  — 
Osmium  ....... 
Oxygen  
Palladium  
Phosphorus  . 
Platinum  
Potassium  .... 

Niobium  
Nitrogenum  .  . 
Osmium  
Oxygenum  
Palladium.  .  .  . 
Phosphorus.  .  . 
Platinum  
Kalium  

Nbv 

Niii  v 
OS"  iv  vi 

on 

Pd  ii  iv 

Piii  v 
Ptii  iv 
K  1 

94. 
14  (14.01) 
199. 
16  (15.96) 
106.2 
31. 
196.7 
39 

7.06 
0.971 
22.48 
1.105 
11.4 
1.83-2.2 
21.5 
0.87 

Gas. 
Solid. 
Gas. 
Solid. 

Rhodium  

Rhodium  

Rhii  iv  vi 

104.1 

12.1 

K 

Rubidium  .... 
Ruthenium.  .. 
Samarium.  .  .  . 

Rubidium..  .. 
Ruthenium  .. 
Samarium.  .  .  . 

Rbi 

Ru  ii  >v  vi 

Sm 

85.47 
103.5 
150. 

1.52 
12.26 

" 

Sc 

44 

n 

Selenium  — 
Silicon  

Selenium  
Silicium  

Se»iTV. 

Si" 

79. 

28 

4.5 
2.39 

" 

Silver  

Argentum.  .  .  . 

AgJ 

107  66 

10.47 

,, 

Sodium    .  . 

Natrium  

Nal 

23 

0.978 

n 

Strontium  .... 
Sulphur  
Tantnlum  .... 
Tellurium  .. 
Terbium  

Strontium  .... 
Sulphurum.  .. 
Tantalum  
Tellurium  
Terbium  

Sri' 

g  ii  iv  vi 

Tav 

Te  ii  iT  vi 

Tb 

87.2 
32. 
182. 
125.? 

2.54 
2.03 
10.4 
6.4 

! 

Thallium 
Thorium  
Tin  

Thallium  
Thorium  
Stannum  

Tl  i  iii 

Th> 

Sn  «  " 

203.6 
231.5 
117.8 

11.85 
11. 
7.29 

• 

Titanium  .... 

Titanium  

Ti  ii  iv 

48. 

i 

Tungsten  
Uranium  
Vanadium  
Yttrium  

Wolframum  .. 
Uranium  
Vanadium 

Wirvi 

Uirvi 

V«IT 

Yil" 

184. 
239.8 
51.5 
89.5 

19.12 
18.7 
5.5 

« 

Zinc  
Zirconium.  .  .  . 

Zincum  
Zirconium..  .. 

Zn» 
Zr" 

65. 
90. 

7. 
4.15 

•' 

CONSERVATION    OF    MATTER.  41 


CHAPTER    V. 


LAWS   OF   COMBINATION— ATOMIC   THEORY. 

49.  Conservation  Of  Matter. — In  order  to  in- 
vestigate fully  the  chemical  actions  described  in  connec- 
tion with  oxygen  it  is  necessary  to  weigh  and  measure 
the  substances.  This  has  been  done  with  the  greatest 
accuracy,  and  the  experiments  have  been  repeated  in  a 
variety  of  ways  with  materials  from  different  sources. 
The  results  which  have  been  obtained  are  always  the 
same  for  the  same  chemical  action.  Thus,  8  grams  of 
oxygen  and  1  gram  of  hydrogen  are  always  obtained  by 
the  decomposition  of  9  grams  of  water ;  8  grams  of 
oxygen  and  100  grams  of  mercury,  by  the  decomposition 
of  108  grams  of  red  oxide  of  mercury;  and  8  grams  of 
oxygen  and  40.8  grams  of  potassic  chloride  from  48.8 
grams  of  potassic  chlorate.  We  see  that  in  chemical 
decompositions  the  sum  of  the  weights  of  the  products  of 
decomposition  is  always  equal  to  the  weight  of  the  substance 
decomposed.  Again,  8  grams  of  oxygen  burn  3  grams 
of  charcoal  to  form  1 1  grams  of  carbon  dioxide  ;  8  grams 
of  oxygen  burn  8  grams  of  sulphur  to  form  16  grams  of 
sulphur  dioxide;  8  grams  of  oxygen  burn  6.2  grams  of 
phosphorus  to  form  14.2  grams  of  phosphorus  pent- 
oxide;  8  grams  of  oxygen  bum  33.1  grams  of  potassium 
to  form  47.1  grams  of  potassic  oxide  ;  8  grams  of  oxygen 
burn  23  grams  of  sodium  to  form  31  grams  of  sodic 
oxide  ;  8  grams  of  oxygen  burn  32.5  grams  of  zinc  to 
form  40.5  grams  of  zinc  oxide.  From  this  statement  it 


42  DEFINITE    PROPOKTIONS EQUIVALENTS. 

is  seen  that  in  chemical  combinations,  also,  the  sum  of 
the  weights  of  the  substances  combining  is  always  equal 
to  the  weight  of  the  substance  formed.  It  may  thus  be 
stated  generally  that  in  chemical  actions  no  matter  is 
lost — put  out  of  existence ;  nor  is  there  any  gain  of 
matter.  The  weights  of  the  elements  entering  into  a 
chemical  action  remain  the  same  at  the  end  of  the 
action  ;  but  the  elements  become  differently  arranged. 
This  law,  founded  on  experiments,  is  the  Law  of  the  Con- 
servation of  Matter.  It  is  the  fundamental  law  of  the 
Science  of  Chemistry. 

50.  Definite  Proportions. — Another  law  can  be 
deduced  from  the  facts  stated  in  Art.  49.     It  is  found 
that  water    is  always  composed  of  hydrogen  and    oxy- 
gen  in  the  proportion   of  1   to   8,   that  mercuric  oxide 
is  always  composed  of  oxygen  and  mercury  in  the  pro- 
portion of  8  to  100,  and  so  with  the  others.     Stated  in 
general  terms  this  is  the  Law  of  Definite  Proportions  : 
Each  chemical  compound  is  always  composed  of  the  same 
elements  and  these  are  always  in  the  same  proportion  in 
this  compound.     In  this  respect  chemical  compounds  differ 
from  mixtures,  which  may  be  made  in  any  proportions. 

51.  Combining  Weights —Equivalents.  -The 

weight  of  oxygen  which  combines  with  1  part  by  weight 
of  hydrogeii  to  form  water  is  called  the  combining  weight 
of  oxygen.  The  weight  of  any  element  which  combines 
with  I  part  of  hydrogen  is  its  combining  weight ;  but,  as 
many  of  the  elements  do  not  form  compounds  with 
hydrogen,  their  combining  weights  must  be  ascertained, 
as  it  were,  at  second  hand.  Thus,  the  weight  of  zinc 
which  combines  with  8  parts  by  weight  of  oxygen  (its 


MULTIPLE    PROPORTIONS.  43 

combining  weight)  is  the  combining  weight  of  zinc. 
The  expression  equivalent,  or  equivalent  weight,  is  used 
in  exactly  the  same  sense  as  combining  weight  when 
applied  to  the  elements  ;  but  it  is  also  applied  to  com- 
pounds. Thus,  the  equivalent  of  water  is  the  quantity 
of  water  which  will  produce,  when  decomposed,  1  part  by 
weight  of  hydrogen,  i  e.,  9.  Again  the  equivalent  of  zinc 
oxide  is  40.5,  since  this  weight  of  zinc  oxide  contains 
8  parts  (an  equivalent)  of  oxygen.  The  equivalent  of  an 
element  or  a  compound  is  the  weight  oj  the  element  or  com- 
pound equal  in  value  in  chemical  action  to  1  part  by 
weight  of  hydrogen.  (Calculate  the  equivalents  of  all 
the  substances  mentioned  in  Art.  49.)  Hydrogen  is 
here,  as  in  most  cases,  taken  as  the  standard  element. 

52.  Multiple  Proportions. — When  charcoal  (car- 
bon) burns  in  a  free  supply  of  oxygen,  8  parts  of  oxygen 
unite  with  3  of  carbon  to  form  carbon  dioxide ;  but  if 
the  supply  be  scant,  as  in  the  centre  of  a  fire,  only  4 
parts  of  oxygen  unite  with  3  of  carbon,  and  carbon  mon- 
oxide is  formed.  Thus,  carbon  combines  with  oxygen  in 
two  proportions,  and  these  are  as  8  to  4,  or  as  2  to  1. 
When  phosphorus  burns  in  oxygen  it  also  forms  two 
compounds,  the  pentoxide  when  the  oxygeii  is  plentiful, 
and  the  trioxide  when  it  is  scant.  In  the  pentoxide  6.2 
parts  of  phosphorus  (Why  use  this  particular  number1?) 
are  combined  with  8  of  oxygen ;  in  the  trioxide  this 
weight  of  phosphorus  is  combined  with  4.8  of  oxygen. 
The  quantities  of  oxygen  are  as  5  to  3.  Sulphur  also 
forms  two  compounds  with  oxygen,  the  relative  quanti- 
ties of  which  are  as  2  to  3.  Many  other  examples  could 
be  given  of  an  element  combining  with  another  in  two 
or  more  proportions  ;  and  it  is  noticed  that  in  every  such 


44  THE    ATOMIC    THEORY. 

case  these  proportions  have  a  simple  arithmetical  relation 
to  each  other,  as  in  the  above  cases.  This  is  stated  in 
the  Law  of  Multiple  Proportions  :  If  two  elements  form 
several  compounds  with  each  other,  then  the  .different 
quantities  of  the  first  which  combine  with  affixed  quantity 
of  the  second  bear  a  simple  ratio  to  each  other,  as  1  to  2, 
2  to  3,  3  to  5,  &c.  (Apply  this  law  to  the  examples 
given.) 

53.  The  Atomic  Theory.— To  explain  these  re- 
markable facts  Daltori  ( 180 1  modified  the  atomic  theory 
of  the  ancient  Greeks.  According  to  this  theory  in 
its  modern  form  the  elements  are  made  up  of  indivisible, 
extremely  minute  particles — atoms.  (What  is  a  theory  1) 
The  atoms  of  the  same  element  are  exactly  alike,  but 
those  of  different  elements  differ  in  weight,  &c.  The 
atoms  of  different  elements  unite  to  form  molecules  of 
compounds,  and  this  union  always  takes  place  1  atom 
with  1  atom,  1  atom  with  2,  2  with  3,  *fec.  For  ex- 
ample, when  zinc  burns  in  oxygen,  1  atom  of  zinc  com- 
bines with  1  atom  of  oxygen,  1,000  atoms  of  zinc  with 
1,000  of  oxygen  and  so  on,  until  the  whole  of  the  zinc  is 
burned,  when,  of  course,  the  oxide  which  is  formed  con- 
tains equal  numbers  of  atoms  of  zinc  and  of  oxygen.  It 
follows  from  this  that  if  the  ratio  of  the  weights  of  the 
atoms  of  zinc  and  oxygen  are  as  65  to  16,  the  ratio  of 
the  whole  quantities  combining  must  be  in  this  propor- 
tion. As  a  matter  of  fact,  zinc  oxide  is  always  com- 
posed of  the  two  elements  in  this  ratio.  In  this  way 
the  atomic  theory  explains  the  law  of  definite  propor- 
tions. The  explanation  of  the  law  of  multiple  propor- 
tions is  equally  simple.  For  example,  sulphur  unites 
with  oxygen  in  two  proportions.  According  to  the 


AVOGADKO'S    LAW.  45 

atomic  theory  the  audition  of  oxygen  to  sulphur  takes 
place  always  by  atoms,  never  parts  of  an  atom ;  in  the 
two  compounds  different  numbers  of  atoms  of  oxygen 
unite  with  an  atom  of  sulphur,  in  one  case  2,  and  in  the 
other  3.  From  this  it  follows  that  the  weights  of  oxygen 
uniting  with  the  atomic  weight  of  sulphur  are  respec- 
tively twice  and  three  times  the  atomic  weight  of  oxygen. 

54.  Avogadro's  Law. — If  equal  volumes  of  hydro- 
gen, oxygen,  marsh  gas,  and  other  gases  generally  are  put 
in  graduated  tubes  over  mercury  and  heated  equally,  they 
are  observed  to  expand  equally,  viz.  :  ^fa  °f  their  volume 
at  Oc  C.  for  every  degree  rise  in  temperature.  Also,  if  the 
same  increase  of  pressure  be  put  on  each  gas,  the  volumes 
are  diminished  equally.  These  facts  point  to  some  close 
similarity  in  equal  volumes  of  gases.  The  hypothesis 
put  forward  by  Avogadro  in  1811  is  that  equal  volumes 
of  all  gases  under  the  same  conditions  of  temperature 
and  pressure  contain  the  same  number  of  particles  (mole- 
cules). This  can  also  be  deduced  mathematically  from 
the  molecular  theory  of  gases. 

55.  Combination  by  Volumes. — When   water 

is  decomposed  by  electricity  (Art.  39)  two  volumes  of 
hydrogen  are  set  free  for  one  of  oxygen.  If  the  gases  be 
mixed  and  caused  to  unite  again,  water  is  formed,  and 
none  of  either  gas  is  left  over.  It  has  been  found  that 
in  all  cases  where  gases  combine  with  each  other,  the 
volumes  combining  have  a  simple  ratio  to  each  other,  as 
1  to  2,  2  to  3,  &c. 

56.  Molecules  and  Atoms. — Two  litres  of  hydro- 
gen unite  with  one  litre  of  oxygen  to  form  water.     If 
the   union  takes  place  at,  say  2UO°  0.,  the  water  is  in 


46  MOLECULES    AND    ATOMS. 

the  gaseous  state.  It  is  then  found  that  the  volume  of 
steam  formed  is  the  same  as  the  volume  of  hydrogen 
used,  and  according  to  Avogadro's  law  contains  the 
same  number  of  particles  (Dalton's  atoms).  Let  us 
suppose  that  the  two  litres  of  hydrogen  contain  two 
millions  of  particles ;  then  one  litre  of  oxygen  contains 
one  million  particles,  and  two  millions  particles  of  steam 
are  formed.  Each  particle  of  steam  contains  oxygen,  so 
that  the  oxygen  must  be  divided  into  two  millions  of 
parts  when  combination  takes  place.  Therefore,  the 
particles  of  oxygen  must  be  divisible.  They  are  not 
atoms,  but  pairs  of  atoms,  which  cannot  be  separated 
except  by  chemical  action.  They  are  called  molecules, 
(little  masses).  The  molecules  of  elements  consist  (with 
few  exceptions)  of  two  like  atoms ;  and  those  of  com- 
pounds, of  two  or  more  unlike  atoms.  The  molecules  of 
substances  are  imagined  as  in  constant  motion  of  some 
sort,  the  motion  being  different  in  solids,  liquids,  and 
gases.  In  solids  the  motion  is  mostly  vibration,  the 
molecules  moving  about  particular  points  ;  in  liquids, 
there  is  a  comparatively  slow  motion  from  place  to  place ; 
and  in  gases  this  motion  from  place  to  place  goes  on 
with  great  velocity,  the  molecules  moving  faster  as  the 
gases  become  hotter.  It  is  the  battering  of  the  mole- 
cules of  an  enclosed  gas  against  the  walls  of  the  enclos- 
ing vessel  which  produces  pressure. 

Definitions. — A  molecule  is  the  smallest  portion  of  any  sub- 
stance which  can  exist  by  itself.  Imagine  a  drop  of  water  divided 
and  subdivided.  A  point  is  at  last  reached  when  no  further 
division  is  possible  by  mechanical  means  ;  and  if  the  division  is 
forced  by  electricity  or  heat,  it  is  no  longer  portions  of  water 
which  are  obtained,  but  two  different  substances — oxygen  and 
hydrogen.  At  this  point  we  have  reached  the  molecule  of 


MOLECULAR    WEIGHTS    OF    GASES.  47 

water.  An  atom  is  the  smallest  portion  of  an  element  which  can 
exist  in  a  molecule.  The  atom  cannot  be  divided  even  by  chemi- 
cal action  ;  it  must  enter  into  combination  as  a  whole. 

57.  Molecular  Weights  of  Gases. — Since  equal 

volumes  of  gases  contain  the  same  number  of  molecules, 
it  follows  that  the  weights  of  equal  volumes  of  gases  are 
in  proportion  to  the  weights  of  single  molecules.  Thus, 
if  a  certain  volume  of  hydrogen  weighs  2  grams,  if  the 
same  volume  of  oxygen  weighs  32,  of  nitrogen  28,  and 
of  carbon  dioxide  44,  these  numbers  may  be  taken  to 
represent  the  weights  of  the  molecules,  or  the  molecular 
weights  of  the  gases.  In  the  case  of  the  elementary 
gases  half  the  molecular  weight  gives  the  weight  of  the 
atom,  or  the  atomic  weight  (why  half?).  In  estimating 
the  specific  weights  of  gases,  the  weights  of  equal  volumes 
are  compared,  one  gas  being  taken  as  standard,  and  its 
weight  taken  as  1.  Hydrogen  is  now  usually  taken  as 
the  standard  gas,  and  it  can  readily  be  shown  that  the 
specific  weight  of  a  gas  is  half  its  molecular  weight. 
(Explain  this.) 

58.  Chemical  Notation— Atomic  Weights.— 

Chemical  symbols  are  letters  used  to  denote  the  atomic 
weights  of  the  elements.  They  are  the  initial  letters  of 
the  Latin  names.  Where  two  or  more  names  begin  with 
the  same  letter,  a  second  is  added  to  distinguish.  For 
example  H  denotes  1  part  by  weight  of  hydrogen  ;  O,  16 
parts  by  weight  of  oxygen ;  C,  1 2  parts  by  weight  of 
carbon  ;  Cl,  35.37  parts  by  weight  of  chlorine,  &c.  These 
numbers  are  obtained  by  experiments,  by  weighing  and 
measuring.  For  example,  it  has  been  found  that  a  litre 
of  oxygen  weighs  16  times  as  much  as  a  litre  of  hydrogen, 
from  which  it  is  concluded  that  a  molecule  of  oxygen 


48  CHKMICAL    NOTATION. 

weighs  Hi  times  as  much  as  a  molecule  of  hydrogen  (Art. 
54),  and  that  the  atom  of  oxygen  must  also  weigh  16  times 
as  much  as  the  atom  of  hydrogen  ( why  ]).  If,  then,  the 
atom  of  hydrogen  be  assumed  to  weigh  1  (1  what1?),  the 
atom  of  oxygen  will  weigh  16.  Again,  16  parts  of 
oxygen  combine  with  65  of  zinc,  and  there  is  evidence 
that  in  this  combination  one  atom  of  oxygen  unites  with 
one  of  zinc.  Hence  it  is  concluded  that  the  atomic  weight 
of  zinc  is  65.  By  similar  methods  the  atomic  weights  of 
all  the  elements  have  been  determined.  There  is  a  method 
of  determining  the  atomic  weights  which  depends  on. 
measuring  the  specific  heats  of  the  elements.  It  was  dis- 
covered by  Dulong  and  Petit  that,  if  the  atomic  weiijht  of 
each  element  is  "multiplied  by  its  specific  heat,  the  product 
is  always  the  same  number  (nearly),  that  is,  6.6.  From 
which  it  follows  that  6.6  divided  by  the  specific  heat 
of  any  element  gives  its  atomic  weight.  This  number  is 
called  the  specific  heat  of  the  elements,  and  Dulong  and 
Petit's  law  may  be  stated  thus  :  The  specific  heat  of  the 
elements  is  a  constant  quantity. 

The  symbols  of  the  atoms  are  combined  into  formulas 
representing  molecules.  Thus  (Art.  56),  since  the  mole- 
cule of  water  contains  2  atoms  of  hydrogen  and  1  of 
oxygen,  its  formula  is  H2O.  This  formula  is  used  to 
express  the  following  :  — 

1.  Water  is  a  compound  of  the  two  elements,  hydrogen  and 
oxygen,  in  the  proportion  by  weight  of  2  to  16.     This  is  estab- 
lished by  experiment. 

2.  The  molecule  of  water  is  made  up  of  3  atoms,  viz. :  2  of  hy- 
drogen and  1   of   oxygen.      This   is  deduced   from  the  atomic 
theory,  which  is  now  strongly  supported  by  experimental  evi- 
dence. 


CHEMICAL    EQUATIONS.  49 

Similai-ly,  the  molecule  of  oxide  of  sodium  is  repre- 
sented by  Na2O  ;  of  oxide  of  zinc,  by  ZnO  ;  of  sulphuric- 
acid,  by  H2SO4,  and  so  on.  In  these  formulas  the  small 
figure  below  the  line  multiplies  the  symbol  immediately 
befoi-e  it.  Thus,  the  molecule  of  sulphuric  acid  contains 
2  atoms  of  hydrogen,  1  atom  of  sulphur,  and  4  atoms  of 
oxygen.  Very  often  the  formulas  are  used  loosely  for 
the  names  of  compounds,  but  it  must  be  always  remem- 
bered that  they  mean  more  than  this.  For  example, 
H2O  means  18  parts  by  weight  of  water,  and  not  simply 
water,  or  any  quantity  of  water.  The  formulas  of  sub- 
stances are  used  to  express  their  molecular  weights,  and, 
given  the  formula  of  a  compound,  its  molecular  weight  is 
found  by  adding  together  the  weights  of  all  the  atoms  in 
the  molecule.  For  example,  the  molecular  weight  of 
sulphuric  acid  is  (2  x  1)  +  32  +  (4  x  16)  =  98.  To 
represent  2,  3,  &c.,  molecules  the  number  is  prefixed  to 
the  formulas  thus,  2ZnO,  3H2O,  4H2SO4,  &c.  Or,  a 
formula  may  be  multiplied  thus,  (H20)3,  (NaCl)5. 

59.  Chemical  Equations. — These  are  short-hand 
representations  of  the  quantities  of  the  substances  taking 
part  in  chemical  actions.  Before  they  can  be  written, 
chemical  actions  must  be  studied  by  weighing  and 
measuring.  For  example,  the  decomposition  of  water 
by  electricity  is  represented  as  follows  : — 2H.2O  =  2H2 
+  02,*  a  short-hand  method  of  stating  that  36  parts  by 


*It  might  be  written  more  simply  H2O  =  H2  +  O,  but  since,  according  to 
theory,  the  molecule  is  the  smallest  portion  capable  of  existing  free,  it  is  by 
many  considered  advisable  not  to  represent  single  atoms  in  equations.  How- 
ever, where  weights  alone  are  being  considered,  this  adherence  to  theory  is 
unnecessary  and  often  cumbrous ;  and,  except  when  considering  the  volumes 
of  gases,  I  shall  always  write  the  equations  in  the  simplest  form. 
5 


50  CHEMICAL    CALCULATIONS. 

weight  (or  2  molecules)  of  water  decompose  into  4  parts 
by  weight  (2  molecules)  of  hydrogen,  and  32  parts  by 
weight  (1  molecule)  of  oxygen. 

(1)  The  equation  should  represent  all  the  substances 
taking  part  in  the  action,  in  the  proportions  in  which  they 
act  on  each  other  or  are  produced  in  the  action.  (2)  The 
sum  of  the  weights  on  the  left  side  of  the  sign  (=)  should 
be  equal  to  the  sum  of  those  on  the  right.  (3)  The  num- 
ber of  atoms  of  each  element  represented  on  the  left 
should  be  the  same  as  that  on  the  right. 

The  decomposition  of  potassic  chlorate  into  oxygen 
and  potassic  chloride  is  represented  by  the  following 
equation  :— 2KC1O3  =  3O2  +  2KC1.  That  is  2  x  122.6 
grams  of  potassic  chlorate  decomposes  into  3x22  grams 
of  oxygen  and  2  x  74.6  grams  of  potassic  chloride.  (Is 
it  necessary  then  to  use  245.2  grams  of  potassic  chlorate 
in  preparing  oxygen  1) 

60.  Chemical  Calculations. — From  a  chemical 
equation  we  can  calculate  readily  the  proportions  of  the 
substances  represented  in  the  equation  ;  and  equations 
are  useful  as  permanent  records  of  chemical  actions,  to 
which  we  can  refer  when  we  wish  to  know  in  what  pro- 
portions to  use  chemical  substances  in  any  reaction. 
Thus,  suppose  it  is  desired  to  prepare  100  grams  of 
phosphorus  pentoxide  by  burning  phosphorus  in  air. 
The  equation  representing  this  action  has  been  made 
out  as  follows: — 2P  +  5O  =  P2O5.  Referring  to  the 
table,  p.  40,  we  see  that  P  represents  31  parts  by 
weight  of  phosphorus.  Also,  P2O5  represents  62  +  80 
=  142  parts  by  weight  of  phosphorus  pentoxide;  and 
62  grams  of  phosphorus  burn  to  form  142  grams  of 


VOLUMES    OF    GASES. 


51 


phosphorus  pentoxide  ;  and  142  :  100  :  :  02  :  43.62, 
or  43.62  grams  of  phosphorus  burn  to  form  100  grams 
of  pentoxide. 

Gases  are  usually  measured  instead  of  weighed.  But, 
since  equal  volumes  of  all  gases  (at  the  same  temperature 
and  pressure)  contain  the  same  number  of  molecules, 
weights  of  gases  in  the  ratio  of  the  molecular  weights 
are  the  weights  of  equal  volumes  of  the  gases.  The 
molecular  weights  of  hydrogen  and  of  oxygen  are  2  and 
3 '2  respectively;  therefore,  2  grams  of  hydrogen  will 
measure  the  same  as  32  grams  of  oxygen  ;  and  so  for  other 
gases.  The  measurement  has  been  made  with  the  greatest 
care  many  times,  and  it  is  found  that  2  grams  of  hydro- 
gen gas  measure  (at  0°  C.  and  760  millimetres  of  mercury 
pressure)  22.33  litres.  The  molecular  weight  in  grams, 
or  g  ram-molecule,  of  any  gas  measures,  at  0°  C.  and 
760  mm.,  22.33  litres. 


GAS. 

Formula. 

Molecular 
Weight. 

Volume  of 
Gram- 
Molecule. 

Hydrogen  

H. 

2 

•\ 

Oxygen  .. 

O., 

32 

| 

Nitrogen  

ML 

28 

22.33 

Ammonia  

NH3 

17 

|     litres. 

Carbon  Dioxide  

CO, 

44 

J 

Marsh  Gas  

CH, 

16 

&c. 

&c. 

&c. 

Calculate  the  volume  of  oxygen  formed  by  heating  50 
grams  potassic  chlorate.     The  equation  is  : 


2KC1O3 

245.2  grams. 


=     2KC1      +      3O2 

149.2  grams.      3  X  22.33  litres. 


52  OZONE. 

That  is,  245.2  grams  potassic  chlorate  give  66.99  litres 
of  oxygen.  Therefore,  50  grains  give  ^fjj.^  x  66.99  = 
13.66  litres. 

In  making  such  a  calculation  we  simply  assume  that 
what  is  true  for  the  molecules  as  represented  in  the 
equation,  is  true  for  weights  which  are  taken  propor- 
tionately to  the  weights  of  the  molecules. 

How  much  sulphur  can  be  burned  by  10  litres  of 
oxygen  measured  at  0°  C.  and  760  mm.  pressure1?  The 
equation  is  : 

S   +  O2  =  S02 

32  grams— 22.33  litres. 

That  is,  32  grams  of  sulphur  use  up  22.33  litres  of  oxy- 
gen. Then,  22.33  :  10  ::  32  :  14.33  ;  i.  e.,  14.33  grams 
of  sulphur  use  up  10  litres  of  oxygen.  What  volume  of 
sulphur  dioxide  is  formed  in  this  action  ]  On  looking  at 
the  equation  we  see  that  a  molecule  of  sulphur  dioxide  is 
formed  from  a  molecule  of  oxygen  ;  therefore,  the  volume 
of  sulphur  dioxide  is  the  same  as  that  of  oxygen,  viz.,  10 
litres. 

61.  Ozone. — This  is  a  peculiar  modification  of  oxy- 
gen produced  chiefly  by  the  action  of  electricity.  It  is  a 
gas  of  a  penetrating  odour.  It  is  one  and  a  half  times 
as  heavy  as  oxygen,  and  its  molecule  must  therefore  con- 
tain three  atoms  (O3),  if  that  of  oxygen  contains  two  (O2). 
Ozone  has  very  powerful  oxidizing  properties,  destroy- 
ing paper,  india-rubber,  &c.,  by  rapid  oxidation.  It  is 
slightly  soluble  in  water,  more  so  in  turpentine.  It  is 
held  to  be  a  valuable  means  of  disinfecting  and  deodoriz- 
ing the  atmosphere,  and  is  supposed  to  exercise  a  bene- 


OZONE.  53 

ticial  influence  on  the  animal  body.  Of  late,  however, 
some  investigators  associate  the  prevalence  of  certain 
diseases  of  the  respiratory  passages  with  an  excessive 
quantity  of  ozone  in  the  air. 

Ozone  is  produced  by  the  passage  of  electricity  through 
air  (or  oxygen) — for  example,  by  flashes  of  lightning.  It 
is  more  abundant  in  pure  than  in  impure  air  ;  in  the  air  of 
the  country  than  in  that  of  the  towns  ;  and  in  the  higher 
than  in  the  lower  strata  of  the  atmosphere.  Its  quantity 
is  said  to  be  smaller  where  cholera  and  other  epidemics 
are  prevalent.  Ozone  is  also  produced  in  small  quantity 
during  the  electrolysis  of  water,  and  during  the  slow 
oxidation  of  phosphorus  in  moist  air.  It  is  also  formed 
about  an  electrical  machine  when  in  motion. — If  a  strip 
of  filter  paper  soaked  in  starch  and  solution  of  potassic 
iodide  be  brought  into  an  atmosphere  containing  ozone, 
the  paper  turns  blue.  Iodine  is  set  free  by  the  ozone,  and 
unites  with  the  starch  forming  the  blue  iodide  of  starch. — 
When  an  element  occurs  in  different  modifications,  these 
are  called  allotropic  modifications.  Ozone  is  an  allotropic 
form  of  oxygen. 


QUESTIONS  AND  EXERCISES. 

1.  What  percentages  of  sulphur  and  oxygen  are  there  in  sul- 
phur dioxide  (S02),  and  in  sulphur  trioxide  (S03)  ? 

2.  What  are  the  molecular  weights  of  the  following  :  Chlorine, 
carbon  monoxide  (CO),  ethylene  C2H4),  nitric  acid  (HN03),  and 
cane  sugar  (C12H2.2011)? 

3.  What  weight  of  mercuric  oxide  is  required  to  give  20  litres 
of   oxygen   measured  at  0°  C.   and  760   millimetres   pressure  ? 
Equation  :  HgO  =  2Hg  +  02. 


54  QUESTIONS    AND    EXERCISES. 

4.  Calculate   the   weights  of    1   litre  of   the   following   gases 
measured  at  0°  C.   and  760  millimetres  of  pressure  :  Oxygen, 
nitrogen,    carbon   dioxide    (C02),    ammonia    (NH3),    acetylene 
(C2H2),  and  sulphur  dioxide  (SO.,). 

5.  What  volume  (at  standard  temperature  and  pressure)  will 
50  grams  of  oxygen  occupy  ? 

6.  What  weight  of  carbon  dioxide  will  fill  a  vessel  of  15  litres 
capacity  ? 

7.  Point  out  the  errors  in  the  following  equation  : — 2FeS04  = 
Fe203  +  S03. 

8.  The  products  of  the  combustion  of  a  candle  are  collected, 
and  are  found  to  weigh  more  than  the  weight  of  candle  burned. 
Is  this  contrary  to  the  law  of  conservation  of  matter  ?     Explain. 

9.  In  what  respects  does  a  chemical  compound  of  iron  and 
sulphur  differ  from  a  mixture  of  these  two  substances  ? 

10.  100  grams  of  oxygen  combine  with  393.75  grams  of  cop- 
per.    What  is  the  equivalent  of  copper  ? 

11.  Iron  forms  two  basic  oxides.     The  percentage  composition 
of  one  is, — iron,  77.78%  ;  oxyyen,  22.22%  ;  and  of  the  other, — 
iron,  70%  ;  oxyyen,  30%.    Apply  the  law  of  multiple  proportions. 

12.  A  litre  of  nitrogen  unites  with  3  litres  of  hydrogen.     The 
molecule  of  the  compound  formed  contains  1  atom  of  nitrogen 
and  3  of  hydrogen.     What  is  the  change  of  volume  when  com- 
bination takes  place  ? 

13.  What  is  the  composition  of  the  substances  represented  by 
the   following   formulas  :— NH3,    H3P04,    S02(OH)2,    Fe2Cl0, 
CO(NH2)2,  and  K4Fe(CN)8  ?     How  many  atoms  in  the  mole- 
cules represented  as  follows  :— C2H4(OH)2,  (NHJ2S04,  FeS04.- 
7H20,  andFea(OH)6? 

14.  What  weights  of  the  following  gases  measure,  at  standard 
temperature  and  pressure,  22.33  litres,  viz.  : — carbon  monoxide 
(CO) ;  ethylene  (C2H4)  ;  hydric  sulphide  (H2S)  ;  phosphine  (PH3), 
and  methylamine  (CrI3.NH._,)  ? 


HYDROGEN.  55 


CHAPTER    VI. 


HYDROGEN. 

62.  Hydrogen. — It  was  noticed  by  very  early  inves- 
tigators that  when  metals  dissolve  in  acids,  an  inflammable 
gas  is  given  off.  but  the  properties  of  this  gas  (inflammable 
air)    were    not    fully   investigated    till    the   advance   of 
pneumatic  chemistry  in  the  hands  of  Priestly  and  Caven- 
dish towards  the  end  of  last  century  (1766   to    1781). 
Cavendish  showed   that  when  it  burns  in  air  it  forms 
water ;    hence  its  name,  which   means,   water-generator. 
Hydrogen   is   found    uncombined    only   in    small   quan- 
tities, in  volcanic  gases,  in  the  gas  which  occurs  with 
petroleum,  and  in  meteoric  iron.     In  combination,  how- 
ever, it  occurs  in  vast  quantities,  forming  £th  by  weight 
of  water,  ^th  of  marsh  gas,  and  a  considerable  part  of 
sugar,  wood,  starch,  and  animal  and  vegetable  substances 
generally. 

63.  Preparation  of  Hydrogen. — Hydrogen  can 

be  prepared  : 

1.  By  the  electrolysis  of  water  : 

H2O  =  H2  +  O. 

2.  By  the  action  of   various   acids  on  certain  of  the 
metals,  particularly  sulphuric  acid  (diluted)  on  zinc  : 

Zinc.  Sulphuric  Acid.      Zinc  Sulphate.        Hydrogen. 

Zn      +      H2SO4     =     ZnSO4      +      H2 
65  98  161  2 


56  HYDROGEN. 

Hydrochloric  acid  may  also  be  used : 

Zinc.       Hydrochloric  Acid.     Zinc  Chloride. 

Zn      +      2HC1  ZnCl2      +      H, 

65  73  136  2 

and  iron  may  be  used  as  the  metal : 

Iron.        Hydrochloric  Acid.    Ferrous  Chloride. 

Fe     +      2HC1  FeCl2     +      H2 

56  73  127  2 

(If  the  weights  of  metals  represented  in  these  equations 
be  taken  in  (/rams,  what  volumes  of  hydrogen  will  be  pro- 
duced ?) 

3.  By  decomposing  water  by  metals.  Sodium  and 
potassium  decompose  water  at  ordinary,  and  even  at  very 
low,  temperatures ;  magnesium  begins  to  decompose  it 
only  at  100°C.;  and  iron,  copper,  and  silver,  only  at  a  red 
heat. 

Experiment  28.— Put  some  small  pieces  of  zinc  in  a  flask 
provided  with  a  gas  delivery  tube,  and  pour  over  the  zinc  some 
sulphuric  acid  previously  diluted  with  about  five  times  its 
volume  of  water  and  cooled.  Collect  the  gas  which  comes  off, 
at  first  in  an  inverted  test-tube  filled  with  water,  and  test  it 
from  time  to  time  with  a  match  until  it  no  longer  explodes,  but 
burns  quietly  (Why  does  it  explode  ?)  Then  collect  the  gas  in 
jars  for  further  experiments.  When  the  gas  ceases  to  come 
off  pour  the  liquid  in  the  flask  into  a  porcelain  basin  and  set 
it  to  evaporate.  Crystallised  zinc  sulphate,  or  white  vitriol 
(ZnS04.7H20),  is  obtained.  Repeat  this  experiment  using  iron 
(in  the  form  of  tacks)  instead  of  zinc.  Hydrogen  is  obtained  as 
before,  and  on  evaporating  the  liquid  crystallised  ferrous  sul- 
phate, or  green  vitriol  (FeS04.7H20),  is  left.  (How  do  these 
cases  of  solution  differ  from  that  of  common  salt  in  water  ?) 


HYDROGEN.  57 

On  consulting  the  equations  representing  these  actions, 
it  is  seen  that  one  atom  of  zinc  sets  free  a  molecule  of 
hydrogen.  In  other  words  an  atomic  of  zinc  or  of  iron 
is  equivalent  to  two  atoms  of  hydrogen.  (What  are  the 
combining  weights  of  zinc  and  iron  1) 

Experiment  29. — Dissolve  scraps  of  zinc,  iron,  and  tin,  in 
small  quantities  of  hydrochloric  acid  (diluted  with  four  times  its 
bulk  of  water),  in  test  tubes  ;  and  observe  that  an  inflammable 
gas  bubbles  off.  Evaporate  the  remaining  liquids  and  examine 
the  salts  obtained  (zinc  chloride,  ZnCl2,  ferrous  chloride,  FeCl2, 
and  stannous  chloride,  SnCl2).  (Write  the  equations). 

If  the  same  weights  of  zinc  and  iron  were  used  in  this 
experiment,  as  in  Experiment  28,  the  same  volume  of 
hydrcgen  would  be  formed.  The  volume  of  hydrogen 
evolved  by  dissolving  any  weight  of  a  metal  is  not  influ- 
enced by  the  acid,  if  there  are  no  secondary  actions. 

Experiment  30. — Stick  a  small  piece  of  metallic  sodium  to 
the  end  of  an  iron  wire  and  push  the  sodium  quickly  under  an 
inverted  test-tube  filled  with  water  (a  small  jar  may  be  used). 
The  sodium  rises  to  the  bottom  of  the  tube  disengaging  a  gas 
which  pushes  the  water  down.  Closing  the  tube  with  the  thumb, 
remove  it,  and  touch  its  mouth  with  a  flame.  The  gas  burns  ;  it 
is  hydrogen.  Pour  a  few  drops  of  red  litmus  solution  into  the 
test-tube  and  shake  it  up.  The  litmus  turns  blue,  showing  the 
presence  of  caustic  soda. 

An  atom  of  sodium  displaces  an  atom  of  hydrogen  from 
water  : 

H2O  +  Na  =  NaOH  +  H. 
18         23  40  1 

(Compare  with  zinc  and  iron.      What  is  the  equivalent, 
or  combining  weight,  of  sodium).     The  remaining  atom 


58  HYDROXIDES. 

of  hydrogen  can.  be  driven  out  of  the  caustic  soda  by 
heating  it  with  a  second  equivalent  of  sodium,  when  sodic 
oxide  (Na20)  is  formed.  (Write  the  equation  for  this 
action).  Similar  actions  may  be  brought  about  with 
potassium  (K)  and  water  j  and  it  is  found  that  39.1  grams 
of  potassium  decompose  18  grams  of  water,  forming  2.2^ 
litres  of  hydrogen  (1  gram),  and  56.1  grams  of  caustic 
potash.  (Write  the  equation). 

64.  Hydroxides. — Caustic  soda  and  caustic  potash 
are,  in  composition,  midway  between  water  and  the 
oxides  of  sodium  and  potassium.  They  are  called  hydrox- 
ides. Other  examples  of  hydroxides  are  quick  lime, 
or  calcic  hydroxide  (Ca(OH)2),  magnesic  hydroxide 
(Mg(OH)2),  and  bismuthic  hydroxide  (Bi(OH)3).  (How 
many  atoms  of  bismuth,  oxygen,  and  hydrogen  in  a  mole- 
cule of  the  last  of  these?)  The  two  atoms  OH  constitute 
a  group,  which,  united  with  atoms  of  the  metals,  forms 
hydroxides  of  the  metals.  Such  a  group  of  atoms,  acting 
together  as  a  single  atom,  and  being  present  in  a  series  of 
similar  compounds,  is  called  a  compound  radical  The 
name  given  to  the  compound  radical  OH  is  hydroxyl 
(=  hydrogen  oxygen  radical). 

60.  Properties  of  Hydrogen.  —  (Symbol,    H  ; 

atomic  weight,  1  ;  molecule  contains  2  atoms  (H2) ;  1  litre 
weighs  0.0806  gram.)  Hydrogen  is  an  invisible  gas,  the 
lightest  substance  known.  Air  is  14.43  times  as  heavy 
as  it. 

Experiment  31. — Take  a  jar  of  hydrogen  from  the  pneu- 
matic trough,  holding  it  upside  and  pour  it  carefully  upward 
into  another  jar  filled  with  air.  To  show  that  the  hydrogen  has 


PROPERTIES  OF  HYDROGEN.  69 

risen  into  the  second  jar  apply  a  flame  to  its  rnouth.     (What 
properties  of  hydrogen  does  this  experiment  illustrate  ?) 

Hydrogen  is  buoyant  in  air  just  as  a  cork  is  buoyant 
in  water.  It  was  formerly  used  for  inflating  balloons, 
but  the  cheaper  coal  gas  has  now  largely  taken  its  place. 
If  soap  bubbles  be  blown  with  hydrogen  they  rise 
rapidly,  and  can  be  set  on  fire  while  in  the  air. 

Experiment  32- — Thrust  up  a  burning  pine  splinter,  or 
small  taper,  into  a  jar  of  hydrogen  held  upside  down.  The 
hydrogen  catches  fire  and  burns,  but  puts  out  the  taper. 

Hydrogen  is  not  a  supporter  of  ordinary  combustion, 
but  is  itself  combustible.  Still,  a  jet  of  oxygen  will 
burn  in  an  atmosphere  of  hydrogen,  just  as  a  jet  of  hy- 
drogen will  burn  in  an  atmosphere  of  oxygen.  This 
shows  that  the  terms  combustible  and  supporter  of  com- 
bustion are  only  relative  and  would  need  to  be  reversed 
in  their  application  if  we  happened  to  be  living  in  an 
atmosphere  of  hydrogen.— When  hydrogen  burns  in  air 
it  unites  with  the  oxygen  of  the  air  to  form  water.  It 
is  a  constituent  of  most  kinds  of  fuel,  as  wood,  coal, 
oils,  &c.,  so  that  water  is  a  constant  product  of  fires  and 
lights.  (Account  for  the  gathering  of  water  on  the  bot- 
tom of  a  kettle  of  cold  water  set  on  the  tire.) 

Experiment  33. — Fill  an  inverted  test-tube  one-third  with 
oxygen  (by  pouring  it  upwards  in  the  pneumatic  trough)  from  a 
jar  of  the  gas,  and  the  remaining  two-thirds  with  hydrogen. 
Close  the  tube  with  the  thumb,  remove  it  from  the  trough,  turn 
it  several  times  to  mix  the  gases,  and  apply  a  flame  to  its 
mouth.  The  mixed  gases  explode  with  violence.  Repeat  the 
experiment  using  air  and  hydrogen  in  the  proportions  of  5  to  2 
by  volume.  The  explosion  is  less  violent. 


60  PROPERTIES    OF    HYDROGEN. 

The  mixture  of  2  volumes  of  hydrogen  and  1  of 
oxygen  is  known  as  knall  gas.  The  gases  may  remain 
mixed  for  any  length  of  time  at  the  ordinary  tempera- 
ture without  combining,  but  as  soon  as  the  temperature 
is  made  high  enough  at  any  point,  combination  begins 
at  that  point,  water  is  formed,  and  enough  heat  is  set 
free  to  raise  the  temperature  of  the  surrounding  uncom- 
bined  gases  to  the  point  at  which  union  takes  place,  so 
that  the  combination  goes  on  spreading  throughout  the 
whole  mass.  This  takes  place  with  great  rapidity,  and 
an  explosion  results.  (Why  mix  the  gases  in  the  pro- 
portions given,  rather  than  in  any  other  f)  One  gram 
of  hydrogen  in  combining  with  oxygen  liberates  34,462 
heat  units.  (What  is  a  heat  unit  1)  Thus,  in  Experi- 
ment 33,  the  water  formed  by  the  combination  of  the 
gases  is  heated  to  a  very  high  temperature ;  it  expands 
suddenly,  and,  losing  heat,  contracts  suddenly,  making  a 
disturbance  in  the  air  which  reaches  our  ears  as  the 
sound  of  the  explosion. — The  flame  of  hydrogen  is  a  very 
hot  one.  In  the  oxyhydrogen  blowpipe  a  mixture  of 
oxygen  and  hydrogen  is  used  to  give  a  flame  of  intense 
heat.  When  directed  upon  a  piece  of  lime,  this  flame, 
although  giving  very  little  light  itself,  heats  the  lime  to  a 
bright  white  heat,  producing  the  lime-light.  Coal  gas  is 
often  used  instead  of  hydrogen. — Hydrogen  can  be  con- 
densed to  a  liquid  by  pressure  at  a  very  low  tempera- 
ture.— Although  hydrogen  and  oxygen  do  not  usually 
combine  at  the  ordinary  temperature  of  the  air,  yet  in 
contact  with  certain  substances  which  have  the  power  of 
condensing  gases  on  their  surface,  the  two  gases  unite  at 
the  ordinary  temperature.  Platinum  has  this  power 
very  strongly,  especially  when  in  the  finely  divided  con- 


VALENCE.  61 

ditions  of  spongy  platinum  and  platinum  black.  If  a 
coil  of  platinum  wire  is  held  in  a  flame  of  hydrogen 
until  it  is  red  hot  so  as  to  drive  the  condensed  gases  off 
its  surface,  it  acquires  the  power  of  setting  fire  to  a  mix- 
ture of  hydrogen  and  air,  even  after  it  has  cooled.  This 
property  of  platinum  is  utilised  in  heating  a  kind  of 
surgical  knife,  which  burns  instead  of  cutting. — Hydro- 
gen is  very  slightly  soluble  in  water.  It  can  be 
breathed,  and  then  gives  a  high  squeaking  tone  to  the 
voice.  It  is  not  poisonous,  but  will  not  support  animal 
respiration. — Hydrogen  is  evolved  during  the  growth  of 
fungi,  and  as  it  is  being  set  free  it  has  the  power  of  con- 
verting arsenic,  antimony,  and  sulphur  into  poisonous 
gaseous  compounds.  This  accounts  to  some  extent  for 
the  poisonous  effects  of  arsenical  dyes  in  wall  papers  and 
carpets. 

66.  Valence. — Two  volumes  of  hydrogen  unite  with 
one  of  oxygen.  This  indicates  a  difference  between 
hydrogen  and  oxygen  atoms.  We  shall  hereafter  study 
cases  in  which  gases  unite  with  hydrogen  in  equal 
volumes,  or,  as  is  deducible,  one  atom  with  one  atom. 
One  atom  of  nitrogen  is  united  with  three  of  hydrogen  to 
form  ammonia ;  and  one  atom  of  carbon  is  united  with 
four  of  hydrogen  in  the  molecule  of  marsh  gas.  On  refer- 
ring to  Art.  63  we  find  that  while  an  atom  of  sodium  dis- 
places an  atom  of  hydrogen,  an  atom  of  zinc  displaces  two. 

The  following  formulas  express  some  of  these  facts  : — 

Hydrochloric  acid.    Water.    Ammonia.    Marsh  gas. 

HC1,          H2O,      H3N,      H4C. 

The  atoms  of  chlorine,  oxygen,  nitrogen,  and  carbon 
differ  in  the  number  of  atoms  of  hydrogen  which  they 


62  DIFFUSION. 

can  take  to  form  a  molecule.  The  power,  varying  in  this 
way,  is  called  atomicity,  valence,  or  quantivalence.  It  is 
measured  for  any  element  by  the  number  of  atoms  of 
hydrogen  with  which  an  atom  of  the  element  will  combine  j 
or  by  the  n  -mber  of  atoms  of  hydrogen  which  an  atom  of 
the  element  will  replace  in  a  compound  (Art.  63).  The 
woids  bivalent,  trivalent,  quadrivalent,  quinquevalent,  &c., 
are  used  to  denote  the  valence  of  atoms  which  combine 
with  or  replace  one,  two,  three,  &c.,  atoms  of  hydrogen. 
The  words  monad,  dyad,  triad,  tetrad,  Jtexad,  &c.,  are  also 
used.  It  will  be  seen  later  that  the  valence  of  an  element 
varies  in  different  compounds,  but  it  nearly  always  in- 
creases by  twos.  Thus,  the  valence  of  phosphorus  is 
sometimes  3  and  sometimes  5,  e.  g.,  PC13  and  PC15. 
(Form  a  table  of  the  elements  mentioned  in  this  article, 
writing  the  valence  of  each  opposite  its  symbol).  The 
valence  of  atoms  is  often  indicated  by  Roman  numerals 
attached  to  the  symbols,  as  H1,  Na1,  Cl1,  O",  Ca11,  Zn11  • 
N1",  Bim,  Pm;  CIV,  SiIV;  Nv,  Pv,  Sbv,  <fcc.  In  writing 
the  formulas  of  compounds  the  valence  of  the  atoms  is 
often  shown  by  lines  joining  the  symbols.  Formulas 
thus  written  are  used  to  picture  the  knowledge  we  have 
of  the  way  in  which  the  atoms  are  grouped  in  the  molecule, 
and  are  called  Graphic,  or  Constitutional  Formulas.  For 

example,  H — O — H,  water ;  Zn~O,  zinc  oxide  ;  N-^HJ 
ammonia  ;  Ca~oZn.  calcic  hydroxide. 

67.  Diffusion. — If  a  lump  of  sugar  be  placed  at  the 
bottom  of  a  glass  of  water,  it  dissolves  ;  but  the  water  is 
not  at  once  sweetened,  if  it  is  kept  still.  After  some 
time  the  taste  of  the  sugar  can  be  detected  at  the  top, 
and  the  sugar  will  at  length  become  evenly  mixed  with 


DIFFUSION.  63 

the  whole  quantity  of  water,  without  the  slightest  ap- 
parent motion  of  the  liquid.  There  must  be  some  motion 
of  the  particles  of  sugar;  in  fact,  the  motion  of  mole- 
cules mentioned  in  Art.  56.  It  is  supposed  that  the 
molecules  of  sugar,  after  dissolving,  move  about  continu- 
ally until  they  are  evenly  diffused.  This  change  of 
place  of  substances  by  molecular  movements  is  called 
diffusion.  In  one  form  or  another  it  goes  on  in  the 
most  important  processes  of  animal  and  vegetable  life ; 
e.g.,  digestion,  assimilation,  respiration,  secretion,  and 
excretion.  Various  words,  osmose,  endosmose,  dialysis, 
&c.,  are  used  to  describe  diffusion  iinder  certain  condi- 
tions, but  the  process  is  essentially  the  same  in  each  case. 

Experiment  34. —  Fill  a  tall  glass  cylinder  nearly  full  of 
water,  and  pour  through  a  funnel  with  long  capillary  tube  lead- 
ing to  the  bottom  a  layer  of  saturated  solution  of  potassic  bi- 
chromate.  If  the  cylinder  be  kept  undisturbed,  the  red  liquid 
will  not  be  evenly  diffused  until  several  months  have  elapsed. 

Diffusion  of  liquids  is  very  slow.   (Will  solids  diffuse?) 

Experiment  35. — Fit  with  a  bored  cork  a  clean,  dry,  porous 
earthenware  vessel  (inner  cell  of  a  galvanic  battery);  push  a  long 
glass  tube  through  the  cork  (it  must  fit  tightly);  and  fix  this 
apparatus  so  that  the  free  end  of  the  tube  may  stand  in  a  vessel 
of  some  coloured  liquid.  Place  over  the  cell  a  glass  beaker  or 
other  convenient  vessel  bottom  up,  and  fill  the  space  between 
with  hydrogen.  The  coloured  liquid  is  at  once  forced  down  the 
tube,  and  bubbles  of  gas  are  driven  out.  Remove  the  outer  ves- 
sel, and  the  liquid  rises  in  the  tube .  (Coal  gas  answers  in  place 
of  hydrogen. ) 

In  the  first  part  of  the  process,  there  is  air  inside  and 
hydrogen  outside  the  porous  cell.  Both  find  their  way 
through  the  pores,  but  hydrogen  diffuses  faster  than  air, 


64  HYDROGEN    DIOXIDE. 

the  volume  of  gases  inside  the  tube  is  increased,  and  the 
liquid  is  forced  down  to  make  room.  (Explain  the 
second  part  of  the  experiment.)  Diffusion  of  gases  is  a 
much  more  rapid  process  than  that  of  liquids.  (Why?) 
Law  of  Diffusion  of  Gases  (Graham):  —  Gases  diffuse  at 
rates  (measured  by  volume]  inversely  proportional  to  the 
sqiiare  roots  of  their  relative  weights.  For  example,  the 
relative  weights  of  hydrogen  and  oxygen  are  1  and  16. 
Their  rates  of  diffusion  are  as  I/IG  to  -j/I,  or  as  4  to  1, 
i.e.,  hydrogen  diffuses  four  times  as  fast  as  oxygen. 

68.  Hydrogen  Dioxide.  —  (H2O2).    This  substance 

is  also  called  peroxide  of  hydrogen,  and  oxygenated  water. 
It  is  present  in  the  air,  and  comes  down  with  rain  and 
snow,  but  only  in  minute  quantities.  It  is  a  compound 
of  hydrogen  and  oxygen  in  the  proportion  of  1  to  16  by 
weight.  The  simplest  formula  representing  this  compo- 
sition is  HO,  but  if  we  try  to  write  this  graphically, 
showing  the  atomicities,  thus  H  —  O,  an  anomaly  ap- 
pears ;  the  oxygen  atom  must  be  represented  as  monad, 
and  there  is  every  reason  to  believe  that  it  is  dyad. 
From  this  and  other  considerations  the  formula  is 
doubled,—  H2O2  ,  or  H-O—  O-H. 

Preparation.  —  Mix  barium  dioxide  with  dilute  sulphuric 
acid,  allow  to  settle,  and  pour  off  the  clear  liquid.  The  sub- 
stances formed  in  this  reaction  are  hydrogen  dioxide  and  baric 
sulphate  : 

H2S04  +  BaO,  =-  H2O2  -f 


(Put  this  equation  into  words,  giving  weights  and  names.)  —  In 
this  action  the  hydrogen  dioxide  is  obtained  diluted  with  water, 
while  the  heavy,  white,  insoluble  baric  sulphate  falls  to  the 


HYDROGEN    PEROXIDE.  65 

bottom.  As  the  dioxide  is  decomposed  by  heat  it  cannot  be 
separated  from  water  by  boiling  away  the  latter  ;  but  the  water 
can  be  removed  by  placing  the  solution  under  the  receiver  of  an 
air  pump  along  with  a  vessel  of  oil  of  vitriol,  which  is  very 
hygroscopic  (attracts  moisture  strongly).  When  the  air  is  pumped 
out  the  water  evaporates  rapidly,  and  is  absorbed  by  the  oil  of 
vitriol. 

Properties. — A  colourless,  oily  liquid,  heavier  than  water 
(specific  weight  1.452),  soluble  in  water ;  very  unstable,  decom- 
posing slowly  at  low  temperatures,  rapidly  when  heated,  into 
water  and  oxygen.  (Write  the  equation.)  It  easily  parts  with 
half  its  oxygen  to  oxidisable  substances ;  i.e.,  it  is  an  oxidising 
agent.  For  this  reason  it  bleaches  many  colouring  matters  ; 
e.  g. ,  that  of  the  hair ;  and  is  commonly  used  for  bleaching 
dark  hair,  old  paintings,  &c.  It  forms  a  froth  when  taken 
into  the  mouth,  and  excites  the  flow  of  saliva.  It  blisters  when 
undiluted.  Its  principal  use  in  medicine  depends  on  its  oxidis- 
ing power. 

Test. — Add  a  few  drops  of  sulphuric  acid  to  a  solution  of 
hydrogen  peroxide,  then  a  little  ether,  and  a  few  drops  of  pot- 
assic  bichromate  solution.  After  being  shaken  the  mixture  turns 
deep  blue.  The  ether  takes  up  the  colouring  matter,  and  forms 
a  layer  at  the  top. 


QUESTIONS  AND  EXERCISES. 

1.  Would  you  expect  to  find  hydrogen  uncombined  in  the 
air? 

2.  Calculate  the  percentage  composition  of  water,  i.  e.,  the 
quantities  of  hydrogen  and  oxygen  in  100  parts. 

3.  What  volume  of  mixed  gases,  measured  at  1000  mm.  pres- 
sure and  10°  C.,  is  produced  by  the  electrolysis  of  100  grams  of 
water  ? 

6 


66  QUESTIONS    AND    EXERCISES. 

4.  50  grams  of  zinc  are  put  into  a  closed  vessel  of  10  litres 
capacity  along  with  1  litre  of  dilute  sulphuric  acid.     What  is 
the  pressure  of  gas  when  the  action  is  completed  ? 

5.  What  weight  of  zinc  must  be  dissolved  in  order  to  give  50 
litres  of   hydrogen  measured  at   3    atmospheres    pressure    and 
16°  0.? 

6.  20  grama  of  iron  are  dissolved  in  hydrochloric  acid,  and  the 
evolved  gas  is  measured  at  750  mm.  pressure  and  20°  C.     What 
is  its  volume  ? 

7.  "One   atom   of   zinc   sets   free  a  molecule  of  hydrogen." 
What  evidence  can  you  bring  in  support  of  this  statement  ? 

8.  What  is  a  hydroxide  ?    Are  there  hydroxides  of  the  non- 
metals  ? 

9.  Define   compound  radical.     What   is   the  meaning  of   the 
word  radical  as  used  here  ?     Should  it  not  be  spelled  radicle  t 

10.  Calculate  the  weight  of  20  litres  of  hydrogen  measured  at 
25°  C.  and  1000  mm.  pressure.     What  volume  of  oxygen  has  the 
same  weight  under  the  same  conditions  ? 

11.  How  does  it  happen  that  the  atomic  weight  of  hydrogen 
is  1? 

12.  Why  does  hydrogen  explode  less  violently  with  air  than 
with  oxygen  ? 

13.  From  the  following  formulas  deduce  the  atomicity  of  bis- 
muth (Bi),   phosphorus   (P),    tin  (Sn),   and  Sulphur  (S) :  HC1, 
Bi013,  PC16,  H20,  Su02,  S0a. 

14.  The  specific  weights  of  water  vapour  and  oxygen  are  9 
and   16  (Hydrogen  =  1).     If  100  c.c.   of  water  vapour  diffuse 
into  a  vacuum  in  a  certain  time,  what  volume  of  oxygen  will 
diffuse  under  the  same  conditions  in  the  same  time  ? 

15.  Hydrogen  dioxide  is  sometimes  called  hydroxyl,     J^xplain. 


THE    ATMOSPHERE.  67 


CHAPTER     VII. 


AIR. 

69.  The  Atmosphere. — This  is  the  gaseous  ocean, 
supposed  to  be  from  50  to  100  miles  deep,  at  the  bottom 
of  which  we  are  living.  Spaces  said  to  be  empty  are 
generally  full  of  air. 

Experiment  36. — Pass  the  tube  of  a  small  glass  funnel 
through  a  bored  cork  having  a  short  glass  tube  passing  through  a 
second  hole  (as  in  a  wash-bottle).  Fit  the  cork  into  a  flask,  close 
the  glass  tube  with  a  finger  and  fill  the  funnel  with  water.  A 
little  runs  into  the  flask  ;  remove  the  finger  from  the  tube,  and 
the  whole  of  the  water  runs  in  freely.  If  the  tube  have  only  a 
small  opening,  the  water  runs  in  slowly,  and  a  flame  held  near 
the  opening  will  be  blown  by  a  stream  of  air  coming  out. 

(What  conclusions  do  you  dra  w  from  this  experiment  ? 
Why  is  it  necessary  in  filling  a  cask  with  water  through 
a  funnel,  to  bore  a  second  hole  in  the  cask  1) 

Air  has  weight,  as  can  be  shown  by  weighing  a  glass 
globe  emptied  by  the  air  pump,  then  allowing  the  air  to 
flow  in,  and  weighing  it  again ;  a  delicate  balance  shows 
an  increase  in  weight.  One  cubic  foot  of  dry  air  at 
standard  temperature  and  pressure  weighs  565^  grains 
(calculate  the  weight,  in  grams,  of  1  litre). 

Since  air  has  weight,  the  atmosphere  presses  upon  the 
surface  of  the  earth,  and  upon  every  object  on  that  sur- 
face. This  pressure  can  be  shown  by  the  barometer.  Fill 
a  glass  tube,  closed  at  one  end,  and  about  35  inches  long, 


68  ATMOSPHERIC    PRESSURE. 

with  mercury ;  close  it  with  the  thumb,  invert  it,  and 
place  the  end  beneath  the  surface  in  a  dish  of  mercury. 
Remove  the  thumb,  and  the  mercury  after  a  few  oscilla- 
tions comes  to  rest  at  about  30  inches  above  the  level  of 
the  surface  in  the  dish  (Torricelli,  1643).  Some  pressure 
on  this  surface  balances  the  weight  of  the  mercury  in  the 
tube,  otherwise  the  mercury  would  fall.  This  pressure  is 
the  weight  of  the  atmosphere.  It  is  exerted  in  every 
direction  on  the  surface  of  any  body  placed  in  it.  (What 
is  there  inside  the  tube  above  the  mercury  ?) 

Experiment  37. — Make  an  experiment  such  as  that  described 
above,  using  water  instead  of  mercury.  The  water  does  not  fall 
at  all  when  the  inverted  tube  is  opened  under  water.  But  if  the 
tube  were,  say,  40  feet  long  the  water  would  fall  and  remain  at 
the  height  of  about  34  feet. 

An  ocean  of  water  34  feet  deep,  or  one  of  mercury  30 
inches  deep,  would  exert  the  same  pressure  upon  the 
earth  as  the  atmosphere  does.  (Calculate  the  specific 
weight  of  mercury).  A  column  of  mercury  30  inches 
high  and  1  inch  square  in  section  weighs  14.7  Ibs. ;  and 
the  pressure  of  the  atmosphere  is  generally  taken  as  15  Ibs. 
on  the  square  inch.  When  a  barometer  is  carried  up  a 
tower  (Blaise  Pascal,  1648),  or  mountain,  the  mercury 
falls  ;  if  it  is  carried  down  a  deep  mine  the  mercury  rises. 
(Explain).  At  the  top  of  Mont  Blanc  it  stands  at  1  6 
inches  only.  When  any  great  portion  of  the  atmosphere 
is  heated,  it  expands  and  part  of  the  air  flows  away  over 
to  colder  regions.  This  diminishes  the  weight  of  the  air 
in  the  heated  region,  and  the  barometer  falls,  until  air 
rushes  in  from  surrounding  regions  to  equalise  the  pres- 
sure. (What  natural  phenomena  does  this  explain  ?) 
Similarly,  cooling  causes  rise  of  the  barometer.  These 


BOYLE'S  LAW.  69 

and  other  causes  produce  variations  in  the  height  of  the 
mercury  in  the  barometer,  so  that,  in  studying  weather, 
careful  attention  must  be  given  to  the  readings  of  the 
barometer.  The  average  height  at  the  sea-level  is  30 
inches,  or  760  millimetres.  This  is  taken  as  the  standard 
pressure  in  calculating  volumes  of  gases. 

70.  Boyle's  Law  (1662  >.— If  a  quantity  of  air  be 
confined  and  the  pressure  upon  it  increased,  its  volume 
is  lessened  ;  if  the  pressure  be  decreased  the  volume  be- 
comes greater.  This  can  be  shown  by  a  cylinder  closed 
at  one  end  and  provided  with  an  air-tight  piston.  If 
one  presses  the  piston  down,  one  feels  a  resistance,  but 
the  volume  is  diminished  ;  on  the  other  hand,  if  the 
piston  is  raised  above  a  certain  distance,  one  feels  that  a 
weight  is  being  lifted  (What  is  the  weight?)  as  the 
volume  increases.  The  law  of  change  of  pressure  with 
change  of  volume  is  as  follows  : — The  volume  of  any 
portion  of  air  is  inversely  proportional  to  the  pressure, 
the  temperature  being  constant.  (Consult  a  work  on 
Physics  for  the  experimental  demonstration.)  When 
any  portion  of  air  is  closed  off  from  the  rest  of  the  at- 
mosphere, the  pressure  upon  it  is  that  shown  by  the 
barometer  at  that  moment,  say  30  inches  of  mercury. 
If  the  pressure  be  doubled,  the  volume  is  induced  to  \ 
the  original ;  if  the  pressure  be  tripled,  to  J,  &c.  On  the 
other  hand,  if  the  pressure  is  reduced  to  £,  the  volume 
is  increased  to  twice  the  original  and  so  on.  Stated  in 
general  terms,  V  :  V1  :  :  P1  :  P,  where  Y  and  V1  repre- 
sent volumes  at  pressures  P  and  P1.  This  law  applies  to 
all  gases  far  removed  from  their  point  of  condensation* 
It  does  not,  for  example,  apply  to  water  at  105°  C. 

*  At  high  pressures  Boyle's  Law  is  only  approximately  true. 


70 


CHARLES     LAW. 


71.  Expansion    of  Air    by  Heat— Charles' 
Law. 

Experiment  38. — Place  a  narrow-necked  flask  with  its  neck 
downwards  and  dipping  under  water.  Pour  hot  water  on  the 
bottom  of  the  flask  ;  the  air  inside  expands  and  part  of  it  is 
driven  out  of  the  flask.  Allow  the  flask  to  cool ;  the  air  con- 
tracts, and  the  water  rises  in  the  neck  of  the  flask. 

The  amount  of  this  expansion  and  contraction  has  been 
accurately  measured  for  air  and  other  gases.  It  is  almost 
exactly  the  same  for  all  gases.  If  we  begin  with  a  portion 
of  gas  at  0°  C  and  heat  it  to  1°,  its  volume  increases  by 
irl^rd  ;  at  2°  its  volume  is  greater  than  at  0°,  by  2f  s^ds, 
and  so  on.  If  the  gas  be  cooled  to  —  1°  its  volume  is 
less  by  ^rd,  at  —  10°  less  by  ^rds,  and  at  —  273° 
(if  we  could  reach  such  a  temperature)  the  volume  would 
be  0.  Hence  —  273°  C.  is  called  the  absolute  zero  of 
temperature,  and  temperatures  reckoned  from  this  point 
are  called  absolute  temperatures.  Evidently  the  absolute 
temperature  corresponding  to  0°  C.  is  273°,  and  gener- 
ally any  temperature  t  of  the  centigrade  scale  is  273  +  t 
of  the  absolute  scale.  The  Law  of  Expansion  of  Gases 
is  as  follows  : — The  volume  of  any  portion  q/  gas  is  pro- 
portional to  the  absolute  temperature,  the  pressure  being 
constant.* 

Example. — A  quantity  of  air  measuring  10  litres  at  20°  C. 
is  heated  to  50°  C.  What  is  its  volume  ? 

273  -f  20  :  273  +  50  :  :  10  :  x 
x=  8aVs°  =  11. 02  litres. 

72.  Measurement   of  Volumes  of  Gases.— 

In  measuring  the  volume  of  any  portion  of  gas  at  dif- 

*  At  high  pressures,  and  near  the  points  of  condensation,  Charles'  Law  is 
only  approximately  true. 


VOLUMES    OF    GASES.  71 

ferent  times,  it  nearly  always  happens  that  both  pressure 
and  temperature  vary,  so  that  it  becomes  necessary  in 
comparing  volumes  to  apply  both  Boyle's  and  Charles' 
laws. 

Example.  —20  litres  of  hydrogen  at  pressure  755  mm.  and 
temperature  15°  C.  occupy  what  volume  at  740  mm.  pressure 
and  20°  C.? 

740  :  755  » 

273  +  15  :  273  +  20   1    : 
x  =  20  x  J$$  x  £|f  =  20.7  litres. 

A  good  check  on  the  correctness  of  the  statement  in 
such  calculations  is  to  remember  the  general  effect  of 
change  of  pi'essure  and  temperature  on  volume,  and  see 
if  the  fractional  factors  in  the  last  equation  increase  or 
decrease  the  number  of  litres.  Thus,  in  the  above  ex- 
ample, the  pressure  is  lessened,  and  therefore  the  volume 
increased.  I  see  that  the  factor  ?£0  is  greater  than  1, 
and  its  effect  is  to  increase  20.  Similarity  with  tem- 
perature. 

Volumes  of  gases  are  best  compared  at  the  standard 
temperature  and  pressure  (0°  C.  and  760  mm.  of  mer- 
cury) ;  and  calculations  of  volumes  of  gases  produced 
in  chemical  actions  are  always  based  on  the  experimental 
facts  that  the  gram-molecule  of  hydrogen  (i.  e.,  2  grams  of 
hydrogen)  measures,  at  0°  C.  and  7(50  mm.  of  mercury 
pressure,  22.33  litres,  and  that  the  gram-molecules  of  all 
other  gases  occupy  the  same  space,  under  the  same  con- 
ditions. (Art.  60.) 

Example  1. — What  volume  of  oxygen  measured  at  15°  C. 
and  740  mm.  is  formed  by  heating  100  grams  manganese  dioxide  ? 

Equation  :  3Mn02  =  Mn304  +  02 

3  x  86.8  22.33 

grams.  litres. 


72  COMPOSITION    OF    AIR. 

260.4  grams  of  the  oxide  give  22.33  litres  oxygen  measured  at 

100 
0°  0.  and  760  mm.     Thus,  100  grains  give  -^~4  x  22.33  litres  at 

0°  and  760;  and  ^r^  x  22.33  x  273  +  Q  x  ^j  =  the  volume  at 
15°  C.  and  740  mm. 

Example  2. — What  weight  of  potassic  chlorate,  when  de- 
composed, will  give  10  litres  of  oxygen  measured  at  20°  C.  and 
1000  mm.  pressure  ? 

Equation  :  2  KC1O3  =  2  KC1  +  3  O2 

2  x   122.6  3  x  22.33 

grams.  litres. 

i.  e.,  245.2  grams  of  the  salt  give  66.99  litres  of  oxygeu  measured 
at  0°  C.  and  760  mm.  In  order  to  compare  this  volume  with  the 
volume  required,  they  must  be  reckoned  at  the  same  temperature 
and  pressure.  It  is  well,  to  avoid  confusion,  to  compare 
volumes  of  gases  always  at  0°  and  760.  In  this  case,  then,  we 
first  reduce  the  10  litres  to  the  volume  at  0°  and  760  : 

273  +  0  1000 

10  x  293TTo  x  Teo  =  vo1-  at  °°  and  760. 

Then,  66.99  :   10  x  ^  x   j^  :  :  245.2  :  x 

10  X  273  X  100  X  245. 2 
x  =  SOS  X  ft    X  66.00 

73.  Composition  of  Air. 

Experiment  39- — Put  a  bit  of  phosphorus,  dried  on  filter 
paper,  into  a  small  porcelain  cup  floated  on  a  flat  cork  in  a  basin 
of  water.  Touch  the  phosphorus  with  a  hot  wire  and  immediately 
cover  it  with  a  beaker.  It  burns  for  some  time,  but  at  length 
goes  out.  As  the  enclosed  gas  cools,  water  rises  in  the  beaker, 
showing  a  lessening  of  volume.  Decant  some  of  the  gas  into  a 
test-tube  (Art.  41),  close  the  test-tube  with  the  thumb,  remove 
it,  and  thrust  a  burning  match  or  splinter  into  the  gas.  The 
flame  is  put  out. 

Further  examination  of  this  gas  has  shown  it  to  be  an 
element,  nitrogen.  The  combustion  of  phosphorus  is  due 


VOLUMETRIC    ANALYSIS    OF    AIR.  73 

to  the  oxyjeti  of  the  air.  (What  has  become  of  these  in 
the  experiment]) — -Air  is  a  mixture  of  the  two  elementary 
gnues,  nitrogen  and  oxygen.  That  it  is  a  mixture,  and 
not  a  chemical  compound,  is  seen  from  the  following 
facts  :  — 

1.  If  air  be  shaken  up  with  water,  oxygen  dissolves 
in  greater  relative  quantity  than  does  nitrogen. 

2.  The  two  gases  can  be  partially  separated  by  dif- 
fusion.    (Which  gas  diffuses  faster  ]) 

3.  If  nitrogen  and  oxygen  be  mixed  in  the  ratio  in 
which  they  are  present  in  air,  no  heat  is  given  out,  and 
there  is  no  other  evidence  of  chemical  action ;  but  the 
mixture  has  the  properties  of  pure  dry  air. 

4.  The   volumes  of  the  gases  are  not  in  any  simple 
ratio.      (What  law  is  here  referred  to?) 

VOLUMETRIC  ANALYSIS  OF  AIR. — This  analysis  is  made 
in  a  graduated  glass  tube  called  a  eudiometer.  The 
tube  is  closed  at  one  end,  at  which  two  platinum  wires 
are  melted  through  the  glass  so  as  nearly  to  meet  within 
the  tube.  The  eudiometer  is  filled  partly  with  mercury, 
inverted  in  a  trough  of  the  same  liquid,  and  the  volume 
of  air  thus  closed  off  is  noted,  as  well  as  the  temperature 
and  pressure.  To  this  air  is  added  about  half  its  volume 
of  pure  hydrogen,  and  volume,  temperature,  and  pressure 
are  once  more  noted.  The  eudiometer  is  now  tightly 
pressed  down  on  a  sheet  of  India-rubber,  and  an  electric 
spark  is  passed  between  the  platinum  wires.  An  ex- 
plosion takes  place  within  the  tube,  and  the  volume  of 
gas  is  seen  to  be  much  reduced.  Volume,  temperature, 


74  ANALYSIS    OF    Alll    BY    WEIGHT. 

and  pressure  are  once  more  read  off.  The  remaining  gas 
is  a  mixture  of  nitrogen  and  hydrogen.  One-third  of  the 
loss  of  volume  after  the  explosion  is  due  to  the  disap- 
pearance of  oxygen  (What  has  become  of  it1?) ;  and  the 
volume  of  oxygen  being  known,  the  remainder  of  the 
original  volume  of  air  is  reckoned  as  nitrogen. 

Example. — Volume  of  air  enclosed,  20  c.c. ;  temperature, 
15°  C. ;  pressure,  750  mm. — Volume  after  addition  of  hydrogen, 
32  c.c. ;  t,  16*  C. ;  p,  718.  Volume  after  explosion,  20  c.c.; 
t,  18°  C. ;  p,  710  mm.  Reduce  these  volumes  to  0°  C.  and 
760  mm.,  and  find  the  decrease  of  volume  after  explosion  to  be 
11. 03  c.c.  Calculate  from  this  that  100  c.c.  of  air  consist  of 
20.96  of  oxygen  and  79.04  of  nitrogen. 

COMPOSITION    BY    VOLUME. 

Oxygen 20.96 

Nitrogen 79.04 

100.00 

ANALYSIS  OF  AIR  BY  WEIGHT. — Air  is  dried  and 
purified,  and  then  allowed  to  flow  slowly  through  a 
weighed  tube  containing  red  hot  copper  filings.  It  loses 
oxygen,  cupric  oxide  (CuO)  being  formed  ;  and  the  nitro- 
gen passes  on  into  a  weighed  vacuous  globe  of  glass. 
he  increase  in  weight  of  the  tube  gives  the  weight  of 
oxygen,  and  the  increase  in  the  globe  the  weight  of 
nitrogen.  The  average  of  many  careful  experiments 
gives  as  the  percentage  composition  of  air  by  weight  : 

Oxygen 22.77 

Nitrogen 77.23 

100.00 


CARBON    DIOXIDE,    ETC.  75 

The  composition  of  air  is  almost  constant  in  all  parts 
of  the  world,  and  in  all  situations.  In  towns  and  in 
foggy  weather  the  percentage  of  oxygen,  by  volume,  may 
sink  to  20.8 ;  and  in  crowded  rooms  it  sometimes  sinks 
much  lower. — The  oxygen  of  the  air  is  necessary  to  sup- 
port life,  fires,  decaying  processes,  &c.  (What  is  the 
use  of  nitrogen  1) 

74.  Other  Substances  in  the  Atmosphere. — 

The    atmosphere    contains    several    substances    besides 
oxygen  and  nitrogen,  but  in  small  and  variable  quantities. 

CARBON  DIOXIDE  (CO2). — Constantly  present  in  air, 
from  4  to  6  volumes  in  10,000.  Its  presence  is  of  great 
importance,  as  it  forms  an  essential  constituent  of  the 
food  of  plants.  The  supply  is  kept  up  by  the  processes  of 
combustion,  respiration,  and  the  decay  of  organic  matter. 

WATER  (H2O). — -This  is  present  in  the  air,  in  the  form 
of  vapour,  in  very  variable  quantities,  depending  on  tem- 
perature and  degree  of  saturation.  A  cubic  metre  of  air 
at  25°  C.  contains  when  saturated  22.5  grams  of  water 
vapour ;  at  0°  C.,  only  5.4  grams.  Dew  is  formed  by 
the  condensation  of  aqueous  vapour  from  the  air  by  con- 
tact with  cooled  surfaces.  Leaves  radiate  heat  much 
faster  than  other  objects,  and  therefore  condense  water 
vapour  much  more  rapidly.  (Why  does  not  dew  form 
on  a  cloudy  or  a  windy  night  ])  Aqueous  vapour  in  the 
air  tempers  the  heat  of  the  sun,  the  rays  of  which  would 
be  unbearably  hot,  if  a  large  fraction  of  the  heat  were 
not  stopped  on  its  passage  through  the  atmosphere. 
Water  vapour  has  in  a  very  high  degree  the  power  of 
absorbing  radiant  heat. 


76  COMBUSTION    IN    AIR. 

COMPOUNDS  OF  NITROGEN.  —  Minute  quantities  of 
oxides  of  nitrogen  are  formed  in  the  atmosphere  by  the 
action  of  electricity.  (On  what1?)  Ammonia  (NH3)  is 
present  in  small  proportions,  generally  as  ammonic  nitrite 
or  nitrate.  These  substances  are  brought  down  to  earth 
in  rain  and  snow,  and  serve  for  plant  food. 

DUST,  <fec. — Under  this  head  are  included  solid  par- 
ticles of  all  sorts,  organic  and  inorganic.  Common  salt 
is  always  present  in  the  air.  If  a  little  clean  rain-water 
is  evaporated  on  a  microscope  slide  and  the  residue  ex- 
amined with  the  microscope,  crystals  of  common  salt 
(NaCl)  can  always  be  seen.  Spores,  or  eggs  of  minute 
plants  and  animals,  are  constantly  present.  Many  of 
them  are  the  eggs  of  ferments  ;  others  are  the  cause,  or 
at  least  the  concomitants,  of  diseases.  The  object  of  the 
antiseptic  spray  in  surgery  is  to  kill  such  living  dust ; 
and  it  is  important  to  note  that  disinfectants  and  anti- 
septics which  are  to  purify  air  must  (unless  the  air  is 
drawn  over  them)  be  volatile.  Air  can  be  purified  from 
dust  by  drawing  it  through  cotton-wool,  and  other  filters. 

Ozone  and  hydrogen  dioxide  have  been  already  men- 
tioned. 

75.  Combustion  in  Air. — Substances  which  burn 
in  oxygen  usually  burn  in  air,  but  not  so  rapidly. 
(Why  1)  The  substances  burned  in  lamps,  candles,  and 
fires  are  composed  mostly  of  carbon  and  hydrogen,  which 
unite  with  oxygen  to  form  carbon  dioxide  and  water. 
Thus,  the  oxygen  of  the  air  is  used  up,  and  in  lamps  and 
stoves,  &c.,  provision  must  be  made  for  a  renewal  of  the 
supply  of  air,— in  other  words,  there  must  be  a  draught. 


RESPIRATION.  77 

The  consumption  of  oxygen  must  be  taken  into  account 
in  considering  questions  of  ventilation.  An  ordinary 
lamp  consumes  as  much  oxygen  in  1  hour  as  a  man  does 
in  5.  (How  does  a  lamp  or  gas  flame  differ  from  a  fire 
in  a  stove  or  grate  in  their  effect  on  the  air  ?) 

76.  Respiration. — In  the  air  cells  of  the  lungs  the 
air  is  separated  from  the  blood  by  a  very  thin  membrane 
through  which  diffusion  goes  on  readily  and  rapidly. 
The  blood  comes  into  the  lungs  charged  with  carbon  di- 
oxide, a  waste  product  which  it  has  gathered  fi-om  the  fur- 
naces of  the  body, — the  muscles,  glands,  «fec.  These  need 
a  continual  supply  of  oxygen  to  feed  the  slow  fires  going 
on  in  them.  Carbon  dioxide  passes  outwards  by  diffu- 
sion, and  is  exhaled ;  oxygen  diffuses  inward  and  is  car- 
ried away  to  the  tissues.  Thus,  the  air  which  returns 
out  of  the  lungs  contains  less  oxygen  and  more  carbon 
dioxide  than  when  it  was  inspired.  In  one  hour  an 
adult  man  consumes  about  29  grams  (how  many  litres "?) 
of  oxygen,  and  breathes  out  about  33  grams  carbon  di- 
oxide. Besides  carbon  dioxide,  water  vapour  and  small 
quantities  of  complex  organic  compounds  are  exhaled. 
Both  the  organic  compounds  and  the  carbon  dioxide  are 
injurious  to  health ;  and  therefore  the  necessity  for  re- 
newing the  air  which  is  breathed.  It  is  estimated  that 
an  average  adult  man  requires  1500  cubic  feet  of  fresh 
air  per  hour. 


QUESTIONS  AND  EXERCISES. 

1.  A  flask  filled  with  water  and  closed  with  a  cork  through 
which  passes  a  narrow  tube  open  at  both  ends  is  held  upside 
down.  The  water  does  not  run  out.  Why  ? 


78  QUESTIONS    AND    EXERCISES. 

2.  Air  is  found  to  become  less  and  less  dense  as  we  ascend. 
Account  for  this. 

3.  Why  do  men  need  to  breathe  faster  at  great  elevations  than 
lower  down  ? 

4.  Why  does  blood  burst  through  any  place  where  the  skin  is 
thin,  when  a  very  great  elevation  is  reached,  as  in  balloons  ? 

5.  What  is  the  pressure  of  the  atmosphere  in  grams  on  every 
square  centimetre  ? 

6.  The  specific  weight  of  alcohol  is  0. 784.     What  would  be 
the  average  height  of  an  alcohol  barometer  ? 

7.  A  quantity  of  air  measures  264  c.c.  at  a  pressure  of  700  mm. 
What    will    it    measure    when    the    pressure   is   increased   to 
1000  mm.? 

8  Calculate  the  weight  of  hydrogen  in  a  vessel  of  10  litres 
capacity,  filled  when  the  barometer  reads  756  and  the  ther- 
mometer 18°  C. 

9.  A  quantity  of  air  under  a  pressure  of  32  inches  of  mercury 
undergoes  a  change  of  pressure  and  increases  from  10  cubic  feet 
to  12.4  cubic  feet.     What  is  the  new  pressure?     (Temperature 
constant. ) 

10.  What   pressure  must   be   used   to  compress   18  litres  of 
oxygen  into  a  vessel  of  6  litres  capacity,  the  oxygen  being  origi- 
nally under  a  pressure  of  785  mm.  ? 

11.  150  c.c.  of  air  at  50°  C.  is  cooled  to- 10°  C.     Calculate  the 
volume  (pressure  being  constant). 

12.  A  quantity  of  air  measures  24  litres  at  15°. ;  the  tempera- 
ture is  reduced  to  —  16°  C.     What  is  now  the  volume  ?     (Pres- 
sure constant.) 

13.  What  volume  of  gas  at  200°  C.  will  measure  320  c.c.  at 
0°  C.?    (Pressure  constant. ) 

14.  What  volume  of  oxygen  measured  at  2000  mm.  and  12°  C. 
is  formed  by  the  decomposition  of  500  grams  potassic  chlorate  ? 


NITROGEN.  79 


CHAPTER    VIII. 


NITROGEN  AND  ITS  COMPOUNDS. 

Nitrogen— (N'"  =  14). 

77.  Occurrence.  —  Forms     about     four-fifths     by 
volume  of  the  atmosphere ;  in  combination,  it  forms  part 
of  all  living  matter,  and  is  therefore  an  essential  consti- 
tuent of  plant  and  animal  food.     Other  compounds  oc- 
curring in  nature  are  nitre  (KNO3^,  Chili  nitre  (NaNO3), 
and  ammonia  (NH3). 

78.  Preparation. — Nitrogen  is    most    readily    pre- 
pared from  air  (Exp't  39) ;  see  Art  73,  "  Analysis  of  Air 
by  Weight."     It  can  also  be  prepared  from  nitre  (salt- 
petre). 

Experiment  40- — Heat  10  grams  iron  filings  in  a  hard-glass 
test-tube  with  \  gram  saltpetre.  Collect  the  evolved  gas  in 
ttie  usual  way.  It  is  nitrogen,  the  "nitre-generator." 

79.  Properties. — An  invisible  gas,  without  smell  or 
taste.     It  is   14  times  as  heavy  as  hydrogen  and  a  little 
lighter  than  air.     (Calculate  its  specific  weight,  air  being 
standard.)     It  is  incombustible  and  not  a  supporter  of 
combustion  (Mate  experiments  to  show  this).     It  is  re- 
markable for  chemical  inactivity,  not  entering   readily 
into  combination   with  other  elements.     Under  the  in- 
fluence of  electric  sparks  or  flashes  it  unites  with  oxygen, 
and,  in  very    small   quantities,    with    hydrogen.     It    is 


80  AMMONIA. 

slightly  soluble  in  water,  about  1£  vols.  in  100  of  water. 
It  has  been  liquefied  at  —  146°  C.  by  a  pressure  of  33 
atmospheres. 

Compounds  of  Nitrogen. 

80.  Ammonia.— (NH;!  =17.) 

Experiment  41. — Heat  apiece  of  dried  meat  in  a  glass  tube 
closed  at  one  end.  It  chars,  and  moisture  collects.  Wet  a  small 
strip  of  filter  paper  with  red  litmus  and  bring  it  in  contact  with 
this  moisture.  The  litmus  is  turned  blue.  Same  result,  if  a  bit 
of  coal,  bread,  or  horn  be  heated. 

The  alkaline  reaction  of  the  moisture  is  due  to  the 
presence  of  a  volatile  alkali,  ammonia.  (What  alkalis 
have  already  been  noticed  1)  It  is  always  formed  when 
animal  or  vegetable  bodies  are  destructively  distilled, 
Coal,  wood,  bones,  and  other  animal  and  vegetable  sub- 
stances are  distilled  on  a  large  scale  in  the  manufacture 
of  gas,  charcoal,  bone-black,  &c.;  and  the  watery  liquid 
(a  by-product)  contains  much  ammonia.  From  gas 
liquor  ammonia  is  prepared  by  heating  with  lime  and 
condensing  the  ammonia  in  cooled  water.  Formerly,  it 
was  prepared  by  distilling  scraps  of  the  horns  of  the 
hart ;  hence  the  popular  name,  "  spirits  of  hart's  horn." 
— Ammonia  is  always  one  product  of  the  decay  of  ani- 
mals and  vegetables,  excrement,  guano,  &c. 

PREPARATION. — Ammonia  is  most  conveniently  pre- 
pared from  one  of  its  salts,  sal  ammoniac  (]STH4C1). 

Experiment  42. — Mix  thoroughly  about  2  parts  of  dry 
powdered  sal  ammoniac  with  1  of  powdered  quick-lime.  Note 
the  smell  of  ammonia.  Put  the  mixture  into  a  hard  glass  test- 


AMMONIA.  81 

tube  or  tlask,  ar.d  arrange  a  gas  delivery  tube  so  as  to  collect  the 
evolved  gas  in  inverted  bottles,  to  the  tops  of  which  the  delivery 
tube  must  reach.  ( Why  not  over  water  ?) 

• 
Equation  : 

Ammonic  Calcic 

I.ime.  Chloride.      '     Chloride. 

CaO  +  2NH4C1  =  CaCl2  +   2NH3  +  H,O. 
56  106.8          110.8        44.66 

grams.  grams.  grams.  litres. 

Ammonia  is  generally  used  in  water.  To  prepare  the 
solution  slaked  lime  is  employed  instead  of  quick  lime, 
and  the  gas  is  passed  into  water,  which  must  be  kept  well 
cooled.  (Why  1) 

PROPERTIES. — Ammonia  is  an  invisible  gas,  with  a 
pungent  smell,  and  a  sharp  alkaline  taste.  It  is  lighter 
than  air.  (How  does  Exp't  42  show  this  1)  (Calculate  its 
specific  weight,  air  being  standard  ;  also  when  hydrogen  is 
standard.)  It  is  liquid  at  15.5°  C.  under  a  pressure  of 
7  atmospheres  (What  is  an  atmosphere  of  pressure  1),  and 
at  —  70°  C.  is  solid.  It  is  very  soluble  in  water,  1148 
vols.  dissolving  in  1  of  water  at  0°  C.  Much  heat  is 
given  out  during  the  process.  (Account  for  this.)  When 
a  solution  of  ammonia  is  warmed  the  gas  is  driven  off 
rapidly. 

Experiment  43- — Fill  a  test-tube  with  ammonia  gas  (Experi- 
ment 42),  close  with  the  thumb,  place  the  mouth  under  water, 
and  remove  the  thumb.  The  water  rises  to  the  top,  if  there  is 
no  air  in  the  tube.  Close  again  with  the  thumb,  remove  the 
tube  and  examine  the  water  in  it,  as  to  its  smell,  taste,  feel, 
and  action  on  red  litmus. 

The  solution  of  ammonia  in  water  (liquor  ammoniae) 
has  a  strong  alkaline  reaction  and  probably  contains 
ammonic  hydroxide  (NH4OH).  It  neutralises  acids,  and 

7 


82  LIQUOR    AMMONITE. 

thereby  forms  salts.  The  dry  gas,  also,  combines  with 
acids,  and  it  is  to  be  observed  that  no  water  is  pro- 
duced when  ammonia  unites  with  acids  to  form  salts, 
e.g.,  NH3  +  HC1  =  NH4di.  These  salts  are  called 
ammonium  salts,  since  they  contain  the  compound  radical 
ammonium  (NH4)  acting  as  a  monad  metal.  Ammonia 
is  not  a  supporter  of  combustion,  and  is  incombustible  in 
air ;  but  a  mixture  of  the  gas  and  oxygen  burns  with  a 
pale  bluish  flame,  and  forms  water,  nitrogen,  and  a  little 
nitric  acid. 

Experiment  44. — Try  the  combustibility  of  ammonia  gas  by 
thrusting  a  burning  match  into  a  jar  of  it. 

COMPOSITION  OF  AMMOXIA. — The  gas  is  decomposed 
by  electric  sparks.  Two  volumes  of  ammonia  give  one 
of  nitrogen  and  three  of  hydrogen.  That  is,  two  mole- 
cules of  ammonia  decompose  into  one  molecule  of  nitrogen 
and  three  of  hydrogen.  From  this  it  is  concluded  that 
the  molecule  of  ammonia  contains  an  atom  of  nitrogen 
and  three  of  hydrogen. 

Liquor  ammonite  fortior  is  a  nearly  saturated  solution  of  am- 
monia in  water,  prepared  by  distilling  in  an  iron  retort  3  Ibs. 
ammonic  chloride  (NH4C1),  with  4  Ibs.  slaked  lime  (Ca(OHj,2), 
and  receiving  the  gas  in  water.  It  is  lighter  than  water 
(sp.  wt.  0.892),  and  contains  32.5  per  cent,  of  ammonia.  It 
must  be  kept  well  stoppered,  otherwise  the  ammonia  escapes. 
As  ammonia  is  less  soluble  at  high  than  at  low  temperatures  this 
solution  should  not  be  allowed  to  get  very  warm,  (What 
would  happen  ?) 

Liquor  ammonite  or  aqua  ammonice  is  a  weaker  solution,  made 
by  mixing  1  pint  of  the  stronger  solution  with  2  pints  distilled 
water.  It  contains  10  per  cent,  of  ammonia. 

Ammonia  has  a  strong  inflammatory  action  on  the  respi- 
ratory and  other  mucous  membranes.  When  breathed 


NITKIC   ACID.  83 

diluted  with  air,  it  stimulates ;  but  its  constant  use  may 
at  length  produce  serious  inflammation.  Since  it  unites 
with  acids  to  form  mild  salts,  it  is  an  antidote  to  acid 
gases.  (What  are  antidotes  for  ammonia  ?). 

Hydroxylamine  (NH30)  is  a  substance  nearly  related  to  am- 
monia. Its  molecule  contains  hydroxyl  (OH)  instead  of  one  of 
the  atoms  of  hydrogen.  If  the  molecule  of  ammonia  is  repre- 
sented by  N-^H>  that  of  hydroxylamine  is  N^— H  •  It  is  pre- 

^\  H  ^*  H 

pared  by  the  action  of  nascent  (just  being  set  free)  hydrogen  on 
nitrogen  dioxide  (NO) :— NO  +  3H  =  NH30. 

81.  Nitric  Acid  (HN03  =  63). 

Experiment  45. — Put  a  little  saltpetre  in  a  test-tube,  and 
drop  on  it  from  a  pipette  a  small  quantity  of  oil  of  vitriol.  Heat 
very  gently,  and  observe  the  drops  of  fuming  liquid  which  gather 
on  the  sides  of  the  tube.  Gather  them  on  a  glass  rod  and  try 
their  action  on  red  litmus. 

The  acid  liquid  is  nitric  acid,  or  aquafortis.  Saltpetre 
is  a  salt  of  the  base,  potassic  hydroxide,  and  nitric  acid. 
When  it  is  acted  upon  by  the  stronger  acid,  the  weaker 
is  displaced,  and  a  new  salt,  potassic  sulphate,  is  formed. 

2KN03  +  H2SO4  =  2HN03  +  K2SO4. 

Nitric  acid  occurs  naturally  only  in  small  quantities ; 
but  its  salts  are  found  in  large  quantities.  Chili  salt- 
petre, or  sodic  nitrate  (NaNO3),  is  most  plentiful,  and  is 
now  used  in  preparing  the  acid. 

PREPARATION. — Distil  at  a  gentle  heat,  from  a  glass 
retort,  equal  weights  of  Chili  saltpetre  and  oil  of  vitriol, 
keeping  the  receiver  well  cooled.  Nitric  acid  distils, 
and  a  white  salt  remains  in  the  retort.  This  salt  has 


84  NITRIC    ACID. 

acid  properties.     In  fact  it  is  still  half  acid,  and  is  called 
acid  sulphate  of  sodiiim,  or  sodic  hydric  sulphate  : 

NaNO3  +  H2S04  =  HNO3  +  NaHS04. 

Why  not  use  double  the  quantity  of  sodic  nitrate,  as 
represented  in  the  preceding  equation  ]  Because,  in 
this  case,  the  action  would  go  on  only  at  such  a  high 
temperature  that  the  nitric  acid  would  be  partially 
decomposed. 

PROPERTIES. — Pure  nitric  acid  is  a  colourless  fuming 
liquid  of  sp.  wt  1.52.  It  has  a  strong  attraction  for 
water,  and  it  is  hard  to  prepare  it  free  from  water.  The 
strong  acid  of  commerce  always  contains  about  10%  of 
water,  and  is  generally  reddish  or  yellowish  from  im- 
purities. Red  fuming  nitric  acid  is  strong  nitric  acid 
coloured  by  nitrogen  trioxide  and  tetroxide.  When 
strong  nitric  acid  is  boiled  it  loses  acid  faster  than  it 
does  water,  until  a  liquid  containing  70%  of  acid  distils 
at  121°  C.  unchanged.  If  a  weaker  acid  be  boiled  it 
loses  water  faster  than  acid  until  it  contains  70°/0  of 
acid. 

Experiment  46. — Put  a  few  cubic  centimetres  of  strong  nitric 
acid  in  a  test-tube,  and  add  an  equal  volume  of  water.  Note  the 
rise  of  temperature.  ( Account  for  it. )  Pour  the  contents  of  the 
tube  into  about  one  litre  of  water,  stir  well,  and  taste. 

Nitric  acid  is  a  powerful  oxidising  agent.  (Calculate 
the  percentage  of  oxygen  which  it  contains). 

Experiment  47. — Drop  a  little  strong  nitric  acid  on  some 
powdered  charcoal  heated  to  redness  in  a  deflagrating  spoon. 
N.B. — The  acid  must  contain  very  little  water. 


NITRIC    ACID.  85 

Experiment  48- — Very  carefully  drop  a  small  piece  of  phos- 
phorus into  a  little  nitric  acid  in  a  small  porcelain  dish.  It  dis- 
solves with  the  evolution  of  reddish  fumes.  If  the  acid  is  very 
strong,  the  phosphorus  catches  fire. 

If  the  solution  from  Experiment  48  be  evaporated, 
phosphoric  acid  is  left.  (How  has  this  acid  been  before 
obtained  ?)  Similarly,  sulphur  can  be  oxidised  to  sul- 
phuric acid  by  boiling  with  nitric  acid.  Turpentine  can 
be  set  on  fire  by  a  mixture  of  nitric  and  sulphuric  acids. 

Experiment  49.  — Dilute  a  few  cubic  centimetres  of  nitric  acid 
with  about  4  times  its  volume  of  water,  and  pour  it  over  a  small 
bit  of  lead  in  an  evaporating  dish.  The  lead  begins  to  dissolve, 
and  red,  strongly-smelling  fumes  come  off.  Apply  heat  to  hasten 
the  action,  and  when  the  lead  is  completely  dissolved  evaporate 
the  solution  on  the  water  bath.  Crystals  of  colourless  plumbic 
nitrate  (Pb(N03)a)  are  left.  Repeat  the  experiment,  using  cop- 
per instead  of  lead.  Fine  blue  crystals  of  cupric  nitrate  are 
formed  : 

3Pb  +  8HNO3  =  3Pb(N03)2  +  2NO  -f  4H20. 
3Cu  +  8HN03"  =  3Cu(N03)2  -f-  2NO  -f  4H2O. 
(Write  down  the  names  and  weights.) 

Nitric  acid  dissolves  most  metals,  forming  nitrates  of 
the  metals,  oxides  of  nitrogen,  and  water.  It  dissolves 
silver,  but  not  gold,  and  can  thus  be  used  to  separate 
silver  from  gold.  If  the  strong  acid  be  used  in  dissolv- 
ing metals,  the  gaseous  product  is  nitrogen  trioxide 
(N2O3);  dioxide  (NO),  monoxide  (N2O),  nitrogen,  and 
even  ammonia,  are  obtained  at  successive  stages  of  dilu- 
tion. As  a  rule,  the  more  violent  (rapid)  the  action  of 
the  acid,  the  less  oxygen  does^it  lose.  Thus,  the  higher 
oxides  of  nitrogen  are  formed  when  the  acid  is  strong, 


86  NITRATES. 

the  temperature  high,  or  the  metal  easily  oxidised. — 
Nitric  acid  is  a  strong  corrosive  poison.  It  '  eats  up '  or- 
ganic tissues  when  it  is  strong,  and  partially  destroys  them 
even  when  it  is  weak.  The  antidotes  are  mild  alkaline 
substances,  as  magnesia.  A  drop  of  the  weak  acid  left  on 
the  skin  for  a  few  moments  colors  it  bright  yellow  (Picric 
acid  is  formed). 

