PRACTICAL AND ANALYTICAL 
CHEMISTRY. 
HENRY TRIMBLE, Ph.M. PRACTICAL AND ANALYTICAL 
CHEMISTRY. 
A COMPLETE COURSE IN CHEMICAL 
ANALYSIS. 
BY 
HENRY TRIMBLE, Ph.M., 
» M 
PROFESSOR OF ANALYTICAL CHEMISTRY IN THE PHILADELPHIA COLLEGE 
OF PHARMACY. 
FOURTH EDITION. 
WITH ILLUSTRATIONS. 
PHILADELPHIA: 
P. BLAKISTON, SON & CO., 
No. 1012 Walnut Street. Copyright, 1892, by P. BLAKISTON, SON & CO. 
PRESS OF WM. F. FELL A CO., 
1220-24 SANSOM STREET, 
PHILADELPHIA. Preface to the Fourth Edition. 
In the preparation of this edition the author has made such 
additions and changes as, in his experience with the use of 
the book in large classes, he has found to be desirable. 
Slight additions have been made to the charts for the 
separation of bases, wherever it has been found that fuller 
explanations were needed. It is very important, however, 
that the student should not be allowed to become too mechani- 
cal in this part of the work, and thereby overlook the reasons 
for carrying out the processes as described. This difficulty 
may be obviated by the teacher frequently questioning the 
student, and in large classes delivering short lectures in which 
the separations are explained and the reactions involved 
written out, so as to emphasize the fact that every step in an 
analytical operation is strictly a chemical process. 
The use of formulas and symbols for the reagents, instead 
of the full chemical name, has been continued in this edition. 
They may be inconvenient at first, but the student soon learns 
them, almost without effort, and thus becomes familiar with 
what he should know, but should not be compelled to mem- 
orize. H. T. 
January, 1892. 
VII Preface to the First Edition 
The increased amount of time devoted by students of 
pharmacy and medicine to analytical chemistry has directed 
more attention to the subject of imparting instruction in this 
science. 
The object of the present volume is to place before the 
student as compact a course as possible, in order to enable 
him to become familiar with the subject in the necessarily 
limited time at his disposal. 
The author’s experience has led him to believe that a 
study of Qualitative Analysis should be preceded by some 
experience in the preparation of the more important gases 
and a few of the salts. Such practice requires the student to 
familiarize himself with the construction of apparatus as well 
as with the processes of filtration, evaporation, crystallization, 
ignition, etc. 
The examples for preparation may be increased at the 
option of the instructor. 
In Part II the student should perform the reactions of each 
group, and then be furnished with a solution containing 
some, or all, of the bases of the group. This should be 
followed by a solution in which he should search for all the 
elements previously studied; such practice being repeated 
until he can correctly determine all the bases present. X 
PREFACE TO THE FIRST EDITION. 
In order to enable the student to see the comparative effect 
of the group reagents readily, a summary has been intro- 
duced at the end of each group. This he should be able to 
write out, without the use of the book, before attempting to 
analyze a group solution. By progressive steps he is thus 
led rapidly on to successfully examine the more complex 
solutions for both bases and acids. 
The grouping of the bases is, to a certain extent, new, but 
it places together those elements which are very closely 
related, and, in addition, adapts each group to the time of 
one lesson ; which may be repeated if desirable. 
In the Part devoted to the gravimetric and volumetric 
analysis, the examples are limited in number; but this much 
is intended to give the student an opportunity to learn the 
methods with the aid of an instructor, so that he may after- 
ward pursue the subject alone, with the aid of a book like 
Fresenius’s “ Quantitative Analysis.” 
In conclusion, it is but just to state that many of the works 
on Qualitative Analysis have been consulted. 
Those of Attfield and Miiter furnished many valuable hints, 
but Fresenius’s admirable work has been used as authority, 
and, to those who wish to pursue the subject in greater detail 
than here offered, it is recommended for reference. 
H. T. 
Philadelphia, August igth, i88j. CONTENTS. 
PAGE 
PART FIRST—PRACTICAL CHEMISTRY, . . .17 
SECTION I.—Preparation and Properties of some Non- 
metals and their Compounds, 17 
Hydrogen, Chlorine, Hydrochloric Acid, Bromine, Iodine, Oxygen, 
Sulphurous Oxide, Nitrogen, Ammonia, Nitric Acid, Carbon 
Dioxide, 17-27 
SECTION II.—Preparation of Salts 28 
Potassium Chloride, Potassium and Sodium Tartrate, Ammonium 
Nitrate, Ammonium Oxalate, Calcium Phosphate, Magnesium 
Sulphate, Magnesium Carbonate, Magnesium Oxide, Alum- 
inium Hydrate, Ferrous Sulphate, Ferric Sulphate, Ferric 
Hydrate, Copper Sulphate, Lead Acetate, . . . 28-32 
PART SECOND-QUALITATIVE ANALYSIS, . . 35 
SECTION I.—Bases, 35 
GROUP I.—Reactions of Potassium, Sodium, Lithium, Ammonium, 35 
Summary of Group I, 37 
Analysis of Group I, 37 
GROUP II.—Reactions of Barium, Strontium, Calcium, Magne- 
sium, 38-39 
Summary of Group II, ........ 40 
Analysis of Group II, ....... 40 
Chart for Analysis of Groups I and II 41 
Precautions and Observations on Chart, 42 
GROUP III.—Reactions of Manganese, Zinc, Cobalt, Nickel, . 42-45 
Summary and Analysis of Group III, 45 
Chart for Analysis of Groups I to III, inclusive, ... 46 CONTENTS. 
PAGE 
GROUP IV.—Reactions of Iron, Cerium, Aluminium, Chromium, 47-49 
Summary and Analysis of Group IV, . . . . .50 
Chart for Analysis of Groups I to IV, inclusive, . . . 51 
Precautions and Observations on Chart, 52 
GROUP V.—Reactions of Arsenic, Antimony, Tin, Gold, Platinum, 52-56 
Summary and Analysis of Group V 57 
Chart for Analysis of Groups I to V, inclusive 58 
Precautions and Observations on Chart, .... 59 
GROUP VI.—Reactions of Mercury(ic), Bismuth, Copper, Cad- 
mium, ........... 59-61 
Summary and Analysis of Group VI, 62 
Chart for Analysis of Groups I to VI, inclusive, ... 63 
Precautions and Observations on Chart, . . . . .64 
GROUP VII.—Reactions of Silver, Mercury(ous), Lead, . . 64-66 
Summary and Analysis of Group VII, . . . . .66 
Chart for Analysis of Groups I to VII, inclusive, ... 67 
Precautions and Observations on Chart, 68 
Special Observation, 69 
SECTION II.—Acids, 70 
GROUP I.—Reactions of Acids Hydrochloric, Hydrobromic, Hydri- 
odic, Hydrofluoric, Hydrocyanic 70-72 
Summary of Group 1 72 
Analysis of Group I, 72 
GROUP II.—Hypochlorous, Chloric, Water, Hydrates, Oxides, Hy- 
drosulphuric, Sulphurous, Sulphuric, Thiosulphuric, Nitric, 
Hypophosphorous, Orthophosphoric, Pyrophosphoric, Meta- 
phosphoric, Boric, Carbonic, Silicic, Hydroferrocyanic, Hy- 
droferricyanic, 73-79 
Analysis of Group II 80 
GROUP III.—Acetic, Valerianic, Stearic, Oleic, Lactic, Oxalic, Suc- 
cinic, Malic, Tartaric, Citric, Carbolic, Benzoic, Salicylic, 
Gallic, Tannic, 82-85 
Analysis of Group III 86 
SECTION III.—Detection of Bases and Acids, . . 87-92 
Chart of Solubilities, 93 CONTENTS. 
XIII 
PAGE 
SECTION IV.—Reactions and Tests of Organic Com- 
pounds, 94 
Alkaloids, 97 
Neutral Principles, 99 
PART THIRD—QUANTITATIVE ANALYSIS, . 103 
SECTION I.—Gravimetric Estimation, .... 103 
Preliminary Directions 103 
Examples in Gravimetric Estimation, Barium, Chlorine, Copper, 
Sulphuric Acid, Potassium, Nitric Acid, Magnesium, Calcium, 
Carbonic Acid, Iron, 104-109 
SECTION II.—Volumetric Estimation, . . . no 
Examples, Oxalic Acid, Sodium Hydrate, Potassium Bichromate, 
Iodine, Sodium Hyposulphite, Silver Nitrate, . . 110-115 
Table of Elements, Symbols and Atomic Weights, . . 115-116 
Index, 117—119 PART FIRST. 
PRACTICAL. PRACTICAL AND ANALYTICAL 
CHEMISTRY. 
PART I.—PRACTICAL. 
SECTION I. 
PREPARATION AND PROPERTIES OF SOME NON-METALS AND 
THEIR COMPOUNDS. 
HYDROGEN, H. 
Preparation.—Place a few fragments of zinc in a flat-bottom 
flask of about one-fourth liter capacity; cover the zinc with 
water and adapt a cork, through which pass two tubes (Fig. i), 
one just reaching through the cork and bent so the long end 
may be dipped under water, the other running directly from 
a short distance above the cork nearly to the bottom of the 
flask, so as to be below the surface of the liquid. The upper 
end of this tube should have a small funnel placed in it, or be 
enlarged by softening in the flame, inserting and revolving a 
file or similar instrument, previously warmed. 
Add, slowly, a small quantity of sulphuric acid, through 
this tube, and notice an immediate effervescence, with the 
escape of bubbles through the water in which the exit tube 
dips. Fill a test tube with water, and, keeping the open end 
under the liquid, bring it over the tube, so as to collect the 
gas. When full, close with the thumb, and, bringing the 
mouth of the tube near a flame, quickly remove the thumb 
and allow the gas to ignite. It will burn quietly if the gas be 
pure, but with a slight explosion if it be mixed with air. 18 
PRACTICAL CHEMISTRY. 
Properties.—This gas is Hydrogen, and its physical prop- 
erties may now be studied by observing that it is insoluble in 
water, and without odor, color or taste. 
Experiment I. Collect a tube full, and holding it, covered, 
in a vertical position, bring a lighted taper a short distance 
above its mouth and remove the cover; the gas will ignite, 
showing its great levity. 
Experiment II. Another tube, similarly filled, is held in 
an inverted position, and the cover removed ; it will be found, 
Fig. i. 
even after the lapse of some time, that the hydrogen at the 
mouth of the tube maybe ignited, thus demonstrating that the 
gas is too light to come down and out the mouth of the tube. 
These two experiments have also demonstrated the combus- 
tibility of the gas, which, when pure, burns quietly, with a 
colorless flame. If, however, it be mixed with air and flame 
applied, a violent explosion ensues. Therefore, the tube from 
the generator should never be brought near a flame until it is 
certain all the air has been expelled. This is determined by 19 
PREPARATION AND PROPERTIES OF NON-METALS. 
trying a test tube full; if it burn quietly, the jet may be lighted. 
This preca7ition should always be observed. 
Experiment III. On bringing a lighted taper to the mouth 
of a tube full of hydrogen, the gas is ignited, but on pushing 
the burning taper up into the gas its flame is extinguished, 
thus showing that while hydrogen is combustible it is not a 
supporter of combustion. 
Experiment IV. Continue the addition of acid to the zinc 
until the latter is nearly all dissolved; disconnect the appa- 
ratus, pour the liquid on a filter, collect the filtrate in a small 
beaker or evaporating dish, concentrate and set aside for 
twenty-four hours, to crystallize. These crystals are zinc 
sulphate, ZnS04, the result of a combination of the sulphuric 
acid and the zinc, as follows:— 
Zn + H2S04 = ZnS04 + Hr 
CHLORINE, Cl. 
Preparation.—In a flask, arranged so that heat may be 
applied (F'ig. 2), place a small quantity of manganese dioxide, 
add hydrochloric acid, agitate well, to moisten all the powder 
on the bottom, and apply heat. A yellowish-green gas is 
evolved, which, being somewhat soluble in water, may be 
collected by dowmvard displacement, that is, by running the 
delivery tube to the bottom of the receptacle, loosely covered ; 
the heavy gas displaces the lighter air. 
If the evolution be moderately active it may be collected 
over warm water, as only a small loss occurs. Care should 
be taken to avoid inhaling the gas, by passing it, when not 
collecting, into a solution of potassium or sodium hydrate. 
The following expresses the reaction in symbols :— 
MnOa -f- 4HCI = MnCl2 -j- 2H20 -f- CI2. 
Properties.—Experiment I. Pass the gas into water; it 
is absorbed ; if this be continued until the water is saturated, it 
will be found to have absorbed about twice its volume of the gas; 
the Aqua Chlori of the Pharmacopoeia is the resulting product. 
Experiment II. A tube full of the gas held with mouth 
upward, and a lighted taper applied, fails to ignite. Push the 20 
PRACTICAL CHEMISTRY. 
taper into the gas; it is extinguished, or only burns with a 
small, dense, smoky flame, the result of a combination of the 
chlorine with the hydrogen of the wax, liberating the carbon. 
Experiment III. Into a tube full of the gas put a piece 
of brightly dyed calico or blue litmus paper, previously 
moistened; it is rapidly bleached. Writing on paper is simi- 
Fig. 2. 
larly decolorized, but printing is not affected, as it contains 
carbon, in the form of lampblack, which is not acted on by 
the gas. If the experiments be made with chlorine which 
has been passed through sulphuric acid to dry it, and the 
materials are not moistened, no decoloration takes place. 
The process of bleaching by chlorine is one of oxidation ; it PREPARATION AND PROPERTIES OF NON-METALS. 
21 
combines with the hydrogen of the water, forming hydro- 
chloric acid, while the liberated oxygen in the nascent state 
readily attacks the coloring matter, water and a colorless com- 
pound resulting. 
HYDROCHLORIC ACID, HC1. 
Preparation.—Collect one test tube full of hydrogen and 
one of chlorine, bring their mouths together (the hydrogen 
tube above with mouth down, as it is lighter), turn over once 
or twice, so as to thoroughly mix, and open their mouths to 
a flame; a sharp report will occur, with the development of 
strongly acid fumes, which will be recognized by future tests 
as hydrochloric acid:— 
H2 + Cl, = 2HC1. 
To prepare a quantity of it, the apparatus used for the 
preparation of chlorine serves best. Put into the flask some 
sodium chloride (common salt), add sulphuric acid slowly, and 
when the evolution of gas ceases apply a gentle heat. Collect 
by downward displacement or over mercury:— 
NaCl + H2S04 = NaHS04 + HC1. 
The above expresses the reaction when an excess of acid 
has been used, which is preferable, as the resulting acid sodium 
sulphate is easily dissolved out of the flask with water. On 
the large scale the following more economical method is 
used:— 
Properties.—The pungent, suffocating odor and freedom 
from color are noted with its production. 
(NaCl)2 + H2S04 = Na2S04 + (HC1)2. 
Experiment I. A test tube or jar of the gas placed with 
the open mouth under water will so rapidly dissolve that the 
liquid rises in the vessel. This solution of the gas in water, 
when of the proper strength, is the Aciduni Hydrocldoricum, 
U. S. P. 
Experiment II. A piece of moistened blue litmus paper 
held in a tube of the gas is instantly reddened. 
Experiment III. A lighted taper applied to the gas fails 
to ignite it, and is extinguished if lowered into it. 22 
PRACTICAL CHEMISTRY. 
Experiment IV. Bring a rod moistened with ammonia 
over the mouth of a tube full of the gas: dense white fumes of 
ammonium chloride are formed. 
The mixture remaining in the flask dissolved in warm 
water, treated with sodium carbonate so long as effervescence 
occurs, to neutralize the excess of sulphuric acid, concen- 
trated, filtered and set aside to crystallize, yields sodium 
sulphate (Glauber salt). 
BROMINE, Br. 
Preparation.—Pass chlorine into a one per cent, solution 
of potassium bromide. The liquid will become reddish-brown 
in color, owing to separation of bromine which remains dis- 
solved in the water. 
Experiment I. To a portion of this solution add a solution 
of starch ; a yellow color will be developed, due to the forma- 
tion of a compound of starch with bromine. 
Experiment II. To the remainder of the solution in a test 
tube add a small quantity of chloroform and agitate. The 
chloroform will remove the bromine from the water and will 
settle to the bottom of the tube as a deep reddish-brown 
liquid, while the water above will become nearly colorless. 
IODINE, I. 
Preparation.—Pass chlorine into a one per cent, solution 
of potassium iodide. The liquid will become dark in color, 
owing to separation of iodine. Stronger solutions would 
result in the precipitation of the iodine. 
Experiment I. To a portion of the iodine solution, diluted 
with water, add solution of starch ; a blue color will form, due 
to the combination of starch with iodine. When a precipitate 
is formed instead of a blue color it is an indication that the 
iodine solution is too strong, and a more dilute one of the 
iodine must be used. The blue color of the starch iodide is 
destroyed by heat and by an excess of chlorine. 
Experiment II. To another portion of the iodine solution 
in a test tube add a small quantity of chloroform and agitate. preparation and properties of non-metals. 
23 
The chloroform will remove the iodine from the water and 
settle to the bottom of the tube as a deep violet-colored liquid, 
and the water above will become colorless. 
OXYGEN, O. 
Preparation.—Place a few crystals of potassium chlorate 
in a clean and dry test tube, adapt a delivery tube long 
enough to reach under the surface of some water near by. 
Apply a steady flame, taking care not to allow any water to 
be drawn back into the tube; as soon as the bubbles of gas 
escape freely and the air has been expelled, bring a test tube 
filled with water over the escaping gas, and collect it. It is 
Oxygen, produced from the potassium chlorate by heat, 
according to the following reaction:— 
KC103 = KC1 + 03. 
Properties.—The appearance and insolubility of the gas in 
water are noted as it is collected. 
Experiment I. The gas is not ignited by the application 
of a lighted taper. The taper, however, will burn with greatly- 
increased energy if it be plunged into the gas. If the flame 
be extinguished and the taper again brought into the gas, 
provided a spark remain, the taper will be rekindled. 
Experiment II. A piece of charcoal, previously ignited, is 
lowered into the gas; a rapid combustion ensues, and the 
charcoal disappears. Pour some lime water into the tube, 
agitate well; a white precipitate of calcium carbonate is pro- 
duced. If this be tried with oxygen previous to the burning 
of the charcoal, no precipitate will be formed. A number of 
other substances, as sulphur, phosphorus, and even iron, when 
once kindled, will burn in oxygen with great brilliancy, form- 
ing characteristic oxides. 
Experiment III. Take two test tubes, one about twice the 
capacity of the other. Fill the larger with hydrogen and the 
smaller with oxygen, bring their mouths together, and, after 
turning once or twice, to thoroughly mix their contents, open 
them and apply flame. A sharp report is caused by the com- 
bination to form water. 24 
PRACTICAL CHEMISTRY. 
If more than two volumes of hydrogen to one of oxygen 
be present, the surplus remains uncombined; if oxygen be 
present in greater proportion, the excess of it remains. 
SUEPHUROUS OXIDE, S02. 
Preparation.—In a flask, arranged as for the preparation 
of chlorine, place a few copper turnings and cover them with 
concentrated sulphuric acid. Apply a gentle heat, gradually 
increasing until effervescence commences, and collect the 
resulting gas by downward displacement, as in the case of 
chlorine. This gas is sulphurous oxide, and its formation 
takes place according to the following equation :— 
Cu + (H2 S04)2 = Cu S04 + S02 + (H20)2. 
Properties.—As will be noticed during its collection, this 
is a colorless, heavy gas, of a suffocating odor. 
Experiment I. A tube full of the gas inverted in a vessel of 
water will show the ready solubility of it in that liquid. 
Experiment II. A piece of moistened blue litmus paper 
held in a tube of the gas will momentarily become red and 
then white, thus showing its active bleaching property. 
Experiment III. A lighted taper plunged into a tube full 
of the gas is quickly extinguished, indicating that it is neither 
combustible nor a supporter of combustion. 
Experiment IV. Pass the gas into water until the latter 
smells strongly of it, then add to the resulting solution a 
solution of potassium permanganate. The purple color of 
the latter will be destroyed as rapidly as added. 
Sulphurous oxide is an active reducing agent, that is, it 
combines with the oxygen of other compounds. In this 
respect its action is directly opposite to that of chlorine. 
NITROGEN, N. 
Preparation.—The usual method of preparing this gas is 
to deprive air of its oxygen, leaving the nitrogen pure. This 
is best accomplished by placing a small fragment of phos- 
phorus on a cork covered with some fireproof material. Float PREPARATION AND PROPERTIES OF NON-METALS. 
25 
the cork and phosphorus on water, ignite the latter and bring 
over it a bell-jar. The phosphorus combines with the oxygen, 
converting it into phosphorus pentoxide, P205, which dissolves 
in the water present, thus leaving the nitrogen pure. A better 
method for obtaining larger quantities is to heat a mixture of 
potassium nitrite and ammonium chloride dissolved in water. 
When the reaction begins, the temperature must be carefully 
watched, in order to prevent the too rapid evolution of the 
gas 
KN02 + NH4C1 = KC1 + 2H20 + n2. 
Properties.—The physical properties have been observed 
during its preparation and collection. In regard to chemical 
properties, it is inert in the free state. Its compounds, how- 
ever, are very energetic. 
Preparation.—In a test tube or evaporating dish mix equal 
quantities of powdered calcium oxide (quicklime) and ammo- 
nium chloride, with a few drops of water; the odor of ammonia 
will be immediately developed :— 
AMMONIA, NH3. 
CaO + (NH4C1)2 = CaCl2 + H20 + (NH3)2. 
In smaller quantities the gas may be recognized by holding 
over the mixture a strip of moistened red litmus paper; it will 
slowly become blue; or similarly hold a glass rod moistened 
with hydrochloric acid ; dense white fumes of ammonium 
chloride will form. To prepare larger quantities of the gas, 
heat the ordinary water of ammonia, which, at a comparatively 
low temperature, gives it off freely. If it be desired to dry 
the gas, it must be passed over quicklime. Ammonia is col- 
lected by upzvard displacement, that is, by passing the delivery 
tube upward into a jar or test tube inverted over it; being 
lighter than air the latter is displaced. 
Properties.—Experiment I. Place a vessel filled with 
ammonia gas, mouth downward, into some water, and agitate 
slightly ; the water will rise in the vessel rapidly, nearly filling 
it, showing the great solubility of the gas in water. The other 
physical properties, as color, odor, etc., have been noted during 
its collection. 26 
PRACTICAL CHEMISTRY. 
Experiment II. On applying a lighted taper to the gas it 
does not burn ; if, however, it be mixed with oxygen it will 
ignite readily. On introducing the taper into the gas it is 
extinguished. 
NITRIC ACID, HN03. 
Preparation.—Place a small quantity of potassium nitrate 
in a test tube, and cover it with strong sulphuric acid. Apply 
a gentle heat; brown, strongly acid fumes are given off. 
Dilute with a little water and add indigo solution; it is 
decolorized. This is a characteristic test for nitric acid. 
To prepare a larger quantity a retort is used, to which is 
adapted a glass receiver. The nitric acid distills over on the 
application of a moderate heat, forming a reddish-yellow liquid, 
which rapidly attacks and destroys organic matter. 
Two reactions may be employed to represent the produc- 
tion of nitric acid, depending on the relative quantity of the 
materials used. In the first case an excess of sulphuric acid 
gives— 
kno3 + h2so4 = khso4 + hno3. 
In the second case just a sufficient amount of sulphuric acid 
is used to decompose the potassium nitrate— 
(KN03)2 + H2S04 = K2S04 + (HN03)2. 
On the large scale sodium nitrate is now employed in place of 
the potassium salt, on account of its lower price. 
CARBON DIOXIDE, C02. 
Preparation.—The flask used in the preparation of hydro- 
gen will serve for making carbon dioxide. A few pieces of 
marble are placed in the flask, covered with water, and hydro- 
chloric acid added. A brisk effervescence ensues, and the gas 
being somewhat soluble in water is collected by downward 
displacement:— 
CaC03 + (HC1)2 = CaCl2 + H20 + C02. 
Sulphuric acid should not be used, as it forms an insoluble 
calcium sulphate which is very difficult to remove from the flask. PREPARATION AND PROPERTIES OF NON-METALS. 
27 
Properties.—Experiment I. Pour some clear lime water 
into a jar of the gas and agitate ; the solution immediately 
becomes cloudy, owing to formation of insoluble calcium 
carbonate. If more gas be passed into the mixture it will 
become nearly clear again, on account of the solubility of 
the precipitate in carbonic acid. 
Experiment II. Add solution of potassium hydrate to a 
jar of the gas, close and shake well; the gas is absorbed by 
the alkali, as may be shown by placing the mouth of the jar 
under water and removing the stopper, when the water will 
rush in, nearly filling it. 
Experiment III. A lighted taper lowered into the gas is 
immediately extinguished. The same result is accomplished 
by opening the vessel some distance above the flame and 
allowing the gas to flow down upon it. This latter experi- 
ment also illustrates the great density of the gas, which is 
twenty-two times heavier than hydrogen. 28 
PRACTICAL CHEMISTRY. 
SECTION II. 
PREPARATION OF SALTS. 
POTASSIUM CHLORIDE, KC1. 
Preparation.—One or two cubic centimeters of hydro- 
chloric acid, diluted with three or four times its bulk of water, 
are placed in a small beaker glass, and potassium carbonate 
added so long as effervescence occurs, and until after boiling 
(to remove C02) the solution is neutral to litmus paper, that 
is, when the blue litmus paper is not changed to red nor the 
red changed to blue. Filter, evaporate to a small bulk and 
set aside to crystallize. The cubic crystals which separate 
after standing twenty-four hours may be collected on filter 
paper and dried at a moderate temperature:— 
K2C03 + (HCl)t = (KC1), + H20 + C02. 
Potassium chloride is rarely prepared in this manner, except for practice, as it 
occurs largely in nature, and is used for preparing many other potassium salts. 
POTASSIUM AND SODIUM TARTRATE, 
KNaC4H406.4H20. 
Preparation.—Heat, in a porcelain capsule, a solution of 
sodium carbonate, and add to it potassium bitartrate until 
effervescence ceases, and the solution (after the escape of C02) 
is neutral to litmus. On filtering and cooling, crystals of 
Rochelle salt are deposited, rapidly and in small crystals if the 
solution be concentrated, but slowly and in much larger ones 
if the solution be dilute:— 
(Rochelle Salt.) 
NajCOj + (KHC4H406)2 = (KNaC4H406)2 + H20 + C02. 
Note on Calculation.—In order to calculate the amount of each salt to use in 
the above process, we notice the number of molecules of each employed, and 
multiply this by the sum of the atomic weights (that is, by the molecular weights). 
In the above case one molecule of Na2COs = 106, and two molecules of 
KHC4H406 = 2 X I88.i =376.2. Therefore every 106 parts of anhydrous 
sodium carbonate require 376.2 parts of potassium bitartrate, to form Rochelle 
salt. If we have 50 grams of anhydrous sodium carbonate and wish to convert 
it into Rochelle salt, we use the following formula: As 106 : 376.2 : : 50 grams : 
number of grams of potassium bitartrate required, = 177.4 grams. PREPARATION OF SALTS. 
29 
AMMONIUM NITRATE, NH4N03. 
Preparation.—Add to about 20 c.c. of dilute nitric acid, in 
a beaker glass, sufficient ammonia water to give it a distinct 
ammoniacal odor ; filter, concentrate, keeping the ammonia in 
slight excess, and set aside in a cool place for crystals to form:— 
Properties.—These crystals contain twelve molecules of 
water of crystallization, which it is desirable to get rid of before 
using the salt. By exposure to a temperature of 1550 C. the 
water gradually escapes, and the fused or granulated salt is 
ready to be converted into nitrogen monoxide (laughing gas), 
which takes place at about 185° C„ according to the following 
reaction:— 
NH4OH + HN03 = NH4N03 + h2o. 
NH4NOj = n2o + (H20)3. 
AMMONIUM OXALATE, (NH4)2C204. 
Preparation.—Dilute 20 c.c. of solution of ammonia with 
twice its bulk of water, add a solution of oxalic acid until 
neutral, concentrate slightly, filter and set aside to crystallize. 
The crystals may be collected on a filter, and another crop 
obtained by concentrating the “ mother liquor ” :— 
Ammonium carbonate is sometimes used for combining with the oxalic acid, 
but the neutralization is not as easily effected, besides it is not desirable on the 
ground of economy. 
(NH4OH)2 + H2C204 = (NH4)2C204 + (H20)2. 
Preparation.—Finely powdered bone ash is digested for 
a short time with diluted hydrochloric acid. The solution is 
filtered, boiled, and filtered again, if necessary. The filtrate is 
treated with ammonia until it smells strongly of it. Collect 
the precipitate on a filter, wash by pouring on warm water 
until the washings are tasteless, and dry at a low temperature. 
The resulting powder is calcium phosphate, which exists in 
the bone ash and is dissolved by hydrochloric acid, forming 
acid calcium phosphate, as follows :— 
CALCIUM PHOSPHATE, Ca3(P04)2. 
From this solution it is precipitated by ammonia, as follows :— 
Ca3(P04)2 + (HC1)4 = CaH4(P04), + (CaCl2)2. 
CaH4(P04)2 + (CaCl2)2 -f (NH4OH)4 = Ca3(P04)2 + (NH4C1)4 + 4H20. 30 
PRACTICAL CHEMISTRY. 
In addition to the ordinary apparatus with which a student 
supplies himself, there is required a wash bottle (Fig. 3), which 
it is well for every student to construct for himself, as it 
furnishes him valuable practice in cutting and bending glass 
tubing. This bottle is used in washing all precipitates, and is 
convenient as a water supply, which may be kept hot, if desired. 
MAGNESIUM SULPHATE, MgS04./H20. 
Preparation.—To about 5 c.c. of dilute sulphuric acid, 
diluted with five or six times its volume of water, heated in a 
Fig. 3. 
capsule, add powdered magnesium carbonate until effervescence 
ceases, and filter. Concentrate and set aside to crystallize. 
(MgC03)4Mg(0H)2.5H20 + (H2S04)5 = (MgS04)5 + (H20)u + 4C02. 
MAGNESIUM CARBONATE, (MgC03)4Mg(0H)2.5H20. 
Preparation.—On mixing solutions of magnesium sulphate 
and sodium carbonate and boiling, we get magnesium car- 
bonate precipitated, while carbon dioxide escapes. The pre- 
cipitate is very variable in its composition, depending on the 
concentration of the solutions. When the U. S. P. product is 
obtained the following equation expresses the reaction :— 
(MgS04)5 + (Na2COs)5 + (H20)6 = 
(MgC03)4Mg(0H)2.5H20 + (Na2S04)5 + C02. 31 
PREPARATION OF SALTS. 
The precipitate washed with hot water and dried, serves for 
the following example of a compound prepared by ignition. 
Preparation.—Heat some of the magnesium carbonate, 
prepared in the above reaction, in a porcelain crucible until, 
on taking out a small portion, placing in a test tube with a 
little water, heating to remove air bubbles, and adding a drop 
or two of hydrochloric acid, no effervescence is produced. 
This will require some time, and great care is necessary to 
determine when the powder fails to give an effervescence with 
the acid. 
MAGNESIUM OXIDE, MgO. 
(MgC03)iMg(0H)25H30 = (MgO)5 + (C02)4 + (H20)6. 
Zinc oxide may be prepared in a similar manner, from zinc carbonate. This 
differs from the magnesium oxide by being yellow while hot, and very pale yellow 
when cold. 
ALUMINIUM HYDRATE, Al2(OH)6. 
Preparation.—To a solution of alum add a solution of 
sodium carbonate and boil. Allow the precipitate to settle, 
decant the clear supernatant liquid on a filter, add more hot 
water to the precipitate and again decant. Collect the pre- 
cipitate on the filter, wash well with hot water and dry; the 
resulting white powder is Aluminii Hydras, U. S. P. 
A12(S04)3,K2S04 + (Na2C03)3 + (H20)8 = 
Al2(OH)6 + K2S04 + (Na2S04)3 + (C02)3. 
FERROUS SULPHATE, FeS04.;H20. 
Preparation.—Add enough dilute sulphuric acid mixed 
with ten volumes of water to some iron filings, or wire, in a 
beaker, to more than cover them. Allow the reaction to 
proceed, assisted by a little heat, until effervescence ceases. 
Filter from the excess of iron, concentrate, filter and crystallize. 
Fe2 + (H2S04)2 = (FeS04)2 + (H2)2. 
These crystals should be rapidly dried and preserved in well-stopped bottles, as 
they quickly become converted into ferric sulphate on exposure to air. 
FERRIC SULPHATE, Fe2(S04)3. 
Preparation.—To a strong solution of ferrous sulphate add 
one-fourth its bulk of sulphuric acid, heat to the boiling point 
and drop in nitric acid as long as effervescence is produced 32 
PRACTICAL CHEMISTRY. 
and until the resulting liquid becomes of a clear reddish 
brown color. 
(FeS04)6 + (H2S04)3 + (HN03)2 = (Fe2(S04)3)3 + N202 + (H20)4. 
This is the Liquor Ferri Tersulphatis of the Pharmacopoeia; and is the most 
convenient compound to use in the preparation of some of the other iron salts. 
FERRIC HYDRATE, Fe2(OH)6. 
Preparation.—Dilute some of the above ferric sulphate 
solution with an equal bulk of water, add solution of am- 
monia until, after stirring, it smells strongly. The resulting 
precipitate is ferric hydrate, the well-known antidote to arsenic. 
When needed for this purpose, it is sufficient to pour the 
mixture on a muslin strainer, wash once or twice until the 
saline taste nearly disappears from the washings, when the 
compound is ready for use. 
Fe2(S04)3 + (NH4OH)6 = + ( (NH4)2S04)3. 