Nitro-glycerine  and  gun-cotton  are  chemical  compounds 
of  glycerine  and  cotton  with  nitric  acid,  in  which  the 
oxygen  of  the  acid  is  ready  to  combine  with  the  carbon 
and  hydrogen  of  the  glycerine  and  cotton.  (Explain 
their  explosiveness.)  Dynamite  is  a  commercial  prepara- 
tion of  nitro-glycerine. 

Tests. — 1 .  Heat  with  some  bits  of  copper.  Red  fumes  are 
given  off. 

2.  Colour  light  blue  with  a  drop  of  sulphate  of  iiidigo,  and 
heat.     The  colour  disappears,  because  the  indigo  is  oxidised. 

3.  Mix  in  a  test-tube  about  equal  volumes  of  strong  sulphuric 
acid  and  solution  of  ferrous  sulphate  (FeSO4),  cool,  and,  holding 
the  test-tube  aslant,  carefully  pour  down  the  side  so  as  to  form 
a  layer  on  the  top  a  dilute  solution  of  nitric  acid.     Either  at 
once  or  after  a  few  moments  a  brown  ring  appears  between  the 
two  layers.     This  is  due  to  the  formation  of  a  brown  compound 
of  ferrous  sulphate  and  nitrogen  dioxide  :  3H2S04  -f-  2HNO-, 
-f  10  FeS04  =  3FeV(S04),  +  2(FeSOJ2NO  +  4H20. 

82.  Nitrates. — Basicity. 

Experiment  50- — Put  a  pipetteful  of  caustic  potash  solution 
in  a  porcelain  basin,  colour  with  litmus,  and  add  dilute  nitric 
acid  (1  of  strong  acid  to  4  of  water)  slowly  with  a  pipette,  stir- 
ring constantly,  until  the  blue  litmus  turns  purple.  Taste  the 
solution.  Evaporate  on  the  water  bath.  Long  prismatic  crys- 
tals of  saltpetre  are  obtained. 


BASICITY.  81 

If  this  experiment  be  made  quantitatively  it  will  be 
found  that  to  get  the  neutral  point,  pure  potash  and  pure 
acid  must  be  mixed  in  the  ratio  by  weight  of  56.1  to  63. 
If  any  other  ratio  be  used,  some  of  the  acid  or  of  the  base 
is  left  over : 

KOH  +  HNO3  =  KNO3  +  H2O. 

56.1  63 

Base  and  acid  give  salt  and  water. 

If  a  similar  experiment  be  made  with  caustic  soda, 
sodic  nitrate,  in  cubical  crystals,  is  formed ;  and  it 
is  found  that  in  this  case  also  the  base  and  acid  unite 
in  only  one  proportion,  viz.,  40  to  63. 

NaOH  -f  HNO3  =  NaNO3  -f  H,O. 
40  63 

40  g.  of  caustic  soda  are  equivalent  to  63  g.  nitric  acid 
(and  to  what  weight  of  caustic  potash  1)  Since  nitric 
acid  acts  on  these  two  bases  in  only  one  proportion  for 
each,  and  forms  only  one  salt  for  each,  it  is  called  a 
monobasic  acid.  The  bases  are  monacid  bases. 

Experiment  51. — Put  some  dilute  nitric  acid  in  a  porcelain 
dish,  colour  with  litmus,  and  add  dilute  solution  of  ammonia 
until  the  litmus  just  turns  blue.  Then  evaporate  to  dryness  on 
the  water  bath,  and  preserve  the  crystals  of  ammonic  nitrate  for 
a  later  experiment : — 

NH3  +  HN03  =  NH4NOS. 

Experiment  52. — Warm  a  little  plumbic  oxide  (litharge)  in  a 
porcelain  basin  with  dilute  nitric  acid.  The  oxide  dissolves  to  a 
colourless  solution.  Evaporate  on  the  water-bath.  Colourless 
crystals  of  plumbic  nitrate  (Pb(NOs)2)  are  left : — 

PbO  +  2HN03  =  Pb(N03)2  -f-  HaO. 
Keep  this  salt  for  a  later  experiment. 


88  SALTS. 

The  molecule  of  plumbic  oxide  is  equivalent  to  two 
molecules  of  nitric  acid.  There  is  a  corresponding  base, 
plumbic  hydroxide,  Pb(OH)2  ,  the  molecule  of  which  is 
also  equivalent  to  two  of  nitric  acid.  It  is  therefore 
called  a  di-aci<t  base.  We  shall  see  later  that  there  are 
also  di-basic  acids. 

(We  have  observed  three  ways  in  which  nitric  acid 
forms  salt.  What  are  they  ]) 

The  nitrates  are,  almost  without  exception,  soluble  in 
water.  They  can  all  be  decomposed  by  heat.  The  nitrates 
of  the  heavy  metals  (copper,  iron,  lead,  &c.,)  give  off  a 
mixture  of  oxygen  and  oxides  of  nitrogen,  when  strongly 
heated.  The  nitrates  of  the  alkali  metals  (sodium,  potas- 
sium, &c.,)  decompose  slowly  only  at  a  bright  red  heat, 
giving  off  one-third  of  their  oxygen  and  leaving  nitrites  : 

KNO3  =  KNO2  4-  O. 

The  nitrates  are  used  in  many  operations  as  oxidising 
agents.  Gunpowder  is  a  mixture  of  potassic  nitrate, 
charcoal,  and  sulphur  (What  becomes  of  the  charcoal  and 
sulphur  during  the  explosion?)  Nitrates  of  barium,  cal- 
cium, and  strontium,  are  used  to  prepare  coloured  fires. 

Tests. — Same  as  in  Art.  81,  but  add  sulphuric  acid  in  (1)  and 
(2).  (Why  ?)  Try  these  tests  with  solution  of  saltpetre. 

83.  Salts. — The  general  nature  of  acids  and  bases  can 
now  be  seen.  They  are  opposite  in  properties,  and,  when 
brought  together,  tend  to  neutralise  each  other,  i.e., 
destroy  each  other's  distinctive  properties.  But  only  the 
stronger  acids  and  bases  do  this  completely.  The  sub- 
stances formed  when  acids  and  bases  act  on  each  other 


NOMENCLATURE    OF    SALTS.  89 

are  mostly  like  common  salt  in  their  properties,  and  are 
thence  called  salts.  Their  relation  in  composition  to  the 
acids  and  bases  from  which  they  are  formed  can  be  best 
shown  as  follows  : — 

H.NO3,       nitric  acid.  K.OH       \potassichydroxide. 

K.NO3,       potassic  nitrate.  K.NO3     /       "        nitrate. 

Na.NO3,      sodic  Na  OH     ^    godic      ^vdroxide 

Pb(N03)2,  plumbic  Na  NO      f       "        nitrate 
Bi(NO3)3,  bismuthic  " 

NH^NO.,,    ammonic    "  Pb(OH)2  \plumbic  hydroxide. 

Pb(NO3)2J        "       nitrate. 

Bi(OH)3  \bismuth  hydroxide. 
Bi(NO3)3J        "        nitrate. 

An  inspection  of  these  two  lists  shows  that  the  nitrates 
differ  from  nitric  acid  by  having  metal  instead  of  hydro- 
gen, and  from  the  bases  by  having  NO3  instead  of  OH 
(hydroxyl).  Salts  may  be  regarded  as  made  up  of 
two  parts,  metal,  and  salt-radical.  The  salt  radical  of 
the  nitrates  is  then  — NO3,  and  we  write  the  formula  of 
any  nitrate  by  adding  to  the  symbol  of  the  metal  n  times 
NO3,  n  being  the  valence  of  the  metal.  Thus,  for  monad 
metals  put  NO3  once  ;  for  dyads,  NO3  twice,  and  so  on. 

Names  of  salts  are  formed  from  the  names  of  the  metals 
(as  adjectives),  and  those  of  the  acids  (as  nouns).  The 
syllabic  -ic  is  generally  added  to  the  root  of  the  name  of 
the  metal.  In  case  the  metal  forms  two  bases,  the  salt 
of  that  containing  the  greater  relative  quantity  of  metal  is 
distinguished  by  the  ending  -ous.*  Thus,  ferrous  nitrate, 
Fe(N03)2;  and  ferric  nitrate,  Fe2(NO3)6.  The  names  of 
acids  generally  end  in  -ic,  and  (with  the  exception  of  the 
haloid  acids  and  a  few  others)  this  ending  is  changed  to 

*  Latin  onum,  abounding  in. 


90  OXIDES   OF    NITROGEN. 

-ate  for  the  name  of  the  salt.  When  the  name  of  the  acid 
ends  in  -ous,  the  name  of  the  salt  ends  in  -ite.  Thus, 

Nitric  acid Nitrate. 

Sulphuric       "    Sulphate. 

Nitrous  "    Nitrite. 

Phosphorous  " Phosphite. 

There  is  a  class  of  acids  containing  no  oxygen,  e.g., 
hydrochloric  acid  (HC1),  hydrobromic  acid  (HBr),  <fcc. 
The  salts  of  these  acids  are  called  chloricZes,  bromides,  &c. 
They  are  named  in  the  same  way  as  oxides  and  sulphides. 

84.  Oxides  Of  Nitrogen. — Nitrogen  and    oxygen 
do  not  readily  combine ;  but  by  indirect  methods  5  com- 
pounds can  be  obtained.     If  electric  sparks  be  passed 
through   dry   air,   a  red  gas,   nitrogen   tetroxide   (NO2), 
appears,  and  this  in  presence  of  water  and  oxygen  forms 
nitric  acid.     It  is  probable  that  considerable  quantities 
of  nitric  and  nitrous  acids  are  formed  in  the  atmosphere 
by  lightning.     The  five  oxides  are  as  follows  : — 

Nitrogen  monoxide N2O 

"       dioxide NO. 

"       trioxide N2O3. 

"       tetroxide NO2  or  N2O4. 

"       pentoxide N2O5. 

If  the  formula  of  the  second  be  doubled  it  will  be  easily 
seen  that  the  proportion  of  oxygen  increases  regularly 
from  1  to  5.  This  is  a  good  illustration  of  the  Law  of 
Multiple  Proportions.  (Apply  it.) 

85.  Nitrogen  Monoxide  (N.2O) — Also  called  laugh- 

iny  gas,  and  nitrous  oxide. 

PREPARATION. — Experiment  53. — Dry  some  ammonic  nitrate 
(Exp't  51)  by  fusing  it  in  a  porcelain  dish.  Break  the  dried 
salt  into  small  lumps  and  heat  some  of  it  in  a  test-tube  provided 


NITROGEN  MONOXIDE.  9l 

with  a  gas-delivery  tube.  The  salt  rnelts  easily  to  a  clear  liquid, 
and  a  continuous  stream,  of  gas  comes  off.  Collect  6  jars  of  the 
gas  over  warm  water,  and  try  its  action  on  a  glowing  match  and 
on  burning  phosphorus. 

PROPERTIES. — Colourless  gas,  of  ethereal  smell  and 
sweetish  taste.  ^Try  it.)  (Calculate  its  specific  weight.) 
At  0°  C.  it  becomes  liquid  under  a  pressure  of  about  30 
atmospheres.  It  is  soluble  in  water  to  the  extent  of 
about  1J  vols.  in  1  of  water  (at  0°  C.).  (Try  solubility 
in  cold  water.)  It  supports  ordinary  combustion  better 
than  air  does.  (Compare  the  percentages  of  oxygen  in 
air  and  nitrogen  monoxide).  The  whole  of  the  oxygen  is 
easily  separated  from  the  nitrogen  by  carbon,  phos- 
phorus, ttc. : 

C  +  2N2O  =  2N2  +  CO2 
2P  +  5N2O  =  5N2  +  PA. 

(Translate  these  equations  into  ordinary  language). 

Experiment  54. — Set  fire  to  a  little  sulphur  in  a  deflagrating 
spoon,  and  at  once  plunge  into  a  jar  of  nitrogen  monoxide.  The 
flame  is  extinguished.  Light  it  again,  and  let  it  burn  vigorously 
before  putting  it  into  the  gas. 

Nitrogen  monoxide  extinguishes  a  weak  flame  of  sul- 
phur, because  the  temperature  is  not  high  enough  to 
start  the  action.  Sulphur  burning  well  in  air,  burns 
more  brightly  in  nitrogen  monoxide.  The  substances 
formed  are  nitrogen  and  sulphur  dioxide  (SO2).  (Write 
the  equation.) 

When  breathed,  nitrogen  monoxide  causes  at  first  a 
peculiar  intoxication  with  insensibility  to  pain.  If 
mixed  with  about  one-fifth  of  its  volume  of  air  it  can  be 


92  NITROGEN    DIOXIDE. 

breathed  for  some  time  ;  but,  as  it  diffuses  into  the  blood, 
and  cannot  supply  oxygen  to  the  blood,  symptoms  of  suf- 
focation at  length  appear.  When  prepared  for  inhala- 
tion, it  should  be  freed  from  acid  fumes  and  nitrogen 
dioxide  by  bubbling  it  through  solutions  of  caustic  soda 
and  ferrous  sulphate. — Salts  of  an  acid,  hyponitrous,  cor- 
responding to  this  oxide,  are  known ;  e.g.,  KNO,  pot. 
hyponitrite. 

86.  Nitrogen  Dioxide. — NO.  (Nitric  Oxide). 

PREPARATION — Experiment  55. — Mix  1  part  strong  nitric 
acid  with  3  of  water  in  a  t.  t.  or  flask,  and  add  some  scraps  of 
copper.  Heat  gently  and  collect  six  jars  of  the  gas  over  water. 

Other  metals,  as  mercury,  silver,  iron,  lead,  &c.,  may 
be  used.  Nitrogen  and  the  monoxide  are  always  pi-esent 
as  impurities. 

PROPERTIES. — An  invisible  gas,  a  little  heavier  than 
air.  (Calculate  specific  weight.)  Smell  and  taste  not 
known.  (Why1?)  It  can  be  condensed  to  a  liquid  at 
-  11°  0.  by  a  pressure  of  104  atmospheres.  It  is  the 
most  stable  of  all  the  oxides  of  nitrogen.  It  is  very 
sparingly  soluble  in  water. 

Experiment  56. — Set  fire  to  a  bit  of  phosphorus  and  plunge  it 
quickly  into  a  jar  of  the  gas.  It  is  extinguished.  Try  again, 
allowing  the  phosphorus  to  burn  brightly  before  putting  it  in  the 
gas.  Try  a  lighted  match,  and  sulphur  burning  strongly.  They 
are  both  put  out. 

Experiment  57-— Pass  up  a  little  oxygen  into  a  jar  of  the 
gas  over  water  ;  red  fumes  appear  and  then  disappear,  the 
volume  of  gas  being  lessened  : 

NO  +  O  =  N02 

3  N02  +  H2O  =  2HN03  +  NO. 


NITROGENT   TRIOXIDE.  93 

Nitrogen  dioxide  supports  only  very  vigorous  combus- 
tion. It  combines  with  oxygen  to  form  nitrogen  tet- 
roxide  (N02).  (What  takes  place  on  contact  with  air?) 

Experiment  58- — Pour  some  solution  of  green  vitriol  into  a 
jar  of  the  gas,  close,  and  shake  up.  The  gas  dissolves,  giving 
a  dark  brown  colour  to  the  solution.  A  weak  compound 
(2FeS04.NO)  is  formed.  This  can  be  decomposed  by  heat,  giv- 
ing pure  nitrogon  dioxide.  (In  what  former  experiment  was 
this  compound  obtained  ?) 

In  writing  the  graphic  formula  of  nitrogen  dioxide, 
either  the  nitrogen  must  be  represented  as  dyad,  N  Q, 
or  one  atomic  bond  must  be  left  unemployed,  —  N~O. 
In  either  case  the  compound  seems  anomalous.  Other 
cases  of  this  kind  warn  us  that  the  idea  of  valence  is  one 
to  be  used  cautiously,  however  convenient  ordinarily. 

87.  Nitrogen  Trioxide. — N203.    It  is  also  called 

anhydrous  nitrous  acid. 

PREPARATON. — Late  experiments  seem  to  show  that 
this  substance  does  not  exist  at  ordinary  temperatures, 
but  decomposes  above  —  18°  C.  It  can  be  prepared  by 
mixing  nitrogen  dioxide  with  one-fourth  its  volume  of 
oxygen,  and  cooling  to  —  20°  C. 

Experiment  59. — Heat  a  little  starch  with  some  nitric  acid 
diluted  with  about  one  volume  of  water.  Red  vapours  are 
evolved,  said  to  be  nitrogen  trioxide.  If  these  are  cooled  below 
—  18°  C.  they  form  a  deep  blue  liquid. 

PROPERTIES. — Nitrogen  trioxide  is  a  blue  liquid  which 
decomposes  at  —  18°  C.  into  the  dioxide  and  tetroxide  : 

NO   =  NO       NO. 


94  NITROUS    ACID. 

Experiment  60- — Repeat  Experiment  59  on  a  larger  scale 
and  lead  the  red  gas  into  a  t.  t.  containing  water  and  surrounded 
by  ice  and  salt.  A  blue  solution  is  formed,  having  acid  proper- 
ties. Add  some  caustic  soda  solution  to  this.  The  colour  dis- 
appears. Evaporate  on  the  water  bath,  and  keep  the  remaining 
salt,  sodic  nitrite. 

88.  Nitrous  Acid.-  HNO,  This  is  the  blue  solu- 
tion of  Experiment  60.  It  has  not  been  obtained  pure, 
and  is  easily  decomposed.  Its  relation  to  nitrogen  tri- 
oxide  is  seen  thus  : 

N2O3  +  H20  =  2HNO2. 

It  is  a  monobasic  acid,  and  forms  salts  called  nitrites. 
The  "  nitrous  acid  "  of  pharmacy  is  impure  nitric  acid. 

NITRITES. — The  nitrites  are  nearly  all  soluble  in  water, 
and  quickly  absorb  oxygen  to  form  nitrates  : — KNO2  -}-  O 
=  KNO3.  Nitrites  are  produced  by  the  decomposition 
of  animal  matter,  and  their  presence  in  water  is  an 
indication  of  pollution. 

Tests. — •!.  Put  a  few  drops  of  dilute  sulphuric  acid  on  a  little 
of  the  salt  from  Expt.  60.  Red  fumes  are  evolved  : 

2NaN02  +  H,S04  =  Na2S04  +  N203  (?)  -f  H20. 

2.  Add  a  little  potassic  iodide  and  starch  paste  to  a  solution  of 
sodic  nitrite ;  to  this  add  a  few  drops  of  dilute  sulphuric  acid. 
The  deep  blue  of  'iodide  of  starch  appears. 

3.  Add  a  little  potassic  permanganate  (KMn04)  to  solution  of 
any  nitrite  ;  then  some  acetic  acid.     The  colour  of  the  perman- 
ganate dissappears.     It  is  reduced  by  the  nitrous  acid.     Try  the 
acetic  acid  without  the  nitrite. 


TETROXIDE    AND    PENTOXIDE   OF    HYDROGEN.  95 

89.  Nitrogen  Tetroxide,  NO2  or  N2O4.— This  is 

also  called  peroxide  of  nitrogen. 

PREPARATION. — Experiment  61. — Dry  some  plumbic  nitrate 
(Expt.  52)  by  gently  heating  in  a  porcelain  dish  ;  transfer  to  a 
hard-glass  t.  t.  and  heat  gradually.  Receive  the  red  gas  in  a  1. 1. 
surrounded  by  a  freezing  mixture  of  snow  and  salt  : 

Pb(NOa)2  ^  PbO  +  N204  +  O. 

(What  is  the  substance  left  in  the  t.  t.  ?  How  can  it  be 
again  converted  into  lead  nitrate  1)  Nitrogen  tetroxide 
can  also  be  prepared  by  cooling  a  mixture  of  nitrogen 
dioxide  and  half  its  volume  of  oxygen  :  NO  -f-  0  =  NO2. 

PROPERTIES. — A  colourless  liquid  at  low  temperatures, 
a  reddish  brown  gas  above  22°  C.,  its  boiling  point. 
With  cold  water  it  forms  nitric  and  nitrous  acids  : 
N2O4  +  H,O  =  HNO3  -f  HNO2.  Similarly,  with  caus- 
tic soda  or  caustic  potash  it  forms  nitrate  and  nitrite 
(and  water).  (Write  equations).  It  is  irritating  when 
breathed,  and  injures  the  mucous  membranes  of  the 
respiratory  passages. 

90.  Nitrogen   Pentoxide.  —  N2O5.     Also  called 

anhydrous  nitric  acid. 

PREPARATION. — By  passing  dry  chlorine  gas  through 
a  tube  containing  dry  silver  nitrate  heated  to  50°  C.;  or 
by  distilling  pure  nitric  acid  with  phosphorus  pentoxide : 

2HN03  +  P2O5  =  2HPO3  +  N2O5. 

PROPERTIES. — A  white  crystalline  solid,  beginning  to 
decompose  at  40°  C.  into  tetroxide  and  oxygen.  (Write 


96  QUESTIONS    AND    EXERCISES. 

the  equation.)  It  is  very  hygroscopic,  and  rapidly  deli- 
quesces in  moist  air.  It  unites  with  water  forming 
nitric  acid  : 

KjO5  +  H2O  =  2HNO.,. 


QUESTIONS  AND  EXERCISES. 

1.  Calculate  the  weights  of  1  litre  of  nitrogen  and  of  ammonia, 
measured  at  0°  and  760. 

2.  Calculate  the  volume  of  air  to  give  20  litres  of  nitrogen. 

3.  What  weight  of  ammonic  chloride  is  required  to   give  5 
litres  of  ammonia  gas  measured  at  17°  C.  and  700  mm.  pressure  ? 

4.  Calculate   the    volume    of   ammonia   gas   at    20°   C.    and 
760  mm.  formed  by  heating  100  grams  ammonic  chloride  with 
lime. 

5.  Hold  the  moist  stoppers  of  the  ammonia  and  hydrochloric 
acid  bottles  near  each  other.     Explain  what  you  observe. 

6.  Warm  some  ammonic   chloride  solution  with  sodic   hyd- 
roxide  solution.     Observe   the   smell.     What   substances   have 
been  formed  ?     Write  the  equation. 

7.  In  determining  the  composition  of  ammonia  (Art.  80),  how 
can  the  hydrogen  be  got  rid  of  in  order  to  measure  the  volume 
of  nitrogen  ? 

8.  What  weight  of  sodic  nitrate  (NaNO3)  is  required  to  pre- 
pare 200  grams  of  nitric  acid?     How   much  sulphuric  acid  is 
used? 

9.  How  much  nitric  acid  can  be  obtained  by  the  decomposi- 
tion of  10  Ibs.  of  saltpetre  ? 


QUESTIONS    AND    EXERCISES.  97 

10.  Lime  water  is  an  antidote  to  poisoning  by  nitric  acid. 
Explain  its  action. 

1 1 .  Calculate  the  weights  of  nitric  acid  required  to  neutralise 
10  grams  each  of  the  following  bases  and  oxides  : — Sodic  hyd- 
roxide   (NaOH),  potassic  hydroxide    (KOH),    calcic   hydroxide 
(Ca(OH)2),  magnesia  (IvlgO),  and  litharge  (PbO). 

12.  Write  the  formulas  for  nitrates  of  the  following  metals, 
referring  to  the  table  of  elements  for  the  valences: — Calcium, 
silver,  iron,  mercury,    copper,    cobalt,   barium,  magnesium,  and 
aluminium. 

13.  What  experiments  already  made  illustrate  the  direct  re- 
placement of  hydrogen  by  metals  ? 

14.  How  many  litres  of  nitrogen  monoxide,  measured  at  15°  C. 
and  730  mm.  pressure,  can  be  prepared  from  100  grams  ammonic 
nitrate  ?     What  weight  of  water  is  formed  ?     What  volume  will 
it  occupy  at  200°  C.  and  760  mm.  pressure  ? 

15.  How  much  ammonic  nitrate  must  be  used  to  fill  a  10  gal- 
lon receiver  with  nitrous  oxide  under  a  pressure  of  60  Ibs.  to  the 
square  inch  when  the  thermometer  reads  10°  C.  ? 

16.  What  volume  of  nitrogen  monoxide  at  16°  C.  and  750  mm. 
will  completely  burn  1  gram  carbon  ? 

C  +  2N20  -=  2N2  -f-  C02. 

17.  What  volume  of  air  will  convert  10  cubic  inches  of  nitro- 
gen dioxide  into  the  tetroxide  ? 


98  SEA    WATER. 


CHAPTER     IX. 


THE  HALOGENS. 

91.  Sea  Water.  —  Sea  water  is  a  dilute  solution  of 
salts,  some  in  much  larger  proportions  than  others.  The 
total  quantity  in  100  parts  is  about  3.5.  The  specific 
weight  of  sea  water  at  0°  C.  is  1.03.  Common  salt 
(NaCl)  is  present  in  the  greatest  proportion  ;  next  come 
magnesic  chloride  (MgCl2),  magnesic  sulphate  (MgS04), 
calcic  sulphate  (CaSO4),  potassic  chloride  (KC1),  and 
magnesic  bromide  (MgBr2).  The  bitter  taste  is  due  to 
the  magnesic  salts.  Magnesic  iodide  (MgI2)  is  present 
in  sea  water  in  very  minute  quantity. 

When  sea  water  is  evaporated,  the  sparingly  soluble 
gypsum  (CaS04.2H2O)  first  crystallises  out,  and  later, 
sodic  chloride.  This  latter  salt  is  a  compound  of  the 
metal  sodium  with  a  non-metal,  chlorine,  which  is 
therefore  called  a  halogen,  or  "  salt  generator."  The 
"  mother  liquor,"  an  intensely  bitter  liquid  (bittern), 
contains  compounds  of  two  elements  very  similar  to 
chlorine,  viz.,  bromine  and  iodine;  and  these  three, 
with  a  fourth,  form  the  group  of  halogens,  which 
are  peculiar  among  the  elements,  as  forming  with 
the  metals  salts  which  contain  no  third  element,  e.g., 
NaCl,  KBr,  MgI2,  CaF2,  <fec.  (Compare  with  KNO3, 
Na2SO4,  <fec.)  Sea  water  is  the  principal  source  of  the 
first  three  of  the  halogens.  —  (In  the  following  sections 
note  resemblances  and  differences  among  the  halogens.) 


CHLORINE.  99 

CHLORINE  AND  ITS  COMPOUNDS. 

92.  Chlorine.— (Cl' =  35.37).  Very  widely  diffused 
in  water  and  on  land  ;  never  free,  but  in  combination  as 
chlorides. 

PREPARATION. — Experiment  62-*— Put  a  little  sodic  chlo- 
ride mixed  with  manganese  dioxide  (Mn02)  in  a  t.t.,  pour 
some  Bulphuric  acid  in,  and  heat  gently.  Chlorine  gas  is  set 
free.  Note  colour  and  smell  (cautioiuly) . 

This  chemical  action  is  most  clearly  represented  in 
three  steps  : 

(1)  H2SO<  -f  NaCl  =  NaHSO,  -f  HC1. 

(2)  MnO,  -f  H2S04  =  MnSO4  +  H2O  -f  0. 

(3)  2HC1  -f  O  =  H2O  -f  C12. 

The  substances  ultimately  formed  are  sodic  hydric  sul. 
phate  (NaHSO4),  manganous  sulphate  (MnSO4),  chlorine, 
and  water.  (Write  a  single  equation  representing  the 
action.  You  know  the  formulas  of  all  the  substances. 
Remember  that  the  equation  must  balance.) 

Experiment  63-* — Heat  gently  a  little  manganese  dioxide 
and  hydrochloric  acid  in  a  test-tube.  Chlorine  is  given  off : 

4HC1  +  Mn02  =  MnCla  +  Cla  +  2HaO. 

This  is  the  method  by  which  chlorine  is  generally  pre- 
pared on  the  large  scale. 

Experiment  64.* — Warm  a  little  hydrochloric  acid  with  a 
small  crystal  of  potassic  chlorate  (KCIO^),  and  note  the  evolu- 
tion of  chlorine :  KC103  -f  6HC1  =  KC1  +  3C12  -f  3H2O. 

All  these  and  similar  methods  may  be  considered  as  so 
many  ways  of  oxidising  hydrochloric  acid  : 


1  Fill  the  test-tube  with  dilute  caustic  soda  when  the  experiment  is  finished. 


100  CHLORINE. 

2HC1  +  O  =  H2O  +  CL, 

and,  indeed,  chlorine  is  prepared  on  the  lai-ge  scale  by 
passing  a  mixture  of  air  and  hydrochloric  acid  over  hot 
bricks. 

PROPERTIES. — A  heavy,  greenish  yellow  gas ;  is  liquid 
at  0°  under  a  pressure  of  six  atmospheres ;  causes  suffoca- 
tion when  breathed ;  has  violent  irritating  action  on 
mucous  membranes,  and  may  cause  catarrh  or  ulceration  ; 
soluble  in  water,  3  vols.  in  1  at  10°  C. ;  the  aqueous  solu- 
tion acts  as  a  powerful  irritant  both  externally  and  inter- 
nally ;  antidotes  are  white  of  egg,  milk,  ammonia,  lime 
water,  soap,  and  other  alkaline  substances.  Chlorine,  both 
as  gas  and  solution,  is  a  powerful  antiseptic  and  deodoriser; 
it  is  an  antidote  to  poisoning  by  prussic  acid  (HGN), 
sulphuretted  hydrogen  (H2S),  and  ammonic  sulphide 
((NH4)2S).  The  action  on  prussic  acid  has  not  been  ex- 
plained ;  that  on  sulphuretted  hydrogen  and  amrnonic 
sulphide  is  explained  by  the  following  equations  : 

H2S  -f  C12  =  2HC1  +  S 
(NHJ2S  +  4C!2  =  8HC1  +  N2  -f  S. 
The  substances  formed  are  not  poisonous. 

Experiment  65. — Prepare  a  little  chlorine  as  in  Experiment 
63,  and  hang  a  narrow  strip  of  moist  turkey-red  cloth  in  the  t.t. 
It  is  bleached  white.  Try  other  vegetable  colours,  e.g.,  a  small 
flower.  Try  a  mineral  colour,  e.g.,  red  lead. 

Chlorine  and  water  bleach  organic  colouring  matters, 
but  not  mineral.     The  presence  of  water  is  necessary. 
If  dry  chlorine  be  used,  it  does  not  bleach.     A  solution 
of  chlorine  in  water  gradually  decomposes  : 
H2O  +  C12  =  2HC1  +  O. 


HYDROCHLORIC    ACID.  101 

This  shows  the  nature  of  the  bleaching  action  of  chlo- 
rine. It  is  really  an  oxidation  by  means  of  the  oxygen 
of  water,  the  hydrogen  combining  with  chlorine  to  form 
hydrochloric  acid.  The  colouring  matter  is  destroyed, — 
generally  oxidised  to  carbon  dioxide  and  water. 

Experiment  66. — Prepare  some  chlorine  water,  by  leading  the 
gas  prepared  as  in  Experiment  63  into  a  bottle  of  water,  using  a 
gas-delivery  tube  bent  twice  at  right  angles  and  reaching  to 
the  bottom  of  the  bottle.  (This  should  be  done  under  a  hood  or 
in  a  draught  cupboard.)  Note  the  colour  of  the  solution.  Try 
its  effect  on  litmus.  Explain. 

Chlorine  combines  directly  with  most  elements  and 
thereby  forms  chlorides. 

Experiment  67-  -Prepare  a  little  chlorine  (Experiment  63) 
and  drop  some  powdered  antimony  into  the  tube.  It  catches 
fire  as  soon  as  it  falls  into  the  gas :  Sb  +  5C1  =  SbCl6.  Try  a 
bit  of  ph  sphorus  in  the  deflagrating  spoon  ;  and  a  lighted  taper. 

Chlorine  supports  the  combustion  of  tallow,  wax,  tur- 
pentine, and  other  organic  substances  containing  a  large 
percentage  of  hydrogen.  It  combines  with  the  hydrogen 
to  form  hydrochloric  acid,  setting  the  carbon  free. 

Test. — Add  some  starch  paste  to  a  dilute  solution  of  potaa- 
sic  iodide  (KI),  and  then  a  few  drops  of  chlorine  water.  The 
deep  blue  iodide  of  starch  is  formed  :  KI  +  Cl  =  KC1  +  I- 

Compounds  of  Chlorine. 

93.   Hydrochloric  Acid  (HC1  =  36.37.)— Also 

called  muriatic  acid,  and  spirit  of  salt.  If  equal  volumes 
of  hydrogen  and  chlorine  be  mixed  in  the  dark  and  then 
exposed  to  diffused  daylight  they  combine  without  any 


102  HYDROCHLORIC    ACID. 

change  of  volume,  forming  a  colourless  gas,  hydrochloric 
acid.  If  exposed  to  direct  sunlight  they  combine  with  a 
violent  explosion. 

PREPARATION. — Experiment  68. — Put  some  drysodic  chlo- 
ride in  a  t.t.,  pour  strong  sulphuric  acid  over  it,  fit  the  gas-delivery 
tube  used  in  Experiment  66,  warm  gently,  and  collect  a  t.  t.  of 
the  gas  by  displacting  the  air.  (You  can  see  when  the  t.  t.  is 
full  by  the  fuming  of  the  gas  at  the  top  ;  and  by  its  overflowing  on 
the  fingers  and  making  them  feel  warm.)  Close  with  the  thumb 
and  open  under  water.  Test,  by  taste  and  litmus,  the  water 
which  rises  in  the  tube. 

Solution  of  hydrochloric  acid  is  prepared  in  this  way 
on  the  large  scale  as  a  by-product  of  the  alkali  manufac- 
ture : 

2NaCl  +  H2SO4  =  2HC1  +  Na2S04. 

(Write  the  names  and  weights.  What  remains  in  the 
t.  t.  in  Experiment  68  1) 

PROPERTIES. — A  colourless  gas,  soluble  in  water,  500 
vols.  in  1  at  0°  C.;  the  solution  is  heavier  than  water. 
The  saturated  solution  becomes  weaker  on  boiling  until 
it  contains  20^  p.  c.  of  the  acid,  then  distils  unchanged 
at  110C  C.  The  specific  weight  of  this  solution  is  1.1 1. 
(What  similar  case  has  been  studied  already  '<)  The 
specific  weight  of  the  gas  is  1.269  (air  =1).  At —  4°  C. 
and  with  a  pressure  of  25  atmospheres  it  condenses  to  a ' 
colourless  liquid. 

Experiment  69. — Colour  some  solution  of  hydrochloric  acid 
with  litmus,  and  add  to  it  a  solution  of  caustic  soda  until  the 
litmus  begins  to  turn  blue.  Taste  the  solution.  (What  does  it 
taste  of?)  Evaporate  to  dry  ness  on  the  water  bath.  Crystals  of 
common  salt  remain:  NaOH  +  HC1  =  NaCl  +  H20.  Try 
the  same  with  potassic  hydroxide. 


CHLORIDES. 


103 


Hydrochloric  acid  is  a  monobasic  acid,  as  can  be  seen 
from  its  formula.  Its  molecule  contains  only  one  atom 
of  hydrogen,  which  cannot  be  partially  replaced  by  atoms 
of  metals.  It  is  called  a  haloid  acid.  Its  aqueous  solu- 
tion dissolves  tin,  iron,  zinc,  aluminium,  and  other  metals. 

CHLORIDES.  —  These  are  formed  by  replacing  the 
hydrogen  of  the  acid  by  metals,  a  monad  replacing  H,  a 
dyad  2H,  a  triad  3H,  &c.  (Write  the  formulas  for 
chlorides  of  silver,  mercury,  lead,  copper,  nickel,  anti- 
mony, and  tin.) 

Experiment  70-  —  Drop  a  little  hydrochloric  acid  solution 
into  solutions  of  argentic  nitrate  (AgN03),  mercurous  nitrate 
(Hg2(N03)2),  and  plumbic  acetate  (Pb(C2H3O1!).J).  Precipitates 
are  formed.  Note  any  differences  in  their  appearance. 

Mercurous  chloride,  and  the  chlorides  of  silver  and 
lead  are  insoluble  in  water  (plumbic  chloride,  sparingly 
soluble)  ;  the  chlorides  of  the  other  metals  are  soluble 
(but  a  few  are  decomposed  on  contact  with  water). 


AgNO3  +  HC1  =  AgCl  +  HNO3. 

The  silver  and  hydrogen  atoms  exchange  places. 
equations  for  the  other  two.) 


(Write 


Hydrochloric  acid  is  irritating  when  breathed  as  a  gas.  The 
strong  solution  is  a  corrosive  poison.  (Alkaline  substances  are 
antidotes  ;  magnesia,  lime  water,  or  aoap  may  be  used.)  A 
very  dilute  solution  is  a  tonic.  Hydrochloric  acid  is  secreted 
into  the  stomach  during  digestion,  in  which  it  plays  an  impor- 
tant part. 

Tests.  —  Add  a  few  drops  of  argentic  nitrate  (AgNO,).  A 
curdy  white  precipitate  (AgCl)  is  formed,  insoluble  in  nitric  acid, 
soluble  in  ammonia  solution.  (Try  with  solution  of  hydrochloric 
acid,  and  also  with  a  chloride.  ) 


104  OXIDES    OF    CHLORINE. 

94.  Oxides  of   Chlorine. — Chlorine    and    oxygen 
do   not    unite   directly,    but   by   indirect   methods   three 
compounds    can    be    obtained,    viz.,   chlorine    monoxide, 
C12O  ;  trioxide,  C12O3 ;  and  tetroxide,   C12O4.      They  are 
dangerously  explosive  compounds  of  little  importance. 

95.  Chlorine  Monoxide.— C12O.     A  yellow  gas 

prepared  by  the  action    of  dry  chlorine  on  dry  mercuric 
oxide  : 

2HgO  +  2C12  =  HgClg-HgO  -f  C12O. 

It  dissolves  in  water  forming  hypochlorous  acid  : 
C12O  +  H2O  =  2HC1O. 

96.  Chlorine  Trioxide. — C1203.      A    greenish    yel- 
low gas,  of  irritating  action  when  breathed,  dangerously 
explosive.      Prepared    by   warming  a  mixture  of  nitric 
acid,  potassic  chlorate,  and  sugar  :   2HC103  -|-  N2O3  = 
C12O3  -j-  2HNO3.     It  dissolves  in  water  forming  chlorous 
acid : 

C1,O3  -f-  H2O  =  2HC1O2. 

97.  Chlorine  Tetroxide.— C1204.     Also  called  per- 
oxide of  chlorine.      It  is  a  yellow  gas  of  not  unpleasant 
odour.     Yery  explosive  when  heated. 

Experiment  71. — Put  a  small  crystal  of  potassic  chlorate  in 
a  t.t. ,  held  by  the  t.t.  holder  or  a  pair  of  forceps.  Add  a  few 
drops  of  strong  sulphuric  acid,  and  heat  gently,  taking  care  that 
the  mouth  of  the  t.t.  is  directed  away  from  everybody.  The  gas 
is  evolved  and  quickly  explodes  with  violence  : 

3KC10S  +  2H,S04  =  KC104  -f  2KHSO<  -f  C12O4  +  H.O. 


HYPOCHLOROUS    ACID.  105 

It  dissolves  in  a  caustic  potash  solution  forming  chlorite  and 
chlorate.  (What  substance  already  studied  is  similar  ?) : 

C1.,O4  +  2KOH  =  KC10,  +  KC108  +  H20. 

98.  Oxygen  Acids  of  Chlorine. — Like  nitrogen 
chloi'ine  combines  with  hydrogen  and  oxygen  in  several 
proportions,  forming  a  series  of  oxygen  acids  : 

Hypochlorous  acid HC10 

Chlorous  acid HCK)a 

Chloric  acid HC103 

Perchloric  acid HC1O4 

The  names  of  these  acids  illustrate  very  well  the  use  of 
the  terminations  -ic  and  -out ;  and  of  the  prefixes  hypo- 
and  per-.  HypocUorous  acid  is  below  chlorous  acid  in 
proportion  of  oxygen.  Perchloric  contains  more  oxygen 
than  chloric. — These  acids  are  all  monobasic,  and  mostly 
unstable. 

99.  Hypochlorous  Acid.— HC1O. 

PREPARATION. — By  the  action  of  chlorine  water  on 
freshly  precipitated  mercuric  oxide  : 

HgO  +  H,O  -f  2C12  =  HgCl2  +  2HC1O. 

PROPERTIES. — It  has  never  been  obtained  except  as 
a  dilute  aqueous  solution.  It  has  a  pleasant,  somewhat 
chlorous,  smell  (that  of  bleaching  powder),  and  bleaches 
powerfully. 

HYPOCHLORITES. — The  salts  of  hypochlorous  acid  are 
important,  particularly  bleaching  powder  : 

(CaCl,  .  Ca(OCl),). 


106  CHLOROUS    ACID. 

Experiment  72- — Generate  some  chlorine  as  in  Experiment 
62,  and  pass  it  for  a  short  time  into  a  beaker  containing  a  very 
dilute  solution  of  sodic  hydroxide  : 

2NaOH  +  C12  =  NaCl  +  NaCIO  +  H,0. 

Sodic  chloride,  sodic  hypochlorite,  and  water  are 
formed.  Keep  the  solution  for  further  experiments. 

Experiment  73 — Moisten  a  small  piece  of  tiirkey-red  cloth 
with  acetic  acid,  and  put  it  into  a  portion  of  the  solution  from 
Experiment  72.  The  colour  is  bleached. 

Experiment  74. — Heat  another  portion  of  the  solution  nearly 
to  boiling,  and  then  try  Experiment  73  with  it.  The  colour  is 
not  bleached. 

Hypochlorites  are  readily  decomposed  by  heat  into 
chlorides  and  chlorates,  one  portion  giving  up  its  oxygen 
to  the  other : 

3KC1O  =  KC1O3  +  2KC1. 

Experiment  75- — Warm  a  third  portion  gently  with  a  little 
hydrochloric  acid  and  note  the  smell  of  chlorine  : 

KC10  +  HC1  =  HC10  +  KC1 
HC10  +  HC1  =  Clt  +  H20. 

Hypochlorous  acid  is  a  very  weak  acid,  its  salts  being 
decomposed  even  by  carbonic  acid.  Hence  the  chlorous 
smell  of  bleaching  powder.  (Explain). 

100.  Chlorous  Acid. — HC1O2.    Cannot  be  obtained 
except  as  a  solution  in  water  (Art.  96).     Its  salts,  the 
chlorites,  are  unstable  and  unimportant. 

101.  Chloric  Acid. — HC1O3.     There  is  no  corres- 
ponding oxide  known.     (What  would  its   formula  be  ? 
Compare  with  nitric  acid.) 


CHLORIC    ACID.  107 

PREPARATION. — By  action  offluosilicic  acid  on  solution 
of  potassic  chlorate.  An  insoluble  salt  of  potassium  is 
formed  (precipitated),  and  chloric  acid  remains  in  solu- 
tion : 

2KC1O3  -f  H2SiF6  =  K2SiF6  -f  2HC103. 

PROPERTIES. — It  forms  a  colourless,  acid  solution,  of 
strong  oxidising  properties.  It  readily  decomposes  into 
chlorine,  oxygen,  and  perchloric  acid,  and  cannot  bo  ob- 
tained free  from  water. 

CHLORATES. — Chloric  acid  is  monobasic,  and  its  salts 
are  all  soluble  in  water.  They  can  be  prepared  by  dis- 
solving the  corresponding  bases  in  chloric  acid ;  but 
those  of  the  stronger  bases  can  be  formed  directly  by  the 
action  of  chlorine.  Thus,  when  chlorine  gas  is  passed 
into  a  hot  solution  of  potassic  hydroxide,  chloride  and 
chlorate  are  formed  (Exp't  74)  : 

6KOH  +  3C12  =  KC1O3  -f  5K01  +  3H2O. 

Similarly  with  sodic,  calcic,  baric,  and  other  hydrox- 
ides. The  weaker  bases,  however,  do  not  react  in  this 
way.  Potassic  chlorate  is  only  sparingly  soluble,  and 
can  thus  be  separated  from  the  more  soluble  chloride. 
(How  1) 

Experiment  76. — Pass  chlorine  gas  through  a  strong  solution 
of  caustic  potash  until  it  begins  to  smell  of  chlorine.  Evaporate 
on  the  water  bath  until  crystals  begin  to  form,  then  set  aside 
to  cool.  Pour  off  the  liquor,  dry  the  crystals  on  filter  paper  and 
keep  them. 

Tests. — 1.  Dry  chlorates  are  known  by  the  formation  of 
chlorine  tetroxide  when  acted  on  by  strong  sulphuric  acid. 

2.  Colour  with  indigo  sulphate,  add  a  few  drops  of  sulphuric 
acid,  and  boil.  The  colour  disappears.  (What  other  acid  an- 
swers to  this  test  ?  How  would  you  distinguish  ?) 


108  PERCHLORIC    ACID. 

102.  Perchloric  Acid. — HC1O4.  This  is  the  most 
stable  of  all  the  oxygen  acids  of  chlorine. 

PREPARATION. — By  distilling  chloric  acid  : 

3HC1O3  =  HC1O4  +  C12  +  2O.,  -f  H2O. 

Also,  by  the  action  of  fluosilicic  acid  on  solution  of  potas- 
sic  perchlorate  : 

2KC1O4  +  H2SiF6  =  K2SiF6  -f  2HC1O4. 

(How  are  the  two  substances  separated  I) 

PKOPEKTIES. — It  is  a  strong  acid,  of  very  great  oxid- 
ising power,  exploding  violently  when  dropped  on  char- 
coal. It  sets  fire  to  wood,  paper,  <fcc.  The  acid  and  its 
salts  are  unimportant  practically. 

PERCHLORATES. — Potassic  perchlorate  is  prepared  by 
heating  potassic  chlorate  until  it  fuses  and  at  length  be- 
comes nearly  solid  again  : 

2KC1O3  =  KC1O4  +  02  +  KC1. 
(How  can  the  two  salts  formed  be  separated?) 


BROMINE  AND  ITS  COMPOUNDS. 

103.  Bromine. — (Br'  =  79.75.)  Found  as  bromides 
in  sea  water,  salt  springs,  and  crude  Chili  saltpetre ;  also 
as  silver  bromide  (AgBr)  in  some  silver  mines.  Certain 
sea-weeds  extract  bromine  and  iodine  from  sea  water  and 
store  it  up  in  their  tissues.  These  two  elements  are  pre- 
pared generally  in  one  operation  from  the  "  mother 
liquor "  of  common  salt,  kelp  (ashes  of  sea-weed),  or 
Chili  saltpetre. 

PREPA RATON. — A  small  sample  of  the  liquor  is  first 
analysed  and  the  quantities  of  iodine  and  bromine  de- 


BROMINE.  109 

termiued.  Then  enough  manganese  dioxide  and  sul- 
phuric acid  to  set  free  the  whole  of  the  iodine  is  added 
to  a  large  quantity  of  liquor  :  2KI  -f-  3H2SO4  -f  MnO2 
=  2KHSO,  +  MnSO4  +  I2  +  2H2O.  The  iodine  is 
distilled  off,  more  manganese  dioxide  and  sulphuric 
acid  are  added  to  set  free  the  bromine,  which  is  then 
distilled  and  condensed  in  separate  receivers  :  2KBr  -f- 
3H2SO4  +  MnO2  =  2KHSO4  +  MnSO4  +  Br,  + 
2H2O. 

Experiment  77. — Mix  about  equal  quantities  of  well  pow- 
dered manganese  dioxide  and  potassic  bromide,  put  in  a  t.t., 
add  a  little  sulphuric  acid,  and  heat  gently.  Bromine  is  evolved, 
condensing  on  the  sides  of  the  t.t.  as  a  dark  reddish-brown 
liquid.  Take  care  not  to  breathe  the  vapour.  Pour  some  of  the 
heavy  vapour  into  a  second  t.t.  containing  a  little  water,  and 
shake  it  up.  The  water  dissolves  the  bromine. 

PROPERTIES. — A  heavy  liquid  (specific  weight  =  3.18), 
almost  black  when  in  mass ;  dark  red,  in  thin  layers. 
It  is  the  only  liquid  element  at  ordinary  temperature, 
excepting  mercury.  It  freezes  at  —  22°  C.,  and  boils  at 
63°  C.,  forming  a  reddish-brown  vapour,  which  has  a  very 
unpleasant  smell  and  an  irritating  action  on  the  mucous 
membranes.  On  account  of  its  low  boiling  point  and 
corrosive  action  bromine  should  be  handled  very  care- 
fully. Its  vapour  should  never  be  inhaled  unless  diluted 
with  much  air. 

Experiment  78. — Try  to  bleach  with  bromine  water  made  in 
Experiment  77. 

Bromine  is  soluble  in  water  to  about  3%.  It  bleaches 
in  the  same  manner  as  chlorine,  but  not  so  powerfully. 

Experiment  79.  —  Dissolve  a  small  crystal  of  potassic  bro- 
mide (KBr)  in  a  test-tube  one-third  full  of  water,  and  add  a  few 


110  HYDKOBROMIC    ACID. 

drops  of  chlorine  water.  Bromine  is  set  free.  Add  a  few  drops 
of  carbon  bisulphide,  close  with  the  thumb  and  shake  the  t.t. 
violently.  Allow  to  stand  a  moment  and  observe.  (What 
property  of  bromine  is  illustrated  ?) 

Chlorine  displaces  bromine  from  combination  with 
metals ;  in  other  words  chlorine  has  greater  chemical 
affinity  (chemism)  for  the  metals. 

COMPOUNDS  OF  BROMINE. 

104.  Hydrobromic  Acid.— HBr. 

PREPARATION. — Experiment  80. — Add  a  few  drops  of 
strong  sulphuric  acid  to  a  crystal  of  potassic  bromide  and  warm. 
(What  action  would  you  expect  ?  What  do  you  observe  ?) 
Hydrobromic  acid  cannot  be  prepared  in  this  way  because  of  its 
decomposition  by  strong  sulphuric  acid ; 

H8S04  +  2HBr  =  Br,  +  S02  +  2H20. 
Compare  hydrochloric  acid. 

Hydrobromic  acid  is  prepared  by  the  action  of  phos- 
phorus and  bromine  on  water.  A,  bromide  of  phos- 
phorus is  first  formed  and  this  is  then  decomposed  by 
the  water : 

P.  tribromide.  Phosphorous  acid. 

PBr3     -f     3H2O  =  3HBr     +     P(OH)3. 

To  prepare  it  on  the  small  scale,  put  a  little  bromine  in  the 
bottom  of  a  t.t ,  then  a  layer  of  pieces  of  glass  on  which  a  few 
small  bits  of  phosphorus  are  laid,  cover  this  with  an  inch  or 
two  of  small  bits  of  glass,  moisten  with  a  drop  or  two  of  water, 
and  distil  with  a  gentle  heat,  receiving  in  a  t.t.  of  water.  A 
solution  of  hydrobromic  acid  is  obtained. 

PROPERTIES. — A.  heavy,  fuming  gas,  very  soluble  in 
water,  forming  a  solution  similar  to  that  of  hydrochloric 


BROMINE  AND  OXYGEN.  Ill 

acid.  It  is  a  monobasic  acid,  and  acts  on  bases  to  form 
salts  called  bromides. 

BROMIDES. — These  are  very  like  the  chlorides  in  ap- 
pearance and  properties.  Bromine  unites  directly  with 
nearly  all  metals.  The  bromides  can  all  be  decomposed 
by  chlorine.  The  potassium  salt  (KBr)  is  the  most  im- 
portant. It  is  much  used  in  medicine.  Argentic  bromide 
(AgBr)  is  used  in  photography.  Ammonic  bromide 
(NH4Br)  is  used  as  a  medicine.  Most  of  the  bromides 
are  soluble  in  water.  A  few  are  insoluble. 

Experiment  81. — To  a  few  drops  of  solution  of  potassic 
bromide  add  a  drop  of  argentic  nitrate.  Argentic  bromide  is 
precipitated : 

AgN03  +  KBr  =  KNO3  +  AgBr. 

Note  its  colour  and  try  the  action  of  ammonia  solution  and  of 
nitric  acid  on  it  (dividing  it  into  two  portions).  Make  the  same 
experiment  with  solutions  of  mercurous  nitrate  (Hga(NO3).j), 
and  plumbic  acetate  (Pb(C2H30.2)J),  instead  of  argentic  nitrate. 
Write  the  equations. 

Tests  for  Bromides. — Experiments  79  and  81.  In  the 
test  with  chlorine  care  must  be  taken  not  to  use  too  much,  as 
there  is  a  colourless  chloride  of  bromine.  Except  in  weak  solu- 
tions the  bromine  can  be  liberated  by  strong  sulphuric  acid. 
(Try  it. )  Free  bromine  colours  starch-paste  orange-yellow. 

105.  Bromine  and  Oxygen. — No  compounds  of 

bromine  and  oxygen  are  known,  but  two  oxygen  acids, 
hypobromous  (HBrO),  and  bromic  (HBrO3),  have  been 
prepared.  They  and  their  salts  (the  hypobromites  and 
bromates)  are  prepared  by  methods  similar  to  those  for 
the  hypochlorites  and  chlorates.  Potassic  hypobromite 
(KBrO)  is  used  in  estimating  urea.  (Write  equations 


112  IODINE. 

for  the  action  of  bromine  on  moist  mercuric  oxide,  and 
on  dilate  and  strong  solutions  of  potassic  hydroxide). 


IODINE  AND  ITS  COMPOUNDS. 

106.  Iodine- — (I'  =  126.5).  Never  found  free  in 
nature,  but  always  as  iodides,  in  small  relative  quanti- 
ties, but  widely  diffused.  Minute  traces  of  iodides  are 
present  in  sea- water.  Certain  plants  (kelp  and  sponges) 
extract  these  and  store  them  in  their  tissues.  Some 
sea  animals  do  the  same,  e.g.,  the  cod.  The  ashes  of 
kelp  contain  from  0.1  per  cent,  to  0.3  per  cent,  of 
iodine,  and  are  the  principal  source  of  this  element. 
The  weed  is  washed  up  on  the  coasts  of  Ireland  and 
Scotland  in  great  quantities.  It  is  gathered,  dried, 
burned  in  shallow  pits  at  a  temperature  not  high  enough 
to  volatilise  the  iodides  and  bromides,  and  then  lixivi- 
ated. The  solution  contains  sodic  carbonate  (Na2CO3), 
chloride  (Nad),  sulphate  (Na2SO4),  bromide  (NaBr), 
iodide  (Nal),  &c.  It  is  evaporated  to  crystallise  the 
carbonate,  chloride,  and  sulphate  ;  and  the  iodine  and 
bromine  are  prepared  as  described  in  Art.  103. 

PREPARATION. — Experiment  82. — Mix  well  about  equal 
quantities  of  potassic  iodide  (KI)  and  manganese  dioxide,  put  in 
a  t.t.,  add  a  little  strong  sulphuric  acid,  and  heat  very  slightly 
for  a  few  minutes.  Violet  vapour  of  iodine  appears  and  con- 
denses on  the  cooler  parts  of  the  tube,  as  a  steely  looking  solid. 
2KI  +  3H2S04  +  Mn02  =  MnS04  +  2KHSO4  +  I2  +  2HaO. 

PROPERTIES. —  Blackish  grey  solid,  somewhat  metallic 
in  appearance  (compare  sulphur,  selenium,  and  tel- 
lurium), opaque  and  crystalline.  Its  odour  suggests  that 
of  rotting  sea- weed.  The  specific  weight  of  the  solid  is 
4.948  (calculate  that  of  the  gas).  It  melts  at  114°  C., 


IODINE.  113 

and  boils  at  200°  C.,  forming  a  splendid  deep-blue  vapour, 
which  is  purple  when  mixed  with  air.  It  volatilises 
slowly  at  ordinary  temperatures  and  must  therefore  be 
kept  in  well  stoppered  bottles.  Both  iodine  and  bromine 
(as  well  as  chlorine)  corrode  cork. 

Experiment  83. — Put  a  few  crystals  of  iodine  in  a  t.t.,  fill 
half-full  of  water,  shake  for  some  time,  and  note  that  a  little  of 
the  iodine  dissolves.  Add  a  small  quantity  of  potassic  iodide, 
and  shake  up.  The  whole  of  the  iodine  dissolves.  A  tri-iodide 
(Klj)  is  formed. 

Iodine  is  sparingly  soluble  in  water,  but  freely  so  in 
solutions  of  certain  salts,  particularly  potassic  iodide. 

Experiment  84. — To  solution  of  potassic  iodide  in  two  test- 
tubes  add  respectively  chlorine  water  and  bromine  water,  then 
a  few  drops  of  carbon  bisulphide  (CS2)  ;  shake  well  and  allow  to 
stand  a  moment.  (What  conclusions  do  you  draw  with  regard 
to  the  relative  chemism  of  the  three  halogens  ?) 

Experiment  85- — Try  the  solubility  of  iodine  in  a  mixture 
of  equal  volumes  of  alcohol  and  water,  and  in  chloroform,  using 
small  quantities  of  each  of  the  substances  mentioned. 

Iodine  has  an  irritating  action  on  the  skin  and  mucous 
membranes,  but  not  so  violent  as  chlorine  and  bromine. 
It  stains  the  skin  yellow,  but  the  stain  disappears  very 
soon,  unless  the  iodine  is  applied  often.  It  promotes 
absorption  and  thus  reduces  swellings,  and  is  much  used 
as  an  external  application. 

TINCTURE  OF  IODINE  is  a  solution  of  iodine  and  potassic 
iodide  in  rectified  spirit  (alcohol  and  water) — 1  pint  of 
spirit  to  |  oz.  iodine  and  ^  oz.  potassic  iodide. 

SOLUTION  OF  IODINE. — Twenty  grains  of  iodine,  thirty 
grains  of  potassic  iodide,  dissolved  in  1  oz.  of  distilled 
water. 

9 


114  HYDRIODIC    ACID. 

Tests. — 1.  With  chlorine  water  and  carbon  bisulphide  as  in 
Experiment  84. 

2.  Free  iodine  colours  starch  paste  deep  blue.  (Try  with 
iodine  water). 

COMPOUNDS  OF  IODINE. 

107.  Hydriodic  Acid,  HI.— 

PREPARATION. — In  the  same  way  as  hydrobromic  acid 
(Art.  104).  (Write  the  equation.)  Hydrogen  and  iodine 
do  not  unite  directly  under  ordinary  circumstances. 

It  can  also  be  prepared  in  solution  by  the  action  of 
iodine  on  sulphuretted  hydrogen  (H2S)  in  presence  of 
water. 

Experiment  86- — Put  a  few  small  pieces  of  ferrous  sulphide 
in  a  t.t.  fitted  with  the  gas-delivery  tube  of  Experiment  66,  add 
some  sulphuric  acid  diluted  with  about  five  times  its  volume  of 
water,  and  pass  the  evolved  gas  into  a  test-tube  of  water  in 
which  is  some  finely  divided  iodine.  The  iodine  disappears  and 
sulphur  is  precipitated  : 

H2S  +  Ia  =  2HI  +  S. 

Test  a  portion  of  the  solution  with  litmus,  add  a  little  iodine 
water,  shake  up,  and  set  aside  for  further  experiments. 

PROPERTIES. — A  heavy  gas  (specific  weight  4.416), 
colourless,  fuming  in  moist  air  owing  to  its  strong  attrac- 
tion for  water,  with  which  it  forms  minute  drops  of  solu- 
tion. It  dissolves  readily  in  water,  forming  a  strongly 
acid  solution  similar  to  those  of  hydrochloric  and  hydro- 
bromic acids.  The  gas  is  easily  decomposed  by  heat,  and 
the  solution  in  water  is  decomposed  by  the  oxygen  of 
the  air : 

•2HI  +  O  =  H20  +  L, 


IODIDES.  115 

(Compare  with  the  stability  of  hydrochloric  acid.)     It  is 
a  monobasic  acid  and  unites  with  bases  to  form  iodides, 

KOH  +  HI  =  KI  +  H2O. 

IODIDES. — Of  these  the  most  important  is  potassic 
iodide  (KI),  of  which  very  large  quantities  are  used  in 
medicine.  The  following  iodides  are  insoluble  in  water 
and  are  bright  in  colour :  Argentic  (Agl),  mercurous 
(Hg.,I2),  mercuric  (HgI2),  cuprous  (Cu.2I2S  and  plumbic 
(PbI2),  the  latter  very  sparingly  soluble. 

Experiment  87. — Put  a  little  of  the  solution  of  hydriodic 
acid  in  four  test  tubes,  and  add  a  drop  or  two  of  argentic  nitrate 
to  one,  of  mercurous  nitrate  to  a  second,  of  mercuric  chloride 
(HgCl2)  to  a  third,  and  of  plumbic  acetate  to  a  fourth.  Note 
the  color,  &c.,  of  the  precipitates  formed  : 

AgNO,  +  HI  =  Agl  +  HNOS. 

Write  equations  for  the  other  three.     Try  the  same  experiments 
with  a  dilute  solution  of  potassic  iodide  (KI). 

Tests. — -These  apply  to  both  free  acid  and  iodides. 

1 .  Add  to  the  solution  a  few  drops  of  argentic  nitrate  solution. 
A  yellow  precipitate  of  argentic  iodide  is  formed.     Divide  this 
into  two  portions .     Test  one  part  with  ammonia  solution  ;  it  is 
whitened,  but  not  dissolved.     Test  the  other  with  nitric  acid ; 
it  is  not  dissolved. 

2.  Add  some  starch  paste  and  a  few  drops  of  chlorine  water 
to  any  solution  containing  an  iodide.     A  blue  colour  appears. 
(Explain). 

To  test  for  chlorides,  bromides,  and  iodides  in  a  mixture, — Distil 
a  small  portion  of  the  carefully  dried  mixture  with  dry  powered 
potassic  bichromate  (K^C^O;),  and  strong  sulphuric  acid.  (The 
apparatus  must  be  dry.)  Receive  the  distillate  in  a  t.  t.  contain- 
ing a  little  water,  add  caustic  soda  to  it,  until  the  colour  of  the 


116  IODINE    AND    OXYGEN. 

bromine  and  iodine  disappears.  A  yellow  colour  remains,  that 
of  sodic  chromate.  This  proves  the  presence  of  a  chloride  in 
the  mixture. — Explanation  :  A  volatile  compound,  chromic 
oxychloride  (CrO^ClJ,  is  formed,  distils,  and  forms  chromic 
acid  with  the  water  : 

CrO-jCl.,  +  2H20  =  H2Cr04  +  2HC1. 

Dissolve  a  small  quantity  of  the  mixture  in  water  in  at.  t., 
add  a  few  drops  of  carbon  bisulphide,  and  then  chlorine  water 
drop  by  drop,  shaking  up  after  each  addition  and  observing  the 
colour  of  the  carbon  bisulphide.  The  violet  of  iodine  appears 
first,  but  disappears  on  the  addition  of  more  chlorine  owing  to 
the  formation  of  an  almost  colourless  compound  (IC1);  after- 
wards, on  the  addition  of  more  chlorine,  the  orange  colour  of 
bromine  appears. 

108.  Iodine  and  Chlorine. — Two  compounds  are 

known,  the  monochloride  (IC1),  and  the  trichloride 
(IC13),  formed  by  direct  union  of  the  elements.  (What, 
then,  is  the  valence  of  iodine  ?) 

109.  Iodine  and  Oxygen. — Only  one    compound 
of  these  two  elements  is  known,   viz.,  iodine  pentoxide 
(I2O5),  prepared  by  carefully  heating  the  corresponding 
acid,  iodic  acid  (HIO3)  : 

2HIO3  =  H2O  +  I205- 

When  heated  more  strongly  it  decomposes  into  its  ele- 
ments. It  is  a  deliquescent  white  solid,  and  unites  with 
water  to  form  iodic  acid. 

IODIC  ACID,  HIO3. — Can  be  prepared  by  oxidising 
iodine  with  strong  nitric  acid,  and  evaporating  the  solu- 
tion to  dryness.  It  is  a  white  crystalline  solid,  freely 
soluble  in  water,  forming  a  strongly  acid  solution.  It 


FLUORINE.  117 

is  a  powerful  oxidising  agent.     When  mixed  with  hydri- 
odic  acid,  it  oxidises  the  latter,  setting  free  iodine  : 

HIO3  +  5  HI  =  3L,  +  3H20. 

IODATES. — Besides  the  normal  iodates  (KIO3,  &c.), 
there  are  peculiar  acid  iodates,  e.g.,  KIO3.HIO3,  and 
KI03.2HIO3. 

Tests  — Add  chlorine  water  and  starch  paste  to  a  solution  of 
an  iodate.  No  blue  colour  appears.  (What  conclusion  do  you 
draw  with  regard  to  chemism  of  iodine  and  chlorine  in  oxygen 
salts?)  Now  add  a  little  sodic  sulphite  or  sulphurous  acid. 
Blue  colour  appears  :  KI03  +  3Na2SO,  =  KI  +  3Na2SO4. 
KI  +  Cl  =  KC1  +  I. 


FLUORINE  AND  ITS  COMPOUNDS. 

110.  Fluorine  (F'  =  19.1).— Is   not    known    free, 
but  only  through  its  compounds,  the  fluorides.     Many 
attempts  have  been  made  to  prepare  it,  but  with  no  suc- 
cess.    It    invariably    attacks    the   vessels    in    which  its 
preparation  is  attempted.     It  seems  to  have  even  stronger 
chemical  affinities  than  chlorine. — The  chief  compounds 
of  fluorine  occurring  in  nature  are  fluor  spar  or  calcic 
fluoride    (CaF2),     and    cryolite    (3NaF.AlF3). — Fluorine 

forms  no  compound  with  oxygen,  and  is  the  only  element 
of  which  this  is  true. 

111.  Hydrofluoric  Acid.— HF. 

PREPARATION. — By  distilling  fluor  spar  with  sulphuric 
acid  in  lead  or  platinum  vessels  and  receiving  in  water  : 

CaF2  +  H2SO4  =  2HF  +  CaSO4, 

Glass  or  porcelain  vessels  cannot  be  used,  as  they  are 
corroded  by  hydrofluoric  acid.      The  pure  acid  is  very 


118  HYDROFLUORIC    ACID. 

dangerous,    and    is    not    often    prepared.      The   aqueous 
solution  is  kept  in  leaden  or  gutta  percha  bottles. 

PROPERTIES. — A  colourless,  fuming,  acid  liquid,  boiling 
at  19.4C  C.  It  is  soluble  in  water,  forming  a  strongly 
acid  solution. 

Experiment  88- — Put  a  little  calcic  fluoride  in  a  small  leaden 
dish,  add  some  sulphuric  acid  and  cover  the  dish  with  a  piece 
of  glass  which  has  been  coated  with  paraffin  wax  through  which 
some  word  has  been  scratched  with  a  knife  or  a  pin.  Set  aside 
for  a  while,  then  warm  the  wax,  and  rub  it  off.  The  word  has 
been  etched  into  the  glass. 

Glass  contains  silica  (Si(X).  This  is  acted  on  by  the 
hydrofluoric  acid  with  the  formation  of  volatile  com- 
pounds : 

SiO,  +  4HF  =  SiF4  +  2H20. 

Hydrofluoric  acid  is  monobasic.      It  forms  salts  called 
fluorides,  analogous  to  chlorides,  bromides,  and  iodides. 

FLUORIDES. — These  are  mostly  colourless  salts,  soluble 
in  water  ;  but  calcic  fluoride  (CaF2),  baric  fluoride  (BaF2), 
and  strontic  fluoride  (SrF2)  are  insoluble.  The  fluorides 
have  a  tendency  to  unite  with  each  other  and  with  hydro- 
fluoric acid  forming  double  fluorides,  e.g.,  3NaF.  A1F3  and 
HF.KF. 

Test.— See  Expt.  88. 


QUESTIONS  AND  EXERCISES. 

1.  Compare  the  halogens  (1)  as  to  their  chemism  for  metals, 
and  (2)  as  to  their  chemism  for  oxygen. 


QUESTIONS    AND    EXERCISES.  119 

2.  What  weight  of  manganese  dioxide  (Mn02)  must  be  used  to 
prepare  chlorine  for  the"  con  version  of  10  Ibs.  of  potassic  hydrox- 
ide (KOH)  into  chlorate  and  chloride  ? 

3.  Calculate  the  specific  weight  of  chlorine  (air  =  1).     Calcu- 
late the  weight  of  1  litre  at  0°  and  760  mm. 

4.  Why  cannot  chlorine  be  collected  over  mercury  or  water  ? 

5.  What  weight  of  pure  hydrochloric  acid  is  there  in  1  litre  of 
the  solution  of  sp.  wt.  1.11  ? 

6.  Compare  the  bleaching  power  of  chlorine  (and  water),  and 
hypochlorous  acid  . 

C12  +  H20  =  2HC1  +  O. 
HC10  HC1  +  O. 

What  weight  of  chlorine  is  equivalent  to  100  grams  of  hypochlor- 
ous acid  ? 

7.  Compare  the  oxygen  compounds  of  chlorine  and  of  nitrogen. 

8.  What  weight  of  sodium  chloride  must  be  used  to  prepare 
chlorine  enough  to  set  free  the  bromine  from  100  oz.  of  potassic 
bromide  (KBr)? 

9.  What  weight  of  potassic  chlorate  (KC103)  contains   one 
equivalent  (in  grams)  of  oxygen  ? 

10.  Write  the  formulas  of  matjnesic,  ammonic,  ferric,  cobaltic, 
and  mercurous  chlorides  ;  of  mercuric,  argentic,  and  baric  iodides  ; 
and  of  ferric  bromide. 

11.  Fluorine  has  never  been   prepared.     How   then   has   its 
atomic  weight  been  determined  ? 

12.  Write  the  formulas  of  the  chlorine  acids  and  calculate  the 
percentage  of  oxygen  in  each.     Apply  the  Law  of  Multiple  Pro- 
portions. 

13.  What  weight  of  hydrochloric  acid  is  equivalent  to   100 
grams  sodic  hydroxide  (NaOH)  ?     To  100  grams  potassic  hydrox- 
ide (KOH)  ?     To  100  grams  calcic  hydroxide  (Ca(OH)2)  ? 


120  THE    SULPHUR    GROUP. 


CHAPTER    X. 


SULPHUR,  SELENIUM,  AND  TELLURIUM. 

112.  The    Sulphur  Group. — The  three    elements 
mentioned  here  form  a  group   closely  related   in    their 
chemical    and    physical    properties.     Each    unites    with 
hydrogen  to  form  a  gas  analagous  to   water  in  formula 
(H2S,   H2Se,  H2Te).     They  unite  with  oxygen  in  two 
proportions  (SO2,  SO3 ;  SeO2,  —  ;  Te02,  TeO3) ;  and  each 
has  two  acids  corresponding  to  the  oxides.     As  in  the 
case  of  the  halogens,  there  are  gradations  in  the  proper- 
ties in  passing    from   one   element  to  the  next  of  this 
group  of  elements.     Thus,   while    sulphur  is  a  distinct 
non-metal,  selenium  is   slightly  metallic,  and  tellurium 
more  so,  in  its  appearance.     The  attraction  for  oxygen 
is  stronger  in   sulphur,  weaker  in   selenium  and  tellu- 
rium. 

SULPHUR  AND  ITS  COMPOUNDS. 

113.  Sulphur  (S»-iT-vi=  31.98). 
OCCURRENCE. — Found  uncombined,  in  great  masses  in 

volcanic  districts,  particularly  in  Italy,  Iceland,  and 
Sicily.  The  compounds  of  sulphur  found  in  nature  are 
very  numerous  and  abundant.  There  are  two  principal 
classes  :  (1)  sulphides,  including  a  great  many  ores  of 
metals,  e.g.,  galena  (PbS),  zinc  blende  (ZnS),  cinnabar 
(HgS),  and  iron  and  copper  pyrites;  and  (2)  sulphates, 
e.g.,  gypsum  (CaSO4.2H2O),  heavy  spar  (BaSO4),  Epsom 


SULPHUR.  121 

salts  (MgSCVTHaO),  and  green  vitriol  (FeSO4.7H2O). 
Sulphur  also  forms  an  essential  constituent  of  animal  and 
vegetable  bodies. 

PREPARATION. — (1)  Crude  sulphur  is  heated  in  shallow 
pits,  and  the  melted  sulphur  is  run  off  from  the  earthy 
impurities.  It  is  further  purified  by  distillation,  the 
vapour  being  led  into  cool  chambers  where  it  solidi- 
fies and  falls  down  as  "  flowers  of  sulphur."  When  the 
walls  of  the  chamber  become  hot  the  sulphur  remains 
liquid  and  is  run  off  into  moulds  roll  brimstone). 

(2)  Iron  pyrites  (FeS2)  when   distilled  gives  off  one- 
third  of  the  sulphur : 

3FeS2  =  Fe3S4  +  S2. 

This  can  be  condensed,  <kc.,  as  above.  (Try  with 
powdered  pyrites  in  a  narrow  glass  tube  sealed  at  one 
end.) 

(3)  Sulphur  is  also  obtained  as  a   by  product  in  the 
process  of  purifying  coal  gas. 

(4)  The  waste  liquor  from  the  alkali  manufacture  con- 
tains sulphur,  and  has  been  used  as  a  source. 

PROPERTIES. — Sulphur  has  three  allotropic  modifica- 
tions, two  crystalline,  and  one  amorphous  (formless). 
These  differ  in  specific  weight  and  other  physical  pro- 
perties. 

Experiment  89. — Heat  carefully  until  completely  fused 
some  pieces  of  roll  sulphur  in  a  porcelain  dish.  Then  take  the 
burner  away,  and  hold  a  cork  in  the  liquid  until  a  crust  forms 
over  the  surface.  Break  two  small  holes  in  this  crust  on  op- 
posite sides,  and  pour  out  the  liquid  part  into  another  dish. 


122  SULPHUR. 

(Why  two  holes?).  Then  cut  out  the  crust  with  a  knife,  and, 
lifting  it  carefully  by  the  cork,  observe  the  beautiful  crystals  of 
sulphur.  Lay  aside  for  a  day,  and  observe  that  the  translucent 
amber  solid  becomes  opaque  and  yellow. 

Sulphur  melts  at  114.5°  C.  to  a  clear  amber  liquid; 
if  allowed  to  cool  at  the  ordinary  temperature  of  the  air 
it  crystallises  in  needle-shaped  crystals  (monoclinic). 
These  are  unstable  and  gradually  change  to  another 
crystalline  form  (rhombic  prisms).  Sulphur  is  therefore 
dimorphous.  Each  monoclinic  crystal  becomes  trans- 
formed into  a  great  number  of  minute  rhombic  crystals, 
and  hence  the  opacity.  This  change  goes  on  gradually, 
and  its  progress  can  be  watched. 

Experiment  90.— Melt  some  sulphur  in  a  dry  t.  t.  (using  the 
holder),  and  continue  heating  it.  The  mobile  amber  liquid  be- 
comes (at  200°  C.)  dark  and  viscid  (will  not  pour).  Continue 
the  heat.  It  becomes  mobile  again,  and  at  length  begins  to  rise 
as  a  heavy  vapour.  Now  pour  it  in  a  thin  stream  into  a  beaker 
of  cold  water.  Plastic  sulphur  is  obtained.  Keep  and  observe 
it. 

Plastic  sulphur  is  amorphous — has  no  particular  form, 
as  crystals  have.  It  can  be  drawn  out  into  threads,  but 
slowly  hardens  and  changes  into  rhombic  sulphur.  This 
latter  is  the  permanent  form  of  sulphur.  Large  crystals 
can  be  obtained  by  evaporating  solutions  of  sulphur  in 
carbon  bisulphide. —  Sulphur  boils  at  450°  C.,  forming  a 
dark  red  vapour.  The  specific  weight  of  this  vapour  is 
peculiar.  If  determined  at  a  temperature  near  its  boil- 
ing point,  it  is  96  (hydrogen  =  1).  The  molecule  must 
then  weigh  192  times  that  of  hydrogen,  and  must  con- 
tain 6  atoms,  since  the  atomic  weight  of  sulphur  is  32. 
The  formula  for  sulphur  at  temperatures  not  much  above 
450°  C.  is  then  S6.  The  specific  weight  at  10-10°  C. 


MILK    OF    SULPHUR.  123 

(boiling  point  of  zinc)  is,  however,  only  32.     (How  many 
atoms  in  the  molecule  at  this  temperature  ?) 

Experiment  91-  — Try  the  solubility  of  flowers  of  sulphur  in 
water,  alcohol,  carbon  bisulphide,  and  paraffin  oil,  by  shaking 
in  test  tubes  with  a  very  large  proportion  of  the  solvent  Use  a 
little  heat  also  before  deciding  in  the  negative. 

Sulphur  burns  in  air  with  a  pale  blue  flame  forming 
sulphur  dioxide  (SO2).  It  oxidises  very  slowly  in  moist 
air  to  sulphuric  acid  H2SO4). 

MILK  OF  SULPHUR  (Lac  sidphuris}. — This  is  sulphur 
in  a  very  finely  divided  state  suspended  in  water.  When 
dried  it  is  called  precipitated  sulphur. 

Experiment  92. — Boil  for  some  time  in  a  t.  t.  a  little  of  the 
flowers  of  sulphur  (5  parts)  with  slaked  lime  (10  parts)  and 
water  (20  parts),  filter,  and  divide  the  filtrate  (the  clear  liquid 
which  runs  through)  into  two  parts.  To  one  part  add  dilute 
hydrochloric  acid  until  the  liquid  is  acid.  Sulphur  is  precipi- 
tated as  a  white  powder.  Allow  it  to  settle,  collect  it  on  a 
filter,  wash  (by  pouring  hot  water  on  the  filter),  and  allow  it  to 
dry.  To  the  second  portion  add  dilute  sulphuric  acid,  until  the 
liquid  is  faintly  acid,  and  proceed  as  before.  Compare  the  two 
specimens  in  appearance.  Burn  a  small  portion  of  each  on  a 
piece  of  mica.  That  prepared  with  sulphuric  acid  leaves  a  white 
residue  ;  the  other  leaves  little  or  none.  Test  a  few  specimens 
obtained  from  druggists. 

When  sulphur  and  lime  water  are  heated  together  a 
persulphide  and  thiosulphate  of  calcium,  are  formed  : 
3Ca(OH)2  +128  =  2CaS5  +  CaS,O,  +  311,0. 

The  action  of  the  acids  is  shown  in  the  following 
equations  : 

2CaS6  +  CaS.,O3  +'6HC1  --=  3CaCl2  +  12S  +  3H20. 
2CaS5  +  CaS2O3  +  3H2SO,  =  3CaS04  -f  12S  +  3H20. 


124  SULPHUR   DIOXIDE. 

Prepared  with  sulphuric  acid,  precipitated  sulphur  is 
very  impure.  Calcic  chloride  (CaCl2)  is  freely  soluble  in 
water  and  is  washed  away  from  the  sulphur ;  while 
calcic  sulphate  (CaSO4)  being  sparingly  soluble,  remains 
mixed  with  the  sulphur.  The  impure  precipitated  sul- 
phur has  a  glistening  look,  and  feels  slightly  gritty. 

SULPHUR  AND  OXYGEN. 

114.  Oxides  of  Sulphur. — Four  oxides  of  sulphur 
are  known.     Their  composition  is  represented  as  follows: 
SO2,  SO3,  S2O3,  and  S2O7.     Of  these  only  the  first  two 
are  of  importance. 

115.  Sulphur  Dioxide. — S02. — Also  called  sulphur- 
ous anhydride.      As  has  been   seen,  sulphur   burns    in 
oxygen  forming  an  oxide  which,  dissolving  in  water,  gives 
an  acid  substance.     This  oxide  is  sulphur  dioxide. 

PREPARATION. — Sulphur  dioxide  can  be  prepared  (1) 
by  burning  sulphur  in  air  (Is  it  obtained  pure  ]) ;  (2)  by 
burning  iron  pyrites  in  air  :  4FeS2  +  11O2  =  2Fe2O3  + 
8SO2 ;  (3)  by  heating  strong  sulphuric  acid  with  char- 
coal:  2H2SO4  +  C  -  COa  +  2SO2  +  2H20;  and  (4) 
by  heating  strong  sulphuric  acid  with  certain  metals, 
e.g.,  copper,  mercury,  silver. 

Experiment  93. — Heat  a  little  iron  pyrites  on  mica.  Notice 
that  it  burns  with  a  blue  flame  and  the  smell  of  "  burning  sul- 
phur." (What  action  has  taken  place  ?  What  substance  is  left  ?) 

Experiment  94. — Put  a  considerable  quantity  of  scraps  of 
copper  in  a  small  flask,  or  a  large  t.  t.  and  pour  in  enough  con- 
centrated sulphuric  acid  to  cover  the  metal.  Fit  with  the  deliv- 
ery tube  of  Expt.  66,  and  heat  carefully  until  a  gas  begins  to  bubble 


SULPHUR    DIOXIDE.  125 

off.  Collect  three  jars  of  it  by  displacement  of  air  (downward  or 
upward?  Calculate  the  sp.  wt.)  Then,  still  heating  very  little, 
allow  the  gas  to  bubble  through  a  bottle  of  distilled  water  cooled 
to  0°  C.  by  snow  or  ice.  After  some  time  a  white  crystalline 
solid  (H2SO,.  14H20)  is  formed. 

(Try  preparation  on  a  small  scale  with  charcoal.) 

PROPERTIES. — An  invisible  gas  of  suffocating  smell 
("burning  sulphur"),  about  2^  times  as  heavy  as  air; 
soluble  in  water  (at  0°  C.,  80  vols.  in  1  ;  at  20°,  40  in  1); 
condenses  to  a  liquid  at  —  8°  and  with  ordinary  atmos- 
pheric pressure. 

Experiment  95. — Put  a  burning  taper  or  match  in  a  jar  of 
the  gas,  and  observe. 

Experiment  96- — Hang  a  strip  of  moist  turkey-red  cotton  in 
another  jar,  close  it,  and  allow  it  to  stand.  The  colour  is 
bleached. 

Sulphur  dioxide  bleaches  only  in  presence  of  water ; 
and  its  bleaching  action  is  the  reverse  of  that  of  chlorine. 
It  is  due  to  the  action  of  hydrogen  from  water : 

SO2  +  2H2O  =  H2SO4  +  H2. 

The  hydrogen  unites  with  the  colouring  matter  forming  a 
colourless  compound.  In  many  cases  the  colour  may  be 
gradually  restored  by  the  oxidizing  action  of  the  air. 
The  action  of  sulphur  dioxide  is  not  so  destructive  as 
that  of  chlorine,  and  it  is  therefore  used  for  bleaching 
delicate  materials,  such  as  silk,  wool,  &c.  Sulphur  diox- 
ide is  a  reducing  agent,  but  most  of  the  reducing  actions 
take  place  in  presence  of  water,  which  is  decomposed 
under  the  double  influence  of  sulphur  dioxide  attracting 
the  oxygen,  and  some  other  substance  attracting  the 
hydrogen. 


126  SULPHUR    TRIOXIDE. 

Experiment  97-  —  Test  the  action  on  litmus  of  the  solution 
prepared  in  Experiment  94.  It  contains  an  acid,  sulphurous  add 
(H2SOS).  Note  the  smell  and  taste  of  the  solution. 

Sulphur  dioxide  forms  a  weak  compound  with  water, 
which  cannot  be  obtained  free  from  water,  except  at  low 
tempei-atures. 

Experiment  98. — Mix  some  sulphurous  acid  with  iodine 
solution  until  the  colour  is  just  discharged.  Note  that  the  odour 
of  both  is  now  absent.  Test  with  litmus.  The  solution  is  still 
acid  :  SO.,  +  Ia  -f  2H,O  =  H2S04  +  2HI.  Do  similarly  with 
chlorine  water. 

Experiment  99. — Colour  some  sulphurous  acid  in  a  porcelain 
dish  with  litmus,  add  solution  of  caustic  soda  until  the  colour 
becomes  blue,  and  evaporate  to  dryness  on  the  water  bath.  A 
white  crystalline  salt,  todic  sulphite,  remains  : 

H,SO,  +  2NaOH  =  Na,SOs  +  2H30. 
Sulphur  dioxide  is  a  good  disinfectant  and  antiseptic. 

Tests. — 1.  Wet  a  piece  of  filter  paper  with  a  mixture  of 
ferric  chloride,  (Fe2Cl6)  and  potassic  ferricyanide  (K3FeC6N6)  in 
solution,  and  hold  it  in  an  atmosphere  containing  sulphur 
dioxide.  It  turns  blue.  The  ferric  yanide  is  reduced  to  fer- 
rous cyanide,  and  this  forms  Prussian  blue  with  ferric  chloride. 

2.  Mix  a  few  drops  of  solution  of  sulphur  dioxide  with  starch 
paste  and  solution  of  iodic  acid  (HI03).  Blue  iodide  of  starch 
is  formed  :  2HIO,  +  5SO2  +  4H2O  =  5H,SO4  +  I2.  If  more 
sulphur  dioxide  be  added  the  colour  disappears.  (Explain.) 

116.  Sulphur  Trioxide  SO3.  —  Also  called  sul- 
phuric anhydride,  or  anhydrous  sulphuric  acid. 

PREPARATION. — (1)  By  passing  a  mixture  of  sulphur 
dioxide  and  oxygen  over  heated  spongy  platinum,  or 
platinised  asbestos  : 

SO,        O        SO. 


OXYGEN  ACIDS  OF  SULPHUR. 


127 


(In  what  proportions  by  volume  do  the  gases  combine1?) 
(2)  By  distilling  pure  sulphuric  acid  with  phosphorus 
pentoxide  (P2O5)  : 

H2SO4  +  PA  =  2HPO3  +  S0:1. 

PROPERTIES.  —  A  transparent  white  crystalline  solid, 
melting  at  16°  C.,  boiling  at  46°  C.  It  readily  combines 
with  water,  forming  sulphuric  acid  (H2S04).  It  hisses 
like  red  hot  iron  when  thrown  into  water,  and  when 
exposed  to  the  air  it  deliquesces.  Its  relation  to  sul- 
phuric acid  is  seen  thus  : 

H2O  +  S03  ==  H2S04. 

117.  Oxygen  Acids  Of  Sulphur.  —  These  are 
very  numerous,  but  only  two  are  of  any  importance  as 
acids,  viz.,  sulphurous  and  sulphuric.  The  others  will 
be  studied  mostly  through  their  salts. 


Dithionic  acid. ...  H2S2O( 

ZVithionic  acid. . .  H2S3O( 

2Vrathionic  acid.  H2S4O6 

Pentathionic  acid.  HoSrO, 


H2S02 
Sulphurous  acid  ....   H2SO3 

Sulphuric  acid H2SO4 

T/wosulphuric  acid . .   H  2  S  2  O  3 


118.  Hyposulphurous  Acid,  H2SO2. — Also  called 

hydrosulphurous  acid.  Prepared  by  the  action  of  zinc 
on  sulphurous  acid  solution  in  closed  vessels  : 

H2SO3  -f  Ha  =  H2SO2  +  H,0. 

(What  is  the  source  of  the  hydrogen  represented  in  the 
equation  1)  It  is  a  strong  reducing  agent,  and  readily 
absorbs  oxygen  from  the  air.  (What  is  formed  1)  It  is 
a  monobasic  acid,  and  forms  salts,  the  hyposulphites,  in 
which  half  the  hydrogen  of  the  acid  still  remains.  The 


128  SULPHUROUS    ACID. 

second  atom  of  hydrogen  in  the  molecule  cannot  be  re 
placed  by  metal. 

SODIC  HYPOSULPHITE  (NaHSO2),  prepared  by  dissolv- 
ing zinc  in  a  solution  of  sodic  hydric  sulphite  (NaHSO3), 
is  used  to  reduce  indigo  in  calico  printing. 

119.  Sulphurous  Acid,  H2SO3. — Has  been  already 
mentioned  (Art.  115).  The  pure  acid  exists  as  a  solid 
at  low  temperatures,  and  the  solution  of  tlie  dioxide  in 
water  has  acid  properties. 

PREPARATION. — By  heating  in  a  glass  flask  1  part  of 
charcoal,  broken  in  small  pieces,  with  7|  parts  concen- 
trated sulphuric  acid.  The  sulphur  dioxide  is  washed  by 
passing  through  a  small  quantity  of  water,  and  is  then 
dissolved  in  cold  water  to  saturation.  The  solution 
must  be  kept  well  stoppered.  It  contains  about  12  °/0 
of  sulphurous  acid,  and  its  specific  weight  is  1.04.  For 
properties  see  Art.  115. 

SULPHITES. — Sulphurous  acid  neutralises  strong  bases 
(Exp't  99)  and  forms  salts,  the  sulphites.  There  are  two 
classes  of  sulphites  :  ( 1 )  normal  sulphites,  containing  no 
hydrogen,  e.g.,  sodic  sulphite  (Na2S03)  ;  and  (2)  acid 
sulphites,  containing  hydrogen,  e.g.,  sodic  hydric  sulphite 
(NaHSO3*.  Sulphurous  acid  is  dibasic;  it  unites  with 
bases  in  two  proportions.  This  subject  will  be  more  fully 
discussed  in  treating  of  sulphuric  acid.— The  sulphites  are 
generally  colourless  salts,  not  very  stable,  readily  absorb- 
ing oxygen  from  the  air  to  form  sulphates.  They  are 
strong  reducing  agents,  and  are  used  as  such  in  the  arts. 
In  medicine,  sodic  sulphites  are  administered  to  introduce 
sulphurous  acid  into  the  stomach.  (What  is  the  action  ?) 


SULPHURIC    ACID.  129 

Tests. — See  Art.  115.  Sulphites  can  generally  be  known  by 
giving  the  smell  of  sulphur  dioxide  when  treated  with  sulphuric 
acid. 

Experiment  100. — To  a  solution  of  sodic  sulphite  in  a  t.  t. 
add  a  few  scraps  of  zinc  and  some  hydrochloric  acid.  Note  the 
smell  of  the  gas  evolved.  Dip  a  strip  of  filter  paper  in  plumbic 
acetate  solution  and  hold  it  over  the  mouth  of  the  t.  t.  It 
blackens.  By  the  reducing  action  of  the  nascent  hydrogen, 
hydric  sulphide  (HaS)  is  formed:  S0a  +  3H2  =  HaS  +2H.,0. 
This  acts  on  the  lead  salt  to  form  plumbic  tulphide  (PbS),  a  black 
substance. 

1 20.  Sulphuric  Acid,  H2SO4.  —  Known  commer- 
cially as  "  oil  of  vitriol,"  or  simply  "  vitriol."  It  was 
formerly  prepared  by  distilling  green  vitriol,  and  "  fum- 
ing sulphuric  acid  "  is  still  prepared  in  this  way.  It  is 
the  most  important  of  all  acids,  and  is  used  in  the  pre- 
paration of  most  other  acids.  About  850,000  tons  are 
manufactured  annually  in  Great  Britain. 

MANUFACTURE. — To  form  a  molecule  of  sulphuric  acid 
it  is  necessary  only  to  unite  a  molecule  of  sulphur  dioxide 
with  an  atom  of  oxygen  and  a  molecule  of  water  : 

S02  +  0  +  H2O    =  H2SO4. 

But  this  union  does  not  take  place  directly  except  under 
the  influence  of  platinum  or  some  other  "  contact  action  " 
substance.  In  practice  it  is  brought  about  by  another 
chemical  action.  Nitrogen  dioxide  (NO)  readily  combines 
bines  with  the  oxygen  of  the  air  and  thus  forms  the  tetrox- 
ide  (NO2).  This  as  readily  gives  up  half  its  oxygen  to  sul- 
phur dioxide  and  water,  and  is  ready  to  unite  again  with 
oxygen  from  the  air.  Thus,  nitrogen  dioxide  acts  as  a 
"  carrier  of  oxygen  "  from  the  air  to  sulphur  dioxide  and 
10 


130  SULPHURIC   ACID. 

water.  (Of  course,  this  carrying  goes  on  between  mole- 
cules, and  not  between  vessels  or  rooms  !  )  Thus  : 

NO  +  O  -  NO2. 

N02  +  H2O  +  SO2  -  H2SO4  +  NO., 
<fec.,     &c. 

Sulphur  dioxide  (from  iron  pyrites  by  burning),  water, 
air,  and  nitric  acid  (from  sodic  nitrate  and  sulphuric 
acid),  are  brought  together  in  large  leaden  chambers. 
The  nitric  acid  immediately  suffers  reduction  to  nitrogen 
dioxide,  and  the  actions  described  above  go  on  continu- 
ously. Theoretically,  one  molecule  of  nitrogen  dioxide 
would  in  time  convert  any  quantity  of  the  materials  into 
sulphuric  acid.  In  practice  there  is  waste,  so  that  a  fresh 
supply  of  nitric  acid  is  needed  from  time  to  time.  (What 
is  left  in  the  chambers  after  the  sulphuric  acid  is  formed  1) 
The  sulphuric  acid,  as  fast  as  it  forms,  falls  in  a  fine 
shower  to  the  floor  of  the  chamber  into  a  layer  of  water. 
It  is  drawn  off  and  concentrated  in  leaden  pans  to  78% 
(sp.  wt.  1.71),  at  which  strength  it  is  sold  as  "  brown  oil  of 
vitriol  "  (B.O.V.)  For  special  purposes  it  is  further  con- 
centrated in  glass  or  platinum  vessels  to  96%  (sp.  wt. 
1.84). 

PROPERTIES. — The  pure  acid  free  from  water  is  a 
crystalline  white  solid  below  10°  C.,  a  colourless  oily 
liquid  above  this.  It  has  such  a  strong  attraction  for 
water  that  it  is  impossible  to  keep  it  pure  unless  it  is 
sealed  up.  Thus  in  practice  we  have  to  deal  with  an  acid 
containing  more  or  less  water.  The  sp.  wt.  of  the  pure 
liquid  acid  is  1.85.  With  many  of  the  properties  of  sul- 
phuric acid  the  student  has  already  become  acquainted. 
It  is  the  strongest  of  acids,  and  its  powerful  attraction  for 


SULPUHRIC    ACID.  131 

water  enables  it  to  decompose  many  substances  which  con- 
tain hydrogen  and  oxygen  in  the  proportions  to  form  water. 

Experiment  101. — Pour  some  sulphuric  acid  on  a  little  sugar 
in  a  porcelain  dish.  The  sugar  blackens  after  a  few  minutes. 
Put  a  drop  of  acid  on  a  piece  of  paper  ;  on  a  piece  of  cotton  cloth. 

Paper  and  sugar  are  composed  of  carbon  united  with 
hydrogen  and  oxygen  in  the  proportions  to  form  water. 
Sulphuric  acid  sets  the  carbon  free  by  taking  the  hydro- 
gen and  oxygen  as  water.  The  dilute  acid  turns  dyed 
cloth  reddish.  If  ammonia  be  applied  the  colour  is 
restored.* 

Experiment  102- — Heat  a  drop  of  sulphuric  acid  on  mica. 
Note  that  it  disappears  entirely,  with  "  fuming." 

Sulphuric  acid  boils  at  338°  C.,  but  undergoes  par- 
tial dissociation  (temporary  decomposition)  into  sulphur 
trioxide  and  water,  recombination  taking  place  in  the 
air.  Hence,  the  appearance  of  fumes. — With  the  excep- 
tion of  gold,  platinum,  and  a  few  others,  sulphuric  acid 
dissolves  all  the  metals,  forming  sulphates.  When  the 
concentrated  acid  is  used  it  acts  as  an  oxidising  agent 
(Exp't  94),  and  sulphur  dioxide  is  a  by-product.  Some 
metals,  as  copper,  silver,  lead,  and  mercury,  are  dissolved 
only  by  the  strong  hot  acid  : 

2Ag  +  2H2SO4  =  Ag2SO4  -f  SO2  -f  2H2O. 

Others,  as  zinc,  iron,  and  aluminum,  are  dissolved  by 
the  strong  acid  with  the  evolution  of  sulphur  dioxide, 
but  can  also  be  dissolved  by  the  dilute  acid,  and  in  this 
case  hydrogen  is  the  gas  evolved  : 

Fe  +  H2SO4  =  FeSO4  -f  H2. 

*  Sulphuric  acid  is  a  corrosive  poison,  charring  the  surfaces  with  which  it 
comes  in  contact.  It  has  often,  owing  to  its  oily  appearance,  been  mistaken 
for  castor  oil.  The  antidotes  are  dilute  alkalis. 


132  SULPHATES. 

Impurities  in  Sulphuric  Acid. 

PLUMBIC  SULPHATE  (PbSO4)  is  very  commonly  present 
in  the  commercial  acid.  It  comes  from  the  leaden  evapo- 
rating pans,  which  are  attacked  when  the  concentration 
is  pushed  too  far.  This  salt  is  soluble  in  the  concen- 
trated acid,  but  is  precipitated  on  diluting  with  water. 

Experiment  103.— Pour  a  little  commercial  oil  of  vitriol  into 
about  four  times  its  volume  of  distilled  water.  A  white  preci- 
pitate indicates  the  presence  of  plumbic  sulphate. 

ARSENIC,  derived  from  the  pyrites,  is  often  present; 
and,  as  it  forms  many  volatile  compounds,  it  is  a  dpaiger- 
ous  impurity.  For  example,  when  arsenical  sulphuric 
acid  is  used  in  the  preparation  of  hydrochloric  acid,  the 
arsenic  distils  over  as  trichloride  (AsCl3),  and  thus 
renders  the  hydrochloric  acid  highly  poisonous.  It  can  be 
tested  for  by  diluting  the  sulphuric  acid  with  water  and 
adding  hydric  sulphide  (H2S),  when  the  yellow  arsenic 
trisulphide  (As2S3)  is  thrown  down. — The  colour  of 
B.  O.  V.  is  due  to  organic  matter. — Sulphur  dioxide,  and 
oxides  of  nitrogen  occasionally  occur  as  impurities. 

SULPHATES. — Sulphuric  acid  is  a  dibasic  acid.  Its 
molecule  contains  2  atoms  of  hydrogen,  and  both  of  these, 
or  one  of  them  only,  can  be  replaced  by  metal.  Thus, 
there  are  two  series  of  sulphates,  normal  and  acid ;  e.  g., 
K2SO4,  normal  potassic  sulphate,  and  KHS04,  acid 
potassic  sulphate.  The  acid  sulphates  are  often  called 
bisulphates  (Art.  121). — The  following  normal  sulphates 
are  insoluble  or  very  sparingly  soluble  in  water:  Plum- 
bic (PbSOJ,  baric  (BaSO4),  and  strontic  (SrSO4).  Calcic 
(CaSO4),  argentic  (Ag2SO4),  and  mercurous  (Hg._,SO4)  are 


NORMAL   AND   ACID   SALTS.  133 

sparingly  soluble.     The  rest  of  the  sulphates  are  soluble, 
mostly  easily  crystallisable  salts. 