This preparation should always be freshly prepared when wanted for use as an 
antidote, as it loses H20 on keeping, becoming a mixture of ferric oxide Fe203 
and hydrate. This change takes place, although more slowly, when the compound 
is kept under water. 
COPPER SULPHATE, CuS04.sH20. 
Preparation.—Heat copper turnings for some time, with 
strong sulphuric acid, in a fume closet, until the reaction 
ceases. Dilute with water, filter and crystallize. 
Cu + (HaS04)j = CuS04 + S02 + (H,0)r 
LEAD ACETATE, Pb(C2H302)2.3H20. 
Preparation.—Lead oxide (litharge) is boiled with three or 
four times its weight of acetic acid, in a capsule, adding water 
from time to time, with more acid, if necessary, until most of 
the oxide has disappeared. Filter, concentrate, keeping the 
solution acid, and set aside to crystallize:— 
The solution or crystals should not be exposed to the fumes of the laboratory, 
for if there be only a small quantity of hydrogen sulphide in the room, they will 
become black. 
PbO + (HC2H302)2 = Pb(C2H302)2 + H20. PART SECOND. 
QUALITATIVE ANALYSIS. PART II.—QUALITATIVE ANALYSIS. 
SECTION I. 
BASES 
GROUP I.—POTASSIUM, SODIUM, LITHIUM, 
AMMONIUM. 
REACTIONS OF POTASSIUM (IC). 
Use a solution of potassium chloride (KCl). 
1. PtCl4 added to some of the above solution of KCl, causes 
a yellow crystalline precipitate of K2PtCl6, soluble in excess 
of water. 
PtCl4 + 2KCI = K2PtCl6. 
The delicacy of this reaction is increased by the addition ot 
alcohol. This precipitate should be preserved to compare 
with a similar one under ammonium. 
2. H2C4H406, in concentrated solution, added to another 
portion of the KCl, produces a white crystalline precipitate of 
acid potassium tartrate—KHC4H4Oe, soluble in excess of 
water, readily in hot water, acids or potassium hydrate. 
H2C4H4Oe 4- KC1 = KHC4H406 + HC1. 
The addition of alcohol and violent agitation facilitate the 
formation of this precipitate. This should be compared with 
a similar one under ammonium. 
3. A fragment of potassium salt on the loop of a platinum 
wire, held in the colorless flame of a Bunsen gas lamp imparts 
a violet color. This reaction is interfered with by the presence 
of sodium salts, which color the flame yellow; organic matter 
also colors the flame violet, and should be removed by ignition 
before testing for potassium. The yellow rays of sodium may 
be destroyed by viewing the flame through blue glass. 
4. Potassium salts are not volatile at a low red heat; at a 
white heat they are slowly volatilized. 36 
ANALYTICAL CHEMISTRY. 
REACTIONS OF SODIUM (Na). 
Use a solution of sodium chloride (NaCl). 
1. Sodium salts color the gas flame yellow; so delicate is 
this reaction that the merest traces are revealed by it. 
2. The salts of sodium are not volatile at a low red heat, 
but slowly volatilize at a white heat. 
REACTIONS OF LITHIUM (Li). 
Use a solution of lithium chloride (LiCl). 
1. Na2HP04 added to a strong solution produces, on boiling, 
a white precipitate of lithium phosphate—Li3P04. 
Na2HP04 -f 3LiCl = Li3P04 + 2NaCl + HC1. 
This reaction takes place more readily when the solution is 
first made alkaline with NH4OH. 
2. Lithium salts impart an intense crimson color to the gas 
flame. This is somewhat interfered with by sodium salts, but 
the yellow color of sodium may be excluded by blue glass, 
which if not too dark will allow the crimson rays of lithium 
to pass through. These must not be confused with the 
violet potassium rays, which will pass through a deep blue 
glass. 
3. Lithium salts do not volatilize at a low red heat, but are 
slowly volatilized at a white heat. 
REACTIONS OF AMMONIUM (NH4). 
Use a solution of ammonium chloride (NH4C1). 
1. PtCl4 produces, in strong solution, a yellow crystalline 
precipitate of ammonium platino-chloride—(NH4)2PtCl6. 
2. NaOH, added to the ammonium chloride solution and 
the mixture heated, causes the evolution of ammonia—NH3. 
This NH3 is detected by the odor; by holding near a glass 
rod moistened with HC1, which will produce dense white 
fumes of NH4C1; or by holding in the mouth of the tube a 
strip of moistened red litmus paper, when it will immediately 
become blue. Care must be taken, in this last test, to pre- 
vent any of the alkaline liquid coming in contact with the 
paper, as it would likewise cause the blue color. BASES, 
37 
3- H2C4H406, added to a concentrated solution, produces a 
white precipitate of acid ammonium tartrate—NH4HC4H406, 
soluble in slight excess of water. 
4. Ammonium salts are volatile at a low red heat. 
SUMMARY OF TESTS WITH SOLUBLE SALTS OF GROUP I, 
K 
Na 
Li 
nh4 
PtCl4 
Yellow Precipitate 
No Precipitate 
No Precipitate 
Yellow Precipitate 
h2c4h4o6 
White Precipitate 
No Precipitate 
No Precipitate 
White Precipitate 
Flame 
Violet 
Y ellow 
Crimson 
None 
Volatility 
Not Volatile 
Not Volatile 
Not Volatile 
Volatile 
DIRECTIONS FOR THE DETECTION OF THE 
BASES IN A SOLUTION CONTAINING 
SOLUBLE SALTS OF GROUP I. 
To a small portion of the solution add NaOH and heat: 
NH3will be given off if ammonium salts are present, and may 
be detected by the odor, or by moistened red litmus paper. 
Evaporate another portion of the solution to dryness, 
transfer to a porcelain crucible, and heat until the white fumes 
of ammonium salts cease to be given off Dissolve the 
residue in a few drops of H20 with a drop or two of HC1, and 
add PtCl4; K, if present, will be precipitated. 
A loop of platinum wire dipped in the original solution 
and held in the colorless gas flame will give evidence of Na 
and Li. 
If Na be present in excess, so as to obscure the Li flame, 
evaporate a portion of the original solution to a very small 
bulk, add Na2HP04 and NH4OH and boil, filter off the Li;jP04, 
wash with hot water containing NH4OH until the washings 
cease to impart a yellow color to the flame, and dissolve the 
precipitate in a few drops of HC1. With this solution Li may 
be detected by the flame test. 38 
ANALYTICAL CHEMISTRY. 
GROUP II.—BARIUM, STRONTIUM, CALCIUM, 
MAGNESIUM. 
REACTIONS OF BARIUM (Ba). 
Use a solution of barium chloride (BaCl2). 
1. H2S04 produces an immediate precipitate of barium 
sulphate—BaS04, insoluble in boiling hydrochloric or nitric 
acid. 
BaCl2 + H2S04 = BaS04 + 2HCI. 
2. K2Cr04, even in dilute solutions, causes a yellow precipi- 
tate of barium chromate—BaCr04, soluble in hydrochloric 
or nitric acid, but insoluble in acetic acid. 
BaCl2 -f K2Cr04 = BaCr04 -}- 2KCI. 
3. (NH4)2C03 precipitates white barium carbonate—Ba- 
C03, soluble in acetic acid. 
BaCl2 + (NH4)2C03 -= BaCOg + 2NH4C1. 
4. (NH4)2HP04 produces a white precipitate of barium 
phosphate—BaHP04, soluble in acetic or hydrochloric acid. 
5. (NH4)2C204 causes a white precipitate of barium oxalate 
—BaC204, slightly soluble in acetic acid. This precipitation 
will not take place in very dilute solutions. 
BaCl2 + (NH4)2HP04 = BaIIP04 + 2NH4C1. 
6. A loop of platinum wire moistened with the solution 
colors the gas flame green when held in it. 
BaCl2 + (NHJ2C204 = BaC204 + 2NH4C1. 
REACTIONS OF STRONTIUM (Sr). 
Use a solution of strontium nitrate (Sr(N03)2). 
i. H2S04 forms a white precipitate of strontium sulphate 
—SrS04 immediately, if the solution be strong, but not until 
after some time, if it be very dilute. 
2. K2Cr04 produces no precipitate in the presence of acetic 
acid, but if the solution be made alkaline with KOH, a yellow 
precipitate, strontium chromate—SrCr04, falls. 
3. (NH4)2C03 produces a white precipitate of strontium 
carbonate—SrC03, soluble in acetic and the stronger acids, 
Na2C03 produces the same precipitate. BASES 
39 
4. (NH4)2HP04 forms a white precipitate of strontium 
phosphate—SrHP04, soluble in acids. 
5. (NH4)2C204 causes the precipitation of white strontium 
oxalate—SrC204, sparingly soluble in acetic acid, but readily 
soluble in HC1. 
6. Strontium salts impart an intense red to the colorless gas 
flame. 
REACTIONS OF CALCIUM (Ca). 
Use a solution of calcium chloride (CaCl2). 
1. H2S04, in moderately dilute solutions, forms a white 
precipitate of calcium sulphate—CaS04, soluble in excess of 
water. 
2. (NH4)2C03 or Na2C03 produces a white precipitate of 
calcium carbonate—CaC03, soluble in acids. This precipi- 
tation is not complete unless the solution is boiled. 
3. (NH4)2HP04 causes the precipitation of calcium phos- 
phate—CaHP04, soluble in acetic and the stronger acids. 
4. (NH4)2C204 produces a white precipitate of calcium oxa- 
late—CaC204, insoluble in acetic acid, soluble in hydrochloric 
or nitric acid. 
5. The salts of calcium color the flame yellowish-red. 
REACTIONS OF MAGNESIUM (Mg). 
Use a solution of magnesium sulphate (MgS04). 
I. (NH4)2C03, forms a white precipitate of magnesium-am- 
monium carbonate—MgC03(NH4)2C03, soluble in NH4C1. 
By preceding the addition of the reagent by that of NH4C1, 
a much smaller quantity will suffice to keep the precipitate 
in solution than will be required to dissolve it after once 
formed. 
MgS04 + 2(NH4)2COs = MgC03(NH4)2C03 + (NH4)2S04. 
2. KOH,NaOH or NH4OH produces a white precipitate of 
magnesium hydrate—Mg(OH)2, soluble in NH4C1. 
3. (NH4)2HP04 with NH4C1 and NH4OH produces a white 
crystalline precipitate of ammonium-magnesium phosphate 
—Mg(NH4)P04, slightly soluble in water, but almost insoluble 
in water containing NH4OH. Violent agitation or stirring 40 
ANALYTICAL CHEMISTRY. 
assists in the formation of this precipitate. (NH4)2HAs04 
under similar circumstances precipitates white Mg(NH4)As04. 
MgS04 + NH4OH + (NH4)2HP04 = Mg(NH4)P04 + (NH4)2S04 + H20. 
The ammonium chloride takes no part in the reaction, except 
to keep magnesium hydrate from precipitating. 
4. Magnesium salts impart no color to the flame. 
SUMMARY OF TESTS WITH SOLUBLE SALTS OF GROUP II. 
Ba 
Sr 
Ca 
Mg 
h2so4 
White Precipitate 
insoluble in acids 
White Precipitate 
insoluble in acids 
White Precipitate 
soluble in excess 
of H20 
No Precipitate 
K2Cr04 
Yellow Precipitate 
insoluble in acetic 
acid 
No Precipitate 
unless alkaline 
No Precipitate 
No Precipitate 
(NH4)2COa 
White Precipitate 
White Precipitate 
White Precipitate 
White Precipitate 
soluble in NH4C1 
(NH4)2HP04 
White Precipitate 
White Precipitate 
White Precipitate 
White Precipitate 
(NH4)2C204 
White Precipitate 
in strong solution 
White Precipitate 
in strong solution 
White Precipitate 
in dilute solution 
No Precipitate 
unlessconcentrated 
nh4oh 
No Precipitate 
No Precipitate 
No Precipitate 
White Precipitate 
DIRECTIONS FOR THE DETECTION OF THE BASES IN A SOLUTION 
CONTAINING SOLUBLE SALTS OF GROUP II. 
Add NH4CI, NH4OH and (NH4)2C03, boil and filter. 
Ppt. Ba, Sr, Ca. 
Wash, dissolve in HC2H302, add K2CrC>4, 
filter. 
Filt. Mg. 
Allow to cool, add (NH4)2HP04, agitate. 
White ppt. if Mg be present. 
Ppt. Ba. 
Yellow. 
Filtrate Sr, Ca. 
Add very dilute H2SC>4. allow to 
stand io minutes, filter. 
Ppt. Sr. 
Confirm by 
flame test in 
original 
solution. 
Filt. Ca. 
Add NH4OH and 
(NH4)2C204, 
white ppt. 
In order to thoroughly acquaint the student with the method 
of analysis by this and subsequent charts, the following expla- 
nation is given, in the belief that a careful study of it, until 
perfectly understood, will enable the student to follow all the 
charts which come after, and which are simply enlargements 
of this scheme. 
To a small quantity of the solution in a test tube add an 41 
BASES 
equal volume of solution of NH4C1, close and invert the tube, 
so as to thoroughly mix the contents, add NH4OH until, after 
mixing, the solution smells distinctly of it, and then add 
(NH4)2C03, boil and filter. 
We suppose all four of the bases under consideration to be 
present until their absence is proven; so by this treatment we 
divide it into two groups. The precipitate consists of Ba, 
Sr and Ca, and the filtrate of Mg. A little more (NH4)2C03 
should be added to this filtrate and the whole again boiled, 
to make sure that all the insoluble carbonates have been 
precipitated. If this is found to be the case, (NH4)2HP04 
is added to the filtrate when cold, and Mg, if present, will 
form a white precipitate. 
The first precipitate, consisting of Ba, Sr and Ca, having, in 
the meantime, been washed by forcing a jet of water from the 
wash bottle on it, is dissolved in HC2H302, and to the solu- 
tion, diluted with an equal bulk of water, K2Cr04 added; this 
again divides the solution into a precipitate and a filtrate. 
The yellow precipitate represents Ba, while the filtrate con- 
tains the Sr and Ca. To this filtrate very dilute H2S04 (made 
by adding a small quantity of H2S04 to about ten times its 
volume of water) is added, and the mixture allowed to stand 
ten minutes, for the SrSO.t to form and subside, then filtered, 
which again gives a precipitate, representing Sr, and a fil- 
trate indicating, after the addition of NH4OH and (NH4)2C204, 
the presence or absence of Ca. 
DIRECTIONS FOR THE ANALYSIS OF A SOLUTION CONTAINING 
SOLUBLE SALTS OF ALL THE PRECEDING ELEMENTS. 
Add NH4C1, NH4OH and (NH4)2C08, boil and filter. 
Ppt. Ba, Sr, Ca. 
Wash, dissolve in HC2H302)* add K2Cr04, 
filter. 
t Filtrate Mg, K, Na, Li, NH4. 
Allow to cool, add (NH4)2HP04, agitate, filter. 
Ppt. 
Mg. 
Filtrate K, Na, Li, NH4. 
Evaporate to dryness, ignite, dissolve 
in a small quantity of H20, add 
Na2HP04, boil, filter. 
Ppt. Ba. 
Yellow. 
Filtrate Sr, Ca. 
Add very dilute H2SC>4, allow 
to stand, filter. 
Ppt. Sr. 
Fill. Ca. 
Add NH4OH and 
(NH4)2C204, 
white ppt. 
Ppt. 
Li. 
Confirm 
by flame 
test. 
Filtrate K, Na, NH4. 
Concentrate, add HC1 and 
PtCLj yellow ppt. = K. 
Test tor Na and NH4 
in original solution. 
* Apply flame test to this acetic acid solution for Sr. 
f When sodium is absent Li may he detected by flame test. 42 
ANALYTICAL CHEMISTRY. 
PRECAUTIONS TO BE OBSERVED IN TIIE PRECEDING CHARTS. 
Ammonium chloride must be added in excess, in order to 
keep the magnesium salts in solution when the hydrate and 
carbonate are added. Ammonium hydrate is added until 
the liquid smells of it. Ammonium carbonate is added as 
long as a precipitate is produced. In order to determine this 
to a certainty, a portion of the filtrate is tested with a little 
more of the reagent, when, if no precipitate occurs, the analysis 
may be proceeded with. This precaution of applying more of 
the reagent to a portion of the filtrate, to prove the complete 
precipitation, should be exercised in every case, as it is im- 
portant to add just sufficient of the reagent to accomplish the 
object, but always to avoid a large excess. 
GROUP III.—MANGANESE, ZINC, COBALT, 
NICKEL. 
REACTIONS OF MANGANESE (Mn). 
Use a solution of manganous sulphate (MnS04). 
1. NH4HS, in neutral or alkaline solution, precipitates the 
flesh-colored manganous sulphide—MnS,which, on exposure 
to air, becomes brown. HC1, HN03 and HC2H302 dissolve 
this precipitate, but it is insoluble in alkalies. NH4C1 facili- 
tates the separation of the precipitate, while the salts of the 
organic acids and excess of NH4OH prevent it. 
2. KOH or NaOH produces a whitish precipitate of man- 
ganous hydrate—Mn(OH)2, insoluble in excess. 
MnS04 + 2K0H = Mn(OH)2 + K2S04. 
3. NH4OH likewise precipitates Mn(OH)2, first white, but 
becoming rapidly brown, partly soluble in excess. This pre- 
cipitation is prevented by the previous addition of NH4C1. 
4. (NH4)2C03 produces a white precipitate of manganous 
carbonate—MnC03, soluble in excess in the cold, but re- 
precipitated on warming. 
5. HN03 and red oxide of lead will, on heating and allow- 
ing the precipitate to subside, impart to the supernatant liquid 
a red color, due to permanganic acid—H2Mn2Og. Hydro- BASES, 
43 
chloric acid and chlorides interfere with this reaction. This 
is known as Crum’s process for detecting manganese. 
6. A fragment of a manganese salt fused on platinum foil 
with K2C03 and KN03 will form a green mass containing 
potassium manganate— K2Mn04. 
7. A borax bead (formed by fusing on the loop of a plati- 
num wire some borax until it becomes a clear glass) with 
manganese, in the oxidizing blowpipe flame, becomes violet 
while hot, and a fine amethyst color on cooling. 
REACTIONS OF ZINC (Zn). 
1. NH4HS produces a white precipitate of zinc sulphide— 
ZnS, insoluble in acetic acid, readily soluble in dilute hydro- 
chloric acid. 
Use a solution of zinc sulphate (ZnS04). 
2. KOH, NaOH and NH4OH give white precipitates of 
zinc hydrate—Zn(OH)2, readily soluble in excess, forming 
zincates as Zn(OK)2. Zn(OH)2 is again precipitated on boil- 
ing. 
3. (NH4)2C03 forms a white precipitate of basic zinc car- 
bonate—(ZnC03)2(Zn(0H)2)3, readily soluble in excess. 
5(NH4)2C03 + 5ZnS04 + 3H20 = 
2ZnC03 3Zn(OH)2 + 5(NH4)2S04 -f 3C02. 
4. K2C03 or Na2C03 produces a similar precipitate, insoluble 
in excess. , 
5. On charcoal, before the blowpipe, metallic zinc volatilizes 
and burns, forming an incrustation of oxide, which is yellow 
while hot, becoming white on cooling; if this coating be 
moistened with a drop of cobaltous nitrate, and again heated 
in the outer flame, it becomes green. 
REACTIONS OF COBALT (Co). 
Use a solution of cobaltous nitrate (Co(N03)2). 
i. NH4HS produces a black precipitate of cobaltous sul- 
phide—CoS, insoluble in acetic acid and cold dilute hydro- 
chloric acid. The precipitation is promoted by the presence 
of NH4C1. 
2NH4HS + Co(N03)2 = CoS + 2NH4N03 + I-I2s. 44 
ANALYTICAL CHEMISTRY. 
2. KOH or NaOH produces a blue precipitate of cobaltous 
hydrate—Co(OII)2, insoluble in excess, and becoming pink on 
boiling or exposure to air. 
3. causes a similar precipitate of Co(OH)2, soluble 
in excess with a red color. Sugar and some other organic 
compounds prevent the precipitation by the alkalies. The 
alkaline carbonates behave like their respective hydrates. 
4. KCN gives a red brown precipitate of cobaltous cyanide 
—Co(CN)2, soluble in excess and reprecipitated by HC1; if, 
however, the solution be boiled with only a few drops of HC1, 
the cobaltous cyanide will not be precipitated on the further 
addition of HC1, on account of the formation of potassium 
cobalti-cyanide—KeCo2(CN),2. This experiment should be per- 
formed in a fume closet, in order to avoid inhaling the fumes of 
hydrocyanic acid. 
5. Salts of cobalt color the borax bead blue before the 
blowpipe. 
REACTIONS OF NICKEL (Ni). 
Use a solution of nickelous sulphate (NiS04). 
1. NH4HS forms a black precipitate of nickelous sulphide 
—NiS, insoluble in acetic acid- and cold dilute hydrochloric 
acid. The precipitation is promoted by the presence of 
NH^Cl. 
2. KOH or NaOH produces a green precipitate of nickel- 
ous hydrate—Ni(OH)2, insoluble in excess. 
3. NH4OH gives a similar in excess, with 
a blue color. Sugar and some other organic compounds pre- 
vent the precipitation by the alkalies. The alkaline carbon- 
ates behave like their respective hydrates. 
4. KCN produces a yellowish-green precipitate of nickel- 
ous cyanide—Ni(CN)2, soluble in excess, and reprecipitated 
by HC1 even after boiling, also precipitated, after adding HC1 
and boiling, by KOH. This is used as a method of dis- 
tinguishing and separating nickel and cobalt, the latter not 
precipitating under these circumstances with KOH. This 
experiment should be performed in a fume closet. 
5. The salts of nickel color the borax bead violet while hot, 
and reddish-brown when cold. BASES, 
45 
SUMMARY OF TESTS WITH SOLUBLE SALTS OF GROUP III. 
Mn 
Zn 
Co 
Ni 
nh4hs 
Flesh colored 
precipitate. 
White precipitate 
Black precipitate 
Black precipitate 
KOH 
White precipitate 
insoluble in excess 
White precipitate 
soluble in excess 
Blue precipitate 
insoluble in excess 
Green precipitate 
insoluble in excess 
NH4OH 
White precipitate 
soluble in excess 
White precipitate 
soluble in excess 
Blue precipitate 
soluble in excess 
Green precipitate 
soluble in excess 
(NH4)2C03 
White precipitate 
soluble in excess 
White precipitate 
soluble in excess 
Blue precipitate 
soluble in excess 
Green precipitate 
soluble in excess 
DIRECTIONS FOR THE DETECTION OF THE BASES IN A SOLUTION 
CONTAINING SOLUBLE SALTS OF GROUP III. 
Mn, Zn, Co, Ni. Add NH4OH, then acidify slightly with HC2H302, warm and add H2S until 
the liquid smells strongly; filter. 
Precipitate Zn, Co, Ni. 
Wash, dissolve in HC1 and HNO3, evaporate excess of acid, 
add KOH until distinctly alkaline, filter. 
Filtrate Mn 
Add NH,OH and H2S. 
Flesh-colored ppt. 
Precip. Co, Ni. 
Wash, dissolve in HC1, evaporate excess of acid, 
add KCN until the ppt. first formed re-dissolves, 
boil with a drop or two of HC1 for some minutes, 
and then with more HC1 until all odor of HCN 
disappears and the reaction is acid, then evapo- 
rate excess of acid, add KOH until alkaline, 
filter. 
Filt. Zn. 
Add NH4HS 
white ppt. 
Precip. Ni. 
Green. 
Confirm by dry- 
ing and testing 
with borax bead. 
Filt. Co. 
Evaporate to dryness and test 
with borax bead. 46 
ANALYTICAL CHEMISTRY. 
Mn, Zn, Co, Ni, Ba, Sr, Ca, Mg, K, Na, Li, NH4. 
Add NH4C1, NH4OH and NH4HS, warm, filter. 
Ppt. Mn, Zn, Co, Ni. 
Wash, dissolve in HC1 and HN03, evaporate excess of acid, add NH4OH, 
acidify with HC2H302, add H2S, filter. 
* Filtrate Ba, Sr, Ca, Mg, K, Na, Li, NH4. 
Evaporate to remove NH4HS, add (NH4)2C03, boil, filter. 
Ppt. Zn, Co, Ni. 
Wash, dissolve in HC1 and HNO,, evaporate excess of acid. 
Add KOH until distinctly alkaline, filter. 
Fill. Mn. 
Add NH,OH 
and H2S. 
Flesh-colored 
ppt. 
Ppt. Ba, Sr, Ca. 
Wash, dissolve in HC2H3G2, f 
add K2Cr04, filter. 
Filt. Mg, K, Na, Li, NH,. 
Add (NH4)2HP04, agitate, filter. 
Ppt. Co, Ni. 
Wash, dissolve in HC1, evaporate excess of 
acid and treat with KCN and HC1, according 
to directions on page 45, add KOH 
until alkaline, filter. 
Fill. Zn. 
Add NH4HS, 
warm, allow to 
stand, 
white ppt. 
Ppt. Ba. 
Yellow. 
Filt. Sr, Ca. 
Add very dilute H2S04, allow 
to stand, filter. 
Ppt. Mg. 
White. 
Filt. K, Na, Li, NH4. 
Evaporate to dryness, ignite, 
dissolve in a small quantity 
of H20, add Na2HP04, 
boil filter. 
Ppt. Sr. 
Filt. Ca. 
Add NH4OH anrt 
(NH4)2C204, 
white ppt. 
Ppt. Ni. 
Green. 
Confirm by dry- 
ing and testing 
with borax bead. 
Fill. Co. 
Evaporate to dryness and test 
with borax bead. 
Ppt. Li. 
Confirm by 
flame test. 
Filt. K, Na, NH4. 
Concentrate, add 
HC1 and PtCl4, 
yellow, ppt. = K. 
Test original 
solution for 
Na and NH4. 
* When Ni is present it may remain in solution in the NH4HS, in which case the filtrate will be dark. To avoid this the mixture should be boiled after the addition 
of NH4HS until the excess is driven off, and on filtering yields a colorless filtrate. 
t Apply flame test for Sr. 
DIRECTIONS FOR THE DETECTION OF THE BASES IN A SOLUTION CONTAINING SOLUBLE SALTS OF ALL THE ELEMENTS 
PREVIOUSLY CONSIDERED. BASES. 
47 
GROUP IV.—IRON, CERIUM, ALUMINIUM, 
CHROMIUM. 
REACTIONS OF IRON in ferrous state (Fe)n. 
Use a solution of ferrous sulphate (FeS04). 
i. K4Fe(CN)6, in a neutral or acid solution, gives a white 
(rapidly changing to light blue) precipitate of potassium 
ferrous ferrocyanide—K2Fe2(CN)6, also known as Everett’s 
salt. Alkalies decompose this precipitate, forming ferrous hy- 
drate and a ferrocyanide of the base used. 
K4 Fe(CN)6 + FeS04 = K2Fe2(CN)6 + K2S04. 
2. K6Fe2(CN)12, in a neutral or slightly acid solution, forms 
a dark blue precipitate of ferrous ferricyanide—Fe3Fe2(CN)12, 
known as Turnbull’s blue. If the solutions be very dilute 
there is produced merely a deep blue-green coloration. 
K6Fe2(CN)12 + 3FeS04 = FesFe2(CN)12 + 3K2S04. 
3. KCNS produces no change. 
4. H2S in acid solution does not form a precipitate. 
5. NH4HS, with a neutral or alkaline solution, forms a black 
precipitate of ferrous sulphide, soluble in HC1 or HN03. 
NH4CI promotes the formation of this precipitate. 
6. NH4OH in the absence of NH4C1 produces a dirty green 
precipitate of ferrous hydrate—Fe(OH)2. This precipitate 
rapidly becomes reddish-brown, owing to absorption of oxygen. 
7. KOH produces a dirty green precipitate of ferrous hy- 
drate—Fe(OH)2, similar to that produced by ammonia. Non- 
volatile organic substances, as sugar and some acids, retard 
the precipitations by NH4OH and KOH. 
8. Na2C03 causes a white precipitate of ferrous carbonate 
—FeC03, which rapidly becomes brown, from absorption of 
oxygen. This rapid oxidation is prevented by the use of dis 
tilled water and sugar. 
REACTIONS OF IRON in ferric state (Fe2)VI 
Use a solution of ferric chloride (Fe2Cl6). 
I. K4Fe(CN)6, in a neutral or add solution, produces a dark 
blue precipitate of ferric ferrocyanide—(Fe2)2(Fe(CN)6)3, de- 
composed by alkalies. 
3K4Fe(CN)6 + 2Fe2Cl6 = (Fe2)2(Fe(CN)6)3 +12 KC1. 48 
ANALYTICAL CHEMISTRY. 
2. KsFe2(CN)12 forms no precipitate, but produces a deep 
reddish-brown or olive-green color. The olive-green color is 
due to traces of ferrous salt. 
3. KCNS imparts to acid solutions a deep blood-red color, 
due to the formation of ferric sulphocyanate. This color is 
immediately destroyed by HgCl2. Dilute solutions show this 
reaction best. 
4. H2S forms in acid solutions a white turbidity due to 
separation of sulphur; the ferric salt being at the same time 
reduced to the ferrous condition. 
2Fe2Cl6 + 2H2S = 4FeCl2 -|- 4HCI -|- S2. 
In alkaline solutions this reagent acts as an alkaline sulphide. 
5. NH4HS causes a black precipitate of ferrous sulphide 
—FeS, sulphur separating at the same time. 
6. NH4OH precipitates reddish-brown ferric hydrate— 
Fe2(OH)6; non-volatile organic acids and sugar prevent this 
precipitation. 
7. KOH and NaOH react like NH4OH. 
8. Na2C03 and the other alkaline carbonates precipitate 
ferric hydrate—Fe2(OH)6. 
9. With borax in the oxidizing blowpipe flame, ferrous and 
ferric compounds give dark yellow colored beads while hot, 
and yellow when cold. In the reducing flame the beads change 
to bottle-green. 
REACTIONS OF CERIUM (Ce). 
Use a solution of cerous chloride (Ce2Cl6). 
1. NH4HS causes a white precipitate of cerous hydrate— 
Ce2(OH)6, H2S escaping. 
2. NH4OH produces the same white precipitate of Ce2(OH)6, 
as also do KOH and NaOH. 
3. (NH4)2C204 or H2C204 forms a white precipitate of cerous 
oxalate—Ce2(C204)3. Organic matter does not interfere with 
the formation of this precipitate. 
4. Na2HP04 precipitates white cerous phosphate— 
Ce2(P04)2. 
5. With a borax bead before the blowpipe the salts of cerium 
behave like those of iron. BASES, 
49 
REACTIONS OF ALUMINIUM (Al). 
Use a solution of alum (K2S04A12(S04)3). 
i. NH4HS produces a white precipitate of aluminium hy- 
drate—Al2(OH)6, while H2S escapes. 
6NH4HS + K2S04A12(S04)3 + 6Ha0 = Al2(OH)6 + 3(NH4)2S04 
+ K2S04 + 6H2S. 
2. NH4OH forms a white precipitate of aluminium hydrate 
—Al2(OH)6, insoluble in excess; this is an important distinc- 
tion from zinc. 
3. KOH and NaOH produce a similar white precipitate of 
Al2(OH)6, soluble in excess, forming aluminates of the base, 
as Al2(OK)6. This is not reprecipitated by boiling (distinction 
from chromium). 
4. Na2C03 and the other alkaline carbonates precipitate 
white gelatinous aluminium hydrate—Al2(OH)6, insoluble in 
excess, C02 escaping at the same time. 
3Na2COs + K2S04A12(S04)3 + 3H20 = Al2(OH)6 + 3Na2S04 
-)- K2S04 —(“ 3C02. 
The presence of non-volatile organic acids and sugar prevent 
the complete precipitation in the above reactions. 
REACTIONS OF CHROMIUM (Cr). 
Use a solution of chromic chloride (Cr2Cl6). 
1. NH4HS produces a greenish precipitate of chromic hy- 
drate—Cr2(OH)6, H2S escaping. 
2. NH4OH forms the same greenish precipitate of Cr2(OH)6, 
insoluble in excess. 
3. KOH and NaOH produce the same precipitate soluble 
in excess, but reprecipitated on boiling (distinction from alu- 
minium). 
4. Na2C03 and the other alkaline carbonates precipitate 
green basic carbonates. 
5. A mixture of KN03 1 part and K2C03 2 parts, fused with 
chromium compounds, become yellow from formation of potas- 
sium chromate—K2Cr04. AgN03 and Pb(C2H302)2 are im- 
portant tests for this compound; the former precipitates red 
silver chromate, the latter yellow lead chromate. 
6. With the borax bead in the inner blowpipe flame, chro- 
mium compounds give a green color. 50 
ANALYTICAL CHEMISTRY, 
SUMMARY OF TESTS WITH SOLUBLE SALTS OF GROUP IV. 
Fe(ous) 
Fe(ic) 
Ce 
A1 
Cr • 
K4Fe(CN)6 
White ppt. 
turning blue. 
Deep blue 
precipitate. 
K8Fea(CN)12 
Deep blue 
precipitate. 
Olive-green or 
brownish-red 
color. 
KCNS 
No change. 
Blood-red 
color. 
* 
NH4HS 
Black ppt. 
Black ppt. 
White ppt. 
White ppt. 
Greenish ppt. 
NHtOH 
Dirty green 
ppt. 
Reddish- 
brown ppt. 
White ppt. 
White ppt. in- 
soluble in excess. 
Greenish ppt. in- 
soluble in excess. 
KOH 
Dirty green 
ppt. 
Reddish- 
brown ppt. 
White ppt. 
White ppt. solu- 
ble in excess, not 
reprecipitated 
by boiling. 
Green ppt. 
soluble in excess, 
reprecipitated 
by boiling. 