Tests. — Experiment  104. — Add  a  few  drops  of  solution  of 
baric  chloride  (BaCl2)  or  nitrate  to  dilute  sulphuric  acid.  A 
white  precipitate  (BaSO4)  is  thrown  down.  (Write  the  equa- 
tion). This  is  insoluble  in  nitric  acid.  Try  the  same  with  solu- 
tion of  magnesic  sulphate  (MgS04). 

121.  Normal  and  Acid  Salts. 

Experiment  105. — Put  one  pipetteful  of  dilute  sulphuric  acid 
in  a  porcelain  basin,  colour  with  litmus,  carefully  add  sodic 
hydroxide  solution  until  neutral,  and  evaporate  on  the  water 
bath.  Repeat  this  operation,  but  after  reaching  the  neutral 
point  add  a  second  pipetteful  of  the  acid. 

By  these  experiments  two  different  salts  are  obtained. 
Their  crystalline  form  and  other  physical  properties 
differ ;  and  that  obtained  in  the  second  operation  is  acid 
in  taste,  &c.  In  fact,  it  is  still  half  acid,  as  can  be  seen 
by  referring  to  the  equations  : 

2NaOH  -f  H2S04  =  Na.2SO4  -f  2H20. 
NaOH  -f  H2SO4  =  NaHSO,  -f  H,0. 

Sulphuric  acid  acts  on  caustic  soda  in  two  proportions, 
98  to  80,  and  98  to  40.  With  the  larger  proportion  of 
the  base  a  neutral  salt  is  formed,  and  the  hydrogen  of 
the  acid  is  completely  replaced  by  the  metal.  With  the 
smaller  proportion  an  acid  salt  is  formed,  and  only  half 
the  hydrogen  of  the  acid  is  replaced  by  metal.  Now,  the 
term  "  neutral  salt "  applies  as  long  as  both  acid  and 
base  are  well  marked,  but,  if  either  acid  or  base  is  weak, 
the  properties  of  the  other  predominate,  even  in  a  salt 
formed  with  the  largest  proportion  of  the  weak  acid  or 
base ;  so  that  a  so-called  neutral  salt  may  turn  blue 


134  FUMING    SULPHURIC    ACID. 

litmus  to  red,  or  red  to  blue.  It  has  been  agreed  to 
restrict  the  use  of  the  term  "  neutral  "  to  the  action 
of  substances  on  litmus ;  and  to  use  the  word  normal 
instead. 

Definitions. — A  normal  salt  is  one  in  which  all  the  replace- 
able hydrogen  is  replaced  by  metal. 

An  acid  salt  is  one  containing  replaceable  hydrogen  (i.e.,  re- 
placeable by  metal). 

(Can  monobasic  acids  form  acid  salts  ?) 

122.  Fuming  Sulphuric  Acid,  or  Nordhausen  sul- 
phuric acid. — This  is  of  somewhat  vai-iable  composition. 
It  is  sulphuric  acid  combined  with  the  trioxide ;  gener- 
ally H2S04   +   SO3.      It  is  prepared   from  green  vitriol 
(FeSO4.7H2O)  by  roasting  in  air  to  drive  off  the  water 
of  crystallisation  and  oxidise  the  salt : 

2(FeSO4.7H20)  -f  O  =  Fe2O3.(SO3>2  +  HH,O, 

and  then  distilling.  Enough  moisture  is  obtained  from 
the  air  and  the  apparatus  to  form  a  certain  proportion  of 
sulphuric  acid  : 

Fe2O3.(S03)2  -f  H2O  =  Fe2O3  -f  H2SO4.SO3. 

It  is  used  to  dissolve  indigo,  and  in  the  manufacture  of 
alizarine. 

1 23.  Thiosulphuric  Acid,  H2S2O3.— Formerly  called 
hyposulphurous  acid.     The  name  now  in  use  means  sul- 
pho-sulphuric  acid,  and  refers   to   the    composition,  the 
acid  containing  an  atom  of  sulphur  in  place  of  one  of  the 
four  atoms  of  oxygen.     The  acid  itself  cannot  be  isolated, 
as  it  decomposes  quickly  on  being  set  free  from  its  salts  : 

H2S,O3  -  H,O  +  S  +  SO,. 


THIOSULPHATES.  135 

Experiment  106-  — Add  some  dilute  sulphuric  acid  to  a  solu- 
tion of  sodic  thiosulphate  (Na2S203),  and  watch  carefully.  After 
a  few  seconds  the  solution  becomes  milky.  .Now  note  the  odour- 
(Write  the  equation.) 

THIOSULPHATES. — Thiosulphuric  acid  is  dibasic,  both 
atoms  of  hydrogen  being  replaceable  by  metal.  The  most 
important  salt  is  sodic  thiosulphate,  prepared  by  combin- 
ing sodic  sulphite  with  sulphur  : 

Na2SO3  +  S  =  Na2S2Os, 

or,  in  practice,  by  dissolving  sulphur  in  sodic  hydroxide 
solution  and  saturating  with  sulphur  dioxide. 

Experiment  107- — Boil  a  solution  of  sodic  sulphite  for  some 
time  in  a  porcelain  basin  with  flowers  of  sulphur,  filter,  evaporate 
the  filtrate  to  small  bulk,  and  set  aside  to  crystallise.  Fine 
colourless  crystals  (Na2S203.5H20)  are  formed.  Collect  and 
dry  them  on  filter  paper. 

SODIC  THIOSULPHATE  readily  absorbs  oxygen  from  the 
air  and  becomes  oxidised  to  sulphate.  Its  action  on 
iodine  renders  it  very  useful  in  analytical  chemistry. 

Experiment  108. — Add  a  few  drops  of  sodic  thiosulphate 
solution  to  a  solution  of  iodine.  The  colour  of  the  iodine  dis- 
appears : 

I2  +  2Na2S2O3  =  2NaI  +  NaaS4Oe. 

Thus,  by  means  of  a  solution  of  thiosulphate  of  known 
strength,  the  quantity  of  iodine  in  a  solution  can  be 
determined.  (How  could  a  quantity  of  chlorine  be 
measured  indirectly  by  this  reaction  ?) 

Experiment  109. — Prepare  a  little  argentic  chloride  by  mix- 
ing hydrochloric  acid  and  argentic  nitrate  in  solution.  Shake  up 
well,  allow  the  chloride  to  settle,  pour  off  the  liquid  carefully, 
add  water,  shake,  &c.,  and  repeat  this  operation  three  or  four 


136  HYDRIC    SULPHIDE. 

times.  (This  is  washing  by  decantation.)  Now,  add  some  sodic 
thiosulphate  solution  to  the  argentic  chloride.  The  latter  is  dis- 
solved, forming  a  very  sweet  solution  of  the  double  thiosulphate, 
NaAgS203  : 

AgCl  +  Na2S2O3  =  NaAgS20,  +  NaCl. 

If  any  of  the  argentic  chloride  has  become  darkened,  it  remains 
undissolved. 

Sodic  thiosulphate  is  used  in  photography  to  dissolve 
out  the  undarkened  chloride,  bromide,  or  iodide  of  silver, 
so  as  tojix  the  negative  or  print. — It  is  used  in  medicine 
to  destroy  certain  organisms  in  the  stomach  (sarcinae 
ventriculi).  (What  is  the  action  of  the  gastric  juice  on  it  ]) 

SULPHUR  AND   HYDROGEN. 

124.  Hydric  Sulphide,  H2S. — Commonly  called 
sulphuretted  hydrogen.  It  occurs  in  volcanic  gases  and 
in  "  sulphur  springs."  It  is  a  product  of  the  decay  of 
animal  and  vegetable  bodies,  and  is  one  cause  of  the  at- 
tendant bad  smell  and  poisoning  of  the  air. 

PREPARATION. — Experiment  110. — Pour  a  few  drops  of 
dilute  sulphuric  acid  upon  a  little  ferrous  sulphide  (FeS)  in  a  watch 
glass.  Note  the  odour  of  the  gas  given  off.  Repeat,  with  dilute 
hydrochloric  acid  and  ferrous  sulphide.  Try  also  with  antimony 
trisulphide  (Sb2S3)  and  rather  strong  hydrochloric  acid  solution 
(1  vol.  of  the  strong  solution  to  1  of  water),  warming  them  in  a  1. 1. 

The  gas,  known  by  its  unpleasant  smell,  is  hydric  sul- 
phide (H2S)  : 

Ferrous  sulphate. 

(1)  FeS  +  H2S04  =  H2S  +  FeSO4. 

Ferrous  chloride. 

(2)  FeS  +  2HC1  =  H2S  +  FeCl2. 

Antimony  trichloride. 

(3)  Sb2S3  +  6HC1  =  3H2S  +  2SbCl3. 


HYDRIC   SULPHIDE  137 

Experiment  111. — Heat  a  little  paraffin  wax  and  flowers  of 
sulphur  together  in  a  small  porcelain  dish.  Hydric  sulphide  is 
evolved. 

Paraffin  wax  is  a  mixture  of  hydrocarbons,  or  com- 
pounds of  carbon  and  hydrogen.  Hydric  sulphide  is 
prepared  on  the  large  scale  in  this  way. — When  hydric 
sulphide  is  prepared  from  ferrous  sulphide  it  contains 
hydrogen,  owing  to  the  presence  of  uncombined  iron  in 
the  ferrous  sulphide  of  commerce.  (What  is  the  action  ]) 
The  method  with  antimony  trisulphide  gives  the  pure 
gas. 

PROPERTIES. — An  invisible  gas,  of  very  unpleasant 
odour  (that  of  rotten  eggs),  and  of  sweetish  taste. 

Experiment  112. — Prepare  sulphuretted  hydrogen  with  fer- 
rous sulphide  and  dilute  sulphuric  acid,  evolving  it  in  a  small 
flask  or  bottle,  and  bubbling  it  through  a  small  quantity  of  water 
in  a  icash  bottle  (to  wash  it  free  from  small  drops  of  sulphuric  acid 
carried  up  by  the  gas),  and  then  through  water  in  the  reagent 
bottle  provided.  Gases  are  washed  by  passing  them  through  a 
liquid  in  a  bottle  provided  with  a  twice-bored  cork.  The  tube 
from  the  generating  flask  passes  through  one  hole  and  nearly  to 
the  bottom  of  the  water.  A  second  tube  passes  just  through 
the  other  hole  to  lead  the  washed  gas  away.  Note  that  after  the 
air  is  driven  out,  most  of  the  gas  dissolves  in  the  water.  Some, 
however,  escapes  solution,  and  this  experiment  should  be  made 
in  a  draught  cupboard  or  under  a  hood.  After  the  gas  has  been 
running  for  a  few  minutes,  bubble  a  little  of  it  through  test 
tubes  containing  solutions  of  cupric  sulphate  (CuS04),  tartar 
emetic,  and  stanuous  chloride  (Sn012).  Then  set  fire  to  the  gas 
and  note  its  inflammability.  Evaporate  the  solution  remaining 
in  the  flask,  and  obtain  crystals  of  green  vitriol  (FeS04.7H2O). 

Hydric  sulphide  is  soluble  in  water,  about  3  volumes 
in  1  of  water.  The  solution  has  the  smell,  &c.,  of  the 
gas.  It  slowly  decomposes,  absorbing  oxygen  from  the 


138  HYDRIC    SULPHIDE. 

air  (H.2S  -f-  O  =  H.,O  -\-  S),  and  sulphur  is  deposited 
as  a  white  powder.  The  gas  burns  in  the  air  with  a 
bluish  flame  and  the  formation  of  water  and  sulphur 
dioxide  :  H,S  +  30  =  H2O  -f  SO2.— It  can  be  lique- 
fied at  —  70°  C.  under  the  ordinary  pressure  of  the 
atmosphere. — Hydric  sulphide  is  a  valuable  reagent  (test 
substance)  in  analytical  chemistry.  Either  as  gas  or  in 
solution,  it  precipitates  the  sulphides  of  certain  metals 
from  solutions  of  their  salts,  while  it  does  not  precipitate 
others.  It  precipitates  the  sulphides  of  some  metals  in 
presence  of  a  free  acid,  others  only  in  the  presence  of  a 
free  alkali.  Thus,  all  metals  can  be  divided  into  three 
groups : 

1.  Those  precipitated  as  sulphides  in  presence  of  a  free 
mineral  acid  (Hg,  Ag,  Pb,  Cu,  Bi,  Cd,  Sb,  Sn,  As,  Pt, 
Au,  &c.) 

2.  Those  precipitated  only  in  presence  of  a  free  alkali 
(Fe,  Ni,  Co,  Zn,  Mn,  Al,  Or,  the  last  two  as  hydroxides). 

3.  Those   not  precipitated  at  all   by  hydric  sulphide 
(Mg,  Ca,  Sr,  Ba,  K,  Na,  Li,  &c.) 

Hydric  sulphide  is  very  poisonous,  even  when  largely 
diluted  with  air.  It  produces  lassitude,  headache,  giddi- 
ness, fainting,  and  at  last  death.  The  best  antidote  is 
chlorine  gas  largely  diluted  with  air.  It  can  be  adminis- 
tered by  inhalation  as  evolved  from  a  mixture  of  "  chloride 
of  lime  "  (bleaching  powder)  and  vinegar. 

Experiment  113-  — Mix  some  chlorine  water  with  solution  of 
hydrio  sulphide  in  a  t.  t.  Note  that  sulphur  is  precipitated.  If 
the  right  proportions  are  used,  the  smell  of  both  chlorine  and 
hydric  sulphide  is  destroyed  : 

H2S  +  C12  =  2HC1  +  S. 


SULPHIDES    OF    METALS.  139 

When  hydric  sulphide  is  taken  as  a  medicine  it  is  par- 
tially given  off  through  the  skin.  Hence,  its  efficacy  in 
skin  diseases. 

SULPHIDES  OF  METALS. — Sulphur  combines  directly 
with  most  of  the  metals,  forming  sulphides.  In  these 
compounds  the  sulphur  is  dyad,  as  in  hydric  sulphide. 
The  sulphides  of  the  heavy  metals  can  be  formed  by  pre- 
cipitation with  hydric  sulphide  (see  above) ;  the  sulphides 
of  potassium,  sodium,  calcium,  &c.,  can  not  be  formed  in 
this  way,  but  are  best  prepared  by  reduction  of  the  sul- 
phates by  heating  with  charcoal  : 

K2^O4  +  40  =  K,S  +  4  CO. 

The  sulphides  run  parallel  with  the  oxides  in  composi- 
tion, as  the  following  formulas  show  : 

H2S,  FeS,  Fe.2S3,  Fe3S4,  Ag,S,  Na2S,  As2S3,  <fec. 
H2O,  FeO,  Fe2O3,  Fe3O4,  Ag,O,  Na2O,  As2O3,  <fcc. 

Tests. — Free  hydric  sulphide  blackens  lead  jKiper. 
Pb(C2H9O.J11  -f  H4S  =  PbS  +  2C2H40,. 

Most  sulphides  when  treated  with  hydrochloric  acid  give  the 
smell  of  hydric  sulphide.  Insoluble  sulphides  when  fused  on 
charcoal  before  the  mouth  blow-pipe  with  sodic  carbonate,  give 
a  bead  which  stains  silver  black,  when  placed  upon  it  and  moist- 
ened with  dilute  hydrochloric  acid. 

HYDROGEN  PERSULPHIDE. — The  exact  composition  is 
unknown  ;  perhaps  H2S2,  analogous  to  H2O2.  It  is  pre- 
pared by  pouring  a  strong  solution  of  an  alkaline  persul- 
phide  (e.  g.,  K2S5),  into  dilute  hydrochloric  acid.  It  is 
an  oily  liquid,  readily  decomposing  into  sulphur  and 
hydric  sulphide. 

125.   Sulphur   and   the    Halogens.  — Sulphur 

unites  directly  with  chlorine  in  several  proportions 
(S2Clo,  S,C14,  and  SC14).  The  chlorides  are  all  decom- 


140  SELENIUM — TELLURIUM. 

posable  by  water. — Sulphur  combines  with  bromine  in  only 
one  proportion  (S2Br2) ;  with  iodine  it  forms  two  com- 
pounds (S2I2  and  SI6).  The  moniodide  (S2I2)  is  prepared 
by  heating  together  four  parts  of  iodine  with  one  of 
flowers  of  sulphur.  It  is  a  crystalline  greyish-black  solid, 
insoluble  in  water,  but  decomposed  by  boiling  with  it. 
It  is  soluble  in  glycerine.  It  is  used  in  medicine  in  the 
form  of  an  ointment. 

126.  Selenium  (Se"-iT-vi  =   79).— Rather   rare,   and 
found  only  in  small  quantities,  in  sulphur  ores,  &c.    It  was 
discovered  by  Berzelius  (1817),  in  the  deposit  found  in  cer- 
tain sulphuric  acid  chambers.    Selenium  has  two  allotropic 
forms,  one  of  which  is  somewhat  metallic  in  appearance. 

COMPOUNDS. — These  are  similar  in  composition  to  those 
of  sulphur ;  seleniuretted  hydroyen,  or  hydric  selenide 
(H2Se);  selenium  dioxide  (SeO2);  selenious  acid(H2SeO3>; 
and  selenic  acid  (H2SeO4),  are  similar  in  properties  to  the 
corresponding  compounds  of  sulphur ;  but  no  trioxide  is 
known  with  certainty.  Selenic  acid  differs  from  sulphuric 
acid  in  being  easily  reduced  to  selenious  acid.  Metallic 
selenium  conducts  electricity,  and  its  conductivity  is  in- 
creased by  exposure  to  light.  It  is  so  sensitive  to  the 
influence  of  light,  that  an  electrical  instrument  has  been 
constructed  which  shows  the  slightest  variations  in  the 
intensity  of  the  light  falling  on  the  surface  of  the  selenium 
which  is  placed  in  the  electric  circuit. 

127.  Tellurium  (Tei!- iv- vi  =  128).— This  is  also  a  rare 
element,    being  generally  found    combined   with    silver, 
gold,  or  bismuth.*    It  forms  compounds  with  hydrogen 
and  oxygen,  parallel  with  those  of  sulphur  and  selenium, 
but  tellurium  dioxide  (Te(X)  has  weak  basic  properties. 


QUESTIONS    AND    EXERCISES.  141 

QUESTIONS  AND  EXERCISES. 

1.  Show  how  the  oxides  of  sulphur  illustrate  the  law  of  mul- 
tiple proportions. 

2.  How  much  oxygen  is  there  to  1  part  by  weight  of  hydrogen 
in  each  of  the  oxygen  acids  of  sulphur  ?     How  much  sulphur  to 
1  of  hydrogen  ? 

3.  What  weight  of  sulphuric  acid  is  equivalent  to  100  g.  sodic 
hydroxide  (KOH)  ?    To  100  g.  calcic  hydroxide  (Ca(OH)2)  ? 

4.  196  g.  sulphuric  acid  in  solution  is  mixed  with  150  g.  sodic 
hydroxide.     Is  the  solution  neutral,  acid,  or  alkaline  ? 

5.  What  is  the  valence  of  sulphur  in  sulphur  dioxide  ?     In 
sulphur  trioxide  ?     In  hydric  sulphide  ?     In  sulphur  hexiodide 
(SI6)  ?     In  Sulphuric  acid  (S03  ZQH)  ? 

6.  What  volume  of  chlorine  is  necessary  to  decompose  1  cubic 
foot  of  hydric  sulphide  gas  ? 

7.  Define  normal  and  acid  salts  with  regard  to  the  equivalence 
of  the  acids  and  bases  forming  them. 

8.  Write   the  formulas  for   the  following  normal  sulphates : 
Sodic,  argentic,  ammonic,  ferric,  cupric,  zincic,  bismuth,  chromic, 
and  mercuric. 

9.  Write  the  formulas  for  the  following  compounds  :  potassic 
hydric   sulphite,    calcic  thiosulphate,    sodic   argentic  thiosulphate, 
zincic  hyposulphite,  mercuric  sulphide,  and  argentic  sulphide. 

10.  Moisten  a  strip  of  filter  paper  with  plumbic  acetate  and 
hold  it  in  the  mouth  of  the  bottle  of  hydric  sulphide  solution. 
It  is  blackened.     Explain. 

11.  What  weight  of  ferrous  sulphide  (FeS)  must  be  used  to 
saturate  2  litres  of  water  with  hydric  sulphide  ? 

12.  Balance  the  following  equations  : 

(1)  SbCl3  +  H2S  =  Sb2S3  +  HC1. 

(2)  H2S04  +  Fe203  =  Fe2  (SO4)3  +  H2O. 

(3)  SO2  +  I  +  H2O  =  H2SO4  +  HI. 

(4)  FeS04  =  Fe203  +  SO2  +  SO3. 


142  PHOSPHOUUS. 


CHAPTER     XI. 


PHOSPHORUS  AND  ITS  COMPOUNDS. 

128.  Phosphorus  (P"llT  =  31).— The  name  phos- 
phorus (light-bearer)  was  at  first  given  to  substances 
which  shone  in  the  dark,  particularly  to  Bononian  phos- 
phorus (barium  sulphide).  The  substance  at  present 
called  phosphorus  was  discovered  about  1678  by  an  al- 
chemist, Brandt,  while  distilling  the  residue  left  by 
evapoi'ating  urine.  It  created  a  great  deal  of  interest 
on  account  of  its  luminosity  and  inflammability  ;  and, 
besides,  it  was  held  to  be  a  very  valuable  medicine.  A 
few  years  after  Brandt's  discovery,  Boyle  discovered  the 
secret  of  its  preparation  from  urine,  having  had  only  the 
vague  information  that  it  was  prepared  from  the  human 
body.  At  present  the  phosphorus  of  commerce  is  pre- 
pared from  bones  and  from  mineral  phosphates. 

OCCURRENCE. — Phosphorus  is  never  found  in  nature 
uncombined.  (Why1?)  It  occurs  in  phosphates,  e.  g., 
phosphorite  (Ca3(P04)2),  apatite  (3Ca3(PO4)2  -j-  CaCl2), 
and  coprolites,  or  fossil  dung  of  serpents.  It  is  present  in 
all  fertile  soils  and  is  an  essential  constituent  of  the  food 
of  plants.  Animals  take  it  into  their  bodies  along  with 
their  food  and  build  it  up  especially  into  bones,  which 
consist  very  largely  of  calcic  phosphate  (Ca3(P04)2).  It 
is  also  found  in  sea-water,  and  in  all  parts  of  the  earth's 
crust. 


PHOSPHORUS.  143 

PREPARATION. — -Bones  are  either  distilled,  bone-black 
and  bone-oil  being  the  products ;  or  they  are  heated 
under  pressure  with  water  to  extract  the  animal  matter 
as  gelatine.  In  the  former  case,  the  bone-black  is  used 
in  sugar-refining  and  then  burned  to  bone-ash,  which  is 
impure  calcic  phosphate.  In  the  latter  case  the  calcic 
phosphate  is  left  from  the  extraction  process.  In  either 
of  these  ways,  then,  or  from  some  mineral  source,  calcic 
phosphate  is  obtained.  This  is  treated  with  dilute  sul- 
phuric acid,  when  two-thirds  of  the  calcium  is  precipi- 
tated as  calcic  sulphate,  while  an  acid  phosphate  goes 
into  solution  : 

Ca3(PO4)2  +  2H2SO4  =  2CaSO4  +  CaH4(PO4)2. 

The  solution  is  drawn  off,  evaporated  to  dryness,  and 
strongly  heated.  It  loses  water,  and  calcic  metaphos- 
phate  remains  : 

CaH4(PO4)2  =  Ca(-PO3)2  -f  2H2O. 

This  is  mixed  with  charcoal  and  sand  (SiO2),  and 
strongly  heated  in  earthenware  retorts.  The  object  of 
the  charcoal  is  to  remove  oxygen  from  the  phosphate ; 
the  sand  forms  a  silicate  of  calcium,  and  thus  aids  in 
setting  free  the  phosphorus,  which  distils,  and  is  collected 
under  water  : 

Ca(PO3)2  +  50  +  SiO2  =  CaSiO3  +  5CO  -f  2P. 

It  is  purified  by  melting  under  water  and  pressing  through 
chamois  leather,  or  by  oxidising  the  impurities  by  means 
of  a  mixture  of  potassic  bichromate  (K2Cr207)  and  sul- 
phuric acid. 

PROPERTIES. — Phosphorus  has  several  allotropic  modi- 
fications, of  which  two  only  will  be  described. 


144  YELLOW    PHOSPHORUS. 

(1)  COMMON  OR  YELLOW  PHOSPHORUS  is  a  translucent, 
waxy  solid,  of  sp.  wt.  1.8.  It  melts  at  44°  C. 

Experiment  114.  —  Put  a  small  piece  of  phosphorus  in  a  por- 
celain dish  half-filled  with  water  and  heat  the  water.  Notice 
that  the  phosphorus  melts  before  the  water  becomes  uncomfort- 
ably hot.  Note  the  smell,  &c.,  of  the  phosphorus.  Is  it 
heavier  or  lighter  than  water  ?  Does  it  dissolve  in  water  ? 

If  heated  in  an  atmosphere  of  some  inactive  gas  (nitro- 
gen) it  boils  at  290°  C.  The  specific  weight  of  its 
vapour  is  4.3  (air  =1).  (How  many  atoms  in  the 
molecule  1)  If  heated  in  the  air  it  catches  fire  at  60°  C. 
It  catches  fire  at  lower  temperatures  if  exposed  to  the 
oxidising  action  of  the  air  for  some  time.  Very  slight 
friction  often  sets  it  on  fire ;  and  as  it  burns  very  vigor- 
ously, and  causes  frightful  wounds  if  it  comes  in  contact 
with  the  skin  while  burning,  it  should  always  be 
handled  under  water.  It  should  be  kept  in  well  stop- 
pered bottles  of  water  in  a  cool,  dai-k  place. — Yellow 
phosphorus  is  soluble  in  carbon  bisulphide  (CS2),  in 
fats,  oils,  and  slightly  in  ether  and  oil  of  turpentine. 
If  the  solution  in  carbon  bisulphide  be  evaporated,  crys- 
tals of  phosphorus  are  obtained. — Yellow  phosphorus 
is  very  poisonous,  and  as  it  is  often  used  in  making 
matches,  phosphorus-poisoning  cases  are  not  rare.  The 
fatal  dose  is  as  low  as  1^  grains.  The  symptoms  are 
pain  in  the  stomach,  nausea,  vomit  with  garlic  odour. 
The  antidote  is  blue  vitriol  in  3-grain  doses  (dilute  solu- 
tion) every  5  minutes  until  vomiting  ensues. 

Experiment  115. — Put  a  bit  of  clear  phosphorus  in  a  small 
quantity  of  cupric  sulphate  solution.  It  is  blackened. 

When  a  poisonous  substance  is  taken  into  the 
stomach  it  begins  at  once  to  pass  into  the  circulation. 


RED    PHOSPHORUS.  145 

The  object  of  an  antidote  is  to  prevent  this,  either  by 
decomposing  the  poison  (if  it  is  a  compound),  or  by 
rendering  it  insoluble,  and  therefore  incapable  of  absor- 
ption into  the  circulation.  In  the  present  case,  the 
phosphorus  is  combined  with  copper  (Cu3P2),  forming  an 
insoluble  black  compound. 

2.  RED  OR  AMPHOROUS  PHOSPHORUS  is  prepared  from 
common  phosphorus  by  heating  it  for  some  time  at  240°C., 
a  small  quantity  of  iodine  being  added  to  hasten  the 
change.  It  is  an  amorphous  red  solid  of  specific  weight 
2.2  ;  insouble  in  carbon  bisulphide,  &c.;  does  not  fuse  at 
250°,  catches  fire  only  at  260°,  and  is  not  poisonous.  It 
is  prepared  on  the  large  scale  in  the  manufacture  of 
matches,  to  replace  the  poisonous  common  phosphorus. 

Tests.  —  Phosphorus  can  be  known  by  its  odour  of  garlic,  and 
and  by  its  luminosity  in  the  dark.  To  detect  small  quantities* 
e.g.,  in  cases  of  poisoning,  the  substance  is  distilled  with  water 
in  a  dark  room,  when  the  phosphorescence  appears. 

PHOSPHORUS  AND  OXYGEN. 

129.  Oxides  Of  Phosphorus.—  Two  are  known 
with  certaint  : 


Phosphorus  pentoxide 
Phosphorus  trioxide 


and  the  existence  of  a  third  is  probable  : 

Phosphorus  monoxide  ........  P2O  (?) 

Derived  from  these  are  three  acids  : 

Phosphoric  acid  .............  H3PO4 

Phosphorous  acid    ..........  H3PO3 

Hypophosphorous  acid  .......  H3PO2 

11 


146  OXIDES    OF    PHOSPHORUS. 

130.  Phosphorus  Pentoxide. — P206. —  Already 

noticed.  Prepared  by  burning  phosphorus  in  air.  It  is  a 
snow-like  solid,  very  hygroscopic,  and  hence  used  to  dry 
gases,  and  in  separating  water  from  acids.  In  moist  air 
it  quickly  deliquesces  forming  metaphosphoric  acid  : 

H2O  +  P2O5  =  2HPO3. 

If  put  into  water  it  hisses  and  dissolves,  forming  at  first 
metaphosphoric  acid  ;  but,  gradually  at  ordinary  tempera- 
tures, and  quickly  if  heated,  it  combines  with  more  water 
to  form  orthophosphoric  acid  : 

HP03  +  H20  =  H3P04. 

131.  Phosphorus  Trioxide. — P2O3.— Prepared  by 
burning  phosphorus  slowly  with  a  scant  supply  of  air. 
If  heated  in  air  it  combines  with  oxygen  and  forms  the 
pentoxide : 

PA  +  o2  =  PA. 

It  combines  with  water  to  form  phosphorous  acid : 
PA  +  3H20  =  2H3P03. 

132.  Phosphoric    Acid.  —  Phosphorus    pentoxide 
unites  with   water  in   three  proportions,  forming  three 
different  acids,  each  of  which  is  called  phosphoric  acid, 
but  prefixes  are  used  to  distinguish  them.     The  ordinary 
acid  is  that  formed  with  the  greatest  proportion  of  water, 
and  is  called  or^Aophosphoric  acid  : 

3H20  +  PA  -  2H3PO4. 

(1)  ORTHOPHOSPHORIC  ACID,  H3PO4. — This  is  gener- 
ally called  simply  phosphoric  acid. 


PHOSPHORIC    ACID.  147 

PREPARATION. — Bone  ash  or  mineral  phosphate  is  dis- 
solved in  nitric  acid  : 

Ca3(PO4),  +  4HN03  =  CaH4(PCV2  +  2Ca(N03)2. 

Plumbic  acetate  (PbA2)  is  added  to  the  solution,  when 
plumbic  phosphate  is  precipitated  : 

CaH4(PO4)2   -f  3PbA2  =  Pb3(PO4)2  -f  CaA2  -f  4HA. 

The  object  of  the  first  reaction  is  to  get  the  phosphate 
and  impurities  in  solution  ;  that  of  the  second  is  to 
remove  the  phosphoric  acid  from  its  impurities,  and  at 
the  same  time  obtain  a  phosphate  from  which  the  acid 
can  be  easily  obtained  in  one  operation.  The  plumbic 
phosphate  is  washed,  suspended  in  water,  and  decomposed 
by  sulphuretted  hydrogen  : 

Pb3(PO4)2  +  3H2S  =  2H3PO4  +  3PbS. 

Phosphoric  acid  dissolves,  and  plumbic  sulphide  remains 
undissolved.  (How  is  the  process  completed  1) 

Experiment  116. — -Put  a  small  piece  of  phosphorus  in  a  por- 
celain dish  and  cover  it  with  dilute  nitric  acid.  Warm  gently. 
The  phosphorus  melts  and  oxidises  with  the  evolution  of  red 
fumes.  (What  are  these  ?)  Evaporate  to  dryness,  add  a  few 
drops  concentrated  nitric  acid,  and  evaporate  again  to  dryness. 
Phosphoric  acid  remains.  Dissolve  in  water  and  keep. 

Red  phosphorus  is  better  for  this  experiment,  as  it  can 
be  oxidised  with  concentrated  nitric  acid,  which  acts 
explosively  on  yellow  phosphorus. 

PROPERTIES. — A  thick  syrup,  or  hard,  transparent, 
colourless  crystals  ;  soluble  in  water,  giving  a  strongly 


148  PHOSPHATES. 

acid   solution.     (Try  it   by  taste  and   litmus.)     It  is  a 
tribasic  acid,  and  forms  three  series  of  salts,  e.  g. : 

(1)  Na3PO4,  trisodic  phosphate. 

(2)  Na2HPO4,  disodic  hydric  phosphate. 

(3)  NaH2PO4,  sodic  dihydric  phosphate. 

(To  which  of  these  classes  do  the  phosphates  mentioned 
in  PREPARATION  belong  ?) 

Phosphoric  acid  is  used  in  medicine  in  dilute  solution 
as  a  tonic,  <fec. 

PHOSPHATES  (orthophosphates). — Of  the  normal  and  the 
mon-acid  phosphates,  only  those  of  the  alkali  metals 
(Na,  K,  NH4,  &c.),  are  soluble  in  water.  The  di-acid 
phosphates  are  all  soluble. 

Tests. — 1.  Add  a  drop  or  two  of  argentic  nitrate  to  a  solu- 
tion of  sodic  phosphate  (NaaHP04).  A  yellow  precipitate  is 
formed.  Divide  into  3  parts,  and  add  to  one  part  ammonia  ;  to 
the  second,  dilute  nitric  acid  ;  and  to  the  third,  acetic  acid.  The 
precipitate  is  soluble  in  each.  (Could  this  test  be  used  in  a  solu- 
tion containing  hydrochloric  acid  ?). 

2.  Add  ammonic  chloride,  magnesic  sulphate,  and  ammonia  to 
a  solution  of  sodic  phosphate,   and  stir  with  a  glass  rod.     A 
granular    white    precipitate    of    magnesic    ammonic    phosphate 
(Mg.NH4.POJ  is  formed  : 

NaaHP04  +  NHS  +  MgS04  =  MgNH4P04  +  Na.,S04. 

3.  Baric  chloride  (BaCl2)  gives  a  white  precipitate,  soluble  in 
nitric  acid. 

Phosphates  are  present  dissolved  in  urine,  and  when 
the  urine  becomes  alkaline  by  decomposition  (Explain.) 
crystals  of  microcosmic  salt,  or  sodic  ammonic  hydric 
phosphate  (Na.NH4.H.P04)  are  formed. 


METAPHOSPHORIC    ACID.  149 

(2)  METAPHOSPHORIC  ACID. — HPO3. — Formed  when 
phosphorus  pentoxide  is  dissolved  in  cold  water ;  but  is 
generally  prepared  by  heating  orthophosphoric  acid  to  a 
red  heat : 

F3P04  =  HPO3  +  H2O. 

PROPERTIES. — It  is  a  hard  glassy  colourless  solid 
(glacial  phosphoric  acid) — the  common  phosphoric  acid 
of  commerce.  It  dissolves  in  water,  but  gradually  com- 
bines with  the  water  to  form  ortho-phosphoric  acid. — It 
is  a  monobasic  acid  (To  what  acid  is  it  similar  in  compo- 
sition]), and  forms  salts  called  metaphosphates,  e.g., 
NaP03,  Mg(P03)2,  &c. 

Tests. — 1.  Add  argentic  nitrate  to  a  solution  of  the  acid  or 
a  salt.  A  gelatinous  white  precipitate  is  formed .  It  is  soluble 
in  nitric  acid. 

2.  Add  some  white  of  egg  to  a  solution  of  metaphosphoric 
acid,  or  to  a  metaphosphate  acidified  with  acetic  acid.  The 
white  of  egg  is  coagulated.  Try  with  orthophosphate. 

(3)  PYROPHOSPHORIC    ACID. — H4P2O7    (  =  P2Ofi    -f- 
2H20). 

PREPARATION. — Heat  orthophosphoric  acid  till  the 
temperature  rises  to  215°  C. : 

2H3P04  =  H4P207  +  H2O. 

PROPERTIES. — A  sott  glassy  liquid  ;  when  heated  with 
water,  forms  orthophosphoric  acid.  It  is  a  tetrabasic 
acid.  Its  salts  can  be  formed  by  heating  acid  orthophos- 
phates  of  the  class  M2HP04.  (Write  the  equation.) 
The  pyrophosphates  are  changed  to  acid  ortho-salts  when 
they  are  boiled  with  water. 


150  PHOSPHOROUS    ACID. 

Tests. — 1.  White  precipitate  with  argentic  nitrate,  soluble 
in  nitric  acid. 

2.  Pyrophosphoric  acid  does  not  coagulate  white  of  egg. 

133.  Phosphorous  Acid.— H8PO3. 

PREPARATION. — This  acid  is  formed  when  phosphorus 
oxidises  slowly  in  moist  air ;  but  it  is  best  prepared  by 
the  action  of  phosphorus  trichloride  (P013)  on  water  : 

PC13  +  3H2O  =  P(OH)3  +  3HC1. 

This  action  of  water  on  a  chloride  is  a  very  frequent  one. 
Observe  that  hydroxyl  ( —  OH)  replaces  chlorine  ( —  Cl). 

PROPERTIES. — A  white  solid  with  garlic  odour.  Oxi- 
dises readily  to  phosphoric  acid.  When  heated  it  de- 
composes into  phosphoric  acid  and  gaseous  phosphoretted 
hydrogen  (PH3)  : 

4H3P03  =  3H3PO<  +  PH3. 

There  are  many  decompositions  of  compounds  containing 
oxygen,  such  that  one  part  gains  oxygen  at  the  expense 
of  another  part.  (Mention  one  already  described.) — 
The  phosphites  are  unimportant.  Only  two  of  the  three 
atoms  of  hydrogen  seem  to  be  readily  replaceable  by 
metal. 

Tests. — 1.  Add  argentic  nitrate  to  a  solution  of  a  phosphite. 
Metallic  silver  is  precipitated. 

2.  Add  solution  of  baric  chloride.     It  gives  a  white  precipitate 
soluble  in  nitric  acid. 

3.  Heat  a  little  of  the  substance  on  mica  and  smell.     The 
odour  of  phosphoretted  hydrogen  is  observed. 

4.  Add  lime  water  to  solution  of  phosphorous  acid  or  a  phos- 
phite.    A  white  precipitate  is  thrown  down  (CaHP03). 


HYPOPHOSPHOROUS   ACID.  151 

134.  Hypophosphorous   Acid.  —  H3P02.  —  The 

acid  itself  is  of  little  importance.  It  can  be  prepared  from 
its  barium  salt  by  treating  with  sulphuric  acid,  (Why 
use  the  barium  salt  in  particular?)  It  is  a  monobasic 
acid,  of  very  strong  reducing  power,  and  quickly  absorbs 
oxygen  from  the  air. 

HYPOPHOSPHITES.— Experiment  117.— Heat  a  small  piece 
of  phosphorus  in  a  t.  t.  with  milk  of  lime.  Connect  the  t.  t. 
with  a  gas-delivery  tube  dipping  under  water.  A  gas  having 
the  smell  of  decaying  fish  is  evolved,  and  catches  fire  as  soon  as 
it  comes  in  contact  with  the  air.  Continue  the  heating  until 
the  phosphorus  is  dissolved,  and  no  more  gas  appears.  The 
solution  contains  calcic  hypophoiphite  (Ca(H4PO,)2).  Filter, 
pass  into  the  clear  solution  a  current  of  carbon  dioxide  to  preci- 
pitate the  uncombined  calcic  hydroxide  as  calcic  carbonate 
(CaCOs),  then  filter  again. 

The  actions  which  take  place  are  represented  thus  : 
3Ca(OH)2  +  8P  +  6H2O  =  3Ca(H2PO2)2  +  2PH3. 
Ca(OH)2  +  CO2  =  CaCO3  -f-  H2O. 

The  hypophosphites  of  sodium,  potassium,  barium,  and 
strontium,  may  be  prepared  in  the  same  way.  The 
hypophosphites  are  powerful  reducing  agents.  They 
must  be  kept  out  of  contact  with  air.  otherwise  they 
become  oxidised  to  phosphites  and  finally  to  phosphates. 
They  are  used  in  medicine,  and  are  supposed  to  produce 
the  same  specific  effects  as  free  phosphorus ;  but  this  is 
very  doubtful. 

Tests. — 1.  Argentic  nitrate  gives  a  white  precipitate  which 
quickly  turns  brown  and  then  black  : 

4AgN03  +  H3P02  +  2H20  =  4HN03  +  H3PO4  +  4Ag. 
2.  Baric  chloride  gives  no  precipitate, 


152  PHOSPHINE. 

3.  Zinc  and  hydrochloric   acid  (nascent  hydrogen)  give  the 
smell  of  phosphoretted  hydrogen. 

4.  Evaporate  to  dryness,  heat  residue  on  mica  ;  odour  of  phoa- 
phoretted  hydrogen  : 

2H3P02  =  PH3  +  H3P04. 

Note. — Try  these  tests  •with  the  solution  of  calcic  hypophos- 
phite  (Experiment  117) ;  also  with  a  sample  of  "  hypophosphites  " 
from  the  druggist's. 

The  "  hypophosphites  "  of  medicine  is  a  preparation  of 
calcic  hypophosphite. 

135.  Phosphorus  and  Hydrogen.  —  There  are 

three  compounds  known : 

Gaseous  phosphoretted  hydrogen,  or  phosphine.  .PH3 

Liquid  "  "  P2H4 

Solid  "  "  P4H2 

Only  the  first  will  be  described. 

136.  Phosphine    (gaseous  phosphoretted  hydrogen}. 
PH3. — Prepared  as  in  Experiment  117,  but  better  with 
a  strong  solution  of  potassic  or  sodic  hydroxide  : 

3KOH  +  4P  +  3H20  =  3KH2PO2  +  PH3. 
It  is  formed  by  the  decay  of  fish  and  other  animal  and 
vegetable  bodies.  It  is  supposed  to  be  the  cause  of  that 
dancing  phosphorescent  light  which,  under  the  name  of 
ignis  fatuus,  or  will  o'  the  wisp,  wanders  over  marshy 
and  swampy  places. 

PROPERTIES. — An  invisible  gas,  of  strong  smell  (that 
of  rotting  fish ),  very  poisonous,  even  when  largely  diluted 
with  air.  It  readily  burns  in  air  (spontaneously  inflam- 
mable when  containing  vapour  of  the  liquid  compound, 
as  in  Experiment  117)  : 

PH3  +  202  =  H3PO,. 


CHLORIDES    OF    PHOSPHORUS.  153 

In  chemical   characters   it   resembles  ammonia,  uniting 
with  acids  to  form  salts,  e.  g.  : 

Phosphonium  iodide. 

PH3  +  HI  =  PH4I. 
PHOSPHORUS  AND  THE  HALOGENS. 

137.  Phosphorus  and  Chlorine. — Two  com- 
pounds are  known  : 

Phosphorus  trichloride PC13 

"  pentachloride PC15 

PHOSPHORUS  TRICHLORIDE  (PC13),  is  prepared  by  pass- 
ing dry  chlorine  gas  over  heated  amorphous  phosphorus. 
It  is  a  colourless  liquid,  fuming  in  moist  air,  decomposed 
by  water  (Art.  123).  It  boils  at  76°  0. 

PHOSPHORUS  PENTACHLORIDE  (PC15),  is  prepared  by 
treating  amorphous  phosphorus  with  dry  chlorine  at  a 
low  temperature,  until  no  more  is  absorbed.  It  is  a  white 
solid,  of  pungent  smell,  fumes  in  moist  air,  being  decom- 
posed by  water  in  two  stages  : 

P.  oxychloride. 

(1)  PC15  +  H2O  =  POC13  +  2HC1. 

(2)  POC13  +  3H2O  =  PO(OH)3  +  3HC1. 

In  the  first  stage  two  atoms  of  chlorine  are  replaced  by 
an  atom  of  oxygen,  and  phosphoryl  chloride  (or  phos- 
phorus oxychloride),  and  hydrochloric  acid  are  formed. 
This  is  the  action  with  a  small  proportion  of  water.  In 
the  second  stage  the  three  remaining  atoms  of  chlorine  are 
replaced  by  three  hydroxyl  groups,  and  phosphoric  acid 
is  formed,  along  with  a  further  quantity  of  hydrochloric 
acid.  From  this  it  is  seen  that  phosphoric  acid  is  the 
hydroxide  of  a  trivalent  radical  phosphoryl  (PO^).  Oxy- 
gen acids  generally  can  be  represented  as  hydroxides,  e.  g. : 
N02-OH,  S02(OH)2. 


154  QUESTIONS   AND    EXERCISES. 

138.  Phosphorus  and  Bromine,  &c. — With  the 

other  halogens  phosphorus  combines  forming  compounds 
the  formulas  of  which  are  : 

PBr3,  PBr5;  P2I4,  PI3 ;  and  PF5. 

These  compounds,  excepting  the  last,  are  prepared  by 
direct  union  of  the  elements  ;  and  all  ai-e  of  great  value 
in  preparing  halogen  derivatives  of  carbon  compounds. 
By  means  of  these  (as  well  as  the  chlorides),  the  oxygen 
and  hydroxyl  of  carbon  compounds  can  be  replaced  by 
the  halogens.  (Compare  the  action  of  phosphorus  penta- 
chloride  on  water.) 


QUESTIONS  AND  EXERCISES. 

1.  Compare  nitrogen  and  phosphorus,  as  to  their  compounds. 

2.  In  what  dehydrating  processes  is   phosphorus  pentoxide 
used? 

3.  When  an  orthophosphate  of  the  class  MH2P04  is  heated, 
water  is  driven  off.     What  is  left  ?     Write  the  equation. 

4.  Mix  a  solution  of  sodic  phosphate  (Na2HPO4)  with  one  of 
plumbic  acetate.     What  result  ?     Try  with  phosphite  and  hypo- 
phosphite. 

5.  When  an  orthophosphate  of  the  class  M2HPO4  is  heated, 
water  is  driven  off.     What  is  left  ?     Write  the  equation. 

6.  What  is  the  basicity  of  phosphorous  acid  ? 

7.  By  what  tests  can  the  acids  hypophosphorous,  phosphorous, 
and  phosphoric  be  distinguished  from  each  other  ? 

8.  "The  atomicity  of  phosphorus  is  3  in  some  compounds  and 
5  in  others."     Apply  this  statement. 

9.  Calculate  the  wt.  of  1  litre  of  phosphorus  vapour  at  400°  C., 
and  760  mm.  pressure. 

10.  Calculate  the  wt.  of  1  litre  of  phosphorus  vapour  at  1040* 
C.,  and  760  mm.  pressure. 

11.  What  ia  the  percentage  of  phosphorus  in   pure   apatite 
(Cas(P04)2.CaCl2)? 


ARSENIC.  155 


CHAPTER    XII. 


ARSENIC  AND  ITS  COMPOUNDS. 

139.  Arsenic  (As'"'T=  75). — Arsenic  is  found  free 
in  nature  in  crystalline  masses.  The  ores  from  which 
the  elementary  substance  is  principally  derived  are  mis- 
pickel  (Fe2As2S2),  arsenical  iron  (FeAs2),  and  arsenical 
ores  of  cobalt,  nickel,  and  copper.  Iron  pyrites  often 
contain  arsenic. 

PREPARATION. — Chiefly  by  sublimation  from  mispickel. 
The  ore  is  heated  in  earthenware  retorts,  and  the  metal 
condensed  in  sheet  iron  and  earthen  cylinders  : 

Fe2As2S2  =  2FeS  +  2As. 

It  can  be  prepared  on  the  small  scale  by  heating  the  tri- 
oxide  (As-jOj)  with  charcoal. 

Experiment  118- — Mix  a  small  quantity  of  white  arsenic 
(Asa03)  with  powdered  charcoal,  put  in  a  narrow  hard-glass  tube 
closed  at  one  end  (matrass),  drop  in  a  small  splinter  of  charcoal, 
and  heat  the  tube  slowly  in  the  Bunsen  flame,  applying  the  flame 
first  to  the  part  of  the  tube  in  which  the  charcoal  splinter  is.  A 
mirror  of  arsenic  forms  as  a  ring  around  the  inside  of  the  tube. 
The  charcoal  has  reduced  the  oxide  of  arsenic  : 

As2O3  +  30  =  2As  +  SCO. 

PROPERTIES. — A  steel  grey  solid,  of  somewhat  metallic 
appearance.  Sp.  wt.  =  5.727.  It  sublimes  at  356°  C. 
It  can  be  melted  by  heating  under  pressure.  Sp.  wt. 


156  ARSENIC    TRIOXIDE. 

of  vapour  =10.2  (air  =1).  (Calculate  the  number  of 
atoms  in  the  molecule.)  It  oxidises*  in  moist  air ;  and 
burns  brightly  when  heated  in  air  or  oxygen,  forming  the 
trioxide  (As2O3).  Can  also  be  oxidised  by  heating  with 
nitric  or  sulphuric  acid.  Tt  is  not  soluble  in  hydrochloric 
acid.  Arsenic  and  all  its  volatile  compounds  are  very 
poisonous.  Lead  is  hardened  in  the  manufacture  of  shot 
and  bullets  by  alloying  it  with  a  small  proportion  of 
arsenic. 

140.  Arsenic  and  Oxygen. — Two  oxides  are 

known,  viz.  : 

Arsenic  trioxide As203 

Arsenic  pentoxide As2O5 

They  are  both  acid-forming  oxides,  and  this  marks  arsenic 
as  a  non-metal,  although  it  has  some  metallic  characters. 

141.  Arsenic  Trioxide  (As2O3). — Also  called  (in- 
correctly) arsenious  acid.      This  is  the  "  white  arsenic  " 
of  the  shops. 

PREPARATION. — By  roasting  arsenical  pyrites  and  other 
ores  of  arsenic  in  a  good  current  of  air  and  receiving  the 
fumes  in  cool  chambers  to  condense  them.  These  crude 
"  flowers  of  arsenic  "  are  purified  by  resublimation. 

PROPERTIES. — A  white  solid,  either  a  crystalline  pow- 
der, or  a  porcelain-like  cake  ;  of  sweetish,  metallic  taste  ; 
sparingly  soluble  in  water.  (Carefully  study  the  appear- 
ance of  specimens.)  It  is  volatile,  and  sublimes  readily. 
(Try  a  little  in  a  1. 1.) 

Experiment  119. — Boil  a  little  arsenic  trioxide  with  distilled 
water  in  a  t.  t.  It  dissolves  only  to  a  small  extent.  Prove  this 
by  evaporating  a  little  of  the  clear  liquid  in  a  watch  glass.  Add 


ARSENIC    PENTOXIDE.  157 

about  £  of  its  volume  of  hydrochloric  acid  to  the  water,  and  heat 
again.  The  arsenic  trioxide  dissolves  readily.  Keep  the  solution. 
Experiment  120. — Try  the  solubility  of  arsenic  trioxide  in 
solutions  of  sodic  hydroxide,  ammonia,  nitric  acid,  and  sulphuric 
acid. 

Arsenic  trioxide  has  very  weak  basic  properties.  It 
forms  with  acids  no  salts  which  are  not  decomposed  by 
water.  Its  acid  forming  properties  are  distinct.  (What 
proof  of  this  have  you  had  ?)  It  is  used  in  the  manufac- 
ture of  pigments,  glass,  &c.,  and  in  calico  printing. — 
One  grain  (0.06  gram)  of  arsenic  trioxide  is  a  dangerous 
dose.  Doses  of  2  to  4  grains  are  nearly  always  fatal ;  but 
by  use  the  system  becomes  able  to  withstand  its  action. 
In  one  case  more  than  five  grains  were  eaten  without  ill 
effect. 

142.  Arsenic  Pent OXide.— As2O6.— Is  unimportant. 
Can  be  prepared  by  heating  arsenic  acid  (H3AsO4)  nearly 
to  a  red  heat :  2H3AsO4  =  3H2O  -|-  As2O5.     (Compare 
with  phosphoric  acid).     At  a  red  heat  it  begins  to  de- 
compose into  oxygen  and  arsenic  trioxide. — It  is  a  deli- 
quescent white  solid,  uniting  with  water  to  form  arsenic 
acid. 

143.  Arsenious    Acid— H3AsO3    (hypothetical).— 
This  acid   is  known  only   in  solution  and  through  its 
salts.     Its  relation  to  arsenic  trioxide  is  the  same  as  that 
of  phosphorous  acid    to    phosphorus   trioxide.     It  is   a 
tribasic  acid ;  but  arsenites  are  known  corresponding  to 
a  monobasic  acid,  HAsCv     (Compare  HJMO.,). 

ARSENITES. — Experiment  121. — Heat  powdered  arsenic  tri- 
oxide with  solution  of  sodic  hydroxide  until  no  more  will  dis- 
solve, filter,  dilute  the  nitrate,  and  preserve  it  for  the  following 
experiments. 


158  ARSENIOUS    ACID    AND    ARSENITES. 

This    solution   contains  an   arsenite  of  sodium ;   say 

»/  *  J 

Na3AsO3. 

Experiment  122. — To  a  little  of  the  solution  add  a  few  drops 
of  solution  of  ferric  chloride  (FeaCl6).  A  reddish  brown  precipi- 
tate appears : 

Fe2Cl6  +  2Na3AsO3  =  Fe2(AsO3)2  +  GNaCl. 
Test  its  solubility  in  dilute  hydrochloric  acid,  and  in  ammonia. 

This  precipitate  of  ferric  arsenite  is  harmless  when  taken 
into  the  stomach,  provided  the  contents  of  the  stomach 
are  not  acid.  Hence,  freshly  precipitated  ferric  hydroxide 
(Fe2(OH)6)  is  the  best  antidote  to  poisoning  by  arsenic. 
It  unites  with  the  acid  of  the  stomach  to  form  ferric 
chloride  and  precipitates  the  arsenic  as  arsenite.  The 
ferric  hydroxide  can  be  prepared  by  the  action  of  am- 
monia on  ferric  chloride,  but  must  then  be  carefully 
washed,  since  ferric  arsenite  is  somewhat  soluble  in  am- 
monia. It  is  better  prepared  by  the  action  of  calcined 
magnesia  on  solution  of  ferric  chloride  ;  or  a  mixture  of 
sodic  carbonate  and  ferric  chloride  may  be  given. 

Experiment  123. — Add  some  cupric  sulphate  solution  to 
sodic  arsenite  in  a  t.  t.  A  green  precipitate  (Scheele's  green)  is 
thrown  down  : 

CuSO4  +  Na2HAs03  =  CuHA8O3  +  Na2SO4. 

This  green  is  often  used  for  colouring  wall  papers,  <fcc., 
and  has  been  the  cause  of  many  cases  of  poisoning. 
Schweinfurth  green  is  prepared  from  Scheele's  green  by 
boiling  it  with  acetic  acid.  (Try  it).  Paris  green  is  a 
variable  mixture  of  these  substances  with  others  used 
merely  as  make-weights,  or  diluents. 

Experiment  124. — Add  a  few  drops  of  argentic  nitrate  solu- 
tion to  sodic  arsenite.  A  yellow  precipitate  of  argentic  arsenite 
i»  formed.  Test  its  solubility  in  nitric  acid  and  in  ammonia. 


ARSENIC    ACID    AND    ARSENATES.  159 

144.  Arsenic  Acid. — H3AsO4. — Prepared  by  action 
of  oxidising  agents  and  water  on  arsenic  trioxide. 

Experiment  125- — Put  a  little  arsenic  trioxide  in  a  small 
porcelain  basin,  add  a  small  quantity  of  strong  nitric  acid,  and 
heat  on  the  water  bath.  Red  fumes  are  evolved  (N2O3) ;  evapo- 
rate to  a  syrup,  and  keep.  This  is  arsenic  acid  containing  a  little 
water : 

2HNO3  +  As2O3  +  2H2O  =  2H3AsO4  +  N2O3. 

Orthoarsenic  acid  (H3AsO4)  is  a  strong  tribasic  acid. 
Pyroarsenic  (H4As.X)7),  and  metarsenic  (HAsO3)  acids 
can  be  prepared  by  heating  the  ortho-acid. 

ARSENATES  are  very  like  the  corresponding  phosphates, 
being  generally  exactly  the  same  in  crystalline  form  (iso- 
morphous).  Disodic  hydric  arsenate  (Na2HAsO4.7H20) 
is  used  in  medicine.  It  is  a  white,  crystalline,  soluble 
salt,  prepared  by  fusing  together  arsenic  trioxide  (10 
parts),  sodic  nitrate  (8£  parts),  and  dry  sodic  carbonate 
(5^  parts) ;  dissolving  in  35  parts  of  distilled  water, 
filtering,  and  evaporating. 

Experiment  126. — Dissolve  the  acid  from  Experiment  125  in 
*  little  water,  colour  with  litmus,  and  add  solution  of  caustic 
soda  until  neutral.  With  this  solution  make  the  tests  for 
orthophosphoric  acid  (Art.  122). 

Arsenates  respond  to  the  same  tests  as  phosphates, 
but  argentic  nitrate  gives  a  chocolate  coloured  precipi- 
tate (Ag3AsO4). 

Experiment  127. — To  solution  of  sodic  arsenate  add  a  little 
sodic  acetate  (NaC2H302)  and  then  ferrous  sulphate  (FeSO4). 
A  white  precipitate  of  ferrout  arsenate  (Fe3(As04)2)  is  thrown 
down.  It  quickly  turns  greenish  owing  to  absorption  of  oxygen 
from  the  air  : 

2Na2HAsO4  +  2NaC2H302  +  3FeSO4  = 

Acetic  acid. 

Fe3(AaO4)2  +  3Na2SO4  +  2HC2H3O2. 


160  SULPHIDES    OF    ARSENIC. 

Ferrous  arsenate  is  used  in  medicine.  Ferrous  phos- 
phate (Fe3(PO4)2)  can  be  prepared  in  the  same  way 
from  sodic  phosphate.  (Try.) 

145.  Arsenic  and  Sulphur.— Three  sulphides  are 
known  : 

Arsenic  disulphide,  or  realgar As2S2 

Arsenic  trisulphide,  or  orpiment As2S3 

Arsenic  pentasulphide As3S5 

Experiment  128- — Treat  some  solution  of  arsenic  trioiide  in 
dilute  hydrochloric  acid  with  hydric  sulphide.  A  yellow  preci- 
pitate (As2S3)  is  formed  : 

As2O3  +  3H2S  =  As2S3  -f  3H2O. 

Test   its   solubility   in   sodic   hydroxide,  in   ammonic   sulphide 
((NH4)aS),  and  in  strong  hydrochloric  acid. 

Experiment  129. — Dissolve  a  little  arsenic  trioxide  in  sodic 
hydroxide  and  treat  the  solution  with  hydric  sulphide.  No  pre- 
cipitate is  formed.  (Why  is  arsenic  trisulphide  not  formed,  as 
in  Experiment  128  ?) 

Experiment  130. — Test  some  solution  of  sodic  arsenate  with 
hydric  sulphide.  No  precipitate  is  formed.  Acidify  with 
hydrochloric  acid  and  repeat  the  test.  A  yellow  precipitate 
(AsaSs  -f-  S2)  is  formed.  Try  its  solubility  as  in  Experiment 
128. 

Two  of  the  sulphides  of  arsenic  (As2S3  and  As2S6)  com- 
bine with  the  sulphides  of  the  alkali  metals  to  form 
soluble  sulphur  salts,  in  which  sulphur  plays  the  part  of 
the  oxygen  of  ordinary  salts.  This  explains  the  solu- 
bility of  these  sulphides  in  alkaline  sulphide  solutions ; 
for  example,  in  Experiment  128,  when  arsenic  tri- 
sulphide dissolves  in  ammonic  sulphide,  ammonic  sulph- 
arsenite  ((NH4)3AsS3)  is  formed. 


MARSH'S  TEST.  161 

146.  Arsenic  and  Hydrogen. — Arsine,  or  arseni- 
uretted  hydrogen  (AsH3)  is  the  only  compound  of  impor- 
tance. 

Experiment  131- — Fit  a4-ounce  wide-mouthed  bottle  with 
a  cork  twice  bored  and  provided  with  a  thistle  funnel  passing 
nearly  to  the  bottom,  and  a  hard  glass  tube  passing  just  through 
the  cork,  both  tubes  fitting  tightly.  The  hard  glass  tube  must 
be  bent  at  right  angles  and  drawn  out  at  the  point. .  Cover  the 
bottom  of  the  bottle  with  granulated  zinc,  push  in  the  stopper 
tightly,  pour  in  through  the  funnel  enough  dilute  sulphuric  acid 
to  fill  the  bottle  about  one-fourth.  Be  sure  that  the  funnel  dips 
below  the  surface  of  the  acid.  Hydrogen  is  evolved  and  issues 
from  the  glass  jet.  After  two  or  three  minutes  wrap  a  towel 
around  the  bottle  and  light  tKe  hydrogen.  (Why  not  at  once  ?) 
Hold  a  piece  of  cold  porcelain  in  the  flame  for  a  second  or  two  and 
see  if  it  is  blackened.  If  it  is,  there  is  probably  arsenic  in  the 
zinc.  If  not,  pour  through  the  funnel  a  few  drops  of  aqueous 
solution  of  arsenic  trioxide.  The  flame  of  hydrogen  very  soon 
becomes  livid  bluish.  Try  with  the  cool  porcelain.  A  metallic 
spot  is  formed.  Make  several  of  these,  and  then  set  the  apparatus 
in  a  draught  or  out  of  doors.  Try  the  effect  on  the  spots,  of  (1) 
ammonic  sulphide,  (2)  hydrochloric  acid,  (3)  bleaching  powder 
solution.  Results  : 
(1)  (2) 

(3) 

This  is  Marsh's  test  for  arsenic.  For  very  delicate 
cases,  the  hydrogen  should  be  dried  by  a  soda-lime  tube. 
Explanation  :  Nascent  hydrogen  reduces  arsenic  trioxide, 
forming  arsine  and  water  : 

As,O3  +  6H2  =  2AsH3  +  3H2O. 

Arsine  (an  invisible  gas)  is  decomposed  by  the  heat  of 
the  hydrogen  flame,  and  the  metallic  arsenic  is  con- 
densed on  the  porcelain. 

Arsine    is    formed    in    vegetable    matters    containing 
arsenic  when  certain  minute  fungi  are  growing  in  them, 
12 


162  ARSENIC    AND    HALOGENS. 

e.g.,  in  arsenical  wall  papers.  The  air  of  a  room  may 
thus  become  contaminated  by  this  intensely  poisonous 
gas.  One  small  bubble  of  the  pure  gas  has  been  known 
to  kill  a  man. 

147.  Arsenic  and  Halogens. 

ARSENIC  TRICHLORIDE  (AsCl;s)  is  a  colourless  volatile 
liquid,  formed  by  the  action  of  sulphuric  acid  on  a  mix- 
ture of  arsenic  trioxide  and  sodic  chloride.  This  explains 
its  presence  in  hydrochloric  acid  prepared  by  means  of 
arsenical  sulphuric  acid.  It  dissolves  in  water,  at  the 
same  time  decomposing  : 

2AsCl3  +  3H2O  =  As2O3  +  6HC1. 

ARSENIC  TRI-IODIDE  (AsI3)  is  used  in  medicine  It 
is  prepared  by  carefully  heating  arsenic  and  iodine  to- 
gether. (In  what  proportion  1).  In  Donovan's  solution, 
it  is  combined  with  mercuric  iodide  (AsI3  -\-  HgI2),  and 
forms  a  clear  colourless  solution.  The  other  halogen 
compounds  are  of  little  importance  (AsBr3,  AsF3,  and 
AsF5). — Clemen's  "  bromide  of  arsenic  "  (so  called)  is  a 
solution  of  arsenic  acid  and  hydrobromic  acid,  prepared 
by  the  action  of  bromine  on  arsenic  trioxide. 

148.  Tests  for  Arsenic.— (See  Experiments  118, 
128,  130,  and  131.) 

1.  To  a  solution  of  arsenic  trioxide  in  dilute  hydrochloric  acid 
add   a  drop   of  ammonio-cupric  sulphate   (prepared  by  adding 
ammonia  to  cupric  sulphate  solution  gradually,  until  the  preci- 
pitate at  first  formed  is  just  redissolved).     A  green  precipitate 
(Scheele's  green)  appears. 

2.  Reimctis  Test. — Put  a  thin  strip  of  bright  copper  in  some 
hydrochloric   acid   solution  of  arsenic  trioxide,   and  boil  (in  a 
t.  t.).     Arsenic  is  deposited  on  the  copper.     Remove  the  copper, 
wash  it  well  (without  rubbing),  dry  it  by  holding  it  over  a  flame 


QUESTIONS    AXI)    EXERCISES.  163 

(not  too  near),  put  it  in  a  narrow  dry  t.  t.  and  heat  it.  Arsenic 
and  the  trioxide  are  sublimed  upon  the  tube.  Try  this  test 
with  some  green  wall  papers,  by  digesting  them  with  strong 
hydrochloric  acid,  diluting  with  water,  and  then  using  the  solu- 
tion as  directed  above. 


QUESTIONS  AND  EXERCISES. 

I.  Compare  arsenic  and  phosphorus  (a)  as  to  physical  proper- 
ties, and  (b)  as  to  their  compounds. 

'2.   Show  the  analogy  between  oxygen  and  sulphur,  as  seen  in 
the  compounds  of  arsenic. 

3.  In  what  proportion  should  ferric  chloride  and  sodic  car- 
bonate be  mixed  so  as  to  produce  ferric  hydroxide  : 

Fe2Cl6  +  3Na2CO    +  3H2O  =  Fe2(OH)6  +  GNaCl  +  3CO2. 

4.  What  data  have  you  for  concluding  that  arsenic  is  triad 
and  pentad  ? 

5.  How  much  mispickel   will   (theoretically)  give   10  Ibs.   of 
arsenic. 

6.  In  what  respect  does  the  molecule  of  arsenic  deviate  from  the 
general  rule  ?    What  other  element  shows  the  same  deviation  ? 

7.  Is  arsenic  a  metal  or  a  non-metal  ? 

8.  What  is  the  meaning  of  "  isomorphous  "  ? 

9.  In  testing  for  arsenic  with  hydric  sulphide  the  solution  is 
made  acid.     Why  ? 

10.  In  Marsh's  test  oxidising  agents  must  be  absent.     Why  ? 

I 1 .  Balance  the  following  equations  : 

(1)  As  4-  HNO3  =  H3AsO4  +  NO2  +  H2O. 

(2)  As2O3  +  HNaO  =  Na3AsO3  +  H2O. 

(3)  NH4OH  +  Fe2Cl6  =  NH4C1  +  Fe2(OH)6. 

(4)  Na2HAs04  +  AgN03  =  Ag3AsO4  +  NaNO3  +  HNO3. 

12.   Write  the  formulas  for  the  following  compounds  :  Sodic 

i('>lti/i/rii-  (tr.'H'iiiiti',   citj>ri<-  Jii/f/ri<-  ur.ir  a  ;/c,   ,W<Y   thiuarsenite,  and 
iiiDninnic  ortkoarsenate. 


164  CARBON. 


CHAPTER    XIII. 


CARBON   AND   ITS   COMPOUNDS. 

H9.  Carbon  (C  iv  =r  12).— This  element  has  three 
allotropic  modifications  :  (1)  diamond,  (2)  graphite,  (3) 
charcoal  (lamp  black,  &c.)  Diamond  and  graphite  are 
crystalline  ;  chai'coal  is  amorphous. 

1.  DIAMONDS  are  found  in  pebbly  deposits  of  ancient 
rivers,  in  South  Africa,  Australia,  South  America,  &c. 
It  is  the  purest  form  of  carbon,  and  its  nature  remained 
long  unknown,  until  it  was  proved  to  be  carbon  by  burn- 
ing it  in  oxygen. — An  impure  black  variety,  carbonado, 
is  used  for  polishing  the  diamonds. — Diamond  is  one  of 
the  hardest  substances  known.     When  pure  it  is  colour- 
less.    Specific   weight  =  3.5   to   3.6.      Its    crystals   are 
modifications  of  the  cube.     Its  great  lustre  is  due  to  its 
strong  refractive  power  on  light.     At  a  high  temperature 
it  burns  in  air  or  oxygen  forming  carbon  dioxide  (COo). 
It  has  never  been  prepared  artificially,  although   many 
attempts  have  been  made. 

2.  GRAPHITE  is  found  in  lumps  in  granite,  &c.,  well 
crystallised ;  or  in  obscurely  crystalline  masses  of  plum- 
bago, or  black  lead.     The  best    black  lead   is   found    in 
Cumberland,  England.     Graphite  can  be  artificially  pre- 
pared by  crystallising  carbon  from  molten  iron.    It  often 
occurs  in  pig  iron.      It  is  not  pure  carbon,   but  leaves 
considerable    ash    on  burning. — It   is  a  black   metallic- 


CARBON.  165 

looking  substance,  greasy  to  the  touch  ;  specific  weight 
=  2  to  2.6.  It  burns  with  great  difficulty  at  a  high 
temperature  to  form  carbon  dioxide.  It  is  used  for  lead 
pencils,  crucibles,  for  polishing  gunpowder  grains,  <fec., 
and  as  a  lubricant. 

3.  AMORPHOUS  CARBON.  —  Lamp-black,  gas-carbon, 
coke,  charcoal,  and  animal  black  are  impure  amorphous 
carbon.  Lamp-black  is  prepared  by  burning  turpentine, 
&c.,  with  a  scant  supply  of  air  and  collecting  the  soot 
on  woollen  cloths.  It  consists  of  particles  of  carbon 
covered  with  tar.  It  is  used  in  the  preparation  of  Indian 
ink,  black  paint,  and  printers'  ink,  and  to  give  a  grey 
shade  to  calico. — Gas  carbon  gathers  in  the  upper  part 
of  the  retorts  in  which  coal  is  distilled  in  the  manufac- 
ture of  coal  gas.  It  is  very  hard  and  resonant,  and  con- 
ducts electricity.  It  is  used  as  the  negative  element  in 
electric  batteries,  and  in  the  manufacture  of  micro- 
phones.— Coke  remains  in  the  bottom  of  the  gas  retorts 
when  the  process  of  distilling  is  complete.  It  is  hard  and 
dense,  and  burns  (at  a  high  temperature)  without  smoke 
and  giving  an  intense  heat.  It  is  used  in  smelting  opera- 
tions and  as  fuel  in  engines. — Charcoal  is  prepared  by  the 
destructive  distillation  of  wood  or  bones.  In  the  latter 
case  it  is  called  animal  charcoal,  or  bone  black. 

Experiment  132.—  Heat  a  small  bit  of  dry  wood  in  a  mat- 
trass,  applying  a  match  to  the  open  end.  A  combustible  yas  is 
evolved,  tar  and  water  gather  on  the  sides,  and  charcoal  re- 
mains as  the  skeleton  of  the  wood.  Try  the  same  experiment 
with  a  bit  of  bone,  horn,  or  dried  meat.  Similar  results  are  ob- 
tained. 

Charcoal  has  the  power  of  absorbing  gases  and  causing 
their  oxidation  when  they  are  oxidisable.  A  piece  of 


1GG  CHARCOAL. 

charcoal  made  from  the  shell  of  the  cocoa-nut  will  absorb 
171  times  its  volume  of  ammonia  gas  and  100  times  its 
volume  of  hydric  sulphide.  It  must  first  be  heated  to 
expel  the  gases  already  condensed  in  its  pores.  On  ac- 
count of  this  property  it  is  used  as  a  filter  for  air,  being 
made  into  so-called  "respirators."  Evidently  the  char- 
coal in  these  respirators  should  be  often  renewed,  or 
taken  out  and  heated  i-ed-hot,  in  order  to  expel  the  con- 
densed gases. 

Experiment  133.— Fill  a  filter  paper  (placed  in  a  funnel) 
about  two-thirds  full  of  bone-black.  Colour  some  water  with 
indigo,  heat  it,  and  pour  it  on  the  filter. 

Charcoal  has  the  property  of  extracting  colouring  mat- 
ters from  solution.  Bone-black  is  used  for  this  purpose 
in  sugar  refining. 

Experiment  134- — Fill  a  filter  with  fresh  ground  charcoal, 
and  pour  some  hydric  sulphide  water  on  it.  If  the  smell  has 
not  been  destroyed  when  the  water  rans  through  pour  it  on  again. 

Charcoal  has  the  power  of  extracting  offensive  animal 
and  vegetable  matters  from  water.  It  is  used  for  filter- 
ing water.  Filters  should  be  renewed  very  often.  The 
charcoal  should  either  be  replaced  or  be  taken  out  and 
heated  red-hot  to  destroy  the  substances  extracted  from 
the  water. — Charcoal  poultices  are  used  for  purifying 
foul  wounds.  They  are  very  efficacious. — Coal  is  impure 
amorphous  carbon.  It  is  the  remains  of  primeval  forests 
which  have  undergone  a  slow  process  of  decay.  This 
process  is  now  going  on  in  peat  bogs.  Anthracite  is  very 
pure,  containing  as  high  as  94  °/0  carbon.  Bituminous, 
or  soft  coal  contains  much  hydrogen,  oxygen,  and  nitro- 
gen, and  only  up  to  75  %  carbon.  Lit/nite  is  coal  in 


COMPOUNDS    OF    CARBON.  167 

the  process  of  formation.     It  contains  up  to  70  %  carbon. 
Jet  is  a  hard  variety  of  coal  which  takes  a  good  polish. 

150.  Compounds  Of  Carbon. — Carbon  occurs  in 
nature  combined  as  well  as  free.  As  shown  in  Experi- 
ment 13'2,  it  forms  an  essential  constituent  of  animal 
and  vegetable  bodies.  In  fact,  it  is  the  element  of  organic 
tissues. — It  is  found  in  vast  quantities,  as  carbonates, 
e.  g.,  calcic  carbonate  (limestone,  marble,  chalk,  &c.), 
magnesic  carbonate  (in  mountain  limestone) ;  also  in  the 
air  as  carbon  dioxide  (CO.,) ;  in  marsh  gas  (CH4),  petro- 
leum, shale,  ike. — The  number  of  the  compounds  of 
carbon  is  so  vast  that  a  special  branch  of  chemistry, 
Organic  Chemistry,  is  devoted  to  their  considera- 
tion. Many  of  them  exist  already  formed  in  the 
bodies  of  plants  and  animals,  and  in  the  crust  of 
the  earth ;  but  many  are  the  product  of  the  labora- 
tory. The  great  number  of  the  carbon  compounds  is 
due  to  the  property  which  carbon  atoms  have  of  unit- 
ing with  each  other  so  as  to  form  groups  of  carbon 
atoms  capable  of  acting  as  the  nuclei  of  molecules. 
Thus,  carbon  and  hydrogen  can  unite  in  a  great 
number  of  different  proportions,  the  simplest  com- 
pound being  that  having  one  carbon  atom  in  the  inole- 

H  H       H 

I  I          I 

cule,  H — C — H,  the  next  containing  two,  H — C — C — H, 
I  I       I 

H  H       H 

H       H       H 

I          I          I 

the    next    three,    H — C — C — C — H,    and    so    on.       The 
i        I        I 
H       H       H 

hydrogen     atoms    can     be     replaced     by   other    atoms 


168  CARBON    DIOXIDE. 

and    groups,    and    thus    derivatives    :m>    formed.      Tims, 
H  II 

H— C— 01,  H— 0— OH,  &c. 

I  I 

H  H 

151.  Sources  of  Carbon  Compounds.— Plants 

and  animals  supply  compounds  of  carbon  ready  formed, 
and  the  most  fruitful  natural  sources  are  the  remains  of 
organised  bodies.  In  fact,  if  we  exclude  some  of  the 
naturally  occuring  carbonates,  it  may  be  stated  that  all 
carbon  compounds  not  of  artificial  origin  are  the  products 
of  vegetable  or  animal  organisms. 

Petroleum  supplies  a  great  number  of  compounds  of 
carbon  and  hydrogen,  mostly  belonging  to  the  class  of 
paraffins  (CnH2n  +  2)- — By  the  distillation  of  coal,  coal 
tar  is  obtained.  This  once  waste  product  has  developed 
during  the  last  30  years  into  an  almost  inexhaustible 
source  of  new  and  valuable  compounds.  From  it  are 
manufactured  the  aniline  and  other  beautiful  colours, 
artificial  essences  and  perfumes,  and  substances  which  in 
some  degree  serve  as  substitutes  for  quinine,  &c. 
Lately,  saccharine,  said  to  be  220  times  as  sweet  as 
sugar,  has  been  made  from  coal-tar  products. — Bone-oil 
and  wood-tar  are  sources  of  organic  compounds. 

CARBON   AND  OXYGEN. 

152.  Carbon  Dioxide,  GO,.  —  Commonly    called 
carbonic  acid. 

OCCURRENCE. — Free  in  the  air  and  in  the  crust  of  the 
earth,  issuing  from  fissures,  as  in  the  celebrated  Grotto 
del  Cane.  It  is  found  combined  injhe  carbonates.  It 


CARBON    DIOXIDK.  169 

is  ;i  product  of  the  combustion,  decay,  and  fermentation 
of  organic  substances.  It  is  also  given  out  by  animals 
in  respiration. 

PREPARATION. — Experiment  135.— Put  some  lumps  of  mar- 
ble or  limestone  in  an  8-ounce  bottle  provided  with  a  gas-delivery 
tube  bent  twice  at  right  angles  so  as  to  collect  a  heavy  gas  by 
displacement  of  air.  Fill  the  bottle  about  one-third  with  dilute 
hydrochloric  acid,  and  collect  several  bottles  of  the  gas  by  dis- 
placement of  air,  covering  them  with  pieces  of  glass.  Try  to 
light  the  gas  as  it  issues  from  -the  tube.  Bubble  some  of  it 
through  blue  litmus.  Allow  it  to  biibble  through  lime  water. 
Evaporate  the  solution  left  in  the  bottle  and  get  crystals  of 
calcic  chloride  (CaCl2.6H"2O). 

CaCO3  +  2HC1  =  CaCl2  +  H2O  +  CO2. 

Carbon  dioxide  can  also  be  prepared  by  heating  lime- 
stone to  a  red  heat : 

CaCO3  =  CaO  +  CO2. 

PROPERTIES. — A  heavy,  invisible  gas,  of  slightly  pun- 
gent smell,  and  sharp,  sour  taste.  Specific  weight  = 
(Calculate.)  It  is  not  a  strong  poison,  but  has 
a  narcotic  action,  and  does  not  support  animal  life.  It  can 
be  liquefied  at  — 78.2°  C.  under  the  atmospheric  pressure. 
At  higher  temperatures  it  can  be  liquefied  by  increased 
pressure ;  and  when  the  pressure  is  removed  the  liquid 
boils,  part  of  it  becoming  frozen  into  a  snow-like  solid. 
(Explain  this.)  It  is  somewhat  soluble  in  water;  at 
0°  C.,  1.8  vols.  in  1  of  water;  at  15°  C.,  1  vol  in  1. 
The  solution  is  slightly  acid,  and  probably  contains  car- 
bonic acid  (H.,CO3). 

Experiment  136- —  Put  alighted  match  into  a  bottle  of  car- 
bon dioxide.  Put  another  bottle,  mouth  downwards,  in  recently 
boiled  cold  water,  and  let  it  stand  a  few  minutes  with  occasional 


170  rAHI'.ONIC    A(!ID. 

shaking.  Try  the  same  experiment  with  solution  of  caustic 
soda.  Pour  a  bottle  of  gas  into  a  bottle  tilled  with  air,  testing 
with  a  lighted  match. 

Experiment  137. — Bubble  air  from  the  lungs  through  clear 
lime  water  by  means  of  a  glass  tube. 

Carbon  dioxide  is  absorbed  by  strong  bases,  with  . 
which  it  forms  salts,  the  carbonates.  It  is  being  con- 
stantly given  off  into  the  air  from  the  bodies  of  animals, 
but  the  quantity  in  the  air  does  not  increase.  Vege- 
tables use  it  as  fast  as  animals  excrete  it.  Carbon  diox- 
ide is  present  in  the  air  to  the  extent  of  about  4  volumes 
in  10,000.  In  poorly  ventilated  rooms  the  proportion 
much  exceeds  this,  and  the  relative  quantity  of  gas  is 
a  good  measure  of  the  purity  of  the  air.  The  danger, 
however,  is  not  so  much  from  carbon  dioxide  as  from 
other  waste  products  of  the  body  expired  with  the  air 
and  exhaled  from  the  general  surface  of  the  body.  It  is 
these  which  give  the  stale  smell  to  a  poorly  ventilated 
room. — On  account  of  its  high  specific  weight  carbon 
dioxide  collects  in  depressions  such  as  wells  and  cellars. 
Accidents  often  occur  from  workmen  descending  into 
wells  and  brewery  vats  filled  with  the  gas.  A  good  pre- 
caution is  to  lower  a  light  before  descending,  although 
air  which  will  support  combustion  often  contains  enough 
carbon  dioxide  to  produce  death. 

153.  Carbonic  Acid  and  Carbonates.  —  The 

acid  is  not  known  except  in  solution.  It  is  dibasic,  and 
very  weak,  not  forming  salts  with  the  weaker  bases, 
such  as  ferric  hydroxide,  aluminic  hydroxide,  &c. — There 
are  two  series  of  carbonates,  (1)  normal,  e.g.,  Na.2CO3, 
and  (2)  acid,  e.g.,  NaHCO3.  The  carbonates  of  the 
alkali  metals  (K.,COy.Na.;CO3,  &c.)  are  soluble  in  water; 


i  \i!i:o\   MDNOXIDK.  171 

all  other  normal  carbonates  arc  insoluble.  But  many  of 
the  insoluble  normal  carbonates  combine  with  carbonic 
acid  to  form  unstable,  sparingly  soluble  acid  salts.  Thus, 
calcic  carbonate  (CaCO.,),  magnesic  carbonate  (MgCO3), 
ferrous  carbonate  (FeC03).  &c.,  dissolve  in  water  con- 
taining carbonic  acid.  The  solubility  is  increased  by 
pressure  of  the  gas. 

Tests. — 1.   Carbon  dioxide  renders  lime  water  turbid: 
Ca(OH)2  +  CO2  =  CaCO3  +  H2O. 

2.  Carbonates  effervesce  with  hydrochloric  acid.  (Try  with 
several  of  the  carbonates. )  If  the  evolved  gas  be  poured  into  a 
t.  t.  containing  clear  lime  water,  it  renders  it  turbid  on  shaking 
up. 

154.  Carbon  Monoxide,  CO. — Also  called  carbonic 
oxide. 

PREPARATION.— Experiment  138.— Carefully  heat  a  little 
dry  powdered  potassic  ferrocyanide  (K4FeC0Nn)  with  about  10 
times  its  weight  of  concentrated  sulphuric  acid,  and  apply  a  light 
to  the  mouth  of  the  tube.  A  gas  is  evolved  which  burns  with  a 
lambent  blue  flame  : 

K4FeC0N6  +  6H2S04  +  6H20  ----  2K2SO4  + 

3(NHJ2SO4   +   FeSO4    +   6CO. 

Carbon  monoxide  is  also  formed  when  oxalic  acid 
(C.,H.,O4)  is  heated  with  sulphuric  acid  : 

C2H2O4  =  CO  +  CO2  +  H2O. 

It  is  formed  when  carbon  burns  in  a  scanty  supply  of 
air,  as  in  the  inner  parts  of  a  fire. 

PROPERTIES. — A  colourless  gas,  of  very  faint  odour  ; 
s|>.  \vt.  =  0.1)66.  It  burns  with  a  pale,  blue  flame 


172  CARBON    BISULPHIDE. 

(CO  +  O  -  CO.,),  as  seen  OH  the  top  of  a  coal  fire.  It 
can  be  condensed  to  a  liquid  at  — 139.5°  by  a  pressure 
of  35.5  atmospheres.  It  is  sparingly  soluble  in  water 
(3  in  100) ;  easily  in  ammoniacal  solution  of  cuprous 
chloride  (Cu2Cls).  It  is  a  deadly  poison,  combining  with 
the  hpemolglobin  of  the  blood  and  thus  preventing  its 
aeration.  As  it  often  escapes  combustion  in  coal  and 
charcoal  fires,  it  contributes  greatly  to  the  poisonous 
condition  of  ill-ventilated  rooms.  Open  charcoal  fires 
(braziers)  are  especially  dangerous. 

155.  Carbon  Bisulphide,  CS,.— 

PREPARATION. — By  passing  sulphur  vapoxir  through 
long  tubes  filled  with  red  hot  charcoal. 

PROPERTIES. — A  mobile  colourless  liquid,  of  pleasant 
ethereal  smell  when  pure,  but  generally  having  a  disgust- 
ing odour  due  to  impurities.  Sp.  wt.  =  1.29.  It  does 
not  mix  with  water. 

Experiment  139. —  Pour  a  little  carbon  bisulphide  into  a  t.  t. 
of  water.  Note  the  smell.  Shake  up  with  the  water.  Heat  a 
few  drops  of  carbon  bisulphide  gently  in  another  t.  t.  It  boils 
readily.  Put  a  light  to  the  mouth  of  the  t.  t. 

Carbon  bisulphide  boils  at  46°  C.,  forming  a  heavy 
vapour,  very  explosive  when  mixed  with  air  : 

CS2  +  302  =  CO2  +  2SO2. 

It  evaporates  quickly  when  exposed  to  the  air  and  lowers 
the  temperature.  (Try  a  drop  on  the  hand.)  It  mixes 
in  all  proportions  with  ether  and  alcohol  ;  arid  dissolves 
fats,  oils,  india-rubber,  phosphorus,  iodine,  bromine,  &c. 


HYDROCARBONS.  173 

A  solution  of  india-rubber  in  carbon  bisulphide  is  used 
for  cementing  rubber  goods.  Its  vapour  mixed  with 
nitrogen  dioxide  (NO)  burns  with  a  dazzling  white 
light  used  in  photography.  —  Carbon  bisulphide  is 
poisonous.  It  is  antiseptic,  and  is  used  for  preserving 
meat,  <fec. 

Sulphocarbomc  arid,  H2C8a. — Also  called  thiocarbonic  acid. 
This  acid  and  its  salts  are  analogous  to  carbonic  acid  and  the 
carbonates. 

Potasmc  xnlphocarhunatp,  K.^CSg,  can  be  prepared  by  dissolving 
carbon  bisulphide  in  solution  of  potassic  sulphide : 

K2S  +  CS2  =  K2CS3. 

From  this,  sulphocarbonic  acid  can  be  separated  as  an  unstable 
oily  liquid  by  the  action  of  hydrochloric  acid  : 

K2CS3  +  2HC1  =  2KC1  +  H2CS3. 

It  soon  decomposes  into  carbon  bisulphide  and  hydric  sulphide : 
H2CS3  =  H2S  +  CS2. 

CARBON  AND  HYDROGEN. 

156.  Hydrocarbons. — The  number  of  these  is  so 
great  that  only  a  very  small  proportion  of  them  can  be 
described. — The  simplest  compound  of  carbon  and  hydro- 
gen is  that  containing  1  carbon  atom  and  4  hydrogen 

H 

I 

atoms  in  the  molecule,  H — C — H.     A  molecule  contain- 

I 

H 

ing  2  carbon  atoms  can  have  at  most  6  hydrogen  atoms,  for 
each  of  the  carbon  atoms  has  4  bonds,  and  1  of  these  is 


174  PARAFFINS. 

employed  in  holding  the  carbon  atoms  themselves  together. 

H       II 

The  graphic  formula  for  this  molecule  is  H — C  — C — H. 

I       I 

H       H 
H       H       H 

The  next  would  be  H — C — C — C — H,   and   so   on,    the 

•I       i       I 

H       H       H 

molecules  increasing  regularly  by  CEL, ;  so  that  a  general 
formula,  CNH.2N  +  2,  can  be  written.  Very  many  mem- 
bers of  this  series  are  known.  Most  of  them  are  found 
in  petroleum  oil  and  in  the  gas  accompanying  it.  It  is 
called  the  Paraffin  Series  : 

Marsh  gas,  or  methane.. .  .  CH4  Gas 

Ethane C2H6  ...    .Gas 

Propane C3H8  Gas 

Butane C4Ha  0  Boils  at    1°  C. 

Pentane C5H12  .....        "       38°  0. 

Hexane C6H14  .....        "       70°  C. 

&c.  &c. 

The  different  members  of  this  series  resemble  each 
other  in  chemical  properties.  Paraffin  oil  (kerosene)  is 
a  mixture  of  the  liquid  members.  Paraffin  wax,  used  in 
making  candles,  is  a  mixture  of  solid  paraffins,  the  higher 
members.  These  substances  are  obtained  by  the  frac- 
tional distillation  of  petroleum,  which  is  a  product  of  the 
decomposition  of  primeval  forests. 

Besides  the  paraffin  series  there  are  several  other  series 
of  hydrocarbons  containing  a  less  proportion  of  hydro- 
gen, the  Olefines  having  the  general  formula  CNH.,N, 
—  (C2H4,  C3HC,  &c.) ;  the  Acetylene  Series,  CNH.,N_.j, 


MARSH    GAS.  175 

—  (C.,H2,  ttc.),  and  so  on.  In  each  of  these  series  the 
difference  between  consecutive  members  is  GET,,  and  the 
members  of  one  series  have  a  general  likeness  to  each 
other  in  chemical  properties. 

l<r>7.  Marsh  Gas,  or  Methane,  CH4. 

OCCURRENCE. — In  the  mud  of  marshy  and  other  stag- 
nant pools,  rising  in  bubbles  when  the  mud  is  stirred. 
These  bubbles  can  be  collected  in  a  bottle  by  displace- 
ment of  water,  and  are  found  to  be  combustible.  It  is 
also  found  along  with  petroleum  in  vast  quantities. 
The  natural  and  artificial  gas  springs  of  America  and 
Asia  give  off  great  quantities  of  a  mixture  of  hydrogen, 
methane,  and  ethane,  with  small  quantities  of  other 
hydrocarbons.  In  some  parts  of  Pennsylvania,  this 
natural  gas  is  used  for  illuminating  and  other  purposes. 
The  sacred  fires  of  Baku,  in  Persia,  have  been  burning 
for  ages.  Marsh  gas  is  the  dreaded  "  fire  damp  "  of  coal 
miners.  It  is  also  present  in  considerable  quantity  in 
coal  gas. 

PREPARATION.— Experiment  140.— Mix  well  one  part  dried 
sodic  acetate  (NaC9HaOx)  with  four  parts  soda-lime  (a  mixture  of 
sodic  and  calcic  hydroxides),  and  heat  in  a  hard  glass  t. 1.,  collect- 
ing the  gas  in  the  usual  way  over  water.  A  mixture  of  lime  and 
sodic  carbonate  remains : 

NaC2H3O2  +  NaOH  =  Na2CO3  +  CH4. 

PROPERTIES. — A  colourless  gas,  inodorous,  and  harm- 
less when  breathed.  Sp.  wt.  =  0.555.  It  burns  in  air 
with  a  pale  blue  flame,  and  does  not  support  ordinary 
combustion. 

Experiment  141- — Thrust  a  lighted  match  into  a  jar  of  the 
gas  held  upside  down.  Pour  the  gas  from  another  jar  upwards 
into  a  jar  filled  with  air. 


176  MARSH    GAS. 

Methane  is  very  explosive  when    mixed  with  air  or 

oxygen  : 

CH4  +  2O2  =  CO2  +  2H2O. 

Frightful  explosions  occur  in  coal  mines.  For  the  pro- 
tection of  miners  Sir  Humphrey  Davy  invented  the 
"  Davy  Lamp,"  in  which  the  flame  is  surrounded  by  wire 
gauze.  The  metal  conducts  away  the  heat  of  the  flame 
rapidly,  and  thus  keeps  the  explosive  mixture  next  to  it 
below  its  point  of  ignition.  The  flame  cannot  pass 
through  the  gauze. 

Experiment  142. — Hold  a  piece  of  wire  gauze  above  a  jet  of 
gas,  apply  a  light  above  the  gauze  and  observe  that  the  flame 
does  not  pass  below. 

Methane  is  sparingly  soluble  in  water.  It  is  very 
stable,  like  all  the  paraffins  (parum  affinis  =  having 
little  affinity),  resisting  the  action  of  nitric,  sulphuric,  and 
hydrochloi'ic  acids,  and  of  other  vigorous  chemical  sub- 
stances. It  cannot  be  got  to  unite  with  any  element  or 
compound,  except  by  losing  one  or  more  of  its  atoms  of 
hydrogen.  The  hydrogen  can  be  replaced  in  fourths  by 
the  action  of  chlorine  : 

Monochlonnethane. 

CH4       +     Cla  =  CH3C1    +   HC1. 

Dichlormethane. 

CH3C1    +   C12  =  CH2C12  +  HC1. 

Trichlormethane. 

CH2C12  +   C12  =  CHC13    +  HC1. 

Tetraehlormethane. 

CHC13    +    C12   =  CC14       +   HC1. 

The  molecule  of  marsh  gas  is  incapable  of  taking  up 
atoms  except  by  replacement.  It  is  a  saturated  com- 
pound, i.e.,  the  carbon  is  saturated  with  hydrogen.  The 
paraffins  are  the  saturated  hydrocarbons. 


CHLOROFORM.  177 

1 58.  Chloroform,  CHC13. — Formerly  mentioned  as 
trichlormethanf. 

PREPARATION. — By  distilling  at  a  gentle  heat  aqueous 
solution  of  bleaching  powder  with  common  alcohol 
(C,H60). 

Experiment  143- — Rub  together  one  part  bleaching  powder 
and  four  of  hot  water,  add  one-tenth  part  strong  alcohol  and 
warm  gently.  Note  smell,  &c. 

PROPERTIES. — A  colourless  liquid,  of  pleasant  smell, 
and  burning  taste.  It  evaporates  quickly  when  placed 
on  the  hand  (Try  it) ;  and  boils  at  61°  C.  Sp.  wt.  =  1 .525. 

Experiment  144- — Put  two  or  three  drops  of  chloroform  in  a 
t.  t.  of  water  and  shake.  It  dissolves,  giving  its  odour  to  the 
water.  Add  a  considerable  quantity  and  try  to  dissolve  it.  Try 
with  alcohol  instead  of  water  ;  also  with  ether. 

Chloroform  is  slightly  soluble  in  water,  but  mixes  in 
all  proportions  with  ether  and  alcohol.  It  is  a  good 
solvent  for  iodine,  bromine,  phosphorus,  &c. — (Is  it  in- 
flammable 1) 

When  breathed  as  a  vapour  chloroform  causes  insensi- 
bility. It  is  an  anaesthetic. — (Sir  James  Simpson,  1848.) 
It  is  now  used  in  surgical  operations  very  extensively, 
although  ether  is  preferred  by  many  surgeons.  Its  low 
boiling  point  renders  it  easy  to  administer  as  a  vapour. 
It  is  also  administered  in  the  liquid  form  as  a  medicine. 

IMPURITIES  IN  CHLOROFORM. — Great  care  should  be 
taken  in  selecting  chloroform  for  use  as  a  medicine  or  an 
anaesthetic.  The  commercial  article  often  contains  dele- 
terious substances,  e.g.,  hydrochloric  acid,  chlorine,  and 
chlorides  of  other  organic  compounds.  Pure  chloroform 
does  not  turn  blue  litmus  red  or  give  a  precipitate  with 
13 


178  ETHYLENK. 

argentic  nitrate  (impurities,  HC1  and  01);  it  does  not 
colour  a  mixture  of  potassic  bichromate  and  dilute  sul- 
phuric acid  green  (alcohol,  &c.) ;  it  is  not  coloured  brown 
by  potassic  hydroxide  or  sulphuric  acid  ;  it  does  not  cause 
bright  sodium  to  tarnish,  even  when  heated  (C2H4C12.  &c.). 
If,  when  allowed  to  evaporate  on  a  clean  watch  glass, 
chloroform  leaves  a  strongly  smelling  residue,  the  alcohol 
from  which  it  was  prepared  contained  fusel  oil. 

Experiment  145- — Try  these  tests  with  samples  from  the 
druggists. 

lodoform  (CHI3)  is  a  yellow  crystalline  solid  prepared  by 
adding  iodine  to  a  mixture  of  common  alcohol  and  solution  of 
sodic  carbonate  heated  to  60°.  It  is  insoluble  in  water,  but 
soluble  in  alcohol  and  ether. 

159.  Ethylene,  C.2H4.— Also  called  "olefiant  gas.' 
Is  a  constituent  of  coal  gas. 

PREPARATION. — By  the  action  of  potassic  hydroxide  on 
monochlorethane  (C2H5C1)  : 

C2H5C1  +  KOH  =  C2H4  +  KC1  +  H2O. 
Or,  by  heating  alcohol  with  strong  sulphuric  acid  : 
C2H60  =  C2H4  +  H20. 

PROPERTIES. — It  is  the  first  member  of  the  oleftne  series. 
It  is  a  colourless  gas,  distinguished  from  the  paraffins  by 
the  facility  with  which  it  unites  directly  with  other  sub- 
stances, without  losing  hydrogen.  Thus,  it  unites  with 
an  equal  volume  of  chlorine  to  form  chloride  of  ethylene 
(C2H4C12\  an  oily  liquid  : 

C2H4  +  C12  =  C2H4C12. 
Its  molecule  is  unsaturated,  and  we  express  this  fact  by 

H       H 

I         I 

the  graphic  formula,  H — C  =  C — II,  in  which  the  carbon 


ISOMRKISM.  179 

atoms  are  represented  as  united  by  two  bonds.      One  of 
these  can  be  divided  so  as  to  receive  2  atoms  of  chlorine, 

H       H 

I         I 

H — C  -  C — H.    The  ole/tnes  are  all  unsaturated  compounds. 

I       I 
ci     ci 

160.  Isomerism. — -As  has  been  mentioned,  methane 
is  capable  of  entering  into  those  chemical  actions  only  in 
which  its  hydrogen  atoms  are  replaced  by  other  atoms  or 
compound  radicals.  Thus,  an  atom  of  chlorine  replaces 
one  of  hydrogen  to  form  monochlormethane.  This  re- 
placement has  been  made  in  a  great  many  different  ways, 
and  the  same  substance  is  always  obtained.  It  has  been 
concluded  that  the  four  hydrogen  atoms  are  exactly  alike 
in  relative  position,  so  that  it  does  not  matter  which  is 
replaced.  It  is  very  important  to  get  this  idea  clearly 
at  the  outset.  It  has  been  found  also  that  only  one  sub- 
stitution product  of  ethane  having  the  formula,  C2H5C1, 
can  be  obtained ;  and  it  is  therefore  concluded  that 

H     H 

I       I 

the  six  hydrogen  atoms  in  ethane,  H — C — 0 — H,  are  alike 

I        ! 
H       H 

in  position.  But,  when  a  second  atom  of  hydrogen  is 
replaced  by  chlorine,  two  substances  are  obtained,  identi- 
cal in  chemical  composition  (C.,H4C1.,),  but  differing  in 
properties.  One  of  these  is  the  substance  mentioned  in 
in  Art.  151  as  ethylene  chloride.  The  only  other  differ- 
ent arrangement  of  the  atoms  is  represented  as  follows  : 

H     H 

II 
H — C  —  C — ci,  both   chlorine    atoms    being    attached    to 

I       i 

H    ci 


180  ACETYLENE. 

the  same  carbon  atom.     This  second  compound  can  be 

H 

I 

prepared    from    aldehyde    (H — C — 0  =  O),    by    replacing 

I       I 

H       H 

oxygen  by  chlorine  ;  and  it  is  distinguished  from  the  first 
by  the  name  ethylidene  chloride.  These  two  substances 
are  alike  in  composition,  but  different  in  properties.  This 
is  a  case  of  isomerism. 