Na3C03 
White ppt. be- 
coming dark. 
Reddish- 
brown ppt. 
White ppt. 
White ppt. 
Green ppt. 
DIRECTIONS FOR THE DETECTION OF THE BASES IN A SOLUTION 
CONTAINING SOLUBLE SALTS OF GROUP IV. 
Evaporate a portion of the solution to dryness, fuse on platinum foil with a mixture of 2 parts 
Na2C03 and 1 part KN03, boil with water and filter. 
Residue, Fe. Ce. 
Dissolve in concentrated H2S04 5 drops, and 
CsH6OH 20 drops, dilute with water, 
filter and divide in two parts. 
Filtrate, Al, Cr. 
Yellow if Cr be present, divide in two parts. 
Fe. 
Test with 
K4Fe(CN)8. 
Test original solution 
with K4Fe(CN)6 
and K6Fe2(CN)12 for 
Fe(ous) and Fe(ic). 
Ce. 
Add HX6H607, 
NH40H and 
Na2HP04, 
white ppt. 
Al. 
Add N H 4C1 and warm, 
white gelatinous 
ppt. 
Cr. 
Add HC2H302 and 
Pb(C2H302)2 
yellow ppt. 51 
BASES. 
Fe, Ce, Al, Cr, Mn, Zn, Co, Ni, Ba, Sr, Ca, Mg, K, Na, Li, NHi. 
Add NH4CI, N H4OH, NH4HS, warm gently, or boil if Ni is suspected, filter. 
Ppt. Fe, Ce, Al, Cr, Mn, Zn, Co, Ni. 
Wash, dissolve in HC1 and HNO3, evaporate excess of acid, add NH4OH in excess, filter. 
Filtrate Ba, Sr, Ca, Mg, K, Na, Li, NH4. 
Evaporate to remove NH4HS, add (NH4)2C03, boil, filter. 
Ppt. Fe, Ce, Al Cr. 
Wash, dry, fuse on foil with NaoC03 and KNO3, 
boil with HoO, filter. 
Filt. Mn, Zn, Co, Ni. 
Acidify with HC2H3O2, add H2S, filter. 
Ppt. Ba, Sr, Ca. 
Wash, dissolve in HC2H302, 
add K2Cr04, filter. 
Filt. Mg, K, Na, Li, NH4. 
Add (NH4)2HP04, 
agitate, filter. 
Ppt. Fe, Ce. 
Wash, dissolve in concentrated 
H2S04 and C2H5OH, dilute, filter 
and divide the filtrate in two 
portions. 
Fill. Al, Cr. 
Yellow if Cr be present, 
divide in two 
portions. 
Ppt. Zn, Co, Ni. 
Wash, dissolve in HC1 and 
HN03, evaporate excess of 
acid, add KOH until alkaline, 
filter. 
Filt. Mn. 
Add 
NH4OH 
and H2S, 
flesh-colored 
ppt. 
Ppt. 
Ba. 
Yellow. 
Filt. Sr, Ca. 
Add very dilute 
h2so4, 
allow to stand, filter. 
Ppt. 
Mg 
White 
Filt. 
K, Na, Li, NH4. 
Evaporate to dryness, 
ignite, dissolve in 
a small quantity of 
H02, add 
Na2HP04, boil, 
filter. 
Fe. 
Test with 
K4Fe(CN)6. 
Test original solu- 
tion for Fe(ous) 
and Fe(ic). 
Ce. 
Add H3C6H507, 
NH4OH and 
Na?HP04, 
white ppt. 
Al. 
Add 
NH4CI 
and 
warm, 
white 
gelati- 
nous ppt. 
Cr. 
Add 
HC2H302 
and 
Pb(C3H302)2 
yellow ppt. 
Ppt. Co, Ni. 
Wash, dissolve in 
HC1. Evaporate 
excess of acid, and 
treat with KCN 
and HC1 according 
to chart, page 45, 
add KOH until 
alkaline, filter. 
Filt. Zn. 
Add 
NH4HS, 
white ppt. 
Ppt. 
Sr. 
Filt. Ca. 
Add 
nh4oh 
and 
(NH )2C2°4. 
White ppt. 
Ppt. 
Li. 
Con- 
firm by 
flame 
test. 
Filt. 
K, Na, NH4. 
Concentrate, 
add HC1 and 
PtCl4, yellow 
ppt. == K. 
Test original 
solution for 
Na and NH4. 
Green. 
Con- 
firm 
borax 
bead. 
Filt. Co. 
Evapo- 
rate to 
dryness 
and test 
with borax 
bead. 
DIRECTIONS FOR THE DETECTION OF THE BASES IN A SOLUTION CONTAINING SOLUBLE SALTS OF ALL THE ELEMENTS 
HERETOFORE CONSIDERED. 52 
ANALYTICAL CHEMISTRY. 
PRECAUTIONS AND OBSERVATIONS ON THE PRECEDING CHART. 
1. It is particularly desirable not to filter immediately after 
adding NH4HS in the first precipitation; otherwise, the Mn 
will not be thoroughly precipitated. 
2. When the first precipitate is dissolved in HC1 and HN03, 
great care must be exercised to thoroughly oxidize the Fe by 
boiling with the HN03, otherwise, the precipitation by NH4OH 
will not be complete. 
3. The precipitate of aluminium hydrate is difficult to 
observe, as it floats in the solution instead of falling to 
the bottom. Warming the solution will usually render it 
visible. 
4. When phosphoric acid is present, the members of Group 
II precipitate with Group IV, and require an entirely different 
method of separation. It is desirable, however, for the student 
first to familiarize himself with these simpler soluble salts, and 
undertake the more difficult cases of salts insoluble in water 
after the acids have been considered. See page 88. 
GROUP V.—ARSENIC, ANTIMONY, TIN, GOLD, 
PLATINUM. 
REACTIONS OF ARSENIC (As). 
(a) ARSENIOUS COMPOUNDS. 
Use a solution of As203, in water. 
i. H2S passed into the solution produces a yellow color, 
but no precipitate until HC1 is added, when a yellow precipi- 
tate of arsenious sulphide—As2S3, falls. This precipitate is 
insoluble in strong HC1, but soluble in NH4HS, NH4OH and 
(nh4)2co3 
As20„ + 3H2S = As2S3 + 3H20. 
2. NH4HS causes the formation of arsenious sulphide, 
which remains in solution as ammonium sulpharsenite— 
(NH4)3AsS3. On the addition of HC1 arsenious sulphide is 
precipitated. 
2(NH4)3AsS3 + 6HC1 = 6NH4C1 + 3H2S + AsaS3. 
3. AgN03 produces no precipitate until a few drops of 
dilute ammonia solution are added, when a yellow precipitate BASES 
53 
of silver arsenite—Ag3As03, falls, soluble in HN03 and in 
NH4OH. 
4. CuS04, under similar circumstances, produces a yellow- 
ish-green precipitate of cupric arsenite—CuHAs03. 
(b) ARSENIC COMPOUNDS. 
Use a solution of sodium arsenate (Na2HAsO(). 
1. H2S causes, in acid solution only, a yellow precipitate of 
arsenious sulphide—As2S3, mixed with sulphur. This re- 
action takes place slowly, but is accelerated by heat. 
5H2S + 2Na2HAs04 + 4HCI = As2S3 + S2 + 4NaCl + 8H20. 
2. NH4HS produces no precipitate, but forms arsenic 
sulphide—As2S5, which remains in solution as ammonium 
sulpharsenate—(NH4)3AsS4. Upon the addition of HCl 
As2S5 is precipitated, and not As2S3 and S. 
2(NH4)3AsS4 + 6HC1 = 6NH4C1 + 3H2S + As2S5. 
3. AgN03 with a small amount NH4OH, produces a choco- 
late-colored precipitate of silver arsenate—Ag3As04, soluble 
in HN03 and NH4OH. 
4. CuS04, under similar circumstances, forms a bluish-green 
precipitate of cupric arsenate—CuHAs04. 
The following tests are applicable to both arsenious and 
arsenic compounds. 
i. Marsh's Test.—Generate hydrogen in the usual way, 
allowing it to escape through a glass tube drawn out at the 
end so as to form a small orifice (Fig. 4). In very exact cases, 
the gas should be dried by passing over calcium chloride. 
When all the air has been expelled (which should be deter- 
mined by collecting a small test tube full and holding its 
mouth to a flame; if the gas burn quietly, without explosion, 
it is pure), ignite the escaping gas; it should burn with a 
colorless or yellow flame; in the latter case it is due to the 
sodium in the glass. A piece of cold porcelain—a small cru- 
cible lid is best—is pressed down on the flame ; there should 
be no deposit on it. Add now, through the funnel tube, a 
solution of arsenic, washing it down with a little water. The 
flame will become of a pale blue color, due to the formation 
of hydrogen arsenide—H3As. On bringing the crucible lid 54 
ANALYTICAL CHEMISTRY. 
into the flame now, a blackish-brown deposit, with metallic 
lustre, will form on it. This deposit is readily soluble in a 
solution of sodium or calcium hypochlorite. 
2. Reinsch's Test.—Boil some strips of copper with dilute 
HC1; if no discoloration of the copper takes place, the 
arsenic solution may be added. The copper immediately 
becomes coated with an iron-gray metallic film. Pour off 
the liquid, dry the copper by holding it in the lamp with 
the fingers so it may not become too hot, place in a clean, 
Fig. 4. 
dry, narrow test tube, and heat gently, when a white ring 
of As203 will form on the tube above the copper, readily 
distinguished by the characteristic octahedral shape of the 
crystals. 
3. Fleitman's Test.—Generate hydrogen in a test tube with 
zinc and solution of potassium hydrate; moisten a piece of 
filter paper with one drop of solution of silver nitrate, place 
it over the mouth of the tube and heat; there should be no 
coloration of the spot on the paper. Now add some corn- BASES, 
55 
pound of arsenic; the silver nitrate'will immediately become 
black, owing to production of metallic silver. 
H3As + (AgNOs)6 + (HaO)s = H3As03 + (HN03)6 + (Ag2)3. 
Before the blowpipe, on charcoal, arsenic volatilizes, with 
the characteristic odor of garlic. 
REACTIONS OF ANTIMONY (Sb). 
Use a solution of tarter emetic (KSbOC4H4Ofi) acidified with HC1. 
1. H2S forms an orange precipitate of antimonous sul- 
phide—Sb2S3, soluble in NH4HS, and in concentrated HC1, 
but insoluble in (NH4)2C03. 
2. NH4HS produces an orange precipitate of antimonous 
sulphide, readily soluble in excess, forming ammonium 
sulph-antimonite—(NH4)3SbS3, from which HC1 again pre- 
cipitates Sb2S3. 
3. KOH or NaOH precipitates white, bulky antimonous 
hydrate—Sb(OH)3, soluble in excess. 
4. NH4OH precipitates the same compound insoluble in 
excess. 
4. Marsh's Test gives the same result as with arsenic; the 
black spot, however, is insoluble in sodium or calcium hypo- 
chlorite solution, but soluble in NH4HS. 
6. Reinsch's Test causes a deposit on copper, as with arsenic, 
but when heated in a tube there is formed a white amorphous 
ring, which is readily distinguished from the crystalline one 
of arsenic. 
7. Fleitman’s Test gives no result with antimony com- 
pounds. 
8. On charcoal, with Na2C03, before the blowpipe, a metallic 
globule of antimony is produced, while characteristic fumes 
of the oxide are given off. 
REACTIONS OF TIN (Sn). 
(a) STANNOUS COMPOUNDS. 
Use a solution of stannous chloride (SnCl2). 
I. H2S precipitates dark brown stannous sulphide—SnS, 
soluble in concentrated HC1 and in (NH4)2S, insoluble in 
NH4HS. 56 
ANALYTICAL CHEMISTRY. 
2. KOH or NaOH precipitates white stannous hydrate— 
Sn(OH)2, soluble in excess ; on boiling this solution SnO pre- 
cipitates. 
3. NH4OH precipitates the same compound, insoluble in 
excess. 
4. HgCl2 causes a white precipitate of Hg2Cl2, converting 
the SnCl2 into SnCl4. 
This precipitate blackens on the addition of NH4OH. 
2HgCl2 -)- SnCl2 = Hg2Cl2 + SnCl4. 
(b) STANNIC COMPOUNDS. 
Use a solution of stannic chloride (SnCl4). 
1. H2S produces a yellow precipitate of stannic sulphide 
—SnS2, soluble in NH4HS and in concentrated HC1. 
2. KOH or NaOH precipitates white stannic acid—H2Sn03, 
soluble in excess ; on boiling no reprecipitation takes place— 
distinction from stannous salts. 
3. NH4OH produces the same precipitate, insoluble in excess. 
4. Heated on charcoal, before the blowpipe, with Na2C03, 
metallic tin is formed, with the production of a white incrust- 
ation of the oxide. 
REACTIONS OF GOLD (Au). 
Use a solution of auric chloride (AuC13). 
1. H2S precipitates black auric sulphide—Au2S3, insoluble 
in HC1, soluble in (NH4)2S. 
2. H2C204 or FeS04 precipitates metallic gold as a finely 
divided brown powder. 
3. SnCl2 mixed with SnCl4 (prepared by SnCl2 and chlo- 
rine water) produces a purple-red precipitate or coloration 
(Purple of Cassius), consisting of the mixed oxides of gold 
and tin. 
4. Heated on charcoal, before the blowpipe, metallic gold 
is produced. 
REACTIONS OF PLATINUM (Pt). 
Use a solution of platinic chloride (PtCl4). 
i. H2S causes a brown precipitate of platinic sulphide 
PtS2, insoluble in HC1, soluble in (NH4)2S. BASES. 
57 
2. KC1 produces a yellow crystalline precipitate of potas- 
sium platinic chloride—K2PtCl6. 
3. NH4CI forms a similar precipitate of (NH4)2PtCl6. The 
precipitation in both cases is facilitated by the addition of 
alcohol. 
4. Zn, Fe and some other metals precipitate metallic 
platinum. 
5. Heating on charcoal, before the blowpipe, produces the 
metal. 
SUMMARY OF TESTS WITH SOLUBLE SALTS OF GROUP V. 
As 
Sb 
Sn 
Au 
Pt 
H2S 
Yellow 
ppt. 
soluble in 
NH4HS. 
Orange-red 
ppt. 
soluble in 
NH4HS. 
Brown or Yellow 
soluble in 
(NH4)2S. 
Black 
ppt. 
soluble in 
(NH4)2S. 
Brown 
ppt. 
soluble in 
(NH4I2S. 
KOH 
No change. 
White 
ppt. 
soluble in 
excess. 
White 
ppt. 
soluble in 
excess. 
No change. 
With excess 
of HC1 
yellow 
ppt. 
NH4OH 
No change. 
White 
ppt. 
insoluble in 
excess. 
White 
PP* 
insoluble in 
excess. 
Red ppt. 
fulminating 
gold. 
With excess 
of HC1 
yellow 
ppt. 
Marsh’s Test. 
Black spot 
with metallic 
lustre, soluble 
in Ca(OCl)2. 
Black, sooty 
spot, soluble 
in NH4HS, 
insoluble in 
Ca(OCl)2. 
Reinsch’s Test. 
White 
sublimate of 
octahedral 
crystals. 
White 
amorphous 
sublimate. 
Fleitman’s Test. 
Dark spot 
of metallic 
silver. 
No change. 
DIRECTIONS FOR THE DETECTION OF THE BASES IN A SOLUTION 
CONTAINING SOLUBLE SALTS OF GROUP V. 
Add HC1 and H2S, collect the precipitate, transfer to a dish, boil with 
concentrated HC1; filter. 
Ppt. As, Au, Pt. 
If yellow, As only is present; if dark, wash, 
digest with (NH4)2C03, filter. 
Filt. Sb, Sn. 
Dilute with H20, boil in a dish with a strip 
of platinum foil, and a small piece of zinc 
so the metals touch. The platinum will 
be coated with black Sb, while the Sn 
will be deposited as a black sediment. 
This black sediment is dissolved in strong 
HC1 and tested with HgCl2, white ppt. 
if Sn be present. Treat the foil with a 
few drops of HNOa, dissolve in solution 
H2C4H40„, add H2S, orange ppt. if Sb 
be present. 
Ppt. Au, Pt. 
Dissolve in HC1 and HNO3. 
Divide in two parts. 
Fill. As. 
Acidify 
with HC1, 
yellow 
ppt. 
Au. 
Add SnCl2 purple. 
Confirm by 
testing original 
solution. 
Pt. 
Add KC1, yellow 
ppt. Confirm by 
testing original 
solution. 58 
ANALYTICAL CHEMISTRY. 
As, Sb, Sn, Au, Pt, Fe, Ce, Al, Cr, Mn, Zn, Co, Ni, Ba, Sr, Ca, Mg, K, Na, Li, NH4. 
Add HC1 and H2S until, after warming, the liquid smells strongly, filter. 
Ppt. As, Sb, Sn, 
Au, Pt. 
Wash, boil with con- 
centrated HC1, 
filter. 