161.  Acetylene,  C2H2.  —  This  compound  is  a 
colourless  gas  formed  by  the  incomplete  combustion 
of  gases  containing  hydrocarbons.  It  has  a  strong  dis- 
agreeable smell,  which  is  easily  observed  in  the  smoke 
of  a  candle  just  blown  out.  It  is  produced  in  consider- 
able quantity  in  a  Bunsen  burner  "burning  below." 
(Try  it.)  It  can  also  be  prepared  by  the  action  of  alco- 
holic solution  of  caustic  potash  on  ethylene  chloride  or 
bromide : 

U2H4C12  +  2KOH  =  C2H2  +  2KC1  +  2H2O. 

It  is  poisonous  when  breathed,  uniting  with  the  haemo- 
globin of  the  blood  as  carbon  monoxide  does.  As  it  is 
formed  by  lamps  when  the  flame  is  turned  low,  this 
practice  should  be  discountenanced,  especially  in  sick 
rooms. 

Acetylene  can  be  formed  from  the  elements  by  passing 
a  powerful  current  of  electricity  between  carbon  poles 
in  an  atmosphere  of  hydrogen.  This  is  an  important 
synthesis. 


POTASSIC    FKI!KorYAN]l>K.  ISl 

CAKBON  AND  MTKO<:KN. 

162.  Cyanogen  Compounds. — Nitrogen   is    pre- 
sent in  all  organised  bodies.     If  nitrogenous  matters  such 
as  horn,  flesh,  <fcc.,  be  heated  with  sodium,  a  compound  is 
obtained    composed    of    sodium,    carbon,    and    nitrogen 
(NaCN),  arid  called  sodium  cyanide.     When  nitrogenous 
substances  are  heated  with  potassic  carbonate  (K.,CO3), 
and  iron  scraps,  a  fused  mass  is  obtained  which  yields  on 
lixiviation  a  yellow  crystalline  salt,  potassic  f err ocyanide 
(K4Fe(CN)6).     This  salt  gives   a  fine  blue  colour  (Prus- 
sian blue)  with  ferric  salts,  and  is  the  starting  point  in 
the  preparation   of  the  cyanogen  (blue-generating)  com- 
pounds.    All  these  compounds  contain  the  monad  radical 
— ON,  cymogen,  which  can  be  set  free,  but,  like  an  atom, 
unites    immediately   with   another    to    form   a   molecule 
(CN).,.     The  substance  cyanogen  (C.,N2)  can  be  prepared 
by  heating  'mercuric  cyanide  (Hg(CN).>).     It  is  a  very 
poisonous  gas,  having  chemical  properties  like  those  of 
chlorine  ;  and  its  compounds  with  metals  are  called  cyan- 
ides, e.g. : 

KCN,  Fe  CN)2,  Ca(CN)2,  &c. 

163.  Potassic  Ferrocyanide.— K4Fe(CN)6.3H,O. 

Commercially  known  as  yellow  prussiate  of  potash. 

PREPARATION. — Fuse  refuse  animal  substances  with 
potassic  carbonate,  and  lixiviate  the  fused  mass.  A  solu- 
tion containing  potassic  cyanide  (KCN)  is  obtained.  To 
this  add  freshly  precipitated  ferrous  carbonate  (FeCO3). 
It  dissolves  : 

6KCN  +  FeC03  =  K4Fe(CN)e  +  K2CO3. 
The  solution  is  evaporated  and  the  two  salts  separated  by 


182  IIVDKOCVAMC    ACID. 

crystallisation,  the  potassic  rnrbonate  being  much  the 
more  soluble  of  the  two. 

PROPERTIES. — A  salt  crystallising  in  large  yellow  pris- 
matic crystals.  (Kxainine  a  crystal,  noting  its  "feel," 
taste,  the  ease  with  which  a  splinter  can  be  split  off,  and 
the  flexibility  of  the  splinter.)  It  is  soluble  in  about  4 
times  its  weight  of  water.  (Dissolve  a  small  quantity  and 
taste  the  solution.)  It  is  not  poisonous,  and  in  large 
doses  acts  as  a  mild  purgative. 

Experiment  146. — Heat  slightly  a  small  crystal  of  potassic 
ferrocyauide  in  a  t.  t.  It  falls  to  a  powder  and  water  is  con- 
densed on  the  side  of  the  tube.  (From  what  source  does  the 
water  come '!)  Heat  more  strongly,  and  the  salt  turns  brown. 
It  has  been  decomposed  into  potassic  cyanide  (KCN),  carbide  <>f 
iron  (FeC2),  and  nitrogen.  (Write  the  equation.)  Keep  the 
residue. 

Potassic  ferrocyanide  belongs  to  the  class  of  double 
cyanides,  and  is  composed  of  potassic  and  ferrous  cyan- 
ides, 4KCN.Fe(CN).,.  But  these  are  so  united  that  no 
ordinary  test  shows  the  presence  of  iron  ;  and  the  salt 
does  not  possess  the  characters  of  a  simple  cyanide.  It 
is  a  salt  of  ferrocyanic  add  (H4Fe(CN)6).  All  soluble 
simple  cyanides  are  deadly  poisons. 

16t.  Hydrocyanic     Acid,     HCK  —  Originally 

called  prussic  acid.  This  is  one  of  the  most  deadly 
and  sudden  poisons  known.  The  vapour  of  the  pure  acid 
causes  almost  instant  death,  and  even  when  largely 
diluted  with  air  it  causes  headache  and  giddiness  for 
some  hours  if  inhaled  for  a  few  seconds.  Unfortunately 
some  persons  are  not  able  to  detect  its  subtle  odour — that 
of  crushed  cherry  stones.  Experiments  with  it  must  be 


CY  AMI  IKS.  183 

made  very  cautiously. — Antidotes, — ammonia  and  chlo- 
rine water. 

PREPARATION. — Distil  2|  oz.  potassic  ferrocyanide  dis- 
solved in  10  oz.  water;  with  1  oz.  sulphuric  acid  diluted 
with  3  oz.  distilled  water.  Receive  the  distillate  in  8 
oz.  distilled  water  until  there  is  17  fluid  oz.  Make  up 
to  20  oz.  with  distilled  water.  This  solution  contains 
2  %  hydrocyanic  acid ;  it  is  the  acidum  hydrocyanicum 
dilntum  of  the  British  Pharmacopoeia. 

2K4Fe(CN)6  +  6H2SO4  = 

Everitt's  Salt. 

6KHS04  +  FeK2.Fe(CN)6  +  6HCK 

Experiment  147. — Add  a  little  water  to  the  residue  from 
Experiment  146,  and  filter.  To  the  nitrate  add  some  very  dilute 
sulphuric  acid,  and  carefully  smell.  The  peculiar  odour  of 
prussic  acid  can  be  noticed.  (What  is  dissolved  in  the  water  ? 
Write  an  equation  showing  the  action  of  sulphuric  acid. )  Try 
solution  of  potassic  ferrocyanide  with  the  diluted  sulphuric  acid. 
Hydrocyanic  acid  is  not  set  free. 

PROPERTIES. — The  pure  acid  is  a  colourless  volatile 
liquid.  The  dilute  solution  has  the  smell  of  the  acid, 
and  is  feebly  acid  in  reaction.  It  gradually  decomposes, 
the  acid  uniting  with  water  and  forming  ammonic  for- 
mate (]STH4.CHO2)  : 

HCN  +  2H2O  =  HC02.NH4. 

Another  decomposition  sometimes  goes  on  in  which  an  in- 
soluble brown  substance,  paracyanogen  (C3N3)  is  formed. 

CYANIDES. — Hydrocyanic  acid  is  a  weak  monobasic 
acid,  very  like  hydrochloric  acid  in  its  chemical  proper- 
ties. The  cyanides  of  the  alkali  metals  (K,  Na,  &c.),  of 
barium  calcium,  and  strontium,  and  mercuric  cyanide 


184  DOUBLE    CYANIDES. 

(Hg(CN).j)  are  soluble  in  water.  The  other  simple 
cyanides  are  insoluble  in  water,  but  dissolve  in  solutions 
of  the  alkaline  cyanides  to  form  double,  cyanides. 

Experiment  148. — To  a  solution  of  argentic  nitrate  add  a 
drop  of  solution  of  potassic  cyanide .  A.  white  precipitate  of 
argentic  cyanide  (AgCN)  is  thrown  down  : 

AgN03  +  KCN  =  KN03  +  AgCN. 

Add  more  potassic  cyanide,  and  the  precipitate  is  dissolved, 
owing  to  the  formation  of  a  soluble  double  cyanide  (AgCN.  KCN). 

There  are  two  classes  of  double  cyanides  :  ( I )  those 
easily  decomposed  by  dilute  acids,  and  from  which  hydro- 
cyanic acid  is  set  free,  and  (2)  those  from  which  hydrocy- 
anic acid  is  not  set  free  by  dilute  acids.  To  (1)  belong  the 
double  cyanides  of  potassium  with  silver,  mercury,  &c. 
To  (2)  belong  the  ferrocyanides,  &c.  All  the  soluble  simple 
cyanides  and  the  double  cyanides  of  class  (1 )  are  deadly 
poisons.  The  double  cyanides  of  class  (2)  are  not  poison- 
ous. 

Tests. — 1.  To  a  solution  of  hydrocyanic  acid,  or  of  a  cyanide, 
add  argentic  nitrate.  A  white  precipitate  of  argentic  cyanide  is 
formed.  This  is  insoluble  in  nitric  acid,  and  sparingly  soluble  in 
ammonia.  (Compare  AgCl. )  ( Why  must  a  considerable  quantity 
of  argentic  nitrate  be  added  before  a  permanent  precipitate  is 
obtained  with  potassic  cyanide  ?) 

2.  To  a  solution  of  the  acid  or  of  a  cyanide  add  a  few  drops  of 
ferrous  sulphate  solution,  a  drop  or  two  of  ferric  chloride  solution, 
and  caustic  soda  (caustic  potash,  potassic  carbonate,  lime  water, 
&c.,  will  answer)  :  warm  gently,  and  acidify  with  hydrochloric 
acid.  A  blue  colour  or  precipitate  is  formed  (Prussian  blue). 

(1)  2KCN  +  FeS04  =  Fe(CN).,  +  K2S04. 

(2)  4KCN  +  Fe(CN)2  =  K4Fe(CN)6. 

(3)  3K4Fe(CN)G  -f  2Fe2Clc  =  12KC1  -f-  Fe4(FeCGN6)3. 


CYANIC    ACID.  185 

According  to  equations  (I)  and  (2)  potassic  cyanide  is  converted 
into  potassic  ferrocyanide.  This  takes  place  best  in  presence  of 
an  alkali.  When  the  alkali  is  neutralised  by  an  acid,  insoluble 
Prussian  blue  is  formed.  Both  potassic  ferrocyanide  and  Prus- 
sian blue  are  harmless. 

3.  To  a  drop  or  two  of  solution  of  potassic  cyanide  add  a  little 
ammonic  sulphide  (yellow),  and  evaporate  to  dryness  in  a  porce- 
lain dish  on  the  water  bath.  Add  a  small  quantity  of  ferric 
chloride  solution.  A  blood  red  colour  appears. 

Explanation, — The  sulphur  in  the  yellow  ammonic  sulphide 
unites  with  potassic  cyanide  to  form  potassic  sulphocyanate 
(KCNS).  This,  when  treated  with  ferric  chloride,  forms  ferric 
tulphocyanate  (Fe.2(CNS)6).  The  object  of  evaporating  is  to  get 
rid  of  the  excess  of  ammonic  sulphide.  (Add  a  drop  of  ammonic 
sulphide  to  solution  of  ferric  chloride. ) 

165.  Cyanic  Acid. — HCNO.  The  potassium  salt  of 
this  acid  is  formed  along  with  potassic  cyanide  by  the 
action  of  cyanogen  gas  on  a  solution  of  caustic  potash  : 

(ON),  +  2KOH  ---.  K(CN)  +  K(CN;O  +  H,O. 

(Compare  with  chlorine.}  It  can  also  be  prepared  by 
oxidising  potassic  cyanide  by  fusing  it  with  manganese 
dioxide  (MnO,) :  KCN  +  O  =  KCNO. 

The  acid  itself  is  of  little  importance,  but  its  ammonium 
salt  (NH4.CNO)  possesses  for  us  a  great  interest.  It 
was  from  this  salt  that  Wb'hler,  in  1828,  first  prepared 
urea  artificially,  and  thus  broke  down  the  barrier  between 
organic  and  inorganic  chemistry.  Previously  to  this,  it 
was  supposed  that  the  chemical  compounds  associated 
with  animal  and  vegetable  life  could  not  be  made  in  the 
laboratory,  and  thus  organic  chemistry  was  a  separate 
branch  of  the  science. 


180  URKA. 

Sulphocyanates. — These  are  salts  of  an  acid  (HOTS)  which  is 
related  to  cyanic  in  the  same  way  as  sulphocarbonic  is  to  carbonic 
acid.  The  potassium  salt  (KCNS)  is  prepared  by  fusing  potassic 
cyanide  with  sulphur.  The  ferric  salt  (Fe^CNS),.)  is  of  a  blood 
red  colour,  seen  when  solutions  of  ferric  chloride  and  potassic 
sulphocyanate  are  mixed.  These  salts  are  also  called  thiocyanates. 

166.  Urea. — CO(NH2)2.  This  substance  is  isomeric 
with  ammonic  cyanate ;  and  if  an  aqueous  solution  of 
the  latter  be  heated  it  is  transformed  into  urea. 

OCCURRENCE. — Urea  is  a  waste  product  of  the  human 
and  other  animal  bodies.  It  is  the  vehicle  by  which  the 
waste  nitrogen  is  carried  out  of  the  body  in  the  urine. 

PREPARATION. —  Experiment  149.  —  Evaporate  a  small 
quantity  of  urine  on  the  water  bath  to  about  one-third  its  bulk. 
Add  to  it  about  twice  its  volume  of  concentrated  nitric  acid. 
Pearly  white  scales  of  nitrate  of  urea  are  precipitated.  Evapo- 
rate another  small  portion  of  urea  to  dryness,  warm  the  residue 
with  alcohol,  filter,  and  evaporate  the  filtrate  carefully  to  a 
small  volume.  On  cooling,  it  deposits  colourless  needle-shaped 
crystals  of  urea. 

Urea  can  be  prepared  artificially  from  potassic  cyanide 
by  oxidising  it  to  cyanate  (KCNO),  from  which  ammonic 
cyanate  is  prepared  by  double  decomposition  with  ammonic 
sulphate  ((NH4),SO4) : 

2KCNO  +  (NH4)2SO4  =  K2SO4  +  2NH4  CNO. 

Strong  solutions  of  the  salts  are  mixed,  and  the  sparingly 
soluble  potassic  sulphate  is  precipitated.  The  solution 
at  first  contains  ammonic  cyanate,  but  this  soon  under- 
goes a  change,  especially  if  heated,  by  which  the  atoms  in 
the  molecule  are  rearranged,  and  urea  results  : 

NH4.CNO  =  CO  ? 

IN  1 1  -. 


A.MIDKS.  187 

PROPERTIES. — Urea  is  a  colourless  crystalline  solid, 
soluble  in  its  own  weight  of  cold  water,  and  in  five  parts 
of  alcohol,  but  nearly  insoluble  in  ether.  It  has  a  cooling 
taste  like  that  of  saltpetre.  It  unites  with  aci^Js  just  as 
ammonia  does,  forming  crystalline  salts. 

Experiment  150- — Heat  a  little  pure  urea  with  solution  of 
caustic  soda  and  note  the  smell,  &c. ,  of  ammonia. 

Urea  is  closely  related  to  ammonia.  It  can  be  prepared 
by  the  action  of  ammonia  on  chloride  of  carbonyl  (COCL), 
a  substance  formed  by  the  direct  union  of  chlorine  and 
carbon  monoxide  : 

COC1,  +  2NH3  =  CO(NH,1,  +  2HC1. 

In  this  action  01  is  replaced  by  the  monad  radical  — NH.,, 
called  amidoyen.  Chloride  of  carbonyl  is  related  to 
carbonic  acid  (CO(OH).,),  which  may  be  looked  upon  as 
the  hydroxide  of  carbonyl.  Oxygen  acids  generally  are 
the  hydroxides  of  acid  radicals  ;  and  most  acid  radicals 
can  be  got  combined  with  amidogen,  thus  forming  a  class 
of  bodies  called  amides  from  their  relation  to  a?«monia. 
Urea  is  carbamide.  Amides  readily  combine  with  water 
to  form  ammonium  salts.  Thus,  when  urea  is  heated 
above  100°  C.,  with  water,  ammonic  carbonate  is  formed  : 

CO(NH,V>  +  2H2O  =  CO(ONH4),. 

This  decomposition  can  be  brought  about  more  easily 
by  heating  with  alkalis  or  acids.  (Explain  Experiment 
150.)  It  goes  on  at  the  ordinary  temperature  during 
the  fermentation  of  mine. 

Tests. — 1.  Evaporate  the  liquid  to  small  bulk,  and  treat  with 
concentrated  nitric  acid  as  in  Experiment  149.  The  crystals  of 
urea  nitrate  (CO(NH,).,.HN08)  are  quite  easily  recognized. 


188  URIC    ACID. 

2.  Evaporate,  and  add  a  strong  solution  of  oxalic  acid 
(C2H2O4).  Crystals  of  oxalate  of  urea  are  formed  (2CO(H,,N)o. 
C,HSOJ. 


ESTIMATION  OF  UREA. — The  quantity  of  urea  in  urine 
can  be  readily  determined  by  the  Davy-Knop  method, 
which  depends  on  the  fact  that  urea  is  decomposed  by 
sodic  hypobromite,  and  all  its  nitrogen  set  free  : 

CO(NH2)2  +  3NaOBr  =  CO2  +  N2  +  3NaBr  +  2H2O. 

This  decomposition  is  brought  about  in  an  apparatus 
provided  with  a  graduated  tube  in  which  the  nitrogen  is 
collected  and  measured.  From  the  volume  of  nitrogen 
obtained  from  a  given  quantity  of  urine,  the  quantity  of 
urea  can  be  calculated. 


167.  Uric  Acid. — C5H4N4O;?.  This  substance  is 
found  in  small  quantities  in  Imman  urine,  but  forms  the 
principal  part  of  the  nitrogenous  excrement  of  birds  and 
reptiles.  It  occurs  also  in  the  human  body  as  "  chalk 
stones  "  in  gout,  and  urinary  calculi.  Guano  is  the  ex- 
crement of  sea-fowl.  It  contains  much  hydric  ammonic 
urate  (NH4H.C5H.,N4O3),  and  from  it  large  quantities  of 
uric  acid  were  formerly  prepared  for  the  manufacture  of 
murexide,  a  red  colouring  matter,  now  superseded  by 
aniline  red. — If  human  urine  be  strongly  acidified  with 
hydrochloric  acid  and  set  aside  for  24  hours  small  crys- 
tals of  uric  acid  collect  on  the  sides  of  the  vessel.  The 
buff-coloured  sediment  of  urine  is  generally  uric  acid  or 
the  acid  ammonium  salt.  As  these  compounds  are  about 
ten  times  as  soluble  in  hot  as  in  cold  water,  they  dis- 
appear when  the  urine  is  heated. 


URATES.  189 

PROPERTIES. — A  crystalline  solid,  in  small  scales  or 
plates,  soluble  at  20C  C.  in  14,000  parts  of  water;  at 
100°  in  1,800  parts.  It  is  insoluble  in  alcohol. 

URATES. — Uric  acid  is  dibasic,  but  the  ordinary  salts 
are  the  acid  salts.  Sodic  hydric  urate  (NaHUr)  is 
found  in  the  urine  and  in  gouty  concretions.  It  is 
soluble  at  15°  in  1,100  to  1,200  parts  of  water,  and  at 
100°  in  123  to  125  parts.— The  potassium  salt  (KHUr) 
dissolves  at  20°  in  700  to  800  parts,  and  at  100°  in  70 
to  80  parts  of  water.  It  is  found  as  a  urinary  deposit  in 
fevers. — Acid  lithium  urate  (LiHUr)  is  more  soluble 
than  either  of  the  above,  dissolving  at  20°  in  about  370 
parts,  and  at  100°  in  39  parts  of  water.  The  normal 
salts  are  freely  soluble  in  water. — Solutions  of  urates 
deposit  uric  acid  when  acidified.  This  fact  is  of  great 
importance  in  the  treatment  of  uric  acid  stone,  and  gout. 
Both  these  diseases  are  due  to  excessive  excretion  of  uric 
acid,  and  the  treatment  adopted  is  to  transform  the  uric 
acid  or  acid  salts  into  more  soluble  salts.  The  acidity  of 
the  urine  is  decreased  by  administering  alkaline  car- 
bonates. As  can  be  easily  seen,  lithium  carbonate  is  the 
best. 

Experiment  151. — Try  to  dissolve  a  little  uric  acid  in  water. 
Add  some  caustic  potash  or  soda.  The  acid  dissolves  easily. 
Acidify  the  solution  with  hydrochloric  acid. 

Test. — To  any  substance  containing  uric  acid  add  a  few  drops 
of  strong  nitric  acid,  and  evaporate  to  dryness  in  a  porcelain 
basin.  (If  the  substance  is  a  dilute  solution  it  must  first  be 
evaporated  to"  dryness. )  The  residue  is  red.  Allow  the  basin  to 
cool,  and  add  a  drop  of  ammonia  solution.  A  tine,  purple-red 
colour  appears.  This  is  the  murtxule  test.  Potash  turns  the 
red  to  blue. 


190  ALCOHOLS. 

CARBON,  HYDROGEN,  AND  OXYGP]N. 

168.  Alcohols. —  Allied   to  the  paraffins  is  a  series 
of    compounds    which    differ   in    composition    from    the 
paraffins  by  containing  an  atom  of  oxygen  in  the  mole- 
cule. 

Paraffins.  A  Icohols. 

Methane CH4  Methyl  alcohol.. . .   CH4O 

Ethane C2H0  Ethyl         "      ....   C2H6O 

Propane C3H8  Propyl        "      ....   C3H8O 

Butane C4H10  Butyl         "      ...   C4H10O 

Pentane C5H12    j    Amyl          "      ....   C5H12O 

&c.  &c. 

These  alcohols  can  be  prepared  from  the  monochlor- 
pai*affins  (CH3C1,  &c.)  by  the  action  of  water  : 

CH3.C1  -f-  H,O  —  CH3.OH  +  HC1. 

Further,  when  they  are  acted  on  by  phosphorus  penta- 
chloride  (PC15),  they  lose  an  atom  of  oxygen  and  one  of 
hydrogen,  and  gain  one  atom  of  chlorine  : 

CH4O  +  PC15  =  CH3.C1  +  POC13  +  HC1. 

From  these  facts  it  is  concluded  that  the  molecules  of 
the  alcohols  contain  the  radical  hydroxyl  ( —  OH).  They 
are  thus  hydroxides  of  the  radicals,  CH3,  C2H5,  «kc. 

Methyl  alcohol CH3.OH 

Ethyl        "      C2H5.OH 

Propyl       "      03H7.OH 

&c.  &c. 

169.  Methyl  Alcohol-     CH3.OH.— Occurs   in  na- 
ture combined  with   acids,  as  in   oil  of  winteryreen  or 


METHYL    ALCOHOL.  191 

methyl  srtlicylate.  When  wood  is  destructively  distilled 
in  the  preparation  of  wood -charcoal,  three  other  products 
are  obtained  :  (1)  A  mixture  of  combustible  gases  (CH4, 
&c.),  (2)  a  watery  liquid,  and  (3)  a  tarry  liquid.  From 
wood  tar,  creosote  and  other  substances  are  obtained. 
The  watery  liquid  (pyroligneous  acid)  is  strongly  acid, 
and  from  it  is  obtained  wood  vinegar,  or  impure  acetic 
acid,  and  wood  spirit,  or  methyl  alcohol.  The  acid  liquor 
is  neutralized  with  lime,  and  the  methyl  alcohol  is  then 
distilled  off  and  purified  by  fractional  distillation,  <fec.  It 
is  also  prepared  in  considerable  quantities  by  distilling 
the  waste  ("  vinasse  ")  from  the  beet  sugar  industry. 

PROPERTIKS. — Pure  methyl  alcohol  is  a  colourless 
liquid,  similar  to  common  alcohol  in  smell,  and  boiling  at 
55°.  1.  Specific  weight  =  0.8142.  It  mixes  with  water 
in  all  proportions,  with  contraction  of  volume  and  evolu- 
tion of  heat.  It  burns  in  air  with  a  pale  blue  flame.  It 
is  a  good  solvent  for  fats,  oils,  resins,  &c. 

Experiment  152. — Mix  a  little  methylic  alcohol  with  an 
equal  volume  of  water,  and  note  the  evolution  of  heat. 

Experiment  153- — Set  fire  to  a  little  methyl  alcohol  in  a 
porcelain  dish. 

CH40  +  3O  =  C02  +  2H20. 

Commercial  methyl  alcohol  is  always  impure,  and  has 
an  unpleasant  odour.  It  is  used  as  a  solvent,  in  the 
manufacture*  of  aniline  dyes,  and  in  the  preparation  of 
methylated  spirit,  a  mixture  of  common  alcohol  with  a 
small  percentage  of  crude  methyl  alcohol.  This  is  unfit 
for  use  as  a  beverage,  and  is  imported  and  manufactured 
free  of  duty  in  Great  Britain.  It  is,  however,  often 


192  ETHYL    ALCOHOL. 

purified  and  used  to  adulterate  liquors.  It  is  probable 
that  many  of  the  cheaper  sorts  of  strong  liquors  sold  in 
this  country  are  preparations  of  methylated  spirit. 

Methylic  alcohol  acts  towards  acids  like  a  weak  base, 
forming  salts,  which  are  unstable  in  presence  of  water. 
Thus,  with  hydrochloric  acid  it  forms  methyl  chloride  : 

CH3.OH  +  HC1  =  CH3.C1  -f  H2O. 
Compare  K.OH  +  HC1  =  K.C1  +  H2O. 

The  reaction  is  not  complete  unless  the  water  be  re- 
moved as  fast  as  it  is  formed.  The  radical  methyl  CH;!) 
thus  plays  the  part  of  a  monad  metal.  The  salts  of 
alcohol  radicals  are  called  ethereal  salts.  Many  of  them 
are  volatile  and  have  a  pleasant  ethereal  smell. — When 
methyl  alcohol  is  oxidised  by  a  mixture  of  sulphuric  acid 
and  potassic  bichromate  (K2Cr2O7)  formic  acid  is  pro- 
duced : 

CH4O  +  O2  =  CH,O,  -f  H,O. 

170.  Ethyl  Alcohol. — C2H80.  This  is  common 
alcohol,  the  intoxicating  principle  of  all  spirituous  bever- 
ages. It  is  also  known  as  spirits  of  wine.  It  is  pre- 
pared by  the  fermentation  of  sugar. 

FERMENTATION. — This  is  a  chemical  action  brought 
about  by  minute  plants  and  animals,  called  ferments, 
which  grow  and  multiply  in  the  fermenting  liquid.  Cer- 
tain conditions  are  necessary  to  the  life  of  the  ferments, 
and  fermentation  ceases  when  these  conditions  are 
absent.  (1)  The  fermenting  liquid  must  contain  the 
elements  of  food  for  the  ferment ;  a  solution  of  pure 
sugar  will  not  ferment.  (2)  The  temperature  must  not 


FERMENTATION.  193 

be  much  above  40°  C.  nor  much  below  20°  C.  (3) 
There  must  be  a  proper  proportion  of  water  present. 
(4)  The  products  of  fermentation  must  not  accumulate 
too  much  ;  and  (5)  certain  substances  called  antiseptics, 
which  are  fatal  to  the  life  of  the  ferment,  must  be 
absent. — There  are  several  kinds  of  fermentation  : 

1.  Alcoholic. — Caused    by   yeast  (saccharomyces),  the 
products  being  alcohol  and  carbon  dioxide  : 

C6H12Oe  =  2C2H60  +  2C02. 

2.  Acetous. — Due   to   the   vinegar  plant   (mycoderma 
aceti),  which  transforms  alcohol  by  oxidation  into  acetic 
acid. 

3.  Lactic. — The  lactic  acid  ferment  has  the  power  of 
transforming  sugar  into  lactic  acid.     The  ferment  is  pre- 
sent in  sour  milk. 

4.  Butyric. — Sugar  is  fermented  into  butyric  add  by 
a  ferment  present  in  rotten  cheese. 

The  germs,  or  sporules,  of  these  ferments  are  present 
everywhere  in  the  air,  and  wherever  they  find  a  suitable 
liquid  they  at  once  cause  fermentation  to  begin. 

Experiment  154. — Make  a  solution  of  commercial  grape 
sugar  in  40  or  50  times  its  weight  of  water.  Put  some  yeast  in 
the  solution  and  fill  three  convenient  dishes  with  it,  also  three 
test  tubes.  Invert  the  latter  in  the  dishes.  Keep  No.  1  in  a 
warm  room.  To  No.  2  add  a  little  carbolic  acid  and  set  it  in 
the  same  room  with  No.  1,  but  not  too  near  it.  Set  No.  3  in  a 
cold  place  and  surround  it  with  snow  or  ice.  Fermentation 
soon  begins  in  No.  1,  and  bubbles  of  gas  rise  into  the  test  tube. 
After  a  considerable  quantity  has  collected,  test  it  for  carbon 
ditixide.  Fermentation  does  not  begin  in  Nog.  2  and  3. — Distil 
14 


194  ALCOHOLIC   LIQUORS. 

a  few  drops  of  alcohol  from  ale,  receiving  it  in  a  cool  test  tube. 
Try  taste  and  inflammability. 

Alcohol  is  separated  from  fermented  liquors  by  dis- 
tillation. Common  alcohol  contains  water,  and  often 
fusel  oil,  which  consists  principally  of  amyl  alcohol. 
This  is  removed  by  fractional  distillation  and  by  filter- 
ing through  charcoal.  As  it  is  poisonous  it  is  important 
to  have  it  absent  from  spirits  of  wine  used  in  medicinal 
preparations. — Absolute  alcohol  is  alcohol  containing  not 
more  than  5  %  of  water.  It  is  prepared  by  distilling 
the  commercial  article  from  quicklime,  which  holds  the 
water.  The  last  traces  of  water  are  very  difficult  to  re- 
move.— Much  alcohol  is  now  prepared  from  glucose,  a 
sugar  made  from  starch.  When  the  glucose  is  made 
from  potato-starch,  the  alcohol  contains  a  large  percen- 
tage of  fusel  oil  and  is  highly  poisonous. 

PROPERTIES. — A  colourless  liquid  of  pleasant  smell, 
boils  at  78°. 3,  freezes  at  —  130°. 5.  When  evaporated  on 
the  hand  it  leaves  no  unpleasant  smell  (if  fusel  oil  is  ab- 
sent). It  burns  with  a  pale  blue  flame,  forming  carbon  di- 
oxide and  water.  (Write  the  equation.)  It  is  a  good  solvent 
for  organic  substances,  especially  those  insoluble  in  water ; 
and  is  used  extensively  in  the  preparation  of  tinctures,  &c. 

Experiment  155. — Dissolve  a  litttle  powdered  resin  in  warm 
alcohol,  and  add  water. 

Beers  contain  from  2  %  to  10  %  of  alcohol ;  wines, 
from  8  %  to  30  %  ;  distilled  liquors  (whiskey,  brandy, 
<fec.),  up  to  75  %.  The  different  flavours  are  due  to 
small  quantities  of  substances  present  in  the  saccharine 
juice,  formed  during  fermentation,  dissolved  out  of  the 


ETHYL    SALTS.  195 

wood  of  the  cask,  or  added  by  the  manufacturer.     Proof 
spirit  contains  49  %  by  weight  of  alcohol. 

In  its  chemical  characters  ethylic  alcohol  is  like 
methylic.  It  is  the  hydroxide  of  ethyl  (C2H5),  and  forms 
salts,  e.g.,  ethylic  nitrate  (C.2H6.NO3^,  chloride  (C2H5.C1), 
sulphate  ((CoH5)2SO4).  &c.  (Write  equations  showing 
the  action  of  nitric,  hydrochloric,  sulphuric,  and  phos- 
phoric acids  on  ethylic  alcohol).  Sweet  spirit  of  nitre  is 
a  solution  in  alcohol  of  ethyl  nitrite  (C2H5.NO2). 

Experiment  156- — Mix  some  alcohol  in  a  t.  t.  with  one-tenth 
its  volume  of  concentrated  sulphuric  acid,  a  little  more  of  nitric 
acid,  and  some  scraps  of  copper.  Warm  gently  and  note  the 
smell.  It  is  that  of  ethyl  nitrite  (" nitrous  ether").  Attach  a 
bent  tube  and  distil  a  drop  or  two  into  a  cool  t.  t. 

Acetic  ether  is  ethyl  acetate  (C2H5.C2H3O2). 

Experiment  157- — Pour  a  little  rectified  spirit  (alcohol)  on 
some  dried  sodic  acetate  in  a  t.  t.  Add  a  small  quantity  of 
concentrated  sulphuric  acid  and  distil  into  another  t.  t.  kept 
cool.  Note  the  smell,  &c.  of  the  acetic  ether  obtained. 

Tests  for  Alcohol.— 1.  Heat  a  little  alcohol  (diluted  with 
water)  with  a  few  drops  of  "bichromate  mixture  "  (4H2S04  + 
KaCraOr).  It  is  turned  green,  and  the  pungent  odour  of  acetic 
aldehydt  (C2E40)  can  be  observed: 

3C2H6O  +  4H2SO4  +  K2Cr2O7  = 

Chromic  sulphate. 

3C2H4O  +  K2SO4  +  Cr2(SOJ3  +  7H2O. 

2.  Mix  a  little  alcohol  with  solution  of  sodic  carbonate,  add 
iodine,  and  warm  somewhat.  Yellow  crystals  of  iodoform 
(CHI8)  appear.  This  is  a  very  delicate  test. 


196  AMTLIC    ALCOHOL. 

171.  Higher  Alcohols.— Propylic  (C3H7.OH),  bu- 

tylic  (C4H9.OH),  and  amylic  (C5HU.OH)  alcohols  are 
formed  in  small  quantities  during  the  fermentation  of 
sugars.  They  boil  at  higher  temperatures  than  does 
ethylic  alcohol,  and  are  partially  separated  from  it  in 
the  process  of  distillation. 


AMYLIC  ALCOHOL  (C5Hn.OH)  forms  the  greater  part 
of  "fusel  oil."  It  is  formed  in  large  proportion  by  the 
fermentation  of  glucose  prepared  from  starch  (amylum). 
Hence  its  name.  The  quantity  is  especially  lai'ge  when 
the  glucose  has  been  manufactured  from  potato-starch  ; 
and  amylic  alcohol  is  sometimes  called  "potato-oil." 

It  is  a  colourless  oily  liquid  of  penetrating  oppressive 
odour.  It  boils  at  132°  0.  Specific  weight  =  0.818. 
When  oxidised  by  bichromate  mixture  it  forms  first  an 
aldehyde  (C5H10O),  and  then  valerianic  acid  (C6H10O2). 
It  is  used  to  prepare  sodic  valerianate  (NaC5H9O2).  It 
is  highly  poisonous,  and  its  presence  in  cheap  spirits  is 
the  cause  of  the  furious  intoxication  often  resulting  from 
their  use. 

Amyl  nitrite  (C6H11.NOa)  is  prepared  by  passing  nitrogen 
trioxide  (from  nitric  acid  and  starch)  into  amyl  alcohol.  It  is  a 
light  yellow  liquid,  boiling  at  99°.  Specific  weight  =  0.902. 
It  is  insoluble  in  water,  but  soluble  in  alcohol.  It  has  a  peculiar 
action  when  inhaled,  increasing  the  rate  of  the  pulse  and  caus- 
ing flushing  of  the  face.  If  inhaled  too  long  it  causes  suffocation 
by  preventing  the  oxygen  of  the  air  from  combining  with  haemo- 
globin. 

Experiment  158. — Dissolve  a  little  amyl  alcohol  (or  fusel  oil) 
in  an  equal  volume  of  concentrated  sulphuric  acid  and  cool. 
Add  to  this  some  solution  of  potassic  or  sodic  nitrite  in  half  its 
weight  of  water.  Heat,  and  observe  odour  of  amyl  nitrite. 


ISOMERIC    ALCOHOLS.  197 

Several  artificial  essences  are  prepared  from  amyl  alcohol,  e.g., 
essence  of  jargonelle  pear  is  amyl  acetate  (^>6H.ll.G2S.^Q.2). 

Test. — To  detect  amyl  alcohol  in  spirits  add  a  few  drops  of 
diltite  acetic  acid  and  potastsic  permanganate  (KMnO4)  solution. 
The  per-manganate  is  quickly  decolourised  if  amyl  alcohol  be 
present.  Or,  shake  with  small  crystals  of  potasslc  iodide.  They 
are  coloured  yellow  by  amyl  alcohol. 

172.  Isomeric  Alcohols.  —  There  is  only  one 
alcohol  having  the  formula  CH4O,  and  one,  (XH6O ;  but 
there  are  two  propylic  alcohols.  One  of  them,  when 
oxidised,  gives  first  an  aldehyde  and  then  an  acid — pro- 
pionic  acid — having  the  same  number  of  carbon  atoms  in 
the  molecule ;  the  other  gives  a  ketone  and  then  breaks 
up  into  two  acids  having  fewer  carbon  atoms  in  the 
molecule.  For  this  and  other  reasons,  it  is  believed  that 
the  hydroxyl  is  differently  situated  in  the  molecules  of 
the  two  acids. 

H       H       H 

I          I          I 

Primtiry  propylic  alcohol H — C — C — C — OH 

I       I       I 

H      H      H 

H       H       H 

I         I          I 

Secondary  propylic  alcohol  . .  H — C — C — C — H 

I       I       I 

H     OH     H 

A  third  kind  of  alcohols  is  known  which,  when  oxidised, 
break  up  at  once  into  acids  having  fewer  carbon  atoms 
in  the  molecule.  These  are  tertiary  alcohols  ;  e.g. : 

CH3 

I 
Tertiary  butylic  alcohol CH3 — C — CHS 

o1 

H 


198  AMINES ETHER. 

1.  Primary  alcohols  are  those  in  which   the  hydroxyl 
is  attached  to  a  carbon  atom  which   is  joined  to  only  one 
other  carbon  atom. 

2.  Secondary  alcohols  have  the  hydroxyl  attached  to  a 
carbon  atom  which  is  joined  to  two  other  carbon  atoms. 

3.  Tertiary  alcohols  have  the  hydroxyl  attached  to  a 
carbon  atom  which  is  joined  to  three,  other  carbon  atoms. 

173.  Amines,  or  Substituted  Ammonias.— 

These  are  compounds  related  to  the  alcohols  in  the  same 
way  as  urea  is  to  carbonic  acid  (Art.  166);  and  they 
can  be  prepared  by  similar  methods,  viz.,  by  the  action 
of  dry  ammonia  on  the  chlorides,  bromides,  or  iodides  of 
the  alcohol  radicals.  Thus, 

Methyl  chloride.  Methylainine. 

CH3C1  +  NH3  =  CH3.NH2  +  HC1. 

They  can  also  be  regarded  as  derivatives  of  ammonia, 
formed  by  replacing  a  hydrogen  atom  by  an  alcohol  radi- 
cal. But  a  second,  and  a  third  hydrogen  atom  may  be 
replaced,  so  that  there  are  three  classes  of  amines,  pri- 
mary, secondary,  and  tertiary,  e.g.,  mono-methyl-amine 
(CH3.NH2),  di-methyl-amine  ((CH3)2NH),  and  tri- 
methyl-amine  ((CH3)3N).  Many  of  these  compounds  are 
found  in  nature.  They  resemble  ammonia  in  their  pro- 
perties, uniting  with  acids  to  form  salts,  e.g  ,  C2H5.NH3C1, 
ethyl-ammonium  chloride,  and  dissolving  freely  in  water  to 
a  strongly  alkaline  solution.  They  generally  smell  like 
ammonia. 

174.  Ether. — (C2H5)2O.     Also    known  as  sulphuric 
ether.     Ethers  are  oxides  of  alcohol  radicals,  and  bear 
the  same  relation  to  alcohols  as  the  oxides  of  the  metals 
do  to  their  hydroxides. 


ETHER.  199 

PREPARATION. — Experiment  159. — Mix  4  parts  of  alcohol 
with  9  of  sulphuric  acid,  and  heat  the  mixture  until  it  begins 
to  boil  (at  140°  C.).  Ether  distils  over.  Note  its  odour. 

On  the  large  scale,  as  soon  as  the  mixture  begins  to 
boil,  a  small  stream  of  alcohol  is  allowed  to  flow  into  the 
still,  and  the  temperature  is  kept  at  140°  C.  Ether  and 
water  then  distil  over  together  continuously. 

EXPLANATION. — Alcohol  forms  with  sulphuric  acid, 
ethyl  sulphuric  acid,  or  ethylic  hydric  sulphate 
(C2H5.H.S04) : 

C2H5.OH  +  H2SO4  =  C2H5.H.S04  -f  H2O. 

The  water  distils  as  fast  as  it  is  formed.  Alcohol  decom- 
poses ethyl  sulphuric  acid,  forming  ethyl  ether  and  sul- 
phuric acid  : 

C2H5.OH  -f  C2H5HSO4  =  (C2H5)oO  +  H2SO4. 

The  process  is  continuous,  the  alcohol  being  run  in  as 
fast  as  the  ether  and  water  distil.     As  far  as  the  ulti- 
mate result  goes,  the  action  can  be  represented  thus : 
2C2H5OH  =  (C2H5)20  +  H20, 

as  a  simple  dehydration  of  alcohol. — It  is  purified  by  dis- 
tilling from  calcic  chloride  and  lime  to  remove  water  and 
acids. 

PROPERTIES. — A  colourless  liquid,  bright  and  mobile. 
It  boils  at  34°. 9,  and  the  vapour  forms  an  explosive  mix- 
ture with  air.  Ether  should  never  be  boiled  over  a 
naked  flame,  but  always  in  a  hot  water  bath.  Specific 
weight,  0.736. 

Experiment  160.— Set  fire  to  a  little  ether  in  a  porcelain  dish 
tilled  half -full  of  water. 


200  ALDEHYDES. 

Put  •ome  ether  in  a  t.  t.  and  immerse  in  water  a«  hot  as  the 
hand  can  bear  it.  The  ether  boils . 

Ether  is  sparingly  soluble  in  water,  but  mixes  in  all 
proportions  with  alcohol.  It  is  a  good  solvent  for  fats 
and  oils,  resins,  alkaloids,  <kc.  Water  is  very  slightly 
soluble  in  ether,  and  commercial  ether  generally  leaves  a 
wet  stain  when  evaporated. — Ether  is  a  good  anaesthetic, 
but  not  so  rapid  in  its  action  as  chloroform.  Its  latent 
heat  is  high,  and,  as  it  evaporates  so  readily,  an  ether 
spray  can  be  used  to  deaden  by  cold  the  sensibility  of 
any  part.  Water  can  be  frozen  by  the  rapid  evaporation 
of  ether.  (Try  a  drop  or  two  of  ether  on  the  hand.) 

Commercial  ether  generally  contains  considerable 
alcohol,  which  can  be  removed  by  washing  with  water. 
The  presence  of  alcohol  can  be  detected  by  finding  the 
specific  weight. 

175.  Ald.eh.ydes. — When  primary  alcohols  are  oxid- 
ised slowly,  the  first  substances  formed  contain  two 
atoms  of  hydrogen  less  than  the  alcohols.  They  are 
cfeAycfrogenated,  or  aldehydes. 

Alcohols.  Aldehydes. 


Methyl CH4O 

Ethyl C2H6O 

Propyl C3H80 

Ac. 


Formic...  .  CH,O 


Acetic C2H4O 

Propionic 03H6O 

&c. 


The  atom  of  oxygen  can  be  replaced  by  two  atoms  of 
chlorine ;  and  from  acetic  aldehyde  chloride  of  ethylidene 
(C.)H4C13)  is  thus  formed.  It  is  concluded  that  the  atom 
of  oxygen  is  not  united  with  any  of  the  hydrogen  atoms, 


KETONES CHLORAL.  201 

and  the  formulas  of  the  aldehydes  are  written  graphically 

thus : 

H 
I 
H — C=o,     H — C — C=o,  (fee. 

I  I       I 

H  H       H 

The  group  — C=o  is  characteristic  of  the  aldehydes. — 

I 

H 

(In  what  test  already  made  was  acetic  aldehyde  formed  ?) 

KETONES  are  compounds  analogous  to  the  aldehydes,  and 
are  formed  by  the  oxidation  of  secondary  alcohols.  Thus, 
secondary  propyl  alcohol  (CH3.CHOH.CH3),  yields  on 
oxidation,  di-methyl  ketone,  or  acetone  (CH3.CO.CH3). 
By  further  oxidation  the  ketones  are  broken  up  into 
acids  having  fewer  carbon  atoms  in  the  molecule.  In 
this  they  differ  from  the  aldehydes. 

176.  Chloral. —  CC13 — C=0.     As  can  be  seen  by 

I 

H 

the  formula,  choral  is  derived  from  acetic  aldehyde  by 
the  replacement  of  three  atoms  of  hydrogen  by  chlorine  ; 
it  is  trichloraldehyde. 

PREPARATION. — By  the  action  of  chlorine  upon  alco- 
hol. Aldehyde  is  first  formed  : 

CH3.CH2OH  +  C12  =  CH3.COH  +  2HC1. 
Then,  chlorine  replaces  hydrogen  : 

CH3.COH  +  3C12  =  OC13.COH  +  3HC1. 

PROPERTIES. — A  colourless,  somewhat  oily  liquid  of 
pungent  irritating  odour.  It  boils  at  94°.  Specific  weight 


202  CHLORAL    HYDRATE FATTY    ACIDS. 

=  1.502.  When  mixed  with  §  its  volume  of  water  it 
unites  with  it,  forming  chloral  hydrate  (CCL.CH(OH)2), 
a  white  crystalline  solid. 

Experiment  161. — Mix  a  drop  or  two  of  chloral  on  a  watch 
glass  with  a  drop  of  water.  Crystallisation  takes  place.  Dis- 
solve these  crystals  in  water,  and  note  taste,  &c. 

CHLORAL  HYDRATE  is  soluble  in  alcohol,  water,  and 
ether.  By  the  action  of  strong  bases  it  (as  well  as  chloral) 
is  converted  into  chloroform  and  formic  acid  ; 

Chloral.  Potassic  formate. 

CC13.CHO  +  KOH  =  CHC13  +  HCO2K. 

This  explains  the  formation  of  chloroform  by  the  action 
of  bleaching  powder  on  alcohol.  Bleaching  powder 
yields  chlorine,  and  it  always  contains  calcic  hydroxide. 
— The  substance  generally  sold  as  chloral  is  really  the 
hydrate  or  its  aqueous  solution.  It  should  give  no  pre- 
cipitate with  argentic  nitrate,  and  should  have  an  agree- 
able, somewhat  aromatic,  smell.  Its  taste  is  bitter  and 
astringent. 

When  chloral  hydrate  is  injected  under  the  skin  it  is 
decomposed  by  the  alkali  of  the  blood  into  chloroform 
and  a  formate,  as  shown  in  the  above  equation.  Its 
effects  are  thus  the  same  as  those  of  chloroform.  It  is 
much  used  in  cases  of  sleeplessness. 

177.  Fatty  Acids. — By  further  oxidising  the  alde- 
hydes a  series  of  acids  is  obtained.  They  are  called 
fatty  acids,  because  the  higher  members  are  found  com- 
bined in  fats.  They  differ  from  the  aldehydes  by  having 
one  atom  of  oxygen  more  in  the  molecule. 


FORMIC    ACID.  203 


Alcohols 

Methyl....  CH4(> 
Ethyl.            C2H6O 

Aldehydes. 
Formic  CH20 
Acetic  C2H4O 
Propionic...  C3H6O 
&c. 

Acids. 
Formic  CH2O3 
Acetic  C2H4O2 
Propionic.  C3H6O2 

&c. 

Propyl  ....  C.,H80 
&c. 

These  acids  are  thought  to  contain  hydroxyl,  for,  when 
they  are  treated  with  phosphorus  trichloride  (PC13),  they 
yield  substances  in  which  01  takes  the  place  of  OH,  e.g. : 

3C2H402  +  2PC13  =  3C2H3O.C1  +  P2O3  +  3HC1. 

The    expanded    formulas  of    the    acids   are    H — C=o, 

o 
H 

CH.,.C=o,  <fcc.,  and  the  group  — C=o  is  characteristic. 

!  I 

o  o 

H  H 

It  is  called  carboxyl,  and  the  number  of  carboxyls  in  the 
molecule  of  an  organic  acid  marks  the  basicity  of  the  acid. 

178.  Formic  Acid.— HCO.OH.  First  prepared  by 
distilling  the  bodies  of  red  ants ;  hence  the  name 
(formica,  an  ant).  It  can  be  prepared  by  careful  oxida- 
tion of  methyl  alcohol ;  but  more  conveniently  by  dis- 
tilling oxalic  acid  (C2H204)  with  glycerine  (C3H6O3). 
The  glycerine  undergoes  no  change,  and  the  oxalic  acid 
is  split  into  carbon  dioxide  and  formic  acid  : 

C2H2O4  =  C02  +  CH2O2. 

PROPERTIES. — A  colourless  liquid,  strongly  acid,  boil- 
ing at  99°.  9.  Specific  weight  =  1.22.  When  heated  with 
concentrated  sulphuric  acid  it  breaks  up  into  carbon 
monoxide  and  water.  (Write  the  equation).  It  is  a 
strong  reducing  agent,  and  when  heated  with  argentic 
nitrate  is  oxidised  to  carbon  dioxide,  silver  being  set 


204  ACETIC    ACID. 

free. — It  is  a  monobasic  acid,  and  forms  salts,  the  for- 
mates, all  soluble  in  water,  and  all  having  the  reduc- 
ing properties  of  the  acid. 

Tests. — 1.  Neutral  formates  give  a  red  colour  with  solution 
of  ferric  chloride. 

2.  Formates  or  formic  acid,  when  boiled  with  solution  of  argen- 
tic nitrate,  precipitate  silver. 

179.  Acetic    Acid. —   CH3.C=o.     Acetic   acid    is 

o 
H 

formed  when  acetic  aldehyde  (C2H4O)  is  oxidised;  but 
its  usual  method  of  preparation  is  by  the  direct  oxidisa- 
tion of  ethyl  alcohol  in  the  vinegar  process. 

PREPARATION. — 1.  By  acetous  fermentation  of  wines, 
beers,  &c.  In  many  cases  sugar  is  the  raw  material, 
and  it  undergoes  first  alcoholic  and  then  acetous  fermen- 
tation : 

C2H6O  +  O2  =  C2H4O2  -f  H2O. 

This  gives  a  dilute  solution. 

2.  By  the  destructive  distillation  of  wood,  a  mixture 
of  acetic  acid,  methyl  alcohol,  &c.,  is  obtained.  The 
watery  acid  liquid  is  neutralised  with  slaked  lime 
(Ca(OH)2),  with  which  the  acetic  acid  forms  calcic 
acetate  (Ca(C2H3O2)2)  : 

2(H.C2H302)  +  Ca(OH)2  =  Ca(C2H3O2)2  +  2H2O. 

The  volatile  liquids  are  distilled  off,  and  the  calcic  acetate 
is  purified  by  recrystallisation  and  then  decomposed  by 
sulphuric  acid  : 

Ca(C2H3O2)2  +  H2SO4  =  CaSO4  +  2C2H4O2. 


ACETATES.  205 

Sodic  carbonate  (Na2CO3)  is  often  used  instead  of  lime. 
(Write  the  equation  for  this.)  (How  can  the  acid  be 
separated  from  the  sodium  salt  ?) 

PROPERTIES. — A  colourless  liquid,  of  pungent,  vine- 
gary odour,  and  sharp  acid  taste.  When  free  from  water 
it  boils  at  119°.  When  pure,  it  acts  upon  the  skin 
powerfully,  causing  blistering.  It  is  solid  below  16°. 7, 
but  may  be  cooled  below  0°  in  closed  vessels  without 
causing  it  to  solidify,  being  then  in  a  state  analogous  to 
super  saturation.  Glacial  acetic,  acid  is  the  solid  acid.  Of 
course,  in  warm  weather  it  is  not  solid.  The  specific 
weight  of  the  liquid  is  1.08.  Acetic  acid  dissolves  in 
water  in  all  proportions.  Vinegar  is  an  impure  dilute 
solution  (5  %  to  10  %>.  Sulphuric  acid  is  a  frequent 
adulterant  of  vinegar,  and  can  be  detected  by  its  giving  a 
white  precipitate  with  baric  chloride.  It  sometimes  con- 
tains sulphurous  acid.  (How  test  for  this  ]). — The  acidum 
aceticum  of  the  B.  P.  contains  only  33  %  of  the  pure 
acid.  A  dilute  solution  (4|  %)  is  also  used. 

ACETATES. — Acetic  acid,  like  all  the  acids  of  this 
series,  is  monobasic.  Only  one  of  the  4  atoms  of  hydro- 
gen is  replaceable  by  metal.  The  normal  acetates  are  all 
soluble  in  water.  (Of  what  other  salts  is  this  state- 
ment true?)  Sodic  acetate  (Na.C2H3O2),  zinc  acetate 
(Zn(C2H3O,)2),  and  plumbic  acetate  (Pb.(C2H3O2)2)  are 
the  most  important. 

Experiment  162- — To  a  solution  of  sodic  carbonate  in  water 
coloured  with  a  drop  of  litmus,  add  acetic  acid  (dilute)  until  the 
reaction  is  acid.  Evaporate  and  obtain  crystals  of  sodic  acetate. 
This  salt  readily  forms  a  supersaturated  solution. 

Experiment  163. — Dissolve  a  little  litharge  (PbO)  in  acetic 
acid,  and  evaporate  to  crystallisation.  Examine  the  crystals  as 


206  BUTYRIC    ACID. 

to  taste,  &c.  Redissolve  in  water,  and  boil  with  some  more 
litharge.  It  dissolves  forming  a  basic  acetate  (Pb~"HH  0  ). 

Commercial  acetate  of  lead  (sugar  of  lead)  often  contains 
this  basic  salt.  —  Verdigris  is  basic  cupric  acetate.  It 
forms  as  a  green  rust  on  copper  or  brass  kitchen  utensils 
when  these  are  allowed  to  stand  in  contact  with  vinegary 
articles  of  food,  &c.  As  it  is  very  poisonous,  cases  of 
poisoning  sometimes  occur  in  this  way. — Acetic  acid  dis- 
solves many  of  the  heavy  metals,  e.g.,  iron,  zinc,  lead, 
copper,  &c.,  either  unaided  or  aided  by  the  oxidising 
action  of  the  air.  When  it  is  formed  in  badly  sealed 
cans  of  fruit,  &c.,  it  often  dissolves  the  solder  and  thus 
renders  the  fruit  highly  poisonous. 

Tests. — 1.  Warm  some  sodic  acetate  with  sulphuric  acid  and 
observe  the  vinegary  smell. 

2.  Heat  a  solution  of  an  acetate  with  strong  sulphuric  acid 
and  a  little  alcohol,  and  note  the  smell  of  acetic  ether  (ethyl 
acetate). 

3.  Add  a  few  drops  of   neutral   ferric   chloride   solution   to 
neutral  solution  of   an  acetate.     The  blood-red  colour  of  ferric 
acetate  appears  : 

6NaC2H30.2  +  Fe,016  =  Fe2(C2H302)6  +  6Na01. 

Add  a  few  drops  hydrochloric  acid  to  a  part  of  the  solution  ; 
the  colour  disappears.  Boil  the  remainder  for  some  time  ;  basic 
ferric  acetate  is  precipitated  : 

Fe2(C.H302)6  +  H20  =  Fea(OH)2(C.,H302)4  +  2C.,H403. 

180.  Butyric   Acid.—  C3H7.C=o.     This    acid   is 

o 
H 

combined  with  glycerine  in  butter, — whence  its  name. 
When  butter  turns  rancid,  the  characteristic  odour  is 
that  of  free  butyric  acid.  It  is  prepared  by  the  fermen- 


VALERIANIC    ACID.  207 

tation  of  sugar  by  the  butyric  acid  ferment  of  putrid 
cheese.  It  is  an  oily  liquid  with  chemical  characters 
similar  to  those  of  acetic  acid.  Butyrate  of  sodium 
(Na.C4H7O2)  is  often  present  as  an  impurity  in  valerian- 
ate  of  sodium,  being  formed  by  the  oxidation  of  the 
butylic  alcohol  of  fusel  oil. — Ethyl  butyrate  (C2H5.C4H7O2 ) 
is  the  artificial  essence  of  pineapple. 

181.  Valerianic  Acid. —   C4H9.C=o.    It  is  the 

o 

H 

acid  of  valerian  root.  It  is  now  prepared  from  amylic 
alcohol  (fusel  oil)  by  oxidising  with  bichromate  mixture : 

SCgHnOH  +  2K2Cr2O7  +  8H2SO4  = 
3C5H1002  +  2K2S04  +  2Cr2(S04)3  +  11H2O. 

SODIC  VALERIANATE  (Na.C5H9O2)  is  prepared  by  neu- 
tralising the  acid  with  sodic  hydroxide.  From  this, 
valerianate  of  zinc  is  prepared  by  double  decomposition 
with  solution  of  zinc  sulphate  (ZnSO4).  The  valerianate 
is  sparingly  soluble  and  separates  out  in  pearly  white 
scales : 

ZnSO4  +  2NaC5H9O2  =  Zn(C6H9O2)2  +  Na2SO4. 

182.  Higher  Fatty  Acids. — Fats   and   oils   are 

ethereal  salts  of  glycerine  (an  alcohol)  and  the  higher 
members  of  the  fatty  acid  series.  The  acids  of  common 
fats  and  oils  are  : 

Palmitic  acid C16H31.CO.OH 

Stearic  acid C17H35.CO.OH 

In  this  connection  may  be  mentioned  oleic  acid  (C^Hgg. 
CO. OH)  derived  from  the  olefine  series. — These  acids 


208  GLYCOL OXALIC    ACID. 

are  oily  liquids  or  soft  buttery  solids.  Their  metallic 
salts  are  called  soaps.  Hard  soaps  are  the  sodium,  and 
soft  soaps  the  potassium  salts.  These  are  soluble  in  pure 
water.  Other  salts  are  mostly  insoluble,  e.g.,  the  cal- 
cium, magnesium,  and  lead  salts.  (See  Glycerine.} 

183.  GlyCOl.—  C2H4(OH)2.     Prepared  by  the  action 
of  water  on  ethylene  bromide  (C2H4Br2). 

C2H4C12  +  2H3O  =  C2H4(OH),  +  2HBr. 

(Are  the  hydroxyls  attached  to  the  same  or  to  different 
carbon  atoms  ?) — Glycol  is  a  colourless  liquid,  of  burning, 
sweet  taste.  It  has  the  properties  of  an  alcohol,  forms 
ethereal  salts,  and  when  oxidised  gives  first  an  aldehyde, 
and  then  an  acid  (oxalic  acid.}  It  unites  with  acids  in 
two  proportions,  forming  two  series  of  ethereal  salts  :  (1) 
Those  in  which  one  hydroxyl  is  replaced  by  salt  radicals, 
e.g.,  C2H4.OH.C1,  C2H4.OH.N03,  &c.;  and  (2)  those  in 
which  both  hydroxyls  are  replaced,  e.g.,  C2H4C12, 
C2H4(NO3)2,  &c.  It  is  thus  analogous  to  diacid  bases, 
such  as  calcic  hydroxide  (Ca(OH)2) ;  and  is  therefore 
called  a  diacid  alcohol. 

184.  Oxalic  Acid.— C2H2O4.2H2O. 

OCCURRENCE. — In  juices  of  wood  sorrel,  rhubarb,  sour 

dock,    &c.,    as   hydric  potassic   oxalate    (HKC204)  ;    in 

some   plants   and   in    urinary   calculi,    as    calcic  oxalate 

(CaC,O4);  and  in  guano  as  ammonic  oxalate  ((NH4)2C2O4). 

(What  is  the  basicity  of  oxalic  acid  1} 

PREPARATION. — Oxalic  acid  is  formed  when  sugar, 
starch,  &c.,  are  oxidised  with  concentrated  nitric  acid  ; 
but  is  now  made  from  pine  sawdust  by  roasting  at  about 


OXALIC    ACID.  209 

200°  with  a  mixture  of  potassic  and  sodic  hydroxides. 
The  fused  mass  yields  a  solution  of  sodic  oxalate 
(Na2C2O4),  which  is  decomposed  by  boiling  with  milk 
of  lime  (calcic  hydroxide  stirred  up  with  water).  In- 
soluble calcic  oxalate  is  precipitated  and  sodic  hydroxide 
remains  in  solution  : 

Na2C2O4  +  Ca(OH)2  =  CaC2O4  +  2NaOH. 

The  calcic  oxalate  is  then  drained,  washed  and  decom- 
posed by  dilute  sulphuric  acid  : 

CaC2O4  +  H2SO4  =  CaSO4  +  H2C2O4. 

Calcic  sulphate  is  only  sparingly  soluble,  so  that  most 
of  it  remains  undissolved.  The  solution  of  oxalic  acid  is 
drawn  off,  and  evaporated  to  crystallisation.  It  is  puri- 
fied by  re-crystallisation,  and,  if  this  is  not  done,  it  is 
contaminated  by  a  small  quantity  of  calcic  sulphate. — 
The  method  of  preparing  organic  acids  by  precipitating 
the  calcium  salt  and  then  decomposing  with  sulphuric 
acid  is  very  common.  The  object  of  the  precipitation  is 
twofold  :  (1)  To  separate  the  acid  from  the  other  sub- 
stances dissolved  along  with  it,  and  (2)  to  get  it  com- 
bined with  a  metal  whose  sulphate  is  insoluble.  (Why 
is  this  advantageous  1) 

PROPERTIES. — Oxalic  acid  is  a  white  crystalline  solid, 
of  sharp  acid  taste.  The  crystals  are  long  and  pointed 
(prismatic),  and  somewhat  resemble  those  of  Epsom  salts. 
The  acid  is  soluble  in  10  parts  of  water.  It  is  a  strong 
acid,  and  decomposes  carbonates  with  effervescence.  The 
pure  acid  is  entirely  decomposed  and  dissipated  by  heat. 
If  any  residue  remains,  it  is  impurity. 

H2C2O4  =  H2O  +  CO  +  COa. 
15 


210  OXALATES. 

Experiment  164. — Examine  carefully  some  crystals  of  oxalic 
acid,  noting  shape,  taste,  &c.,  and  comparing  with  Epsom  salts. 
Heat  a  small  quantity  on  mica.  Dissolve  a  little  in  water,  taste 
the  solution,  and  try  its  action  on  sodic  carbonate  and  on  lime 
water. 

Oxalic  acid  in  large  doses  (60  grains  and  upwards)  is, 
like  all  strong  acids,  a  corrosive  poison.  In  smaller 
doses  it  is  a  cumulative  poison.  The  oxalates  of  the 
alkalis  are  also  poisonous.  The  antidotes  are  chalk  and 
water,  magnesia,  and  lime  water ;  their  object  being  to 
form  insoluble  oxalates. 

OXALATES. — Oxalic  acid  is  dibasic.     Its  molecule   is 

CO.  OH 

made  up  of  two  carboxyls,   |  . — There  are  two  classes 

CO.OH 

of  oxalates,  normal  (K2C2O4,  CaC2O4,  &e.),  and  acid 
(KHC2O4,  &c.).  Besides  these  there  are  salts  composed 
of  ordinary  acid  oxalates,  combined  with  a  further  quan- 
tity of  acid,  e.g.,  salt  of  sorrel  (KHC2O4.H2C2O4.2H2O). 
Of  the  normal  oxalates  only  those  of  alkalis  ara  soluble 
in  water : 

Experiment  165- — Prepare  some  ammon'tc  oxalate  by  neutral- 
ising solution  of  oxalic  acid  with  ammonia,  filtering  if  necessary, 
and  evaporating  to  crystallisation.  Make  a  solution  of  this  salt 
for  the  following  experiments.  (Write  the  equation.) 

Experiment  166.— Add  solution  of  calcic  chloride  (CaCla)  to 
solution  of  ammonic  oxalate.  Calcic  oxalate  (CaCa04)  is  pre- 
cipitated. Test  its  solubility  in  acetic  and  in  hydrochloric  acids. 
Repeat  with  baric  chloride  (BaCl2). 

Experiment  167- — Add  solution  of  ammonic  oxalate  to  solu- 
tion of  ferrous  sulphate  (FeS04).  Ferrous  oxalate  is  precipi- 
tated. Note  its  colour,  &c.,  and  write  its  formula. 

Tests. — 1.  (Experiment  166.) 


GLYCERINE.  211 

2.  Dry  oxalic  acid  or  an  oxalate  heated  with  concentrated  sul- 
phuric acid  gives  off  a  mixture  of  carbon  monoxide  and  dioxide  ; 
the  monoxide  can  be  lighted,  and  burns  with  the  characteristic 
bluish  flame. 

3.  Argentic  nitrate  gives  a  white  precipitate  of  argentic  oxa- 
late (Ag2C204).     This  is  soluble  in  nitric  acid.     (Can  this  pre- 
cipitate be  obtained  with  free  oxalic  acid  ?) 

4.  "  With  a  solution  of  sulphate  of  lime,  oxalic  acid  gives  a 
white  precipitate  which  is  soluble  in  nitric  acid,  but  insoluble  in 
the  vegetable  acids."     (B.  P.) 

5.  Heat  a  little  calcic  oxalate  in  a  small  t.  t.,  applying  a  light 
to  the  mouth.     The  flame  of  carbon  monoxide  is  seen.     Add  a 
little  hydrochloric  acid  to  the  cold,  white  residue.    It  effervesces. 

CaC204  =  CaC03  +  CO. 
CaCOs  +  2HC1  =  CaClj  +  CO.,  +  H.,0. 
In  this  way  insoluble  oxalates  can  be  tested  for. 

185.  Higher  Dibasic  Acids. — Oxalic  acid  is  the 
first  member  of  a  series  : 

Oxalic  acid (COOH)2 

Malonic  acid CH2(COOH)2 

Succinic  acid C2H4(COOH)2 

&c.  &c. 

SUCCINIC  ACID  is  formed  in  small  qviantities  during 
the  alcoholic  fermentation  of  sugar.  It  is  found  in  the 
urine  of  the  horse,  and  in  the  fluids  of  hydrocoele  and 
hydatid  cysts.  It  is  prepared  by  distilling  amber. 

186.  Glycerine. — C3H5(OH)3.     Glycerine  forms  the 
alcoholic  (basic)  part  of  the  ethereal  salts  called  fats  and 
oils,  and  is  prepared  from  them  in  the  process  of  soap- 
making. 

'  SOAPS  are  made  by  boiling  fats  and  oils  with  aqueous 
solutions  of  sodic  hydroxide  for  hard  soaps,  and  potassic 


212  SOAPS. 

hydroxide  for  soft  soaps.  The  principal  fats  used  are  mix- 
tures in  various  proportions  of  stearin  (C3H5(C18H350.,)3), 
palmitin  (C3H5(C16H31O2)3),  and  olein  (C3H6(C18H33O,)3). 
Stearin  is  the  chief  constituent  of  tallow.  It  is  solid. 
Palmitin  (a  solid)  is  the  chief  constituent  of  palm  oil, 
and  olein  (a  liquid)  'of  olive  oil.  Human  fat  consists 
mostly  of  palmitin.  In  the  process  of  saponification,  the 
glycerine  is  separated  as  indicated  in  the  following  equa- 
tion, in  which,  for  the  sake  of  simplicity,  F  is  put  as  a 
symbol  for  the  salt  radicals  of  the  fatty  acids  : 

Fat.  Soap.  Glycerine. 

C3H5F3  +  3NaOH  =  3NaF  +  C3H5(OH)3. 

The  soap  is  separated  as  a  curd  by  the  addition  of  com- 
mon salt  to  the  solution,  and  the  glycerine  is  recovered 
from  the  mother  liquor.  In  many  factories  the  fats  and 
oils  are  decomposed  by  heating  under  pressure  with 
water  and  2  to  3  per  cent,  of  sulphuric  acid.  Aqueous 
solution  of  glycerine  and  fatty  acids  are  obtained  in 
separate  layers.  The  glycerine  is  drawn  off  and  purified 
by  distilling  with  steam  at  180°  C.  The  acids  are  neu- 
tralised with  caustic  soda  to  form  soap  : 

C3H5F3  +  3H20  =  C3H5(OH)3  +  3HF. 
HF  +  NaOH  =  NaF~  +  H2O. 

Sodium  and  potassium  soaps  are  soluble  in  water  and  in 
alcohol.  Most  soaps  are  insoluble  in  salt  water,  but 
soap  made  from  cocoanut  oil  and  resin  ("  marine  soap  ") 
is  soluble  in  salt  water.  Lime  and  magnesium  salts  de- 
compose ordinary  soaps  forming  insoluble  lime  and  mag- 
nesium soaps.  Thus,  in  washing  with  hard  water,  there 
is  always  a  waste  of  soap. — Lead  plaster  is  a  lead  soap, 
made  by  heating  litharge  with  olive  oil  (or  lard)  and  a 


GLYCERINE.  213 

little  water.     It  consists  of  plumbic  oleate  and  palmitate 
principally,  with  some  glycerine. 

PROPERTIES  OF  GLYCERINE. — A  thick,  sticky,  colour- 
less liquid,  of  sweet  and  burning  taste.  Specific  weight 
=  1.28.  It  dissolves  in  water  in  all  proportions,  and  is 
hygroscopic.  It  mixes  with  alcohol  in  all  proportions. 
It  is  a  very  good  solvent  for  metallic  oxides,  salts,  <fcc.; 
and  is  used  in  medicine  in  preparing  glyceritum  acidi 
carbolici,  &c.,  which  are  solutions  in  4  fluid  ounces  of 
glycerine,  of  1  ounce  of  carbolic,  gallic,  and  tannic  acids. 
Glycerine  has  antiseptic  properties,  and  borate  of  glyce- 
rine has  been  successfully  employed  in  surgery. — Gly- 
cerine cannot  be  distilled  alone,  but  decomposes  at  280°, 
giving  off  pungent  choking  fumes  of  acrolein  (C3H4O). 
If  water  be  present  part  of  the  glycerine  distils  along 
with  the  water. — Glycerine  is  a  triacid  alcohol,  and  is 
analogous  to  triacid  metallic  bases,  such  as  bismuth  hy- 
droxide (Bi(OH)3).  It  forms  three  series  of  ethereal 
salts,  in  which  one,  two.  and  three  hydroxyls  respectively 
are  replaced  by  acid  radicals ;  e.g.,  C3H5(OH)2.NX)3, 
C3H5(OH).(:N"O3),,  and  C3H6(NO3)3,  the  three  nitrates  of 
glycerine.  The  latter,  trinitrate  of  glycerine,  is  com- 
monly called  nitro-glycerine,  and  is  prepared  by  the  ac- 
tion of  a  mixture  of  concentrated  nitric  and  sulphuric 
acids  on  cooled  glycerine.  When  the  resulting  liquid  is 
poured  into  water,  nitro-glycerine  separates  out  as  a 
heavy  oil : 

C3H5(OH)S  +  3HN03  =  C3H5(NO3)3  +  3H2O. 

Dynamite  is  made  by  absorbing  glycerine  in  a  porous 
siliceous  sand.  It  is  less  dangerously  explosive  than 
nitro-glycerine. — On  account  of  its  attraction  for  mois- 
ture, glycerine  is  valuable  in  surgery  as  an  emollient. 


214  LACTIC    ACID. 

It  is  also  used  for  preserving  fruits,  for  making  copying 
ink,  and  as  a  lubricator. 

IMPURITIES. — Glycerine  is  often  adulterated  with  cane 
sugar  and  glucose.  To  detect  cane  sugar,  dissolve  in 
water,  add  a  few  drops  of  sulphuric  acid,  and  evaporate 
on  the  water  bath.  If  cane  sugar  is  present  it  is  black- 
ened. To  detect  glucose,  heat  with  solution  of  caustic 
soda.  If  glucose  is  present  the  solution  turns  brown. 
(Try  with  samples  of  glycerine.)  Glycerine  should  be 
neutral  to  litmus.  Owing  to  imperfect  purification  it 
sometimes  contains  acids. 

Test. — 1.  If  a  liquid  containing  glycerine  be  made  slightly 
alkaline  with  caustic  soda,  a  borax  bead  dipped  in  it  will  give 
a  green  colour  to  the  Bunsen  flame. 

2.  Heat  a  little  glycerine  with  concentrated  sulphuric  acid, 
and  note  the  smell  of  acrolein. 

178.  Hydroxy-Acid.8. — There  are  organic  acids, 
the  molecules  of  which  have  alcoholic  hydroxyl.  They 
partake  of  the  nature  of  both  alcohols  and  acids,  but  the 
acid  properties  predominate.  They  are  called  hydroxy- 
acids.  Thus,  hydroxy-acetic  acid  has  the  formula 
CH2OH.COOH.  There  may  be  two  hydroxyls  in  the 
molecule,  as  in  the  case  of  tartaric  acid,  which  is  dihy- 
droxy-succinic  acid. 

188.  Lactic  Acid,  C2H4OH-OOOH.— This  is  hy- 
droxy-propionic  acid.  Ordinary  lactic  is  formed  by  fer- 
mentation of  milk,  which  contains  a  fermentable  sugar 
(galactose).  It  is  present  in  the  gastric  juice,  and  in 
pickled  cabbage  and  cucumber.  It  is  generally  prepared 
by  the  lactic  fermentation  of  cane  sugar,  by  means  of 
putrid  cheese.  Zinc  carbonate  is  added  to  neutralise  the 


TARTARIC    ACID.  215 

acid  as  fast  as  it  is  formed. — It  is  a  monobasic  acid. 
Ferrous  lactate  (Fe(C3H503)2)  is  prepared  by  dissolving 
iron  filings  in  warm  dilute  lactic  acid.  Tt  is  used  in 
medicine. 

189.  Tartaric  Acid,  C4H6OS. — This  is  dihydroxy- 

CH.OH— COOH 

succinic  acid,  and  its  expanded  formula  is  i 

CH.  OH— COOH 

It  is  at  the  same  time  a  diabasic  acid  and  a  diacid 
alcohol.  There  are  three  isomeric  tartaric  acids,  and  our 
chemical  theory  is  inadequate  to  explain  their  isomerism. 
Ordinary  tartaric  acid  is  found  in  the  juice  of  grapes, 
berries  of  the  mountain  ash,  cucumbers,  potatoes,  <kc. 

PPEPARATION. — From  tartar  or  argol,  which  is  impure 
potassic  hydric  tartrate  deposited  in  wine  casks  and  vats 
during  fermentation.  (It  is  less  soluble  in  alcohol  than 
in  water.)  From  this  salt,  purified  by  crystallisation, 
the  acid  is  prepared  as  follows :  "  Boil  45  oz.  cream  of 
tartar  (potassic  hydric  tartrate)  with  two  gals,  water; 
add  12£  oz.  prepared  chalk  gradually,  stirring  constantly : 

2KHT  +  CaCO3  =  K2f  +  Caf  +  H2O  +  CO2. 

(f  =  C4H4O6.)  Then  add  13£  oz.  calcic  chloride  in 
2  pints  of  water  : 

K2T  +  CaCl2  =  Caf  +  2KC1. 

Allow  the  calcic  tartrate  to  subside,  pour  off  the  liquid 
(What  does  it  contain?),  wash  the  precipitate  with  dis- 
tilled water  until  tasteless,  and  pour  on  it  1 3  fluid  ounces 
sulphuric  acid  diluted  with  3  pints  of  water.  Boil  for 
half  an  hour  and  filter  : 

CaT  +  H2SO4  =  CaS04  +  HaT. 


216  SEIDLITZ    POWDER. 

Evaporate  the  filtrate  at  a  gentle  heat  to  specific  weight 
1.21,  allow  to  cool,  and  separate  the  deposited  gypsum 
(CaSO4.2H3O).  Again  evaporate  till  a  film  forms  on 
the  surface,  cool,  and  drain  the  crystals  of  tartaric  acid 
which  form."  (B.P.)  (Potassic  hydric  tartrate  is  spar- 
ingly soluble,  the  normal  tartrate  quite  soluble,  calcic 
tartrate  insoluble  in  water.  Explain  the  steps  of  the 
above  process). 

PROPERTIES. — Large  colourless  crystals  or  a  white 
granular  powder,  of  acid  taste,  soluble  in  water  (2  parts 
in  1  of  water),  and  in  alcohol,  but  not  in  ether. 

Experiment  168- — Examine  the  appearance  and  taste  of  a 
crystal  of  the  acid,  then  heat  on  mica.  It  browns  and  chars 
with  the  smell  of  burning  sugar  ;  with  a  stronger  heat  it  burns 
away  completely. 

Experiment  169- — Carefully  add  solution  of  tartaric  acid 
to  some  solution  of  potassic  carbonate  (coloured  with  litmus) 
until  the  solution  is  neutral.  (What  salt  is  present  ?)  Then  add 
more  tartaric  acid,  stirring  all  the  time  with  a  glass  rod.  A 
white  granular  precipitate  of  the  acid  tartrate  forms.  (Write 
equations. ) 

Experiment  170. — To  a  strong  solution  of  hot  sodic  car- 
bonate in  a  porcelain  basin  add,  a  little  at  a  time,  cream  of 
tartar  (KHT)  until  it  causes  no  effervescence.  Evaporate  to 
crystallisation.  Sodic  potassic  tartrate,  or  Rochelle  salt  (KNaT), 
is  formed  : 

Na,CO3  +  2KHT  =  2KNaT  +  H30  +  CO, 

Rochelle  salt  is  one  ingredient  of  seidlitz  powder.  The 
blue  paper  contains  usually  3  parts  Rochelle  salt  and  1 
part  sodic  hydric  carbonate;  and  the  white,  1  part  tar- 
taric acid.  When  the  solutions  are  mixed  the  tartaric 


TARTAR    EMETIC.  217 

acid  decomposes  the  carbonate,  while  the  Rochelle  salt 
takes  no  part  in  the  action  : 

2NaHC03  +  H2f  =  Na2T  +  2H2O  +  2C02. 

(From  this  equation  calculate  the  proportions  of  car- 
bonate and  acid  which  must  be  used  in  order  that  there 
may  be  excess  of  neither.) 

TARTAR  EMETIC  is  antimonyl  potassic  tartrate 
(SbO.K.T),  prepared  by  boiling  cream  of  tartar  with 
antimony  trioxide  : 

Sb2O3  +  2KHT  =  2SbOKT~+  H2O. 

The  radical  antimonyl  (SbO)  acts  the  part  of  a  monad 
metal. 

Tests. — 1.  To  neutral  solution  of  a  tartrate  add  calcic  chlo- 
ride or  to  a  solution  of  tartaric  acid  add  lime  water  ;  calcic  tar- 
trate (CaT)  is  precipitated.  Wash  the  precipitate  on  a  filter, 
break  the  filter,  wash  the  precipitate  into  at.  t.,  and  add  sodic 
hydroxide  ;  the  precipitate  dissolves. 

2.  To  tartaric  acid  or  solution  of  a  tartrate  add  acetic  acid 
and  potassic  acetate  and  stir.     Potassic  hydric  tartrate  is  pre- 
cipitated as  a  white  granular  powder. 

3.  Tartaric  acid  or  tartrates  in  solution,  when  heated  with  a 
considerable  proportion   of  concentrated    sulphuric  acid,    turn 
brown  or  black  at  once. 

4.  To  neutral  or  alkaline  solution  of  a  tartrate  add  a  few 
drops  of  potassic  permanganate  solution,  and  heat.     The  colour 


NOTE. —  Solution  of  tartaric  acid  in  water  undergoes 
slow  decomposition  owing  to  the  growth  of  a  fungus. 
Spirits  of  wine  prevent  this. 


218  CITKIC    ACID. 

190.  Citric  Acid,  C6H8O7.H,O.     This  is  a  tribasic 

(  COOH 

acid  and  monacicl  alcohol.  C,H..OH  J.  COOH.     It  is  found 

(COOH 

in  the  juices  of  limes,  lemons,  currants,  raspberries, 
gooseberries,  &c. 

PREPARATION. — From  the  evaporated  juice  of  unripe 
limes  and  lemons  by  almost  the  same  method  as  that  for 
tartaiic  acid.  One  hundred  parts  of  lemon  yield  5^ 
parts  of  acid. 

Experiment  171. — Squeeze  the  juice  of  a  lemon  upon  a  filter 
and  allow  it  to  run  into  a  porcelain  basin.  Heat  to  boiling  and 
add  prepared  chalk  by  degrees  until  it  does  not  cause  effer- 
vescence : 

2H3(jT+  3CaC03  =  Ca3Ci2  +  3H,,0  +  3CO2. 

Filter,  wash  the  precipitate  with  hot  water  four  or  five  times, 
break  the  paper  and  wash  the  precipitate  through  into  a  porce- 
lain basin.  Add  a  small  quantity  (2  or  3  c.c.)  of  sulphuric  acid, 
boil  gently  for  a  little  while,  filter  and  concentrate  the  filtrate 
to  crystallisation.  Crystals  of  citric  acid  are  obtained  mixed 
with  gypsum.  The  acid  can  be  purified  by  recrystallisation 
from  a  small  quantity  of  hot  water. 

PROPERTIES. — Usually  sold  as  large,  colourless,  sharp- 
pointed  crystals.  These  are  soluble  in  £  their  weight  of 
cold  water,  in  f  of  boiling  water  ;  less  soluble  in  alcohol ; 
still  less  in  ether.  Citrie  acid  melts  at  100°  ;  above  this 
it  loses  water  of  crystallisation,  then  chars  with  the  smell 
of  burning  sugar. 

Experiment  172. — Heat  a  little  solid  citric  acid  in  a  t.  t. 
by  placing  the  t.  t.  in  boiling  water.  The  acid  melts.  Wipe 
the  t.  t.  and  heat  it  gently  in  the  Bunsen  flame.  Note  the 
water  condensing.  Heat  more  strongly.  Note  smell  and  char- 
ring. Heat  a  small  crystal  strongly  on  mica. 


MAGNESIC    CITRATE.  219 

Experiment  173. — Put  a  pipetteful  of  citric  acid  solution  in 
each  of  three  porcelain  dishes,  and  carefully  neutralise  with  so- 
lution of  potassic  carbonate.  To  one  add  a  second  pipetteful  of 
the  acid,  and  to  another  two  pipettefuls.  Number  and  evapo- 
rate the  three  solutions  on  a  water  bath  to  a  syrupy  consistence, 
and  then  set  aside  to  crystallise.  Three  salts  are  obtained,  dif- 
ferent in  appearance.  What  are  they  ?  Taste  them. 

Experiment  174- — Dilute  a  little  of  the  normal  potassic 
citrate  prepared  in  Experiment  172  and  add  to  it  about  an  equal 
volume  of  the  reagent  solution  of  calcic  chloride.  If  no  precipi- 
tate appears,  boil.  If  a  precipitate  appears,  gradually  add 
distilled  water  and  shake  up  until  it  is  dissolved,  and  then 
boil.  Calcic  citrate  (Ca3Ci2.4H20)  is  less  soluble  in  hot  than  in 
cold  water.  (Try  with  citric  acid  and  lime  water.) 

MAGNESIC  CITRATE  (Mg3Ci2.4H2O)  is  a  white  sparingly 
soluble  salt  which  can  be  prepared  by  the  action  of  the 
acid  on  magnesium  carbonate  : 

3MgCO3  +  2H3Cl  =  Mg3CT2  +  3H2O  +  3CO2. 

Effervescing  powders  are  made  by  mixing  the  substances 
in  the  solid  state.  Solids  act  on  each  other  very  slowly 
or  not  at  all ;  but  as  soon  as  they  are  brought  together 
in  solution  action  begins.  The  more  coarsely  granular 
the  solid  ingredients  are,  the  slower  is  the  action  between 
them.  (Why  ?)  "  Granular  effervescing  citrate  of  mag- 
nesia "  is  a  mixture  of  coarse  granules  of  Epsom  salts 
(MgSO4.7H2O),  citric  acid,  tartaric  acid,  and  hydric 
sodic  carbonate  (HNaCO3).  (What  substances  are 
formed  when  it  is  dissolved?) 

Tests. — 1.  (Exp't  174.)  The  precipitate  of  calcic  citrate  is 
insoluble  in  sodic  hydroxide. 

2,  Citric  acid  and  citrates  give  no  precipitate  with  acid  solu- 
tions of  potassium  salts.  (See  tartaric  acid). 


220  CARBOHYDRATES. 

3.  Heat    with    concentrated    sulphuric    acid.     The    solution 
darkens  only  after  some  time. 

4.  To  neutral  or  alkaline  solution  add  a  few  drops  of  potassic 
permanganate  solution  and  heat  slowly.     The  permanganate  is 
reduced  to  the  green  manganate  (K2Mn04.) 

5.  All  solid  citrates  are  charred  by  heat. 

Note. — Citric  acid  is  sometimes  adulterated  with  tartaric  acid. 
This  can  be  detected  by  the  test  with  potassium  salts.  The 
addition  of  alcohol  increases  the  delicacy  of  the  test. 


191.  Carbohydrates. — Under    this    head    are    in- 
cluded three  groups  of  isomeric  compounds,  containing 
either  6  or  12   atoms  of  carbon    in    the    molecule,    to- 
gether  with   hydrogen  or  oxygen  in  the  proportions  to 

Jorm  water.  They  are  all  naturally  occurring  sub- 
stances, many  of  them  being  quite  familiar  vegetable  pro- 
ducts, e.g.,  sugar,  starch,  and  cotton.  (What  compounds 
already  discussed  contain  hydrogen  and  oxygen  in  the 
proportion  to  form  water?). — The  carbohydrates  are  in 
all  probability  poly-acid  alcohols,  and  at  the  same  time 
aldehydes  or  ketones.  There  are  three  groups  : 

1.  SACCHAROSES  (C^H^On). — Comprising  cane  sugar 
(saccharose),  milk  sugar,  malt  sugar  (maltose),  &c. 

2.  GLUCOSES  (C6H12O6). — Grape  sugar  (dextrose),  fruit 
sugar  (levulose),  galactose,  inosite,  &c. 

3.  AMYLOSES   (C6H10O5). — Starch,    dextrin,   cellulose, 
glycogen,  gums,  inulin,  &c. 

192.  Saccharoses,  —  C^H^On.     The   saccharoses 
seem  to  be  analogous   to   ethers — they  all    unite    with 
water  to  form  glucoses  : 

C^H-jaOn  -f  H2O  =  2C6H12O6. 


CANE    SUGAR.  221 

(Compare  this  with  the  relation  of  ethyl  ether  to  ethyl 
alcohol.)  This  takes  place  under  the  influence  of  fer- 
ments, and  by  boiling  with  dilute  acids  or  alkalis. 


1.  CANE  SUGAR  (Cl2H..22On). — Cane  sugar  is  so  called 
because  of  its  manufacture  from  the  juice  of  the  sugar- 
cane. It  is  also  now  largely  made  from  the  sugar-beet. 
It  is  found  also  in  the  sugar-maple,  sorghum,  turnips, 
carrots,  coffee,  walnuts,  hazelnuts,  almonds,  and  in  tht 
blossoms  of  many  plants  along  with  more  or  less  fruit- 
sugar. 

PREPARATION. — The  juice  is  expressed  from  sugar- 
cane or  beet-root-pulp,  and  the  solution  of  sugar  is 
evaporated  in  "  vacuum  pans."  It  is  necessary  to  evapo- 
rate at  moderate  temperatures  in  order  to  avoid  the  change 
into  uncrystallisable  members  of  the  glucose  group. 
The  evaporated  solution  is  allowed  to  crystallise,  and 
the  crystals  are  drained  on  "  centrifugals,"  which  are 
large  revolving  sieves.  Molasses  is  the  mother  liquor 
drained  from  the  raw  cane-sugar.  It  is  much  used  in 
the  manufacture  of  rum,  and  lately  a  process  has  been 
devised  to  obtain  more  sugar  from  it.  It  contains 
cane-sugar,  together  with  inverted  sugar,  or  glucoses  pro- 
duced by  the  action  of  water  on  cane-sugar, — inverted, 
because  the  action  on  polarised  light  is  the  exact  op- 
posite. It  must  be  remembered  that  beet-sugar  and 
cane  sugar,  when  pure,  are  the  same  substance.  Vinas- 
ses,  the  mother  liquor  of  beet -sugar,  is  evaporated  to 
dryness,  and  distilled.  A  variety  of  useful  products  is 
obtained,  viz.,  potassic  carbonate,  ammonia,  methylic 
alcohol,  trimethylamine,  <fcc. 


222  CANE    SUGAR. 

PROPERTIES. — A  hard,  colourless,  crystalline  solid  of 
specific  weight  1.593.  It  dissolves  in  about  one-third  of 
its  weight,  of  cold  water,  and  in  all  proportions  in  boil- 
ing water.  It  is  insoluble  in  absolute  alcohol,  and  in 
alcoholic  liquors  is  soluble  in  proportion  to  the  water 
present.  It  melts  at  160°  C.,  and  at  190°  gradually 
loses  water  and  darkens,  forming  a  substance  called 
caramel  (saccharum  iistum),  much  used  for  colouring 
wines  and  liquors.  When  heated  more  strongly  it  chars, 
giving  off  fumes  having  a  characteristic  odour. 

Experiment  175. — Make  a  solution  in  water  of  pure  cane- 
sugar,  add  to  part  of  it  a  few  drops  of  cupric  sulphate  solu- 
tion, and  then  solution  of  caustic  soda  A  blue  solution  is 
formed.  Heat  this.  If  the  sugar  is  pure  no  change  takes  place. 
Add  a  few  drops  of  sulphuric  acid  to  a  little  of  the  sugar  solu- 
tion and  boil  for  some  time  in  a  porcelain  dish  ;  add  water  as  it 
evaporates.  Then  repeat  the  test  with  cupric  sulphate  and 
caustic  soda.  A  precipitate  is  formed  on  heating,  at  first  yellow 
and  then  red.  This  precipitate  is  cuprous  oxide  (Cu20).  (Feh- 
ling's  test.) 

As  we  shall  see  later  this  reduction  of  cupric  sulphate 
solution  is  brought  about  by  glucoses,  but  not  generally 
by  saccharoses.  The  cane-sugar  has  been  inverted  by 
boiling  with  an  acid  : 

Dextrose.  Levulose. 

C12H22Oii  +  H2O  =  C6H12O6  -J-  06H12O8. 

This  pi'ocess  goes  on  slowly  in  moist  impure  sugars 
(brown  sugars).  Cane-sugar  does  not  at  once  undergo 
alcoholic  fermentation  when  its  solution  is  mixed  with 
yeast.  It  must  first  be  inverted  by  the  action  of  the 
ferment. 

Experiment  176. — To  a  dilute  solution  of  pure  sugar  add 
some  brewer's  yeast,  and  set  in  a  warm  place  for  half  an  hour. 


MILK    SUGAR.  223 

Then  test  the  solution  as  in  Experiment  175.     It  will  be  found 
to  contain  inverted  sugar. 

There  are  many  ferments  which  cause  the  inversion  of 
cane-sugar.  The  saliva  has  this  power  owing  to  the 
presence  of  a  ferment,  ptyalin,  which  however  loses  its 
power  when  the  solution  becomes  acid,  and,  hence,  as 
soon  as  it  reaches  the  stomach.  Oxidising  agents  con- 
vert cane-sugar  into  oxalic  and  other  acids.  Formerly 
oxalic  acid  was  made  by  the  action  of  nitric  acid  on  cane- 
sugar. — Cane-sugar  forms  soluble  compounds  with  many 
substances  which  are  insoluble  or  sparingly  soluble  in 
water.  It  is  thus  useful  for  keeping  in  solution  certain 
metallic  compounds  employed  in  medicine,  e.g.,  slaked 
lime,  calcic  hypophosphite,  <fec. — Oxymel  is  a  mixture  of 
honey  (80  %),  acetic  acid  (10  %),  and  water  10%).  The 
name  means  acid-honey. 

Tests. — 1.  Fehling's  test  as  in  Experiment  175.  Fehling's 
solution  is  made  by  dissolving  34.64  grams  cupric  sulphate  (blue 
vitriol)  in  water,  adding  200  g.  Rochelle  salt  ( KNaT),  600  g.  to 
700  g.  solution  of  caustic  soda  of  specific  weight  1.12,  and  mak- 
ing up  to  1  litre  with  water.  It  must  be  kept  well  stoppered, 
and  put  in  a  cool,  dark  place. 

2.  Evaporate  on  the  water  bath  with  a  few  drops  of  sulphuric 
acid.     Blackening  shows  the  probable  presence   of  cane-sugar. 
Grape  sugar  does  not  blacken. 

3.  Heat  with  solution  of  caustic  soda,  caustic  potash,  or  po- 
tassic  carbonate.      Does  not  turn  brown. 

2.  MILK  SUGAR  (LACTOSE)  (C12HMOU.H,O).— Milk 
sugar  is  found  in  the  milk  of  all  mammalia.  Human 
milk  contains  4  %  to  5  %,  cow's  milk  3  %. 

PREPARATION.— From  whey,  by  evaporating  and  crys- 
tallising on  threads  or  sticks. 


224  MALT    SUGAR. 

PROPERTIES. — A  hard,  colourless  crystalline  solid, 
with  one  molecule  of  water  of  crystallisation.  It  is 
not  so  soluble  in  water  as  cane-sugar,  dissolving  in  6 
parts  of  cold,  or  3  of  hot  water.  It  is  almost  insoluble 
in  alcohol.  Its  taste  is  less  sweet  than  that  of  cane- 
sugar.  Specific  weight  1.534. — Milk-sugar  does  not 
ferment  with  yeast ;  but  it  undergoes  another  fermenta- 
tion with  the  formation  of  alcohol  and  lactic  acid,  as  in 
the  preparation  of  the  fermented  liquor  called  koumiss. 
— Milk  sugar  is  used  to  a  considerable  extent  to  increase 
the  percentage  of  sugar  in  cow's  milk  for  feeding  infants. 
It  is  also  used  in  the  powdered  state  as  a  diluent  of  solid 
medicines,  e.g.,  iodoform. 

Experiment  177. — Try  Fehling's  test  with  milk-sugar.  It 
differs  from  the  other  members  of  this  group  in  that  it  reduces 
cupric  to  cuprous  oxide. 

Experiment  178. — Heat  some  solution  of  milk-sugar  with  an 
alkali.  What  result  ? 

Experiment  179. — Add  a  few  drops  of  argentic  nitrate  solu- 
tion to  a  solution  of  milk-sugar,  then  some  ammonia,  and  warm 
gradually.  Silver  is  deposited  as  a  mirror  on  the  t.  t. 

Tests. — Distinguished  from  cane-sugar  by  Experiments  177 
and  178  ;  from  glucose  by  not  fermenting  with  yeast. 


3.  MALT  SUGAR  (MALTOSE)  (C^H^On). — This  is  also 
sometimes  called  starch-sugar. 

PREPARATION. — By  fermenting  potato-starch  with  air- 
dried  malt,  and  extracting  the  sugar  with  alcohol. 

PROPERTIES. — A   crystalline   solid    freely   soluble  in 

water,  and  to  a  considerable  degree  in  alcohol.     It  can 

be  transformed  into  grape-sugar  by  boiling  with  dilute 

ulphuric  acid.     It  is  used  in  the  preparation  of  caramel. 


DEXTROSE.  225 

193.  Glucoses. — CeH^Og.  Glucoses  are  widely  dis- 
tributed in  nature,  occurring  in  both  plants  and  animals. 
Fruits  contain  cane-sugar  in  the  early  stages  of  their 
ripening,  but  this  gradually  combines  with  water  to  form 
glucoses,  generally  equal  quantities  of  dextrose  and  levu- 
lose.  This  mixture  is  called  invert-svgar. 


1.  DEXTROSE  (C6H12O6).  —  Also  called  glucose  and 
grape-sugar. 

OCCURRENCE. — In  sweet  fruits,  almost  always  accom- 
panied by  an  equal  quantity  of  levulose,  thus  showing 
the  origin  of  these  sugars  from  cane-sugar.  The  latter 
is  also  generally  present,  until  the  fruit  becomes  fully 
ripe.  Dextrose  is  also  present  in  honey,  and  in  the  blood, 
liver,  and  urine  of  man.  In  cases  of  diabetes  mellitus 
the  quantity  in  the  urine  may  reach  10  °/0- 

PREPARATION. — Dextrose  can  be  prepared  artificially 
from  starch,  dextrin,  cellulose,  &c.,  by  the  action  of 
acids.  It  is  manufactured  on  the  large  scale  from  corn 
and  potato  starch,  by  heating  with  dilute  sulphuric  acid, 
generally  under  pressure.  Dextrin,  isomeric  with  starch, 
is  first  formed,  and  this  combines  with  water : 

C6H1005  -f  H20  =  C6H1206. 

Experiment  180. — Boil  some  starch  paste  for  half  an  hour 
with  about  one-fifth  its  volume  of  dilute  sulphuric  acid,  re- 
placing the  water  as  it  evaporates,  and  then  test  the  solution 
for  dextrose  by  Fehling's  test.  Try  Fehling's  test  with  starch. 

The  sulphuric  acid  is  neutralised  with  chalk  or  lime- 
stone (CaC03),  and  the  solution  of  dextrose  is  drawn  off 
and  evaporated  to  crystallisation.     It  is  difficult  to  crys- 
tallise, as  it  tends  to  form  supersaturated  solutions. 
16 


226  GLUCOSE. 

PROPERTIES. — The  glucose  of  commerce  is  either  a 
thick  syrup  ("  mixing  syrup ")  or  a  hard,  white  solid 
("grape  sugar").  Specific  weight  =  1.825.  It  always 
contains  dextrin,  and  other  substances  of  unknown 
composition.  The  dextrin  is  harmless,  but  the  unknown 
substances  have  an  effect  on  the  human  system  similar 
to  that  produced  by  fusel  oil.  Dextrose  is  not  so  sweet 
as  cane-sugar,  the  sweetness  being  as  1  to  1.66  (or  as  3 
to  5).  At  15°  C.  it  dissolves  in  -1.2  parts  of  water.  It 
is  more  soluble  in  alcohol  than  cane-sugar.  It  ferments 
to  alcohol  and  carbon  dioxide  under  the  influence  of 
yeast : 

06H1206  =  2C02  +  2C2H60. 

It  also  undergoes  the  lactic  and  butyric  acid  fermenta- 
tions.— "  Granulated  grape  sugar  "  looks  very  like  cane- 
sugar,  and  is  used  to  adulterate  it.  Many  of  the  cheap 
sugars  contain  10  °/0  to  20  °/0  of  grape-sugar.  It  is 
used  in  the  preparation  of  alcohol,  artificial  wines,  &c. 
Wines  and  liquors  made  from  such  materials  are  poison- 
ous.— Dextrose  has  the  power  of  keeping  in  solution 
some  substances  which  are  insoluble  in  pure  water. 

Experiment  181. — A.dd  a  solution  of  sodic  hydroxide  to  solu- 
tion of  cupric  sulphate.  Cupric  hydroxide  (Cu(OH)2)  is  precipi- 
tated, and  is  not  redissolved  by  excess  of  the  precipitant.  Re- 
peat the  experiment,  with  a  solution  of  cupric  sulphate  contain- 
ing grape-sugar.  A  blue  solution  is  obtained.  Heat  the  solu- 
tion, and  red  cuprows  oxide  is  precipitated,  the  sugar  undergoing 
oxidation.  This  explains  Fehling's  test. 

Experiment  182- — Add  a  few  drops  of  argentic  nitrate  solu- 
tion to  dilute  solution  of  grape-sugar  in  a  1. 1.  and  warm  gently. 
The  inside  of  the  t.  t.  is  silvered. 

Experiment  183. — Heat  a  solution  of  grape-sugar  with  sodic 


LEVULOSE.  227 

hydroxide.     Try  also  with  potassic  hydroxide,  and  with  sodic 
carbonate.     The  solution  of  sugar  turns  brown. 

Tests. — Experiments  i81  and  183  serve  to  distinguish  from 
cane-sugar. 

Note. — Olncosides  are  peculiar  ethereal  salts  of  glucose,  which 
break  up  under  the  action  of  a  ferment  (or  of  dilute  acids)  into 
glucose  and  other  substances,  an  aldehyde  being  very  commonly 
among  the  number. 

2.  LEVULOSE,  <fcc.,  C6H12O6.  Levulose  is  nearly  as 
sweet  as  cane-sugar.  It  was  at  one  time  thought  to  be 
uncrystallisable,  but  it  can  be  crystallised,  although  with 
difficulty.  As  has  been  previously  observed,  it  occurs  in 
ripe  fruits  accompanied  by  an  equal  quantity  of  dextrose. 
It  is  fruit-sugar. — Galactose  is  formed  together  with 
dextrose  by  the  inversion  of  milk-sugar.  It  does  not 
ferment  with  yeast,  but  reduces  Fehling's  solution. — 
Inosite  (C6H12O6.2H2O)  is  found  in  the  heart,  lungs, 
&c.,  of  the  ox,  and  sometimes  in  the  urine  of  man.  It 
does  not  reduce  Fehling's  solution.  It  is  soluble  in  6 
parts  of  water,  and  insoluble  in  alcohol. 


AMYLOSES. 

194.  Starch  (amylum). — C6H10O6.  Found  in  grain, 
potatoes,  arrowroot,  nuts,  and  very  generally  distributed 
throughout  vegetable  tissues.  It  is  the  form  in  which 
plants  store  up  a  reserve  of  food,  just  as  animals  store 
up  fat. — Starch  is  manufactured  from  wheat,  Indian 
corn,  rice,  and  potatoes.  The  materials  are  ground  up 
with  water,  strained  through  sieves,  and  treated  with 
dilute  caustic  soda,  or  fermented.  The  starch  is  allowed 


228  STAEOH. 

to  settle,  washed,  &c.  Sago  is  a  starch  made  from  the 
pith  of  the  sago  palm  of  the  East  Indies.  Arrowroot  is 
starch  prepared  from  the  roots  of  Maranta  arundinacea, 
a  West  Indian  plant.  Tapioca  is  a  similar  preparation 
from  Jatropha  manihot. 

PROPERTIES. — Starch  is  a  white  substance,  forming 
granules  of  peculiar  structure.  Under  the  microscope 
these  granules  show  concentric  layers  around  a  spot  or 
hilum.  The  granules  are  different  for  different  plants, 
and  an  examination  by  the  microscope  at  once  reveals 
the  source  from  which  any  specimen  of  starch  has  come. 
— Starch  is  insoluble  in  water,  but,  when  heated  with  it, 
swells  up  and  forms  a  paste.  Continued  boiling,  or  the 
action  of  acids  or  alkalis,  renders  the  starch  soluble. 

Experiment  184. — Scrape  a  piece  of  potato  and  place  the 
pulp  on  a  muslin  filter.  Pour  a  thin  stream  of  cold  water  on  it, 
catching  the  filtrate  in  a  t.  t.  Allow  the  starch  to  settle,  pour 
off  some  of  the  water,  and  then  heat  to  boiling.  Starch  paste  is 
formed . 

Experiment  185. — To  a  little  cold  starch  paste  (rmidlago 
amyli)  add  a  few  drops  of  iodine  solution.  A  deep  blue  colour 
is  produced.  Heat,  and  the  colour  disappears.  Allow  to  cool 
again. 

Starch  paste  (or  mucilage  of  starch,  as  it  is  called  in 
medicine),  changes  gradually  into  a  solution  of  dextrin. 
This  change  is  hastened  by  boiling  with  dilute  acids,  or 
by  the  action  of  ferments  (malt,  &c.). — Starch  is  used  in 
medicine  as  a  vehicle  for  enemata,  <fcc.  The  blue  starch 
of  the  shops  is  coloured  with  indigo  or  smalt,  and  should 
not  be  used  for  medicinal  purposes. 

Test. — Experiment  185  is  a  very  delicate  test.  The  sub- 
stance should  be  as  cold  as  possible. 


DEXTRIN GLYCOGEN.  229 

Inulin  is  a  compound  (C6H1005)  which  partially  replaces 
starch  in  elecampane,  the  roots  of  dahlia,  &c. 

195.  Dextrin. — C6H10O5.     Dextrin   is   prepared    by 
heating  starch  to  about  250°  C.     It  is  formed  from  starch 
at  lower  temperatures  (up  to  95°  C.)  by  the  action  of 
the  ferment  (diastase)  of  malt,  and  also  by  boiling  with 
dilute  acids. 

Experiment  186. — Boil  in  a  t.  t.  for  several  minutes  a  little 
starch  with  water  and  a  few  drops  of  sulphuric  acid.  A  clear 
solution  is  obtained.  Cool  this  and  test  it  with  iodine. 

Dextrin  is  soluble  in  water,  but  insoluble  in  alcohol. 
By  the  further  action  of  malt  or  dilute  acids  it  is  changed 
first  into  maltose  (C^H^On),  and  then  into  dextrose 
(C6H12<">6). — It  is  made  on  the  large  scale,  and  sold  as 
calcined  farina,  or  British  gum,  a  cheap  substitute  for 
gum  arabic. — Considerable  quantities  of  dextrin  are 
formed  in  the  baking  of  bread.  Toast  is  more  easily 
digested,  therefore,  than  bread.  The  first  stage  in  the 
digestion  of  the  insoluble  starch  is  its  transformation 
into  soluble  dextrin. 

196.  GlycOgen. — C6H100S.     This  name  means  sugar- 
generator.  Glycogen  is  found  in  the  livers  of  most  animals  ; 
it  is  also  present  in  the  blood  and  in  muscles.     Oysters 
contain  a  large  percentage.     It  can  be  extracted  from 
fresh  liver  by  boiling  with  water  or  with  alkalis.     It  is 
precipitated  from  its  aqueous  solution  by  the  addition  of 
alcohol,  in  which  it  is  insoluble. 

PROPERTIES. — A  white  amorphous  powder,  somewhat 
soluble  in  water.  Under  the  action  of  malb,  &c.,  it  com- 
bines with  water  to  form  maltose  and  dextrose.  It  is 


230  GUMS — CELLULOSE. 

generally  supposed  that  the  liver  stores  up  the  surplus 
sugar  of  the  food  in  the  form  of  glycogen.  The  liver 
contains  a  ferment  which  has  the  power  of  transforming 
glycogen  into  maltose  or  dextrose. 

197.  Gums. — C6H10O5.     These  must  be  distinguished 
from  resins,  often  called  gums  in  this  country.     Gum- 
arabic,  gum  tragacanth,  and  bassorin  are    examples  of 
true  gums.     They  are  non-crystalline,  soluble  in  water, 
but  insoluble  in  alcohol.     They  are  converted  into  glu- 
coses by  boiling  with  dilute  acids.      Vegetable  mucilages, 
e.g.,  that  of  linseed,  are  of  a  similar  character. 

1 98.  Cellulose. — C0H1005.     This  is  the  groundwork 
of  vegetable  tissues,  forming  the  walls  of  vegetable  cells. 
Cotton,  hemp,  and  flax  are  nearly  pure  cellulose.     Paper 
is  also  very  pure,  especially  when  prepared  from  cotton. 
Woods  consist  in  large  part  of  cellulose. 

PROPERTIES. — A  white  amorphous  solid,  insoluble  in 
water,  and  in  most  chemical  substances.  It  is,  however, 
soluble  in  ammoniacal  solution  of  cupric  hydroxide 
(Schweizer's  reagent),  in  strong,  hot  solution  of  caustic 
soda  or  potash,  and  in  strong  acids.  These  facts  must 
be  remembered  in  filtering  processes.  It  can  be  con- 
verted into  dextrose  by  dissolving  in  strong  sulphuric 
acid,  diluting,  and  boiling  for  some  time. 

Experiment  187. — Dissolve  a  piece  of  filter  paper  in  a  little 
strong  sulphuric  acid,  dilute,  boil  for  some  time,  and  test  with 
Fehling's  solution. 

Gun-cotton,  pyroxylin,  or  nitro-cellulose,  is  the  trinitrate  of 
cellulate,  C6H7O2(NO3)S,  prepared  by  the  action  of  a  mixture  of 
concentrated  sulphuric  and  nitric  acids  on  cotton  or  wood.  It 
is  very  explosive,  containing  as  it  does  an  oxidisable  and  an 


QUESTIONS    AND    EXERCISES.  231 

oxidising  part.  (Explain). — Celluloid  is  a  preparation  of  nitrates 
of  cellulose  and  camphor,  prepared  by  a  method  similar  to  that 
for  gun-cotton.  It  softens  when  heated  and  can  be  moulded 
into  any  shape. — Collodion  is  a  solution  in  alcohol  and  ether  of 
nitrates  of  cellulose.  It  is  used  in  photography. 


QUESTIONS  AND  EXERCISES. 

1.  Compare  carbon  dioxide  and  carbon  bisulphide,  as  to  their 
compounds. 

2.  What  causes  the  effervescence  of  soda  water  ? 

3.  How  would  you  prove  by  an  experiment  that  carbon  di- 
oxide is  formed  by  the  burning  of  a  candle  ? 

4.  Show  that  most  of  our  materials  for  fires  and  lights  are 
from  a  vegetable  source. ' 

5.  Write  a  graphic  formula  for  acetylene  (C2H2).     With  how 
many  atoms  of  chlorine  would  you  expect  its  molecule  to  com- 
bine ? 

6.  Show  how  green  vitriol,  ferric  chloride,  and  sodic  carbonate 
may  be  used  as  an  antidote  to  poisonous  cyanides. 

7.  Account  for  the  smell  of  ammonia  about  stables  and  water 
closets. 

8.  A  quantity  of  urine  (70  c.  c.)  is  decomposed  by  sodic  hypo- 
bromite,  and  yields  8  c.  c.  of  nitrogen.     How  many  grains  of 
urea  per  gallon  ? 

9.  Write  formulas  for  methyl    nitrate,    sulphate,  and   ortho- 
phosphate  ;  and  for  ethyl  bromide,  acid  sulphate,  and  sulphite. 

10.  What    is    the  practical    distinction    between    primary, 
secondary,  and  tertiary  alcohols  ? 

11.  Calculate  the  specific  weight  of  ether  vapour  (air  =  1). 

12.  Given  zincic  carbonate  and  acetic  acid,  prepare  zincic  acetate. 
Try  it  practically. 

13.  What  chemical  actions  accompany  the  souring  of  milk  ? 

14.  What  is  "  invert  sugar  "  ?     Why  so  called  ? 

15.  Why  is  toast  more  easily  digested  than  bread  ? 


232  COAL    TAR. 


CHAPTER    XIV. 


CARBON  AND  ITS  COMPOUNDS  (CONCLUDED). 

199.  Coal  Tar. — About  thirty  years  ago  coal  tar 
was  an  offensive  waste  product  of  gas  manufacture,  and 
the  problem  was  how  to  get  rid  of  it.     To-day  it  is  the 
substance   for  which  the  coal  is  distilled  in  many  fac- 
tories ;  and  the  gas  manufacturers  regard  it  as  one  of 
the  chief  sources  of  their  revenue.     This  change  is  due 
to  the  discovery  of  the  aniline  dyes,  the  raw  materials 
for  their  manufacture  being  obtained  from  coal  tar. — 
The  tarry  liquid  which  collects  in  the  condensers  of  the  gas- 
works is  redistilled  fractionally.     There  distils  over  first 
a  light  oil  which  floats  on  water ;  this  contains  benzene 
(C6H6),  toluene  (C7H8),  carbolic  acid  (C6H70),  <fec.     Later 
there   distils   a   heavy    oil    containing    xylene   (C8H10), 
mesitylene  (C9H12),  cymene  (C10H14),  naphthalene,  anthra- 
cene, &c. 

200.  The  Benzene,  or  Aromatic,  Series.— We 

have  already  studied  unsaturated  hydrocarbons,  such  as 
ethylene  and  acetylene,  the  marked  characteristic  of 
which  is  the  readiness  with  which  they  combine  with 
such  substances  as  chlorine  to  form  saturated  compounds. 
The  benzene  compounds  are  unsaturated  as  far  as  their 
composition  goes,  but  they  are  characterised  by  the 
difficulty  with  which  anything  is  added  to  their  mole- 
cules. They  yield  substitution,  rather  than  addition, 


BENZENE.  233 

products,  and  in  this  respect  resemble  saturated  hydro- 
carbons.    The  chief  members  of  the  series  ai-e  as  follows  : 

Boiling 
points. 

Benzene,  C6H6 81° 

Toluene,  C7H8 Ill 

Xylenes  (3  isomers),  C8H10 about  140 

fMesitylene ~\ 163 

J  Pseudocumene  j-C9H1 2 151 

[Cumene J 166 

201.  Benzene  (Benzol). — C6H6. 

PREPARATION. — From  the  light  oil  of  coal  tar.  This  is 
washed  first  with  caustic  soda  solution  to  dissolve  out 
carbolic  acid  and  other  acid  impurities,  and  afterwards 
with  dilute  sulphuric  acid  to  remove  basic  substances. 
It  is  then  distilled  fractionally,  and  the  part  coming 
over  between  80°  and  90°  is  used  for  the  preparation 
of  benzene. — It  can  also  be  prepared  by  heating  benzoic 
acid  with  lime  : 

C6HS.COOH  +  CaO  =  C6H6  +  CaCO3. 

This  method  is  analogous  to  that  for  the  preparation  of 
methane. 

PROPERTIES. — A  colourless  liquid  of  specific  weight 
0.89.  It  boils  at  80°.5,  and  has  a  pleasant  aromatic 
odour.  It  burns  with  a  brightly  luminous  flame. 

Experiment  188. — Pour  a  little  benzene  into  a  basin  of 
water  and  set  fire  to  it. 

Experiment  189- — Pour  a  few  drops  of  benzene  into  a  t.  t. 
of  water  heated  nearly  to  the  boiling  point.  Observe  that  the 
benzene  boils. 

Benzene  is  solid  at  0°C. — It  is  a  good  solvent  for  fats, 


234  BENZOL    RING. 

oils,  and  resins. — When  benzene  is  acted  on  by  chlorine, 
substitution  products  are  formed.  There  is  only  one 
monochlor-benzene  (C6H6C1).  If  the  atoms  of  carbon  in 
the  molecule  were  connected  as  they  are  supposed  to 
be  in  the  hydrocarbons  already  studied,  we  should 
expect  the  hydrogen  atoms  to  be  differently  situated  and 
thus  to  permit  the  formation  of  isomeric  derivatives 
by  the  replacement  of  the  hydrogen  atoms.  Now  mono- 
chlor-benzene has  been  prepared  in  a  variety  of  ways, 
some  of  which  lead  to  the  conclusion  that  different 
hydrogen  atoms  have  been  replaced  by  chlorine ;  but 
the  products  are  identical.  It  is  plain,  then  that  the 
six  hydrogen  atoms  are  similarly  situated  in  the  mole- 
cule. Kekule,  of  Bonn,  has  devised  a  structural  for- 
mula showing  this.  It  is  called  the  benzol  ring,  and 
although  it  only  pictures  an  hypothesis,  yet  it  has 
proved  a  powerful  lever  in  the  hands  of  experimenters. 

(6)  H  H(!) 

I  i 

s-^    ^ 


(5)  H— C  C— H  (2) 


Q    Q 

(4)  H  H  (3) 

The  molecule  is  represented  as  being  symmetrical,  each 
carbon  atom  being  united  to  a  hydrogen  atom,  to  another 
carbon  atom  by  a  single  bond,  and  to  a  third  by  a  double 
bond.  (Study  this  formula,  and  observe  that  the  hydro- 
gen atoms  are  represented  as  similarly  situated.) 

When  two  hydrogen  atoms  of  the  benzene  molecule 


SUBSTITUTION    PRODUCTS.  235 

are  replaced,  three  isomeric  di-substitution  products  are 
formed.  If  the  hydrogen  atoms  represented  in  the  above 
formula  be  numbered  consecutively  from  1  to  6,  it  is 
easily  seen  that  (2)  and  (6)  are  similarly  situated  with 
regard  to  (1),  so  that  replacing  (1)  and  (2)  with, 
say,  Cl's,  gives  the  same  formula  as  is  obtained  by  re- 
placing (1)  and  (6).  In  the  same  manner  (3)  and  (5) 
are  similarly  situated;  but,  with  regard  to  (1),  (4)  is 
differently  situated  from  (2),  (3),  (5),  or  (6).  Only 
three  different  formulas  of  di- substitution  products  are 
possible,  and  this  result  is  in  accordance  with  the  facts. 
By  a  series  of  beautiful  experimental  investigations,  the 
relative  positions  of  the  substituted  atoms  have  been 
determined,  and  names  have  been  assigned  accordingly. 
Thus  (1)  (2)  (or  (1)  (6))  di-substitution  products  are 
called  ortho,  e.g.,  ortho-dichlorbenzene  ;  (1),  (3)  products 
are  called  meta,  and  (1),  (4)  para. 

202.  Nitro  -  substitution    Products.  —  When 

strong  nitric  acid  acts  on  benzene,  one  or  more  of  the 
hydrogen  atoms  are  replaced  by  the  monad  radical — NO2 
(nitroxyl).  Thus  : 

(1)  CeBTe  +  HN03  =  C6H5.NO2  +  H2O. 

(2)  G6H6  +  2HN03  =  C6H4.(N02)2  +  2H2O. 

itc.  &c. 

Only  one  mononitrobenzene  is  known,  but  three  isomeric 
dinitrobenzenes  have  been  obtained.  (How  many  tri- 
nitrobenzenes  are  possible  according  to  theory  ]) 

MONONITROBENZENE  (C6H6.NO2)  is  manufactured  on 
the  large  scale  in  the  preparation  of  aniline  (C6H5.NH2). 
It  is  a  yellow,  oily  liquid  smelling  like  oil  of  bitter 
almonds.  It  is  often  sold  as  "  artificial  essence  of  mir- 


236  ANILINE. 

bane,"  or  "artificial  oil  of  bitter  almonds;"  but  as  it  is 
poisonous  and  different  in  its  physiological  action  from 
the  true  oil,  it  should  never  be  used  in  medicine  in 
place  of  the  true  oil. — Mononitrobenzene  boils  at  205°  C. 
Its  specific  weight  is  1.2.  It  is  insoluble  in  water.  By 
the  action  of  nascent  hydrogen  it  is  reduced  to  aniline 
(C6H5.NH2) : 

C6H5.N02  +  3H2  =  C6H5.NH2  +  2H2O 

Nitrobenzene  is  used  to  perfume  cheap  soaps  and  confec- 
tionery, as  well  as  in  the  manufacture  of  aniline.  It 
is  sometimes  mixed  with  oil  of  bitter  almonds  as  an 
adulteration.  Its  presence  can  be  detected  by  reducing 
to  aniline  by  means  of  zinc  and  hydrochloric  acid,  and 
then  treating  with  filtered  solution  of  bleaching  powder, 
with  which  aniline  forms  a  beautiful  purple  colour. 

Experiment  190- — Mix  a  few  drops  of  benzene  with  a  little 
concentrated  nitric  acid.  Warm  very  gently  and  note  the  smell 
of  nitrobenzene. 

203.  Aniline  (Phenylamine).  —  C6H5.NH2.  Was 
first  prepared  from  indigo  \,anil).  It  is  found  in  small 
quantities  in  coal  tar,  wood  tar,  and  bone  oil. 

PREPARATION. — By  the  action  of  nascent  hydrogen  on 
mononitrobenzene.  The  hydrogen  is  evolved  from  iron 
and  acetic  or  hydrochloric  acid.  (Write  the  equations). 

PROPERTIES. — Aniline  is  a  colourless  liquid  of  specific 
weight  1.036.  It  boils  at  184°.5.  It  is  soluble  in  31 
times  its  weight  of  water,  and  the  solution  has  a  weak 
alkaline  reaction.  It  is  more  freely  soluble  in  alcohol. 
— Aniline  unites  with  acids  to  form  salts  : 

C6H5.NH2  +  HC1  =  C6H5.NH3C1. 
Compare  with  NH3  +  HC1  =  NH4C1. 


CARBOLIC   ACID.  237 

Aniline  is  a  substituted  ammonia,  or  amine,  of  the  aroma- 
tic series.  Commercial  aniline  always  contains  toluidine 
(C7H7.NH2\  a  higher  member  of  the  same  series ;  and, 
when  it  is  oxidised  by  means  of  arsenic  acid  or  potassic 
chlorate,  the  beautiful  colour  rosaniline  (C20H19N3)  is 
formed  This  is  the  first  of  the  aniline  dyes,  the  dis- 
covery and  manufacture  of  which  have  revolutionised 
the  dyer's  art.  They  are  comparatively  innocuous  when 
pure,  but  arsenic  acid  is  used  in  their  manufacture,  and 
through  carelessness  or  ignorance  it  is  sometimes  imper- 
fectly separated. 

Tests. — 1.  To  an  aqueous  solution  of  aniline  add  some 
filtered  solution  of  bleaching  powder.  A  purple  colour  results. 

2.  Dissolve  a  drop  of  aniline  in  strong  sulphuric  acid.  (What  is 
formed  ?)  Stir  in  a  porcelain  dish  with  a  drop  of  potassic  bichro- 
mate solution  (K2Cr207).  A  blue  colour  is  produced. 

204.  Carbolic  Acid  (Phenol).  —  C6H5.OH.    This 

is  in  reality  a  tertiary  alcohol,  but  it  has  very  weak 
acid  properties. — It  is  found  in  small  quantities  in  urine. 
It  is  a  product  of  the  distillation  of  wood,  coal,  bones,  &c. 
PREPARATION. — From  the  heavy  oil  of  coal  tar.  This 
is  treated  with  caustic  soda  solution,  which  dissolves  the 
carbolic  acid,  forming  sodic  carbolate,  or  phenolate, 
(C6H5.ONa).  The  solution  forms  a  layer  separate  fiom 
the  oil.  It  is  decanted,  and  is  then  decomposed  by  hy- 
drochloric acid  : 

C6Hs.ONa  +  HC1  =  C6H5.OH  +  NaCl. 

The  carbolic  acid  is  purified  by  distillation. 

PROPERTIES. — A  colourless,  crystalline  solid,  crystal- 
lising in  long  needles.  It  has  a  strong,  somewhat  tarry 
smell,  not  unpleasant  when  the  carbolic  acid  is  pure.  It 


238  CREOSOTE. 

melts  at  35°,  and  boils  at  180°.  It  is  deliquescent, 
uniting  with  a  small  proportion  of  water  to  form  a  liquid. 
This  liquid  is  sparingly  soluble  in  water  (1  in  15),  but 
mixes  with  ether  and  alcohol  in  all  proportions. 

Experiment  191. — To  a  little  solid  carbolic  acid  add  a  drop 
or  two  of  water.  They  combine  to  form  an  oily  liquid.  Add  a 
little  more  water.  The  two  do  not  mix.  (Which  is  the 
heavier  ?)  N  ow  add  much  water,  and  the  carbolic  acid  dis- 
solves. Note  taste  and  smell  of  this  liquid. 

It  is  very  poisonous,  and  the  strong  solution  burns  the 
skin.  It  has  a  burning  taste.  The  antidotes  are  olive  oil, 
castor  oil,  and  a  mixture  of  slaked  lime  with  3  times 
its  weight  of  sugar,  rubbed  up  with  a  little  water. — 
Carbolic  acid  is  extensively  used  as  an  antiseptic  in 
surgery,  &c.  It  was  introduced  into  general  use  by 
Lister,  of  Edinburgh,  who  devised  a  system  by  which 
surgical  operations  were  carried  on  in  an  atmosphere 
containing  no  living  putrefactive  germs.  The  antiseptic 
spray  was  combined  with  a  minute  cleanliness  hitherto 
unknown  in  surgery,  and  in  all  probability  to  this  is  due 

in  a  great  measure  the  success  of  Listerism. 

f 

Test. — Mix  the  liquid  with  one-fourth  its  volume  of  am- 
monia solution  and  add  a  few  drops  of  clear  bleaching  powder 
solution.  The  merest  trace  of  carbolic  acid  will  cause  a  blue 
colour  to  appear. 

205.  Creosote. — Pure  creosote  is  obtained  by  distil- 
ing  beech  and  other  hard  woods.  It  consists  mostly  of  is- 
omeric  cresols,  C6ILjlQgs.  It  resembles  carbolic  acid  in 
smell  and  taste.  Carbolic  acid  is  often  sold  as  creosote  ; 
and  the  so-called  "  coal-tar  creosote "  is  merely  impure 
carbolic  acid. 


PICRIC  ACID.  239 

To  distinguish  carbolic  acid  from  creosote  : 

1.  Dissolve  in  glycerine  and  add  water.     Carbolic  acid 
remains  dissolved.     Creosote  is  precipitated  by  the  water. 

2.  Add  to  the  liquid  to  be  tested  alcoholic  solution  of 
ferric  chloride.     Carbolic   acid   gives   a    brown  colour ; 
creosote,  green. 

3.  With  aqueous  ferric  chloride,  carbolic  acid  turns 
blue ;  creosote  is  unchanged. 

206.  Picric   Acid    ( Trinitrophenol,    or    Carbazotic 
A  cid).—G6IL2(  OH).  (NO2)3. 

PREPARATION. — Experiment  192. — Dissolve  some  carbolic 
acid  in  dilute  nitric  acid  and  add  drop  by  drop  about  half 
the  volume  of  strong  nitric  acid.  Dilute  with  cold  water,  and 
picric  acid  is  precipitated. 

PROPERTIES. — A  yellow  crystalline  solid  sparingly 
soluble  in  cold,  easily  in  hot  water,  and  in  alcohol  or  ether. 
It  has  a  bitter  taste  and  is  poisonous.  It  explodes  when 
heated  rapidly.  Ammonic  picrate  (NH4.C6H2N307)  is 
used  as  an  explosive.  Picric  acid  has  strong  colouring 
power,  and  is  a  good  dye  for  silk  and  wool. 

207.  Benzylic  Alcohol.— C6H6.CH2OH.  —  Found 
in  balsam  of  Tolu  and  Peru,  and  in  storax,  as  ethereal 
salts   of  benzoic  and   cinnamic  acids.     When   oxidised 
it  gives   first  benzoic  aldehyde   (C6H6.COH),  and  then 
benzoic  acid  (C6H6.COOH). 

208.  Benzoic  Aldehyde.— C6H6.C=O.     A  com- 

H 

ponent    of    amygdalin,    a    glucoside    found    in    bitter 


240  BITTER   ALMOND    OIL. 

almonds,  in  laurel  leaves,  cherry  kernels,  &c.  Amyg- 
dalin  ferments  under  the  influence  of  synaptase  (a  fer- 
ment present  in  the  almond),  forming  benzoic  aldehyde, 
hydrocyanic  acid,  and  dextrose  : 

OaoHgyOnN  +  2H20  =  C7H6O  +  HCN  +  2C6H12O6. 

Oil  of  bitter  almonds  is  prepared  by  fermenting  bitter 
almonds.  The  crude  oil  contains,  besides  benzoic  alde- 
hyde, hydrocyanic  acid,  and  is  therefore  very  poisonous. 
It  is  purified  by  distilling  with  milk  of  lime  (to  retain 
the  prussic  acid),  and  fused  calcic  chloride  (to  retain  the 
water).  Pure  benzoic  aldehyde  is  thus  prepared. 

PROPERTIES. — A.  colourless,  bright  liquid  of  pleasant 
aromatic  odour.  Specific  weight  =  1.054.  It  boils  at 
180°  C.  It  dissolves  in  30  times  its  weight  of  water, 
and  mixes  in  all  proportions  with  alcohol  and  ether. — 
It  is  not  poisonous ;  but  the  so-called  "  artificial  oil " 
(nitro-benzene)  is  poisonous.  A  delicate  test  for  nitro- 
benzene in  oil  of  bitter  almonds  is  as  follows : — Treat 
with  zinc  and  dilute  hydrochloric  acid,  filter,  and  add 
solution  of  potassic  chlorate.  A  mauve  colour  appears, 
if  nitro-benzene  is  present. — Benzoic  aldehyde  is  easily 
oxidised  to  benzoic  acid  (C6HS.COOH). 

209.  Benzoic  Acid. — C6H6.CO.OH.  Benzoic  acid 
is  found  in  gum  benzoin,  in  balsams  of  Peru  and  Tolu, 
&c.  It  is  also  present  in  the  urine  of  some  animals. 

PREPARATION. — 1.  By  sublimation  from  gum  benzoin, 
formerly  the  chief  source  of  benzoic  acid. 

2.  From  hippuric  acid  (C9H9O3]S"),  which  can  be  ob- 
tained in  large  quantities  from  the  urine  of  herbivorous 
animals.  This,  when  boiled  with  a  dilute  acid,  decom- 


BENZOIC   ACID.  241 

poses  into  benzole  acid  and  amido-acetic  acid,  or  glycocoll 
(NH2CH2.COOH) : 

C9H9O3N  +  H2O  =  C6H5.COOH  +  NH2.CH2.COOH. 

A  considerable  proportion  of  the  benzole  acid  of  com- 
merce is  derived  from  this  source. 

3.  A  good  deal  of  benzole  acid  is  now  made  by  oxid- 
ising toluene  (C6H5.CH3)  : 

C6H5.CH3  +  3O  =  C6H5.COOH  +  H2O, 

also  from  naphthalene  (C10H8),  which,  when  oxidised, 
yields  a  dibasic  acid,  phthalic  acid  (C6H4(COOH)2). 
This  acid  when  heated  with  lime  breaks  up  into  benzole 
acid  and  calcic  carbonate  : 

CaO  +  C6H4(COOH)2  =  CaCO3  +  C6H6COOH. 

Experiment  193. — Carefully  heat  a  small  bit  of  gum  benzoin 
in  a  t.  t.  Benzoic  acid  sublimes  and  collects  in  feathery  crystals 
on  the  cooler  parts  of  the  tube.  Note  the  odour. 

PROPERTIES. — A  white  solid  in  pearly  scales  or 
feathery  needles,  of  aromatic  odour.  It  melts  at  121°  C. 
It  dissolves  easily  in  alcohol. 

Experiment  194. — Try  to  dissolve  a  little  benzoic  acid  in  a 
small  quantity  of  water.  Heat  the  water.  Taste  the  solution. 
Test  with  litmus.  Heat  a  small  particle  of  benzoic  acid  on 
mica,  and  inhale  the  fumes. 

BENZOATES. — Benzoic  acid  is  monobasic,  and  its  salts 
are  mostly  soluble  in  water. 

Experiment  195. — Colour    some  solution  of   ammonia  in  a 
porcelain  dish  with  litmus,  and  gradually  add  benzoic  acid  to  it 
until  it  is  neutralised.     Evaporate  on  the  water-bath.     White 
17 


242  SACCHARINE. 

crystals  of  ammonic  benzoate  (NH4.C?H4OS)  are  left.  It  is 
necessary  to  add  a  little  ammonia  from  time  to  time  during 
evaporation  so  as  to  keep  the  solution  alkaline.  When  salts  of 
ammonia  are  evaporating  a  little  ammonia  always  escapes. 

A.mmonic  benzoate  is  used  in  medicine.  When  taken 
into  the  system  it  is  eliminated  in  the  urine  as  hippuric 
acid,  rendering  the  urine  strongly  acid. — Sodic  benzoate 
(Na.C7H5O2)  can  be  prepared  in  the  same  way  as  am- 
monic benzoate.  It  is  nsed  as  a  medicine. 

Experiment  196. — Dissolve  in  a  little  water  the  ammonic 
benzoate  prepared  in  Experiment  195.  To  a  small  portion  of 
the  solution  add  dilute  hydrochloric  acid.  Benzoic  acid  is  pre- 
cipitated. To  another  portion  add  neutral  solution  of  ferric 
chloride.  Ferric  benzoate  (Fe^Bz,,)  is  precipitated.  Examine 
it  carefully,  noting  its  colour,  &c. 

Tests. — See  last  experiment. 

210.  Saccharine.    C6H4~SQ  ^:NH.    (Benzoyl  sul- 

phonic  imide. — This  substance  is  related  to  benzoic  acid. 
Benzosulphonic  add  has  the  formula  C6H4~o';;  QTT.  If 

the  two  hydroxyls  are  replaced  by  the  dyad  radical 
=NH,  the  formula  for  saccharine  is  obtained. — This  re- 
markable substance  was  discovered  a  few  months  ago  by 
Dr.  Fahlberg. 

PROPERTIES. — A  white  crystalline  solid,  "  220  times 
as  sweet  as  cane-sugar."  It  has  antiseptic  properties, 
and  is  said  to  be  harmless  when  taken  into  the  system, 
passing  away  in  the  urine  unchanged.  It  is  difficultly 
soluble  in  cold  water,  more  freely  in  hot  water,  in  alco- 
hol, and  in  ether.  It  is  proposed  to  use  it  instead  of 
cane-sugar  in  cases  of  diabetes  mellitus. 


GALLIC   ACID.  243 

211.   Salicylic  Acid  (Ortho-hydroxy-benzoic  acid). 
COC 
OH 

— CO.OCH. 


C6H4 .-.TT      .       Occurs  in   several   plants,  e.g.,    winter- 


green,  as  methyl  scdicylate  (C6H4_Qg     /t±3),  or  "  oil  of 

wintergreen."  It  is  now  prepared  from  carbolic  acid,  as 
well  as  by  the  oxidation  of  salicin,  a  glucoside. — It  is  a 
monobasic  acid,  powerfully  antiseptic,  preferable  in  some 
respects  to  carbolic  acid,  being  not  so  injurious  to  the 
tissues. 

212.  Gallic  Acid  ( '  Trihydroxy-benzoic  add),   C6H2 
(OH)3.COOH.     Is  present  as  a  glucoside    in  oak-galls, 
tea,  various  barks,  <kc. 

PREPARATION. — From  oak-galls,  which  are  powdered, 
moistened  with  water,  and  fermented  5  or  6  weeks. 
The  gallic  acid  is  then  dissolved  out  with  boiling  water. 

PROPERTIES. — A  light,  white,  crystalline  solid.  It 
dissolves  in  100  parts  of  cold,  and  in  3  parts  of  hot, 
water ;  it  is  freely  soluble  in  alcohol,  and  sparingly  in 
glycerine  and  in  ether.  Gallic  acid  is  a  monobasic  acid. 

Tests. — 1.  With  solution  of  ferric  chloride  it  gives  a  blue- 
black  precipitate,  soluble  in  excess. 

2.  It  does  not  coagulate  solution  of  gelatine  or  white  of  egg. 

213.  TanniC    Acid. — C14H10O9.     This  is    an    anhy- 
dride of  gallic  acid,  and  its  molecule  is  formed  from  two 
of  gallic  acid  by  the  substraction  of  a  molecule  of  water  : 

Gallic  acid.  Tannic  acid. 

2C7H605  -  H20  =  C14H1009. 

PREPARATION. — From  gall  nuts  by  steeping  them  in 
water  3  or  4  days  and  then  extracting  with  ether. 


244  TERPENES. 

PROPERTIES. — A  white  or  yellowish  solid,  of  acid, 
astringent  taste,  soluble  in  water,  and  in  a  mixture  of 
water  with  alcohol  or  ether. — Tannin  is  the  glucoside, 
present  in  oak-galls,  &c.,  which  produces  gallic  and 
tannic  acids  by  its  decomposition.  Glucose  is  formed 
at  the  same  time,  and  commercial  tannic  acid  generally 
contains  glucose.  Tannic  acid  forms  gallic  acid  when 
boiled  with  dilute  acids. 

Tests — 1.  With  aqueous  solution  of  ferric  chloride,  a  black 
colour  is  produced  (ink). 

2.  With  ferrous  sulphate,  it  slowly  blackens. 

3.  Tannic  acid  coagulates    solution   of  gelatine  or  white  of 
egg.     This  illustrates  the  process  of  tanning. 

NOTE. — Tannic  acid  is  a  good  local  astringent,  while 
gallic  acid  is  a  remote  astringent. 

214.  Terpenes. — C10H16.  These  form  a  series  of  is- 
omeric  hydrocarbons  found  in  pines,  firs,  and  other  trees. 
— Turpentine  is  an  oily  liquid  (like  honey),  which  flows 
from  incisions  in  firs,  larches,  &c.  When  distilled  it 
yields  "oil,"  or  "spirits"  of  turpentine,  and  a  rosin 
(resin)  which  remains  in  the  retort.  Rectified  oil  of  tur- 
pentine is  prepared  by  mixing  the  crude  oil  with  caustic 
soda  and  distilling  it.  It  is  a  limpid,  colourless  liquid 
of  pungent  smell  and  bitter  taste.  When  exposed  to 
the  air  it  partly  volatilises  and  partly  oxidises  to  resin, 
it  dissolves  resins,  and  is  hence  used  in  preparing  var- 
nishes.— Resins  are  oxidation  products  of  terpenes.  Bal- 
sams are  similar  compounds.  Resins  consist  largely  of 
acids,  and  form  soaps  with  alkalis.  They  are  used  in 
soap-making. 


CAMPHOR.  245 

215.  Camphor.  —  C10H16O.     Is    prepared   by    distil- 
ling camphor   wood   with  water.     The  camphor  laurel 
grows  in  China  and  Japan. 

PROPERTIES. — A  white,  crystalline  solid,  melting  at 
175°.  It  volatilises  slowly  at  ordinary  temperatures. 
It  is  slightly  soluble  in  water  (40  grains  in  a  gallon), 
freely  in  alcohol,  and  in  oils.  Specific  weight  =  0.98. — • 
Liquid  camphor,  or  camphor  oil  (CooH^O)  is  the  essential 
oil  of  the  camphor  laurel.  It  yields  camphor  when  oxid- 
ised by  exposure  to  air. — Borneo  camphor  (C10H1SO)  is 
obtained  from  Dryobalanops  camphor  a.  "  Oil  of  Borneo 
camphor  "  (so-called)  is  in  reality  a  terpene. 

216.  Cinnamic  Acid—  C6H5  -  CH=CH  -  COOH, 
or  C9H8O2.     Found  in  liquid  storax,  and  in  balsams  of 
Peru  and  Tolu.     It    can    be  prepared   by  oxidation   of 
oil  of  cinnamon,  which  consists  largely  of  cinnamic  alde- 
hyde (C9H8O). 

217.  Essential  Oils  (volatile  oils)  — The   essential 
oils   are   the    substances    which    impart    fragrance    and 
flavour  to  different  parts  of  plants.     They  generally  con- 
tain compounds  resembling  camphor  or  turpentine,  along 
with  ethereal  salts  of  benzoic,  cinnamic,  and  other  aro- 
matic acids.     They  are  obtained  from  the  leaves,  flowers, 
fruit,  &c.,  of  plants,  by  distilling  with  water,  or  by  ex- 
pressing.    Oil  of  cinnamon  and  oil  of  cloves  are   good 
examples. 

Experiment  197. — Distil  some  ground  cloves  with  water  and 
obtain  oil  of  cloves.  Receive  in  a  cold  t.  t. 


246  INDIGO. 

218.  IndigO  Blue  (Indigotin).— C16H10N2O2.     Pre- 
pared from   certain  plants    which   contain    a  glucoside, 
indican.     This  yields  indigo  when  fermented  and  oxid- 
ised by  the  air.     Indigo  is  also  made  artificially  by  a 
complicated  process,  including  the  formation  of  cinnamic 
acid. 

PROPERTIES. — A  dark  blue  solid,  insoluble  in  water, 
in  dilute  acids,  alkalis,  and  ether ;  but  soluble  in  paraffin 
oil  and  in  hot  alcohol.  When  indigo  is  heated  it  vola- 
tilises, forming  a  beautiful  purple  vapour.  If  it  is  pure 
(as  it  rarely  is),  there  is  no  residue. — Indigo  is  soluble, 
in  concentrated  sulphuric  acid,  forming  sulphindylic,  or 
sulphindigotic  acid.  A  dilute  solution  of  this  is  used  as 
a  test  for  nitric  and  chloric  acids,  which  oxidise  it  and 
destroy  the  colour. — Indigo  is  reduced  to  indiyo  white 
by  green  vitriol  and  other  reducing  agents.  Indigo 
white  is  soluble  in  water,  a  very  important  fact  for  the 
dyer.  Calicos,  &c.,  are  printed  or  dyed  with  the  colour- 
less solution  of  indigo  white  and  then  exposed  to  the  air. 
The  oxygen  of  the  air  unites  with  the  indigo  white 
forming  the  insoluble  indigo  blue,  and  thus  a  fast  colour 
is  produced. — Indigo  is  often  adulterated  with  Prussian 
blue,  which  leaves  a  bulky  reddish  residue  when  heated. 

Experiment  198. — Heat  a  little  indigo  in  a  porcelain  dish  or 
on  mica. 

Experiment  199. — Try  to  dissolve  a  little  indigo  in  water. 
Add  some  solutions  of  ferrous  sulphate  and  caustic  soda,  and 
heat.  Steep  a  piece  of  white  cotton  in  the  solution  and  expose 
it  to  the  air  for  some  time. 

219.  Naphthalene.— C10H8.  Naphthalene  is  related 


GLUCOSIDES.  247 

to  benzene  in  the  manner  indicated  by  the  structural 
formula  : 

H  H 

I  I 

C  C 

H—C  C  C— H 

I  II  I 

H—C  C  C— H 

C  C 

I  I 

H  H 

It  is  prepared  from  the  heavy  oil  of  coal-tar  by  frac- 
tional distillation. — It  is  a  colourless  solid,  crystallising 
in  thin  plates.  It  melts  at  79°.2  and  boils  at  216°.6. 
It  is  insoluble  in  water,  but  soluble  in  alcohol.  When 
oxidised  it  yields  phthalic  acid,  CgH^COOH)^  from 
which  benzoic  acid  is  manufactured. 

220.  Anthracene. — CUH10.     Is  also  obtained  from 
the  heavy  oil  of  coal-tar.     It  is  the  starting  point  in  the 
preparation  of  alizarin,  or  artificial  madder. 

221.  GlUGOSides. — These   are  substances   found  in 
plants,  which  readily  undergo  fermentation,  with  a  glu- 
cose as  one  product.     The  other  products  of  the  fermen- 
tation are  various,  but  generally  an  aldehyde  or  an  alcohol. 
The  ferment  is  present  in  the  same  plant  with  the  glu- 
coside,  but  appears  to  be  enclosed  in  separate  cells,  and 
thus  does  not  set  up  fermentation  until  the  cells  are 
broken  down  by  grinding,  &c. 

1.  AMYGDALIN  has  been  already  described  (Art.  208). 


248  SALICIN — DIGITALIN. 

2.  SALICIN  (C13H1807).  Found  in  willow  bark  and  in 
poplar  leaves  and  bark. 

PREPARATION. — Boil  willow  or  poplar  bark  with  milk 
of  lime,  filter,  decolourise  with  bone-black,  evaporate  the 
filtrate  to  dryness,  and  extract  the  salicin  with  alcohol. 

PROPERTIES. — A  white  solid,  soluble  in  alcohol  and 
in  hot  water,  but  not  in  ether.  It  dissolves  readily  in 
solutions  of  alkalis  or  of  alkaline  carbonates. 

Experiment  200. — Boil  a  little  salicin  for  some  time  with 
water  and  a  few  drops  of  sulphuric  acid,  and  then  test  for 
glucose. 

When  salicin  ferments,  it  unites  with  water,  thus  form- 
ing salicylic  alcohol  and  dextrose  : 

C13H1807  +  H2O  =  C6H4{CHaOH  +  C6H12O6. 

Experiment  201; — Dissolve  a  little  salicin  in  solution  of 
potassic  carbonate.  Neutralise  with  hydrochloric  acid.  The 
salicin  is  precipitated. 

A  solution  of  salicin  with  potassic  carbonate  is  said  to 
be  valuable  in  cases  of  diabetes. 

Tests. — 1.  Mix  on  a  porcelain  plate  with  concentrated  sul- 
phuric acid.  A  red  colour  is  produced. 

2.  See  Experiment  200. 

3.  Heat  with  potassic  bichromate  and  dilute  sulphuric  acid, 
and  observe  the  odour  of  salicylic  aldehyde.     It  is  like  that  of 
heliotrope. 

3.  DIGITALIN. — A  mixture  of  several  glucosides.  It  is 
the  active  principle  of  foxglove.  When  boiled  with  dilute 
acids  it  yields  glucose  and  digitaliretin. — Digitaliii  is  a 
white,  inodorous  solid,  of  very  bitter  taste.  It  is  almost 


ALKALOIDS.  249 

insoluble  in  water  and  in  ether,  but  soluble  in  alcohol 
and  in  acids.  It  leaves  no  residue  when  burned.  It  is 
poisonous,  one-sixteenth  of  a  grain  being  sometimes  fatal. 

Test. — Mix  with  weak  aqueous  solution  of  dried  ox-bile,  and 
add  concentrated  sulphuric  acid.  A  deep  red  colour  is  formed. 

4.  JALAPIN  (C^H^O^),  glycyrrhizin  (C^HggOg),  and 
helleborin  (C^H^O^)  are  glucosides  found  in  jalap,  liquor- 
ice, and  the  hellebore  respectively. — Cerebrin  is  a  gluco- 
side  present  in  the  brain  and  other  nervous  tissues. 

222.  Alkaloids. — As  the  name  implies,  alkaloids  are 
substances  having  the  properties  of  alkalis,  i.e.,  of  the 
volatile  alkali,  ammonia.  They  are,  in  fact,  amines  the 
molecules  of  which  are  of  complex  structure  and  in 
most  cases  not  yet  known.  They  are  for  the  most  part 
products  of  vegetable  growth.  Their  names  are  gen- 
erally derived  from  those  of  the  plants  from  which  the 
alkaloids  are  obtained,  the  terminations  -ine  and  -ia 
being  used  interchangeably,  e.  g.,  strychnine,  or  strychnia, 
derived  from  Strychnos  nux  vomica. — The  alkaloids  all 
contain  nitrogen,  and  many  of  them  oxygen,  in  addition 
to  carbon  and  hydrogen.  Those  which  contain  no 
oxygen  are  mostly  liquids,  which  can  be  distilled  ;  those 
containing  oxygen  are  crystalline  solids  and  cannot  be 
distilled.  Most  of  them  are  insoluble,  or  sparingly 
soluble,  in  water,  but  unite  with  acids  to  form  soluble 
salts.  The  alkaloids  dissolve  in  alcohol,  ether,  chloro- 
form, benzene,  <fec.  Their  taste  is  generally  intensely 
bitter. — They  are  usually  extracted  from  the  plants  with 
water  or  dilute  acid ;  and  the  solution  is  then  decom- 
posed with  an  alkali.  The  volatile  alkaloids  are  dis- 
tilled, while  the  solid  ones  are  filtered  off. 


250  MORPHINE. 

General  Tests. — 1.  All  alkaloids  give  a  precipitate  with 
phosphomolybdic  acid. 

2.  All  alkaloids  give  a  precipitate  with  solution  of  potassic 
mercuric  iodide  (HgI2.2KI). 

3.  Most  of  the  alkaloids    give   a  precipitate    with   potassic 
iodide  solution  of  iodine, 

223.  Volatile  Alkaloids.— Conine  (C8HU.NH)  is 
the  active  principle  of  poisonous  hemlock  (Conium  macu- 
latum).     It  is  a  colourless  liquid,  boiling  at  168°  C.     It 
dissolves  in  100  parts  of  water,  forming  a  strongly  alka- 
line solution.     It  unites  with  acids,  forming  salts,  e.g., 
C8HU.NH2C1.       It    is    a    narcotic    poison.  —  Nicotine 
(C10HUN2)  is  found  in  tobacco,  from  which  it  can  be  ob- 
tained by  distilling  with  solution  of  caustic  potash.     It 
is  a  narcotic  poison. 

224.  Morphine.  —  C17H19N03.H2O.     This  is  found 
in  opium,  associated  with  narcotine  and  other  alkaloids. 
It  is  combined  with  meconic  acid. 

PREPARATION. — :The  meconate  is  dissolved  out  with 
water  and  treated  with  solution  of  calcic  chloride.  Mor- 
phine chloride  remains  in  solution  while  calcic  meconate 
is  precipitated.  From  this  solution  the  morphine  is  pre- 
cipitated by  ammonia. 

PROPERTIES. — A  white  crystalline  solid,  insoluble  in 
ether,  sparingly  soluble  in  water  and  in  cold  alcohol, 
and  soluble  in  hot  alcohol  and  in  dilute  acids.  It  unites 
with  acids,  forming  crystallisable,  soluble  salts.  The 
chloride  (hydrochlorate,  or  muriate],  sulphate,  and  acetate 
are  used  in  medicine. — Morphine  and  its  salts  are  nar- 
cotic poisons. 


QUININE CINCHONINE.  251 

Tests. — 1.  If  solid,  add  a  little  water ;  if  liquid,  evaporate 
nearly  to  dryness.  Stir  with  a  drop  of  neutral  ferric  chloride 
solution.  A  dirty  blue  colour  appears. 

2.  Moisten  the  solid  substance  with  strong  nitric  acid.  It 
gives  a  bright  orange-red  colour. 

225.  The  Alkaloids  of  Peruvian  Bark  (Cin- 
chona). 

1.  QUININE  (C^H^NA-SH^O).     The  alkaloid  itself  is 
a  white,  crystalline  powder.     It  is  sparingly  soluble  in 
water  (1  in  900),  but  easily  soluble  in  alcohol  and  ether. 
The  solutions  have   an  alkaline  reaction.     Sulphate  of 
quinine  is  the  salt  generally  used  in  medicine,  but  pre- 
parations of  citrate  and  chloride  are  also  used.     The  sul-. 
phate  is  only  sparingly  soluble  in  pure  water,  but  dis- 
solves readily  in  water  containing  sulphuric  acid  (a  solu- 
ble acid  salt  being  formed). 

Experiment  202- — Try  to  dissolve  quinine  or  quinine  sul- 
phate in  water.  Add  sulphuric  acid.  Note  the  fluorescence  of 
the  solution. 

Experiment  203. — Add  an  alkali  to  a  solution  of  quinine 
sulphate.  What  is  precipitated  ? 

Quinine  and  its  salts  are  antipyretic  (lower  the  tempera- 
ture when  taken  internally) ;  and  also  antiperiodic  (pre- 
vent the  return  of  periodic  fevers,  <kc.) 

Test. — Add  chlorine  water  and  then  a  considerable  quantity 
of  ammonia  solution.  A  bluish-green  colour  appears.  If  potas- 
sic  ferrocyanide  be  added  before  the  ammonia,  a  red  colour  ap- 
pears for  a  moment,  but  soon  fades. 

Note. — Quinine  and  its  salts  are  sometimes  adulterated  with 
salicin. 

2.  CINCHONINE  (C^H^N-P)  is  less  soluble  in  alcohol 
than  quinine,  and  is  thus  separated  from  it  in  the  process 


252  STRYCHNINE — COCAINE. 

of  preparation.  It  is  similar  to  quinine  in  its  properties, 
but  is  insoluble  in  ether.  It  is  not  so  good  a  febrifuge 
as  quinine. 

Tests. — Cinchonine  does  not  give  the  greenish-blue  colour 
with  chlorine  water  and  ammonia.  It  is  sometimes  used  to 
adulterate  quinine.  To  detect  it,  make  the  following  test : 
"  Into  a  glass  tube  or  bottle  put  ten  grains  of  the  suspected  salt, 
dissolve  in  10  minims  of  dilute  sulphuric  acid,  and  60  minims  of 
distilled  water  ;  to  this  add  150  minims  of  pwre  ether,  3  minims 
of  spirits  of  wine,  and  40  minims  of  a  solution  of  soda  (1  of 
caustic  soda  in  12  of  water).  Agitate  well  and  set  aside  for 
12  hours,  when,  if  the  slightest  trace  of  quinidine  or  cinchonine 
be  present,  they  will  be  seen  at  the  line  of  separation  between 
the  ether  and  the  solution  of  sodium  sulphate." 

3.  QUINIDINE,  cinchonidine,  &c.,  are  other  alkaloids 
found  in  Peruvian  bark. 

226.  Strychnine. — C21H22N2O2.   Is  found  in  Strych- 

nos  nux  vomica,  along  with  brucine  (C23H26N2O4). — 
Strychnine  is  a  white  crystalline  solid,  sparingly  soluble 
in  water  and  intensely  bitter  in  taste.  It  is  soluble  in 
spirits  of  wine,  but  not  in  absolute  alcohol  or  ether. 
The  solution  generally  used  in  medicine  is  made  with 
hydrochloric  acid,  rectified  spirits,  and  water.  It  is 
very  poisonous. 

Tests. — 1.  Dissolve  the  solid  in  pure  sulphuric  acid  and  stir 
with  a  drop  of  potassic  bichromate  solution.  A  violet  colour  is 
produced,  which  soon  changes  to  red  and  yellow. 

2.  The  solution,   when  treated  with   hydrochloric  acid   and 
mercuric  chloride  (HgCl2),  gives  a  clotted  white  precipitate. 

3.  Pure  strychnine  gives  no  colour  with  nitric  acid.     Brucine 
is  turned  deep  red. 

227.  Oocaine. — C17H21NO4.     Obtained    from    cocoa 
leaves  (Erythroxylon  coca).     The  hydrochlorate  (C17H21 


ARTIFICIAL    ALKALOIDS.  253 

NO4.HC1)  is  used  as  a  local  anaesthetic.  When  a  little 
of  it  is  put  into  the  eye,  it  causes  insensibility  to  pain  in 
that  part,  and  operations  can  thus  be  performed  without 
administering  chloroform  or  ether. 

228.  Atropine,  (C17H23NO3),  and  hyoscyamine 

(C15H23NO3),  are  poisonous  alkaloids  found  in  common 
plants ;  atropine  in  belladonna,  the  thorn-apple,  &c.,  and 
hyoscyamine  in  henbane. 

229.  Artificial  Alkaloids.  —  Several    antipyretic 
alkaloids  have  lately  been  made  by  synthesis  from  coal- 
tar   products.     They   promise   to   be  very  valuable   as 
medicines. 

1.  KAIEINE  (C10H13NO)  is   made   from  aniline  by  a 
series  of  rather  complicated  reactions.     It  unites  with 
acids  to  form  soluble  salts.     The  chloride  (C10HUNO.C1) 
is  sold  as  a  substitute  for  quinine.     It  lowers  the  tem- 
perature in  fevers,  but  this  action  is  only  of  short  con- 
tinuance.    On  the  other  hand,  it  has  none  of  those  un- 
pleasant effects  associated  with  the  antipyretic  action  of 
quinine. 

2.  ANTIPYRINE  (CnH12N"2O)  is  also  manufactured  from 
aniline  as  a  starting  point.     It  is  a  powerful  antipyretic, 
but  not  antiperiodic.     It  is  a  white,  tasteless,  odourless, 
crystalline  solid,  easily  soluble  in  cold  water.     It  is  used 
uncombined  with  acids. 

3.  THALLINE  (C10HnNO5),  a  recently  discovered  arti- 
ficial alkaloid,  is  said  to  be  a  specific  for  yellow  fever. 

There  are  many  alkaloids  prepared  from  coal-tar,  the 
physiological  actions  of  which  have  not  been  investigated. 
The  results  already  obtained  show  the  fruitfulness  of  a 


254  ALBUMINOIDS. 

combination  of  the  chemist's  investigations  with  those  of 
the  physician.  An  illustration  of  the  necessity  for  this 
is  the  fact  that  cocaine  was  known  ten  years  before  the 
peculiar  power  of  its  hydrochlorate  was  discovered. 


230.  Albuminoids,    or   Proteids.  —  These  are 

nitrogenous  substances  of  a  very  complex  character. 
They  are  the  basis  of  all  living  matter.  They  are 
similar  in  composition,  but  vary  slightly.  They  contain 
carbon,  hydrogen,  oxygen,  nitrogen,  sulphur,  and,  gen- 
erally, phosphorus.  Their  percentage  composition  is  as 
follows : 

Carbon from  51. 5  to  54. 5 

Hydrogen "       6.9"     7.3 

Oxygen.' "     20.9  "  23.5 

Nitrogen "     15.2  "  17.0 

Sulphur "      0.3  "     2.0 

Phosphorus "       0.4 

The  proteids  are  built  up  by  plants  out  of  the  simpler 
compounds  which  form  their  food.  Animals  have  no 
power  of  synthesising  them,  and  must  therefore  obtain 
them  ready-formed  from  plants. 

Experiment  204. — Beat  up  the  white  of  an  egg  with  water, 
allow  the  solution  to  settle,  and  pour  off  the  clear  liquid.  Heat 
a  small  quantity.  It  coagulates  before  it  begins  to  boil.  Try 
another  portion  after  adding  a  little  caustic  soda.  Test  other 
portions  with  nitric,  hydrochloric,  and  sulphuric  acids,  and 'with 
alcohol ;  also  with  acetic  and  tartaric  acid.  To  another  portion 
add  a  very  small  drop  of  cupric  sulphate  solution  and  then  a 
large  quantity  of  caustic  soda.  A  purple  colour  is  produced. 
This  reaction  is  characteristic  of  proteids. 

Experiment  205- — To  small  quantities  of  the  solution  of 
white  of  egg  add  solutions  of  plumbic  acetate,  cupric  sulphate, 


QUESTIONS   AND    EXERCISES.  255 

and  mercuric  chloride  (corrosive  sublimate).  Collect  the  pre- 
cipitates on  filters,  and  wash  them  three  or  four  times  by  pour- 
ing distilled  water  on  them.  Then  pour  a  little  ammonic  sul- 
phide solution  on  each  of  them.  They  are  blackened,  showing 
the  presence  of  the  metallic  salts.  This  explains  the  action  of 
white  of  egg  as  an  antidote. 

Proteids  are  all  colloid  substances,  and  have  very  little 
power  of  passing  through  animal  membranes.  Hence, 
albuminuria  indicates  something  radically  wrong  in  the 
kidneys. — Many  proteids  are  soluble  in  pure  water,  but 
others  require  a  small  quantity  of  sodic  chloride  to  keep 
them  in  solution.  After  they  have  been  coagulated, 
they  can  be  dissolved  by  the  action  of  dilute  acids 
aided  by  a  ferment,  as  in  the  process  of  digestion.  Al- 
kalis keep  in  solution  many  proteids  insoluble  in  neutral 
or  acid  solutions.  Thus  milk,  naturally  alkaline,  coagu- 
lates when  it  becomes  sour. — Proteids  are  very  unstable 
chemical  compounds,  readily  undergoing  fermentations 
and  other  chemical  changes. 


QUESTIONS  AND  EXERCISES. 

1.  Compare  the  preparation  of  benzene  with  that  of  marsh 
gas, 

2.  Explain  the  action  of  milk  of  lime  in  purifying  oil  of  bitter 
almonds. 

3.  Write  the  formulas  for  chloride,  sulphate,  and  acetate  of 
morphia.     (Its  molecule   is  equivalent  to   one  molecule  of  am- 
monia). 

4.  How  would  you  test  for  salicin  in  a  specimen  of  quinine  ? 

5.  How  would  you  distinguish  between  specimens  of  morphine 
and  brucine  ? 

6.  Calculate  the  percentage  composition  of  acetylene  (CaHa), 
and  of  benzene  (C6H6). 

7.  Is  carbolic  acid  a  true  acid  ?    (Note. — Ethyl  alcohol  forms 
alcoholates,  C2H6ONa,  &c.) 


256  SILICON. 

8.  The  flowers,  &c.,  of  the  Cherry  Laurel  contain  amygdalin. 
Is  that  plant  poisonous  ? 

9.  What  is  the  relation  of  saccharine  to  benzoic  acid  ?    Would 
you  expect  it  to  be  nutritive  ? 

10.  Why  is  indigo-blue  reduced  to  indigo-white  in  dyeing  ? 
How  is  it  re-oxidised  ? 

11.  Indigo  is  often  adulterated   with   Prussian  blue.     How 
would  you  detect  this  ? 

12.  How   would  you   make  an  aqueous  solution  of  quinine 
sulphate  ? 

13.  A  specimen  of  urine  gives  a  white  precipitate  with  strong 
nitric  acid.     What  is  the  cause  of  it  ? 


CHAPTER   XV. 


SILICON  AND  BORON. 

231.  Silicon  (Siiv  =  28). 

OCCURRENCE. — The  greater  part  of  the  earth's  crust  is 
made  up  of  compounds  called  silicates,  composed  of  sili- 
con, oxygen,  and  metals.  Silicon  is  never  found  free  in 
nature,  but  always  in  the  form  of  silica  (SiO2),  either 
uncombined  (sand,  quartz,  &c.),  or  combined  in  silicates. 
Minute  traces  of  combined  silicon  are  present  in  the 
urine,  blood,  bones,  &c.,  of  man. 

PREPARATION.  —  By  heating  potassic  fluosilicate 
(K2SiF6)  in  iron  tubes  with  potassium  or  aluminium. 
When  potassium  is  used,  the  silicon  is  obtained  as  an 
amorphous  powder ;  with  aluminium,  it  crystallises  : 

2KF.SiF4  +  4K  =  6KF  +  Si. 

PROPERTIES. — Two  allotropic  forms,  one  crystalline, 
the  other  amorphous.  There  is  perhaps  a  third  form 


SILICON    DIOXIDE.  257 

resembling  graphite.  (Compare  carbon.)  Amorphous 
silicon  burns  readily  in  air,  forming  the  dioxide  (SiO.,) ; 
crystalline  silicon  oxidises  only  very  slowly.  Silicon  re- 
sembles carbon  in  its  chemical  characters.  Compounds 
with  hydrogen,  chlorine,  bi-omine,  and  iodine  are  known, 
parallel  with  methane,  &c.  (SiH4,  SiCl4,  SiBr4,  SiI4) 
also  more  complicated  compounds  (Si2Cl6,  <kc.) ;  and 
silicon  seems  able  to  replace  carbon  to  a  certain  extent  in 
many  organic  compounds. 

232.  Silicon  Dioxide. — SiO2.     Also  called   silica, 
and    anhydrous  silicic   acid.     It    is    the   only  oxide  of 
silicon  known. 

OCCURRKNCE. — Quartz  (rock  crystal)  is  pure  crystal- 
lised silica.  Amethyst  is  coloured  quartz.  Sand,  agate, 
flint,  jasper,  chalcedony,  diatoms,  opal,  &c.,  are  less  pure 
forms.  United  with  the  oxides  of  metals,  silica  forms 
the  great  divison  of  minerals,  the  silicates. 

PROPERTIES. — Pure  amorphous  silica,  obtained  by 
heating  silicic  acid  (H2SiO3),  is  a  light,  white  powder, 
insoluble  in  water  arid  in  all  acids  except  hydrofluoric 
(HF).  It  is  easily  soluble  in  alkalis,  even  in  ammonia. 
It  decomposes  sodic  carbonate  when  heated  with  it 
either  dissolved  in  water  or  solid.  Various  forms  of 
silica  are  used  in  glass  and  soap  manufacture,  in  mixing 
mortar  and  cements,  for  gems,  &c.  Agate  is  made  into 
mortars  and  pestles  used  in  tritu  rating  hard  minerals. 

233.  Silicic  Acid  and  Silicates. 

Experiment  206. — Fuse  a  little  clean  white  sand  in  a  por- 
celain  (or  better,  platinum)  crucible   with  about   1£  times  its 
18 


258  SILICIC    ACID. 

weight  of  anhydrous  sodic  carbonate  (Na2OO3).  Continue 
heating  until  a  clear  liquid  is  obtained.  Pour  this  out  on  a 
piece  of  clean  iron,  and,  when  it  has  cooled,  break  it  up  and  dis" 
solve  it  in  water  by  boiling  for  some  time. 

In  this  experiment  sodic  silicate  (Na2SiO3)  is  formed  : 

Na2CO3  +  SiO2  =  Na2SiO3  +  CO2. 
(Have  you  observed  the  escape  of  the  carbon  dioxide  1) 

Experiment  207- — To  a  little  of  the  solution  of  sodic  silicate 
add  dilute  hydrochloric  acid  ;  collect  the  gelatinous  precipitate 
on  a  filter  and  wash  it  with  hot  water.  It  is  silicic  acid.  Heat  a 
little  of  it  on  mica  and  obtain  silica.  Try  the  solubility  of  silicic 
acid  in  alkalis,  and  in  acids.  Write  the  equation  for  the  de- 
composition of  sodic  silicate  by  hydrochloric  acid. 

SILICIC  ACID  has  not  been  obtained  of  any  definite  com- 
position free  from  water.  There  are  numerous  classes  of 
silicates,  corresponding  to  a  great  number  of  theoretical 
acids  ;  e.g.,  ortho-silicates,  salts  of  an  acid,  H4SiO4 ;  and 
meta-silicates,  from  H2SiO3. — Silicic  acid  is  insoluble  in 
water,  but  an  unstable  solution,  which  easily  gelatinises, 
can  be  obtained  by  careful  manipulation  (compare  dia- 
lysed  iron). — The  silicates  of  the  alkalis  are  soluble  in 
water  (soluble  glass) ;  all  other  silicates  are  insoluble. 
Mica,  felspar,  garnet,  serpentine,  and  clays,  are  examples 
of  natural  silicates.  Felspar  is  a  double  silicate  of 
aluminium  and  potassium.  By  the  action  of  air  and 
water  it  is  decomposed  into  potassic  carbonate  and  hy- 
drated  aluminium  silicate,  or  clay.  In  this  way  a  fertile 
soil  is  gradually  formed  by  the  "  weathering"  of  felspar. — 
The  silicates  are  most  conveniently  represented  as  com- 
posed of  oxides  of  metals  united  with  one  or  more  mole- 
cules of  silica,  e.g.,  Na2O.Si02,  3MgO.2SiO2,  «fec. 


PLUOSILICIC   ACID,  259 

Tests. — 1.  If  the  substance  is  solid,  dissolve  a  little  of  it  in  a 
bead  of  molten  borax  on  a  small  loop  of  platinum  wire.  The 
silica  remains  undissolved,  floating  in  the  clear  bead. 

2.  Hydrochloric  acid  gives  a  gelatinous  precipitate  with  solu- 
tions of  silicates.  If  the  solution  is  already  acid,  silica  need  not 
be  tested  for. 

234.  Fluosilicic  Acid.—  H2SiF6. 

PREPARATION. — Experiment  208. — Mix  some  white  sand 
with  about  twice  its  weight  of  powdered  fluor  spar  (CaF2),  put 
the  mixture  in  a  t.  t.  or  flask,  pour  some  strong  sulphuric  acid 
over  it,  and  immediately  fit  with  a  delivery  tube  dipping  into 
a  t.  t.  half -full  of  water.  Heat  gently.  An  invisible  gas  is 
evolved,  bubbles  through  the  water,  and  deposits  a  flaky  white 
substance.  Collect  some  of  this  and  heat  it  on  mica.  It  dries 
up  to  a  light  white  powder.  It  is  silica.  Test  the  reaction  of 
the  water. — The  delivery  tube  soon  becomes  choked  up,  and  the 
experiment  must  then  cease. 

In  this  experiment,  three  chemical  actions  take  place, 
as  represented  by  the  following  equations : 

(1)  H2SO4     +     CaF2  CaS04     +     2HF. 

Silicon 
tetrafluoride. 

(2)  4HF     +     Si02  SiF4     +     2H2O. 

Orthosilcic  Fluosilicic 

acid.  acid. 

(3;  3SiF4    +    4H2O  H4Si04    +    2H2SiF6. 

Silicon  tetrafluoride  is  the  gas  which  bubbles  through  the 
water.  (What  are  the  final  products?  Where  are  they 
at  the  end  of  the  experiment  'i). — By  filtering  off  the 
silicic  acid,  a  solution  of  fluosilicic  acid  is  obtained. 

PROPERTIES. — A  strong  dibasic  acid.  Most  of  the 
fluosilicates  are  soluble  in  water,  but  potassic  fluosilicate 
(K,SiF6),  and  baric  fluosilicate  (BaSiF0)  are  only  sparingly 


260  BORON — BORIC    ACID. 

soluble.  For  this  reason  fluosilicic  acid  is  used  to  sepa- 
rate these  metals  from  others  present  with  them  in  solu- 
tions.— Fluosilicic  acid  is  silicic  acid  with  fluorine  instead 
of  oxygen,  6  atoms  of  fluorine  replacing  3  of  oxygen. 


235.  Boron  (Bm  =  11).— The   chief  compounds  of 
this  element  occurring  in  nature  are  borax  (Na2B4O7. 
10H2O),  and  boric,  or  boracic,  acid  (H3BO3).     Borax  is 
found   principally  in  the  plains  of  Thibet,  and  in  the 
borax  lakes  of  California.     Boric  acid  issues  along  with 
steam  from  fissures  in  the  sides  of  volcanic  hills  in  Tus- 
cany.— The  element,  boron,  is  of  little  importance.     It 
can  be  prepared  by  methods  similar  to  those  used  in  the 
preparation  of  silicon,  which  element  it  resembles  in  its 
properties. 

236.  Boric  Acid. — H3BO3.     Derived  from  the  la- 
goons of  Tuscany,  into  which  the  boric  acid  comes  from 
volcanic  fissures  or  from  holes  bored  for  the  purpose. 

PREPARATION. — The  water  of  these  lagoons  is  evapo- 
rated by  the  heat  of  the  steam  which  issues  from  the 
earth,  until  the  acid  crystallises  out.  The  crystals  are 
collected,  dried,  and  purified  by  recrystallisation.  Before 
the  discovery  of  the  presence  of  boric  acid  in  the  Tuscan 
lagoons,  it  was  prepared  from  borax,  by  treating  a  hot 
saturated  solution  with  strong  hydrochloric  acid.  The 
boric  acid  crystallised  out  on  cooling  : 

Na2B4O7  +  2HC1  +  5H2O  -  2NaCl  +  4H3BO3. 

Since  the  discovery  of  the  borax  lakes  of  California,  boric 
acid  is  again  made  in  this  way 


BOKAX.  HOi 

Experiment  209-  —  Dissolve  some  borax  in  strong  hydro- 
chloric acid  and  cool.  Boric  acid  crystallises  out.  Note  the  ap- 
pearance of  the  crystals. 

PROPERTIES. — A  white  solid,  in  small  scales  of  a  pearly 
lustre,  and  of  somewhat  soapy  feel. 

Experiment  210- — Heat  a  small  quantity  of  boric  acid  in  a 
dry  t.  t.  It  decomposes  into  water  and  boron  trioxide  (B2O3) : 

2H3BOS  =  3H.,O  +  B203. 

Experiment  211. — Try  the  solubility  of  boric  acid  in  water, 
cold  and  hot ;  also  in  alcohol.  Pour  the  alcoholic  solution  into 
a  porcelain  dish  and  set  fire  to  it.  Note  the  green  colour  of  the 
flame.  Boric  acid  volatilises  with  alcohol,  and  is  decomposed  in 
the  hot  flame,  imparting  the  green  colour  to  it. --Taste  the 
aqueous  solution,  and  try  its  action  on  blue  litmus. 

Boric  acid  is  tribasic,  but  the  borales,  like  the  silicates, 
are  very  various  in  their  composition,  and  are  best  re- 
presented as  compounds  of  metallic  oxides  with  boron 
tiioxiiie,  e.g.,  3MgO.4B.,O3. 

237.  Borax,  or  Sodic  Bi-borate,  Na./).2B2O3.10H2O. 
— The  crude  salt,  formerly  largely  imported  into  Venice 
and  Amsterdam  from  Thibet,  is  called  lineal.  From 
this  the  borax  of  commerce  was  obtained  by  boiling  with 
lime  to  rid  it  of  grease,  and  then  recrystallisiug  the  salt. 
Borax  is  now  prepared  mostly  by  fusing  boric  acid  with 
sodic  carbonate  : 

4H3BO3  +  Na2CO3  =  Na2B4O7  +  COa  +  6H2O. 

PROPERTIES. — Borax  is  generally  sold  in  large  colour- 
less crystals.  These  effloresce  in  air.  Borax  is  soluble 
in  water,  and  the  solution  is  alkaline.  It  is  insoluble  in 
alcohol. 


262  TESTS    FOR    BORIC    ACID. 

Experiment  212. — Heat  a  small  piece  of  borax  in  a  dry  t.  t. 
Water  condenses  in  the  tube  (Whence  has  it  come?),  and  the 
borax  swells  up.  Heat  a  little  borax  on  a  platinum  wire,  until 
it  fuses  to  a  clear  bead,  then  allow  it  to  cool,  moisten  it,  touch 
it  to  some  powdered  cupric  sulphate,  fuse  it  again,  and  note  the 
colour  when  it  has  cooled. 

When  borax  is  heated,  it  loses  its  water  of  crystallisa- 
tion, and,  if  the  temperature  is  high  enough,  fuses  to  a 
clear  liquid.  At  high  temperatures  the  borax  bead  dis- 
solves metallic  compounds,  forming  borates  which  give 
characteristic  colours  to  the  beads.  Borax  is  used  in 
this  way  to  test  small  quantities  of  solid  substances. 
— Borax  is  soluble  in  glycerine,  and  a  glycerine  of  borax 
is  used  as  a  lotion  and  gargle. 

Tests  for  Boric  Acid  and  Borates. — 1.  Mix  the  sub- 
stance with  strong  sulphuric  acid  and  alcohol  in  a  porcelain  dish, 
and  sst  the  alcohol  on  fire.  It  burns  with  a  green  flame.  Or, 
fuse  some  of  the  substance  on  a  platinum  wire,  moisten  it  with 
strong  sulphuric  acid  or  glycerine,  and  hold  it  in  the  Bunsen 
flame.  It  gives  a  green  colour  to  the  flame. 

2.  Acidify  the  solution  with  hydrochloric  acid,  dip  in  it  a 
strip  of  turmeric  paper,  and  dry  the  paper  at  a  gentle  heat.     It 
is  turned  brown. 

3.  Add  a  few  drops  of  baric  chloride  (BaCl2)  to  a  solution  of 
a  borate.     A  white  precipitate   is  formed.     It    is   soluble   in 
hydrochloric  acid. 


QUESTIONS  AND  EXERCISES. 

1.  Compare  silicon  and  carbon,  (a)  with  regard  to  chemical, 
and  (6)  with  regard  to  physical,  properties. 

2.  What  is  soluble  glass  ? 

3.  Explain  the  term  fluosilicic  acid. 

4.  Balance  the  following  equation  : 

K2SiF6  -f  Al  =  KF  +  A12F6  +  Si. 


METALS.  263 

5.  It  is  found  that  alkaline  solutions  eat  away  glass.    Explain. 

6.  Mention  an  acid  and  a  base  which  will  dissolve  sand. 

7.  How  can  it  be  proved  that  silica  is  an  acid-forming  oxide  ? 
9.  With   what  does  the  delivery-tube  become  choked  up  in 

Experiment  208  ? 

9.  Boric  acid  (H3B03)   is   tribasic.     Write  the   formula   for 
normal  sodic  borate.     Is  this  the  ordinary  salt  ? 

10.  Borax  is  an  acid  salt,  and  yet  its  solution  is   alkaline. 
Account  for  this. 

11.  How  can  borax  be  used  to  test  for  glycerine  ? 

12.  In  the  preparation  of  boric  acid  from  borax,  sodic  chloride 
is  formed.     How  is  the  boric  acid  separated  from  it  ? 


CHAPTER    XVI. 


THE   METALS. 

238.  General   Characters.  -The    only   character- 
istic common  to  all  metals  is  the  power  of  taking  the 
place  of  the  hydrogen  of  acids  to  form  salts,  or  in  other 
words,  the  power  of  forming  bases.     But  most  metals 
have  a  bright  and  peculiar  appearance   called  metallic 
lustre ;   as    a   rule,  they    are    specifically    heavier   than 
water ;  they  are  good  conductors  of  heat  and  electricity ; 
and  most  of  them  are  malleable  and  ductile. 

239.  Ores  Of  Metals. — The  noble  metals  (i.e.,  the 
least     oxidisable)     are     generally     found     uncombinrd 
(native) ;    native    copper  also    occurs.     Sometimes    the 
baser  metals  are  found  free  ;  but  in  most  cases  they  occur 
combined   with    other  elements,    from   which    they   are 
separated   by    the    processes    of    metallurgy.       Common 
metals  such  as  lead,  iron,  tin,  &c.,  are  found  very  gener- 


264  METALLURGY. 

ally  as  sulphides  and  oxides.  In  many  cases  salts  of 
the  metals  are  found  in  nature.  If  the  metals  of  the 
alkalis  be  excepted,  it  may  be  stated  as  a  general  rule 
that  metallic  ores  are  insoluble  in  water.  — The  processes 
by  which  metals  are  obtained  from  their  ores  are  various. 
A  common  way,  applicable  to  oxides  and  some  oxygen 
salts,  is  to  heat  to  a  high  temperature  with  some  form  of 
carbon  (coal,  charcoal,  &c.).  In  many  cases  sulphides 
are  first  roasted  to  convert  them  into  oxides,  and  then 
reduced  by  means  of  coal  or  charcoal.  In  other  cases 
sulphides  are  partially  oxidised  at  a  comparatively  low 
temperature,  and  then  heated  to  a  higher  temperature,  at 
which  the  oxygen  and  remaining  sulphur  combine,  set- 
ting the  metal  free.  Lead  and  copper  are  obtained  in 

this  way : 

(1)  PbS  +  30  =  PbO  +  S02. 

(2)  PbS  +  2PbO  =  3Pb  +  SO2. 

1.  Metals  reduced  by  heating  the  ores  with  coal,  char- 
coal, &c.  :    Na,  K,  Rb,  Sn,  Cd,  Zn,  Fe,  Mn,  Sb,  O,  Ni. 

2.  By  partial  ooddatian  and  subsequent  fusion :  Pb, 
Cu,  Bi. 

3.  By  heating  with  sodium  or  potassium  :  Al,  Mg,  Ca, 
Be. 

4.  By  electrolysis  of  fused  salts  :  Ba,  Ca,  Sr,  Li,  Cs. 

5.  By  distilling  in  a  current  of  air  :  Hg. 

6.  Native :  Cu,  Ag,  Au,  Pt,  Hg,  &c. 

This  is  only  a  general  statement,  and  not  intended  to 
be  exhaustive. 

240.  Alloys.  — -  Metals   usually  combine  with   each 
other  when  fused  together.     In  many  cases  the  combina- 


ALLOYS.  265 

tion  takes  place  in  definite  proportions,  and  chemical  com- 
pounds are  formed.  In  other  cases,  the  proportions  may 
be  varied,  and  it  is  not  easy  to  decide  whether  or  not  there 
is  any  chemical  action.  Compounds  of  metals  with  each 
other  are  called  alloys.  The  properties  of  alloys  are  not 
the  mean  between  those  of  the  metals  present.  Alloys 
generally  melt  at  lower  temperatures  than  any  of  the 
constituents.  Thus,  Rose's  fusible  metal  (tin,  1  part; 
lead,  1  part;  and  bismuth,  2  parts;  melts  at  95°  C.  ; 
while  tin  melts  at  235°,  lead  at  334°,  and  bismuth  at 
270°.  Many  alloys  are  in  common  use,  e.g. 

Solder,  tin  and  lead  in  various  proportions. 

Brass,  65  parts  zinc,  and  137  parts  copper. 

Gold  coin,  1 1  parts  gold  and  1  part  copper. 

Silver  coin,  40  parts  silver  and  3  parts  copper. 

Gun  metal,  9  parts  copper  and  1  part  tin. 

Bell  metal,  4  parts  copper  and  1  part  tin. 

Type  metal,  4  parts  lead  and  I  part  antimony. 

Britannia  metal,  copper,  zinc,  tin,  antimony. 

Pewter,  4  parts  tin,  and  1  part  lead. 

Bronze,  tin,  copper,  and  zinc. 

German  silver,  5  parts  copper,  2  parts  nickel,  and  2 
parts  zinc. 

Wood's  fusible  metal  (melting  at  68°),  8  parts  lead,  5 
parts  bismuth,- 4  parts  tin,  and  3  parts  cadmium. 

Amalgams  are  alloys  of  mercui-y  with  other  metals. 

241.  Compounds  Of  Metals-— The  compounds  of 
metals  are  classified  into  (1)  binary  compounds,  including 
oxides,  sulphides,  chlorides,  bromides,  iodides,  fluorides, 


266  OXIDES    OF    METALS. 

&c.  (Note  the  ending — ide) ;  and  (2)  oxygen  salts,  &c., 
such  as  sulphates,  sulphites,  nitrates,  phosphates,  &c. 

242.  Oxides  Of  Metals. — Each  metal  has  at  least 
one  basic  oxide  ;  several  have  acid-forming  oxides ;  and 
many  have  oxides  which  are  neither  base-forming  nor 
acid-forming,  hence  called  indifferent  oxides. 

1.  BASIC  OXIDES. — These  may  be  classified  as  follows: 

(a)  General  formula  M^O  : — Li2O,  Na20,  K2O,  Rb2O, 
CS2O ;  Ag2O,  Hg20,  Cu20 ;  T12O,  Au2O. 

(b)  General  formula  MO :— CaO,   SrO,   BaO  ;  PbO, 
HgO,  CuO,  SnO;  FeO,  MnO,  ZnO,  NiO,   CoO,  CrO  ; 
PtO,  &c. 

(c)  General  formula   M^OZ : — Sb2O3,    Bi2O3,    Au203, 
T1203 ;  Fe203,  A12O3,  Mn2O3,  Co203,  Cr2O3. 

This  includes  nearly  all  the  basic  oxides  (the  rarer 
metals  being  left  out  of  consideration). — It  will  be 
observed  that  some  metals  have  more  than  one  basic 
oxide.  Some  dyad  metals  (Hg,  Cu)  have  oxides  in 
which  they  seem  to  play  the  part  of  monad  metals. 
There  is  reason  to  believe,  however,  that  the  molecules  of 
the  corresponding  chlorides  must  be  represented  as  fol- 
lows: Hg2Cl2,  and  Cu2Cl2.  The  metals  must,  then,  be 
dyad  in  these  compounds,  and  the  structural  formulas 

Hgv.        Cu  v 

for  the  oxides   and  chlorides    are      I       O,    I    ^O,    and 

Hg^      Cu^ 

Hg— Cl    Cu— Cl 

I  ,    |          . — Gold  and  thallium  are  monad  and  triad. 

Hg— 01    Cu— Cl 

Thus  gold  forms  aurous  oxide,  Au20,  and  auric  oxide, 
Au2O3. — Iron,  cobalt,  and  manganese,  have  two  basic 
oxides,  e.g.  ferrous  oxide,  FeO,  and  ferric  oxide,  Fe2O3. 
In  the  lower  oxides  these  metals  are  dyad.  The  specific 


OXIDES    OF    METALS.  267 

weight  of  ferric  chloride  in  the  gaseous  condition  shows 

Fe 

that  its  formula  is  Fe2Cl6,  or  graphically 


Fe  =O 

oxide  must  then  be  represented  as         ^>  O,  and  in  the 

Fe  =0 

ferric  compounds  iron  is  tetrad.  So  with  manganese, 
cobalt,  aluminium,  and  chromium,  which  have  also  basic 
sesquioxides.  The  first  group  of  sesquioxides,  Sb,O3,  <fec., 
correspond  to  chlorides,  SbCl3,  &c.,  so  that  the  metals 
must  be  represented  as  triad,  e.g.  O  —  Sb  -  O  -  Sb  =  0. — 
There  are  hydroxides  corresponding  to  most  of  the  basic 
oxides. 

2.  ACID-FORMING  OXIDES. — The  basic  oxides  of  imper- 
fect metals,  such  as  antimony,  have  also  weak  acid  charac- 
ters ;  but  there  are  also  distinct  acid-forming  oxides  of 
these  metals,  e.g.  antimony  pentoxide,  Sb2O5.     Chromium, 
manganese,  <fec.,  also  form  acids.     The  acid-forming  oxides 
always  contain  a  greater  proportion  of  oxygen  than  the 
basic,  and,  as  a  rule,  readily  give   up   their  oxygen   to 
reducing  agents,  becoming  transformed  to  basic  oxides. 

3.  INDIFFERENT  OXIDES. — These  are  such  oxides  as 
manganese  dioxide  (MnO2),  having  neither  acid  nor  basic 
properties.     They  generally  contain  a  greater  proportion 
of  oxygen  than  the  basic  oxides,  and  are  hence  often 
called  joer-oxides. 

243.  Sulphides  of  Metals.— These  correspond 
closely  to  the  basic  oxides  of  the  metals,  e.  g.,  Li2S,  Na2S, 
<fec.  The  imperfect  metals  have  sulphides  corresponding 
to  their  acid-forming  oxides,  e.  g.,  Sb2S5,  SnS2,  &c.  As  a 


268  CHLORIDES,    &C. — OXYGKN    SAL'IV 

rule,  when  a  basic  oxide  of  a  metal  is  soluble  in  water, 
the  corresponding  sulphide  is  also  soluble, — and  so  for 
insolubility.  Thus,  both  the  oxides  and  the  sulphides 
of  lithium,  sodium,  potassium,  &c.,  are  soluble  in  water. 

244.  Chlorides,    &c.  —  There  are    chlorides  corres- 
ponding to  the  basic  oxides  of  the  metals.     In  writing 
the  formulas  of  chlorides,  it  must  be  remembered  that  2 
atoms  of  chlorine    replace    1    atom    of   oxygen.     Thus, 
given  the  formula  of  aluminic  oxide,  A12O3,  the  formula 
for   aluminic  chloride  is  written,   A12C16.     It  could    be 
written  more  simply,  A1C13,  but,  for  reasons  similar  to 
those  stated  above  for  ferric  chloride,  the  simpler  formula 
is  doubled.     The  formula  for  bismuth  trioxide  is  Bi2O3 ; 
replacing  O3  by  C16,  we  get  as  the  formula  for  the  chlor- 
ide, Bi.2Cl6,  but  its  specific  weight  in  the  gaseous  state 
shows  that  its  molecule  contains  only  half  the  number  of 
atoms  represented  here,  so  that  the  formula  for  bismuth 
trichloride  is  BiCl3. 

245.  Oxygen    Salts.  —  The    oxygen    salts    of   the 
metals  have  been  already  noticed  along  with  the  various 
acids.     To  derive  the  formula  of  an  oxygen  salt  from  the 
formula  of  an  oxygen  acid,  it  is  necessary  to  know  the 
atomicity  of  the  metal,  and,  also,  whether  or  not  two 
atoms  of  the  metal  play  the  part  of  a  single  atom  as  in 
the  case  of  mercurous  salts  (Hg2(NO3)2,  &c.),  and    the 

ferric,  salts  (Fe2(NO3)6,  &c.).  For  example,  given  the 
formula  of  sulphuric  acid  as  H2S04,  and  knowing  the 
atomicity  of  calcium  to  be  2,  it  is  easy  to  write  the  for- 
mula for  calcic  sulphate,  viz,,  CaSO4.  Again,  bismuth 
is  triad ;  and,  therefore,  1  atom  of  bismuth  replaces  3  of 
hydrogen.  In  order  to  replace  the  hydrogen  of  sul- 


ANALYSIS PRECIPITATION.  269 

phuric  acid  by  bismuth,  we  must  take  3H2SO4.  The 
6  atoms  of  hydrogen  are.  equivalent  to  2  of  bismuth  ; 
and,  therefore,  the  formula  for  bismuth  sulphate  is 
Bi2(S04)3.  The  atomicity  of  iron  in  the  ferric  salts  is  4, 
but  2  atoms  of  iron  are  united  in  the  molecules  of  ferric 
salts  so  that  their  joint  atomicity  is  6;  i.e.,  Fe2  replaces 
6H.  The  formula  for  ferric  sulphate  is  thus  Fe2(SO4)3. 
—To  many  oxygen  salts  there  are  corresponding  sulphur 
salts,  e.g.,  K2CS3. 

246.  Classification  of  Metals  —  Analysis.  - 

The  method  of  classification  to  be  adopted  here  is  that 
which  is  employed  in  the  process  of  examining  unknown 
substances  to  discover  the  elements  of  which  they  are 
composed.  These  processes  constitute  analysis  in  the 
broad  sense  of  the  term.  The  substances  are  not  always 
— indeed,  not  generally — decomposed  into  their  ele- 
ments, but  such  evidence  is  obtained  as  enables  the  ana- 
lyst to  be  certain  of  the  presence  of  the  elements.  This 
is  qualitative  analysis.  If  the  quantities  of  the  elements 
in  a  compound,  or  the  quantities  of  the  elements  and 
compounds  in  a  mixture,  are  to  be  determined,  this  is 
done  by  quantitative  analysis.  A  great  deal  of  the  work 
in  both  qualitative  and  quantitative  analysis  consists  in 
preparing  insoluble  compounds  of  the  metals  (and  acids) 
by  precipitation  from  solutions.  In  order  to  understand 
the  operations  of  analysis  it  is  necessary  to  know  the 
solubilities  of  chemical  substances.  Such  knowledge  is 
also  of  great  importance  to  the  prescriber  of  medicines. 
A  physician  who  is  not  well  acquainted  with  this  part  of 
chemistry  is  very  likely  to  produce  "  muddy  mixtures." 
PRECIPITATION. — When  two  chemical  compounds  are 
brought  together  in  solution,  there  is  usually  chemical 


270  CLASSIFICATION    OF    METALS. 

action,  consisting  in  exchange  of  parts  of  the  molecules. 
If  this  exchange  causes  the  formation  of  an  insoluble 
substance,  the  latter  is  precipitated,  or  thrown  to  the  bottom. 

Experiment  213. — Mix  solutions  of  cupr/c  mlpliate  (CuSOJ 
and  baric  chloride  (BaCl2).  A  white  precipitate  of  baric  sul- 
phate (BaSOJ  falls,  while  cuprir  chloride  (CuCl2)  remains  in 
solution. 

Such  a  chemical  action  is  called  a  double  decomposition, 
because  the  two  salts  decompose  each  other,  the  metals 
changing  places  : 

CuSO4 )          j  BaSO4 
BaCl2   j    :  :    (CuCl2 

GROUP  REAGENTS. — Insoluble  compounds  can  gener- 
ally be  obtained  by  precipitation.  We  have  seen  in 
studying  groups  of  compounds,  such  as  chlorides,  sul- 
phides, carbonates,  &c.,  that  each  group  may  be  classified 
into  (1)  soluble,  and  (2)  insoluble.  Thus,  there  are  three 
insoluble  chlorides  (PbCl2,  AgCl,  Hg2Cl2),  and  the  rest 
are  soluble.  If  a  complex  solution  containing  salts  of  all 
the  metals  were  treated  with  hydrochloric  acid,  these 
three  chlorides  would  be  formed  and  precipitated.  Thus, 
a  separation  would  be  effected  of  lead,  silver,  and  mer- 
cury from  the  rest  of  the  metals.  Hydrochloric  acid 
is  a  group  reagent ;  and,  in  analysing  substances,  the 
first  step  is  to  determine  by  the  use  of  group  reagents  to 
what  group  or  groups  the  substances  under  examination 
belong. 

GROUPS  OF  METALS. 

I.  Lead*  silver,  and  mercury  (mercunms  salts). 
Chlorides  precipitated  by  hydrochloric  acid. 

*  Plumbic  chloride  is  sparingly  soluble,  so  that  lead  appears  in  Groups 
and  II. 


QUESTIONS    AND    EXERCISES.  271 

II.  Lead,  mercury  (mercuric  salts),  copper,  cadmium, 
bismuth,   antimony,    [arsenic \,   tin,  gold,  platinum,   and 
some   rare   metals.     Sulphides   precipitated    by    hydric 
sulphide  from  neutral  or  acid  solutions. 

III.  Iron,   chromium,    aluminium,    zinc,   manganese, 
cobalt,  and  nickel.     Sulphides  precipitated  only  in  pres- 
ence of  an  alkali.     Ammonic  sulphide  is  the  group  re- 
agent. 

IV.  Calcium,  strontium,  and  barium.     Precipitated  as 
carbonates  from  solutions  (in  presence  of  ammonic  chlo- 
rides) by  ammonic  carbonate. 

V.  Magnesium.     Precipitated  as  phosphate,  by  sodic 
phosphate.     Magnesium  is  often  included  in  Group  IV. 

VI.  Lithium,    sodium,  potassium,    ammonium,    rubi- 
dium, and  caesium.     Salts  mostly  soluble. 

Groups  I.,  II.,  and  III.,  include  the  common  metals 
in  everyday  use.  They  are  the  heavy  metals.  Group  IV. 
is  made  up  of  the  metals  of  the  alkaline  earths.  Group 
VI.  includes  the  metals  of  the  alkalis. 


QUESTIONS  AND  EXERCISES. 

1.  Metals  generally  feel  cooler  than  other  substances.     Why 
is  this  ? 

2.  What  is  an  ore  ? 

3.  Are  alloys  compounds  or  mixtures? 

4.  What  class  of  metallic   compounds  does  the  ending  -ide 
mark  ? 

5.  Why  write  the  formula  for  chromic  chloride  Cr^Clg,  and 
not  CrCl3  ? 

6.  Mercury  is  a  dyad  metal,  but  it  has  an  oxide  HgsO.    How 
is  this  explained  ?     What  non-metal  forms  numerous  compounds 
in  the  same  way  ? 


272  METALS   OF   GROUP    I. 

7.  Bismuth  has  an  oxide,   the  formula   of    which  is  Bi.,O3. 
Ferric  oxide  is  represented  by  Fe203.     The  first  is  called  l/tx- 
nit/tk  trio.rii/c  ;  the  second,  iron  sesquioXide.     Is  there  any  rea- 
son for  this  difference  of  momenclatare  ? 

8.  Write  the  formulas  for  the  chlorides  corresponding  to  SnO, 
O20S,  K2O,  Bi203,  BaO,  CuO,  Cu2O,  Ag2O,  and  Au.203. 

9.  The  formula   for  oxalic  acid   is    H2C.,04.     It   is  dibasic. 
Write  the  formulas  for  the  normal  oxalates  of  barium,  sodium, 
ammonium,  iron  (ferrous  and  ferric),  chromium,  and  copper. 

10.  Will  the  group  reagent  of  Group  II.  precipitate  the  mem- 
bers of  Group  I.  as  sulphides  ?     Try. 

11.  Have  you  observed  any  regularity  in  the  atomicities  of 
the  groups  of  metals?     Write  the  formulas  for  the  oxides  of 
Groups  IV.  and  VI. 


CHAPTER     XVII. 


METALS  OF  GROUP  I. 
Lead,  Silver,  and  Mercury. 

247.  General  Characters.— The  metals  of   this 

group  are  heavy  and  soft  (mercury  is  liquid).  They  are 
easily  reduced  from  their  ores  by  heating  with  charcoal. 
Their  sulphides  are  black  and  insoluble  in  water  and 
in  dilute  acids.  The  oxides  are  earthy  compounds  in- 
soluble in  water.  The  chlorides  (except  mercuric  chlo- 
ride), bromides,  iodides,  carbonates,  and  phosphates,  are 
insoluble  in  water. 


LEAD  (Plumbum). 

248.  Lead.— (Pbn=  206.4.  Specific  weight  =11. 352. 
Melting  point  =  334°  C.    Specific  heat  =  0.0315).    The 


LEAD.  273 

chief  ore  of  lead  is  galena  (PbS).  It  is  very  common  in 
Canada,  occurring  in  crystalline  masses,  of  a  brilliant, 
metallic  appearance. 

Experiment  214.— Examine  a  specimen  of  galena,  noting  its 
colour,  hardness  (scratch  with  a  knife),  specific  weight,  &c. 
Mix  a  little  of  the  powdered  mineral  with  sodic  carbonate, 
moisten,  and  heat  it  on  charcoal  before  the  blow-pipe,  or  on  a 
charred  splinter  in  the  reducing  flame  of  the  Bunsen  burner. 
Extract  the  metallic  bead,  and  examine  it  carefully  as  to  its 
hardness,  malleability,  &c.  It  is  lead. 

Galena  generally  contains  silver,  sometimes  only  in 
small  traces,  but  often  in  considerable  proportion. 

PREPAKATION. — The  galena  is  partially  oxidised  by 
heating  in  air,  and  then  more  strongly  heated  to  set  the 
lead  free.  (Art.  239). 

Experiment  215- — Put  a  piece  of  zinc  in  a  solution  of  plum- 
bic acetate,  and  allow  it  to  remain  for  some  time.  Pour  off  the 
liquid,  dry  the  metal  which  remains,  and  melt  it  by  heating  it 
in  a  closed  porcelain  crucible  with  a  little  charcoal  dust.  Ex- 
-  amine  it.  It  is  lead. 

PROPERTIES. — A  heavy,  dull  metal,  soft,  tough,  easily 
tarnished.  A  small  quantity  of  antimony  or  arsenic 
alloyed  with  it  renders  it  hard  and  brittle.  Lead  is 
easily  set  free  from  compounds  in  solution  by  the  action 
of  iron,  zinc,  &c.,  as  in  Experiment  215.  This  method 
is  sometimes  employed  for  impure  ores. 

Experiment  216. — Warm  some  small  scraps  of  lead  with 
dilute  nitric,  dilute  hydrochloric,  and  dilute  sulphuric  acids. 
Divide  the  solution  obtained  with  nitric  acid  into  two  portions. 
To  one  add  some  hydrochloric,  to  the  other  a  little  sulphuric 
acid.  Heat  bits  of  lead  with  strong  hydrochloric  and  sulphuric 
acids. 

19 


•274:  OXIDES    OP    LEAD. 

Experiment  217. — Put  pieces  of  bright  lead  into  4  test 
tubes,  labelled  (1),  (2),  (3),  and  (4).  In  (1)  put  distilled  water, 
in  (2)  water  containing  ammonic  nitrate,  in  (3)  vinegar,  and  in 
(4)  tap  water.  Set  aside  for  a  day  and  then  test  the  liquid  con- 
tents of  the  tubes  for  lead. 