Filt. Fe, Ce, Al, Cr, Mn, Zn, Co, Ni, Ba, Sr, Ca, Mg, K, Na, Li, NH4. 
Add NH4C1, NH4OH until after agitation its odor is perceptible, and NH4HS, warm, filter. 
Ppt. Fe, Ce, Al, Cr, Mn, Zn, Co, Ni. 
Wash, dissolve in HC1 and HN03, evaporate excess of acid, 
add NH4OH in excess, filter. 
Filt. Ba, Sr, Ca, Mg, K, Na, Li, NH4. 
Evaporate to remove NH4HS, add (NH4)2 C03, boil, filter. 
Ppt. As, 
Au, Pt. 
Separate 
by the 
preceding 
Chart 
for 
Group V. 
Fill. 
Sb, Sn. 
Separate 
by the 
preceding 
Chart 
for 
Group V. 
Ppt. Fe, Ce, Al, Cr. 
Wash, dry, fuse on foil with 
Na2CG3 and KNOa, boil with 
H20, filter. 
Filt. Mn, Zn, Co, Ni. 
Acidify with HC2H302, add H2S, filter. 
Ppt. Ba, Sr, Ca. 
Wash, dissolve in 
HC2H302, add K2Cr04, 
filter. 
Filt. Mg, K, Na, Li, NH4. 
Add (NH4)2 HP04, agitate, 
filter. 
Ppt. Zn, Co, Ni 
Wash, dissolve in HC1 and HN03, 
evaporate excess of acid, add KOH 
until distinctly alkaline, filter. 
Filt.Mn. 
Add 
NH,OH 
and 
H2S. 
Flesh- 
colored 
ppt. 
Ppt. 
Mg. 
White. 
Filt. K, Na, Li, NH4. 
Evaporate to dryness, 
ignite, dissolve in a 
small quantity of 
H20, add Na2HP04, 
boil, filter. 
Ppt. 
Fe, Ce. 
Separate 
according 
to Chart, 
page 50. 
Filt. Al, Cr. 
Yellow if Cr be 
present, divide in two 
parts. 
Ppt. 
Ba. 
Yellow. 
Filt. Sr, Ca. 
Add very dilute 
H2S04, allow to 
stand, filter. 
Ppt. Co, Ni. 
Wash, dissolve in HC1, 
evaporate excess of acid, treat 
with KCN and HC1 according 
to Chart, page 45, add 
KOH until alkaline, filter. 
Filt. Zn. 
Add 
NH4HS. 
White 
ppt. 
Al. 
Add 
NH4C1. 
White 
ppt. 
Cr. 
Add 
hc2h3o2 
and 
Pb(A)2. 
Yellow 
ppt. 
Ppt. Sr. 
Filt. Ca. 
Add 
NH4OH 
and 
('!&.*) 
White 
ppt. 
Ppt. Li. 
Con- 
firm by 
flame 
test. 
Filt. 
K, Na, NH4. 
Concentrate, 
add HC1 and 
PtCl4. 
Yellow 
ppt. = K. 
Test original 
solution for 
Na and NH4. 
Ppt. Ni. 
Green. 
Confirm 
by 
borax 
bead. 
Filt. Co. 
Evaporate to dryness 
and test with borax 
bead. 
DIRECTIONS FOR THE DETECTION OF THE BASES IN A SOLUTION CONTAINING SOLUBLE SALTS 
OF ALL THE ELEMENTS HERETOFORE CONSIDERED. BASES, 
59 
PRECAUTIONS AND OBSERVATIONS ON THE PRECEDING CHART. 
A great deal of time may be saved by carefully noting the 
color of the precipitate produced by H2S. If stannic salts 
are absent and the precipitate yellow, As only is present; if 
orange, Sb is present and As may be ; when such an orange 
precipitate is obtained and Sn(ic) is absent, the readiest 
method of separation is to wash, and add to the precipitate 
(NH4)2C03. As will be dissolved, and may be detected in 
the filtrate by adding HC1, while the Sb remains on the filter. 
When the precipitate is dark, Au and Pt should be sought 
for in the original solution, as well as separated by the chart. 
Only a small piece of zinc is necessary to effect the separa- 
tion of Sn and Sb. 
GROUP VI.—MERCURY (IC), BISMUTH, COPPER, 
CADMIUM. 
REACTIONS OF MERCURY as mercuric salt (Hg(ic)). 
Use a solution of mercuric chloride (HgCl2). 
1. H2S or NH4HS produces, when in small proportion, a 
whitish precipitate of (HgS)2HgCl2; a further addition of the 
reagent, together with the application of heat, causes the 
formation of a black precipitate of mercuric sulphide—HgS, 
insoluble in either HC1 or HN03, but soluble in a mixture of 
the two. 
2. KOH or NaOH produces a yellow precipitate of mer- 
curic oxide—HgO ; unless the reagent be in excess, a brown 
precipitate of a basic salt is formed. 
2K0H -f HgCl2 = HgO + 2KC1 + H20. 
3. NH4OH precipitates white mercur-ammonium chloride 
—NH2HgCl. 
This precipitate is readily soluble in HC1 and in HC2H302. 
4. K2Cr04 produces a red precipitate of mercuric chro- 
mate—HgCr04. 
5. KI precipitates mercuric iodide—Hgl2, first yellow, but 
rapidly becoming scarlet. This precipitate is readily soluble 
in excess of KI or HgCl2. 
6. SnCl2, in small quantity, in the presence of HC1, pre- 
2NH4OH + HgCl2 = NH2HgCl + NH4C1 + 2H20. 60 
ANALYTICAL CHEMISTRY. 
cipitates mercurous chloride—Hg2Cl2. On the addition of 
a larger quantity of the reagent, the mercurous chloride is 
reduced to the metal, which may be collected into a globule. 
7. Na2C03 produces a reddish-brown precipitate of basic 
carbonate—HgC03(Hg0)3. 
8. Copper turnings boiled with the solution acidified with 
HC1 become quickly coated with metallic mercury. The 
presence of HN03 must be avoided. If these coated turnings 
are carefully dried at a low temperature, placed in the 
bottom of a narrow test tube and heated, a ring composed of 
globules of mercury will form on the cool part of the tube. 
These minute globules may be formed into larger ones by 
pressing them together with a glass rod, until they may be 
seen with the naked eye. 
9. Before the blowpipe, HgO breaks up into Hg and O. 
HgS, under similar circumstances, sublimes unchanged. 
REACTIONS OF BISMUTH (Bi). 
Use a solution of bismuth nitrate (Bi(N03)3). 
1. H2S or NH4HS produces a black precipitate of bismuth 
trisulphide—Bi2S3, insoluble in dilute acids and alkalies, 
soluble in boiling HN03. 
2. KOH, NaOH or NH4OH forms a white precipitate of 
bismuth hydrate—Bi(OH)3, converted by boiling into the 
yellow oxide—Bi203. 
3. K2Cr04 precipitates yellow bismuth chromat e— 
Bi2(Cr04)3. 
4. KI forms a brown precipitate of bismuth iodide—Bil3, 
soluble in excess of the reagent. 
5. Na2C03 precipitates white oxycarbonate of bismuth— 
(Bi0)2C03.H20. 
+ 2Bi(N03)3 + H20 = (Bi0)2C03.H20 -f 6NaN03 + 2C02. 
6. H20 in excess, when there is not an excess of free acid, 
precipitates bismuth subnitrate—Bi0N03.H20. 
Bi(N03)3 + 2HaO = Bi0N03.H20 + 2HN03. 
When the chloride is so diluted the oxychloride—BiOCl, 
separates. BASES. 
61 
7. Bismuth on charcoal, before the blowpipe, forms a hard 
bead of metal, with a characteristic incrustation of oxide— 
deep orange-yellow while hot, pale when cold. 
REACTIONS OF COPPER (Cu). 
1. H2S and NH4HS precipitate black cupric sulphide— 
CuS, insoluble in dilute acids and alkalies, slightly soluble in 
NH4HS, and entirely dissolved by boiling HN03. This pre- 
cipitation is prevented by KCN. 
Use a solution of cupric sulphate (CuS04). 
2. KOH or NaOH produces a light blue precipitate of 
cupric hydrate—Cu(OH)2, insoluble in excess, and converted 
by boiling into black cupric oxyhydrate—(CuO)2Cu(OH)2. 
In the presence of non-volatile organic acids this precipita- 
tion does not take place, but a blue color results. 
2K0H + CuS04 = Cu(OH)2 + K2S04. 
3Cu(OH)2 = (CuO)2Cu(OH)2 + 2il20. 
3. NH4OH in small quantity forms a greenish-purple precipi- 
tate, readily soluble in excess, forming tetra-ammonio-cupric 
sulphate—(NH3)4CuS04.H20. 
4NH4OH + CuS04 = (NH3)4CuS04.H20 -f 3H20. 
4- K4Fe(CN)6 precipitates reddish-brown cupric ferro- 
cyanide—Cu2Fe(CN)6. 
5. Metallic Fe or Zn precipitates red metallic Cu. 
6. In the outer blowpipe flame copper salts color the borax 
bead green while hot, blue when cold. In the inner flame, 
after moistening with SnCl2, it becomes red, owing to formation 
of Cu20. 
REACTIONS OF CADMIUM (Cd). 
Use a solution of cadmium sulphate (CdS04). 
I. H2S or NH4HS precipitates yellow cadmium sulphide 
—CdS, soluble in hot HNO;j and dilute H2S04, but insoluble 
in NH4HS or KCN. 
2. KOH or NaOH produces a white precipitate of cadmium 
hydrate—Cd(OH)2, insoluble in excess. 
3. NH4OH causes the same white precipitate of Cd(OH)2, 
soluble in excess. 62 
ANALYTICAL CHEMISTRY. 
4. Na2C03 produces a white precipitate of cadmium car- 
bonate—CdC03, insohible in excess, but slightly soluble in 
ammonium salts, entirely soluble in NH4OH. 
5. On charcoal, before the blowpipe, the salts of cadmium 
are reduced to metal and volatilize, forming a brownish in- 
crustation of oxide. 
SUMMARY OF TESTS WITH SOLUBLE SALTS OF GROUP VI. 
Hg(ic) 
Bi 
Cu 
Cd 
HaS or NH4HS 
Black ppt. 
insoluble in HNOs. 
Black ppt. 
Black ppt. 
soluble in KCN. 
Yellow ppt. 
insoluble in KCN. 
KOH 
Yellow ppt. 
White ppt. 
Blue ppt. 
White ppt. 
NH.OH 
White ppt. 
White ppt. 
Blue ppt. 
soluble in excess. 
White ppt. 
soluble in excess. 
Na2C03 
Reddish-brown ppt. 
White ppt. 
Blue ppt. 
White ppt. 
DIRECTIONS FOR THE DETECTION OF THE BASES IN A SOLUTION 
CONTAINING SOLUBLE SALTS OF GROUP VI. 
Add H2S, collect the precipitate, boil it with HNOa, (sp. gr. 1.20); filter. 
Ppt. Hg. 
Black. 
Confirm by Cu, accord- 
ing to 8, page 60. 
Fill. Bi, Cu, Cd. 
Add NH4OH in excess; filter. 
Ppt. Bi.* 
White. 
Fill. Cu, Cd. 
Add H2S, collect and wash ; boil with dilute 
H2S04 (one part of concentrated acid 
and five parts water); filter. 
Ppt. Cu.” 
Black. 
Filt. Cd. 
Dilute, add H2S, yellow ppt. 
* To confirm the presence of Bi, wash this precipitate, pour on it a few drops of hydrochloric 
acid and allow the solution thus formed to run into a test tube of water, when if the precipitate 
was Bi, a white precipitate of bismuth oxychloride will form. 63 
BASES. 
Hg(ic) 
Bi, Cu, Cd, As, Sb, Sn, Au, Pt, Fe, Ce, Al, Cr, Mn, Zn, Co, Ni, Ba, Sr, Ca 
Add HC1 and HaS, filter. 
» Mg 
K, Na, Li, NH4. 
Ppt. Hg, Bi, Cu, Cd, As, Sb, Sn, Au, Pt. 
Wash, digest with (NHJaS, filter. 
Filt. Fe, 
Ce, Al, Cr, Mn, Zn, Co, Ni, Ba, Sr, Ca, Mg, K, Na, Li, NH 
Add NH4C1, NH4OH, NH4HS, warm, filter. 
4* 
Ppt. Hg, Bi, Cu, Cd. 
Wash, boil with HN03, 
filter. 
Filt. 
As, Sb, Sn, 
Au, Pt. 
Acidify with 
Ppt. Fe, Ce, Al, Cr, Mn, Zn, Co, Ni. 
Wash, dissolve in HC1 and HN03, evaporate excess of acid, 
add NH4OH in excess, filter. 
Filt. Ba, Sr, Ca, Mg, K, Na, Li, NH4. 
Evaporate to remove NH4HS,add (NH4)2C03, 
boil, filter. 
Ppt. 
Hg. 
Black. 
Filt. Bi, Cu, Cd. 
Evaporate excess of 
acid, add NH4OH in 
dilute HC1. 
Collect and 
wash the pre- 
cipitate (throw 
away the 
filtrate), boil 
the ppt. with 
concentrated 
HC1, filter. 
Ppt. Fe, Ce, Al, Cr. 
Wash, dry, fuse on foil 
with Na2C03 and 
KN03, boil with H2Q, 
Filt. Mn, Zn, Co, Ni. 
Acidify with HC2H302, add H2S, filter. 
Ppt. Ba, Sr, Ca. 
Wash, dissolve in 
HC2H302> add K,Cr04, 
Filt. Mg, K, Na, Li, NH, 
Add (NH4)2HP04, 
agitate, filter. 
Con- 
excess, filter. 
niter. 
firm 
niter. 
Ppt. Zn, Co. Ni. 
Filt. Mn. 
byCu. 
Wash, dissolve in HC1 and 
Add 
Ppt. 
Filt. K,Na,Li,NH4. 
Ppt. 
Bi. 
White 
Filt. Cu, Cd. 
Add HjS. 
Ppt. 
Fe, 
Filt. Al, Cr. 
HNO3, 
evaporate excess of acid, 
add KOH, filter. 
NH40H 
and 
H2S. 
Ppt. 
Ba. 
Yel- 
Filt. Sr, Ca. 
Add very dilute 
H2S04, allow to 
stand, filter. 
Mg. 
Evaporate to dryness, 
ignite, dissolve in 
Con- 
firm 
wash, boil with 
dilute H2S04, 
Ppt. 
As,Au, 
Pt. 
Sepa- 
rate 
chart 
page 
57- 
Filt. 
Sb, Sn. 
Sepa- 
rate 
,by 
chart 
page 
57- 
Ce. 
Sepa- 
present. Uivide in 
two parts. 
Pp 
t. Co, Ni. 
Filt. Zn. 
Flesh- 
colored 
low. 
INa2nru4, boil and 
filter. 
as on 
filter. 
rate 
Wasl 
1, dissolve in 
Add 
ppt. 
page 
accor- 
ding 
to 
chart 
page 
5°. 
HC1, evaporate excess 
NH4HS. 
Ppt. 
Filt. Ca. 
62. 
Ppt. 
Cu. 
Black. 
Filt. Cd. 
Dilute, 
add 
H„S, 
Al. 
Add 
NH4C1 
White 
ppt. 
Cr. 
Add 
HC2H302 
and 
Pb(A)2. 
of acid, treat with 
KCN and HC1,accord- 
ing to chart page 45, 
add KOH, filter. 
White 
ppt. 
Sr. 
Add NH4OH 
and 
(NH4)2C204, 
white ppt. 
Ppt. 
Li. 
Con- 
firm 
by 
flame 
test. 
Filt. 
K, Na, NH4. 
Concentrate, 
add HC1 and 
PtCl4, yellow 
ppt. = K. 
Test for Na 
and NH4 
in original 
solution. 
yellow 
ppt. 
Yellow 
ppt. 
PNPh- 
Green. 
Con- 
firm by 
borax 
bead. 
Filt. Co. 
Evaporate to 
dryness and 
test with borax 
bead. 
DIRECTIONS FOR THE DETECTION OF THE BASES IN A SOLUTION CONTAINING SOLUBLE SALTS 
OF ALL THE ELEMENTS HERETOFORE CONSIDERED. 64 
ANALYTICAL CHEMISTRY. 
PRECAUTIONS AND OBSERVATIONS ON TIIE PRECEDING CHART. 
1. Unless excess of H2S be used, and the solution warmed, 
Hg will not be thoroughly precipitated. 
2. When Sn, Au and Pt are present, the yellow ammonium 
sulphide (NH4)2S must be used to dissolve Group V. When 
Cu is present some of it may be dissolved, so that any dark 
precipitate in that group should be tested for Cu. 
3. When the precipitate insoluble in NH4HS is boiled with 
HNO;{, a black mass of sulphur is sometimes obtained. 
This is readily distinguished from Hg, because it floats as 
one mass on the liquid, while the mercury collects at the 
bottom as a heavy, black precipitate. 
GROUP VII.—SILVER, MERCURY(OUS), LEAD. 