Lead  is  attacked  and  dissolved  by  distilled  water  (and 
rain  water)  owing  to  the  action  of  the  dissolved  oxygen 
and  carbonic  acid.  Water  containing  ammonium  salts 
(especially  ammonic  nitrate)  dissolves  lead.  Water  con- 
taminated by  sewage  generally  contains  ammonium  salts, 
and  is  therefore  dangerous  on  this  account  as  well  as  on 
others.  (Does  vinegar  dissolve  lead?)  Water  contain- 
ing lime  and  magnesia  salts  does  not  attack  lead,  so  that 
ordinary  river  and  well  waters  may  be  carried  safely 
through  lead  pipes. 

249.  Oxides  of  Lead. — Lead  forms  several  oxides 
(Pb2O,  PbO,  Pb.2O3,  Pb3O4,  PbO,).  Of  these  the  impor- 
tant ones  are  lead  monoxide  (PbO),  and  red-lead  (Pb3O4). 

1.  LEAD  MONOXIDE  (PbO). — Also  called  litharge,  or 
massicot,  according  to  the  method  by  which  it  is  pre- 
pared. 

Experiment  218- — Heat  some  thin  shavings  of  lead  in  an 
open  porcelain  crucible,  They  oxidise  to  a  greyish  yellow  sub- 
stance. This  is  the  monoxide.  Remove  it  and  try  the  solu- 
bility of  portions  of  it  in  dilute  acetic,  nitric,  hydrochloric,  and 
sulphuric  acids.  (Do  you  notice  any  change  with  hydrochloric 
and  sulphuric  acids  ?) 

Litharge  is  used  for  giving  a  glaze  to  earthenware, 
and  in  making  flint  glass.  It  is  also  used  in  preparing 
red-lead,  lead  acetate,  nitrate,  &c.  "  Drying  oils  "  are 
prepared  by  boiling  the  raw  oils  with  litharge. 

Experiment  219- — Heat  some  lead  monoxide  with  solution 
of  sugar.  It  is  dissolved. 


SALTS    OP    LEAD.  275 

Experiment  220. — Try  the  solubility  of  lead  monoxide  in 
water,  and  in  solutions  of  sodic,  potassic,  and  calcic  hydroxides. 

2.  RED-LEAD,  or  MINIUM  (Pb3O4). — This  is  prepared 
by  heating  litharge  or  massicot  in  air  until  it  becomes 
further  oxidised  : 

3PbO  +  O  =  Pb3O4. 

It  is  a  heavy  red  powder,  used  as  a  paint,  and  in  the 
manufacture  of  flint  glass.  It  is  often  adulterated  with 
brick  dust,  ferric  oxide,  <kc. 

Experiment  221. — Heat  a  little  red-lead  in  a  porcelain  dish 
and  note  any  changes. 

Experiment  222. — Warm  a  some  red  lead  with  dilute  nitric 
acid  to  which  a  little  sugar  has  been  added.  It  is  completely 
dissolved,  if  pure,  a  solution  of  plumbic  nitrate  being  obtained. 
(What  is  the  object  of  the  sugar  ?  What  becomes  of  the  oxygen 
in  excess  of  3PbO  ?) 

250.  Salts  of  Lead. — There  are  two  basic  oxides  of 
lead,  PbO  and  Pb2O3,  but  the  sesquioxide  forms  unstable 
salts  of  no  importance      The  ordinary  salts  of  lead  are 
derived  from  the  monoxide,  PbO ;  and  in  these  salts  Pb 
takes  the  place  of  2H.     Moist  lead  oxide,  as  well  as  the 
hydroxide,  turn  red  litmus  blue.     Salts  of  lead  have  a 
sweetish  metallic  taste. 

251.  Plumbic    Acetate,  or  Sugar  of   Lead, 

Pb(C,H3Oo)2-3H2O.  Prepared  by  dissolving  litharge  or 
massicot  in  acetic  acid.  (See  Experiment  163.) 

Experiment  223.  —Examine  carefully  a  specimen  of  sugar  of 
lead.  (Why  called  sugar  ?)  Dissolve  it  in  warm  distilled  water. 
(It  is  soluble  in  1£  parts.)  Test  it  with  blue  litmus.  Note  the 
odour  of  the  solution.  If  the  solution  is  turbid,  add  acetic  acid 
until  it  is  clear. 


276  PLUMBIC    NITRATE WHITE    LEAD. 

GOULARD'S  EXTRACT  is  a  solution  of  sub-acetate,  or 
basic  acetate,  of  lead,  Pb(C2H3O2)2.PbO,  made  by  boiling 
solution  of  the  normal  acetate  with  litharge. 

Experiment  224. — Boil  some  solution  of  plumbic  acetate  for 
some  time  with  litharge  in  a  porcelain  dish.  Filter,  and  test  a 
little  of  the  filtrate  with  red  litmus.  Leave  the  rest  exposed  to 
the  air.  It  becomes  turbid  owing  to  the  formation  of  plumbic 
carbonate.  Try  its  action  on  gum  arabic  mucilage. 

252.  Plumbic  Nitrate,  Pb(NO3)2.— The  nitrate  of 
lead  has  already  been  prepared  in  several  experiments. 

PREPARATION — By  dissolving  litharge  or  massicot  in 
warm  dilute  nitric  acid,  filtering,  and  evaporating  to 
crystallisation  : 

PbO  +  2HN03  =  Pb(NO3)2  +  H20. 

PROPERTIES. — A  white  crystalline  solid,  of  astringent 
metallic  taste.  It  dissolves  in  twice  its  weight  of  water, 
but  is  only  sparingly  soluble  in  alcohol. — It  is  an  irritant 
poison.  It  has  been  used  as  a  disinfectant  and  deodor- 
iser. Its  action  as  a  deodoriser  is  due  to  the  fact  that 
it  reacts  with  sulphuretted  hydrogen  to  form  the  sulphide 
of  lead  (PbS).  Ledoyeris  Disinfecting  Fluid  contains 
a  drachm  of  lead  nitrate  to  an  ounce  of  water. 

Experiment  225.— Shake  up  a  solution  of  plumbic  nitrate 
with  an  equal  volume  of  hydric  sulphide  water.  Note  the  ab- 
sence of  bad  smell  in  the  solution.  What  is  the  black  pre- 
cipitate formed  ? 

253.  White  Lead.— 2PbCO3Pb(OH)2,  a  basic  car- 
bonate of  lead. 

PREPARATION. — 1.  Dutch  process.  Expose  sheets  of 
lead  to  the  conjoint  action  of  the  fumes  from  vinegar  and 


PLUMBIC    CHLOBIDE.  277 

the  carbon  dioxide  from  decaying  tan-bark.     An  acetate 
is  formed  first,  and  this  is  decomposed  by  carbon  dioxide. 
2.  Milner's    process.     Grind    litharge    with    common 
salt  and  water : 

PbO  +  H20  +  NaCl  =  Pb  j  g^  +  NaOH. 

Pass   in    carbon   dioxide    till    the    solution    is    neutral. 
The  basic  carbonate  is  formed. 

PROPERTIES. — A  soft,  heavy,  white  powder,  insoluble 
in  water,  but  easily  dissolved  by  dilute  nitric  or  acetic 
acid  with  effervescence  of  carbon  dioxide.  It  is  poisonous. 

Experiment  226- — Pour  some  hydric  sulphide  water  on  a 
little  white  lead  in  a  porcelain  dish.  Explain  the  blackening. 

Experiment  227.— Dissolve  a  little  white  lead  in  acetic  acid. 
(What  substances  are  formed?  Write  the  equation.)  To  por- 
tions of  the  solution  add  hydrochloric  acid,  solution  of  potassic 
iodide,  and  dilute  sulphuric  acid  respectively.  (Note  the  ap- 
pearance of  the  precipitates  and  write  the  equations. ) 

White  lead  is  used  as  a  paint.  It  is  used  in  medicine 
in  the  form  of  an  ointment  (unguentum  plumbi  carbon 
atis).  It  is  very  often  adulterated  with  baric  sulphate 
(BaSOJ,  and  gypsum ;  but,  as  both  are  insoluble  in 
nitric  or  acetic  acid,  they  are  easily  detected. 

254.  Plumbic  Chloride.  -  PbCl2.  Occurs  in  nature 
as  horn  lead. 

Experiment  228. — To  a  little  solution  of  lead  acetate  add 
excess  of  hydrochloric  acid.  Plumbic  chloride  is  precipitated. 
Filter,  and  transfer  the  precipitate  to  a  t.  t.  ;  fill  the  t.  t.  half- 
full  of  pure  water  and  heat  to  boiling.  The  chloride  dissolves 
completely,  if  there  is  enough  water.  Allow  to  cool,  and  observe 
crystals. — Add  to  the  solution  a  drop  or  two  of  hydrochloric 
acid  to  see  that  it  gives  no  further  precipitate.  Now  add  a  few 
drops  of  sulphuric  acid,  and  white  sulphate  of  lead  (PbS04) 


278  PLUMBIC    IODIDE LEAD    PLASTER. 

is  precipitated.     (What  conclusion  do  you  draw  as  to  the  solu- 
bility of  plumbic  chloride  ?) 

Plumbic  chloride  is  also  sparingly  soluble  in  pure  water. 

255.  Plumbic  Iodide. — PbI2.      (In    what   experi- 
ment has  this  been  already  formed  1) 

Experiment  229.— To  some  solution  of  plumbic  acetate  add 
solution  of  potassic  iodide  (KI),  collect  the  precipitate  of  plum- 
bic iodide  on  a  filter,  and  wash  it  with  cold  water.  Note  its 
colour,  &c.  Transfer  it  to  at.  t.,  fill  half -full  of  water,  boil,  and 
allow  to  cool.  The  iodide  dissolves  in  hot  water,  and  crystal- 
lises beautifully  on  cooling. 

Plumbic  iodide  is  used  in  medicine  in  the  forms  of 
plaster  and  ointment. 

256.  Lead  Plaster. — This  consists  of  oleate  of  lead 
and  lead  salts  of  other  fatty  acids. 

PREPARATION. — Boil  at  a  gentle  heat  for  4  or  5  hours, 
4  pounds  of  litharge  with  1  gallon  of  olive  oil  and  3^  pints 
of  water,  adding  water  as  it  evaporates.  The  glycerine 
dissolves  in  the  water,  and  the  oleate  of  lead  forms  an 
insoluble  gummy  mass.  Lard  is  sometimes  used  instead 
of  olive  oil. 

257.  Plumbic  Sulphate.— PbS04. 

Experiment  230- — Warm  a  little  litharge  with  concentrated 
sulphuric  acid.  It  dissolves.  Cool  the  solution  carefully  and 
then  dilute  with  water.  Plumbic  sulphate  is  precipitated. 

Plumbic  sulphate  is  soluble  in  strong,  but  insoluble  in 
dilute,  sulphuric  acid. 

Experiment  231- — Mix  solutions  of  magnesic  sulphate  and 
plumbic  acetate.  Plumbic  sulphate  is  precipitated  : 

PbA2  +  MgSO4  =  PbSO4  +  MgA2. 


LEAD    POISONING.  279 

Try  its  solubility  in  hot  water,  and  in  nitric  and  hydrochloric 
acids. 

Commercial  oil   of  vitriol  often   contains  sulphate  of 
lead.     It  is  precipitated  on  dilution. 

258.  Commercial  Preparations  of  Lead. 

1.  Basic  Carbonate. — White  lead,  flake  white,  ceruse, 
mineral  white,  Newcastle  white,  and  Nottingham  white. 
Also  some  hair  dyes. 

2.  Sulphate. — Miniature  painter's  white,  white  preci- 
pitate of  lead. 

3.  Chromates. — Chrome  yellow  (PbCrO4),  chrome  red 
(PbCrO4.PbO). 

4.  Basic  Chloride. — Turner's  yellow,  or  Cassella  yel- 
low (PbCl2.7PbO),  and  Pattinson's  white  (PbCl.OH). 

5.  Litharge  and  various  salts  of  lead  are  constituents 
of  hair  dyes. 

6.  Minium  (Pb3O4).— Red  lead. 

259.  Lead  Poisoning. — Nearly  all  compounds   of 
lead  are  poisonous ;  but  the  insoluble  compounds  are  at 
least  very  slow  in  their  action.     The  soluble  compounds 
are    irritant   poisons.     The    antidotes    are    soluble    sul- 
phates,  such   as    Epsom    salts.     Lead    is    a    cumulative 
poison.     Minute  quantities  taken  repeatedly  remain  in 
the  system,  and  at  length  accumulate  to  such  an  extent 
as  to  produce   symptoms  of  poisoning,  as   in   painter's 
colic.     The   treatment    in   such    cases  is   to    administer 
antidotes  and  to  remove  the  lead  from  the  skin  by  re- 
peated sulphur  baths,  which  convert  the  soluble  salts  into 
plumbic  sulphide,  and  this  can  be  rubbed  off  the  skin. 


280  SILVER. 

260.  Tests. 

1.  If  the  solution  is  not  too  dilute,  hydrochloric  acid  gives  a 
white  precipitate  insoluble  in  a  further  quantity  of  the  acid  and 
unchanged  by  ammonia. 

2.  Sulphuretted  hydrogen  gives  a  black  precipitate  (PbS)  in- 
soluble in  solution   of    ammonic    sulphide,    partly    dissolved, 
partly  whitened  by  strong  nitric  acid. 

3.  Sulphuric  acid  gives  a  white  precipitate  (PbS04). 

4.  Potassic  iodide  (KI)  gives  a  yellow  precipitate  (PbI2). 

5.  Potassic  bichromate  (K2Cr207)  gives  a  yellow  precipitate 
(PbCr04). 

6.  Insoluble  compounds  are  detected  by  reducing  on  charcoal 
with  sodic   carbonate,   dissolving   the   metallic   bead   in   nitric 
acid,  and  making  the  above  tests. 


SILVER  (Argentum). 

261.  Silver.  —  (Ag'  =  107.66.     Specific    weight  = 
10.5.      Melting    point  =  1040°    C.       Specific    heat  = 
0.057.) 

OCCURRENCE. — Native,  often  in  large  masses ;  the 
principal  ores  are  silver  glance  (Ag2S),  ruby  silver 
(Ag3SbS3),  silver  copper  glance  (Ag2S.Cu.2S),  and  horn 
silver  (AgCl). 

PREPARATION. — 1.  Silver  is  extracted  from  argenti- 
ferous lead  by  Pattinson's  process,  which  consists  in 
melting  the  lead  and  then  cooling  it  slowly.  Pure  lead 
crystallises  and  sinks  to  the  bottom.  This  is  ladled  out, 
until  the  remaining  molten  lead  contains  a  considerable 
percentage  of  silver,  when  the  process  is  finished  by 
cupellation.  The  lead  is  melted  in  bone-ash  cupels,  or 
shallow  vessels,  and  subjected  to  a  blast  which  oxidises 
the  lead  and  leaves  the  silver. 


ELECTROPLATING.  281 

2.  Amalgamation   Process. — The    silver  is    extracted 
with  mercury,  which  is  distilled,  leaving  the  silver  free. 

3.  The  silver  is  dissolved  out  of  its  ores    by   acids, 
sodic  thiosulphate,  &c.,  and  then  precipitated  by  scraps 
of  copper. 

PROPERTIES. — A  pure  white  metal,  the  best  conductor 
of  heat  and  electricity  known  ;  very  tough  and  ductile  ; 
and  very  soft  when  pure. 

NOTE. — Keep  all  residues  from  experiments  with  silver. 

Experiment  232. — Cut  small  pieces  of  silver  from  a  silver 
coin,  and  try  their  solubility  in  dilute  hydrochloric,  nitric,  and 
sulphuric  acids.  Try  the  strong  acids  with  the  aid  of  heat. 

Silver  is  soluble  in  nitric  acid,  argentic  nitrate  being 
formed.  It  is  insoluble  in  hydrochloric,  and  in  dilute 
sulphuric  acid ;  but  it  dissolves  in  strong,  hot  sulphuric 
acid,  with  the  formation  of  argentic  sulphate  (Ag2SO4) 
and  sulphur  dioxide : 

2Ag  +  2H2S04  =  Ag2SO4  +  2H2O  +  SO2. 

Silver  is  easily  reduced  from  its  compounds. 

Experiment  233. — Put  a  scrap  of  zinc  in  a  small  quantity  of 
argentic  nitrate  solution  in  a  porcelain  basin.  After  a  short 
time,  a  dark  powder  is  formed  in  its  place.  Fuse  this  on  char- 
coal with  the  blow-pipe.  A  bright  silver  bead  is  obtained. 
Dissolve  it  in  nitric  acid,  and  keep  the  solution  to  test. 

Silver  is  much  used  for  plating  inferior  metals.  This 
is  now  generally  done  by  means  of  electricity,  in  the 
process  known  as  electro-plating.  The  object  to  be 
plated  is  fastened  to  the  negative  wire  of  a  galvanic 
battery,  and  immersed  in  an  aqueous  solution  of  the 
double  cyanide  of  silver  and  potassium  (AgCN.KCN), 


282  LUNAR    CAUSTIC. 

through  which  the  electric  current  is  passed.  The  silver 
is  deposited  in  an  even  layer  on  the  object. — Silver  is 
easily  tarnished  by  sulphur  or  hydric  sulphide,  argentic 
sulphide  (Ag2S)  being  formed.  This  can  be  removed  by 
washing  with  solution  of  ammonia,  or  of  sodic  thio- 
sulphate. 

262.  Oxides  Of  Silver.— There  are  three  (Ag4O, 
Ag2O,  Ag2O2),  but  only  one  of  importance,  viz.,  argentic 
oxide,  Ag2U. 

Experiment  234. — To  solution  of  argentic  nitrate,  add  a 
little  caustic  soda.  Argentic  oxide  is  precipitated.  Filter, 
wash,  dry,  and  heat  on  mica.  The  oxygen  is  driven  off  and 
metallic  siver  remains  : 

(1)  2AgNO,  +  2NaOH  =  2NaN03  +  Ag.O  +  H20. 

(2)  Ag20  =  2Ag  +  O. 

Argentic  oxide  is  used  in  medicine.  It  is  less  liable 
than  argentic  nitrate  to  colour  the  skin. 

263  Salts  of  Silver.— The  only  salt  used  in  medi- 
cine is  the  nitrate. 

ARGENTIC  NITRATE  (AgNO3),  or  lunar  caustic,  is  pre- 
pared by  dissolving  pure  silver  in  dilute  nitric  acid  with 
the  aid  of  a  gentle  heat,  evaporating  to  dry  ness,  and 
fusing. 

PROPERTIES.  —  A  colourless  solid ;  sold  either  in 
sticks  or  as  tabular  crystals.  It  has  a  strong  metallic 
taste,  and  is  a  violent,  irritant  poison.  Antidotes — 
soluble  chlorides,  especially  common  salt. 

Experiment  235. — To  a  little  argentic  nitrate  solution  add 
solution  of  common  salt : 

AgN03  +  NaCl  =  AgCl  +  NaN03. 


MERCURY.  283 

Test  the  solubility  of  the  precipitate  (AgCl)  in  nitric  acid,  and 
in  ammonia. 

Argentic  nitrate  is  very  soluble  in  water  (2^  parts  in 
1),  and  in  alcohol.  The  solid  substance  is  used  as  a 
caustic.  If  nitrate  of  silver  be  administered  in  small 
doses  for  a  long  time,  it  may  produce  permanent  colora- 
tion of  the  skin.  It  darkens  the  skin  when  applied  ex- 
ternally. 

264.  Tests. 

1.  To  a  small  quantity  of  argentic  nitrate  solution  add  hydro- 
chloric acid.     A  curdy  white  precipitate  (AgCl)  forms.     Divide 
it  into  three  portions.    To  one  add  nitric  acid  ;   no  change.    To 
another  ammonia  ;  dissolved.     Let  the  third  stand  ;  it  darkens. 

2.  Sulphuretted  hydrogen  gives  a  black   precipitate  (Ag2S), 
insoluble  in  ammonic  sulphide. 

3.  Heat  with  solution  of  ferrous  sulphate.     Metallic  silver  is 
precipitated.     The  ferrous  sulphate  is  oxidised. 

4.  Potassic    bichromate   (K2Cr207)    gives   a    bright,    reddish 
purple  precipitate  of  argentic  chromate  (Ag,CrO4). 

5.  Potassic  iodide  (KI)  gives  a  yellow  precipitate  of  argentic 
iodide  (Agl),  insoluble  in  nitric  acid,  and  whitened,  but  not  dis- 
solved, by  ammonia.     Potassic  bromide  (KBr)  similarly.     These 
salts  of  silver,  especially  the  latter,  are  used  in  photography. 

Note. — Many  salts  of  silver  are  insoluble  in  water,  and  as 
they  have  characteristic  colours  and  other  properties,  argentic 
nitrate  is  used  as  a  group  reagent  in  testing  for  acids. 


MERCURY  (Hydrargyrum). 

265.  Mercury   (Hg"  =  199.8.     Specific    weight  = 
13.595.     Melting  point  =     -  39°  C.     Boiling    point  = 
357°  C.     Specific  heat  =.  0.0319.)     The  principal  ore  of 
mercury  is  cinnabar  (HgS),  from  which  the  mercury  is 


284  AMALGAMS. 

obtained  by  roasting  in  a  current  of  air  and  condensing 
the  vapours  of  mercury  in  a  series  of  cool  chambers.  It 
is  purified  by  distillation. 

NOTE. — Keep  all  residues  from  experiments  with  mer- 
cury. 

Experiment  236- — Put  a  little  mercuric  sulphide  (vermilion) 
in  a  small  hard-glass  tube  open  at  both  ends.     Hold  the  tube 
aslant  and  heat  the  sulphide  until  it  disappears.     Metallic  mer- 
cury collects  in  the  upper  part  of  the  tube  : 
HgS  -f  02  =  Hg  +  S02. 

PROPERTIES. — Mercury  is  commonly  called  quick  silver, 
which  means  living  silver.  It  is  brilliant  when  pure 
and  does  not  readily  tarnish  in  air.  If  it  is  impure 
(containing,  lead,  antimony,  &c.),  it  soon  gets  a  grey 
coating,  and  when  allowed  to  run  over  white  paper  leaves 
a  "  tail." — When  rubbed  up  with  chalk,  fats,  and  other 
substances,  it  becomes  partially  oxidised  and  very  finely 
divided.  It  is  in  this  way  that  mercury  ointments  and 
pills  are  made. — The  specific  weight  of  gaseous  mercury 
is  6.93  (air  =1).  (Calculate  the  number  of  atoms  in 
the  molecule).  Mercury  evaporates  slowly  even  at  low 
temperatures,  and  it  should,  therefore,  always  be  kept  in 
closed  vessels. — On  account  of  its  high  specific  weight  it 
is  used  for  barometers,  areometers,  &c. 

AMALGAMS. — Mercury  unites  with  all  metals  but  iron 
to  form  amalgams.  From  these  the  mercury  can  be 
driven  oft  by  heat.  This  property  is  utilised  in  silvering 
by  means  of  a  silver  amalgam.  Cadmium  amalgam 
(Hg2Cd)  is  very  brittle,  heavier  than  mercury,  and  has 
the  property  of  hardening  gradually.  It  is  used  for 
filling  teeth.  Copper  amalgams  are  brittle  at  ordinary 
temperatures,  but  soften  at  100°  C.  They  are  used  for 
filling  teeth  and  for  sealing  bottles,  &c. 


CALOMEL.  285 

266.  Mercurous  Compounds.  —  In    these    com- 
pounds two  atoms  of  mercury   act  as  a  dyad  radical, 

Hg— 

I       .     The  salts  are  insoluble,  or  sparingly  soluble,  in 
Hg— 

water,  and  can  be  converted  by  oxidation,  &c.,  to  mer- 
curic salts. 

267.  Mercurous  Nitrate.— Hg2(NO3)2. 

Experiment  237- — Pour  some  dilute  nitric  acid  over  a  glo- 
bule of  mercury,  and  allow  it  to  stand  for  some  time.  The 
metal  gradually  dissolves,  forming  a  solution  of  mercurous 
nitrate.  Keep  this  for  further  examination. 

Mercurous  nitrate  is  soluble  in  water,  but  with  a 
large  proportion  of  water  it  is  partially  decomposed,  with 
the  formation  of  a  yellow  basic  nitrate  : 

Hg,(N03)2  +  H20  =  Hg2.N03.OH  +  HNO3. 

268.  Mercurous  Chloride  (Hg2Cl2).    Also  called 

calomel.     This  name  is  from  the  Greek,  and  means  beau- 
tiful black. 

Experiment  238 . — To  a  small  quantity  of  mercurous  nitrate 
solution  (Exp't  237)  add  hydrochloric  acid.  A  white  precipitate 
of  mercurous  chloride  is  thrown  down  : 

Hg2(NO3)2  +  2HC1  =  Hg2Cl2  +  2HN03. 
To  this  add  lime  water  until  the  chloride  is  blackened.     This  is 
the  black  wash  of  the  Pharmacopoeia  : 

Hg2Cla  +  Ca(OH),,  =  Hg20  +  H20  +  CaCl2. 

Calomel  is  usually  prepared  by  subliming  a  mixture 
of  mercuric  sulphate  (HgSO4),  mercury,  and  sodium 
chloride : 

HgSO4  +  Hg  +  2NaCl  =  Hg2Cl2  +  Na2SO4. 
The  sodic  sulphate  is  not  volatile. 


286  MERCUROUS    IODIDE. 

PROPERTIES. — A  white,  tasteless,  inodorous  powder, 
insoluble  in  water,  and  in  acids,  ('fry  with  a  specimen 
of  calomel.)  It  is  turned  black  by  alkalis,  owing  to  the 
formation  of  mercurous  oxide  (Hg2O),  the  active  consti- 
tuent of  black  wash.  Precipitated  calomel  is  more  active 
as  a  medicine  than  that  prepared  by  sublimation,  owing 
to  its  finer  state  of  division ;  but  great  care  should  be 
taken,  in  preparing  it  by  precipitation,  that  no  basic 
nitrate  be  present.  This  would  dissolve  readily  in  the 
acid  juices  of  the  stomach  and  might  cause  mercurial 
poisoning.  In  order  to  guard  against  this,  warm  the 
precipitated  calomel  with  dilute  hydrochloric  acid,  filter, 
and  wash  well. — Calomel  may  be  administered  in  doses 
up  to  6  grains. 

269.  Mercurous  Iodide,  Hg2I2,  green  iodide  of 
mercury,  or  proto-iodide  of  mercury.  —  A  dull  green 
powder,  prepared  by  rubbing  together  iodine  and  mer- 
cury (In  what  proportions'?),  occasionally  moistening 
with  spirits  of  wine. 

Experiment  239- — Add  solution  of  potassic  iodide  to  solution 
of  mercurous  nitrate.  A  greenish  precipitate  of  mercurous 
iodide  appears  : 

2KI  +  Hg.(N03)2  =  HgaI2  +  2KN03. 

Mercurous  iodide  readily  changes  into  mercuric  iodide 
(HgI2)  and  mercury.  As  the  mercuric  salt  is  more 
readily  dissolved  than  the  mercurous  (and  therefore 
more  active  as  a  poison  \  care  should  be  taken  that  it  be 
absent  from  the  preparation.  The  mercurous  iodide' 
should  be  kept  away  from  the  light,  which  tends  to 
bring  about  this  decomposition  : 

Hg2I2  =  Hg  +  HgI2. 


MERCURIC   NITRATE.  287 

270.  Tests  for  Mercurous  Compounds. 

1.  White   precipitate    with   hydrochloric   acid,    blackened   by 
ammonia : 

Hg2Cl2  +  2NH3  =  Hg2Cl.NH2  +  NH4C1. 

2.  Insoluble  compounds  disappear  when  heated  on  mica,  and 
are  blackened  by  caustic  soda. 

271.  Mercuric    Compounds.  —  Tn    these    com- 
pounds a  single  atom  of  mercury  acts  as  a  dyad  (Hg     ). 
They  are  more  soluble  than  the  mercurous  compounds, 
and  are  generally  deadly  poisons. 

272.  Mercuric  Nitrate,  Hg(NO3)2,  is  prepared  by 
boiling  mercury  with  excess  of  nitric  acid  until  the  solu- 
tion no  longer  gives  a  precipitate  with  hydrochloric  acid. 
(What  is  the  precipitate  1) 

Experiment  240. — Heat  a  globule  of  mercury  in  a  porcelain 
dish  with  nitric  acid  diluted  with  an  equal  volume  of  water. 
From  time  to  time  take  out  a  drop  of  the  solution  with  a  glass 
rod,  and  mix  it  with  a  drop  of  hydrochloric  acid.  Continue 
heating,  adding  more  acid  if  necessary,  until  the  drop  remains 
clear  on  mixing  with  hydrochloric  acid.  Evaporate  to  dryness. 
Examine  the  residue,  dissolve  part  of  it  in  water,  and  test  it 
with  lime  water.  It  turns  yellow  : 

Hg(N03)2  +  Ca(OH)2  =  Hg(OH)2  +  Ca(N03)2. 

Heat  a  little  of  the  dry  salt  carefully  in  a  dry  porcelain  capsule. 
Red  fumes  are  evolved,  and  red  oxide  of  mercury,  or  mercuric 
oxide  (HgO),  remains. 

An  acid  solution  of  this  salt  is  used  in  medicine  as  a 
caustic.  It  is  an  irritant  poison. 

273.  Mercuric  Sulphate,  HgSO4,  is  prepared  by 
dissolving  mercury  in  hot,  strong  sulphuric  acid  : 

Hg  +  2H2S04  =  HgSO4  +  2H2O  +  SO2. 
(Calculate  the  proportions  to  be  used.) 


288  CORROSIVE    SUBLIMATE. 

Experiment  241. — Heat  a  drop  of  mercury  in  a  small  por- 
celain vessel  with  about  7  or  8  times  its  volume  of  concentrated 
sulphuric  acid,  stirring  constantly.  The  mercury  dissolves. 
When  cold,  test  a  small  portion  of  the  salt  with  caustic  soda. 
It  is  turned  yellow  : 

HgS04  +  2NaOH  =  Hg(OH)2  +  Na2S04. 

Mercuric  sulphate  is  used  in  the  preparation  of  cor- 
rosive sublimate  (HgCL). 

274.  Mercuric  Chloride,  HgCl2,  or  corrosive  sub- 
limate, is  prepared  by  subliming  a  mixture  of  inercimc 
sulphate  and  common  salt,  a  little  manganese  dioxide 
being  added  to  make  sure  of  the  absence  of  mercurous 
salts.  (How  1)  -. 

HgS04  +  2NaCl  -  HgCl2  +  Na2SO4 

PROPERTIES. — Heavy  colourless  crystals,  having  a 
biting  metallic  taste  (Examine  a  specimen) ;  soluble  in 
water  (7  parts  in  100),  more  so  in  alcohol,  and  still 
more  so  in  ether.  It  volatilises  more  readily  than 
calomel  when  heated.  The  medicinal  dose  is  ^  to  £ 
of  a  grain.  In  larger  doses  it  is  a  violent  poison.  The 
antidotes  are  white  of  egg,  flour,  &c.  Stannous  chloride 
(SnCl2)  may  be  given  : 

SnCl2  +  2HgCl2  =  SnCl4  +  Hg2012 
(How  does  this  prevent  the  poisonous  action?) 

Experiment  242. — Mix  solutions  of  mercuric  and  stannous 
chlorides.  What  is  the  precipitate  formed  ? 

Experiment  243. — To  solution  of  corrosive  sublimate  add 
lime  water.  A  yellow  precipitate  is  formed  : 

HgCl2  +  Ca(OH)2  =  Hg(OH).  +  CaCl,. 
This  is  the  yellow  wash  of  the  Pharmacopreia. 


MERCURIC    IODIDE.  289 

Corrosive  sublimate  is  an  excellent  antiseptic,  even  in 
very  dilute  solution,  and  is  largely  used  instead  of  car- 
bolic acid. 

275.  Mercuric    Oxide. — HgO.       There    are    two 
varieties,  the  differences  depending  on  the  fineness  of 
division.     Red  oxide  of  mercury  is  prepared  by  heating 
mercury  with  mercuric  nitrate  : 

Hg  4-  Hg(N03)2  =  2HgO  4-  2NOa. 

It  is  a  heavy,  red  powder,  soluble  in  those  acids  which 
form  soluble  mercuric  salts.  (What  is  the  effect  of  a 
strong  heat  upon  it  1) — The  yellow  oxide  is  prepared  by 
precipitation.  A  solution  of  mercuric  chloride  is  treated 
with  solution  of  caustic  soda,  and  the  precipitate  of  mer- 
curic hydroxide  (Hg(OH)2)  is  washed,  and  dried  at 
100°.  This  method  of  preparing  the  hydroxides  and 
oxides  of  the  heavy  metals  is  very  generally  employed. 
In  several  cases,  the  soluble  base  used  as  a  precipitant  is 
ammonia  instead  of  soda. 

276.  Mercuric  Iodide,  HgI2.  Also  called  red  iodide 
of  mercury. 

Experiment  244.  — To  solution  of  mercuric  chloride  add  solu- 
tion of  potassic  iodide  drop  by  drop.  A  yellowish  precipitate 
appears,  but  immediately  turns  red  : 

2KI  +  HgCl2  =  HgI2  4-  2KC1. 

Add  more  potassic  iodide.  The  precipitate  of  mercuric  iodide 
redissolves,  a  soluble  double  salt  (HgI2.2KI)  being  formed. 
This  solution,  with  caustic  soda  added  to  it,  constitutes  Nes»- 
ler's  reagent,  a  very  delicate  test  for  ammonia,  used  in  water 
analysis.  Try  it  with  a  drop  of  ammonia  dissolved  in  about  half 
a  litre  of  distilled  water.  A  brown  colour  or  precipitate  ap- 
pears : 

NH3  4-  2HgI2  4-  3NaOH  =  NHg,I  4-  3NaI  4-  3H2O. 
20 


290  WHITE    PRECIPITATE. 

Mercuric  iodide  is  a  brilliant  crystalline  powder  of  a 
vermilion  colour.  It  turns  yellow  when  heated  care- 
fully. It  is  very  slightly  soluble  in  water,  sparingly  in 
alcohol,  quite  freely  in  ether  and  in  aqueous  solution  of 
potassic  iodide.  It  resembles  corrosive  sublimate  in  its 
poisonous  action. 

277.  Infusible  White  Precipitate,  or  mercuric 
ammonium  chloride,  NHgH2Cl. 

Experiment  245.—  Add  ammonia  solution  to  solution  of 
mercuric  chloride.  A  white  precipitate  falls  :_ 

HgCl2  +  2NH3  =  NHgHaCl  +  NH4C1. 

278.  Fusible  White  Precipitate,  or  mercuric  di- 
ammonium  chloride,  Hg(NH3Cl)2,  is  prepared  by  adding 
solution  of  mercuric  chloride  to  a  boiling  solution  of  am- 
monic  chloride  and   ammonia. — These  two  compounds 
differ  as  their  names  indicate.     The  infusible  precipitate 
decomposes    without  •  fusing  when   heated ;    the   fusible 
precipitate  fuses  and  then  decomposes. 

279.  Mercuric  Sulphide,  HgS,  is  found  in  nature 
as  a  heavy  red  mineral,  cinnabar.     It  can  be  prepared 
by  subliming  mercury  and  sulphur  together,  when  it  is 
obtained  as  a  red  or  black  powder,  vermilion,  or  JEthiops 
mineral.     Or,  it  can  be  prepared  as  a  black  powder  by 
precipitation. 

Experiment  246- — Add  hydric  sulphide  to  solution  of  mer- 
curic chloride.  A  white  precipitate  appears  (2HgS.Hg012),  but 
this  rapidly  passes  through  shades  of  colour  until  it  becomes 
black  (HgS).  (Write  the  equation).  Divide  the  precipitate 
into  three  parts  and  try  its  solubility  (1)  in  ammonic  sulphide, 

(2)  in  boiling   dilute  nitric  acid,   and  (3)  in  aqua  regia.      Re- 
sults (1) (2)   

(3)   


MERCURIAL    POISONING.  291 

280.  Tests  for  Mercuric  Compounds. 

1.  See  Experiment  246. 

2.  Sodic  hydroxide  solution  gives  with  solutions  of  mercuric 
salts  a  yellow  precipitate  (Hg(OH)2)  not  redissolved  when  more 
of  the  reagent  is  added. 

3.  Stannous  chloride  (SnCl2^  gives  a  white  precipitate  (Hg2Cl2), 
turning  grey  (Hg)  with  more  of  the  reagent. 

4.  A  bright  copper  wire  is  silvered  when  put  in  a  mercuric 
solution.     This  applies  to  mercurous  solutions  also. 

General  Test  for  Insoluble  Mercury  Compounds. 
— Mix  with  dry  sodic  carbonate  and  charcoal  powder,  and  heat 
in  a  matrass.  Metallic  mercury  is  obtained  as  a  mirror  in  the 
tube. — All  mercury  compounds  volatilise  when  strongly  heated. 

281.  Mercurial  Poisoning. — The    soluble    com- 
pounds   of   mercury    are    violent    irritant    poisons.     A 
characteristic  symptom  is   the   increased  flow  of  saliva 
(salivation}.     The  insoluble  compounds  are  not  so  poison- 
ous, but  are  still  dangerous.     Even  metallic  mercury  in 
a  finely  divided  state  will  cause  symptoms  of  poisoning, 
especially  when  it  is  breathed  as  a  vapour.     Therefore, 
it  is  dangerous  to  boil  mercury,  or  to  sublime  compounds 
of  mercury,  into  the   atmosphere  of  an  inhabited  room. 
As  mercury  and  its  compounds  are  very  extensively  used 
in  arts  and  manufactures,  as  well  as  in  many  patent  and 
quack   medicines,   numerous    cases  of  .poisoning   occur. 
The  antidotes  are  albuminous  substances,  such  as  white 
of  egg,  flour,  <fec.,  and  stannous  chloride. 


292  QUESTIONS    AND    EXERCISES. 

QUESTIONS  AND  EXERCISES. 

1 .  Is  there  any  reason  why  rain  water  should  not  be  stored  in 
lead-lined  tanks  ? 

2.  In  what  liquids  is  lead  soluble,  and  in  what  insoluble  ? 

3.  Mention  some  liquids  which  will,  and  some  which  will  not 
dissolve  litharge. 

4.  A  white  crystalline  solid  is  given  you.     How  would  you 
determine  whether  it  is  sugar  of  lead? 

5.  What  proportions  of  plumbic  acetate  (Pb(C2H30.2)2.3H2O) 
and  litharge  (PbO)  must  be  combined  to  form  the  subacetate  of 
lead? 

6.  Why  should  solution  of  subacetate  of  lead  be  kept  well 
closed  from  the  air  ? 

7.  What  chemical  action  takes  place  when  white  lead  is  taken 
into  the  stomach  ? 

8.  In  what  acids  is  mercury  soluble  ? 

9.  Why  is  mercurous  chloride  called  calomel  ? 

10.  What  chemical  compounds  are  present  in  black  wash  ?    In 
yellow  wash  ? 

11.  Write  formulas  for  mercurous  sulphate,  mercurous  bromide, 
and  mercuric  cyanide. 

12.  How  would  you  prepare  a  specimen  of  mercuric  hydroxide"! 

13.  Black  mercuric  sulphide  becomes  red  when  rubbed.     Can 
you  account  for  this  ? 

14.  Explain  the  action  of  stannous  chloride  as  an  antidote  to 
poisoning  by    corrosive    sublimate.      Would    stannic    chloride 
(SnClJ  do  ? 

15.  Argentic   cyanide   is   insoluble  in  water,   but   soluble  in 
solution  of  potassic  cyanide.     What  happens  when  a  drop  of 
argentic  nitrate  solution  is  added  to  a  considerable  quantity  of 
potassic  cyanide  and  shaken  up  with  it  ? 


COPPER.  293 

CHAPTER     XVIII. 


METALS  OF  GROUP  II. 

[Lead,    Mercury],    Copper,    Cadmium,    Bismuth;    Anti- 
mony, Tin,  [Arsenic],  Gold,  Platinum,  dec. 

282.  General  Characters — The  sulphides  of  these 

metals  are  precipitated  from  acid  solutions  by  hydric  sul- 
phide, and  also  by  ammonic  sulphide,  but  some  of  them  are 
redissolved  by  a  further  quantity  of  this  latter  reagent. — 
The  hydroxides  are  precipitated  from  solutions  of  salts 
by  alkaline  hydroxides,  e.g.,  sodic  hydroxide.  —  The 
oxides  can  be  prepared  by  heating  the  hydroxides.  They 
are  earthy  substances  insoluble  in  water,  but  soluble  in 
acids. — The  carbonates  and  phosphates  are  earthy  sub 
stances  insoluble  in  water. 

This  group  is  subdivided  into  two  : 

A.  Metals  having  sulphides  insoluble  in  solutions  of 
alkaline  sulphides  :    [Lead],  copper,   cadmium,  and  bis- 
muth.    In  analysis,  mercuric  salts  are  included  in  this 
class. 

B.  Metals  having  sulphides    soluble   in    solutions  of 
alkaline  sulphides  :    Antimony,  tin,  [arsenic],  gold,  plati- 
num, and  some  rare  metals. 


A. 

COPPER  (Cuprum). 

283.  Copper  (Cu"  =  63.1.     Specific  weight  =  8.92. 
Melting  point  =  1090°  C.     Specific  heat  =  0.0952). 


294  COPPER. 

OCCURRENCE. — Free;  sometimes  in  great  masses.  The 
ores  of  copper  are  very  numerous.  The  commonest  one 
is  copper  pyrites,  or  chalcopyrite  (CuFeS2).  Copper  is 
found  in  the  liver  and  kidneys  of  man  and  of  domestic 
animals. 

PREPARATION. — The  smelting  of  copper  is  a  very  com- 
plicated process,  but  it  is  the  same  in  principle  as  that 
of  lead,  viz.,  a  partial  oxidation  of  the  sulphide  and  sub- 
sequent fusion  at  a  higher  temperature.  Poor  ores  are 
worked  up  by  dissolving  in  acids  and  precipitating  with 
iron. 

Experiment  247- — Dip  the  point  of  a  knife  blade  in  a  little 
solution  of  cupric  sulphate.  It  becomes  coated  with  copper  : 

CuS04  -f  Fe  =  FeS04  +  Cu. 

PROPERTIES. — A  red  metal,  heavy,  very  tough,  malle- 
able, and  ductile.  Next  to  silver  it  is  the  best  conductor 
of  heat  and  electricity.  Its  solubility  in  nitric  and  sul- 
phuric acids  has  been  already  proved.  (See  Arts.  86 
and  115). 

Experiment  248. — Try  the  solubility  of  copper  in  hydro- 
chloric acid,  strong  and  dilute.  Put  a  piece  of  bright  copper 
wire  in  dilute  acetic  acid,  so  that  it  is  half  covered  with  the 
acid.  Leave  it  for  a  day,  and  then  examine  it.  Test  the  acid 
for  copper.  Verdigris,  or  basic  cupric  acetate,  has  been  formed. 

Experiment  249- — Put  pieces  of  bright  copper  wire  in  test 
tubes  or  beakers  containing  (1)  distilled  water,  (2)  dilute  solu- 
tion of  common  salt,  (3)  butter  or  fat,  and  (4)  solution  of  sugar. 
After  24  hours,  examine  the  condition  of  the  copper  and  test 
the  solutions  for  copper. 

Copper  is  gradually  dissolved  by  water,  especially 
when  the  water  contains  ammonium  salts  or  chlorides. 
Vinegar,  fats,  oils,  and  syrups,  with  the  aid  of  air  and 


BLUE    VITRIOL.  295 

moisture,  dissolve  copper.  Thus  occur  frequent  cases  of 
poisoning,  from  copper  kitchen  utensils,  taps  for  liquors, 
<fec. 

284.  Compounds  Of  Copper. — Copper  resembles 
mercury  in  forming  two  classes  of  compounds,  (1)  cup- 
rous, in  which  two  atoms  of  copper  act  as  a  dyad  radical 

Cu— 

I  ,  and  (2)  cupric,  in  which  one  atom  of  copper  replaces 
Cu — 

two  of  hydrogen.— The  cuprous  salts  are  so  easily  oxidised 
that  it  is  difficult  to  keep  them.  Cuprous  oxide  (Cu2O) 
has  been  already  noticed.  (See  Sugars).  Cuprous 
chloride  (Cu2Cl2)  is  remarkable  as  being  a  solvent  for 
acetylene  and  carbon  monoxide. 

285.  Cupric  Sulphate.— CuSO4.5H2O.   Also  called 
blue  vitriol  and  blue  stone. 

PREPARATION. — Copper  pyrites  is  roasted  so  as  to  form 
cupric  oxide  (CuO)  and  ferric  oxide  (Fe2O3).  The  cupric 
oxide  is  then  dissolved  out  by  hot  sulphuric  acid,  in 
which  ignited  ferric  oxide  is  insoluble.  The  solution  is 
evaporated,  and  the  salt  crystallised.  The  commercial 
salt  always  contains  a  little  ferrous  sulphate  (FeSO4. 
7H2O).  Blue  vitriol  is  obtained  as  a  secondary  product 
in  the  refining  of  silver  by  precipitation  on  copper.  (In 
what  way  has  it  been  already  prepared "?) 

PROPERTIES. — Cupric  sulphate  is  generally  sold  in 
large  blue  crystals.  These  are  soluble  in  water  (2  parts 
in  5). 

Experiment  250. — Carefully  heat  a  crystal  of  cupric  sulphate 
in  a  t.  t.  It  turns  white,  and  water  gathers  on  the  sides  of  the 
tube.  The  crystal  falls  to  a  powder,  because  it  has  lost  its 
water  of  crystallisation.  When  the  t.  t.  is  cool,  pour  a  few 


296  OUPRIC    OXIDE. 

drops  of  water  on  the  anhydrous  salt.  Note  signs  of  heat. — 
Anhydrous  cupric  sulphate  is  used  in  testing  liquids  for  water. 
It  turns  blue  when  acted  on  by  water. 

Experiment  251. — Dissolve  a  little  cupric  sulphate  in  water, 
and  test  the  solution  with  blue  litmus  paper.  The  basic  part  of 
the  salt  is  comparatively  weak.  Taste  the  solution.  Test  it 
for  sulphuric  acid. 

Blue  vitriol  is  used  in  medicine  as  a  caustic,  and  also 
as  an  emetic.  In  small  doses  (up  to  2  grains)  it  is  not 
poisonous,  but  acts  as  a  tonic  and  astringent.  In  larger 
doses  it  is  poisonous,  unless  it  exerts  its  emetic  action. 
Antidotes,  white  of  egg,  <fec.  Cupric  sulphate  is  an  anti- 
dote to  phosphorus. 

286.  Cupric  Oxide. — CuO.     This  is  the  black  oxide 
of  copper. 

Experiment  252.— Heat  a  little  powdered  cupric  sulphate 
strongly  on  mica.  Black  oxide  of  copper  is  left  : 

CuSO4.5H20  =  CuO  +  SO3  +  5H.20. 

Cupric  oxide  is  a  black  hygroscopic  powder,  soluble 
in  acids,  insoluble  in  water.  It  is  used  in  organic 
analysis  to  supply  oxygen  to  oxidisable  substances. 

287.  Commercial  Preparations  of  Copper. 

1.  Scheele's  green.     (See  Arsenic.} 

2.  Schweinfurth    green,    or     emerald     green.       (See 
Arsenic.} 

3.  Brighton  green,  a  mixture  of  impure  cupric  acetate 
and  chalk. 

4.  Brunswick    green,   oxychloride  of    copper,  or  car- 
bonate of  copper  mixed  with  chalk. 


CADMIUM.  297 

5.  Mountain  green,  or  mineral  green,  is  a  native  car- 
bonate of  copper. 

6.  Green  verdites,  a  mixture  of  cupric  oxide  and  car- 
bonate with  chalk. 

7.  Verdigris,  a  basic  acetate. 

Many  alloys  of  copper  are  used,  e.g.,  brass,  tombac, 
Muntz  metal,  bronze,  &c.     (See  Art.  238.) 

288.  Tests. 

1.  Acidify  a  solution  of  cupric  sulphate  with   hydrochloric 
acid,   and  add  hydric  sulphide.     A  black  precipitate  of  cupric 
sulphide  (CuS)  falls : 

CuSO4  +  HZS  =  CuS  +  H2S04. 

Filter,  wash,  and  test  the  solubility  of  a  portion  of  the  precipi- 
tate in  yellow  ammonic  sulphide.  It  is  insoluble  (in  reality, 
sparingly  soluble).  Heat  another  portion  in  a  porcelain  dish 
with  dilute  nitric  acid.  It  is  dissolved.  Heat  a  third  portion 
with  dilute  sulphuric  acid.  It  is  undissolved. 

2.  Add  ammonia  solution  gradually  to  cupric  sulphate  solu- 
tion.    A  light  blue  precipitate  is  first  formed.     When  more  am- 
monia is  added,  this  dissolves  to  a  deep  blue  solution  containing 
a  compound,  CuS04.4NH3. 

3.  Potassic  ferrocyanide  gives  a  reddish-brown  precipitate,  or, 
with  very  dilute  solutions,  a  reddish  colour. 

4.  Insoluble  compounds  may  be  tested  by  heating  a  little  of 
the  substance  on  a  platinum  wire  with  a  borax  bead.     The  bead 
is  green  while  hot,  blue  when  cold.     If  it  be  moistened  with 
solution  of  stannous  chloride  and  heated  in  the  inner  (reducing) 
zone  of  the  Bunsen  flame,  it  becomes  coppery  red  when  cold. 


CADMIUM. 

289.  Cadmium.— (Cd11  =  111.6.  Sp.  wt.  -  8.5. 
Melting  point  =  315°C.  Boiling  point  =  860°C.  Sp. 
heat  -  0.0567). 


298  CADMIC    NITRATE. 

OCCURRENCE. — Along  with  zinc  ores.  In  smelting 
zinc  ores,  cadmium  volatilises  first  (Compare  boiling 
points  \  burns  when  it  reaches  the  air,  and  the  oxide 
(CdO)  collects  as  a  brown  dust  in  the  flue  of  the  furnace. 

PREPARATION.— The  impure  cadmium  oxide  is  dis- 
solved in  hydrochloric  acid,  and  the  cadmium  is  then 
precipitated  as  sulphide  £CdS)  by  sulphuretted  hydrogen, 
the  zinc  salt  remaining  in  solution.  The  sulphide  is  dis- 
solved in  strong  hydrochloric  acid,  and  the  carbonate 
(CdCO3)  is  obtained  by  precipitating  with  sodic  carbonate. 
By  heating  the  carbonate,  pure  cadmium  oxide  is  formed, 
and  this  is  then  reduced  by  heating  in  iron  tubes  with 
charcoal. 

PROPERTIES. — A  white  metal,  somewhat  like  tin.  It 
crackles  when  bent.  It  is  harder  than  tin,  malleable  and 
ductile,  and  takes  a  good  polish.  Cadmium  is  soluble  in 
hot  dilute  hydrochloric  acid,  and  in  sulphuric  acid.  It 
is  very  easier  dissolved  by  nitric  acid. 

(1)  Cd  +  2HC1  =  CdCl2  +  H2 

(2)  Cd  +  H2S04  =  CdSO4  +  H2 

(3)  3Cd  +  8HNO3  =  3Cd(NO3)2  -f  2NO  +  4H2O. 

290.  Compounds   Of  Cadmium. — Cadmium    has 
only  one  oxide   (CdO),  a  brown  solid,  mentioned  above- 
The  salts  of  cadmium  are  mostly  colourless,  and  resemble 
those  of  zinc  both  in  chemical  properties  and  in  physio- 
logical action.     Solutions  of  the  normal   salts  have  an 
acid  reaction. 

291.  Cadmic  Nitrate,    Cd(NO3)2,   is   prepared  by 
dissolving  the  metal  in  nitric  acid  and  evaporating  the 
solution.     It  is  a  deliquescent  white  salt,  used  for  pre- 
paring other  salts,  and  in  chemical  experiments. 


CADMIC    IODIDE.  299 

292.  Oadmic  Sulphate,  3CdSO4.8H2O,  is  prepared 
from  the  nitrate  or  chloride. 

Experiment  253. — Pour  solution  of  cadmic  nitrate  or  chloride 
into  sodic  carbonate  solution  : 

Cd(N03)2  +  Na2C03  =  CdC03  +  2NaN03. 

Filter  off  the  precipitate  of  cadmic  carbonate,  wash  it,  and  dis- 
solve it  in  dilute  sulphuric  acid,  taking  care  not  to  use  too 
much.  Evaporate  to  crystallisation,  and  examine  the  crystals  : 

CdOO3  +  H4S04  =  CdSO4  +  H4O  +  CO.,. 

This  method  of  passing  from  one  soluble  salt  of  a  metal  to 
another  is  often  employed.  Sometimes  the  hydroxide  is  pre- 
cipitated instead  of  the  carbonate. 

Sulphate  of  cadmium  is  a  colourless  salt,  resembling 
sulphate  of  zinc  in  its  physiological  actions,  but  it  is 
more  powerful.  It  is  used  as  a  wash  for  diseases  of  the 
eye.  It  is  soluble  in  about  one  and  a  half  times  its 
weight  of  water. 

293.  Cadmic  Iodide,  CdI2,  is  prepared  by  digesting 
the  metal  with  iodine  and  water,  until  the  colour  of  the 
iodine    disappears.     A    solution    is  obtained   which,  on 
evaporation,  deposits  thin  pearly  plates  of  the  iodide.     It 
is  soluble  in  about  an  equal  weight  of  water. — Cadmium 
iodide  is  used  in  medicine  in  the  form  of  an  ointment. 
It  is  also   used  in   photography  in  preparing  the  sensi- 
tive   plates,  as  it  is  one  of  the  few   iodides    soluble  in 
alcohol  and  ether.     A  solution  in  alcohol  and  ether  is 
mixed  with   a  collodion  solution  and   spread  in  a  thin 
layer  upon  the  plate.     The  liquids  quickly  evaporate  and 
leave  a  thin  layer  of  collodion  impregnated  with  cadmium 


300  BISMUTH. 

iodide.      When  this  is  dipped  in  a  bath  of  argentic  nitrate, 
double  decomposition  takes  place  : 

CdI2  +  2AgNO3  =  2Agl  -f  Cd(NO3)3. 
This  is  the  sensitive  plate. 

294.  Tests. 

1.  To  a  cadmium  solution  add  hydric   sulphide.      A  yellow 
precipitate  of  cadmic  sulphide  (CdS)  is  formed.     This  is  insoluble 
in  yellow  ammonic  sulphide.     It  resembles  arsenic  trisulphide 
(As-^Sg)  and  stannic  sulphide  (SnS.J,  but  these  are  soluble  in 
ammonic  sulphide.     Cadmic   sulphide  is  soluble  in  hot  dilute 
sulphuric  and  nitric  acids. 

2.  To  a  solution  of  a  cadmium  salt  add  ammonia  gradually. 
A  white  precipitate  of  cadmic  hydroxide,  Cd(OH)2,  appears,  but 
redissolves  in  more  of  the  reagent,  forming  a  colourless  solution. 


BISMUTH. 

295.  Bismuth.  (Bi111^-  =  210.—  Sp.  wt.  =  9.8.— 
Melting  point  =  270°.— Sp.  heat  =  0.0305.) 

OCCURRENCE. — Rather  rare,  chiefly  in  the  free  condi- 
tion. It  is  found  often  with  ores  of  cobalt. 

PREPARATION. — Generally  as  a  by-product  in  smalt 
works,  being  separated  in  the  metallic  state  from  the 
sulphide  of  cobalt  by  smelting  with  iron  scraps.  The 
commercial  metal  nearly  always  contains  arsenic.  For 
medicinal  use  it  must  be  freed  from  this  by  melting  with 
a  little  saltpetre,  which  oxidises  the  arsenic,  but  does  not 
attack  the  bismuth. 

PROPERTIES. — A  hard,  lustrous,  brittle  metal,  of  a 
reddish  tint.  It  expands  on  solidifying,  and  gives  this 
property  to  its  alloys,  some  of  which  are  used  in  stereo- 
typing. It  decomposes  steam  at  a  red  heat.  It  oxidises 


BISMUTH    NITRATE.  301 

slowly  in  the  air.  Bismuth  is  not  dissolved  by  cold 
dilute  hydrochloric  or  sulphuric  acid. 

Experiment  254. — Dissolve  a  little  bismuth  in  hot,  strong 
sulphuric  acid  : 

2Bi  +  6H2SO4  =  Bi2(SO4)3  +  3SO2  +  6H2O. 

It  dissolves  easily  in  nitric  acid  and  in  aqua  regia. — 
Alloys  of  bismuth,  lead  and  tin,  made  to  melt  at  par- 
ticular temperatures,  are  used  as  safety  plugs  for  boilers. 

296.  Compounds  Of  Bismuth.— Bismuth  unites 
with  oxygen  in  four  proportions   (Bi.2O2,  Bi2O3,  Bi2O4, 
Bi2O5),   but  only   one  of  the  oxides,  viz.,  the  trioxide, 
(Bi2O3)  is  of  importance  medicinally.     It  is  a  basic  oxide. 
The  pentoxide  (Bi2O5)  is  acid-forming.     The  others  are 
indifferent. 

297.  Bismuth  Trinitrate,  Bi(N03)3.3H2O.    Also 

called  nitrate  of  bismuth. 

Experiment  255. — Dissolve  a  little  bismuth  in  nitric  acid 
diluted  with  about  three-fourths  its  volume  of  water : 
Bi  +  4HN03  =  Bi(N03)3  +  NO  +  2H20. 

Evaporate  the  solution  to  crystallisation.  Colourless  delique- 
scent crystals  are  obtained.  Any  arsenic  present  remains  dis- 
solved in  the  mother  liquor. 

298.  Bismuth  Subnitrate,  or  basic  nitrate  of  bis- 
muth, BiNO3(OH)2f.      Also  called  white  bismuth. 

PREPARATION.— Experiment  256. — Add  a  few  drops  of 
water  to  the  crystals  of  trinitrate  obtained  in  Experiment  255, 
so  as  to  form  a  solution,  and  pour  it  into  a  beaker  of  distilled 
water.  A  white  precipitate  falls  : 

Bi(NO3)3  +  2HaO  =  BiN03(OH).  -f  2HNO3 


t  The  composition  of  this  salt  varies  with  the  amount  of  water  used  in  its 
preparation. 


302  BISMUTH    TRIOXIDE. 

(Many  salts  of  weak  bases  can  be  thus  decomposed  by  water.) 
Filter  off,  wash,  and  dry  the  precipitate.      Keep  the  filtrate. 

PROPERTIES. — A  heavy  white  powder,  insoluble  in 
water,  but  soluble  in  moderately  strong  nitric  acid.  It 
is  again  precipitated  from  this  solution  by  the  addition 
of  water. 

Experiment  257- — Dissolve  a  little  bismuth  subnitrate  in  a 
few  drops  of  nitric  acid  and  a  drop  or  two  of  water.  Then  add 
more  water. 

Experiment  258. — Dissolve  a  little  bismuth  subnitrate  in 
sulphuric  acid  diluted  with  an  equal  volume  of  water.  To  this 
add  a  few  drops  of  ferrous  sulphate  solution.  A  black  colour  or 
precipitate  is  formed  (Bi202). 

Subnitrate  of  bismuth  is  much  used  in  medicine.  It 
is  not  poisonous,  but  it  sometimes  contains  the  tri-nitrate 
or  arsenic  compounds,  and  then  gives  rise  to  symptoms 
of  poisoning.  It  should  be  carefully  distinguised  from 
the  trinitrate.  It  is  also  used  as  a  cosmetic,  and  to  give 
an  iridescent  glaze  to  porcelain. 

299.  Bismuth  Trioxide,  Bi2O3,  is  prepared  by  boil- 
ing subnitrate  of  bismuth  with  solution  of  caustic  soda : 

2BiNOs(OH)2  +  2NaOH  =  2NaNO3  +  Bi2O3  +  3H2O. 
The  precipitate  is  collected  on  a  filter,  washed,  and  dried. 
PROPERTIES. — A  lemon-yellow  powder,  insoluble  in 
water,  but  soluble  in  nitric  and  hydrochloric  acids,  with 
the  formation  of  the  trinitrate  and  the  trichloride  (BiCl3) 
respectively.  It.  is  a  basic  oxide. — It  is  used  instead 
of  the  subnitrate  in  many  cases,  and  is  to  be  preferred 
on  account  of  its  purity. 

300.  Bismuthyl  Carbonate.— 2(BiO)2CO3 .  H2O. 


BISMUTHYL    CARBONATE.  303 

PREPARATION.— Experiment  259- — Pour  solution  of  bis- 
muth trinitrate  into  cold  solution  of  ammonic  carbonate.  A 
white  precipitate  falls  : 

3(NH4)aC08  +  2Bi(NO,),  = 

(BiO)2C03  +  6NH4NOS  +  2C02. 

Collect  the  precipitate  on  a  filter  and  wash  it. 

PROPERTIES. — This  salt,  commonly  called  carbonate  of 
bismuth,  is  a  white  powder,  insoluble  in  water,  but  soluble 
in  nitric  acid  with  effervescence.  It  is  also  soluble  in 
sulphuric  acid,  and  the  solution  should  not  respond  to 
the  test  for  nitric  acid.  If  it  does,  the  carbonate  has 
contained  subnitrate.  The  carbonate  of  bismuth  is  often 
administered  in  place  of  the  subnitrate,  on  account  of  its 
more  ready  solubility  in  the  acid  juices  of  the  stomach, 
and  also  because  of  its  antacid  properties  : 

(BiO)2C03  +  6HC1  =  2BiCl3  -f  C02  +  3H2O. 

In  this  salt  the  radical  -  Bi  =  O  plays  the  part  of  a 
monad  metal.  Other  salts  of  bismuthyl  are  known,  e.g., 
bismuthyl  chloride  (BiOCl),  formed  when  the  trichloride 
of  bismuth  is  acted  on  by  much  water : 

BiCl3  +  H2O  =  BiOCl  +  2HC1. 

301.  Tests. 

1.  To  the  filtrate  from  Experiment  256  add  hydric  sulphide. 
A  browrf-black  precipitate  of  bismuth  trisulphide  (Bi2Ss)  falls  : 

2Bi(N03)3  +  3H2S  =  Bi2S3  +  6HNO3. 

« 
It  is  insoluble  in  yellow  ammonic  sulphide,  but  soluble  in  hot 

dilute  nitric  acid. 

2.  Solutions  of  bismuth  salts  give  with  ammonia  a  white  pre- 
cipitate (Bi(OH)3),  not  dissolved  by  more^f  the  reagent.     If  this 
precipitate  be  filtered  off  and  dissolved  in  as  little  as  possible  of 


304  ANTIMONY. 

hydrochloric  acid,  the  solution  is  turned  milky  on  the  addition 
of  much  water.  This  precipitate  is  not  dissolved  by  tartaric 
acid. 

3.  Potassic  iodide  gives  with  bismuth  solutions  a  brown  pre- 
cipitate (Bils). 

4.  Caustic  potash  gives  a  white  precipitate  (Bi(OH)3),  insoluble 
in  excess. 

5.  Insoluble  bismuth  compounds  can  be  tested  by  dissolving 
them  in  nitric  acid  and  diluting  with  water.     (Experiment  256.) 
The  precipitate  obtained  is  insoluble  in  tartaric  acid.     (Com- 
pare Antimony.) 


B. 

ANTIMONY  (Stibium). 

302.  Antimony  (Sbui-v-  =  120.  Sp.  wt.  =  6.7  to 
6.86.  Melting  point  =  425°C.  Sp.  heat  =  0.0523).— 
The  chief  ore  of  antimony  is  stibnite  (Sb2S3).  This 
occurs  in  black,  shining,  crystalline  masses. 

PREPARATION. — Stibnite,  purified  by  fusion,  is  reduced 
by  heating  with  iron  : 

Sb2S3  +  3Fe  =  2Sb  +  3FeS. 

PROPERTIES. — A.  silvery  metal,  generally  in  masses  of 
laminated  crystals,  hard  and  brittle ;  it  can  be  ground  to 
a  powder  in  the  mortar.  When  heated  in  air,  it  burns, 
forming  the  tetroxide  (Sb204).  It  is  soluble  in  nitric 
acid  and  in  aqua  regia.  It  is  also  soluble  in  hot,  strong 
sulphuric  acid,  antimonic  sulphate  being  formed  : 

2Sb  +  6HaSO4  -  Sb2(SO4)3  +  3S02  +  6H2O. 

In  this  salt  antimony  plays  the  part  of  a  trivalent  metal. — 
Antimony  black  is  finely  divided  antimony  prepared  by 
reducing  the  metal  from  a  solution  of  the  chloride  by 


ANTIMONY    TRISULPHIDE.  305 

means  of  zinc.  It  is  used  for  giving  to  plaster  casts,  &c., 
the  appearance  of  iron  or  steel. — Antimony  is  a  con- 
stituent of  many  use.ful  alloys.  (Art.  240.) 

303.  Compounds    of    Antimony.  —  Antimony 

combines  with  oxygen  in  three  proportions.  The  trioxide 
(Sb2O3)  is  both  basic  and  acid-forming ;  the  pentoxide 
(Sb,O5)  is  acid-forming,  antimonic  acid  (H3SbO4)  and  the 
antimonates  being  similar  to  the  corresponding  compounds 
of  phosphorus  and  arsenic ;  the  tetroxide  vSb2O4)  is  also 
acid-forming. 

304.  Antimony  Trisulphide,  Sb,S3,  is  found  in 

nature  as  sfibnite.  The  mineral  is  purified  by  fusion, 
and  is  used  as  the  starting  point  in  ther  preparation  of 
antimony  compounds.  It  is  also  called  black  antimony 
and  crude  antimony. — It  is  a  greyish-black  crystalline 
powder,  soluble  in  hot,  strong  hydrochloric  acid,  with 
evolution  of  hydric  sulphide. 

Experiment  260 — Heat  a  little  black  sulphide  of  antimony 
with  strong  hydrochloric  acid.     Keep  the  solution  : 

Sb2S3  +  6HC1  =  2SbCl3  +  3H2S. 

Experiment  261-  —  Boil  a  little  sulphide  of  antimony  with 
solution  of  sodic  hydroxide.     It  dissolves  : 

2Sb2Ss  +  4NaOH  =  NaSbO2  +  3NaSbS2  +  2H2O. 
Add  dilute  sulphuric  acid  until  a  precipitate  appears  : 
NaSb02  +  3NaSbS3  +  2H2SO4  = 

2Na2SO4  +  2Sb2S3  +  2H2O. 

The  substances  formed  when  antimony  trisulphide  is  dissolved 
in  caustic  soda  are  sodic  antimonite  (NaSbO.J  and  sodic  sulphanti- 
monite  (XaSbS2).  When  the  acid  is  added  the  trisulphide  is 
again  precipitated,  but  it  is  orange-red.  Prepared  in  this  way 
it  always  contains  a  little  of  the  trioxide,  and  is  called  sulphurated 
antimony,  or  golden  sulphide  of  antimony. 
21 


306  POWDER    OF    ALGAROTH. 

305.  Antimony  Trichloride  (SbCl3)  is  prepared 

as  in  Experiment  260.     By  evaporating  such  a  solution 

the   trichloride  is  obtained   as   a  crystalline,  colourless, 

solid,  melting  at  72°  ("  butter  of  antimony  ").  It  is  very 
deliquescent. 

Experiment  262. — Dilute  a  small  part  of  the  solution  pre- 
pared in  Experiment  260.      A  white  precipitate  of  antimony  I 
chloride,  or  powder  of  Algaroth  (SbOCl)  is  formed  : 
Sb013  +  H2O  =  SbOCl  +  2HC1. 

Filter,  and  test  the  filtrate  for  antimony  and  for  hydrochloric 
acid.  To  two  other  portions  add  considerable  quantities  of 
strong  hydrochloric  acid  and  tartaric  acid  respectively,  and  dilute 
them  as  before.  To  a  strong  solution  of  antimony  trichloride 
add  a  little  dilute  hydrochloric  acid.  A  precipitate  forms  ;  the 
dilute  acid  has  the  same  effect  as  water.  Add  more  acid  and 
the  precipitate  is  redissolved. 

Antimony  trichloride  is  a  powerful  caustic.  It  is  a 
corrosive  poison,  acting  like  a  strong  solution  of  hydro- 
chloric acid.  (Explain  this.) 

306.  Antimony   Trioxide    (Sb2O3),    also    called 

flowers  of  antimony,  is  prepared  by  digesting  powder  of 
Algaroth  (SbOCl)  with  sodic  carbonate,  and  washing 
with  hot  water. — It  is  a  greyish-white  powder,  insoluble 
in  water,  sulphuric  and  nitric  acids,  but  readily  soluble 
in  alkalis,  and  in  hydrochloric  and  tartaric  acids  : 

Sb2O3  +  6HC1  =  2SbCl3  +  3H2O. 

When  heated  strongly  in  air  it  fuses  and,  absorbing 
oxygen,  becomes  changed  to  the  tetroxide  (Sb2O4). — 
Antimonial  powder  is  a  mixture  of  1  part  of  antimony 
trioxide  with  2  parts  of  calcic  phosphate  (Ca3(PO4)2). 

307.  Tartar  Emetic,  SbO.K.C4H4Oe,  is  antimonyl 
potassium  tartrate,  already  mentioned  (Art.  189). 


TARTAR    EMETIC.  307 

Experiment  263-  — Dissolve  some  tartar  emetic  in  as  little  as 
possible  of  hot  water.  Dilute  the  solution.  It  does  not  turn 
milky.  The  explanation  is  as  follows  : — The  antimony  is  already 
combined  as  a  soluble  antimoiiyl  compound.  (What  experiment 
above  does  this  explain  ?) 

It  will  have  been  observed  that  antimony  shares  with 
bismuth  the  tendency  to  form  so-called  oxysalts,  in  which 
a  radical,  in  this  case,  antimonyl,  -  Sb  =  O,  acts  the 
part  of  a  univalent  atom. 

Experiment  264- — Examine  a  specimen  of  tartar-emetic  care- 
fully. Taste  it.  Heat  a  small  portion  on  mica  or  platinum. 
It  chars  and  burns,  leaving  a  white  solid. 

In  large  doses  tartar  emetic  is  a  poison.  Antidotes, 
freshly  precipitated  ferric  hydroxide,  tannic  acid,  or  any 
vegetable  infusion  containing  tannin,  e.g.,  tea. 

308.  Tests. 

1.  To  a  solution  of  an  antimony  salt,  acidified  with  hydrochloric 
acid,  add  hydr'u:  sulphide.     An  orange  precipitate  falls.     Filter 
off  the  precipitate  and  teat  the  solubility  of  parts  of  it  in  yellow 
ammonic   sulphide  and   in  hot  strong   hydrochloric   acid.      It 
dissolves  in  both.      Dilute  the  hydrochloric  acid  solution. — The 
ammonic   sulphide  solution   contains   ammonic    sulphantimonite 
(NH4SbS2).     Treat  it  with  hydrochloric  acid,  and  the  sulphide 
is  reprecipitated. 

2.  Test   for   antimony  as  in   Marsh's  test   for   arsenic  (Art. 
146).      Antimony   has    a    gaseous    compound    with    hydrogen, 
stibinf,  or  antimoniuretted  hydrogen  (SbHs).     Spots  are  obtained, 
similar  to  those  of  arsenic,  but  they  are  turned  orange  by  am- 
monic  sulphide,    and    are   not  dissolved   by  bleaching   powder 
solution.     (Compare  Arsenic. ) 

3.  Put  a  scrap  of  zinc  in  a  solution  of  tartar  emetic,  collect 
the  precipitated  antimony,  and  test  its  solubility  in  hot  hydro- 
chloric acid.     It  is  insoluble.     (Compare  Tin.) 


308  TIN. 

4.  Strong  solutions  of  antimony  trichloride  give  a  white  pre- 
cipitate with  water  or  dilute  hydrochloric  acid,  soluble  in  tar- 
taric  acid  ;   but  this  test  does  not  answer  with  tartar  emetic  or 
alkaline  solutions. 

5.  Insoluble  antimony  compounds  can  be  dissolved  in  hydro- 
chloric or  nitric  acid,  and  the  solution  then  treated  as  above. 


TIN  (Stannum). 

309.  Tin  (Snliiv  ==  117.8.     Specific  weight  =  7.739. 
Melting  point  =  235°.     Specific  heat  =  0.0548).— Tin 
is  prepared  almost  exclusively  from  tin-stone  (SnO2),  by 
smelting  the  purified  ore  in  a  blast  furnace  with  anthra- 
cite : 

SnO2  +  20  =  Sii  +  2CO. 

The  impure  metal  is  purified  by  liquation,  i.e.,  by  melt- 
ing gradually.  The  pure  metal  melts  first  and  flows 
away  from  its  impurities.  Commercial  tin  may  contain 
arsenic,  antimony,  bismuth,  zinc,  lead,  copper,  and  iron. 
PROPERTIES. — A  bright  white  metal,  crackling  when 
bent,  harder  than  lead,  softer  than  gold.  It  is  malleable 
and  ductile  at  100°  C.  It  does  not  tarnish  readily  in 
air,  and  is  therefore  used  for  covering  sheet  iron  in  the 
manufacture  of  "  tin  "  utensils.  It  is  soluble  in  hydro- 
chloric and  dilute  nitric  acids,  and  is  oxidised,  but  not 
dissolved,  by  strong  nitric  acid. — Tin  forms  some  useful 
alloys.  (See  Alloys.)  Tin  amalgam  is  used  for  silvering 
mirrors. 

Experiment  265. — Put  a  bit  of  zinc  in  an  alkaline  solution 
of  tin  (SnCl2  and  NaOH).  Tin  is  gradually  deposited  in  crystals. 

310.  Compounds  Of  Tin. — Tin   forms  two   series 
of  compounds:    (1)  Stanruws  (SnO,  SnS,  SnCl,,  &c.)  in 


STANNIC   OXIDE.  309 

which  the  element  is  dyad  ;  and  (2)  stannic  (SnO2,  SnS2, 
SnCl4,  &c.),  in  which  it  is  tetrad.  Stannous  oxide  (SnO) 
is  basic,  forming  salts  with  acids,  e.g.,  Sn(NO3)2,  SnSO4, 
&c.  Stannic  oxide  (SnO2)  is  acid-forming  (also  weakly 
basic),  and  the  stannates  are  analogous  to  the  carbonates 
and  the  silicates,  e.g.,  sodic  stannate,  Na2SnO3. 

311.  Stannic  Oxide  (SnO2)  is  found  in  nature  as 
tin-stone. 

Experiment  266. — Pour  some  strong  nitric  acid  on  a  few 
scraps  of  tin.  Red  fumes  are  evolved,  and  a  white  powder  is 
formed.  This  is  stannic  acid  (H2Sn03)  : 

Sn  +  4HNO3  =  H2SnO3  +  4NO2  +  H2O. 

Try  its  solubility  in  hydrochloric  acid,  and  in  caustic  soda.  It 
dissolves  in  both,  forming  stannic  chloride  (SnCl4)  and  sodic 
stannate  (Na2Sn03)  respectively. — Heat  a  little  stannic  acid  on 
mica.  A  white  powder  remains  (yellow  when  hot).  This  is 
stannic  oxide  (Sn02).  Try  its  solubility  in  hydrochloric  acid 
and  in  caustic  soda. 

312.  Stannous  Chloride  (SnCl2) 

Experiment  267. — Dissolve  some  scraps  of  tin  in  hydro- 
chloric acid  diluted  with  an  equal  volume  of  water.  Stannous 
chloride  is  formed  : 

Sn  +  2HC1  =  SnCl2  +  H2. 

Put  in  more  tin  and  evaporate  on  the  water  bath  to  crystallisa- 
tion. "Tin  Salt"  (SnCl2.2H20)  is  obtained — Dissolve  a  little 
of  this  salt  in  a  small  quantity  of  water.  It  forms  a  clear  solu- 
tion. Add  more  water  ;  it  becomes  turbid,  owing  to  the  for- 
mation of  a  basic  chloride  : 

SnCl2  +  H2O  =  SnCLOH  +  HC1. 

The  same  precipitate  is  formed  when  a  solution  of  stannous 
chloride  is  exposed  to  the  air : 

3SnCl2  +  O  +  H2O  =  SnCl4  +  2SnClOH. 


310  STANNIC    CHLORIDE. 

Stannous  salts  have  an  astringent  metallic  taste.  They 
are  easily  oxidised,  and  must  be  kept  from  the  air. 
They  are  powerful  reducing  agents,  precipitating  gold, 
silver,  and  mercury  from  their  solutions. — The  nitrate 
(Sn(NO3)2)  and  sulphate  (SnS04)  can  be  prepared  by 
dissolving  tin  in  the  dilute  acids.  Solutions  of  stannous 
salts  are  acid  in  reaction.  They  are  poisonous  ;  anti- 
dote, solution  of  ammonic  carbonate. 

313.  Stannic  Chloride  (SnCl4)  can  be  prepared  by 
passing  dry  chlorine  gas  over  tin  foil,  or  by  distilling 
tin  with  mercuric  chloride  (Hg012).  It  can  also  be  pre- 
pared in  solution  as  follows  : 

Experiment  268. — Boil  solution  of  stannous  chloride  with 
nitric  and  hydrochloric  acids,  using  only  a  small  quantity  of  the 
substances.  Keep  the  solution. 

PROPERTIES. — A  colourless,  heavy,  fuming  liquid.  It 
solidifies  when  mixed  with  one-tlm-d  its  weight  of  water, 
forming  "butter  of  tin"  (SnCl4.5H2O).  With  much 
water  it  is  decomposed  : 

SnCl4  +  3H2O  =  H2SnO3  +  4HC1. 


Many  compounds  of  tin  are  similar  in  composition  to 
compounds  of  silicon,  e.g.,  potassic  fluo-stannate  (K2SnFs), 
and  the  stannates  (Na2SnO3,  &c.). 

Closely  allied  to  tin  are  three  rare  metals,  titanium, 
zirconium,  and  thorium. 

314.  Tests. 

Stannic  Salts. 

1.  To  a  solution  of  stannic  chloride  add  hydric  sulphide.  A 
yellow  precipitate  (SnS2)  falls  : 

SnCl4  +  2H2S  =  SnS2  -f-  4HC1. 


TESTS    FOR   TIN.  311 

This  precipitate  is  soluble  in  ammonic  sulphide  and  in  strong,  hot, 
hydrochloric  acid.  (Filter  it  off  and  try. ) — Stannic  sulphide^nSy) 
combines  with  alkaline  sulphides  to  form  sulphostannates,  e.g. : 

SnS2  +  (NH4)2S  =  (NH4)2SnS3. 

These  are  soluble  salts,  easily  decomposed  by  acids.  (Try  with 
hydrochloric  acid.) 

2.  Caustic  soda  gives  with  stannic  salts  a  white  precipitate 
(H2Sn03),  soluble  in  excess  of  the  reagent. 

3.  To  a  little  stannic  chloride  solution  in  a  porcelain  dish  add 
a  scrap  of  zinc  and  warm.     Wash  the  precipitated  tin,  dissolve 
it  in  warm  hydrochloric  acid,  and  add  to  the  solution  a  drop  of 
mercuric  chloride.     A  white  precipitate  of  calomel  (Hg.,Cl2)  or 
a  grey  precipitate  of  mercury  is  formed  : 

2HgCl2  +  SnCl2  =  Hg2Cl2  +  SnCl4. 
HgCl2  +  SnCl2  =  Hg  +  SnCl4. 
This  test  applies  also  to  stannous  salts. 
Stannous  Salts. 

1.  To  solution  of  stannous  chloride  add  hydric  sulphide.     A 
brown  precipitate  of  stannous  sulphide  (SnS)  is  formed.     Collect 
on  a  filter  or  wash  by  decantation,  and  test  its  solubility  in 
yellow  ammonic  sulphide.     It  dissolves.     To  the  solution  add 
hydrochloric  acid.     Stannic  sulphide  is  precipitated.—  Explana- 
tion :      Yellow  ammonic   sulphide    contains   a    persulphide    of 
ammonium  (NH4)2S2  ;  this  unites  with  stannous   sulphide  to 
form  ammonic  sulphostannatc : 

(NH4)2S2  +  SnS  =  (NH4)aSnS3, 
which  is  decomposed  by  hydrochloric  acid  as  follows  : 

(NH4)2SnS3  +  2HC1  =  2NH4C1  +  SnS2  +  H3S. 

2.  Caustic  soda  gives  with  stannous  solutions  a  white  pre- 
cipitate (Sn(OH)2)  soluble  in  excess. 

3.  Same  as  for  stannic  salts  ;  but  stannous  salts  give  a  white 
(Hg2Cla)  or  grey  (Hg)  precipitate  with  mercuric  chloride,  with- 
out the  preliminary  treatment  with  zinc,  &c. 

Insoluble  tin  compounds  can  be  got  into  solution  by 
reducing  on  charcoal  and  dissolving  in  hydrochloric  acid. 


312  GOLD. 

GOLD  (Aurum). 

315.  Gold.     (AuLUi-  =  196.2.— Sp.  wt.  =  19.265.— 
Melting  point  =  1037°.— Sp.  heat  =  0.03244.) 

OCCURRENCE. — Gold  is  usually  found  free  or  alloyed 
with  silver,  platinum,  <fcc.  It  is  found  associated  with 
quartz  and  pyrites  mostly. 

PREPARATION. — It  is  obtained  from  quartz  by  crushing, 
and  washing  away  the  lighter  mineral.  It  is  obtained 
from  auriferous  sand  and  gravel  by  simple  washing. 
Gold  is  extracted  from  pyrites,  &c.,  by  treating  with 
aqua  regia,  and  precipitating  the  gold  with  green  vitriol. 

PROPERTIES. — Gold  is  of  familiar  appearance,  and  need 
not  be  described.  It  is  softer  than  silver,  and  must  be 
hardened  by  alloying  with  copper  before  it  is  suitable  for 
use.  It  is  the  most  malleable  and  ductile  of  all  metals. 
A  gold  wire  can  be  drawn  so  fine  that  10,000  feet  (about 
1|  mile)  weigh  only  15  grains.  Gold  is  not  attacked  by 
any  simple  acid,  excepting  selenic  (H2SeO4).  It  is  attacked 
by  caustic  potash,  caustic  soda,  and  saltpetre.  The  best 
solvent  for  gold  is  aqua  regia,  which  dissolves  it  as  tri- 
chloride (AuCl3).  Gold  is  very  easily  precipitated  from 
solutions  by  reducing  agents  such  as  stannous  chloride, 
ferrous  sulphate,  oxalic  acid,  inercurous  nitrate,  &c. 
Purple  of  Cassius  is  formed  by  precipitating  with  stan- 
nous chloride.  It  consists  of  finely  divided  gold  and 
some  compound  of  tin. 

316.  Compounds    Of    Gold.       Gold    forms    two 
oxides,  aurous  (Au2O),  and   auric  (Au2O3).     Both   are 
decomposed  into  gold  and  oxygen  at  250°C.     Both  are 
basic,  but  auric  oxide  is  also  a  weak  acid-forming  oxide. — 
The  aurous  salts  (AuCl,  Aul,  &c.)  are   very  unstable, 


PLATINUM.  313 

readily  decomposing  into  gold  and  auric  salts.  Potassic 
aurous  cyanide  (KCN.AuCN),  is,  however,  quite  stable, 
and  is  used  in  electro-gilding. — Auric  chloride  (Au013) 
unites  with  hydrochloric  acid  to  form  chlorauric  acid 
(HAuCl4).  This  is  the  solution  generally  used  and 
called  terchloride  of  gold. 

317.  Tests. 

1.  Solutions  of  gold  acidified  with  hydrochloric  acid  give  a 
brown  precipitate  (Au2S.,),  soluble  in  ammonic  sulphide. 

1'.  Solutions  containing  gold  give  a  purple  colour  or  precipitate 
with  stannous  chloride  or  ferrous  sulphate. 


PLATINUM,  &c. 

318.  Platinum  (Ptiiiv  =  196.7.— Sp.  wt.  =  21.5.— 
Melting  point  about  2600°.— Sp.  heat  =  0.03243). 

OCCURRENCE. — Platinum  is  found  in  the  metallic  state, 
alloyed  with  palladium,  osmium,  iridium,  <fec.  It  is 
found  in  the  Ural  Mountains,  S.  America,  Australia, 
Borneo,  and  California.  It  has  lately  been  discovered  in 
British  Columbia.  It  is  generally  present  in  small 
quantity  in  gold  and  silver. 

PREPARATION. — The  crude  platinum  is  dissolved  in 
aqua  regia,  and  precipitated  as  ammonium  chloroplati- 
nate,  (NH4^2  PtCl6,  by  ammonic  chloride.  From  this 
compound  it  is  obtained  by  heating  : 

(NH4)2PtCl6  =  2NH4C1  +  Pt  +  2C12. 
It  generally  contains  about  2  %  of  iridium. 

PROPERTIES. — A  tin-white  metal,  soft,  very  heavy  and 
malleable.  It  fuses  only  at  an  intense  white  heat.  It 
has  the  power  of  condensing  gases  on  its  surface,  and,  in 


314  PLATINUM. 

the  finely  divided  state  of  spongy  platinum  and  platinum 
black,  it  is  used  to  bring  about  the  oxidation  of  alcohols 
to  aldehydes,  sulphur  dioxide  to  trioxide,  <fec.  Platinum 
is  attacked  by  few  chemical  substances.  It  combines, 
however,  with  the  halogens,  and  is  dissolved  slowly  when 
heated  strongly  with  caustic  potash,  caustic  soda,  potassic 
nitrate,  or  potassic  cyanide.  It  is  attacked  by  oxides 
and  sulphides  of  easily  reducible  metals,  such  as  lead, 
copper,  bismuth,  &c.,  as  well  as  by  these  metals  them- 
selves. These  substances  should  never  be  heated  in  a 
platinum  crucible. 

USES. — Platinum  is  of  the  greatest  importance  in 
chemistry.  "  Without  platinum  the  composition  of  most 
miner-als  would  have  yet  remained  unknown." — (Liebig.) 
It  is  used  for  crucibles,  as  it  resists  the  action  of  most 
chemicals.  It  is  also  used  in  the  manufactui'e  of  chemical 
balances  and  many  other  instruments  of  precision.  The 
surgeon  employs  it,  heated  by  a  current  of  electricity  or 
by  alcohol  vapour,  using  it  instead  of  a  knife. 

319.  Compounds  Of  Platinum.  Platinum  forms 
two  series  of  compounds,  platinous  (PtO,  PtCl2,  PtS,  <fec.), 
and  platinic  (PtO2,  PtCl4,  &c.).  The  platinous  compounds 
are  unimportant.  When  platinum  is  dissolved  in  aqua 
regia  the  solution  contains  chloroplatinic  acid  (H2PtCl6). 
This  is  the  solution  generally  called  "  solution  of  tetra- 
chloride  of  platinum."  It  is  a  well-marked  dibasic  acid, 
and  forms  characteristic  sparingly  soluble  salts  (chloro- 
platinates)  of  potassium,  ammonium,  amines,  and  alka- 
loids. These  salts  are  of  a  golden  yellow  colour,  and  are 
sparingly  soluble  in  water. — As  the  sodium  salt(Na2PtC!6) 
is  very  soluble  in  water,  chloroplatinic  acid  is  used  in 
analysis  to  separate  sodium  from  potassium. 


PALLADIUM,    <fec.  315 

320.  Tests. 

It  is  precipitated  along  with  the  other  members  of  this  group 
by  hydric  sulphide.  Platinic  sulphide  (PtS2 )  dissolves  in  ammonic 
sulphide. — A  solution  containing  platinic  salts  gives  a  golden 
yellow  precipitate  with  ammonic  chloride. — Solid  substances 
containing  platinum  are  extracted  with  aqua  regia  and  tested  as 
above . 

321.  Palladium  is  similar  to  platinum,  but  is  soluble 
in  nitric  acid.     It  is  used  instead  of  gold  by  dentists. — 
Indium  is   alloyed    with    platinum    to    make    standard 
weights  and  measures.     The  alloy  is  very  hard,  and  as 
elastic  as  steel. — Osmium  forms  a  remarkable  acid-forming 
oxide    (O&O4),    which    is   volatile   and    very    poisonous. 
Osmic  acid  (H2OsO6)  is  much  used  in  practical  histology. 
It  stains  fats  black.     An  alloy  of  osmium  and  iridium 
(osmiridium)  is  used  for  tipping  gold  pens.     Osmium  is 
the  heaviest  substance   known   (sp.   wt.  =  22.477).     It 
has    never    been    fused. — Tungsten    and    molybdenum, 
although  not  nearly  allied  to  the  metals  of  this  group, 
are  mentioned  here  because  their  sulphides  are  precipitated 
by  hydric  sulphide  and  dissolve  in  ammonic  sulphide. 


QUESTIONS  AND  EXERCISES. 

1.  "The  hydroxides  of  Group  II.  are  precipitated  from  solu- 
tions of  salts  by  alkaline  hydroxides."  Illustrate  this  statement 
by  examples. 

Illustrate  from  the  members  of  this  group  the  Law  of  Dulong 
and  Petit. 

3.  What  experimental  proof  can  you  bring  to  show  that  iron 
has  a  stronger  attraction  for  salt  radicals  than  copper  has,  i.e.,  is 
a  more  positive  radical  ? 

4.  How  would  you  prepare  crystals  of  cupric  nitrate  ? 


316  QUESTIONS    AND    EXERCISES. 

5.  Why  is  it  dangerous  to  eat  food  which  has  been  for  some 
time  in  contact  with  brass  ? 

6.  Is  there  any  danger  in  using  brass  taps  for  vinegar   and 
cider  casks  ?     Explain. 

7.  How  can  cadinic  chloride  be  prepared  from  cadmic  nitrate  ? 
(Both  salts  are  soluble.) 

8.  The  specific  weight  of  cadmium  vapour  is  3.94  (air  =  1). 
Calculate  its  molecular  weight.     How  many  atoms  in  the  mole- 
cule of  cadmium  ? 

9.  Why   should   we   expect   bismuth   carbonate  to  be   more 
soluble  in  the  gastric  juice  than  the  subnitrate  ? 

10.  Compare  bismuth  and  antimony  (1)  as  to  the  properties  of 
the  elements  themselves,  (2)  as  to  their  compounds. 

11.  Bismuth  trinitrate   is   poisonous,   and   large  doses  cause 
symptoms  of  nitric  acid  poisoning.     Explain  this. 

12.  How  would  you  distinguish  by  a  test  bismuth  subnitrate 
from  bismuthyl  carbonate  ? 

13.  Compare  antimony  and  arsenic  with  regard  to  their  com- 
pounds. 

14.  Balance  the  following  equations  : 

(1)  SbCl2OH  +  H2O  =  Sb2O3  +  HC1. 

(2)  Sba03  +  KHC4H4O6  =  SbOKC4H4Oa  +  H2O. 

(3)  SbCl3  +  H2S  =  Sb2S3  +  HC1. 

(4)  Sb2O3  +  H2  =  SbH3  +  H2O. 

15.  What  substances  are  formed  when  tin  is  dissolved  in  very 
dilute  nitric  acid  ? 

16.  When  solution  of  stannous   chloride  (SnCl2)  is  added  to 
auric  chloride  (AuCl3),  metallic  gold  is  precipitated  and  stannic 
chloride  remains  in  solution.     Write  the  equation. 

17.  What   is   the   chemical  composition   of    "butter  of  tin," 
"  butter  of  antimony,"  "  butter  of  arsenic,"  and  "  tin  salt  ?  " 

18.  Why  is  it  necessary  to  use  ammonic  sulphide  containing 
excess  of  sulphur  to  dissolve  stannous  sulphide  ? 

19.  A  solution  of  chlorauric  acid  in  water  oxidises  ferrous  sul- 
phate.    How  is  this  possible  ?     Does  chlorauric  acid  contain  any 
oxygen? 


METALS    OP    GROUP    III.  317 


CHAPTER     XIX. 


METALS  OF  GROUP  III. 

Iron,  Chromium,  Aluminium  ;  Zinc,  Manganese,  Cobalt, 
and  Nickel.     [Rare  metals  of  the  Cerium  class  J\ 

322.  General  Characters. — The  metals  of  this 

group  ai%e  mostly  reducible  from  their  oxides  by  smelt- 
ing with  charcoal,  but  are  more  difficult  to  reduce  than 
those  of  the  preceding  group.  With  two  exceptions 
(aluminium  and  zinc)  they  have  two  oxides  having  the  gen- 
eral formulas  MO  and  M2O3.  Besides  these,  chromium 
and  manganese  have  well-mai'ked  acid  forming  oxides. 
Their  sulphides  are  not  precipitated  by  hydric  sulphide 
from  acid  solutions  ;  but  are  precipitated  by  alkaline  sul- 
phides (two  exceptions  will  be  noted  further  on).  Both 
sulphides  and  oxides  are  insoluble  in  water,  but  soluble  in 
dilute  acids  (nickel  and  cobalt  sulphides  with  difficulty). 
The  hydroxides,  carbonates,  and  phosphates  are  insoluble 
in  water,  and  can  all  be  prepared  by  precipitation.  The 
hydroxides  are  easily  changed  to  oxides  by  heat.  The 
sulphates,  chlorides,  nitrates,  and  acetates  are  soluble  in 
water,  and  their  solutions  have  an  acid  reaction.  The 
sulphates  form  characteristic  double-sulphates  with  those 
of  the  alkali  metals,  e.g.,  KjSO^FeSO^CHjO.  The 
alums  form  an  interesting  group  of  these  double  sulphates. 
For  convenience  of  analysis  this  group  is  subdivided 
into  :  A.  Metals  forming  stable  hydroxides,  M2(OH)6,  in- 
soluble in  ammonia :  iron,  chromium,  and  aluminium ; 


318  IRON. 

and  B.  Metals,  the  most  basic  hydroxides  of  which  have 
the  general  formula  M(OH)2  and  are  soluble  in  am- 
monia :  zinc,  manganese,  cobalt,  and  nickel.  In  analysis 
the  cerium  metals,  as  well  as  uranium,  zirconium,  and 
thorium  fall  here. 


A. 
IRON  (Ferrum). 

323.  Iron  (Feliiv  =  56.— Specific  weight  =  7.844. 
Melting  point  =  1600°.  Specific  heat  =  0.11379). 

OCCURRENCE. — Rarely  free,  and  then  generally  of 
meteoric  origin.  Meteors  composed  mostly  of  iron  have 
been  found  weighing  several  tons.  The  principal  ores 
of  iron  are  red  hcematite  (Fe2O3),  brown  haematite 
(2Fe2O3.3H2O),  magnetic  iron  ore  (Fe304),  spathic  iron 
ore  (FeCOj),  clay  ironstone  (FeCO3,  with  clay  or  sand), 
and  black  band  (FeCO3,  with  coal). — Iron  is  present  in 
the  sun  and  in  many  fixed  stars.  —  Iron  is  an  essential 
constituent  of  the  bodies  of  plants  and  animals.  In  the 
latter  it  is  chiefly  found  in  the  haemoglobin  of  the  blood. 

SMELTING  OF  IRON.  —  The  ores  are  first  calcined,  if 
necessary,  to  drive  off  water,  &c.,  and  then  placed  in  a 
tall,  somewhat  spindle-shaped,  furnace  (blast  furnace*}, 
with  alternate  layers  of  coal  and  limestone.  The  lime- 
stone combines  with  the  siliceous  impurities  to  form  a 
fusible  slag.  The  coal  burning  at  the  bottom  of  the 
furnace  (where  a  blast  of  hot  air  feeds  the  combus- 
tion) forms  carbon  dioxide,  which,  passing  upward, 
unites  with  carbon  to  form  carbon  monoxide.  Ferric 
oxide  (Fe2O3)  is  then  reduced  by  the  carbon  monoxide  : 

FeO    +  3CO  =  2Fe  +  3CO. 


IRON.  319 

This  is  repeated  with  the  successive  layers.  The  iron,  as 
it  falls  toward  the  bottom  of  the  furnace,  combines  with 
carbon  and  silicon  and  becomes  more  fusible.  It  melts 
and  falls  into  the  hearth  of  the  furnace,  the  slag  forming 
a  layer  above  it.  The  molten  iron  is  run  off  into  chan- 
nels made  of  sand,  and  solidifies  into  bars  of  pig  iron. 
Pig  iron  contains  from  76  %  to  96  %  of  pure  iron,  from 
1  %  to  20  %  of  manganese,  and  from  1  %  to  7  %  of  car- 
bon. It  melts  with  comparative  ease,  and  is  used  in 
manufacturing  stoves,  &c.,  by  the  process  of  casting.  It 
is  hence  called  cast  iron.  It  is  more  brittle  than  pure 
iron  and  lighter. —  Wrouyht  iron  is  prepared  from  cast 
iron  by  removing  the  carbon  and  silicon.  This  is  done 
by  subjecting  the  molten  metal  to  a  hot  oxidising  blast 
in  the  process  of  puddling.  In  the  Bessemer  process  the 
iron  is  kept  hot  by  the  oxidation  of  its  impurities. 
Wrought  iron  contains  up  to  0.3  %  of  carbon.  It  is 
heavier  than  cast  iron  and  very  tenacious  and  malleable. 
— Steel  is  intermediate  in  composition  between  cast  and 
wrought  iron.  It  is  now  generally  prepared  by  the 
Bessemer  process,  modified  more  or  less.  The  impurities 
are  burned  away  until  the  composition  is  that  of  wrought 
iron.  Then  enough  cast  iron  is  added  to  bring  up  the 
percentage  of  carbon  to  about  1.5  %. — Sponyy  iron,  the 
ferrum  redactum  of  the  Pharmacopoeia,  is  prepared  by 
heating  pure  ferric  oxide  in  a  current  of  hydrogen  : 

Fe2O3  +  3H2  -  2Fe  +  3H20. 

It  is  valuable  as  a  medicine,  and  as  a  filter  for  water, 
since  it  possesses  the  power  of  destroying  impurities. 

PROPERTIES. — Pure  iron  is  almost  as  white  as  silver. 
It  is  the  most  tenacious  of  all  metals,  except  nickel  and 


320  OXIDES    OF    IRON. 

cobalt.  Iron  is  soft  at  a  red  heat,  and  can  be  welded 
at  a  white  heat.  Borax,  sand,  &c.,  are  used  to  clear 
away  the  oxide  from  the  surfaces  to  be  welded.  This 
they  do  by  uniting  with  it  to  form  fusible  slags  (silicate 
and  borate  of  iron),  which  are  easily  scraped  off.  Iron 
is  attracted  by  magnets,  and  can  be  made  magnetic  by 
the  influence  of  electricity  or  of  other  magnets.  It  does 
not  rust  in  dry,  but  does  in  moist,  air.  The  rust  is  a 
compound  of  ferric  oxide  with  ferric  hydroxide  (Fe2O3. 
Fe.2(OH)6).  The  rusting  of  iron  can  be  prevented  by 
coating  it  with  tin,  or  with  a  layer  of  the  black  oxide 
(Fe3O4)  by  exposing  it  to  the  action  of  steam  at  650° 
(Barff's  process).  Galvanised  iron  is  covered  with  a 
layer  of  zinc,  which  protects  the  iron  by  rendering  it 
electro-neyative. — Iron  dissolves  in  dilute  acids,  and  thus 
forms  ferrous  salts  and  hydrogen  (Expt's  28  and  29). 
It  is  also  slowly  eaten  away  by  water  in  the  presence  of 
air.  Part  of  the  iron  dissolves  as  acid  carbonate,  and 
part  of  it  forms  rust.  Caustic  soda  and  potash  prevent 
this. 

Experiment  269. — Put  a  piece  of  bright  iron  wire  in  a  beaker 
of  tap  water,  and  another  in  water  containing  caustic  soda. 
Examine  after  24  hours. 

324.  Compounds  Of  Iron.  —  Iron  forms  three 
oxides :  ferrous  (FeO),  ferric  (Fe2O3),  and  the  black 
(Fe304),  or  ferroso-ferric  oxide  (FeO.Fe/>3).  The  latter 
is  formed  when  iron  is  heated  strongly  in  air,  and  is  the 
chief  constituent  of  the  black  scales  of  the  smithy.  It 
can  also  be  prepared  by  adding  caustic  soda  to  a  solution 
containing  a  ferrous  and  a  ferric  salt,  and  then  applying 
heat.  It  then  forms  a  black  precipitate,  attracted  by  the 
magnet.  Ferrous  and  ferric  oxides  are  basic,  and  thns 


FERROUS    SALTS.  321 

there  are  two  classes  of  iron  salts  to  be  considered.     In 
the   Pharmacopeia,   the  names  of  the  ferric   salts    are 
distinguished   by  the  syllable  per-,  e.  g.,  perchloride  of 
iron,  instead  of  ferric  chloride. 

\.  Ferrous  salts  correspond  to  the  oxide  FeO,  in  which 
Fe  takes  the  place  of  H2.  They  are  easily  oxidised,  and 
it  is  difficult  to  preserve  them  unchanged.  They  are 
light  green,  or  colourless ;  have  a  sweetish,  inky,  as- 
tringent taste  ;  and  are  powerful  reducing  agents. 

2.  Ferric  salts  correspond  to  the  oxide  FejO3,  and  are 
prepared  mostly  by  the  oxidation  of  the  corresponding 
ferrous  salts.  They  are  colourless  when  anhydrous,  but 
when  hydra  ted  they  are  yellow  or  brown.  They  have 
an  astringent,  chalybeate  taste,  and  can  be  reduced  to  fer- 
rous salts  by  the  action  of  nascent  hydrogen,  &c. 


FERROUS  SALTS. 

325.  Ferrous   Sulphate   (FeS04 .  7H2O).  —  Also 

called  green  vitriol  and  copperas  (copper  rose,  from  its 
supposed  identity  with  the  green  rust  of  copper). 

PREPARATION.  —  By    slow   oxidation    of  iron    pyrites 
(FeS2)  piled  in  heaps  and  exposed  to  the  weather  : 

FeS2  +  7O  +  H2O  =  FeSO4  +  H2SO4. 

The   acid    drainage   from   these    heaps    is   treated    with 

scrap  iron  : 

H2SO4  +  Fe  =  FeSO4  +  H2. 

On  evaporation,  crystals  of  green  vitriol  are  obtained. 

Experiment  270. — Dissolve  a  few  iron  tacks  in  dilute   sul- 
phuric acid.     (What  is  the  black  substance  remaining?)     Keep 
the  solution  for  further  experiments. 
22 


322  FERROUS    SULPHATE. 

PROPERTIES. — A  green  crystalline  substance  (Examine 
a  specimen  carefully),  often  rust-coloured  on  the  surface 
from  the  oxidising  action  of  the  air.  It  is  soluble  in 
water  (70  parts  in  100),  insoluble  in  alcohol. 

Experiment  271. — To  a  portion  of  the  solution  from  Experi- 
ment 270  add  caustic  soda.  A  precipitate  falls,  which  is  at 
first  white,  but  rapidly  turns  green.  It  is  ferrous  hydroxide 
(Fe(OH)2)  : 

FeSO4  +  2NaOH  =  Fe(OH)2  +  Na2SO4. 

Close  the  1. 1.  with  the  thumb,  and  shake  vigorously.  Note  that 
the  thumb  is  pushed  inwards  showing  that  the  pressure  inside 
has  decreased,  and  that  the  hydroxide  has  become  rust-coloured : 

2Fe(OH)2  +  O  +  H20  =  Fe2(OH)6. 

If  ferrous  hydroxide  is  precipitated  in  an  atmosphere 
free  from  oxygen,  it  is  white.  When  heated,  it  loses 
water,  and  forms  ferrous  oxide  (FeO),  a  black  powder : 

Fe(OH)2  =  FeO  +  HaO. 

Experiment  272- — Boil  a  little  of  the  solution  of  ferrous 
sulphate  with  nitric  acid.  Red  fumes  are  evolved,  and  the  solu- 
tion becomes  red,  owing  to  the  formation  of  ferric  sulphate 
(Fez(SOJ3)  and  nitrate  lFe2(N08)6) : 

6FeSO4  +  8HNO3  = 

2Fe2(S04)3  +  Fe2(N03)6  +  2NO  +  4H2O. 

Add  caustic  soda  to  this  solution  ;  a  red  precipitate  of  ferric  hy- 
droxide is  thrown  down : 

Fe3(S04)3  -f  GNaOH  =  Fe2(OH)6  +  3Na2SO4. 

Experiment  273- — Pour  a  small  quantity  of  hot  saturated 
solution  of  ferrous  sulphate  into  an  equal  volume  of  alcohol. 
Shake  or  stir.  Ferrous  sulphate  is  precipitated  in  a  granular 
condition  (Ferri  sulphas  granulata). 


FERROUS    ARSENATE.  323 

Experiment  274- — Heat  a  crystal  of  green  vitriol  very  gently 
in  a  t.  t.  It  loses  water  of  crystallisation  and  falls  to  a  white 
powder  (Ferrl  sulphas  exsiccata). 

In  large  doses  ferrous  sulphate  may  act  as  poison, 
but  in  smaller  doses  it  is  a  useful  medicine. — Mohr's 
salt  is  ammonio-ferrous  sulphate  (NH4)2SO4.FeS04.6H2O. 
It  does  not  become  oxidised  as  readily  as  green  vitriol. 

326.  Ferrous  Carbonate  (FeC03). 

Experiment  275. — Heat  some  solution  of  ferrous  sulphate  to 
boiling,  and  add  to  it  solution  of  ammonic  carbonate  (best  made 
with  recently  boiled  water).  A  precipitate  falls,  which  is  at  first 
white  but  rapidly  becomes  green  by  oxidation.  It  is  ferrous 
carbonate  (FeC03) : 

(NH4)2C03  +  FeS04  =  FeCO3  +  (NH4)2S04. 

Ferrous  carbonate  is  so  easily  oxidised  that  its  pre- 
paration is  attended  with  some  difficulty.  In  medicine 
it  is  used  mixed  with  sugar  (Ferri  carbonas  saccharata). 
Griffith's  mixture  is  another  preparation  of  ferrous  car- 
bonate. They  must  be  kept  in  well  stoppered  bottles. 
Ferrous  carbonate,  on  account  of  its  easy  solubility  in 
the  gastric  juice  and  its  mild  action,  is  a  favourite  pre- 
scription of  iron. 

327.  Ferrous  Arsenate  (Fe3(As04)2). 

PREPARATION. — By  adding  solution  of  ferrous  sulphate 
to  one  of  sodic  arsenate  (Na-jHAsOJ  mixed  with  sodic 
acetate,  ferrous  arsenate  is  precipitated  as  a  greenish 
readily  oxidisable  substance  : 

2Na2HAsO4  +  2NaC2H3O2  +  3FeSO4  = 

Fe3(AsO4)2  +  3Na2S04  +  2HC2H3O2. 

It  is  always  partially  oxidised  in  the  process  of  prepara- 
tion. 


324  FEKROUS    IODIDE. 

328.  Ferrous  Phosphate  (Fes(PO4)2). 

PREPARATION. — This  compound  is  prepared  in  much 
the  same  way  as  ferrous  arsenate. 

Experiment  276. — To  solution  of  ferrous  sulphate  add  some 
sodic  acetate,  and  then  sodic  phosphate.  Ferrous  phosphate  is 
precipitated.  (Write  the  equation.) 

The  object  of  adding  sodic  acetate  in  these  processes 
is  to  provide  that  the  acid  which  is  set  free  in  the  reac- 
tion shall  not  be  a  solvent  for  the  phosphate.  Note 
that,  while  sodic  phosphate  is  an  acid  salt,  the  iron  salt 
is  normal.  If  the  sodic  acetate  were  not  added,  some  of 
the  phosphate  would  remain  unprecipitated,  being  soluble 
in  sulphuric  acid.  But  it  is  insoluble  in  acetic  acid. 

PROPERTIES.  —  Similar  in  appearance  to  the  arsenate 
(Art.  144),  but  inclining  to  blue  in  colour.  It  is  in- 
soluble in  water,  but  soluble  in  hydrochloric  acid. — To 
distinguish  the  phosphate  from  the  arsenate,  dissolve  in 
hydrochloric  acid  and  test  with  hydric  sulphide.  The 
arsenate  gives  a  yellow  precipitate,  the  phosphate  none. 

329.  Ferrous   Iodide    (Fel,). — This    is    the   green 
iodide  of  iron,  prepared  by  warming  together  3  parts  of 
of  iodine,  1^  of  iron,  and  12  of  water,  until  the  iodine  dis- 
appears, then  boiling,  filtering,  &c.     It  is  a  deliquescent 
green  salt. — Ferrous  Bromide  (FeBr2)  is  prepared  simi- 
larly.— F&rrous    Chloride    (FeCl2)    has     been     already 
noticed  (Exp't  29). 


FERRIC  SALTS. 


330.  Ferric    Chloride  (Fe.2Cl6).— Also  called  per- 
chloride  of  iron. 


FERRIC    CHLORIDE.  325 

PREPARATIONS — Experiment  277-  —  Dissolve  some  iron 
tacks  in  dilute  hydrochloric  acid  in  a  porcelain  dish  with  the 
aid  of  a  gentle  heat.  Let  there  be  insufficient  acid  to  dissolve 
the  tacks  completely.  Filter  the  solution.  (What  is  the  black 
substance  ?)  Add  a  small  quantity  of  nitric  acid  and  a  little 
hydrochloric  acid,  and  heat  quickly  until  red  fumes  are  evolved. 
Evaporate  on  the  water  bath. 

Ferrous  chloride  (FeCU  is  formed  by  dissolving  iron 
in  hydrochloric  acid  : 

Fe  -+-  2HC1  =  FeCl2  +  H2. 

This   is  changed  to  ferric  chloride   by  the  addition  of 
chlorine  produced  by  the  action  of  nitric  on  hydrochloric 
acid  : 
6FeCl2  +  6HC1  +  2HNO3  =  3Fe2Cl6  +  2NO  +  4H2O. 

PROPERTIES. — The  solution  obtained  in  Experiment 
277,  when  evaporated  to  a  syrup  and  allowed  to  cool, 
solidifies  to  a  yellowish  mass  of  the  hydrate,  Fe^lg. 
12H20. — Anhydrous  ferric  chloride  can  be  prepared  as  a 
steel-black  deliquescent  solid  by  heating  iron  in  a  current 
of  dry  chlorine  gas. — Ferric  chloride  dissolves  in  water, 
forming  a  dark  red  solution  which  becomes  yellow  on 
dilution.  The  solution  has  an  astringent  taste.  Ferric 
chloride  is  soluble  in  alcohol,  but  the  solution  (tincture 
of  iron)  tends  to  deposit  ferric  hydroxide,  and  has  no 
virtues  to  recommend  it  above  the  cheaper  aqueous  solu- 
tion.— Ferric  chloride  dissolves  ferric  hydroxide,  and 
thus  forms  soluble  basic  salts,  milder  in  their  action  than 
the  normal  salt.  By  dialysis  of  ferric  chloride,  soluble 
ferric  hydroxide  (dialysed  iron)  can  be  obtained.'  It  is 
an  excellent  antidote  for  arsenic  poisoning,  and  is  the 
best  preparation  of  iron  for  a  delicate  stomach.  It  un- 
fortunately gelatinises  after  a  time. 


326  FERRIC    NITRATE. 

Experiment  278. — Make  a  small  boat  of  parchment  paper, 
fill  it  about  one-fourth  with  dilute  solution  of  ferric  chloride, 
float  it  in  a  basin  or  beaker  of  distilled  water,  and  leave  it  for 
24  hours.  Examine  the  water  for  hydrochloric  acid,  by  taste, 
litmus,  &c.  Taste  the  iron  solution.  It  has  lost  much  of  its 
astringency.  Boil  some  of  it  in  a  t.  t.  Ferric  hydroxide  is  pre- 
cipitated. 

In  this  experiment  water  decomposes  ferric  chloride : 
Fe2016  +  6H20  =  Fe2(OH)6  +  6HCL 

The  crystalloid  hydrochloric  acid  passes  through  the 
membrane,  while  colloid  ferric  hydroxide  remains. 

331.  Ferric  Sulphate  (Fe.2(SO4)3).— This  salt  has 

been  already  noticed  (Exp't  272).  If  it  is  desired  to 
prepare  the  pure  sulphate,  sulphuric  acid  must  be  added 
according  to  the  equation  : 

6FeSO4  +  3H2S04  +  2HN03  =  3Fe2(SO4)3  +  2NO  +  4H2O. 

It  forms  a  reddish-brown  solution,  of  strongly  acid  reac- 
tion.— Ferric  sulphate  unites  with  potassic  sulphate  and 
water  to  form  iron  alum,  K2SO4.Fe2(SO4)3.24HoO. 

332.  Ferric  Nitrate  (Fe2(NO3)6).—  When  iron  dis- 
solves in   cold  very  dilute   nitric   acid,  ferrous   nitrate 
(Fe(NO3)2),  and  ammonic  nitrate  are  formed  : 

4Fe  +  10HNO3  =  4Fe(NO3)2  +  NH4NO3  +  3H2O. 

But  when  the  action  is  hastened  by  heat  or  by  using  a 
stronger  acid,  ferric  nitrate  is  one  product,  and  some 
oxide  of  nitrogen  another,  e.g.  : 

2Fe  -f  8HNO3  =  Fe2(N03)6  +  2ND  +  4H2O. 

It  forms  a  solution  of  a  reddish-brown  colour,  the  ferri 
pernitratis  liquor  of  the  Pharmacopeia. 


FKKRIC    HYDROXIDE.  327 

333.  Ferric  Hydroxide  (Fe2(OH)6).     This  is  pre- 
pared as  in  the  second  part  of  Experiment  272,  but  am- 
monia is  generally  used  instead  of  caustic  soda.     It  is  a 
reddish    brown    substance    which    gradually    undergoes 
change  even  when  kept  in  water.     It  loses  water  when 
dried  in  the  air,  forming  a  hydroxide  of  the  composition 
Fe2O3.H2O.       The  same    decomposition    goes   on    under 
water,  and  the  dehydrated  compound  is  less  active,  e.g. 
it  does  not  combine  with  arsenic  trioxide. 

Experiment  279. — Heat  a  small  quantity  of  ferric  hydroxide 
on  mica  until  it  is  converted  into  the  oxide  (Fe203).  Try  to 
dissolve  this  in  hydrochloric  acid.  (Is  the  hydroxide  soluble  in 
hydrochloric  acid?) 

Experiment  280. — Add  solution  of  sodic  carbonate  to  one  of 
ferric  chloride.  Ferric  hydroxide  is  precipitated.  It  is  so  weak 
a  base  that  it  does  not  form  salts  with  weak  acids  : 

Fe2016  +  3Na2CO3  +  3H2O  = 

Fe2(OH)6  +  6NaCl  +  3CO2. 
(Did  you  observe  the  evolution  of  carbon  dioxide  ?) 

334.  Ferric   Oxide   (Fe2O3).     This    compound   has 
been  already  noticed  several  times. 

Experiment  281. — Heat  a  few  crystals  of  green  vitriol 
strongly  in  a  porcelain  crucible.  Ferric  oxide  remains  as  a  red 
powder,  called  colcothar,  crocus,  rouge,  or  Venetian  red.  Try  to 
dissolve  some  of  it  in  strong  acids.  It  is  insoluble. 

Ignited  ferric  oxide  should  never  be  used  in  medicine 
instead  of  the  hydroxide  described  in  Art.  333,  as  it 
is  insoluble  in  acids,  and  therefore  useless  for  prescription 
as  an  iron  preparation. 

335.  The    "  Scale "    Compounds   of  Iron.— 

Certain  organic  acids  (citric,  tartaric,  <fec.,)  prevent  the 
precipitation  of  ferric  hydroxide  by  ammonia. 


328  THE    "  SCALE  "    COMPOUNDS. 

Experiment  282. — To  a  solution  of  ferric  chloride  add  tar- 
taric  acid,  and  then  ammonia  until  the  liquid  is  alkaline.  No 
precipitate  forms. 

This  is  due  to  the  formation  of  a  soluble  basic  tartrate 
of  iron  and  ammonium.  If  ferric  hydroxide  is  dissolved 
in  tartaric  acid,  ferric  tartrate  is  formed.  When  ammonia 
is  added  to  this  and  the  whole  evaporated  to  dryness,  a 
basic  salt  is  obtained  in  red  amorphous  scales.  Similarly 
with  citric  acid.  Quinine  and  other  alkaloids  are  added 
to  these  scale  preparations.  They  contain  variable 
quantities  of  iron,  and  are  rather  difficult  to  prepare. 

336.  Tests. 

Ferrous  Salts. 

1.  Solutions  of  ferrous  salts  give  with  ammonia  a  greenish 
precipitate  (Fe(OH)2),  turning  rust  coloured  when  shaken  up 
with  air. 

2.  Ammonic  sulphide  gives  a  black  precipitate  (FeS),  soluble 
in  dilute  hydrochloric  acid  : 

FeS  +  2HC1  =  FeCl2  +  H2S. 

3.  Potassic  ferrocyanide  (K4:Fe(CN)6)  gives  a  white  or  light 
blue  precipitate  of  potassic  ferrous  ferrocyanide  : 

K4.Fe(CN)6  +  FeS04  =  K2Fe.Fe(CN)6  +  K2SO4. 
This  quickly  turns  blue  by  oxidation,  forming  Prussian  blue. 

4.  Potassic  ferricyanide  (K3Fe(CN)6)  gives  a  deep  blue  pre- 
cipitate (or   colour,  according  to  the  strength  of  the  solution) 
(KFe.  Fe(CN)  6).     (Turnbull's  Blue. ) 

5.  Ferrous  solutions,  when  heated  with  nitric  acid,  turn  red, 
and  will  then  give  a  reddish-brown  precipitate  with  ammonia. 

Ferric  Salts. 

1.  Ammonia  gives  a  reddish-brown  precipitate  of  ferric  hy- 
droxide, insoluble  in  excess. 


CHROMIUM.  329 

2.  Ammonic  sulphide  gives  a  black  precipitate  of  ferrous  sul- 
phide mixed  with  sulphur  : 

Fe2Cl6  +  3(NH4)2S  =  2FeS  +  S  +  6NH4C1. 

This  is  soluble  in  dilute  hydrochloric  acid,  the  sulphur  remain- 
ing undissolved. 

3.  Potassic  ferrocyanide  gives  a  deep  blue  precipitate  of  Prus- 
sian blue. 

4.  Potassic  ferricyanide  gives  a  greenish-brown  colour,  but  no 
precipitate. 

5.  With  hydric  sulphide  ferric  salts  give  a  white  precipitate  of 
sulphur.     FeaClg  +  HaS  =  2FeCl2  -f  2HC1  +  S. 

6.  Insoluble  iron  compounds  are  tested  by  the  borax  bead, 
which  with  iron  is  yellow  in  the  oxidising,  colourless  or  blue  in 
the  reducing  flame. 


CHROMIUM. 

337.  Chromium  (Or  "•  1T-VI  =  52.4).— The  metal  it- 
self is  of  no  importance.     The  principal   ore  is  chrome 
ironstone,  or  chromite  (FeO.Cr.2O3). — The  name  Chromium 
is  derived   from  a  Greek  work  meaning  colour,  because 
all  chromium  compounds  are  coloured.     The  metal  is  re- 
duced from    its    ores  with  great  difficulty.     Chromium 
ores  are  sometimes  added  to  iron  ores,  because  chromium 
imparts  great  hardness  to  steel.    Chromium  steel  requires 
to  be  worked  at  comparatively  low  temperatures. 

338.  Compounds  of  Chromium.  —  Chromium 

unites  with  oxygen  in  three  portions,  forming  two  basic 
oxides, — chromovs  (CrO),  and  chromic  (Cr2O3);  and  one 
acid-forming  oxide, —  chromium  trioxide  (CrO3). — The 
chromous  salts  (CrCl2,  CrSO4,  <fec.)  are  very  unstable, 
becoming  oxidised  even  more  readily  than  ferrous  salts.  — 
The  chromic  salts  (Cr2Cl6,  Cr2(SO4)3,  &c.)  are  similar  to 
ferric  salts  in  composition  and  properties.  They  are  the 


330  POTASSIC    BICHROMATE. 

ordinary  salts  of  chromium. — Besides  these  two  series  of 
compounds  there  are  the  chromates,  salts  of  chromic  acid 
(H2Cr04). 

339.  Potassic    Bichromate    (K2Cr2O7  =  .  K2O. 

2CrO3). — This  salt  is  the  starting  point  in  preparing 
chromium  compounds. 

PREPARATION. — Chrome  ironstone  is  roasted,  ground, 
and  heated  with  lime  and  potassic  carbonate,  with  con- 
stant stirring  so  as  to  allow  oxidation  to  go  on.  The 
object  of  the  lime  is  to  economise  alkali,  and  to  prevent 
fusion,  so  that  the  air  may  penetrate  into  the  mass : 

Cr203  +  2K2CO3  +  30  =  2K2CrO4  +  2CO2. 
Cr2O3  +  2CaO  +  3O  =  2CaCrO4. 

When  oxidation  is  complete,  the  potassic  and  calcic 
chromates  are  dissolved  in  water,  and  to  the  solution 
potassic  sulphate  (K2SO4)  is  added  to  precipitate  calcium  : 

K3SO4  +  CaCrO4  =  CaSO4  +  K2CrO4. 
To  the  strong  solution  of  potassic  chromate  thus  obtained 
sulphuric  acid  is  added  to  form  the  less  soluble  bichromate: 

2K2CrO4  +  H2SO4  =  K2Cr2O7  +  K2SO4  +  H2O. 

The  object  of  this  operation  is  to  separate  the  chromium 
salt  from  the  impurities  present  in  the  solution. 

PROPERTIES. — Potassic  bichromate  crystallises  in  large 
garnet  red  prisms,  soluble  in  water  (8  parts  in  100).  It 
is  a  strong  oxidising  agent,  especially  in  the  presence  of 
acids.  An  instance  of  this  has  been  already  described 
(Aldehyde). 

Experiment  283. — To  a  solution  of  potassic  bichromate 
acidified  with  sulphuric  acid  add  sulphuretted  hydrogen.  The 
red  colour  changes  to  green,  and  sulphur  is  precipitated  : 


INSOLUBLE    CHROMATES.  331 

K2Cr207  +  4H2SO4  +  3H2S  = 

K2S04  +  Cr2(S04)3  +  7H20  +  3S. 

In  this  action  chromium  trioxide  is  an  oxidising  agent, 
and  the  hydrogen  of  hydric  sulphide  is  oxidised  to  water. 
Leaving  the  acid  out  of  consideration  the  action  can  be 
represented  more  simply : 

2Cr03  +  3H2S  =  Cr2O3  +  3H2O  -f  3S. 

Experiment  284. — To  50  c.  c.  of  saturated  solution  of  potassic 
bichromate  add  5  or  10  drops  of  sulphuric  acid,  and  then  sul- 
phurous acid  until  the  colour  is  bright  green.  Evaporate  to  a 
small  bulk  and  set  aside.  The  solution  contains  chrome  alum 
(K2S04.Cr!i(SOJ3.24HaO): 

K2Cr207  +  H2S04  +  3S03  =  K2S04.Cr2(SO4)3  +  H2O. 

Potassic  bichromate,  on  account  of  its  oxidising  power, 
is  a  corrosive  poison.  The  antidote  is  ferric  chloride, 
which  forms  the  sparingly  soluble  ferric  chromate. 
Alkaline  sulphites  should  also  be  antidotal.  (Why1?) 

Potassic  bichromate  is  used  in  the  preparation  of  other 

chromates.     The   following  are   insoluble  in  water  but 

soluble    in     dilute     acids  :     baric     (BaCrO4),     plumbic 

(PbCrO4),    argentic    (A.g2CrO4),    mercurous    (Hg2CrO4), 

ferric  (Fe2(Cr04)3),  and  others. 

Experiment  285. — Add  solution  of  potassic  bichromate  to 
solutions  of  the  following  salts,  note  the  colour  of  the  precipitates 
and  test  their  solubility  in  acetic  and  nitric  acids,  viz. ,  baric 
chloride  (BaCl2),  plumbic  acetate  (Pb(C2H302)2),  argentic  nitrate 
(AgN03),  and  mercurous  nitrate  (Hg  2(N03)  2).  Normal  chromates 
are  precipitated,  and  acid  is  set  free  in  each  case.  (Write  the 
equations.) 

Plumbic  chromate  (PbCr04)  is  used  as  a  paint  (chrome 
yellow).  Chrome  red  is  a  basic  chromate  of  lead  pre- 
pared by  boiling  the  normal  chromate  with  lime  water. 


332  CHROME    ALUM. 

Experiment  286. — Dip  a  piece  of  white  cotton  in  dilute  solu- 
tion of  pi  imbic  acetate,  soak  it  well,  wring  it,  and  then  dip  it  in 
dilute  solution  of  potassic  bichromate.  It  is  dyed  yellow,  and 
the  colour  is  fast,  being  precipitated  within  the  fibres  of  the 
cotton.  Boil  the  cotton  with  lime  water  ;  it  becomes  orange- 
red. 

340.  Chromic  Acid. 

Experiment  287. — To  a  cold  saturated  solution  of  potassic 
bichromate  add  one  and  a  half  times  its  volume  of  concentrated 
sulphuric  acid,  taking  care  to  stir  well.  Set  in  a  cool  place. 
Chromium  trioxide  (Cr03)  separates  out  in  beautiful  red  crystals. 

Chromium  trioxide  dissolves  in  water,  forming  a  strongly 
acid  solution,  but  no  definite  acid  hns  been  separated 
from  this.  The  normal  chromates  are  analogous  to  the 
sulphates  in  composition,  e.g.  K2CrO4,  PbCrO4,  &c.,  so 
that  chromic  acid  may  be  supposed  to  have  the  formula 
H2CrO4.  Potassic  bichromate  is  then  an  anhydrous 
acid  salt. 

341.  Chrome     Alum     (K2S04.Or2(SO4)8.24H2O). 
This  salt  is  prepared  as  in  Experiment  284.     It  is  a  by- 
product in  some  operations  in  which  potassic  bichromate 
is  used  as  an  oxidising  agent,  e.g.  in  the  manufacture  of 
alizarine.     It  is  soluble  in  water,  and  crystallises  from 
cold  solutions  in  violet  crystals  isomorphous  with  those 
of  common  alum.     If  a  crystal  of  common  alum  be  placed 
in  a  saturated  solution  of  chrome  alum  it  grows  by  addi- 
tion of  layers  of  the  chrome  alum.     Solutions  of  chrome 
alum  undergo  a  peculiar  change  on  being  heated.     The 
colour  changes  from  violet  to  green,  and    this  solution 
do§s  not  crystallise.     It   slowly   returns    to   its   former 
condition.      This  property   is    common   to    all    chromic 


CHROMIC    HYDROXIDE.  333 

salts. — Chrome  alum    is  used    in  tanning,    dyeing,  and 
calico-printing. 

34-2.  Chromic  Hydroxide  (Cr2(OH)6). 

Experiment  288-—  Add  ammonia  to  a  solution  of  chrome 
alum.  Collect  the  precipitate  of  chromic  hydroxide  on  a  filter 
and  wash  it.  Try  the  solubility  of  portions  of  it  in  hydrochloric 
acid,  caustic  soda,  and  ammonia.  Heat  part  of  it  on  mica. 

Chromic  hydroxide  is  of  a  dirty  green  colour.  It  dis- 
solves in  hydrochloric  acid,  forming  chromic  chloride 
(Cr2Cl6).  It  is  also  soluble  in  caustic  soda,  forming  a 
green  solution  from  which  it  is  reprecipitated  by  boiling. 
When  heated  it  loses  water,  and  chromium  sesquioxide 
(Cr.2O3)  remains.  This  oxide  is  used  as  a  paint  (chrome 
green).  Guignefs  green  has  the  composition,  Cr2O3.2H2O. 
It  is  also  sold  as  chrome  green. 

343.  Tests. 

Chromales. 

1.  Baric  chloride  gives  a  yellow  precipitate  insoluble  in  acetic 
acid,  soluble  in  nitric  acid. 

2.  Acidified  solutions  of  chromates  are  turned  green  by  hydric 
sulphide. 

8.  Solutions  of  chromates  are  reduced  by  ammonic  sulphide, 
which  precipitates  chromic  hydroxide,  so  that  this  group  reagent 
is  a  test  for  chromates  as  well  as  for  chromic  salts. 

4.  Insoluble  chromates  can  be  tested  for  by  the  borax  bead, 
to  which  they  give  an  emerald  green  colour. 

Chromic  Salts. 

i.  Ammonia  precipitates  chromic  hydroxide,  insoluble  in 
excess. 

•2.  Ammonic  sulphide  gives  a  dirty  green  precipitate  of  hy- 
droxide : 

Cr2(S04)3  +  3(NHJ2S  +  3H2O  = 

Cr2(OH)6  +  3(NH4)2S04  +  3H2S. 


334  ALUMINIUM. 

3.  Caustic  soda  gives  a  green  precipitate  soluble  in  excess,  re- 
precipitated  by  boiling. 

4.  Insoluble  chromium  compounds  can  be  detected  by  the 
borax  bead,  or  by  heating  in  the  oxidising  flame  with  a  sodic 
carbonate  bead,  to  which  they  give  a  yellow  colowr,  due  to  the 
formation  of  sodic  chromate.     If  the  bead  be  dissolved  in  water 
the  yellow  colour  appears  strongly. 


ALUMINIUM. 

344.  Aluminium  (Aliv  =  27.3.— Sp.  wt.  =±  2.67.— 
Melting  point  =  700°.— Sp.  heat  =  0.2143). 

OCCURRENCE. — Compounds  of  aluminium  form  a  very 
considerable  proportion  of  the  earth's  crust,  being  found 
in  clay,  granite,  gneiss,  mica,  felspar,  &c. 

PREPARATION. — From  bauxite,  a  hydroxide  of  alum- 
inium and  iron.  Aluminium  chloride  (A12C16)  is  obtained 
by  a  series  of  operations,  and  from  this  the  metal  is  set 
free  by  sodium.  Lately  the  metal  has  been  obtained 
more  economically  by  electrolysis. 

PROPERTIES. — A  light,  tin-white  metal,  malleable, 
ductile,  and  sonorous.  When  pure  it  does  not  tarnish  in 
air,  but  the  impure  metal  soon  tarnishes ;  and  this  is  one 
difficulty  in  the  way  of  the  economical  manufacture  of 
the  metal.  It  decomposes  water  at  100°,  and  dissolves 
easily  in  most  acids  and  alkalis.  Obviously,  it  cannot 
be  used  for  cooking  utensils.  It  is  very  useful  wherever 
lightness  and  durability  are  required,  as  in  optical  instru- 
ments, &c.  Aluminium  bronze  is  an  alloy  of  aluminium 
with  90  %  of  copper.  It  has  the  appearance  and  many 
of  the  qualities  of  gold. 


ALUMINA.  335 

345.  Alumina  (A1203).     This  is  the   only  oxide  of 
aluminium. 

OCCURRENCE. — As  corundum,  ruby,  sapphire,  emery, 
<fcc. ;  and  combined  in  many  silicates,  &c, 

PREPARATION. — Experiment  289.— To  a  solution  of  alum 
add  ammonia.  Collect  on  a  filter,  and  wash,  the  gelatinous  pre- 
cipitate of  aluminic  hydroxide,  A12(OH)6  : 

6NH4OH  +  A12(S04)3  =  A12(OH)6  +  3(NH4)2S04. 

Heat  a  portion  of  the  precipitate  on  mica.     It  decomposes  into 
water  and  alumiuic  oxide  : 

Al3(OH)e  =  A12O3  +  3H2O. 

PROPERTIES. — A  white  powder,  insoluble  in  acids  after 
it  has  been  ignited.  Crystalline  alumina  is  next  to 
diamond  in  hardness,  and  in  the  form  of  emery  is  used 
in  grinding  and  polishing  hard  substances. 

Aluminic  hydroxide,  A12(OH)6,  is  a  weak  base,  and 
also  a  weak  acid. 

Experiment  290. — To  a  solution  of  alum  add  caustic  soda,  a 
little  at  a  time.  Aluminic  hydroxide  is  precipitated  and  re- 
dissolved  : 

A12(OH)6  +  6NaOH  =  Al2(ONa)6  +  6H3O. 

Aluminic  hydroxide  dissolves  in  solutions  of  caustic 
soda  and  caustic  potash,  forming  aluminates.  It  does  not 
dissolve  in  solution  of  ammonia. — It  has  the  power  of 
extracting  colouring  matters  from  solution,  and  is  used 
as  a  clarifier,  decolouriser,  and  as  a  mordant  in  dyeing. 

346.  Aluminic  Salts.     Aluminium  salts  are  mostly 
colourless,  and,  when  the  acid  is  a   strong  one,  of  an 
astringent,  acid  taste.     They  resemble  ferric  and  chromic 
salts  in  composition,  e.g.  A12C16,  A12(S04)3,  A12(NO3)6,  &c. 


336  ALUMS. 

Like  chromium  and  iron  (in  ferric  compounds),  aluminium 
does  not  form  salts  of  such  weak  acids  as  carbonic. 

347.  Alums.  Aluminic  sulphate  (A12(SO4)3)  com- 
bines with  potassic  sulphate  (K2S04)  and  water  to  form 
potash  alum,  K2SO4.  A12(SO4)3.24H2O.  It  also  unites 
with  ammonic  sulphate  to  form  ammonia  alum 
(NH4)2SO4.A12(SO4)3.24H2O.  The  alums  are  a  group  of 
compounds  similar  in  properties  and  composition,  and 
exactly  alike  in  crystalline  form.  A  crystal  of  any  one 
alum  will  increase  in  size  when  placed  in  a  saturated  solu- 
tion of  any  other ;  the  alums  are  isomorphous.  A  gen- 
eral formula  may  be  written  thus  : 

M2S04.M'2(S04)3.24H20. 

M  =  K,  NH4,  Na,  Rb,  Cs,  Ag,  or  Tl. 
M'=  Al,  Fe,  Or,  In,  or  Ga. 

PREPARATION. — Potash  and  ammonia  alums  are  the 
ones  in  common  use.  They  are  prepared  mostly  from 
shale,  which  contains  clay  (a  silicate  of  aluminium)  and 
iron  pyrites  (FeS2).  The  shale  is  burned  in  heaps,  when 
aluminic  and  ferrous  sulphates  are  formed.  By  lixivia- 
tion  a  solution  of  aluminic  siilphate  is  obtained.  To  this 
potassic  or  ammonic  sulphate  is  added,  and  the  alum  is 
obtained  by  evaporation,  and  purified  by  recrystallisation. 

PROPERTIES. — Potash  and  ammonia  alums  are  colour- 
less solids,  generally  sold  in  large  crystals.  They  are 
exactly  alike  in  appearance,  and  can.  only  be  distin- 
guished by  a  chemical  test.  They  have  an  acid,  sweetish, 
astringent  taste.  They  are  soluble  in  water  (12  parts  in 
100),  the  potash,  a  little  more  so  than  the  ammonia, 
alum.  They  effloresce  slowly  in  air,  owing  to  the  action 


PORCELAIN.  337 

of  the  ammonia  of  the  air  in  forming  basic  salts. — Potash 
alum  is  now  the  one  generally  sold. 

Experiment  291. — Heat  a  crystal  of  potash  alum  in  at.  1. 1 
and  observe  the  loss  of  water  of  crystallisation.  The  white 
powder  which  remains  is  anhydrous  alum,  or  burnt  alum  (alumen 
ustum). 

348.  Aluminic  Sulphate  (A12(SO4)3.18H2O).  This 

salt,  known  as  concentrated  alum,  or  alum  cake,  is  pre- 
pared on  the  large  scale  by  the  action  of  sulphuric  acid 
on  certain  clays,  the  products  being  aluminic  sulphate 
and  silicic  acid.  It  replaces  the  more  expensive  alum  in 
many  of  the  uses  to  which  that  substance  has  been  put. 

349.  Porcelain,    &C.     Porcelain    is    made    from    a 
pure  white  clay  (kaolin,  or  china-clay),  a  hydrated  silicate 
of  aluminium   (Al203.2SiO2.2H2O).     In   the    process    of 
burning,  the  water  is   driven  off.     The  glaze  is  felspar, 
borax,  bone  ash,  or  red  lead. — Stoneware,  earthenware, 
and  common  pottery,  are  made  from  impure  clays. 

350.  Tests. 

1.  Ammonia   precipitates    aluminic    hydroxide,    insoluble   in 
excess. 

2.  Ammonic  sulphide  gives  the  same  precipitate  : 

A12(S04)3  +  3(NH4)2S.  +  3H20  = 

A12(OH)6  +  3(NH4)2S04  +  3H2S. 

3.  Caustic  soda  precipitates  aluminic  hydroxide,  but  redissolves 
it.     (Experiment  290.)     From  this  solution  aluminic  hydroxide 
is  precipitated  by  ammonic  chloride  : 

Ala(ONa)6  +  6NH4C1  =  A12(OH)8  +  6NaCl  +  6NH3. 

4.  Insoluble  aluminium  compounds  are  detected  by  moistening 
with  cobaltous  nitrate,  Co(N08)2,  and  heating  with  the  blowpipe 
on  charcoal.     A  deep  blue  colour  is  imparted. 

23 


338  ZINC. 

B 
ZINC. 

351.  Zinc   (Zn11  =  64.9.— Sp.    wt.   =  6.9.— Melting 
point  =  433°.— Boiling    point  =  1040°.— Sp.     heat  = 
0.09555). 

OCCURRENCE. — The  principal  ores  of  zinc  are  calamine, 
or  zinc  spar  (ZnC03),  zinc  blende  (ZnS),  franklinite 
(ZnO.Fe2O3),  and  red  zinc  ore  (ZnO). 

PREPARATION. — The  ore  is  roasted,  and  reduced  by 
heating  with  pounded  coal  in  clay  retorts.  The  metal 
distils  over  and  is  condensed  in  iron  tubes : 

ZnO  +  C  =  Zn  +  CO. 

Commercial  zinc  contains  lead,  carbon,  iron,  &c.,  as  im- 
purities. Arsenic  is  often  present,  a  fact  to  be  remem- 
bered in  making  Marsh's  test. 

PROPERTIES. — Zinc  is  of  a  bluish  white  colour  when 
pure.  When  heated  strongly  in  air  it  burns,  forming 
zinc  oxide  (ZnO).  Commercial  zinc  is  brittle  at  ordinary 
temperatures,  but  is  very  malleable  and  ductile  between 
100°  and  150°.  It  is  easily  oxidisable,  and  dissolves 
readily  in  acids.  Alloyed  with  copper  it  forms  brass. — 
Zinc  is  used  in  the  manufacture  of  various  utensils,  in 
galvanizing  iron,  in  generating  electricity,  in  preparing 
hydrogen,  &c. 

352.  Zinc    Oxide    (ZnO). — Zinc    forms    only    one 
oxide,  known  in  commerce  as  zinc  white. 

PREPARATION. — (1)  By  boiling  zinc  and  burning  its 
vapour. — (2)  The  British  Pharmacopoeia  directs  it  to  be 
prepared  by  heating  zinc  carbonate  in  a  loosely  covered 


ZINC    CHLORIDE.  339 

crucible,  until   a  portion  taken  out  does   not  effervesce 

with  acids  : 

ZnC03  =  ZnO  +  CO2. 

This   method    of   preparing   oxides    of    metals    is   often 
employed. 

PROPERTIES. — A  soft,  white  powder,  tasteless,  and  in- 
odorous. It  is  insoluble  in  water.  It  is  used  in  medicine 
and  as  a  paint. 

Experiment  292. — Heat  a  small  quantity  of  zinc  carbonate 
in  a  porcelain  crucible  for  15  minutes.  Test  a  part  of  it  with 
dilute  sulphuric  acid.  Note  that  the  oxide  is  yellow  while  hot. 

353.  Zinc  Salts.     The  salts  of  zinc  are  colourless, 
unless  the  acid  is  coloured.     The  soluble  salts  have  an 
acid  reaction  in  solution,  and  a  nauseoxis  metallic  taste. 
They  are  poisonous  and  act  as  emetics. 

354.  Zinc  Chloride  (ZnCl2).—  This  salt  is  known 
as  "  butter  of  zinc." 

PREPARATION. — The  soluble  salts  of  zinc  are  usually 
prepared  from  the  commercial  metal,  and  the  method  of 
preparation  is  similar  for  most  of  the  pharmaceutical 
preparations. 

Experiment  293- — Dissolve  some  granulated  zinc  in  dilute 
hydrochloric  acid  by  warming  in  a  porcelain  dish.  Boil,  filter, 
and  add  chlorine  water  to  convert  ferrous  to  ferric  chloride. 
Add  zinc  carbonate  until  a  brownish  precipitate  appears : 

3ZnOO3  +  Fe2Cl6  +  3H2O  = 
3ZnCl2  +  Fe2(OH)6  +  3CO2. 

Filter  again,  and  evaporate  until  oily.     Set  aside  for  a  day  and 
observe. 

PROPERTIES. — A  white,  fusible  solid,  very  deliquescent. 
It  has  strong  caustic  properties,  and  the  anhydrous  salt 


340  ZINC    SULPHATE. 

chars  sugar,  &c.  It  dissolves  easily  in  water,  alcohol, 
and  ether. — It  is  used  for  weighting  cotton  goods.  In 
surgery  it  is  employed  as  a  caustic  and  antiseptic. — 
Burnett's  disinfecting  fluid  is  a  solution  of  zinc  chloride. 
It  is  very  poisonous,  and  has  sometimes  been  swallowed 
by  mistake.  The  chemical  antidotes  are  chalk,  magnesia, 
sodic  carbonate,  &c.  (Explain  the  action  of  these  anti- 
dotes.) 

355.  Zinc     Sulphate    (ZnSO4.7H2O),— generally 
known  as  white  vitriol. 

PREPARATION. — Dissolve  zinc  in  dilute  sulphuric  acid, 
and  then  proceed  as  with  zinc  chloride. 

PROPERTIES. — A  white  crystalline  salt,  very  much 
resembling  Epsom  salts  (MgS04.7H2O),  with  which  it  is 
isomorphous.  (These  two  salts  can  be  readily  distin- 
guished by  their  taste.)  It  effloresces  when  exposed  to 
air.  It  is  an  irritant  poison  when  taken  in  large  doses. 
The  antidotes  are  albumen,  tannin  solutions  (tea,  &c.), 
and  sodic  carbonate. — Zinc  sulphate  dissolves  readily  in 
water  (1  part  in  2  of  water). 

356.  Zinc  Carbonate  (ZnC03).     The  normal  car- 
bonate of  zinc  is  difficult  to  prepare.     Zinci  carbonas  of 
the  Pharmacopoeia  is  a  basic  salt  (ZnC03.2ZnO.3H2O). 

Experiment  294. — Add  sodic  carbonate  to  solution  of  zinc 
sulphate.  Basic  zinc  carbonate  is  precipitated.  Note  its  appear- 
ance. (What  gas  is  evolved  ?)  Filter,  wash,  test  solubility  in 
acids.  (Has  this  salt  any  taste  ?) 

357.  Zinc  Acetate  (Zn(C2H3O2)2.H20). 

PREPARATION. — Experiment  295-  --Dissolve  zinc  carbonate 
in  acetic  acid  until  no  more  will  dissolve,  filter  if  necessary, 


MANGANESE.  341 

evaporate  to  small  bulk,  adding  a  few  drops  of  acetic  acid  from 
time  to  time,  and  set  aside  to  crystallise. 

PROPERTIES. — A  colourless  solid,  crystallising  in  thin, 
pearly  plates.  It  is  soluble  in  water,  and  the  solution 
has  a  sharp,  unpleasant,  metallic  taste. 

358.  Tests. 

1.  Ammonia    gives    a    white    precipitate   of    zinc    hydroxide 
(Zn(OH)2),  soluble  in  excess. 

2.  Ammonic  sulphide  gives  a  white  precipitate  of  zinc  sulphide 
(ZnS) : 

ZnSO4  +  (NH4)2S  =  ZnS  +  (NH4)2S04. 

This  is  soluble  in  dilute  hydrochloric  acid. 

3.  Caustic  soda  precipitates  zinc  hydroxide,  and  then  redis- 
solves  it : 

ZnSO4  +  2NaOH  =  Na2SO4  +  Zn(OH,)2. 
Zn(OH)2  +  2NaOH  =  Zn(ONa/)2  +  2H2O. 

This  solution  gives  no  precipitate  with  ammonic  chloride  (because 
zinc  hydroxide  is  soluble  in  ammonia),  but  gives  a  white  pre- 
cipitate of  zinc  sulphide  when  treated  with  hydric  sulphide. 

4.  Zinc  compounds  insoluble  in  water  can  be  dissolved  in 
dilute  sulphuric  or  hydrochloric  acid,  and  tested  as  above. 


MANGANESE. 

359.  Manganese  (Mnu-lT-YL  =  54.8.— Sp. wt.  =  8.—) 
Compounds  of  manganese  are  widely  distributed  in  small 
quantities.  They  give  the  colour  to  many  otherwise 
colourless  minerals,  e.g.  many  silicates ;  and  are  found  in 
minute  quantities  in  both  plants  and  animals. — The 
chief  ores  of  manganese  are  pyrolusite,  or  black  oxide  of 
manganese  (MnO2),  braunite  (Mn2O3),  hausmannite 
(Mn3O4),  psilomelane  (BaO.MnO2),  and  rhodocrozite 


342  MANGANESE    DIOXIDE. 

(MnCO3).  The  metal  is  of  little  importance.  It  can  be 
prepared  by  reducing  manganous  oxide  (MnO)  with 
charcoal  at  a  very  high  temperature.  It  decomposes 
warm  water.  Its  presence  in  iron  renders  that  metal 
very  hard. 

360.  Oxides  Of  Manganese.     Manganese  unites 
with   oxygen  in  four  proportion,  forming  two  basic, — 
manganous  oxide  (MnO),  and  manganic  oxide  (Mn2O3) ; 
and    two   indifferent,    oxides, — red   oxide    of  manganese 
(Mn3O4\  and  manganese  dioxide  (MnO2).     It  unites  with 
oxygen   and    hydrogen    to    form    two    acids — manganic 
(H2MnO4),  and  permanganic  (HMnO4).     There  is  some 
evidence  of  the  existence  of  an   oxide   (Mn2O7)  corres- 
ponding to  permanganic  acid  : 

2HMnO4  =  H2O  +  Mn2O7. 

361.  Manganese    Dioxide    (MnO2).     This   sub- 
stance is  found  in  large  quantities  as  the  mineral  pyro- 
lusite,   or  black  oxide  of   manganese.     It  is    the   most 
important  compound   of  manganese,  being   used  in  the 
manufacture  of  bleaching  powder  and  glass,  and  in  the 
preparation  of  oxygen,  chlorine,  potassic  permanganate, 
&c. 

PROPERTIES. — A  brownish  black  solid.     When  heated, 
it  gives  off  one-third  of  its  oxygen  : 

3Mn02  =  Mn304  +  O2. 

Strong  hot  sulphuric  acid  causes  it  to  lose  half  its 
oxygen : 

MnO2  +  H2SO4  =  MnSO4  +  O  +  H2O. 

The  action  of  hydrochloric  acid  has  been  already  studied. 
(Art.  92.)-^- Manganese  dioxide  is  sometimes  fraudulently 


MANGANATES.  343 

mixed  with  coal  dust  and  charcoal  powder.     This  makes 
it  explosive  when  heated  : 

2MnO2  +  C  =  2MnO  +  CO2. 

362.  Manganous  Salts.     In  these  salts  manganese 
is  bivalent  (MnCl2,  MnSO4,  &c.).     The  soluble  salts  can 
be  prepared   by  dissolving  manganese  dioxide   or  man- 
ganous  hydroxide  (Mn(OH)2)   in    the    acids ;    and    the 
insoluble  salts  by   precipitation.     Manganous  salts  are 
usually  pink  or  rose-coloured,  and  are  not  easily  oxidised 
to   manganic   salts.     (Compare  with  ferrous  and  ferric 
salts. ) 

Experiment  296- — Heat  a  little  manganese  dioxide  with  con- 
centrated sulphuric  acid.  Observe  the  evolution  of  oxygen.  A 
solution  of  manganous  sulphate  (MnS04)  is  obtained.  There  is 
usually  a  reddish  residue  of  ferric  oxide,  nearly  always  present  in 
manganese  dioxide  as  an  impurity.  The  solution  contains  ferric 
sulphate  as  an  impurity.  This  can  be  removed  by  heating  the 
impure  manganous  sulphate  to  redness,  and  thus  decomposing 
the  ferric  sulphate.  Manganous  sulphate  can  then  be  dissolved 
out. 

363.  Manganic  Salts.     Unlike  the  corresponding 
salts  of  iron  and  chromium,  manganic  salts  very  readily 
lose  oxygen,  &c.,  and  become  converted  into  manganous 
salts.     For  example, 

Mn2(SO4)3  +  H20  =  2MnS04  +  O  +  HaSO4. 

364.  ManganateS-     The  manganates  are  similar  to 
the  sulphates  and  chromates  in  composition,  and  are  also 
in  many  cases  isomorphous  with  these  salts.     The  most 
important  manganates  are  potassic  (K2MnO4),  and  sodic 
(Na2MnO4). 

PREPARATION. — Experiment  297. — Fuse  manganese  dioxide 


344  PERMANGANATES. 

in  a  porcelain  basin  with  one  and  a  half  times  its  weight  of  caustic 
potash,  and  stir  for  some  time  with  a  glass  rod  so  as  to  expose 
the  mass  to  the  oxidising  action  of  the  air  : 

2MnO2  +  O2  +  4KOH  =  2K2MnO4  +  2H2O. 

Lixiviate  the  blue  mass  with  water.  A  solution  of  potassic  man- 
ganate  is  obtained.  (Keep  the  solution.)  By  using  caustic  soda 
a  solution  of  sodic  manganate  can  be  prepared  in  the  same  way. 

PROPERTIES. — The  magnanates  of  potassium  and 
sodium  form  dark  green  solutions  from  which  the  solid 
substances  can  be  obtained  by  evaporation.  They  are 
stable  only  in  the  presence  of  free  alkali.  (See  Exp't  298.) 

Condys  green  disinfecting  fluid  is  an  alkaline  solution 
of  sodic  manganate,  generally  containing  permanganate 
in  small  proportion.  The  manganates  are  powerful 
oxidising  agents,  and  to  this  they  owe  their  valuable 
disinfecting  and  deodorising  properties. — Manganic  acid 
(H2MnO4)  is  not  known  apart  from  its  salts. 

365.  Permanganates.  These  are  salts  of  perman- 
ganic acid  (HMnO4).  The  potassium  and  sodium  salts 
are  of  most  importance. 

PKEPARATION. — Experiment  298. — To  the  clear  solution 
of  potassic  manganate  (Exp't.  297)  add  carefully  dilute  sulphuric 
acid  until  the  colour  changes  to  purple.  This  takes  place  as  soon 
as  the  free  alkali  is  neutralised  : 

3K2Mn04  +  2H2O  =  2KMn04  +  Mn02  +  4KOH. 

If  this  solution  is  evaporated,  it  deposits  first  potassic  sulphate, 
and  then  potassic  permanganate  (KMn04). 

PROPERTIES. — Potassic  permanganate  crystallises  in 
needle-shaped  crystals  of  a  dark  purple  colour,  and  a 
somewhat  steely  lustre.  It  is  soluble  in  water  (I  part 
in  15)  and  has  enormous  colouring  power. 


PERMANGANATES.  345 

The  permanganates  are  strong  oxidising  agents,  and 
are  of  great  value  as  disinfectants  and  deodorisers. 
Condy's  red  fluid  is  a  solution  of  more  or  less  pure  sodic 
permanganate.  (NaMn04). 

Experiment  299. — To  a  solution  of  potassic  permanganate 
add  sulphurous  acid.  The  colour  disappears  : 

2KMnO4  +  5802  -f  2H20  =  K3S04  +  2MnS04  +  2H2S04. 

Test  this  solution  for  sulphuric  acid. — Decolorise  acidified  solu- 
tions of  potassic  permanganate  with  other  reducing  agents,  e.g., 
ferrous  sulphate,  hydric  sulphide,  ammonic  sulphide,  &c.  (Write 
the  equations). 

Potassic  and  sodic  permanganates  oxidise  many  or- 
ganic substances,  especially  those  which  are  offensive  and 
noxious. 

Experiment  300- — Bubble  the  air  from  the  lungs  through  a 
dilute  solution  of  potassic  permanganate,  using  a  glass  tube. 
The  purple  colour  disappears  and  a  reddish  precipitate  of  hy- 
drated  dioxide  of  manganese  (MnO(OH)a)  is  thrown  down. — Re- 
peat the  experiment,  first  acidifying  the  solution  with  dilute 
sulphuric  acid.  The  colour  is  discharged,  and  no  precipitate  ap- 
pears. 

366.  Tests. 

Manganout  Salts. 

1.  Ammonia  gives  a  white  precipitate  of  manganous  hydroxide 
(Mn(OH2)),  soluble  in  excess. 

2.  Ammonic  sulphide  gives  a  salmon-coloured  precipitate  of 
manganous  sulphide  (MnS),  soluble  in  dilute  hydrochloric  acid. 

3.  Caustic    soda    gives    a    white    precipitate    of    manganous 
hydroxide,  insoluble  in  excess,  and  oxidising  to  brown  manganic 
hydroxide  (Mnu02(OH)2)  when  shaken  up  with  air  : 

2MrtOHJ9  +  O  =  Mn202(OH)2  +  H2O. 

4.  Insoluble  compounds  can  be  tested  by  heating  with  a  sodic 


346  COBALT. 

carbonate  bead  in  the  oxidising  zone  of  the  Bunsen  flame.  The 
bead  is  coloured  green  by  the  formation  of  sodic  manganate. 
The  borax  bead  is  coloured  amethyst  by  manganese  compounds. 

Manganates  and  Permanganates. 

1.  ffydric  sulphide  reduces  them,  and  precipitates  manganous 
sulphide  if  the  solution  is  alkaline. 

2.  Ammonic  sulphide  reduces  them,  and  precipitates  manganous 
sulphide. 


COBALT. 

367.  Cobalt  (Co  u-iv-  =  58.6.—  Sp.    wt.    =  =   8.5.— 
Melting  pt.  =   1100°.— Sp.  heat  =  0.10696.)— Cobalt 
occurs  combined  with  nickel,  iron,  arsenic,  and  sulphur.  It 
is  always  accompanied    by   nickel.      It  is  also  found  in 
meteoric  iron,  and  is  present  in  the  atmosphere  of  the 
sun.     It  is  an  unimportant  metal.     Cobalt  ores  are  used 
chiefly  in  the   manufacture  of  smalt,  a   powdered  blue 
glass  used  as  a  paint. 

368.  Oxides  Of  Cobalt.     Cobalt  forms  three  oxides, 
parallel  with  those  of  iron.     They  are  cobaltous  (CoO), 
cobaltic  (Co2O3),  and    cobaltoso-cobaltic   oxide   (Co3O4,  or 
CoO.Co2O3).     They  are  all  stable. 

369.  Salts  Of  Cobalt.     There   are   two   series   of 
salts,  cobaltous  and  cobaltic.    The  latter  are  very  unstable, 
being   readily    reduced  to  the    former. — Cobaltous  salts 
(CoCl2,  CoSO4,  Co(NO3)2,  &c.)  are  violet  or  blue  when 
anhydrous,  but  rose  coloured  when  hydrated. 

370.  Cobaltous    Nitrate   (Co(NO3V6H20)  is  a 
reddish,  crystalline  salt,  prepared  by  dissolving  cobaltoso- 
cobaltic  oxide  or  cobaltous  carbonate  in  nitric  acid,  and 


NICKEL.  347 

evaporating  the  solution.  It  is  freely  soluble  in  water, 
and  the  solution  has  an  acid,  astringent  taste. — Cobaltous 
nitrate  is  used  in  testing  substances  by  means  of  the 
blowpipe.  (Art.  350,  4).  It  is  decomposed  by  a  strong 

heat : 

3Co(N03)2  =  Co304  +  6NOa  +  O2. 

371.  OobaltOUS   Chloride  (CoCl2.6H2O)  is  a  rose- 
coloured  salt  prepared  by  the  same  method  as  that  used 
for  the  nitrate.     Its  solution  has  an  acid  reaction,  and 
is  used  as  a  sympathetic,  or  invisible  ink. 

Experiment  301. — Write  with  a  solution  of  cobaltous  chloride 
so  dilute  that  the  writing  is  invisible  when  dry.  Hold  the 
paper  near  a  flame.  The  writing  appears  in  blue  characters. 
(Explain.) 

372.  Tests. 

1 .  A  mmonia  gives  a  blue  precipitate  (basic  salt)  soluble  in  ex- 
cess to  a  brownish  solution. 

2.  Ammonic  sulphide   gives   a  black   precipitate  of  cobaltous 
sulphide  (CoS),*  insoluble  in  cold  dilute  hydrochloric  acid,  but 
soluble  in  aqua  regia. 

3.  Caustic  soda  gives  a  blue  precipitate    insoluble  in  excess 
and  turning  reddish  when  shaken  up  with  air. 

4.  Insoluble  compounds  are  tested  by  the  borax  bead,  to  which 
oobalt  gives  a  deep  blue  colour. 


NICKEL. 


373.  Nickel  (Niu- 1T-  =  58.6.     Specific  weight  =  8.9. 
Melting  point  =  1500°.     Specific  heat  =  0.10863). 
Nickel    is  found   generally  along  with    cobalt.     The 

*  In  reality  the  hydro-sulphide,  Co(SH).,. 


348  SALTS    OF    NICKEL. 

principal  ore  is  kupfer-nickel  (NiA.s).      Nickel  ores  gen- 
erally contain  cobalt,  copper,  iron,  and  arsenic. 

It  is  a  white  metal,  somewhat  like  steel  in  appearance, 
hard,  ductile,  and  not  easily  oxidised  by  air.  It  decom- 
poses steam  slowly  at  a  red  heat,  and  is  soluble  in  dilute 
acids.  On  account  of  its  permanence  in  the  atmosphere 
it  is  used  for  electroplating  other  metals.  —  Nickel  coins 
are  made  of  an  alloy  of  75  parts  of  copper  with  25  of 
nickel.  This  alloy  is  hard  and  wears  well.  —  German 
silver  is  an  alloy  of  copper,  nickel,  and  zinc  in  various 
proportions.  It  is  harder  than  copper,  but  more  easily 
attacked  by  acids. 

374.  Oxides  of  Nickel.  Nickel  forms  two  oxides, 
nickelous  (NiO),  and  nickelic  (Ni2O3).  Nickelous  oxide 
is  basic,  while  nickelic  is  indifferent.  (Compare  with 
iron,  &c.)  The  former  occurs  in  nature  as  Bunsenite. 
It  can  be  prepared  by  strongly  igniting  nickel  nitrate 


Ni(NO3)2  =  NiO  +  2N02  +  0. 

It  is  a  green  powder,  permanent  in  air,  and  soluble  in 
acids. 

375.  Salts  Of  Nickel.  —  There  is  only  one  series  of 
nickel  salts,  and  they  are  derived  from  the  basic  oxide, 
NiO.     They  are  mostly  green  in  colour,  but  are  yellow 
when  deprived  of  water  of  crystallisation.     The  noluble 
salts  (NiCl2,  NiSO4,  Ni(NO3)2,  &c.)  form  acid  solutions  of 
an  astringent  taste.     They  are  prepared  usually  by  dis- 
solving the  metal  in  dilute  acids. 

376.  Nickel  Sulphate  (NiSO4.7H80)   is   a  green 


CERIUM.  349 

salt,  isomorphous  with  Epsom  salts,  prepared  by  dissolving 
nickel  in  dilute  sulphuric  acid  : 

Ni  +  H2SO4  =  NiS04  +  H2. 

It  is  soluble  in  water  (2  parts  in  5).  It  combines  with 
ammonic  sulphate  to  form  ammonio-nickel  sulphate 
(NiSO4.(NH4)2SO4.6H2O),  used  in  nickel-plating.  In 
this  compound,  as  in  Mohr's  salt,  and  other  double  sul- 
phates of  the  same  class,  a  molecule  of  alkaline  sulphate 
replaces  one  of  the  seven  molecules  of  water  of  crystal- 
lisation. 

377.  Tests. 

1.  Ammonia  gives  a  green  precipitate  (Ni(OH)2),  soluble  in 
excess  of  ammonia  to  a  blue  solution. 

2.  Ammonic  sulphide  gives  a  black  precipitate  of  nickel  hydro- 
sulphide  (Ni(SH)2),  insoluble  in  cold  dilute  hydrochloric  acid. 
This  precipitate  is  somewhat  soluble  in  excess  of  ammonic  sul- 
phide, which  should  therefore  be  sparingly  used  in  precipitating 
solutions  containing  nickel.     It  is  better  to  use  freshly  prepared 
sulphide  free  from  excess  of  sulphur. 

3.  Caustic  soda  gives  a  green  precipitate  of  nickeloits  hydroxide 
(Ni(OH)2)  insoluble  in  excess. 

4.  Insoluble  nickel  compounds  are  tested  by  the  borax  bead, 
to  which  they  give  a  reddish-brown  tint.     The  test  is,  however, 
easily  obscured  by  the  presence  of  other  metals. 

378.  Cerium  (Ce  '"•  =  14 1 .2).— This  is  the  most  impor- 
tant of  the  Cerite  metals,  a  group  of  rare  metals,  allied 
to  the  group  under  consideration.     They  have  recently 
been  discovered  in  large  quantities  and  may  assume  con- 
siderable   practical    importance    in    medicine.     Cerium 
forms  two   oxides  (Ce203  and  CeO2),  both   basic.     The 
cerous  salts  are  of  most  importance  in  medicine.    'Gerous 


350  QUESTIONS    AND    EXERCISES. 

nitrate  (Ce(NO31)3.6H2O),  and  cerous  oxalate  (Ge2'C2O4)3) 
are  used.  The  nitrate  is  easily  soluble  in  water,  the 
oxalate  very  sparingly.  They  are  both  colourless. 


QUESTIONS  AND  EXERCISES. 

1.  In  what  experiments  already  made  has  green  vitriol  been 
one  product  ? 

2.  What  weight  of  70  %  nitric  acid  is  required  to  oxidise  1  lb. 
green  vitriol  to  ferric  sulphate  ?     How  much  pure  sulphuric  acid 
must  be  added  ? 

3.  Why  would  you  expect  ferrous  carbonate  to  be  less  irritat- 
ing to  the  stomach  than  ferrous  sulphate  ? 

4.  How  much  iron  in  one-eighth  of  a  grain  of  ferric  arsenate  ? 
To  what  weight  of  ferric  chloride  is  it  equivalent  ? 

5.  What  causes  the  red   fumes  in  the  preparation   of  ferric 
chloride  ? 

6.  Why  must  ferric  hydroxide  be  freshly  precipitated  when 
used  as  an  antidote  to  arsenic  ? 

7.  In  what  respects  are  iron  and  chromium  alike  ?    In  what 
respects  do  they  differ  ? 

8.  The  third  equation  in  Art.  339  represents  calcic  sulphate 
(CaS04)  as  being  precipitated  from  aqueous  solution.     Is  this 
strictly  correct  ? 

9.  Is  there  any  resemblance  between  chromium  and  sulphur  ? 

10.  Chrome  yellow  dissolves  in  solution  of  caustic   potash. 
What  substances  are  formed  ? 

11.  Does  chrome  alum  contain  aluminum?     Why  is  it  called 
"  alum"  ?     Write  the  formulas  of  all  the  alums  which  have  been 
mentioned. 

1 2.  Why  cannot  aluminium  be  used  for  making  cooking  uten- 
sils ? 

13.  How  would  you  distinguish  aluminium  bronze  from  gold  ? 

14.  Write  the  formula,  for  c<emim  alum  ? 


METALS   OP    GROUP    IV.  351 

15.  What  is  the  composition  of  burnt  alum  ? 

16.  Can  alums  be  represented  by  a  simpler  formula  than  that 
given  in  Art.  347  ? 

17.  Common  alum  is  a  good  antidote  to  lead  salts.     Explain. 

18.  How  would  you  distinguish  a  solution  of  aluminic  from 
one  of  zinc  sulphate  ? 

19.  What  is  zinc  white  ? 

20.  What  resemblance  is  there  between  manganese  and  sul- 
phur ? 

21.  Calculate  the  percentage  of  oxygen  in  potassic  permanga- 
nate.    What  weight  of  potassic  permanganate  will  oxidise  the 
sulphurous  acid  obtained  by  burning  10  g.  of  sulphur  ? 

22.  Explain  the   explosive   nature  of  a  mixture  of   charcoal, 
manganese  dioxide,  and  potassic  chlorate  ? 

23.  Write  a  short  essay  on  the  general  resemblances  and  dif- 
ferences among  the  metals  of  Group  III.  with  regard  to  the  salts 
which  they  form. 

24.  Explain  the  antiseptic  properties  of  a  solution  of  potassic 
permanganate.     Will  it  disinfect  the  atmosphere  ? 

25.  How  would  you  prepare  cobaltous  nitrate  from  cobaltous 
sulphate  ? 

26.  What  differences  have  you  noticed  in  the  chemical  charac- 
ters of  nickel  and  cobalt  ? 


CHAPTER   XX. 


METALS  OF  GROUPS  IV.  AND  V. 

Calcium,  Strontium,  Barium  ;  Magnesium. 

379.  General  Characters. — Calcium,   strontium, 

barium,  and   magnesium  are  the  metals  of  the  alkaline 
earths.     Formerly,  any  insoluble  earthy  substance  which 


352  METALS    OF    GROUP    IV. 

remained  unchanged  when  heated  was  called  an  earth;  so 
that  substances  such  as  liine,  silica,  phosphates,  <fcc., 
were  classed  under  this  term.  Then  it  was  noticed  that 
some  of  these  earths,  viz.,  lime,  strontia,  baryta,  and  mag- 
nesia were  somewhat  soluble  in  water,  and  had  alkaline 
properties.  They  were  therefore  called  alkaline  earths, 
and  were  held  to  be  elements  until  metals  were  obtained 
from  them.  The  metals  have  very  strong  chemism  and 
are  difficult  to  separate  from  their  compounds.  Cal- 
cium, strontium,  and  barium  are  prepared  by  the  elec- 
trolysis of  their  fused  chlorides  or  cyanides.  Magne- 
sium can  be  prepared  in  the  same  way,  but  is  prepared 
on  the  large  scale  by  reducing  its  chloride  by  means 
of  sodium.  These  metals  oxidise  so  readily  that  they 
must  be  protected  from  the  action  of  the  air  by  naph- 
tha. They  are  bright,  easily  fusible,  and  decompose 
water  at  ordinary  temperatures.  They  are  all  dyad,  and 
each  has  only  one  basic  oxide  (CaO,  SrO,  BaO,  MgO). 
These  oxides  unite  directly  with  water,  forming  spar- 
ingly soluble  hydroxides  (Ba(OH)2,  Sr(OH)2,  CavOH)2, 
and  Mg(OH)2 ; — in  the  order  of  their  solubility  begin- 
ning with  the  most  soluble).  The  sulphides  (CaS,  SrS, 
BaS,  MgS)  are  decomposed  by  water,  and  cannot  be  pre- 
pared by  precipitation.  They  are  prepared  by  heating 
the  sulphates  with  charcoal,  e.g.  : 

CaSO4  +  4C  =  CaS  +  4CO. 

The  carbonates  (CaCO3,  SrCO3,  BaC03,  MgC03),  sulphates 
(MgSO4,  CaSO4,  SrSO4,  BaSO4,  in  the  order  of  their 
solubility  : — magnesic  sulphate  is  freely  soluble.),  and 
phosphates  (Ca3(PO4)2,  &c.)  are  insoluble  (or  sparingly 
soluble)  in  water.  The  chlorides,  nitrates,  &c.,  are 
soluble. 


LIME.  353 

CALCIUM. 

380.  Calcium  (Caij-  =  40). 

OCCURRENCE.  —  Always  in  combination.  Its  com- 
pounds occur  in  vast  quantities,  and  include  limestone, 
marble,  chalk,  coral,  dolomite,  gypsum,  apatite,  &c.  It 
is  present  in  natural  waters,  in  the  bodies  of  plants  and 
animals,  and  in  the  sun  and  fixed  stars. 

381.  Calcic  Oxide  (CaO). — Calcium  combines  with 
oxygen  in  two  proportions  (CaO,   CaO2),  but  the  mon- 
oxide (CaO)  is  the  only  important  oxide. 

PREPARATION. — It  is  prepared  on  the  large  scale  by 
strongly  heating  limestone.  Impure  calcic  oxide  (quick 
lime)  is  obtained  : 

CaCO3  =  CaO  +  C03. 

Experiment  302. — Wrap  several  times  around  a  small  frag- 
ment of  calc  spar  a  platinum  wire,  one  end  of  which  is  fastened 
in  a  handle  of  glass  (by  fusing  the  glass  and  sticking  the  wire 
into  it).  Thrust  the  calc  spar  into  the  centre  of  a  Bunsen  flame, 
and  hold  it  for  two  or  three  minutes  just  above  the  point  of  the 
central  bluish  green  zone.  Remove  it  and  observe  the  change  in 
its  appearance.  Put  a  drop  of  water  on  it,  and  observe  the 
change.  Wash  it  into  at.  t.,  shake  it  up  with  water,  and  add  a 
little  red  litmus.  See  whether  calc  spar  affects  red  litmus. 

PROPERTIES. — Calcic  oxide  is  a  white  solid,  not  fused 
by  the  intense  heat  of  the  oxy -hydrogen  flame. 

Experiment  303. — Place  a  lump  of  good  quick  lime  on  a 
clean  iron  or  porcelain  plate,  and  pour  over  it  one-third,  its 
weight  of  water.  Note  the  heating  of  the  mass.  Whence 
comes  the  heat,  ?  This  process  is  called  the  slaking  or  slacking  of 
lime  : 

CaO  +  H2O  =  Ca(OH)2. 

(Keep  the  slaked  lime  for  Experiment  305). 
24 


354  SLAKED    LIME. 

Experiment  304.  — Leave  a  small  lump  of  quick  lime  exposed 
to  the  air  for  two  or  three  days.  Note  any  changes  in  its  ap- 
pearance. Test  it  for  carbonic  acid. 

Quick  lime  absorbs  moisture  and  carbon  dioxide  from 
the  air  and  becomes  changed  at  length  to  calcic  carbon- 
ate. Quick  lime  is  often  used  to  keep  the  air  of  an 
apartment  dry  and  pure. 

382  Calcic   Hydroxide   (Ca(OH)2),  also    called 

slaked  lime,  is  prepared  as  in  Experiment  303.  Much 
heat  is  given  out  during  the  combination  of  the  lime  and 
the  water.  Fires  have  been  caused  by  the  accidental 
slaking  of  large  quantities  of  quick  lime. 

Experiment  305. — Put  about  equal  quantities  of  recently 
slaked  lime  in  two  test  tubes  ;  to  the  one  add  about  100  parts  of 
water,  and  to  the  other  100  parts  of  water  and  2  of  sugar. 
Shake  both  for  some  time  and  observe  that  the  sweetened  water 
dissolves  much  more  slaked  lime  than  the  pure  water  does. 
Filter  off  the  solution  in  pure  water,  and  try  its  taste  and  action 
on  red  litmus. 

Calcic  hydroxide  dissolves  in  water  to  the  extent  of  1 
part  in  700  of  water.  The  solution  is  called  lime  water. 
Calcic  hydroxide  is  more  soluble  in  cold  than  in  hot 
water.  (Heat  a  little  of  the  filtered  solution  prepared  in 
Experiment  305.)  Sugar  increases  the  solubility  of  lime 
in  water.  Saccharated  solution  of  lime  contains  about 
1  part  by  weight  of  calcic  hydroxide  and  2  of  cane  sugar 
dissolved  in  20  of  water. 

Experiment  306. — Leave  a  few  cubic  centimetres  of  lime- 
water  in  an  open  vessel  for  two  or  three  days,  stirring  it  from 
time  to  time.  Then,  try  its  action  on  red  litmus.  Pour  off  the 
liquid,  add  a  few  drops  of  hydrochloric  acid  to  the  sediment, 
and  observe  the  result.  What  is  the  sediment  ? 


CALCIC    CARBONATK.  355 

Lime  water  is  used  in  medicine  as  an  antacid,  &c.,  and 
as  an  antidote  to  poisoning  by  acids,  particularly  oxalic. 
— Milk  of  lime  is  water  shaken  up  with  more  sLiked 
lime  than  it  can  dissolve. 

383.  Calcic  Carbonate  (CaCO3). 

OCCURRENCE. — Gale  spar  is  nearly  pure  crystallised 
calcic  carbonate.  It  is  colourless  and  transparent. 
Aragonite  is  another  crystalline  form  of  calcic  carbonate. 
This  compound  is  therefore  dimorphous;  it  crystallises 
in  two  forms.  Calcic  carbonate  is  found  more  or  less 
pure,  as  marble,  limestone,  chalk,  coral,  &c.  It  forms  a 
considerable  part  of  the  mass  of  egg-shells  and  of  the 
shells  of  molluscs. 

PREPARATION. — Calcic  carbonate  occurs  so  abundantly 
in  nature  that  it  is  not  necessary  for  most  purposes  to 
prepare  it  artificially.  In  medicine,  however,  it  is  pre- 
pared by  precipitation  (calcis  carbonas  precipitata)  in 
order  to  obtain  it  in  a  fine  state  of  division. 

Experiment  307.  —  To  a  hot  solution  of  calcic  chloride 
(CaCl2),  add  solution  of  sodic  carbonate  until  no  more  precipitate 
is  formed  on  further  addition  of  the  carbonate.  Filter  and  wash 
the  precipitate  : 

Na2C03  +  CaCl2  =  2NaCl  +  CaC03. 

PROPERTIES. — Precipitated  calcic  carbonate  is  a  white 
powder,  slightly  granular,  insoluble  in  water,  soluble 
with  effervescence  in  hydrochloric  and  other  acids  (Art. 
152).  Its  solubility  in  water  containing  carbonic  acid 
has  been  already  referred  to  (Art.  153).  It  is  used  as 
an  antidote  to  poisoning  by  acids,  and  also  to  correct 
acidity  of  the  stomach. 


356  CALCIC    SULPHATE. 

384.  Calcic  Chloride  (Ca012.6H2O). 

PREPARATION. — Experiment  308. — Dissolve  in  a  porcelain 
dish  a  few  fragments  of  calc  spar  in  hydrochloric  acid,  and 
evaporate  the  solution  on  the  water  bath  until  it  forms  a  pellicle 
on  the  surface.  Then  heat  it  over  the  Bunsen  burner,  placing 
the  dish  so  far  above  the  flame  that  it  may  not  be  touched  by  it. 
When  the  salt  is  quite  dry,  apply  the  flame  to  the  dish  until 
moisture  ceases  to  come  off.  Leave  this  fused  calcic  chloride 
(CaCl2)  exposed  to  the  air  for  24  hours,  and  observe  the  result. 

PROPERTIES. — Calcic  chloride  is  a  colourless  salt  very 
soluble  in  water.  The  fused  chloride  contains  no  water 
of  crystallisation.  It  is  very  hygroscopic,  and  is  used 
for  drying  gases.  Solution  of  calcic  chloride  is  used  as 
a  test  reagent,  and  also  as  a  medicine. 

385.  Calcic  Sulphate  (CaS04). 

OCCURRENCE. — Crystallised  as  gypsum,  alabaster,  or 
selenite  (CaSO4.2H2O),  and  anhydrite  (CaSO4).  It  is 
present  in  sea  and  other  waters. 

PREPARATION. — Experiment  309. — To  solution  of  calcic 
chloride  add  dilute  sulphuric  acid,  filter,  and  wash  the  pre- 
cipitate with  hot  water.  (Write  the  equation.)  The  precipitate 
is  gypsum. 

PROPERTIES. — Gypsum  is  a  white  solid  of  sp.  wt.  2.3. 

Experiment  310. — Heat  a  small  quantity  of  gypsum  in  a  t.  t. 
and  observe.  Heat  a  larger  quantity  in  a  porcelain  basin  until 
moisture  ceases  to  escape.  Let  cool,  and  mix  the  anhydrous 
salt  with  water  on  a  piece  of  broken  porcelain  or  smooth  wood. 
(What  proportion  of  water  should  be  used  to  form  gypsum  ?) 
Observe  change  after  a  few  minutes. 

Anhydrous  calcic  sulphate  is  called  plaster  of  Paris, 
or  sometimes  simply  plaster.  It  is  used  as  a  cement,  for 
making  plaster  casts,  &c.,  and  in  surgery  for  making  stiff 


BLEACHING    POWDER.  357 

bandages.  -  Gypsum  is  less  soluble  in  hot  than  in  warm 
water.  It  is  most  soluble  at  35°C. 

Experiment  311. — Heat  a  saturated  solution  of  gypsum  to 
boiling. 

One  part  of  gypsum  requires  500  of  water  to  dissolve 
it. 

386.  Bleaching  Powder.  ("  Chloride  of  Lime") 
This  substance  has  been  already  referred  to  at  p.  105. 
It  is  probably  a  mixture  of  calcic  hypochlorite  and 
chloride  in  molecular  proportions,  Ca  OC1)2  -f-  CaCl2 ; 
but  it  always  contains  calcic  hydroxide. 

PREPARATION. — Slaked  lime  is  spread  on  a  series  of 
shelves  in  large  closed  compartments,  and  exposed  to  the 
action  of  chlorine  gas.  The  chlorine  is  genei-ated  by  the 
action  of  crude  hydrochloric  acid  on  manganese  dioxide  : 

4HC1  +  Mn02  =  MnCl2  +  C12  +  2H2O. 

Its  action  on  the  slaked  lime  may  be  represented  as 
follows  : 

2Ca(OH)2  +  2C12  =  Ca(OCl)2.CaCl2  +  2H2O. 

PROPERTIES. — A  white  powder,  having  a  chlorous 
smell.  The  odour  is  due  to  the  liberation  of  hypochlorous 
acid  by  the  action  of  the  carbon  dioxide  and  moisture  of 
the  air : 

Ca(OCl)2  +  H2O  +  CO2  =  CaCO3  +  2HC1O. 

It  is  one  of  the  best  known  deodorisers  and  disinfectants, 
and  is  especially  useful  because  of  the  volatility  of  hypo- 
chlorous  acid,  a  substance  which  deodorises  and  disinfects 
the  atmosphere  into  which  it  escapes. 

Experiment  312. — Examine  a  specimen  of  bleaching  powder, 


358  CALCIC    PHOSPHATE. 

noting  its  colour  and  odour.  Shake  up  about  10  grains  of  it  with 
50  to  100  cubic  centimetres  of  water  and  filter.  The  filtrate 
contains  calcic  chloride  and  hypochlorite.  It  is  the  Liquor  calcis 
chloratce  of  the  pharmacopoeia.  Add  a  few  drops  of  it  to  a 
solution  of  sulphate  of  indigo.  Try  with  acidified  litmus  solution, 
and  with  potassic  bichromate  solution. 

Experiment  313. — Pour  some  dilute  sulphuric  and  hydro5 
chloric  acid  over  small  portions  of  bleaching  powder  in  watch 
glasses,  and  note  evolution  of  chlorine  : 

CaCl2  +  Ca(OCl)a  +  2H2S04  =  2CaS04  -f  2C12  -f  2H..O. 
CaCl2  +  Ca(OCl)2  +  2HC1  =  2Ca012  +  2C12  -f  2H20. 

When  bleaching  powder  is  heated,  the  hypochlorite  is  decom- 
posed into  chlorate  and  chloride  : 

6Ca(OCl)2  =  Ca(C103)2  +  5CaCl2. 

387.  Calcic  Phosphate  (Ca3(PO4)2).  This  com- 
pound has  been  already  studied.  (Arts.  128  and  132.) 

OCCURRENCE.' — It  is  found  in  the  minerals  apatite 
(3Ca3(PO4),  +  Ca[Cl2,  F2]*),  phosphorite  (Ca3(POJ2), 
sombrerite  (Ca3(PO4)2.2H2O),  &c.,  and  in  bones.  Bone 
ash  contains  about  80  °/0  of  calcic  phosphate. 

PREPARATION. — Experiment  314.— Digest  bone  ash  with 
dilute  hydrochloric  acid,  filter,  dilute  to  double  volume,  add 
ammonia  to  alkaline  reaction,  filter,  and  wash  the  precipitate 
with  hot  distilled  water. 

The  hydrochloric  acid  acts  on  the  phosphate  to  form 
calcic  chloride  and  calcic  tetrahydric  phosphate : 

Ca3(P04)2  +  4HC1  =  CaH4(P04)2  +  2CaCl2. 

V 

The  ammonia  reprecipitates  the  calcic  phosphate  : 
CaH4(P04)2  +  2CaCl2  +  4NH3  =  Ca3(PO4)2  +  4NH4C1. 

*  When  symbols  are  put  within  square  brackets  and  separated  by  a  comma, 
it  signifies  that  the  elements  are  present  in  varying  proportions. 


MORTARS  AND  CEMENTS.  359 

It  is  plain  that  any  impurities  in  the  bone  ash  which  are 
soluble  in  hydrochloric  acid  may  be  precipitated  along 
with  calcic  phosphate. 

PROPERTIES. — A  light,  white  powder,  insoluble  in 
water,  soluble  in  dilute  hydrochloric,  nitric,  or  phosphoric 
acid.  From  this  solution  calcic  phosphate  is  again  pre- 
cipitated when  an  alkali  is  added.  (Experiment  314.) 
Thus,  an  acid  solution  of  calcic  phosphate  might  be  mis- 
taken for  an  aluminium  salt;  but  the  former  gives  a 
white  precipitate  (calcic  oxalate)  on  the  addition  of  sodic 
acetate  and  ammonic  oxalate,  while  the  latter  does  not. 

388.  Mortars  and    Cements. — Ordinary    mortar 
is  a  mixture  of  slaked  lime  (Oa(OH).,),  sand  (SiO2),  and 
water.     It  hardens  by  the  formation  of  calcic  carbonate, 
carbon  dioxide  being  absorbed  from  the  air.     The  process 
of  hardening  continues  for  years.     There  appears  to  be 
no  combination  of  the  lime  with  the  silica. — Hydraulic 
mortar,  or  Roman  cement,  is  prepared  by  carefully  heat- 
ing a  mixture  of  lime  and  clay.     It  hardens  under  water, 
and  the  hardening  seems  to  be  due  to  the  formation  of 
silicate  of  lime  and  alumina.     Portland  cement  is  made 
by   carefully   heating  a  levigated   mixture  of  clay  and 
chalk.     These  cements  deteriorate  when  kept  exposed  to 
the  air,  because  the  lime  in  them  unites   with  carbon 
dioxide. 

389.  Tests. 

1.  Solutions  of  calcium  salts  give  a  white  precipitate   with 
ammonic  carbonate.     If  the  solution  is  acid  it  must  be  neutralised 
with  ammonia.     The  precipitate  is  soluble  with  effervescence  in 
nitric  acid. 

2.  Sodic  phosphate  gives  a  white  precipitate  with  neutral  or 


360  STRONTIUM. 

alkaline  solutions  of  calcic  salts  (in  practice  the  solution  is  made 
alkaline  with  ammonia)  : 

3CaCl2  +  2Na2HP04  +  2NH3  = 

Ca3(PO4)2  +  4NaCl  +  2NH4C1. 

3.  Ammonic  oxalate  gives  a  white  precipitate  of  calcic  oxalate 
even  with  very  dilute  solutions  of  calcium  salts  : 

CaCl2  +  (NH4)2C204  =  CaC204  +  2NH4C1. 

This  precipitate  is  soluble  in  dilute  nitric  or  hydrochloric  acid, 
but  is  insoluble  in  acetic  acid. 

4.  Calcium  compounds  insoluble  in  water  can  be  tested  by 
dissolving  in  hydrochloric  acid,  adding  sodic  acetate  (so  that  the 
acidity  may  be  due  to  acetic  acid),  and  then  adding  ammonic 
oxalate.     Or,    they   may   be   tested   by  moistening   them   with 
strong  hydrochloric  acid  and  bringing  them  by  means  of  a  plati- 
num wire  into  the  Bunsen  flame.    A  brick  red  colour  is  imparted 
to  the  flame  by  calcium  compounds. 


STEONTIUM. 

390.  Strontium  (Sr11-  =  87.2).  —  Compounds  of 
strontium  are  found  in  considerable  quantities  in  nature. 
The  most  commonly  occurring  are  celestine  (SrSO4),  and 
strontianite  (SrCO3).  The  metal  is  prepared  by  electro- 
lysis.— Strontium  compounds  have  not  yet  found  a  place 
in  the  pharmacopoeia.  The  nitrate  (Sr(NO3)2)  and  other 
salts  are  used  in  making  fire-works.  They  give  a  fine 
red  colour  to  flames. 

Experiment  315. — Mix  a  little  strontic  nitrate  with  powdered 
charcoal  and  sulphur,  put  the  mixture  on  a  piece  of  mica  or  por- 
celain and  touch  it  with  a  hot  wire.  (What  causes  the  rapid 
combustion  ?) 

The  soluble  salts  are  prepared  by  dissolving  strontia- 
nite in  the  acids,  or  by  reducing  celestine  to  strontic  sul- 


BARIUM.  361 

phide  (SrS),  and  dissolving  this  in  the  respective  acids. 
Strontic  oxide  (SrO),  and  hydroxide  (Sr(OH)2)  are  simi- 
lar to  the  corresponding  calcium  compounds,  but  the 
hydroxide  is  more  soluble  than  calcic  hydroxide.  Stron- 
tium compounds  are  mostly  colourless.  The  sulphate 
(SrSO4)  is  less  soluble  than  calcic,  but  more  so  than 
baric,  sulphate. 

391.  Tests. 

1.  Am/nonic  carbonate  gives  a  white  precipitate  (Sr003)  with 
neutral  or  alkaline  solutions  of  strontium  salts.     The  precipitate 
is  soluble  with  effervescence  in  dilute  nitric  acid. 

2.  Sodic  phosphate  (with  ammonia)  gives  a  white  precipitate 
(Sr3(P04)2). 

3.  Ammonic   oxalate   gives   a  white   precipitate    of    strontic 
oxalate  (SrC204),  soluble  in  nitric  and  sparingly  soluble  in  acetic 
acid. 

4.  Calcic  sulphate  gives  a  white  precipitate  (SrS04)  after  some 
time : 

CaS04  +  SrCl,  =  SrS04  +  CaCla. 

5.  Insoluble  strontium  compounds  can  be  tested  by  moistening 
with  strong  hydrochloric  acid  and  bringing  on  a  platinum  wire 
into  the  Bunsen  flame.     A  carmine  red  colour  is  imparted  to  the 
flame.     The  test  is  more  delicate  if  the  substance  be  held  for  a 
few  moments  in  the  reducing  flame,  then  moistened  with  hydro- 
chloric acid,  and  held  in  the  oxidising  flame. 


BARIUM. 

392.  Barium  (Ba11-  =  136.8).— The  principal  com- 
pounds of  barium  found  in  nature  are  witherite  (BaC03), 
heavy  spar  (BaSO4),  and  psilotnelane  ([Mn,  Ba]O.MnO2). 
The  metal  can  be  prepared  by  electrolysis  of  the  fused 
chloride  or  cyanide. 


362  BARIUM  CHLORIDE. 

393.  Oxides  Of  Barium. — Barium,    like    calcium 
and  strontium,  unites  with  oxygen  in  two  proportions, 
forming  the  monoxide  (BaO),  and  the  dioxide  (BaO2). 

1.  Bariiim   monoxide  (BaO)   is  prepared  by  heating 
the  nitrate  : 

Ba(N03)2  =  BaO  +  2NO2  +  O. 

It  unites  with  water,  and  forms  barium  hydroxide,  or 
baryta  (Ba(OH)2).  This  compound  is  more  soluble  than 
either  calcic  or  strontic  hydroxide.  Baryta  water,  used 
in  analysis,  is  a  solution  of  barium  hydroxide.  The 
hydroxide  is  prepared  on  the  large  scale  from  heavy  spar, 
for  use  in  sugar  refining.  It  forms  an  insoluble  com- 
pound with  cane  sugar. 

2.  Barium  dioxide  (BaOa)  is  a  white  solid  prepared  by 
heating  the  monoxide  to  a  red  heat  in  a  current  of  oxygen. 
It  is  used  in  the  preparation  of  hydrogen  dioxide. 

394.  Baric  Chloride  (BaCL,2H20). 

PREPARATION. — By  dissolving  baric  carbonate  or  sul- 
phide in  hydrochloric  acid  : 

BaCO3  +  2HC1  =  BaCl2  +  HaO  +  CO2. 

It  is  also  prepared  from  heavy  spar,  which  is  first  reduced 
to  sulphide  (BaS),  and  then  decomposed  by  hydrochloric 
acid  : 

BaSO4  +  40  =  BaS  +  4CO. 

BaS  +  2HC1  =  BaCl2  +  H2S. 

PROPERTIES. — A  white  crystalline  solid,  soluble  in 
water  (35  parts  in  100),  sparingly  soluble  in  alcohol. 

Experiment  316. — Examine  carefully  a  specimen  of  barium 
chloride,  noting  taste,  &c.  Dissolve  it  in  a  little  distilled  water, 
and  to  a  small  portion  of  the  solution  add  about  twice  the  volume 


BARIUM    NITRATE.  363 

of  strong  nitric  acid.  Baric  chloride  is  only  sparingly  soluble  in 
strong  nitric  acid.  To  another  portion  add  magnesic  sulphate. 
Baric  sulphate  is  precipitated  : 

MgS04  +  BaCl2  =  BaS04  +  MgCl2. 

Baric  chloride,  in  common  with  the  other  soluble  com- 
pounds of  barium,  is  very  poisonous.  Solution  of  mag- 
nesium sulphate  is  a  good  antidote.  (How  does  it  act1?) 

395.  Baric  Nitrate   (Ba(NO3)2).     Is   prepared    by 
the  same  methods  as  those  used  in  preparing  the  chloride. 
It  is  much  used  in  pyrotechny,  being  mixed  with  char- 
coal and  sulphur  to  form  "  green  fire." 

Experiment  317. — Mix  carefully  with  little  friction  small 
quantities  of  baric  nitrate,  sulphur,  and  ground  charcoal.  Put 
the  mixture  on  a  piece  of  mica,  a  flat  stone,  or  a  piece  of  porce- 
lain, and  touch  it  with  a  red  hot  wire. 

396.  Tests. 

1.  Ammonic  carbonate  gives   a   white    precipitate    (BaC03), 
soluble  with  effervescence  in  dilute  nitric  acid. 

2.  Sodic  phosphate  (with  ammonia)  gives  a  white  precipitate 
(Ba3(POJ2). 

3.  Ammonic  oxalate  gives  a  white  precipitate  (BaC^OJ,  soluble 
in  dilute  nitric,  sparingly  soluble  in  acetic  acid. 

4.  Calcic  sulphate  gives  at  once,  a  white  precipitate  (BaS04). 
Baric  sulphate  requires  400,000  parts  of  water  for  its  solution. 

5.  Insoluble   compounds   are  tested   for   as   in  Art.   391   (5). 
Barium  compounds  give  a  greenish  colour  to  the  flame. 


MAGNESIUM. 

397.  Magnesium.  (Mgu=23.95). 
OCCURRENCE. — Magnesium  compounds  are  very  abuu- 


364  MAGNESIUM    SULPHATE. 

dant  and  widely  distributed.  Mountain  limestone,  or 
dolomite,  ([Mg,  Ca]  C03.)  forms  whole  ranges  of  moun- 
tains. Other  commonly  occurring  compounds  are  mag- 
nesite  (MgCO3),  kieserite  (MgSO4.H2O),  carnallite  (MgCl2. 
KC1.6H2O),  spinelle  (MgO.  A12O3),  asbestos  ([Mg,Ca]SiO3), 
Epsom  salts  (MgSO4.7H2O)  &c.  Magnesic  phosphate 
(Mg3'PO4)2)  is  found  in  the  bones,  &c.,  of  animals. 

PREPARATION. — By  heating  together  a  mixture  of  mag- 
nesium chloride,  calcium  fluoride,  and  sodium.  The 
sodium  displaces  the  magnesium  : 

MgCl2  +  2Na  =  Mg  +  2NaCl. 
The  metal  is  purified  by  distillation. 

PROPERTIES.- — A  soft,  silvery  metal ;  it  tarnishes  in 
moist  air;  sp.wt.  =  1.75  ;  melts  at  red  heat,  and  boils 
above  1040°  ;  when  heated  in  air  it  catches  fire  and 
burns  with  a  dazzling  white  light.  The  magnesium  light 
is  used  in  photography,  signalling,  &c. 

Experiment  318.  Burn  a  piece  of  magnesium  wire,  dissolve 
the  white  ash  (MgO)  in  a  drop  of  dilute  sulphuric  acid,  evaporate 
on  a  watch  glass,  and  obtain  crystals  of  Epsom  salts. 

398.  Magnesic  Sulphate  (MgSO4.7H,O).— Mag- 
nesium sulphate  is  generally  sold  as  Epsom  salts,  crystal- 
lised with  seven  molecules  of  water  to  the  molecule  of 
sulphate.  It  occurs  in  nature  crystallised  with  one 
molecule  of  water,  as  kieserite  (MgS04.H2O.) 

PREPARATION. — Epsom  salts  are  now  prepared  mostly 
from  kieserite,  by  heating  in  water.  The  sparingly 
soluble  kieserite  apparently  combines  with  more  water 
and  becomes  changed  to  the  much  more  soluble  heptahy- 
drated  salt.  This  is  then  purified  by  crystallisation. 


MAGNESIUM    CARBONATE.  365 

PROPERTIES. — A  colourless,  crystalline  solid,  isomorph- 
ous  with  zinc  sulphate  (ZnS04.7H2O),  which  it  resembles 
very  closely  in  appearance ;  it  has  an  unpleasant  bitter 
taste.  It  is  soluble  in  water,  35  parts  in  100. 

Experiment  319-  Compare  specimens  of  Epsom  salts  and 
crystallised  zinc  sulphate,  noting  general  appearance  and  taste. 
Dissolve  them  separately  in  water,  and  add  ammonium  chloride 
and  a  drop  of  ammonia  to  each. 

Magnesium  sulphate  is  used  in  medicine  as  a  purgative. 
It  is  a  "  saline  purgative,"  i.e.,  it  acts  by  promoting  the 
diffusion  of  water  into  the  intestinal  canal. 

399.  Magnesic  Carbonate  (MgCO3).— Found  in 

nature  in  large  quantities,  as  magnesite.  It  was  formerly 
used  in  preparing  Epsom  salts.  The  carbonate  of  mag- 
nesia of  the  pharmacopoeia  is  a  basic  salt  (4MgO.3CO2. 
5H2O). 

PREPARATION. — By  the  action  of  solution  of  sodic  car- 
bonate on  solution  of  magnesic  sulphate. 

Experiment  320-  To  a  hot  solution  of  magnesium  sulphate 
add  hot  solution  of  sodic  carbonate.  Basic  carbonate  of  mag- 
nesium is  precipitated.  Filter,  wash  the  precipitate,  and  dry  it 
in  a  glass  or  porcelain  dish.  Note  the  appearance  of  the  dried 
salt.  It  is  the  heavy  carbonate  of  magnesia,  or  magnesia  alba 
ponderosa.  Repeat  the  experiment,  but  use  cold  solutions,  and 
boil  for  a  few  minutes  after  precipitating.  The  dried  precipitate 
is  very  light.  It  is  magnesia  alba  levis. 

PROPERTIES. — The  carbonates  of  magnesia  are  white 
powdery  solids  insoluble  in  pure  water,  but  soluble  in 
dilute  acids  (with  effervescence),  in  solutions  of  am- 
monium salts,  and  in  water  containing  carbonic  acid. 

Experiment  321. — To  solution  of  magnesic  sulphate  add 
amuionic  carbonate  ;  a  white  precipitate  falls  (unless  the  am- 


366  MAGNESIA. 

monic  carbonate  consists  mostly  of  the  acid  salt,  in  which  case  a 
little  ammonia  must  be  added).  Add  ammonic  chloride  ;  the 
precipitate  is  dissolved. 

Experiment  322-  — Put  a  little  magnesia  alba  in  a  consider- 
able quantity  of  distilled  water  in  a  flask  or  beaker,  and  pass 
carbon  dioxide  through  the  water  for  some  time.  If  the  car- 
bonate does  not  dissolve  completely,  filter.  Boil  a  little  of  the 
clear  solution  ;  it  becomes  milky  owing  to  the  precipitation  of 
magnesic  carbonate. 

The  solubility  of  magnesic  carbonate  increases  with 
the  pressure  of  the  carbon  dioxide.  Liquor  magnesice 
carbonatis  is  a  solution  made  by  using  carbon  dioxide 
under  considerable  pressure.  It  contains  about  1 3  grains 
of  "  carbonate  of  magnesia  "  to  the  fluid  ounce. 

400.  Magnesic  Oxide  (MgO). — Magnesium  forms 
only  one  oxide.  It  is  known  in  pharmacy  as  magnesia, 
or  magnesia  usta. 

PREPARATION. — By  heating  strongly  the  carbonate  of 
magnesia.  A  light  or  heavy  magnesia  is  obtained  cor- 
responding to  the  carbonate  used. 

Experiment  323. — Heat  a  small  quantity  of  magnesia  alba 
for  some  time  in  a  porcelain  crucible.  Try  the  action  of  hydro- 
chloric acid  on  the  residue.  It  dissolves  without  effervescence, 
if  the  heating  has  been  continued  long  enough: 

MgO  +  2HC1  =  MgCl2  +  HaO. 

PROPERTIES. — A  light  white  powder,  almost  insoluble 
in  water  (1  part  dissolves  in  about  55,000) ;  soluble  in 
acids,  salts  of  magnesium  being  formed  ;  tasteless  ;  when 
moistened,  turns  red  litmus  blue.  It  is  thus  a  well- 
marked  basic  oxide.  It  is  a  good  antidote  to  arsenic 
poisoning ;  it  forms  insoluble  arsenite  or  arsenate,  and 
neutralises  the  acid  of  the  gastric  juice. 