REACTIONS OF SILVER (Ag). 
Use a solution of silver nitrate (AgN03). 
I. HC1 or soluble chlorides precipitate white, curdy silver 
chloride—AgCl, insoluble in HN03, soluble in NH4OH, form- 
ing ammonio-silver chloride—(AgCl)2(NH3)3, from which 
the chloride is again precipitated by acids. 
HC1 + AgNOj = AgCl + HNOs. 
aAgCl + 3NH4OH = (AgCl)2(NH3)3 + 3H20. 
2. H2S or NH4HS produces a black precipitate of silver 
sulphide—Ag2S, insoluble in dilute acids and in alkalies, 
soluble in boiling HN03. 
3. KOH or NaOH forms a grayish-brown precipitate of 
silver oxide—Ag20, insoluble in excess, but soluble in 
NH4OH. 
4. NH4OH in small quantity, precipitates silver oxide, 
soluble in excess. 
2K0H + 2AgN03 = Ag20 + 2KNO3 + H20. 
5. K2Cr04 produces a red precipitate of silver chromate— 
Ag2Cr04, soluble in concentrated HN03 and in NH4OH. 
6. KI and KBr produce precipitates of silver iodide—Agl, 
yellow, insoluble in NH4OH, and silver bromide—AgBr, 
yellowish-white, slowly soluble in NH4OH. 
7. KCN precipitates white silver cyanide—AgCN, soluble 
in excess and in concentrated HN03. BASES, 
65 
8. Heated with Na2C03 on charcoal, before the blowpipe, 
compounds of silver form a bright, metallic button, soluble in 
HN03. 
REACTIONS OF MERCURY as mercurous salt (Hg(ous)). 
Use a solution of mercurous nitrate (Hg2(N03)2). 
1. HC1 or soluble chlorides precipitate mercurous chloride 
—Hg2Cl2, converted by strong HN03 into a mixture of 
HgCl2 and Hg(N03)2, also becoming black on the addition 
of NH4OH, forming NH2Hg2Cl. 
2. H2S or NH4HS precipitates a mixture of Hg with HgS. 
3. KOH or NaOH produces a black precipitate of mer- 
curous oxide—Hg20, insoluble in excess. 
4. NH4OH causes a black precipitate of mercurous- 
ammonium nitrate—NH2Hg2N03. 
5. K2Cr04 forms an orange precipitate of mercurous 
chromate—Hg2Cr04. 
6. KI precipitates green mercurous iodide—Hg2I2. 
7. Before the blowpipe, mercurous salts volatilize, some 
being converted into mercuric salt and mercury, both of 
which sublime. 
2NH4OH + Hg2(N03)2 = NH2Hg2N03 + NH4N03 + 2H20. 
REACTIONS OF LEAD (Pb). 
Use a solution of lead acetate (Pb(C2H302)2). 
1. HC1, or soluble chlorides, produce a white precipitate of 
lead chloride—PbCl2, soluble in hot water. 
2. H2S or NH4HS precipitates black lead sulphide—PbS, 
insoluble in HC1, soluble in hot HN03. 
3. KOH or NaOH produces a white precipitate of lead 
hydrate—Pb(OH)2, soluble in large excess, forming potas- 
sium or sodium plumbate—K2Pb02 or Na2Pb02. 
4. NH4OH precipitates white basic lead hydrate. 
5. K2Cr04 produces a yellow precipitate of lead chromate 
—PbCr04, soluble in KOH and in strong HN03. 
6. KI forms a yellow precipitate of lead iodide—Pbl2, 
soluble in boiling water. 
7. H2S04 produces a white precipitate of lead sulphate— 
5 66 
ANALYTICAL CHEMISTRY. 
PbS04, insoluble in acids, but soluble in solution of ammo- 
nium acetate or tartrate. 
8. Na2C03 precipitates white basic lead carbonate— 
(PbC03)2Pb(0H)2. 
3^C03 + 3Pb(C2H302)2 + H20 = (PbC03)2Pb(0H)2 
-)- 6NaC2H302 -f- C02. 
9. Before the blowpipe, on charcoal, lead compounds are 
converted into a malleable globule of the metal, with the 
formation of some yellow oxide. 
SUMMARY OF TESTS WITH SOLUBLE SALTS OF GROUP VII. 
Ag 
Hg(ous) 
Pb 
HC1 
White ppt 
soluble in NH4OH. 
White ppt. turning 
black with NH4OH. 
White ppt. 
soluble in hot HaO. 
H2S and NH4HS 
Black ppt. 
Black ppt. 
Black ppt. 
KOH 
Brown ppt. 
Black ppt. 
White ppt. 
NH4OH 
Brown ppt. 
Black ppt. 
White ppt. 
Na2C03 
Brown ppt. 
Black ppt. 
White ppt. 
K2Cr04 
Red ppt. 
Orange ppt. 
Yellow ppt. 
KI 
Yellow ppt. 
Green ppt. 
Yellow ppt. 
DIRECTIONS FOR THE DETECTION OF THE BASES IN A SOLUTION 
CONTAINING SOLUBLE SALTS OF GROUP VII. 
Add HC1, collect, wash and pour on the filter boiling H20. 
Ppt. Ag, Hg(ous). 
Pour on the filter NH4OH. 
Filt. Pb. 
Add dilute H2S04, white ppt. 
Ppt. Hg(ous). 
Black. 
Filt. Ag. 
Acidify with HNOa, 
white ppt. Ag, Hg(ous), Pb, Hg(ic), Bi, Cu, Cd, As, Sb, Sn, Au, Pt, Fe, Ce, Al, Cr, Mn, Zn, Co, Ni, Ba, Sr, Ca, Mg, K, Na, Li, NH-. 
Add HC1, filter. 
Ppt. Ag, Hg(ous), Pb. 
Wash, pour boilingHaO 
on the filter. 
Filt. Hg(ic), Bi, Pb, Cu, Cd, As Sb, Sn, Au, Pt, Fe, Ce, Al, Cr, Mn, Zn, Co, Ni, Ba, Sr, Ca, Mg, K, Na, Li, NH4. 
Add H2S until, after warming and agitating, it smells strongly, filter. 
Ppt. Hg, Bi,Pb, Cu, Cd, As, Sb, Sn, Au, Pt. 
Wash, digest with (NH4)2S, filter. 
Filt. Fe, Ce, Al, Cr, Mn, Zn, Co, Ni, Ba, Sr, Ca, Mg, K, Na, Li, NH4. 
Add NH4C1, NH4OH, NH4HS, warm, filter. 
Ppt. 
Ag, Hg(ous). 
Pour on the filter 
NH4OH. 
Filt. Pb. 
Add 
H2S04, 
white 
ppt. 
Ppt. Hg. Bi, Pb, Cu, Cd. 
Wash, boil with HN03, filter. 
Filt. 
As, Sb, Sn, 
Au, Pt. 
Acidify with 
dilute HC1. 
Collect and 
wash the pre- 
cipitate (throw 
away the 
filtrate), boil 
the ppt. with 
concentrated 
HC1, filter. 
Ppt. Fe, Ce, Al, Cr, Mn, Zn, Co, Ni. 
Wash, dissolve in HC1 and HN03, evaporate excess of acid, add 
NH40H in excess, filter. 
Filt. Ba, Sr. Ca, Mg, K, Na, Li, NH4. Evaporate 
to remove NH4HS,add (NH4)2C03, boil, filter. 
Ppt. 
Hg(ous). 
Black. 
Fill. 
Ag. 
Add 
HN03, 
white 
ppt. 
Ppt. 
Hg. 
Black. 
Filt. Bi, Pb, Cu, Cd. 
Add two drops of dilute 
H2S04, evaporate to 
small bulk, dilute with 
HaO, filter. 
Ppt. Ba, Sr, Ca. 
Wash, dissolve in 
HC2H3O2,addK2Cr04, 
filter. 
Filt. Mg, K, Na, Li, NH4. 
Add (NH4)„HP04, 
agitate, filter. 
Ppt. Fe, Ce, Al, Cr. 
Wash, dry, fuse on foil 
with Na2C03 and KN03, 
boil with H20, filter. 
Filt. Mn, Zn, Co, Ni. 
Acidify with HC2H302, add H2S, filter. 
Ppt. Zn, Co, Ni. 
Wash, dissolve in HC1 and 
HNOj, evaporate excess of 
acid, add KOH, filter. 
Filt. Mn. 
Add 
NH4OH 
and 
H2S. 
Flesh- 
colored 
ppt. 
Ppt. 
Mg. 
Filt. K, Na, Li, NH4. 
Evaporate to dryness, 
ignite, dissolve in 
solution of Na2 H P04, 
filter. 
Ppt. 
Ba. 
Yel- 
low. 
Filt. Sr, Ca. 
Add very dilute 
H2S04, allow to 
stand, filter. 
z- 
Filt. Bi, Cu, Cd. 
Add NH4OH in 
excess, filter. 
Ppt. 
Fe, 
Ce. 
Sepa- 
rate 
accor- 
ding 
to 
chart 
page 
S°- 
Filt. Al, Cr. 
Yellow if Cr be 
present. Divide 
in two parts. 
Ppt. 
As, Au, 
Pt. 
Separs 
ch< 
page 
Filt. 
Sb,Sn, 
te by 
irt 
57- 
Ppt. Co, Ni. 
Wash, dissolve in HC1, 
evaporate excess of 
acid, treat with KCN 
and HC1 according to 
chart page 4s, add 
KOH, filter. 
Filt. Zn. 
Add 
NH4HS 
white 
ppt. 
Ppt. 
Bi. 
White 
Con- 
firm 
as on 
page 
62 
Filt. Cu, Cd. 
Add H?S. 
Collect, boil 
the ppt. with 
dilute H2S04, 
filter. 
Ppt. 
Sr. 
Filt. Ca. 
Add NH4OH 
and 
(NH4)2C204, 
white ppt. 
Ppt. 
Li. 
Con- 
firm 
aby 
name 
test. 
Filt 
K, Na, NH4. 
Concentrate, 
add HC1 and 
PtCl4, yellow 
ppt. = K. 
Test for Na and 
NH4 in the ori- 
ginal solution. 
Al. 
Add 
NH4C1 
white 
ppt. 
Cr. 
Acidify 
with 
hc2h3o2 
and add 
Pb(A)2, 
yellow 
ppt. 
Ppt. Ni. 
Green. 
Confirm 
by borax 
bead. 
Filt. Co. 
Evaporate 
to dryness 
and test 
with borax 
bead. 
Ppt. 
Cu. 
Filt. Cd. 
Dilute, 
addH,S, 
yellow 
ppt. 
DIRECTIONS FOR THE DETECTION OF THE BASES IN A SOLUTION CONTAINING SOLUBLE SALTS 
OF ALL THE ELEMENTS HERETOFORE CONSIDERED. 68 
ANALYTICAL CHEMISTRY. 
SPECIAL PRECAUTIONS TO BE OBSERVED IN THE EXAMI- 
NATION FOR BASES. 
1. HC1 may precipitate basic salts of Bi and Sb; these are 
readily distinguished from the bases of Group VII by dis- 
solving on the further addition of HC1. It is best to make 
a preliminary test on a small quantity of the solution by the 
addition of an equal volume of dilute HC1. If no precipitate 
is formed, or if at first formed, is redissolved, the addition of 
an excess of HC1 to the portion taken for analysis may be 
avoided. 
2. The precipitation with HC1 should be performed in the 
cold, in order to prevent the PbCl2 from dissolving. 
3. Before commencing the analysis of a solution its action 
on litmus paper should always be noted. If alkaline a pre- 
cipitate may be caused by HC1, on account of the presence of 
substances soluble in alkalies, as As2S3 in (NH4)2C03, and 
AgCl in NH4OH. Or silicic acid from alkaline silicates may 
be precipitated, in which case it is necessary to precede the 
addition of HC1, by evaporation to dryness with HN03, filter- 
ing out the insoluble silica, and proceeding with the filtrate in 
the usual way. 
4. The addition of H2S should be slow, and interrupted 
from time to time to warm and agitate the solution, and so 
continued until it smells strongly of the reagent. 
5. When the presence of HN03 is suspected, the filtrate 
after the precipitation of Group VII, and before the addition 
of H2S, should be evaporated to dryness to expel the HN03, 
dissolved in water and, if necessary, acidified with HC1. BASES. 
69 
SPECIAL OBSERVATION. 
The separation into groups is accomplished by the following 
reagents, which are called group reagents:— 
HC1 precipitates AgCl, Hg2Cl2, PbCl2, 
HgS, PbS, Bi2S3, CuS, CdS, insolu-1 
ble in (NH4)2S, J 
. Group VI. 
H2S 
As2S3, Sb2S3, SnS, SnS2, Au2S3, "1 
PtS2, soluble in (NH4)2S, J 
. Group V. 
nh4hs 
FeS,Ce2(OH)6,Al2(OH)6,Cr2(OH)6, 
hydrates insoluble in NH4OH, j 
MnS,ZnS,CoS, NiS, hydrates soluble 
in NH4OH, } 
. Group IV. 
. Group III. 
(NH4)2COs “ 
(NH4)2HP04 “ 
BaC03, SrCOs, CaC03, \ 
MgNH4P04, J 
Not precipitated in a 
group. 
}. K, Na, Li, NH4, 
When one of these reagents fails to produce a precipitate, 
it indicates the absence of that group, and the next group 
reagent is immediately added. 70 
ANALYTICAL CHEMISTRY. 
SECTION II. 
ACIDS. 
GROUP I.—HYDROCHLORIC ACID, HYDROBRO- 
MIC ACID, HYDRIODIC ACID, HYDROCY- 
ANIC ACID, HYDROFLUORIC ACID. 
REACTIONS OF HYDROCHLORIC ACID (HC1). 
Use a solution of potassium chloride (KC1). 
1. AgN03 produces a white, curdy precipitate of silver 
chloride—AgCl, insoluble in HN03, readily soluble in 
NH4OH. 
This precipitate should be preserved to compare with those 
of AgBr and Agl. 
2. Hg2(N03)2 causes a white precipitate of mercurous 
chloride—Hg2Cl2, insoluble in HN03, blackening on the 
addition of NH4OH. 
3. Pb(C2H302)2 forms a white crystalline precipitate of lead 
chloride—PbCl2, soluble in 33 parts of boiling water. 
4. Warming with H2S04 and Mn02 causes the evolution 
of Chlorine, recognized by its odor and color. 
5. On warming with H2S04, hydrochloric acid is given off, 
recognized by its odor and intensely acid reaction ; also by 
the dense white fumes of NH4C1 produced by holding a rod 
moistened with NH4OH near the mouth of the tube. 
REACTIONS OF HYDROBROMIC ACID (HBr). 
Use a solution of potassium bromide (KBr). 
1. AgN03 produces a yellowish-white precipitate of silver 
bromide—AgBr, insoluble in HN03, slowly soluble in 
NH4OH. 
This precipitate should be compared with those of AgCl 
and Agl. 
2. HgCl2 forms a white precipitate of mercuric bromide— 
HgBr2, soluble in a large quantity of water. 
3. Hg2(N03)2 precipitates yellowish-white mercurous bro- 
mide—Hg2Br2. ACIDS. 
71 
4. H2S04, when concentrated, causes the evolution of red 
vapors of bromine; this occurs more readily in the presence 
of Mn02. 
5. Chlorine water and starch paste cause a yellow color, 
due to formation of starch bromide; in dilute solutions it is 
necessary to agitate the mixture with ether or chloroform, 
which will separate, carrying the bromine in solution, with a 
yellowish-red color. 
REACTIONS OF HYDRIODIC ACID (HI). 
Use a solution of potassium iodide (KI). 
1. AgN03 produces a yellowish precipitate of silver iodide 
—Agl, insoluble in HN03 and almost insoluble in NH4OH. 
2. Hg2(N03)2 precipitates green mercurous iodide—Hg2I2. 
3. HgCl2 causes a red precipitate of mercuric iodide—Hgl2, 
soluble in excess of either reagent. 
4. Pb (C2H302)2 produces a yellow precipitate of lead iodide 
—Pbl2, soluble in boiling water. 
5. Chlorine water and starch paste form a blue color of 
starch iodide, which color disappears on heating and returns 
on cooling, also destroyed by excess of chlorine water. 
6. A concentrated solution of CuS04 produces a white pre- 
cipitate of cuprous iodide—Cu2I2,—which is colored dark by 
free iodine according to the following reaction:— 
2CuS04 -f- 4KI = Cu2I2 + I2 + 2K2S04. 
When a reducing agent, as FeS04, is added at the same 
time, the liberation of iodine is prevented and the true color 
of the precipitate may then be seen. 
2CuS04 + 2KI + 2FeS04 = Cu2I2 -f K2S04 + Fe2(S04)3. 
This reaction is sometimes useful in separating iodine from 
chlorine and bromine. The reagent in this case should be a 
strong solution of one part CuS04 and three parts FeS04. 
HYDROFLUORIC ACID (HF). 