QUESTIONS    AND    EXERCISES.  367 

401.  Tests. 

1.  Ammonic  carbonate  gives  a  white   precipitate,    soluble  in 
acids,  and  in  solution  of  ammonium   chloride.     The  carbonates 
of  barium,  strontrum,  and  calcium  are  not  soluble  in  solution  of 
ammonium  chloride. 

2.  Sodic  phosphate  (with  ammonia)  gives  a  granular  white  pre- 
cipitate (Mg.NH4.P04),  which   forms  in  lines  on  the   walls  of 
the  t.  t.  when  the  solution  is  stirred  with  a  glass  rod : 

MgS04  +  NHS  -f  Na.HP04  =  Mg.NH4.P04  -f  Na2S04. 

3.  Compounds  insoluble  in  water  are  dissolved  in  hydrochloric 
acid  and  tested  by  (1)  and  (2). 


QUESTIONS  AND  EXERCISES. 

1.  In  what  way  does  calcic  carbonate  act  as  an  antidote  to 
acids  ?    To  zinc  chloride  ? 

2.  What  substances  does  quick  lime  absorb  from  the  atmos- 
phere ? 

3.  What  weight  of  water  will  10  Ibs.  of  quick  lime  absorb  ? 
What  volume  of  air  at  20°C  saturated  with  water  vapour  will 
supply  this  amount  of  water  ? 

4.  Explain  the  "  setting"  of  plaster  of  Paris. 

5.  Why  does  bleaching  powder  become  moist  when  exposed  to 
the  air  ?     Why  must  it  be  kept  in  well  closed  vessels  ? 

6.  Calculate  the  weight  of  limestone  to  furnish  1  ton  of  lime. 
What  volume  of  carbon  dioxide  would  be  driven  from  it  into  the 
atmosphere  in  the  process  of  burning  ? 

7 .  What  substances  are  formed  by  the  action  of  hydrochloric 
acid  on  barium  sulphide  (BaS)  ?     Write  the  equation. 

8.  How  would  you  distinguish  a  specimen  of  zinc  sulphate 
(white  vitriol)  from  one  of  Epsom  salts  ? 

9.  When  potassic  hydroxide  solution  is  added  to  solution  of 
magnesic  sulphate  a  white  precipitate  falls.     What  is  it  ?     Write 
the  equation, 


368  METALS    OF    THE   ALKALIS. 

10.  How  would  you  prepare  baric  nitrate  from  baric  chloride  ? 

11.  Baric  sulphate  is  often  used  by  painters  as  a  substitute  for 
white  lead.     Why  is  it  preferable  ? 

12.  What  substances  are  formed  when  solutions  of  sodic  sul- 
phate and  baric  nitrate  are  mixed  ? 


CHAPTER    XXL 


METALS  OF  GROUP  VI. 

Lithium,   Sodium,   Potassium,  Rubidium,  Ccesium, 
[Ammonium]. 

402.  General  Characters.     The  metals  of  this 

group  form  freely  soluble  hydroxides,  called  alkalis  (from 
alkali,  the  Arabic  name  for  the  plant  from  the  ashes  of 
which  potash  was  first  prepared).  The  metals  of  the 
alkalis  form  a  series,  similar  in  properties,  but  showing 
that  gradation  always  found  in  chemical  series.  Thus 
the  hydroxides  increase  in  basic  character  from  lithium 
to  caesium ;  the  oxidisability  of  the  metals  increases  in 
the  same  order,  while  their  melting  points  increase  from 
caesium  to  lithium.  The  metals  are  all  univalent,  and 
are  difficult  to  separate  from  their  compounds.  Each 
forms  only  one  basic  oxide  (Na20,  K20,  Li2O,  &c.),  which 
can  be  obtained  by  burning  the  metal  in  air.  The  oxides 
combine  with  water  to  form  hydroxides  (NaOH,  KOH, 
&c.),  from  which  the  water  cannot  be  driven  off  by  heat. 
The  salts  of  the  alkali  metals  are  mostly  soluble  in  water. 
— The  metals  will  be  treated  generally  in  the  order  of 
their  importance. 


SODIUM    CHLORIDE.  369 

SODIUM. 

403.  Sodium   (Na'  =  23.  .  Sp.   wt.  =  0.97.     Melting 
point  ^95°.6). 

OCCURRENCE. — Compounds  of  sodium  are  universally 
diffused,  so  that  it  is  very  difficult  to  get  any  substance 
free  from  them.  The  compounds  occurring  in  greatest 
abundance  are  sodic  chloride  (NaCI),  and  Chili  nitre 
(NaN03). 

PREPARATION. — By  strongly  heating  in  iron  tubes  a 
mixture  of  sodic  carbonate,  slack  coal,  and  chalk : 

Na2CO3  +  2C  =  2Na  +  SCO. 

The  sodium  distils,  is  received  in  iron  pots,  and  immedi- 
ately sealed  up  to  prevent  oxidation. 

PROPERTIES. — A  bright,  white  metal,  soft,  and  very 
easily  tarnished.  In  moist  air  it  remains  bright  only  a 
few  seconds.  It  readily  burns  in  air,  and  forms  sodium 
monoxide,  NaaO,  along  with  some  dioxide,  Na^Oj  (Exp't 
26).  It  decomposes  water,  even  in  the  form  of  ice,  the 
products  being  sodic  hydroxide  and  hydrogen  (Exp't  30): 

Na  +  H2O  -  NaOH  +  H. 

Sodium  forms  liquid  and  solid  amalgams  with  mercury. 
A  liquid  amalgam  is  used  in  extracting  gold  from  quartz. 

404.  Sodic  Chloride  (NaCl). 

OCCURRENCE. — In  masses  in  the  earth,  as  rock  salt ;  in 
sea- water  to  the  extent  of  2.6  %  ;  in  salt  springs ;  and 
in  small  quantities  very  widely  diffused. 

PREPARATION. — (1)  By  mining,  as  in  Poland,  Ger- 
many, &c.;  (2)  by  allowing  water  to  run  down  shafts  ex- 
tending into  the  salt  beds,  and  pumping  up  and  evapor- 
25 


370  SODIUM    SULPHATE. 

ating  the  solution  formed  ;  (3)  by  evaporating  the  water 
from  salt  springs;  and  (4)  by  evapoi-ating  sea- water. 
This  latter  operation  is  often  conducted  in  shallow  bays 
(salterns]  in  which  the  water  is  left  to  evaporate  by  the 
heat  of  the  sun ;  hence  the  name  bay  salt.  The  mother 
liquor  (bittern)  contains  chloride,  bromide,  sulphate,  and 
iodide,  of  magnesium,  sodium,  potassium,  and  calcium  ; 
and  is  used  as  a  source  of  bromine  and  iodine.  Com- 
mon salt  contains,  as  impurities,  sodium  and  calcium 
sulphate,  and  magnesium  chloride. 

PROPERTIES. — Colourless  cubical  crystals.  Specific 
weight  =  2.16  ;  melting  point  =  776°;  soluble  in  water, 
the  solubility  increasing  slowly  with  the  temperature ; 
100  parts  of  water  at  5°C  dissolve  35.63  parts  of  the 
salt_at  50°,  37  parts— and  at  100°,  39. 16  parts.— Sodium 
chloride  forms  an  essential  constituent  of  blood  and  other 
animal  liquids.  It  seems  to  be  necessary  to  keep  in 
solution  certain  albuminous  compounds  which  are  insolu- 
ble in  pure  water. — It  is  used  in  medicine  as  an  antidote 
to  nitrate  of  silver.  (What  is  the  action  1) 

405.  Sodic  Sulphate  (Na2SO4). 

PREPARATION. — From  common  salt  by  the  action  of 
oil  of  vitriol ;  enormous  quantities  are  in  this  way  pre- 
pared as  "  salt  cake,"  in  the  manufacture  of  soda  (sodium 
carbonate).  Common  salt  is  treated  with  sulphuric  acid 
in  furnaces  so  constructed  that  half  the  salt  is  decom- 
posed at  a  comparatively  low  temperature  : 

2NaCl  +  H2S04  =  NaHS04  +  HOI  +  NaCl. 

and  the  remainder  by  transferring  the  mixed  salts  to  a 
hotter  part  of  the  furnace  : 

NaHS04  +  NaCl  =  Na2S04  +  HC1. 


SODIUM    CARBONATE.  371 

The  hydrochloric  acid  escapes  up  a  flue,  and  is  dissolved 
by  water  descending  in  a  shower.  Thus  is  obtained  the 
impure  hydrochloric  acid  of  commerce. — Sodic  sulphate 
is  sometimes  a  by-product  in  the  manufacture  of  nitric 
acid,  but  the  acid  sxilphate  (NaHSO4)  is  generally 
obtained. 

Experiment  324- — Warm  a  small  quantity  of  common  salt  in 
a  porcelain  basin  with  about  an  equal  weight  of  sulphuric  acid, 
dissolve  the  residue  in  water,  and  evaporate  to  crystallisation. 
The  crystals  have  the  composition  represented  by  Na2SO4.  IOH2O. 
They  are  called  Glauber's  salt.  Drain  them  and  leave  them 
exposed  to  the  air  for  a  day. 

PROPERTIES. — Crystallised  sodic  sulphate  (Na2SO4. 
10H2O)  is  a  colourless  salt,  of  bitter  saltish  taste.  It  is 
efflorescent,  and  is  soluble  in  water  (35.96  parts  in  100 
at  15°).  Its  solubility  is  greatest  at  34QC,  and  is  only 
42  in  100  at  100°.  It  is  very  sparingly  soluble  in 
alcohol. — Sodic  sulphate  is  much  used  in  Europe  as  a 
saline  purgative.  It  is  also  used  as  an  antidote  in  cases  of 
poisoning  by  salts  of  lead  and  barium.  (How  does  it  act  '£) 

406.  Sodic   Carbonate    (Na2CO3.lOH2O).     Also 

called  carbonate  of  soda,  soda,  and  washing  soda. 

PREPARATION. — 1.  By   LEBLANC'S    PROCESS,   in   three 


(a)  The  salt  cake  process.     (Art.  393.) 

(b)  The  black  ash  process,  in  which  the  salt  cake  is 
strongly  heated  with  limestone  and  slack  coal.    A  certain 
proportion  of  quick  lime  is  also  generally  added.     The 
following   equations    represent    the    principal    chemical 
actions  which  take  place  : 

Na2SO4  +  40  =  Na2S  +  4CO. 
Na2S  +  CaCO3  =  Na2CO3  +  CaS. 


372  THE    AMMONIA    SODA    PROCESS. 

Calcic  sulphide  is  insoluble,  and  the  sodium  carbonate  is 
extracted  from  the  "  black  ball  "  by 

(c)  Lixiviation.  The  strong  solution  thus  obtained  is 
evaporated  and  the  soda  allowed  to  crystallise.  It  may 
be  purified  by  recrystallisation. 

2.  By  the  AMMONIA-SODA  process.  A  saturated  solu- 
tion of  common  salt  is  saturated  with  ammonia,  and  then 
a  current  of  carbon  dioxide  is  passed  through  it.  Sodic 
hydric  carbonate  is  precipitated  : 

NH3  +  CO2  +  NaCl  +  H2O  =  HNaCO3  +  NH4C1. 

The  ammonia  is  recovered  by  distilling  the  mother  liquor 
with  magnesia  : 

MgO  +  2NH4C1  =  MgCl2  +  2NH3  +  H2O. 

The  magnesia  is  recovered  (as  hydroxide)  by  treating  the 
chloride  with  superheated  steam  : 

MgCl2  +  2H2O  =  Mg(OH)2  +  2HC1. 

Sodic  hydric  carbonate  is  easily  converted  into  the  car- 
bonate by  the  action  of  heat : 

2NaHCO3  =  Na2CO3  +  H2O  +  CO2. 

PROPERTIES. — Sodic  carbonate  is  sold  in  large  crystal- 
line lumps,  or  in  the  form  of  a  white  powder.  The 
crystals  are  efflorescent,  and  are  readily  soluble  in  water 
(60  parts  in  100).  The  solution  is  alkaline  in  reaction, 
and  neutralises  strong  acids. 

Experiment  325- — Heat  gently  a  crystal  of  washing  soda  in 
a  t.  t.  and  note  the  escape  of  water.  Dissolve  a  fragment  in 
water  and  test  with  red  litmus.  Add  sulphuric  acid  to  a  solu- 
tion of  washing  soda. 

Experiment  326.  —Heat  some  washing  soda  in  a  porcelain 


CAUSTIC    SODA.  373 

capsule  until  the  liquid  at  first  formed  dries  up  to  a  cake.  This 
is  dried  carbonate  of  soda.  Powder  and  preserve  it  in  a  stop- 
pered bottle. 

407.  Sodic  Hydric  Carbonate    (NaHCO3).  - 

Also  called  bicarbonate  of  soda,  and  baking  soda. 

PREPARATION. — By  exposing  washing  soda  crystals  to 
the  action  of  carbon  dioxide  evolved  by  the  action  of 
hydrochloric  acid  on  marble  or  limestone  : 

Na2CO3.10H20  +  CO2  =  2NaHC03  +  9H2O. 

It  is  also  prepared  from  common  salt  by  the  ammonia 
soda  process  (Art.  406). 

PROPERTIES. — A  white  powder,  soluble  in  water,  and 
in  hydrochloric  and  other  acids,  with  much  efferves- 
cence. It  is  decomposed  by  heat  (Art  406.)  It  is  much 
less  soluble  in  water  than  the  normal  carbonate,  100 
parts  of  water  at  15°,  dissolving  only  10.5  parts. 

408.  Sodic  Hydroxide    (NaOH).  —  Also   called 

sodium  hydrate  and  caustic  soda. 

PREPARATION. — Experiment  327. — Dissolve  some  washing 
soda  in  about  six  times  its  weight  of  water,  heat  to  boiling  in  a 
porcelain  dish,  add,  a  little  at  a  time,  slaked  lime  equal  to  about 
half  the  weight  of  the  washing  soda,  keeping  the  liquid  at  the 
boiling  point,  and  adding  water  as  it  boils  away.  Allow  to 
settle,  pour  off  the  clear  liquor,  and  test  it  with  hydrochloric 
acid.  It  should  give  no  effervescence.  Test  the  precipitate 
with  hydrochloric  acid.  It  effervesces  : 

Ca(OH),,  +  Na2C03  =  CaC03  +  2NaOH. 

This  solution  (liquor  sodce)  prepared  on  the  large 
scale  in  this  way  is  boiled  down  in  iron  pots  until  a  red 
heat  is  attained,  when  the  molten  caustic  soda  is  run 


374  SODIUM    NITRATE. 

into  iron  cylinders  and  sealed  up. — Caustic  soda  is  also 
formed  when  sodium  decomposes  water.     (Exp't  30). 

PROPERTIES. — A  brittle  white  solid,  of  specific  weight 
2.13.  It  melts  at  a  dull  red  heat,  but  is  not  decomposed 
until  an  intense  white  heat  is  attained.  It  is  delique- 
scent, and  when  exposed  to  the  air  soon  becomes  changed 
to  carbonate.  It  dissolves  readily  in  water,  with  the 
evolution  of  heat. 

Experiment  328. — Put  a  small  piece  of  caustic  soda  in  a 
porcelain  dish,  leave  it  a  few  days,  and  then  test  it  with  hy- 
drochloric acid. 

Experiment  329. — Pour  a  little  water  upon  a  piece  of  caustic 
soda  in  a  t.  t.  Note  the  heat.  Add  more  water,  and  note  the 
taste  and  action  on  the  skin  of  the  resulting  solution. 

Caustic  soda  has  a  strong  corrosive  action  on  animal 
tissues.  It  is  therefore  very  poisonous. — Solution  of 
caustic  soda  dissolves  glass  and  porcelain.  This  goes  on 
gradually  even  with  very  dilute  solutions. — Caustic  soda 
has  many  uses.  Combined  with  fatty  acids  it  forms 
hard  soaps ;  and  on  account  of  its  solvent  action  on  fats, 
&c.,  it  is  used  in  cleansing  rags,  grass,  <kc.,  in  the  manu 
facture  of  paper.  In  medicine  caustic  soda  is  used  as  an 
escharotic.  In  the  chemical  laboratory  it  ia  often  used 
to  precipitate  insoluble  metallic  hydroxides. 

409.  Sodic  Nitrate  (NaNO3).- — Also  called  cubic 
nitre  and  Chili  saltpetre.  This  salt  is  found  in  large 
quantities  in  Peru  and  Bolivia.  It  is  obtained  by  min- 
ing, and  purified  by  solution  and  crystallisation.  Its 
crystals  when  perfect  are  nearly  cubical.  It  is  very 
soluble  in  water  (84  parts  in  100),  and  somewhat  deli- 
quescent. In  dissolving,  it  renders  much  heat  latent. 


SODIUM    PHOSPHATE.  375 

This  accounts  for  its  cooling  taste.  It  is  used  as  a  fer- 
tiliser, and  in  the  manufacture  of  nitric  acid  and  inferior 
blasting  powders. 

410.  Sodic  Sulphite  (Na,S03.7H2O). 

PREPARATION.  —  Divide  a  solution  of  sodium  carbonate 
into  two  equal  parts.  Saturate  one  with  sulphur  di- 
oxide, add  the  other,  and  evaporate  to  crystallisation. 
Sodic  hydric  sulphite  (NaHSO3)  is  first  formed  : 

Na2CO3+2SOa  +  H2O  =  2NaHSO3  +  CO2. 
This  is  then  converted  into  the  normal  sulphate  : 
2NaHSO3  +  Na2CO3  -  2Na2SO3  +  C02  +  H2O. 

PROPERTIES. — A  colourless  crystalline  salt,  soluble  in 
water  ( 1  part  in  4).  The  solution  is  alkaline.  Both  the 
normal  and  the  acid  sulphite  are  used  in  medicine. 
They  are  used  especially  to  destroy  sarcince  ventriculi, 
minute  parasitic  plants  sometimes  present  in  the  stomach. 
The  acid  of  the  gastric  juice  sets  free  sulphur  dioxide, 
which  has  an  antiseptic  action  : 

Na2S03  +  2HC1  =  2NaCl  +  S02  +  H2O. 

Sodic  sulphite  must  be  kept  well  stoppered,  as  it  absorbs 
oxygen  from  the  air  and  becomes  oxidised  to  sulphate. 

411.  Sodic  Phosphate  (Na^HPO4.12H2O).— This 

is  the  "common"  or  "rhombic"  phosphate  of  soda.  It 
was  first  prepared  from  urine,  and  is  present  in  the 
blood  and  in  other  animal  liquids. 

PREPARATION. — Bone  ash  is  decomposed  with  sul- 
gulphuric  acid  : 

Ca3(P04)2  +3H2SO4  =  3CaSO4  +  2H3PO4. 


376  SODIUM    BROMIDE. 

The  solution  of  phosphoric  acid  is  separated  from  the 
sparingly  soluble  gypsum,  and  treated  with  sodic  car- 
bonate : 

H3P04  +  Na2C03  =  Na2HPO4  -f-  H2O  +  C02. 

The  solution  is  then  evaporated  to  crystallisation. 

PROPERTIES.  —  A  colourless  salt,  soluble  in  water  (14 
parts  in  100).  Its  solution  is  alkaline  in  reaction.  Its 
taste  closely  resembles  that  of  common  salt. 

Experiment  330-  —  Dissolve  a  little  sodic  phosphate  in  water. 
Note  the  taste  of  the  solution.  Test  with  red  litmus.  Add  a 
few  drops  of  argentic  nitrate  to  a  little  of  the  solution.  Note 
the  yellow  precipitate  (Ag3P04)  : 

3Na2HP04  +  6AgN08  =  2Ag3P04  +  6NaN03  +  H3PO4. 

412.  Sodic    Bromide  (NaBr).—  This    salt    is   pre- 
pared by  dissolving  bromine  in  solution  of  caustic  soda  : 

GNaOH  +  3Br2  =  5NaBr  +  NaBrO3  +  3H2O. 

evaporating  to  dryness,  and  heating  with  a  little  char- 
coal to  decompose  the  bromate  : 

2NaBr03  +  30  =  2NaBr  +  3CO2. 

It  is  a  soluble,  colourless,  crystalline  salt,  similar  to  the 
potassium  salt.  It  is  proposed  to  use  it  in  medicine  in 
place  of  potassium  bromide,  since  it  is  equally  efficacious, 
and  causes  none  of  the  unpleasant  symptoms  resulting 
from  the  continued  use  of  the  potassium  salt. 

413.  Sodic  Sulphide  (Na.2S).  —  Can   be    prepared 
impure  (liver  of  sulphur)  by  heating  sodic  sulphate  with 
charcoal  : 

40  -  NaS  +  4CO. 


To  prepare  solution  of  sodic  sulphide,  take  two  equal 


GLASS.  377 

quantities  of  a  solution  of  caustic  soda,  saturate  the  one 
with  hydrogen  sulphide  : 

NaOH  +  H2S  =  NaSH  +  H2O, 
and  then  add  the  other  : 

NaSH  +  NaOH  =  Na,S  +  H2O. 

The  same  method  is  used  for  preparing  solutions  of  po- 
tassium and  ammonium  sulphides.  —  Solution  of  sodium 
sulphide  is  used  to  precipitate  and  dissolve  sulphides  of 
heavy  metals.  It  combines  with  the  sulphides  of  arsenic, 
antimony,  &c.,  to  form  soluble  sulphur  salts;  e.g.: 

Sb2S5  =  2Na3SbS4. 


414.  Glass.  —  Glass  is  a  mixture  of  silicates.  The 
materials  used  are  (1)  silica,  in  the  form  of  quartz, 
ignited  flint,  white  sand,  or  red  sand  ;  (2)  alkali,  purified 
potashes,  refined  soda  ash,  or  salt  cake  ;  and  (3)  calcic 
carbonate,  &c.  —  calcspar,  marble,  chalk,  or  limestone,  and, 
for  flint  glass,  red  lead  or  litharge.  There  are  sevei-al 
varieties  of  glass  : 

1.  Bohemian  Glass.  —  Silicates  of  potassium  and  calci- 
um.    It  fuses  with  difficulty  and   resists  the  action   of 
chemicals  better  than  the  other  kinds  of  glass. 

2.  Window,  or  Grown,  Glass.  —  Silicates  of  sodium  and 
calcium.     It  is  more  fusible  than  Bohemian  Glass,  and 
more  easily  acted  on  by  chemicals. 

3.  Bottle  Glass.  —  Silicates  of  sodium  and  calcium,  but 
made  from  cheap  materials.     Its  colour  is  due  to  iron 
compounds. 

4.  Flint  Glass,  Crystal,  or  Strass.  —  Silicates  of  potas- 
sium and  lead.     It  is  heavy,  fusible,  and  has  a  bright 


378  POTASSIUM. 

lustre.     It  is  used  for  ornamental  purposes,  and  one  vari- 
ety (paste)  is  used  for  imitating  diamonds. 

The  properties  of  glass  make  it  very  useful  in  chemical 
operations.  It  is  transparent,  not  readily  fused,  and 
only  slowly  and  sparingly  soluble  in  most  chemical 
substances.  It  can  be  fused  at  a  red  heat,  and  can  then 
be  moulded  into  any  desired  form. 

415.  Tests. 

Very  few  salts  of  sodium  are  insoluble,  so  that  the  tests  are 
mostly  negative.  Sodic  metantimonate  is  insoluble,  and  a  solu- 
tion of  the  potassium  salt  is  sometimes  used  as  a  test  for  sodium; 
but  the  test  chiefly  relied  upon  in  analysis  is  the  exclusion  of 
other  metals,  and  the  yellow  colour  given  to  the  Bunsen  flame 
when  sodium  compounds  are  brought  into  it. 


POTASSIUM. 

416.  Potassium  (K1  =  39.04.  Specific  weight  = 
0.875.  Melting  point  =  62°.5). 

OCCURRENCE. — Is  found  almost  universally,  but  always 
in  combination.  It  forms  from  1.5  to  3.1  %  of  granite,  and 
occurs  in  many  double  silicates.  Potash  felspar  is  a 
double  silicate  of  potassium  and  aluminium  (K2O.A1208. 
6Si02),  which  "  weathers  "  and  thus  forms  clay.  The 
chief  source  of  potassium  compounds  is  the  mineral  de- 
posits of  Stassfurth,  which  contain  silvine  (KC1),  carnal- 
lite  (KCl.MgCl2.6H2O),  &c.  Potassium  compounds  are 
present  in  sea  water,  in  soils,  and  in  plants  and  animals. 

PREPARATION. — In  the  same  way  as  sodium,  but 
special  precautions  must  be  taken  on  account  of  the 
formation  of  an  explosive  compound  of  potassium  and 
carbon  monoxide. 


POTASSIC    CARBONATE.  379 

PROPERTIES. — A  silver  white  metal,  resembling  sodi- 
um, but  it  is  more  easily  oxidised.  Its  chemical  pro- 
perties are  very  like  those  of  sodium,  but  more  pro- 
nounced. It  is  used  in  preparing  boron,  silicon,  magne- 
sium, &c. 

417.  Potassic  Carbonate. — (K2CO3).  Also  called 

carbonate  of  potash,  potashes,  and  salt  of  tartar  (was  form- 
erly prepared  by  igniting  '  cream  of  tartar '). 

PREPARATION. — 1.  From  potassic  chloride  (KC1),  and 
sulphate  (K.2SO4),  by  the  same  process  as  that  employed 
for  sodium  carbonate  (p.  371). 

2.  From    wood  ashes,  by   leaching,    evaporating,  and 
calcining.     The   residue   is  impure  potassium  carbonate 
(potashes).      It  is  purified    by   recrystallisation,  and   is 
then  called  pearl  ash. 

3.  From  the  waste  liquors  of  the  beet  sugar  industry, 
by  evaporation  and  repeated  crystallisation. 

4.  From    the    washings    of   sheep's  wool,   technically 
called  "  suint."     The  washings  are  evaporated  to  dryness 
and  distilled.     An  illuminating  gas,  and  an  ainmoniacal 
liquor,  are  obtained.     The  fixed  residue  is  lixiviated,  &c., 
for  potassium  carbonate.     Commercial  carbonate  of  pot- 
ash is  rarely  pure.     It  can  be  purified  by  treating  a  sat- 
urated aqueous  solution  with  carbon  dioxide,  and  collect- 
ing and  heating  the  acid  carbonate  thus  precipitated  : 

K2C03  +  C02  +  H2O  =  2KHC03. 
2KHCO3  =  K2C03  +  H20  +  CO2. 

PROPERTIES. — A  white  granulated  powder,  or  a  pasty 
mass,  very  deliquescent,  very  soluble  in  water  (106.4 
parts  in  100),  sparingly  soluble  in  alcohol.  The  aqueous 


380  BICARBONATE    OF    POTASH. 

solution  has  an  alkaline  reaction,  the  properties  of  the 
strong  base  being  only  partially  neutralised  by  the  weak 
acid.  Potassic  carbonate  is  used  in  medicine  as  an  ant- 
acid, and  to  promote  the  solution  of  uric  acid  stones.  Its 
taste  is  very  disagreeable.  It  is  also  used  in  the  manu- 
facture of  soft  soap,  crystal  glass,  potassium  ferrocyanide, 
bichromate,  and  cyanide. 

Experiment  331. — Scatter  a  few  grains  of  dry  potassium 
carbonate  upon  a  sheet  of  paper  and  examine  after  a  few  hours. 
Make  a  solution,  and  try  the  taste  and  action  on  red  litmus. 

418.  Potassic  Hydric  Carbonate  (KHCO3).— 

Also  called  bicarbonate  of  potash. 

PREPARATION. — By  passing  a  current  of  carbon  diox- 
ide for  several  hours  through  a  cold  saturated  solution  of 
potassium  carbonate.  The  sparingly  soluble  acid  carbon- 
ate is  precipitated. 

PROPERTIES. — Colourless  crystals,  of  a  saltish  taste ; 
soluble  in  water  (25  parts  in  100),  giving  a  slightly  alka- 
line solution,  decomposed  by  heat,  even  when  in  solution  : 

2KHC03  =  K2C03  +  H20  +  C02. 

This  salt  has  none  of  the  corrosive  action  of  the  normal 
carbonate.  It  is  used  as  an  antacid  and  antilithic.  Li- 
quor potassoe  effervescent,  or  potash  water,  is  a  solution 
of  potassium  bicarbonate  into  which  carbon  dioxide  has 
been  introduced  under  a  pressure  of  seven  atmospheres. 

419.  Potassic    Hydroxide  (KOH),  also    called 

potassium  hydrate,  and  caustic  potash. — This  compound  is 
prepared  from  potassium  carbonate  by  the  same  method 


POTASSIC    CHLORATE.  381 

as  that  used  for  sodic  hydroxide,  which  it  resembles  in 
its  properties.  It  is,  however,  more  strongly  corrosive 
in  its  action  on  the  skin  &c.  Potash  lye  is  an  aqueous 
solution  of  impure  potassic  hydroxide. 

420.— Potassic    Chlorate   (KC1O3),    also   called 

chlorate  of  potash. 

PREPARATION. — Milk  of  lime  is  saturated  with  chlo- 
rine, and  part  of  the  water  is  evaporated  : 

6Ca(OH)2  +  6C12  =  Ca(ClO3)2  +  5CaCl2  +  6H2O. 

Potassium  chloride  is  added  (In  what  proportion  ?),  and 
the  solution  is  boiled  down  and  allowed  to  cool,  when 
potassic  chlorate  crystallises  out  : 

Ca(ClO3)2  +  2KC1  =  CaCl2  +  2KC1O3 

In  this  way  the  whole  of  the  potassium  is  obtained  as 
chlorate.  The  old  method  is  wasteful  (Expt.  76). 

PROPERTIES. — Colourless,  flat  crystals,  or  a  white 
granular  powder.  It  has  a  cooling  acid  taste,  and  is 
sparingly  soluble  in  water  (6  parts  in  100).  When 
heated  to  352°  C.  it  decomposes  into  oxygen,  chloride, 
and  perchlorate  : 

2KC1O3  =  O2  +  •  KC1  +  KC1O4. 

At  a  higher  temperature  the  whole  of  the  oxygen  is 
driven  off.  It  is  a  powerful  oxidising  agent,  but  is  not 
capable  of  supplying  oxygen  to  the  blood. 

Experiment  332. — Carefully  mix  some  dry  sugar  with  about 
one-fourth  its  weight  of  powdered  potassic  chlorate,  place  the 
mixture  on  a  stone  or  a  piece  of  poroelain  and  touch  with  a  glass 
rod  dipped  in  concentrated  sulphuric  acid.  (Explain  the  action). 

421.  Potassic  Nitrate,  (KNO3). — Also  called  salt- 
petre, and  nitre. 


382  POTASSIC    BROMIDE. 

OCCURRENCE. — As  efflorescence  on  the  soil  in  hot  dry 
countries,  such  as  Bengal  and  Egypt.  Has  been  obtained 
by  lixiviating  certain  porous  rocks  (hence  the  name,  sal 
petrce).  Its  formation  in  the  soil  is  due  to  the  slow  oxi- 
dation of  nitrogenous  matter  in  the  presence  of  potassic 
carbonate  or  silicate.  Nitric  acid  is  first  formed  and  this 
decomposes  the  potassium  carbonate,  &c. 

PREPARATION. — Tt  is  prepared  from  the  soil  incrusta- 
tion by  lixiviation  and  crystallisation.  Much  saltpetre 
is  now  prepared  from  potassic  chloride  by  dissolving  along 
with  an  equivalent  (Calculate  the  proportions)  of  sodium 
nitrate  in  hot  water  until  the  specific  weight  is  1.5.  Sodi- 
um chloride  is  precipitated,  and  potassium  niti-ate  sepa- 
rates out  when  the  solution  cools  : 

KC1  +  NaNOj  =  KNO3  +  Nad. 
(Compare  the  solubilities  of  these  four  salts.) 

PROPERTIES. — Long  colourless  crystals,  of  a  cooling 
bitter  taste,  soluble  in  water  (26  parts  in  100).  It  is  an 
oxidising  agent  and  plays  this  part  in  gun  powder  and  in 
many  coloured  fires.  Its  uses  in  medicine  depend  princi- 
pally on  its  cooling  properties. 

422.  Potassic  Bromide,  (KBr). 

PREPARATION. — Experiment  333.— Gradually  add  dilute 
solution  of  potassic  hydroxide  to  a  drop  of  bromine  under  water, 
until  the  colour  of  the  bromine  disappears.  Evaporate  the  solu- 
tion to  dryness  in  a  porcelain  dish  and  ignite  the  residue.  Potas- 
sic bromide  remains  : 

6KOH  +  3  Br2  =  KBr03  +  5KBr  -f  3H20. 

KBr03  =  KBr  -f  30. 
Dissolve  in  a  little  hot  water  and  crystallise. 

PROPERTIES. — Translucent,  colourless,  cubical  crystals, 


POTASSIC    IODIDE AMMONIUM.  383 

resembling  those  of  potassic  iodide,  but  not  so  porcelain- 
like  ;  taste,  shai-p  and  saline ;  readily  soluble  in  water, 
somewhat  sparingly  in  alcohol. 

423.  Potassic  Iodide, 'KD. 

PREPARATION. — Experiment  334.  —  Repeat  Expt.  333, 
using  iodine  instead  of  bromine.  Potassic  iodide  is  obtained. 
(Write  the  equations). 

PROPERTIES. — Potassic  iodide  is  very  like  the  bromide 
in  its  properties.  It  crystallises  in  cubes,  which  are 
opaque  and  porcelain-like  when  deposited  from  a  hot 
solution,  but  clear  when  crystallised  cold.  It  is  very 
soluble  in  water  (140  parts  in  100 '.  (Examine  carefully 
and  compare  crystals  of  potassium  bromide  and  iodide). 

424.  Tests. 

1.  Tartaric  acid  gives  a  white  crystaline  precipitate  of  potassic 
hydric  tartrate,  especially  on  stirring  with  a  glass  rod. 

2.  Platinum  tetrachloride  gives  a  yellow  crystalline  precipitate 
(K2PtCl8).     Make  this  test  by  stirring  together  a  drop  or  two  of 
the  solutions  on  a  watch  glass. 

3.  Potassium  compounds  give  a  violet  colour  to  the  Bunsen 
flame. 


AMMONIUM. 

425.  Ammonium  Salts. — These  are  compounds 
containing  the  radical  ammonium  (NH4 — ),  which  acts 
the  part  of  a  monad  atom.  They  are  generally  prepared 
by  neutralising  the  appropriate  acids  with  solution  of 
ammonia.  This  solution  may  be  supposed  to  contain 
ammonium  hydroxide  (NH4OH),  which  acts  towards 
acids  as  sodic  hydroxide  does,  e.g.  : 

NH4.OH  +  HC1  =  NH4.C1  -f  H2O. 
Compare     KOH  +  HC1  =     K.C1  +  H2O. 


384  AMMONIC    CHLORIDE. 

The  ammonium  salts  resemble  those  of  sodium  and  potas- 
sium, especially  the  latter,  but  can  all  be  decomposed 
by  heat. 

426.  Ammonic  Sulphate,  ((NH4)2SO4). 

PREPARATION. — By  heating  gas  liquor  with  lime,  and 
receiving  the  evolved  ammonia  in  dilute  sulphuric  acid. 
The  solution  thus  obtained  is  evaporated  to  crystallisation. 

PROPERTIES. — A  colourless  salt,  generally  in  small 
crystals,  soluble  in  water  (75.5  parts  in  100).  When 
heated  strongly  it  volatilises  completely. 

Experiment  335- — Heat  a  little  ammonic  sulphate  in  a  dry 
glass  tube.  Dissolve  another  portion  in  water  and  test  it  for 
sulphuric  acid. 

Ammonic  sulphate  is  used  as  a  fertiliser.  It  is  the 
starting  point  in  the  manufacture  of  other  ammonium 
salts. 

427.  Ammonic   Chloride,   (NH4C1),  also   called 

sal  ammoniac. 

PREPARATION. — Experiment  336-  —  Neutralise  a  small 
quantity  of  dilute  hydrochloric  acid  with  solution  of  ammonia 
and  evaporate  to  dryness.  Ammonium  chloride  remains  : 
NH3  +  HC1  =  NH4C1. 

Ammonic  chloride  is  also  prepared  by  subliming  a 
mixture  of  the  sulphate  and  common  salt : 

(NH4)2SO4  +  2NaCl  =  2NH4C1  +  Na2SO4. 
It  is  purified  by  re-sublimation. 

PROPERTIES. — Colourless  crystals,  either  in  tough 
masses  of  flexible  fibres,  or  in  grains.  It  has  a  sharp 
cooling  taste,  and  is  freely  soluble  in  water  (35  parts  in 
100).  When  the  aqueous  solution  is  boiled,  ammonia 


AMMONIUM    CARBONATE.  385 

escapes  and  the  solution  becomes  acid.  Ammonium  chlo- 
ride renders  latent  a  great  deal  of  heat  when  dissolving. 
It  is  an  excellent  cooling  agent. 

Experiment  337 — Dissolve  some  ammonic  chloride  in  a  little 
water  and  note  the  low  temperature  produced.  Heat  a  small 
quantity  of  the  solid  in  a  dry  glass  tube.  It  sublimes  and  leaves 
no  residue,  if  it  is  pure. 

428.  Ammonic  Carbonate. — The  normal  carbon- 
ate (NH4)2CO3)  is  difficult  to  prepai'e  and  keep.  It  loses 
ammonia  and  becomes  converted  into  the  acid  carbonate 
(NH4HCO3).  The  "  sesquicarbonate  "  of  commerce  is  a 
compound  of  the  acid  carbonate  with  ammonium  car- 
bamate  (NH4.CO,.NH2). 

PREPARATION. — By  subliming  a  mixture  of  chalk  and 
sal  ammoniac  or  ammonium  sulphate  : 

CaC03  +  (NH4)2S04  =  (NH4)2CO3  +  CaSO4. 

The  salt  which  sublimes  is,  however,  not  the  normal  car- 
bonate represented  in  the  equation,  but  the  "  sesquicar- 
bonate "  ^so  called)  : 

NH4HC03.NH4NH2C02 

Experiment  338- — Heat  a  mixture  of  ammonium  sulphate 
and  ground  limestone  in  a  dry  test  tube.  Scrape  out  a  little  of 
the  sublimate,  observe  its  odour,  and  note  that  it  effervesces 
with  hydrochloric  acid. 

PROPERTIES. — Commercial  carbonate  of  ammonia  (sal 
volatile)  is  sold  in  hard  translucent  crystalline  masses. 
It  smells  of  ammonia,  and,  if  exposed  to  the  air,  is  soon 
changed  to  the  acid  carbonate  by  losing  ammonia  and 
gaining  water  : 

NH4.NH2.CO2  +  H2O  =  NH4HCO3  +  NH3. 

It  is  soluble  in  water  (27.5  parts  in   100),  and  the  solu- 
26 


.5KG  AMMONIUM     SUU'IIIDK. 

tion  is  alkaline.      Aromatic  spirit  of  ammonia  is  a  prepa- 
ration of  the  carbonate. 

429.  Ammonic  Phosphate,  (NH4)2HPO4. 

PREPARATION. — By  neutralising  solution  of  phosphoric 
acid  with  ammonia  solution,  and  crystallising  : 

2NH3  +  H3P04  =  (NH4)2  HP04 

PROPERTIES. — Colourless  crystals,  soluble  in  water, 
insoluble  in  alcohol. 

430.  MicroCOSXnic  Salt. — Is  hydric  sodic  ammonic 
phosphate  (H.Na.NH4!PO4.4H20.)  first  noticed  as  crystal- 
lising from  concentrated  urine.     It  is  a  colourless  crystal- 
line   salt  prepared  by   mixing    hot    strong   solutions  of 
sodium  and  ammonium  phosphates  and  allowing  to  crys- 
tallise.    By  heat  it  decomposes  as  follows  : 

HNaNH4PO4.4H2O  =  5H2O  +  NH3  +  NaP03. 
The  non-volatile  sodic  metaphosphate  remains. 

431.  Ammonic  Sulphide,  ((NH4)2S.) 
PREPARATION. — In  the  some  way  as  solution  of  sodic 

sulphide  (Art.  413). 

PROPERTIES. — Forms  a  colourless  solution  which  gradu- 
ally turns  yellow,  owing  to  the  absorption  of  oxygen, 
which  sets  free  sulphur.  Ammonium  sulphide  solution 
dissolves  sulphur,  and  thus  forms  "  yellow  ammonium 
sulphide,"  used  in  analysis  to  dissolve  stannous  sulphide. 

Experiment  339. — Add  hydrochloric  acid  to  a  few  drops  of 
yellow  ammonium  sulphide,  and  observe  the  precipitation  of 
sulphur  and  the  evolution  of  hydric  sulphide  : 

(NH4)aS»  +  2HC1  =  2NH4C1  -f-  H2S  +  aS. 
Ammonic  sulphide  is  poisonous. 


LITHIUM.  387 

432.  Tests. 

1.  Heat  in  a  test  tube  with  caustic  soda  solution,  observe  the 
odour,  and  hold  over  the  mouth  of  the  test  tube  a  glass   rod 
moistened   with   dilute    hydrochloric    acid.     White   fumes   are 
formed  : 

NH3  +  HC1  =  NH4C1. 

(What  is  the  object  of  the  caustic  soda  ?} 

2.  Platinum  tetrachloride  gives  a  yellow  precipitate 
((NH4)2PtCl6).     (Art.  424  (2)). 


LITHIUM. 

433.  Lithium  (Li1  =  7.01.—  SpeciBc  weight  =  0.59.— 
Melting  point  =  180°.)  The  metal  ia  prepared  by  elec- 
trolysing the  fused  chloride  (LiCl).  It  resembles  sodium 
and  potassium  in  its  properties,  but  is  much  lighter,  and 
has  not  so  strong  an  attraction  for  oxygen.  Lithium 
compounds  are  widely  diffused,  but  in  small  quantities. 
They  are  found  in  most  mineral  waters,  and  in  river  and 
spring  water  generally.  The  compounds  of  lithium  re- 
semble those  of  sodium  and  potassium,  but  the  hydroxide 
(LiOH),  carbonate  (Li2CO3\  and  phosphate  (Li3PO4),  are 
much  less  soluble  in  water. 


434.  Lithic   Carbonate,  (L^COs).  —  This   is    the 

most  important  compound  of  lithium,  as  it  is  much  used 
in  medicine  in  treating  gout,  stone,  &c. 

PREPARATION.  —  Experiment  340-  —  To  a  small  quantity  of 
solution  of  ammonium  carbonate  in  liquor  ammonia;  add  a  small 
quantity  of  a  strong  solution  of  lithium  chloride.  Lithium  car- 
bonate is  precipitated  : 

2L1C1  +  (NH4),00,  =  Li.2C03  +  2NH4CL 
Warm,  filter,  and  wash  with  cold  water. 


.M     C.KSHM. 

PROPERTIES. — Lithium  carbonate  is  a  white  crystalline 
powder,  sparingly  soluble  in  water  (0.78  parts  in  100  . 
It  is  more  soluble  if  carbonic  acid  be  added,  as  the  bi- 
carbonate (LiHCO;!)  is  formed,  of  which  5.25  parts  dis- 
solve in  100  of  water.  When  this  solution  is  exposed  to 
the  air,  it  loses  carbon  dioxide  and  the  normal  carbonate 
is  precipitated.  (What  other  carbonates  behave  simi- 
larly 1) 

Lithium  carbonate  is  prescribed  for  gout,  stone,  &c.  It 
is  preferable  to  the  potassium  salt.  (Art.  167). 

435.  Tests. 

Lithium  compounds  can  be  recognised  by  the  spectrum  of  the 
beautiful  red  colour  which  they  give  to  the  Bunsen  flame. 

436.  Rubidium,  (Kb1  =   85.2).— Compounds  of  this 
rather  rare  alkaline  metal  are  found  in  mineral  springs, 
and  in  some  minerals.     The  metal  can  be  prepared  by  the 
same  process  as  that  for  the  preparation   of  sodium  and 
potassium.      It  is  like  these  in  properties,  but  has  greater 
chemism  for  oxygen  than  potassium  has.     Its  compounds 
are   similar  in   composition  and   properties   to  those  of 
potassium,  e.g.,  Rb.2O,  RbOH,  KbCl,  &c. 

437.  Caesium    (Gsi-  =  132.5). — Caesium   compounds 
were   discovered  in    1860  by  Bunsen,  by  means  of  the 
spectroscope ;   and    this  discovery  was  the  first   fruit  of 
spectrum  analysis.     When  white  light  is  passed  through 
the  edge  of  a  wedge-shaped  piece  of  glass  (prism],  it  is 
spread  out  in   such  a  way  that  the  waves  of  different 
lengths  fall  on  different  parts  of  a  retina  receiving  them. 
The  sensation  is  one  of  a  band  of  coloured  lights  ranging 
from  red  to  violet.      Such  a  band  is  called  the  spectrum 


QUESTIONS    AND    KXKKC1SES.  389 

and  in  the  spectrum  each  colour  has  its  fixed  place. 
Now,  each  element  in  the  state  of  a  hot  vapour  gives  a 
light  corresponding  to  pai-ticular  lines  in  the  spectrum  ; 
and  if,  when  looking  through  a  prism  (appropriately  ar- 
ranged in  a  spectroscope},  we  see  in  a  flame  bands  or 
lines,  we  can  recognise  these  as  being  due  to  the  presence 
of  some  known  element.  Bunsen,  when  looking  at  a 
flame  in  which  was  volatilising  the  solid  residue  of  the 
water  from  a  mineral  spring,  saw  lines  produced  by  no 
known  element.  He  thus  discovered  the  metals  caesium 
and  rubidium. — Caesium  was  prepared  in  1881  by  Carl 
Setterberg  by  electrolysing  the  fused  cyanide  (CsCN). 
It  melts  at  the  temperature  of  the  hand,  and  exceeds 
rubidium  in  its  chemism  for  oxygen.  Its  compounds  are 
very  like  those  of  rubidium  and  potassium,  e.g.,  Cs.2O, 
CsOH,  CsCl,  Cs,CO3,  &c. 


QUESTIONS  AND  EXERCISES. 

1 .  Calculate  the  percentage  of  water  in  washing  soda. 

2.  How  much  dried  sodium  carbonate  is   equivalent  to    100 
grains  of  washing  soda  ? 

3.  What  substances  are  antidotes  to  caustic  soda  and  caustic 
potash  ? 

4.  "Sodium  nitrate  gives,  weight  for  weight,  more  nitric  acid 
than  potassic  nitrate  does."     Show  the  truth  of  this  statement. 

5.  How  would  you  distinguish  (practically)  sodium  carbonate 
from  potassium  carbonate. 

6.  Solution  of  normal   sodic  sulphite  is  alkaline  in  reaction. 
How  do  you  account  for  this  ? 

7.  Show  that,  weight  for  weight,  lithium  carbonate  will  dis- 
solve a  larger  quantity  of  uric  acid  stone  than  will  potassium 
carbonate. 


390  ELECTRICITY. 

8.  How  much  potassium  chlorate  (KC103)  can  be  obtained 
from  10  Ibs.  of  caustic  potash  (KOH) — (1)  by  the  method  of 
Art.  101,  and  (2)  by  that  of  Art.  420? 

8.  What  would  you  use  as    an  antidote  to  poisoning  by  am- 
monium sulphide  ? 

9.  Arrange  the  metals  of  the  alkalis  ( 1 )  in  the  order  of  their 
atomic  weights,  and  (2)  in  the  order  of  their  chemism  for  oxygen. 


CHAPTER    XXII. 


ELECTRICITY. 

438.  Voltaic  Batteries. — If  a  plate  of  copper  held 
by  a  glass  rod  (an  insulator)  be  touched  by  a  zinc  plate 
similarly  held,  and  then  withdrawn,  the  two  plates  will 
be  found  to  be  in  a  peculiar  condition,  such  that  they 
produce  exactly  opposite  effects  on  electrified  bodies. 
What  the  copper  attracts  the  zinc  repels,  and  vice  versa, — 
the.  zinc  is  positively,  the  copper  negatively,  electrified. 
The  condition  is  one  analogous  to  a  difference  of  temper- 
ature. If  a  drop  of  sulphuric  acid  be  placed  between 
the  plates,  a  similar  condition  is  produced.  And,  if  the 
dry  ends  of  the  plates  be  now  connected  by  a  copper 
wire,  so  as  to  complete  the  circuit,  the  copper  and  zinc 
tend  to  assume  the  same  electrical  condition  by  a  trans- 
ference of  electricity  through  the  liquid  from  the  zinc  to 
the  copper.  At  the  same  time  sulphuric  acid  is  decom- 
posed into  H2  and  SO4,  the  hydrogen  appearing  at  the 
surface  of  the  copper  plate,  and  the  salt  radical  (SOt) 
attacking  the  zinc  plate  and  forming  zinc  sulphate.  While 
this  takes  place  in  the  liquid  there  passes  round  the  rest 
of  the  circuit  a  disturbance  of  some  sort,  called,  for  lack 


ELECTRICITY.  391 

of  a  better  term,  a  current  of  electricity.  The  source  of 
this  form  of  energy  (which  can  be  used  to  drive  ma- 
chinery, to  heat  a  wire,  <fec.)  is  the  chemical  action  be- 
tween the  zinc  and  acid.  If  no  electricity  were  produced, 
e.g.  if  the  zinc  were  put  alone  into  sulphuric  acid,  an 
amount  of  heat  equivalent  to  the  electricity  would  appear 
in  the  liquid  for  the  same  amount  of  zinc  dissolved.  This 
is  the  simplest  form  of  voltaic  cell.  The  zinc  is  called  the 
positive  element,  the  coppei',  the  negative  element.  As  a 
rule,  when  any  two  substances  are  brought  together  in 
this  way,  the  one  assumes  the  positive  condition  and  the 
other  the  negative.  The  chemical  elements  can  thus  be 
arranged  in  an  electro-chemical  series,  beginning  with  the 
most  positive,  and  ending  with  the  most  negative  ele- 
ments. In  this  series,  any  element  is  positive  as  regards 
succeeding,  but  negative  as  regarding  preceding,  elements. 

Electro-chemical  Series  :  Cs,  Rb,  K,  Na,  Li,  Ba,  Sr, 
Ca,  Mg,  Al,  Mn,  Zn,  Fe,  Ni,  Co,  Cd,  Pb,  Sn,  Bi,  Cu,  Ag, 
Hg,  Pt,  Au,  H,  Si,  Te,  Sb,  C,  B?  Or,  As,  P,  I,  Br,  01,  F, 

N,  Se,  S,  O. 

439.  Electrolysis. — Most  compounds  which  can  be 
got  in  the  liquid  condition,  either  by  fusion  or  by  solu- 
tion, can  be  decomposed  by  a  current  of  electricity  allowed 
to  pass  through  the  liquid  between  two  wires  dipping  into 
it.  Compounds  decomposable  by  electricity  are  called 
electrolytes.  In  such  decompositions,  one  portion  of  the 
compound  is  set  free  at  the  wire  (the  negative  pole)  com- 
ing from  the  positive  element  of  the  battery,  and  the 
other  at  that  (the  positive  pole)  connected  with  the  nega- 
tive element.  If  the  liquid  is  a  metallic  compound  the 
metal  always  appears  at  the  negative  pole  (Art.  261),  and 
the  non-metal,  or  negative  radical  (NO3,  SO4,  &c.)  at  the 


392  ELECTRICITY. 

positive  pole.  The  positive  part  of  a  compound  goes  with 
the  current.  The  wires  or  plates  used  for  conducting 
the  current  into  and  out  of  the  liquid  are  generally 
called  electrodes.  The  positive  electrode  is  often  eaten 
away  by  the  negative  product  of  electrolysis.  Plat- 
inum resists  in  most  cases,  but  is  attacked  by  chlorine, 
<fec.  Various  secondary  actions  occur  when  solutions  are 
electrolysed.  For  example,  when  a  solution  of  sodium 
sulphate  is  electrolysed,  sodium  is  set  free  at  the  negative 
pole,  but  immediately  decomposes  water,  and  thus  hydro- 
gen and  caustic  soda  are  the  final  products.  At  the  nega- 
tive pole,  oxygen  and  sulphuric  acid  appear,  since  the 
radical  SO4  cannot  exist  by  itself.  Generally,  when  solu- 
tions of  alkaline  salts  are  electrolysed,  an  alkali  and 
hydrogen  appear  at  the  negative  electrode,  while  an  acid 
and  oxygen  appear  at  the  positive.  The  nascent  hydrogen 
and  oxygen  may  exert  reducing  or  oxidising  action  on 
the  other  substances  present  in  the  solution. — The  liquids 
of  the  human  body  contain  salts  of  sodium  and  potas- 
sium, so  that  when  a  current  of  electricity  passes  through 
any  part,  alkali  collects  around  the  negative,  and  acid 
around  the  positive,  needle.  The  alkali  exerts  its  well- 
known  solvent  action  on  the  tissues,  while  the  acid  co- 
agulates the  albuminous  substances  and  thus  causes  the 
needle  to  become  more  or  less  firmly  fixed  in  the  tissues. 
Advantage  is  taken  of  these  phenomena  in  the  destruc- 
tion of  tumors,  <fec.,  by  electrolysis  of  the  diseased  tissue. 


SYSTEMATIC    TKSTlMi.  393 


CHAPTER     XXIII. 


ANALYSIS— TOXICOLOGY. 

440.  Systematic  Testing.  —  Most  of  the  sub- 
stances which  are  met  with  in  practice  can  be  arranged 
under  the  three  heads  of  bases,  acids,  and  salts.  In  sys- 
tematic testing  of  unknown  substances,  two  cases  may 
occur:  (1)  the  substance  may  be  pure — a  chemical  in- 
dividual, or  (2)  it  may  be  mixed — including  two  or  more 
chemical  individuals.  The  first  case  is  the  only  one 
which  admits  of  treatment  here ;  the  case  of  mixtures 
requires  more  space  than  we  have  at  our  disposal.  For 
purposes  of  systematic  testing,  the  metals  are  classified 
as  at  p.  270,  and  the  first  step  is  to  determine,  by  the 
use  of  group  reagents,  to  which  of  the  six  groups  the 
metal  whose  base  or  salt  is  under  examination  belongs. 
This  determined,  further  testing  shows  which  metal  of 
the  group  is  present.  The  acids  can  be  arranged  in  four 
groups : 

1.  Organic  acids,  which  char  when  heated  : — Tartaric, 
citric,  succinic,  benzoic,  &c. 

2.  Inorganic  acids,  the  barium  salts  of  which  are  in- 
soluble or  sparingly  soluble  : — Sulphuric,  carbonic,  phos- 
phoric, oxalic,  boriCf  sulphurous,  chromic,  &c.     The  group 
reagent  is  barium  nitrate. 

3.  Inorganic  acids,  the  silver  salts  of  which   are  in- 
soluble :  —Ferrocyanic,   ferricyanic,    hydrocyanic,    hypo- 
chlorous  {chloride  precipitated),  hydriodic,  hydrobromic, 


D94  DISSOLVING    THE    SUBSTANCE. 

hydrochloric,    thiosulphuric,    nitrous,  [boric,    oxalic,   sul- 
phurous].    Argentic  nitrate  is  the  group  reagent. 

4.  Acids  which  give  no  precipitate  with  the  group 
reagents  :  —  Nitric,  chloric,  acetic,  and  (in  sufficiently 
dilute  solutions)  oxalic,  boric,  sulphurous,  and  nitrous. 

NOTE. — In  testing  for  the  acids,  the  nature  of  the 
metal  (already  discovered)  must  be  considered.  For 
example,  a  solution  of  plumbic  acetate  might  char  on 
evaporation  and  ignition.  On  the  addition  of  calcic 
chloride,  a  white  precipitate  would  appear  at  once.  This 
might  lead  to  the  conclusion  that  the  acid  is  tartaric, 
whereas  the  precipitate  is  plumbic  chloride.  In  such 
cases,  the  metal  must  be  removed  as  carbonate,  by  boil- 
ing with  solution  of  sodic  carbonate,  filtering,  and  neu- 
tralising part  of  the  filtrate  with  nitric  acid,  to  test  for 
all  acids  but  nitric  and  chloric,  and  another  with  hydro- 
chloric acid,  to  test  for  these  two  acids. 

441.  Dissolving  the  Substance. — Try  the  solu- 
bility of  small  portions  of  the  substance  in  water,  in  hy- 
drochloric acid,  in  nitric  acid,  and  in  aqua  regia.  If  a 
solution  is  obtained  in  any  of  these  solvents,  make  the 
analysis  according  to  the  following  tables.  If  the  sub- 
stance is  insoluble  in  water,  but  soluble  in  hydrochloric 
acid,  it  may  be  a  phosphate,  oxalate,  or  citrate,  in  which 
case  it  would  be  re-precipitated  by  ammonia  (Table  B), 
and  might  then  be  mistaken  for  alumina.  In  such 
cases,  these  three  acids  must  be  tested  for  separately  (see 
articles  132,  184,  190).  If  the  substance  is  insoluble  in 
all  the  above-mentioned  solvents,  it  is  pi'obably  one  of 
the  following : — Baric  sulphate,  strontic  sulphate,  silica, 
calcic  fluoride,  alumina,  stannic  oxiJe.  argentic  chloride, 


PRACTICAL    HINTS — CHEMICAL    TOXICOLOGY.  395 

plumbic  sulphate,  carbon,  ferric  oxide.  These  are  all 
white,  excepting  the  last  two.  They  may  be  tested  for 
by  special  tests,  described  in  the  preceding  pages  (see 
Barium,  &c.) 

442.  Practical  Hints. — Use  small  quantities,  both 
of  reagents  and  of  liquids,  to  be  tested.     Add  the  re- 
agents a  little  at  a  time.     Remember  that  in  chemical 
actions,  equivalents  of  the  substances  must  be  used  in 
order  to  complete  the  actions. 

Excess  of  a  reagent  means  more  than  enough  to  com- 
plete the  chemical  action  which  the  reagent  brings  about. 
It  does  not  necessarily  mean  a  large  quantity.  Chemi- 
cal actions  take  time,  and  when  solutions  are  very  cold 
the  time  is  longer.  This  is  a  matter  to  be  considered 
in  the  more  delicate  tests. — Work  slowly,  and  handle 
apparatus  gently.  Clean  apparatus  as  soon  as  possible 
after  using  it,  as  it  is  more  difficult  to  clean  after  standing 
for  some  time.  Explain  every  test,  and  write  equations 
where  possible. 

443.  Chemical    Toxicology.  -  -  Poisonous    sub- 
stances may  be  conveniently  classified  as  follows  : — 

1.  Corrosives. — Corrosive  sublimate,  concentrated  acids 
(sulphuric,  nitric,  hydrochloric,  oxalic,  &c.),  alkaline  sub- 
stances (caustic  potash,  caustic  soda,  ammonia,  and  the 
carbonates  of  these  bases),  acid,  alkaline,  and  corrosive 
salts  of  the  metals  (potassium    bi-sulphate,  alum,  anti- 
mony trichloride,  silver  nitrate,  <fcc.),  and  carbolic  acid. 

2.  Irritants. — Arsenic  compounds,  dilute  acids,  phos- 
phorus, many  metallic  salts,  e.g.,  those  of  antimony,  lead, 
zinc,    copper,    and    chromium,    and    many    organic    sub- 


39G  CHKM1CAL    TOXICOLOGY. 

stances,  e.g.,  elaterium,  gamboge,  aloes,  colocynth,  croton 
oil,  cantharides,  &c. 

3.  Xeurotics. — Prussic    acid,    opium    (including    the 
opium    alkaloids,    e.g.,    morphine),    strychnine,    aconite, . 
belladonna,  &c. 

4.  Gaseous  Poisons. — Chlorine,  bromine,  hydrochloric 
acid,  hydrofluoric  acid,  sulphur  dioxide,  nitrogen  oxides, 
ammonia,   carbon    dioxide,    carbon    monoxide,  coal    gas 
(carbon  monoxide  and  acetylene),  sulphuretted  hydrogen, 
anaesthetics,   vapours  of  hydrocarbons   (e.g.,   of  mineral 
naphtha). 

Tests  for  most  of  these  substances  have  been  described 
in  the  preceding  pages.  These  tests  can  be  applied  with- 
out any  difficulty  where  a  definite  chemical  substance  is 
to  be  examined,  but  in  many  cases,  the  substance  to  be 
tested  for  poisons  is  a  complicated  mixture,  such  as  the 
contents  of  a  stomach,  or  some  article  of  diet.  In  such 
cases,  special  methods  must  be  used,  in  order  to  separate 
the  poison  from  the  organic  matter  which  would  obscure 
the  tests.  A  very  useful  method  is  that  of  dialysis, 
since  all  poisonous  substance  diffuse  through  a  moist 
membrane,  while  albuminous  substances  do  not.  The 
following  method  may  be  employed  in  testing  for  the 
common  poisons.  Boil  with  hydrochloric  acid  for  some 
time,  filter,  and  heat  a  portion  of  the  clear  filtrate  for 
half  an  hour  with  a  small  piece  of  bright  copper.  Any 
mercury,  arsenic,  or  antimony  present  will  be  deposited 
on  the  copper.  Remove  the  copper,  wash,  and  dry  it 
carefully,  and  heat  it  in  a  narrow  glass  tube  held  aslant. 
Mercury  forms  a  metallic  coat  on  the  tube,  arsenic  oxid- 
ises and  forms  a  white  crystalline  deposit  at  some  dis- 
tance from  the  metal,  while  antimony  forms  a  white  de- 


C'HKMK'AL    TOXI('i)l,(i(;v.  3D" 

])osit  near  the  metal.  Marsh's  test  (pp.  161  and  307) 
may  be  applied  to  other  portions  of  the  liquid.  To  test 
for  lead  and  copper,  pass  sulphuretted  hydrogen  through 
the  warm  liquid  for  some  time.  A  black  precipitate 
may  be  PbS  or  CuS.  Filter,  wash,  dissolve  the  precipi- 
tate in  aqua  reyia,  and  test  as  at  pp.  280  and  297. 
Zinc  is  tested  for  by  treating  the  liquid  with  excess  of 
ammonia,  filtering,  and  adding  sulphuretted  hydrogen  to 
the,  filtrate.  A  white  precipitate  (ZnS)  indicates  zinc. 
Chromium  is  detected  as  at  p.  333. — Acid  and  alkaline 
substances  can  be  tested  for  by  observing  whether  a  large 
quantity  of  alkali  or  acid  is  required  to  render  the  sub- 
stance neutral.  For  hydrochloric,  nitric,  and  sulphuric 
acids,  see  pp.  103,  86,  and  133.  As  small  quantities  of 
the  salts  of  these  acids  are  naturally  present  in  articles 
of  food,  &c.,  only  large  quantities  should  be  looked  upon 
as  abnormal. — To  test  for  oxalic  acid,  precipitate  the 
clear  liquid  with  lead  acetate,  collect  the  precipitate, 
wash,  mix  with  water,  decompose  with  sulphuretted 
hydrogen  (PbC2O4  -f  H2S  =  PbS  -f  H2C2O4),  and  test 
the  filtered  liquid  for  oxalic  acid  (p.  210). — Prussic  acid 
can  often  be  recognised  by  its  smell,  especially  when  the 
substance  is  treated  with  a  little  sulphuric  acid.  The 
tests  at  p.  184  can  be  made  by  mixing  the  suspected  sub- 
stance with  sulphuric  acid  in  small  porcelain  dishes,  and 
placing  over  these  inverted  dishes  or  watch  glasses 
moistened  with  argentic  nitrate,  caustic  soda  solution, 
and  ammonium  sulphide  respectivel) .  The  acid  is  vola- 
tilised upon  these,  and  the  tests  are  completed  as  at  p. 
184. — Any  substance  containing  free  phosphorus  is  lum- 
inous in  the  dark.  The  luminosity  is  especially  apparent 
when  the  substance  is  distilled  alonsr  with  water  in  the 


3!)<S  CHEMICAL    TOXICOLOGY. 

dark.  A  luminous  ring  appears  in  the  neck  of  the  con- 
denser.— The  examination  of  an  organic  mixture  for 
poisonous  alkaloids  is  a  process  too  extensive  for  treat- 
ment here.  They  can  be  tested  for  separately  as  at  pp. 
250,  251,  and  252. 


NOTE. — The  following  ANALYTICAL  TABLES  are  to  be  used  in 
examining  solutions  containing  not  more  than  one  metal  and  one 
acid. 


ANALYTICAL    T.M'.LKS. 


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APPENDIX. 


TABLE  OF  SOLUBILITIES. 

100  parts  (by  iceight)  of  water  at  15°C  dissolve,  of 


NAMES. 

FORMULAS. 

PARTS. 

CjH.O,.. 

all  proportions. 
25.  OO 
all  proportions, 
all  proportions. 
11.4 
14.3 
12.3 
59.7 
27.5 
35. 
0.666 
200 
4.22 
75.5 

loot* 

0.001  (1) 
0.00 

loot 

0.033 

loot 

3.33 
0.0071 
45. 
0.0263 
3.4 
8.1 
0.00 
0.00025 
0.1  (?) 
0.5 
332 

loot 

lOOt 
0.00 
3.00 
5.3 
3.226 
140 
58 
0.0018 
75. 
0.1(7) 

C.H12O  

"      ethyl    

C..H.O... 

cO)  ...: 

<NH4.)2A1,  (SO4)4.  24H2O 
K2Cr2(SO.),.  24H.O  
K2A12  (S04)4.  24H2O  

"     chrome  
"     potash  

N,ll,,C,b,,.. 

Nk,.tf  !..'... 

"         chloroplatinate  

(NH.),PtCl6  

NH4NO3 

"         oxalate  '.  
"         sulphate  

(NH4)2C,04.H20  
(NH4),S04..  ..  

SbCl3  

Sb,O,  

AgCl  

AgNO3  

Ag2O  

Arsenic  acid  

H3AsO4  

ASnO,       

BaCO3   

BaCl,.2H,O  

fluosilicate  

BaSiF,  

hydroxide  

Ba(OH)2  

nitrate  

Ba(NO,)2  

phosphate  

Ba,(PO4)2  

sulphate  

BaSO.  •... 

C6H8.    . 

Benzoic  acid  
"      aldehyde  

CgHj.COOH  
C.jH.s.COH  
Bid,  

Bi(NO,),.3H2O 

"        subnitrate  
Boraeic  acid  

BiNOJOH),  
H,BO,  

Borax  
Bromine  

Na2B4O,.10H2O  
Br  

Cadmic  chloride  

CdCl,  

"       sulphate  

CdSO4  

Calcic  carbonate  

CaCO3  

**     chloride  

CaCl,  

"     citrate  

Ca^CaHsO-MI^O...... 

lOOt  means  very  soluble. 


406 


APPENDIX. 


NAMES. 

FORMULAS. 

PARTS. 

Ca(OH)., 

0  137 

CaC2O4"2H.,O  

0.00 

Ca.,(PO.;2 

0  00 

CaSO4  2H2O 

0  238 

CaC4H4O6.4H.,O 

0  016 

0  1 

c  *  H  "(Sn 

6  67 

cs  ° 

0  1 

CO2..     . 

0  19744 

CO  

0  00305 

C2HC1,O.H2O  

100  (">> 

Cl  .  .  .  .  !  

0.7515 

CHC1,  

0  1  ('>.) 

C6HHO-  H2O 

133  00 

CoCl2.6H.,O  

loot 

Co(NO,).,,"6H.,O 

loot 

CoSO4.7H2O  "  .    . 

33  5 

Cu(NO,)2  3H2O  

loot 

"      sulphate  

CuSO4.5H2O  
C6H10O.  

39.5 

loot 

Ether  (ethylic)                       

11  00 

Fe2  cf  2 

loot 

Fe2iSO4),  

loot 

FeSO..7H.,O  

70 

Fe(NH4)2(SO4)26H,O.  .. 

19 

C,H6O5  H2O 

1  00 

HC1 

75  72 

HCN  

all  proportions. 

H  ,  

0.00017 

"         peroxide  

H.,O2  

all  proportions. 

Iodine  

0.182 

CHI, 

0  01  (">) 

Li2C'O.,  

0.8 

LiCl  

78 

"     urate  
Magnesic  chloride  
"        oxide  

Li2C.H2N40.,  
MgCl,,  
MgO  

10-  (I) 

130. 
0.0018 

"        am.  phosphate  
"        sulphate  

MgNH4PO4.6H2O  
MgSO4.7H2O  
MnCO, 

0.0067 
67.5 
0  013 

MnCl2  4H2O         .... 

200  00 

MnSO,.4H2O 

130  00 

Marsh  gas  

CH4  

0.00286 

Mercuric  chloride  

Hg.Cl.,  

6.98 

'  '        oxide  

HgO  

0.00 

Mercurous  chloride  
"         nitrate  

Hg2Cl2  
Hg2'NO3)2  

0.00 

loot 

"         oxide  

Hg2O  

0.00 

HgoSO4 

0.2 

C,,H.JNO..H.O  

0.05 

CJ7H1S(NOVHC1  

5.00 

Nickel  sulphate  
Nitric  acid  ... 

NiSO4.7H2O  
HNO,  

67.0 
all  proportions. 

Nitrobenzene.           

C6H-'.NO.,  

0.1  (?)       4 

Nitrogen  

"        monoxide  

N  O 

0.1533 

Oxalic  acid  .  .  . 

C.,H.,O..2H.,O  .  .  . 

11. 

APPENDIX. 


407 


NAMES. 

FORMULAS. 

PARTS. 

O  

0.0057 

loot 

Pb  C.,H,O.,  ...H.,O. 

59  00 

PbBr2  .... 

0  05  (  ') 

PbCO3     

0.002 

PbCl  

0.9 

0.082 

Pb  NO.,  

53.00 

"        sulphate  

PbSO4.  ."  

0.0044 

Potassio  bichromate  

K.,Cr.,0  
KBr  "   

10.3 
100  ?) 

K,CO,   

106.4 

KHCOn 

25  00 

'        chlorate  

KC1O,'.  

6.00 

K.PtCl,, 

1.10 

'       chloride.   

KC1  

32.8 

'        chromate  ... 

K.,Cr04  
KCN                    

62.0 

loot 

'        fluosilicate  
'        ferri<"vanide  

K2SiF,  

0.12 
40.5 

'        ferrocvanide  

K4FeiCN  vSH.O.  •• 

30.00 

"        hydroxide  

KOH  

213. 

"        iodide  

KI  

140. 

"        nitrate  

UNO,  

26. 

K2C.,O4.H2O 

33. 

"        h  vdric  oxalate  

KHC.,O4.H.,O    

4.8 

"        permanganate  

KMnO4         

6.6 

"        sulphate  

K»SO4  

9.2 

'  '        tartrate  

151.5 

"       hydric  tartrate  

KHC4H4OH  

0.453 

KNaC4H,Ofi.4H2O     . 

59. 

"        urate   

2.24 

KHC5  H"  N  0 

0  125 

Quinine  ;  .  . 
"        sulphate  
'  '       bi  -sulphate  

C.,0H24S2O2.3H26  
C20H.4N.,02'.,.H,S04... 
C2,,H.,4N2O.,.H2S64  
C-H.O, 

0.05 
0.14 
10.0 
0.5  (">) 

Sodic  acetate  

NaC,H'1O2.3H2O      

30. 

carbonate  

Na.,£ov10H2O.. 

60. 

hydric  carbonate  

NaHCO3.  .       .     . 

10.5 

NaClO3 

95 

chloride  

Nad  

35.76 

hydroxide  

NaOH  

loot 

NaNO., 

84. 

phosphate  
sulphanthnonate  
sulphate  
thiosulphate  
Stannic  chloride  

Na2HPO4.12H20  
Na,SbS4.9H,O  
Na2SO4.10H*O  
Na.,S2O3.5H2O  
SnCl4         

14. 
34.5 
35.96 
116. 

loot 

Stannous  chloride  
Starch  

SnC1..2H2O  

270. 
0  00 

Strontic  carbonate  

sfco'3°  5 

0.00554 

"        chloride  

SrCl, 

5000 

"        hydroxide  

SnOHu 

0457 

"        nitrate  
"        phosphate  
"        sulphate  

Sr.NO,),  
Sr,<P04U  
SrSO4                 

55. 
0.00 
0.0145 

Strychnine  

C21H2.,N2O.,  

0.05 

408 


APPENDIX. 


NAMES. 

FORMULAS. 

PARTS. 

Strychnine  chloride  

C21H22N2O2.HC1  

5.  (!) 

300 

CgH^Od!* 

75 

"      Milk  

20. 

0.00 

Sulphuretted  h3'drogen   .... 

H2S 

0.4918 

Sulphur  dioxide  

SO2                

14.33 

Sulphuric  acid  

H2SO4         

all  proportions. 

Tartar  emetic  

2tKSbOC4H4O6  .H.,O.     . 

7.0  (1) 

Tartaric  acid  

138- 

C10H1B 

0.05  (?) 

Urea  

CON.,H4  

100.00 

Uric  acid  

C,H4N40,  
ZnCO, 

0.0067 
0.00 

"    chloride  

ZnCl., 

loot 

"    oxide  

ZnO  " 

0.00 

"    sulphate  

ZnSO4.7H2O  

50.5 

INDEX. 


A. 

Acid,  sulphurous,  35,  125,    126,    128. 

;etates,  205. 

"    tannie,  243. 

^etone,  201. 

"    tartaric,  214,  215. 

>etylene,  180. 

"    thiosulphuric,  134. 

*etylene  series,  174. 

"    uric,  188. 

:id,  acetic,  193,  204. 

"     valerianic,  196,  207. 

antiinonic,  305. 

Acids,  35,  153 

arsenic,  159 

"        dibasic,  132. 

arsenious,  156,  157. 

haloid,  103. 

benzoic,  239,  240. 

Air,  composition  of,  72. 

boric,  or  boracic,  260,  262. 

Air,  impurities  of,  75. 

bromic,  111. 

Albuminoids,  254. 

butyric,  193,  207. 

Alcohol,  amylic,  194,  196. 

fluosilicic,  259. 

"        benzvlic,  239. 

formic,  192,  202,  203. 

"        butyiic,  196. 

gallic,  243. 

ethylic,  177,  192,  193,  195. 

hippuric,  240,  242. 

"        methylic,  190. 

hydriodic,  114. 

"       propylic,  196. 

hydrobromic,  110. 

"        salicylic,  248. 

hydrochloric,  101. 

Alcohols,  190. 

hydrocyanic,  182,  183. 

diacid,  208. 

hydrofluoric,  117,  267. 

"        isomeric,  197. 

hypobromous,  111. 

Aldehyde,  acetic,  195,  200. 

hypochlorous,  104,  105. 

"        benzoic,  239. 

hypophosphorous,  151. 

"        cinnamic,  245. 

hvposulphuruus,  127. 

"        salicylic,  248. 

iodic,  116. 

"        valerianic,  196. 

lactic,  193,  214. 

Aldehydes,  200, 

manganic,  344. 

Alkaline  earths  352. 

meconic,  250. 

Alkaline  reaction,  80. 

metaphosphoric,  146,  149. 

Alkalis,  368. 

muriatic,  101. 

Alkaloids,  249, 

nitric,  83. 

"        artificial.  253. 

nitrous,  94. 

Allotropy,  53. 

oleic,  207. 

Alloys,  264. 

orthophosphoric,  146. 

Alum,  ammonia,  336. 

osmie,  315. 

"        chrome,  332. 

oxalic,  171,  203,  208,  211. 

"        iron,  326. 

palmitic,  207. 

"        potash,  336. 

perchloric,  108. 

Alums,  332,  336. 

phosphoric,  35,  85,  146,  147. 

Alumina,  335. 

phosphorous,  146,  150. 

Aluminium,  334. 

picric,  86,  239. 

"        bronze,  334. 

propionic,  197. 

"       chloride,  334. 

prussic,  100,  182. 

"        sulphate,  337. 

pyrophosphoric,  149. 

Amalgams,  284. 

salicylic,  243. 

Amides,  187. 

selenic,  140,  312. 

Amines,  198,  249. 

selenious,  140. 

Ammonia,  76,  80. 

silicic,  257,  258. 

Ammonio-cupric  sulphate,  162. 

stearic,  207. 

Ammonium,  82,  383. 

succinc,  211.                   [131,134. 

"        benzoate,  242. 

sulphuric,    85,    127,    129,     130, 

"        bromide,  111. 

410 


INDEX. 


Ammonium  carbonate,  385. 

chloride,  384. 
'        cyanate,  185. 
'        formate,  183. 
'        hydroxide,  383. 
'        nitrate,  87,  90. 
'        oxalate,  210. 

phosphate,  386. 
'        picrate,  239. 
'        sulphate,  384. 

sulphide,  100,  386. 
Amygdalin,  239,  247. 
Amyl  nitrite,  196. 

"      acetate,  197. 
Amyloses.  220,  227. 
Anaesthetics,  177,  200. 
Analysis,  138,  269,  393. 
Aniline,  236. 

dyes,  232,  237. 
Anthracene,  247. 
Antidotes,  145,  254. 
Antimony,  101,  304. 
butter  of,  306. 
pentoxide,  305. 
"        sulphate,  304. 
"        tetroxide,  304,  305. 
'•        trichloride,  306. 

trioxide,  217,  305,  306. 
"        trisulphide,  136,305. 
Antimoniuretted  hydrogen,  307. 
Antlmonyl,  217,  307. 
"        chloride,  306 
"        potassic  tartrate,  217,  306. 
Antipyrine,  253. 
Antiseptics,  193,  238. 
Apatite,  142,  358. 
Aqua  Fortis,  83. 
Argentic  salts  (nee  Silver). 
Argol,  215. 
Aromatic  series,  232. 
Arsenic,  61,  155. 
"        antidotes  to,  158. 

flowers  of,  156. 
"        pentoxide,  156. 
"        sulphides,  160. 
"        trichloride,  132,  162. 
"       tri-iodide,  162. 
"        trioxide,  155,  156,  157,  159. 
Arsenical  wall  paper,  162. 
Arseniuretted  hydrogen,  161. 
Atmosphere,  67,  68. 
Atomic  heat,  48. 

weights,  47,  48. 
"       theory,  44. 
Atomicity,  62. 
Atoms,  44.  47. 
Atropine,  253. 
Avogadro's  Law,  45. 


B. 


Balsams,  244. 
Barium,  361. 


Barium  dioxide,  64,  362. 

"        fluosilicate,  259. 

"        nitrate,  363. 

"      '  sulphate,  132,  361. 
Barometer,  67. 
Bases,  36,  37. 
Basicity,  86,  87,  203. 
Beer,  194. 
Beet  sugar,  379. 
Benzene,  233 
Benzene  series,  232 
Bismuth,  300. 

"        carbonate,  303. 

"        chloride,  302. 

"        oxide,  301. 

"        nitrates,  301. 

"        trioxide,  302. 

trisulphide,  303. 
Bismuthyl  carbonate,  302. 

chloride,  303. 
Black  ash,  371.      , 

"        wash,  285. 
Blue  vitriol,  295. 
Bleaching,  100,  106,  125. 

powder,  105,  138,  177,202,357 
Boiling  points,  23. 
Bone  ash,  143,  147,  358. 
black,  14S. 
oil,  143,  168. 
Bones,  142. 
Borates,  261. 
Borax,  260,  261. 
Boron,  260. 

"        trioxide,  201. 
Boyle's  Law,  69. 
Brimstone,  121. 
Bromides,  111. 
Bromine,  98,  108. 
Brucine,  252. 
Brunswick  green,  296. 
Burnett's  disinfecting  fluid,  340. 
Butter,  206. 

c. 

Cadmium,  297. 
"        amalgam,  284. 

compounds,  298,  299. 
Caesium,  388. 
Calcium,  353, 

acetate,  204. 
"        carbonate,  355,. 
"        chlorate,  381. 
"        chloride,  356. 

citrate,  218,  219. 
"        fluoride,  117,  118. 
"        hydroxide,  354. 
"        hypophosphite,  151. 
"        oxalate,  209. 
"        oxide,  353. 
•'        phosphate,  358. 

sulphate,  98   132,  356. 
"        thiosulphate,  123. 


INDEX. 


411 


Calomel,  285. 

Calculations,  chemical,  50. 
Camphor,  245. 
Caramel,  222 
Carbohydrates,  220. 
Carbon^  164. 

bisulphide,  113,  144,  172. 
"        compounds,  167. 

dioxide,  43.  75,  77,  168. 
"        monoxide,  43,  171. 
Carbonates,  170,  171. 
Carbonyl  chloride,  187. 
Carboxyl,  203. 
Celluloid,  231. 
Cellulose,  230. 

Cerium  and  its  compounds,  349. 
Charcoal,  35,  165,  166. 
Chalk,  355. 
Charles'  Law,  70. 
Chemical  action,  1,  6. 
Chemism,  38. 
Chemistry,  1. 

"        organic,  167,  185. 
Chloral,  201. 
Chlorates,  107. 
Chlorides,  101, 103. 
Chlorine,  98,  99. 
"        oxides,  104. 
"        oxygen  acids,  105. 
Chloroform,  177,  178,  202. 
Chromates,  331,  332. 
Chrome  alum,  331,  332. 

"        yellow,  279, 
Chromium,  329 
"        chloride,  333 

hydroxide,  333. 
"        sesquioxide,  333. 
"        trioxide,  332. 
Chromyl  chloride,  116. 
Cinchonine,  251. 
Clay,  258,  337. 

Clemens'  "  bromide  of  arsenic,"  162. 
Coal,  166. 
Coal  tar,  232. 
Cobalt,  346. 

"        chloride,  347. 
"        nitrate,  346. 
Cocaine,  252 
Collodion,  231. 
Combination  by  volume,  45. 
Combining  weights,  42,  57. 
Combustion,  34,  38,  59,  76. 
Compounds,  6. 
Conine,  250. 

Conservation  of  matter,  41. 
Condy's  fluids,  344,  345. 
Copper,  293. 
Copperas,  321 
Coprolites,  142. 
Corrosive  sublimate,  288. 
Cream  of  tartar,  215. 
Creosote,  191,  238,  239. 
Cresols,  23S. 


Crystallisation,  3,  24,  27. 
Crystalloids  and  colloids,  326. 
Cupric  nitrate,  85. 

'•        oxide,  296. 

"        sulphate,  295. 
Cuprous  chloride,  172,  295. 
iodide,  115. 
oxide,  222,  295. 
Cyanides,  181,  183,  184. 
Cyanogen,  181. 

D. 

Decomposition,  doable,  270. 

Definite  Proportions,  Law  of,  42,  44. 

Deliquescence,  27. 

Dextrin,  225,  229, 

Dextrose,  225,  248. 

Dialysis,  63,  326. 

Diamonds,  164. 

Diffusion,  62,  63. 

Digitalin,  248. 

Dimorphism,  355. 

Dissociation,  131. 

Distillation,  3,  80. 

Donovan's  solution,  162. 

Dulong  &  Petit,  Law  of,  48. 

Dynamite,  86,  213. 


Earths,  352. 

Effervescing  powders.  219 
Efflorescence,  27. 
Electricity,  29,  390. 
Electrolysis,  392. 
Electroplating,  281,  313. 
Elements  and  Compounds,  5,  6. 
Elements,  Table  of  the,  39. 
Elutriation,  2. 
Emery,  335. 
Epsom  salts,  120,  364. 
Equations,  chemical,  49. 
Equivalents,  42. 
Essence  of  mirbane,  235. 
Essences,  artificial,  197. 
Ether,  acetic,  195. 

"        ethyl  (sulphuric),  198,  200. 

"        nitrous,  195. 
Ethers,  198. 
Ethylene,  178. 
Evaporation,  21,  23,  24. 
Explosion,  60. 


Fatty  acids,  202,  207. 
Fats  and  oils,  207. 
Fehling's  test,  223  226. 
Ferments,  76,  192. 
Fermentation,  192,  193. 
Ferric  arsenite,  158. 

chloride,  158,  324. 

hydroxide,  158,  325,  327. 


412 


INDKX. 


Ferric  nitrate,  322,  326. 

oxide,  320,  327. 

sulphate,  322,  326. 

sulphocyanate,  186. 

tartrate,  328. 
Ferrous  ai  senate,  159,  323. 

bromide,  324. 

carbonate,  181,  323. 

iodide,  324 

lactate,  215. 

nitrate,  326. 

oxalate,  210. 

oxide,  320,  322. 

phosphate,  160,  324. 

sulphate,  321. 

sulphide,  136. 
Filtration,  2. 
Flame,  176. 
Fluorides,  117,  118. 
Fluorine,  117. 
Fluor  spar,  117,  259. 
Fluosilieates,  259. 
Formulas,  48,  49,  62. 
Freezing  and  melting,  18. 
Freezing  mixtures,  26. 
Fusel  oil,  194,  196,  197. 
Fusion,  4. 

G. 

Galena,  120,  273. 
Gas  liquor,  80. 
Gas,  defiant,  178> 
Gases,  2. 

"        Law  of  Diffusion,  64. 

"        Molecular  weight  of,  47. 

"        Solubility  in  water,  34. 

"       volume  of,  51,  70,  71,  72. 
German  silver,  348. 
Glass,  274,  377. 

"        etching  of,  118. 

"        soluble,  258. 
Glauber's  salt,  371. 
Glucose,  194,  220,  225,  226,  244. 
Glucoses,  225. 
Glucosides,  227,  247. 
Glycerine,  203,  211,  213. 

"        of  borax,  262. 
Glycogen,  229. 
Glycol,  208. 
Gold,  312. 

Goulard's  extract,  276. 
Gram-molecule,  51,  71. 
Graphite,  164. 
Green  vitriol,  321. 
Group  reagents,  270. 
Guano,  188. 
Gum,  British,  229. 

"        benzoin,  240,  241. 
Gums,  230. 
Gun-cotton,  86,  230. 
Gun-powder,  88. 
Gypsum,  24,  98,  120,  356. 


H. 

Halogens,  98. 
Heat,  12. 

"        expansion  by,  14. 

"        latent,  20,  23,  26. 

"        specific,  21. 
Hydrocarbons,  173. 

•'        saturated,  176. 

"        unsaturated,  178. 
Hydrogen,  55,  58,  60,  61. 
dioxide,  64,  65. 

"        persulphide,  139. 

sulphide,  136. 
Hydroxides,  58,  190. 
Hydroxy -acids,  214. 
Hydroxyl,  58,  190. 
Hydroxylamine,  83. 
Hypochlorites,  105, 106. 
Hypophosphites,  151,  152. 
Hyposulphites,  127. 


Indigo,  128,  134,  246. 

Infusion,  27,  28. 

Iridium,  315. 

Ink,  244,  347. 

Inosite,  227. 

Inulin,  228. 

Iodides,  112,  115. 

Iodine,  98,  112,  113.  126,  135. 

"        chlorides,  116. 

"        pentoxide,  116. 
lodoform,  178,  195. 
Iron,  56,  318. 

"        dialysed,  325. 

"       galvanised,  320. 

"        rust,  320. 

"        pyrites,  120,  121,  124, ]55,  321. 
Isomerism,  179,  180. 
Isomorphism,  159,  332,  336. 


Jalapin,  249, 
Jet,  167. 


J. 


K. 


Kairine,  253. 
Kaolin,  337. 
Kelp,  108,  112. 
Ketones,  197,  201. 

L. 

Lactose,  223. 
Lamp-black,  165. 
Latent  heat,  20,  23. 
Laughing  gas,  90. 
Lead,  272. 

"        acetates,  103,  206,  275,  276. 

"        black,  164. 

"        bromide,  111. 


INDEX. 


413 


Lead  chloride,  103,  277. 

Methvl  salicvlate,  191,  243. 

"        fhromate,  331. 

Methylated  spirit,  191. 

hydroxide,  88. 

Microcosmic  salt,  148,  386. 

"        iodide,  115. 

Minium.  274. 

nitrate,  85,  87,  95,  27(i. 

Mixtures,  separation  of,  2. 

oxides.  274,  27S. 

Mohr's  salt,  323. 

'        phosphate,  147. 

Molasses,  221. 

plaster,  212,  278. 

Molecules,  43,  46. 

poisoning,  279. 

Morphine,  250. 

red,  274.  275. 

Mortars  and  cements,  359. 

'        sulphate.  132,  278. 

Mucilages,  vegetable,  230.              [90. 

"       white,  276. 

Multiple  proportions,  Law  of,  43,  44, 

Ledoyen's  disinfecting  fluid,  276. 

Levulose.  225,  227. 

N. 

Lime,  chloride  of,  357. 
'•        milk  of,  355. 
"        quick.  353. 
"        slaked,  354. 

Naphthalene,  246. 
Nessler's  reagent,  289. 
Neutralisation.  88. 

Liquids,  2. 
Listerism.  238. 

Nickel,  347.  348. 
Nicotine,  250. 

Litharge,  205,  274. 
Lithium  and  compounds,  387. 

Nitrates,  85,  86,  88. 
Nitre,  (gee  Saltpetre). 
Nitrites,  94. 

1 

Nitrobenzene,  235,  240. 

Nitrogen,  72,  79. 

• 
Maceration,  27. 

"        compounds  in  air,  76. 
dioxide,  92,  93,  129,  173. 

Magnesia,  366. 

"        monoxide,  90,  91. 

"        alba,  365. 

"        pentoxide,  95. 

Magnesium,  363. 

tetroxide  (peroxide),  95. 

'         carbonate,  365. 

"        trioxide,  93. 

"        citrate,  219. 
"        oxide,  366. 

Nitroglycerine,  86,  213. 
Notation,  chemical,  47. 

"        sulphate,  98,  364. 

Manganates.  343. 

0. 

Manganese,  341. 
"        dioxide,  33,  99,  341,  342. 

Oil  of  bitter  almonds,  236,  240. 

"        salts,  343. 

"        cinnammon,  245. 

Marsh  gas,  175. 
Marsh's  test,  161. 

"        clov  63^45. 
"        vitriol,  129,  130. 

Massicot,  274. 

"        wintergreen,  243. 

Matter,  three  states  of,  2. 

Oils,  drying,  274. 

Mercurial  poisoning,  291. 
Mercuric  chloride,  288. 

Oils,  essential,  245. 
Olefines,  174,  178. 

"        iodide,  115.  162,  289. 

Olein,  212. 

"       nitrate,  287. 

Orpiment,  160. 

"        oxide,  32,  112,  287,  289. 

Osmium,  315. 

sulphate,  285,  287,  288. 
"        sulphide.  290. 
Mercurous  bromide,  111. 

Osmose.  63. 
Oxalates,  210. 
Oxides,  35. 

"        chloride,  103,  285. 

Oxidising  agents,  65. 

"        iodide,  115,  286. 

Oxygen,  32,  33,  34,  73. 

"        nitrate,  103,  285. 

Oxymel,  223. 

"        sulphate  132. 

Oxy  -salts,  307. 

Mercury,  283. 

Ozone,  52. 

Metallurgy,  263. 

Metals,  39,  263. 

. 

"        classification  of,  269. 

Palladium,  315. 

"        compounds  of,  265. 

Palmitin,  212. 

Methane,  175,  176. 

Paraffin  oil,  174. 

ethyl,  192. 

"        wax,  137,  174. 

«'        amine,  198. 

Paraffins,  168,  174. 

.'        chloride,  198. 

Pnris  green,  158. 

414 


1NDKX. 


Pearl  ash,  379. 
Percolation,  28. 
Permanganates,  344. 
Peroxides,  267. 
Petroleum,  Hi*. 
Phenol,  237. 
Phosphates,  142,  148. 
Phosphine.  152. 

Phosphoretted  hydrogen,  150,  152. 
Phosphorus,  35,  142,  143. 
"        oxides  of,  145. 
"        pentachloride,  153,  190. 
"        pentoxide,  43,  146. 
"        tests  for,  145. 
"        tribromide,  110. 

trichloride,  150,  153. 
"        trioxide,  43, 146. 
Phosphoryl,  153. 
Photography,  136,  299. 
Plaster  of  Paris.  356. 
Platinum,  60,  313. 
Plumbago,  164. 

Plumbic  compounds  (gee  Lead). 
Poisons,  395. 
Porcelain,  337. 
Potash  lye,  381. 

Potassium,  378.  [330. 

"        bichromate,  33,  65,  115,  143, 
"        bromide,  110,  382. 
carbonate,  379. 
acid  carbonate,  380. 
chlorate,  32,  99,  104,  107,  381. 
chloride,  98. 
cyanide,  182. 
ferricyanide,  328. 
ferrocyanide,  171,  181. 
fluosilieate,  107,  256,  259. 
hydroxide,  380, 
hypobromite,  111. 
iodide,  112,  115,  278,  383. 
manganate,  343. 
nitrate,  381. 
perchlorate,  108. 
permanganate,  344. 
sulphate,  83. 
sulphocarbonate,  173. 
sulphocyanate,  185. 
Potato  oil,  196." 
Pottery,  337. 
Powder  of  Algaroth,  306. 
Precipitation,  269. 
Proteids,  254. 

Prussian  blue,  181,  246,  328. 
Ptyalin,  223. 
Pure  substances,  2. 


Quantivalence,  62. 
Quicksilver,  284. 
Quinine,  251. 


R. 

Radicals,  compound,  58. 
Realgar,  160. 
Reducing  agents,  12.r>. 
Reduction,  36. 
Resins,  244. 
Respiration,  77. 
Rochelle  salt,  216. 
Rosaniline,  237. 
Rouge,  327. 
Rubidium,  388. 

s. 

Saccharine,  168,  242. 

Saccharoses,  220. 

Sago,  228. 

Sal  ammoniac,  80,  384. 

Salicin,  243,  248,  251. 

Salt  cake,  370. 

Salt  in  air,  76. 

"        of  sorrel,  210. 

"        of  tartar,  379. 

"        radical,  89. 
Saltpetre,  24,  79,  86,  381. 

"        Chili,  83,  374. 
Salts,  37,  88,  89. 

"        ethereal,  192. 

"        normal  and  acid,  133. 

"        oxygen,  268. 

"        sulphur,  269. 
Sal  volatile,  385. 
Saponiflcation,  212. 
Scale  compounds.  327. 
Scheele's  green,  158. 
Sea  water,  98. 
Seidlitz  powder,  216. 
Selenium,  140. 
Silica,  256.  257,  259. 
Silicates,  256,  257. 
Silicon,  256. 

"        dioxide,  257. 

"        tetrafluoride,  259. 
Silver,  280. 

"        bromide,  108,  111. 
chloride,  103,  135. 

"        cyanide,  281. 

"        iodide,  J15. 

"        nitrate,  103,  282. 

"        sulphate,  132,  281. 

"        sulphide,  282. 
Smalt,  300,  346. 
Soaps,  208,  211,  212. 
Soda,  baking,  373 

"        caustic,  373. 

"        washing,  371. 
Sodium,  369. 

"        acetate,  175,  206. 

"        antimonite,  305. 

"        arse  n  ate,  159. 


INDKX. 


415 


Sodium  arsenite,  158. 

Sujrar,  milk,  223. 

"        benzoate,  242. 

of  lead,  206,  275. 

"        bihorate,  261. 

Sulphates,  120,  132. 

"        bromide,  376. 

Sulphides,  120,  139. 

"        carbolate,  237. 

Sulphites,  128. 

carbonate,  112,  371 

Sulphostannates;  311. 

"        bi-carbonate,  373. 

Sulphur,  35,  120. 

chloride,  369. 

"        bromides,  140. 

'        cvanide,  181. 

"        chlorides,  139. 

'        hydroxide,  57,  373. 

dioxide,  123,  124. 

hypochlorite,  106. 

flowers  of,  121. 

'        hyposulphite,  !:>. 

iodides.  140. 

"        manganate,  343. 

liver  of,  376. 

nitrate,  83,  374. 

milk  of,  123. 

"        nitrite.  94. 

"        oxides  of,  124. 

'         oxalate,  209. 

"        oxygen  acid  of,  127. 

'        oxides,  369. 

"        precipitated,  123. 

'        phosphate,  375. 

salts,  160,  269. 

'        silicate,  258. 

"        trioxide,  126. 

'        sulphantimonite,  305. 

Sulphuretted  hydrogen,  100,  114,  136. 

sulphate,  370. 

Symbols,  chemical,  47. 

'        bi-sulphate,  84. 

Synthesis.  180. 

sulphide,  376. 

sulphite,  126.  375. 
'        thiosulphate   135 

T. 

valerianate,  196,  207. 

Tannin.  244. 

Solids,  -2. 

Tartar,  215. 

Solubility,  25. 

Tartar  emetic,  217,  306. 

Solution,  3,  24. 

Tellurium,  140. 

"        chemical,  25. 

Temperature,  15,  16. 

latent  heat  of,  26. 

"        absolute,  70. 

"        saturated,  25. 

"        critical.  34. 

"        supersaturated,  26. 

"        of  ignition,  37. 

Specific  heat,  21. 

Terpenes,  244. 

"        weight,  8,  9,  10. 

Thalline.  253. 

Spectroscope,  389. 

Thermometers,  16,  17. 

Spectrum  analysis,  388. 

Thiosulphates,  135. 

Spirit,  proof,  195. 

Tin,  308. 

Spirit,  sweet,  of  nitre,  195. 

Tin,  butter  of,  310. 

Spontaneous  ignition,  38. 

Tin'salt,  309. 

Stannic  chloride.  309,  310. 

Tinctures,  28. 

"        oxide,  309. 

-    Tincal,  261. 

sulphide,  311. 

Toxicology,  chemical,  395. 

Stan  nous  chloride,  288,  309. 

Turnbull's  Blue,  328. 

"        oxide,  309. 

Turpentine,  244. 

"       sulphide,  311  . 

Starch,  227,  228. 

u. 

Stearin,  212. 

Strontium,  360. 

Urates,  189. 

"        hydroxide,  361. 

Urea,  111,  186,  187,  188. 

"        nitrate,  360. 

Urine,  188. 

"        oxide.  361. 

sulphate,  132,  361. 

V. 

Strychnine,  252. 

Sublimation,  3. 

Valence,  61,  62. 

Substitution,  179,  232. 

Ventilation,  170. 

Sugar,  beet,  221,  37i». 

Verdigris.  206,  294,  297. 

"        cane,  221. 

Vermilion,  290. 

fruit,  227. 

Vinegar,  204. 

"        grape,  225. 

Vitriol,  blue,  144,  295. 

im-ert,  221,  225. 

green,  121,  134,  321. 

"        malt,  224. 

white,  340. 

416 


INDKX. 


w. 

Water,  action  on  lead,  274. 

"        analysis,  289. 

"        boiling  point,  22. 

"        composition  of,  30. 

."        decomposition  of,  28,  29. 
distilled,  12. 

"        expansion  of  by  heat,  14. 

"        freezing  of,  19. 

"        in  air,  75. 

"        latent  heat  of,  20. 
Weights,  combining,  42. 

"        and  measures,  6,  7. 
White  precipitate,  290. 
Wines,  194. 
Wine,  spirits  of,  192. 


Wood,  distillation  of,  191,  204. 
"        spirit,  191. 
tar,  168. 

Y. 

Yellow  wash,  288. 


Zinc,  56,  57.  338. 
"        acetate,  340. 
"        butter  of,  339. 
"        carbonate,  340. 
"        chloride,  339. 
"        oxide,  37,  338. 

sulphate,  37,  207,  340,  365. 


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