The evolution of intensely irritating fumes of this acid, on the addition of H2S04 
to calcium fluoride, which etch glass, is sufficiently characteristic. 72 
ANALYTICAL CHEMISTRY. 
REACTIONS OF HYDROCYANIC ACID (HCN). 
Use a solution of potassium cyanide (KCN). 
1. AgN03 produces a white precipitate of AgCN, soluble 
in KCN, sparingly soluble in NH4OH,and insoluble in dilute 
HN03. 
2. NH4HS evaporated with KCN to dryness, on a water 
bath, will give, on dissolving in water and adding Fe2Cl6, a 
deep blood-red color, due to formation of sulphocyanate. 
3. Fe2Cl6 and FeS04, then NaOH until a precipitate is pro- 
duced, heat and add HC1, will give a deep blue ppt. of 
ferric ferrocyanide (Prussian blue). 
4. HC1 or any other acid causes an evolution of HCN, 
readily recognized by its odor. Great care should be exercised 
in carrying out this test in order to avoid inhaling more than a 
trace of the fumes. 
SUMMARY OF TESTS WITH SOLUBLE SALTS OF GROUP I, 
HC1 
HBr 
HI 
HCN 
AgNOg 
White ppt. 
soluble in N H4OH 
Y ellowish-white 
ppt. slowly soluble 
in NH4OH. 
Yellowish ppt. 
insoluble in 
nh4oh. 
White ppt. 
sparingly soluble 
in NH4OH. 
Hga(N03)a 
White ppt. 
Y ellowish-white 
ppt. 
Greenish ppt. 
White ppt. 
HgCla 
No ppt. 
White ppt. 
Red ppt. 
No ppt. 
Pb(CaH3Oa)a 
White ppt. 
White ppt. 
Yellow ppt. 
White ppt. 
Chlorine water 
and 
starch solution. 
No color. 
Yellow color. 
Blue color. 
No color. 
DIRECTIONS FOR THE DETECTION OF THE ACIDS IN A SOLU- 
TION CONTAINING SOLUBLE SALTS OF GROUP I. 
Add HN03 and boil for some time; the HCN is driven off. 
This operation should be performed with caution, so as to avoid 
inhaling these poisonous fumes. 
After all the HCN has been driven off, add AgN03 and 
shake well; collect the precipitate of AgCl, AgBr, Agl on 
a filter, and, after washing, pour on the filter NH4OH ; AgCl 
will be dissolved, and may be detected in the filtrate by acidi- 
fying with HN03. ACIDS, 
73 
To detect HBr and HI take another portion of the original 
solution ; add a few drops of starch solution, and then, slowly, 
chlorine water, until, after agitation, the liquid smells of it. 
If there is a blue color HI is present. Continue the addition 
of chlorine water until the blue color is. destroyed, when, if 
the solution is yellow, HBr is indicated. This may be still 
further verified by agitating with chloroform, which will, after 
separating, assume a yellowish-red color if HBr is present. 
HYPOCHLOROUS ACID. 
CHLORIC 
HYDROXYL. 
HYDROSULPHURIC ACID. 
SULPHUROUS 
SULPHURIC “ 
THIOSULPHURIC 
NITRIC “ 
GROUP II. 
HYPOPHOSPHOROUS ACID. 
ORTHOPHOSPHORIC 
PYROPHOSPHORIC 
METAPHOSPHORIC 
BORIC 
CARBONIC 
SILICIC 
REACTIONS OF HYPOCHLOROUS ACID (HCIO). 
Use a solution of sodium hypochlorite (NaCIO). 
1. HC1 causes the evolution of chlorine, readily recognized 
by its odor and by the bleaching action it exerts on litmus 
paper, or solution of indigo. 
2. AgN03 produces a white precipitate of silver chloride 
—AgCl, soluble in NH4OH. 
3AgN03 + 3NaC10 = 2AgCl + AgC10s + 3NaN03. 
3. Pb(C2H302)2, with an excess of the solution, forms at first 
a white precipitate which, on heating, turns brown, owing to 
formation of lead peroxide—Pb02. 
REACTIONS OF CHLORIC ACID (HC103). 
Use a solution of potassium chlorate (KC103). 
I. H2S04 causes the evolution of yellow chlorine tetroxide 
—C1204, having a characteristic odor. When the dry salt is 
used this reaction takes place with some violence, or when 
organic matter is present, with explosion. 
2. On heating the dry salt in a crucible oxygen is given off, 74 
ANALYTICAL CHEMISTRY. 
readily detected by holding over the mass the glowing end of 
a match or taper. 
3. AgN03 produces no precipitate. This is an important 
distinction from HC1 and HCIO. When they are present they 
may be removed by just enough AgN03 to precipitate them, 
and the remaining solution tested for HC103by evaporating to 
dryness, heating to redness, dissolving in water and testing 
with AgN03, when, if a chlorate was present, a precipitate of 
AgCl will form, on account of the chlorate having been con- 
verted by heat into chloride. 
4. A modification of the above test may be used to dis- 
tinguish a chlorate from a nitrate, as the latter gives many 
of the reactions of the former. The residue, after heating a 
nitrate, will not give a precipitate when dissolved in water and 
tested with AgN03. 
5. AgN03 with a drop or two of strong H2S04, will cause 
a white precipitate of silver chloride, AgCl, owing to the 
decomposition of the chlorate by the strong acid. 
WATER (H20), HYDRATES AND OXIDES. 
Water is distinguished by having no odor or taste, not 
changing litmus, and evaporating without residue. 
The soluble hydrates KOH, NaOH, LiOH, Ca(OH)2, 
Sr(OH)2, Ba(OH)2, yield solutions with water which change 
red litmus paper blue. The insoluble hydrates give off steam 
when heated in a dry tube. 
The soluble oxides K20, Na20, Li20, CaO, SrO, BaO, are 
known by forming hydrates with water. The insoluble oxides 
are recognized by giving negative tests for acids when dis- 
solved and tested in the usual way. 
HYDROSULPHURIC ACID (H2S). 
Sufficient evidence of the presence of this acid is afforded 
by the characteristic odor. In the case of sulphides, first 
adding H2S04 and warming if necessary. A trace may be 
detected by holding over the mouth of the tube a piece of 
filter paper moistened with lead acetate, which will become 
black (lead sulphide) in the presence of H2S. ACIDS. 
75 
REACTIONS OF SULPHUROUS ACID (H2S03). 
Use a solution of sodium sulphite (Na2S03). 
1. In solution, uncombined, the free acid is recognized by 
the odor of burning sulphur, by strong bleaching action, by de- 
colorizing potassium permanganate solution, and by causing 
the evolution of hydrogen sulphide when added to a mixture 
of zinc and hydrochloric acid. 
2. Sulphites are distinguished by the characteristic odor of 
S02 on the addition of a strong acid. 
3. Salts of Ag, Hg or Pb, produce precipitates which 
blacken on heating, owing to formation of sulphides. 
4. BaCl2 with neutral solutions forms a white precipitate of 
barium sulphite—BaS03, soluble in HC1. 
5. With zinc and hydrochloric acid hydrogen sulphide is 
evolved. 
6. When to a dilute solution of a sulphite, zinc sulphate and 
sodium nitroprusside are added, a red color is produced, or 
when the solution is more concentrated, a flocculent reddish 
precipitate falls ; a few drops of potassium ferrocyanide render 
the reaction more distinct. This test will serve to distinguish 
sulphites from thiosulphates, as the latter do not cause this 
red color. 
REACTIONS OF SULPHURIC ACID (H2S04). 
Use dilute H2S04, or an alkali sulphate. 
1. BaCl2 produces a white precipitate of barium sulphate 
—BaS04, insoluble in boiling concentrated acids. 
2. Pb(C2H302)2 causes a precipitate of white lead sulphate, 
insoluble in dilute acids, but soluble in hot concentrated acids. 
Alcohol increases the delicacy of this reaction. 
3. A sulphate fused on charcoal with Na2C03, the fused 
mass placed on a bright silver coin and moistened with a drop 
of dilute HC1 will cause a black stain, due to formation of 
silver sulphide. This reaction is especially adapted to the 
detection of insoluble sulphates. 76 
ANALYTICAL CHEMISTRY. 
REACTIONS OF THIOSULPHURIC (HYPOSUL- 
PHUROUS) ACID (H2S203). 
Use a solution of sodium thiosidphate (Na2S203). 
1. H2S04 causes the evolution of sulphurous oxide—S02, 
recognized by the odor. A deposit of sulphur takes place 
at the same time, which is an important distinction from 
sulphites. 
2. AgN03 produces a white precipitate of silver thiosul- 
phate—Ag2S203, soluble in excess. After a time (immediately 
on heating) the precipitate becomes dark and then black, silver 
sulphide and sulphuric acid being formed. 
3. BaCl2 produces a white precipitate soluble in excess of 
H20, and decomposed by HC1. 
4. Added to a mixture of zinc and hydrochloric acid, 
hydrogen sulphide is evolved. 
REACTIONS OF NITRIC ACID (HN03). 
Use a solution of potassium nitrate (KN03). 
1. H2S04, on heating, will cause the nitric acid to volatilize. 
If copper turnings be added with the sulphuric acid colorless 
nitrogen dioxide—N202 will be given off, which in contact 
with air will form red nitrogen tetroxide—N204, readily 
recognized by the color and odor. If alcohol be added 
to the mixture, the characteristic odor of nitrous ether is 
developed. 
2. FeS04, acidified with H2S04, added in a test tube, so as 
to form a layer on a solution of nitrate, acidified with H2S04, 
will cause a dark layer to form at the line of contact. 
3. Indigo solution, strongly acidified with H2S04, is decolor- 
ized by a nitrate. 
4. Zinc and potassium hydrate cause the reduction of the 
acid to ammonia, which may be detected in the usual way 
This is a valuable test for distinguishing nitric acid in the 
presence of chloric acid. 
5. Heated on charcoal, deflagration takes place, the charcoal 
burning at the expense of the oxygen of the nitrate. ACIDS, 
77 
HYPOPHOSPHOROUS ACID (HH2P02). 
1. By ignition the hypophosphites are resolved into spontaneously inflammable 
hydrogen phosphide—PH3, and phosphate. 
2. AgN03 produces at first a white precipitate of silver hypophosphite— 
AgPI2P02, which soon becomes black, owing to the formation of metallic silver. 
3. HgCI2 in excess causes a white precipitate of mercurous chloride, Hg2Cl2 ; 
this takes place more rapidly on warming, and in the presence of HC1. 
REACTIONS OF ORTHOPHOSPHORIC ACID (H3P04). 
Use a solution of sodium phosphate (Na2HP04). 
1. AgN03 causes a light yellow precipitate of silver phos- 
phate— Ag3P04, soluble in HN03 and NH4OH. 
2. Fe2Cl6, in presence of sodium acetate, produces a yel- 
lowish-white, gelatinous precipitate of ferric phosphate— 
Fe2(P04)2. An excess of Fe2Cl6 must be avoided. 
3. (NH4)2Mo04, in neutral or acid solution, causes a yellow 
precipitate to separate slowly, which is ammonium phospho- 
molybdate—(NH4)3PO4(MoO3)10 + 2H20, insoluble in HN03, 
soluble in NH4OH. The precipitate forms more rapidly in 
the presence of NH4N03. 
4. Magnesia mixture (consisting of MgS04)NH4Cl,NH40H) 
causes a white precipitate of ammonium magnesium phos- 
phate—MgNH4P04 -f- 6HaO. Agitation facilitates the forma- 
tion of this precipitate. 
5. BaCl2, in neutral solution, produces a white precipitate of 
barium hydrogen phosphate—BaHP04. 
6. Albumen (white of egg) does not cause a precipitate. 
REACTIONS OF PYROPHOSPHORIC ACID (H4P207). 
Use a solution of sodium pyrophosphate (Na4P207), 
1. AgN03 precipitates white silver pyrophosphate—Ag3- 
P207, soluble in HN03 and NH4OH. 
2. MgS04 causes a white precipitate of magnesium pyro- 
phosphate—Mg2P207, soluble in excess of the reagent; from 
this solution a mixture of NH4OH and NH4C1 does not re- 
precipitate it in the cold. This reaction may be used to 
distinguish ortho- from pyrophosphoric acid. 78 
ANALYTICAL CHEMISTRY. 
3. Neither (NH4)2Mo04, nor albumen produces a precipitate. 
4. BaCl2, in neutral solution, precipitates white barium 
pyrophosphate— Ba2P207. 
REACTIONS OF METAPHOSPHORIC ACID (HP03). 
Use a solution of sodium metaphosphate (NaP03). 
1. AgN03 produces a white precipitate of silver meta- 
phosphate—AgP03, soluble in HN03 and NH4OH. 
2. Albumen forms a white precipitate with the free acid, 
and with the salts on the addition of acetic acid. This is an 
important distinction from the ortho- and pyro-acids. 
3. Neither (NH4)2Mo04 nor magnesia mixture produces a 
precipitate; should one form with the latter reagent, it is 
readily soluble in NH4C1. 
4. BaCl2 forms a white precipitate of barium metaphos- 
phate—Ba(P03)2. 
5. Solutions of the meta- and pyro-acids in water are con- 
verted into the ortho-acid by boiling. 
REACTIONS OF BORIC ACID (H3B03). 
Use a solution of sodium borate (Na2B407). 
1. AgN03 produces a white precipitate of silver borate, 
soluble in HN03. 
2. BaCl2 precipitates white barium borate, soluble in ex- 
cess of water and in NH4C1. 
3. H2S04 or HC1 causes the separation of the acid H3B03, 
in crystalline form, from strong solutions. 
4. Alcohol added to the acid, and ignited, burns with a 
characteristic green flame. In the case of salts, the addition 
of alcohol is preceded by that of a few drops of H2S04, in order 
to liberate the free acid. This test is most successful when the 
acid and alcohol are added to the dry salt and not to the solu- 
tion. Salts of copper when present should be removed before 
this test is applied, as they likewise impart a green color to the 
alcohol flame. ACIDS. 
79 
REACTIONS OF CARBONIC ACID (H2C03). 
Use a solution of sodium carbonate (Na2C03). 
1. All free acids except HCN and H2S decompose carbon- 
ates with effervescence. The escaping gas passed into a 
solution of Ba(OH)2 or Ca(OH)2 causes a white precipitate. 
2. BaCl2, in neutral solution precipitates white barium 
carbonate—BaC03, soluble in acids. 
3. CaCl2 precipitates white calcium carbonate—CaC03, 
soluble in acids with effervescence. 
REACTIONS OF SILICIC ACID (H4Si04). 
Use a solution of sodium silicate, (Na4Si04). 
1. Insoluble silicates are determined by fusing on platinum 
foil some of the fine powder with Na2C03, treating the fused 
mass with HaO and HC1, evaporating to dryness and redis- 
solving in H20, when the silica, Si02 will remain as a fine 
white precipitate. 
2. Before the blowpipe, with a bead of microcosmic salt, 
silica forms the so-called silica skeleton, which is very charac- 
teristic. 
3. Soluble silicates give gelatinous precipitates of silicic 
acid—H4Si04, on the addition of H2S04 or HC1; on evapo- 
rating this to dryness with a little HC1, and redissolving in 
H20, silica—Si02, remains. 
4. NH4C1 also precipitates H4Si04 when added to a soluble 
silicate. 
REACTIONS OF HYDROFERROCYANIC AND 
HYDROFERRICYANIC ACIDS. 
These acids have been sufficiently characterized under iron. 80 
ANALYTICAL CHEMISTRY. 
DIRECTIONS FOR THE DETECTION OF THE ACIDS IN A SOLU- 
TION CONTAINING SOLUBLE SALTS OF GROUP II. 
I. Try the solution with litmus paper ; if alkaline, hydrates, 
carbonates, borates, silicates and, possibly, phosphates, may be 
present. If the solution is acid, neutralize with NaOH before 
applying the following tests :— 
II. Evaporate a portion of the solution to dryness, and add 
concentrated H2S04. The following acids will give charac- 
teristic reactions, and may be recognized by the further appli- 
cation of tests previously given :— 
HCIO and HC103 give the odor of chlorine. They are 
readily distinguished by the chlorate deflagrating with char- 
coal, by its liberating oxygen on heating, and by the yellowish- 
green gas, C1204, which is given off. The addition of H2S04 
to a chlorate should be performed with very small quantities 
and with great care, on account of the tendency of the chlorine 
tetroxide to decompose with explosive violence. The most 
characteristic difference of HCIO is the white precipitate of 
AgCl which it gives with AgN03. 
H2S is recognized by its peculiar odor, and by blackening 
a piece of filter paper moistened with solution of lead 
acetate. 
H2S03 and H2S203 give off S02, readily recognized by its 
odor and bleaching property. They are distinguished from 
each other by H2S04 producing with a concentrated solution 
of thiosulphate a white precipitate of sulphur in addition to 
the liberation of S02. 
HN03 is detected by the peculiar acid fumes, which become 
red when metallic copper is added with the H2S04. All the 
special tests, previously given, of this acid should be applied 
before a conclusion is reached concerning the presence or 
absence of it. 
H3B03 is easily detected when, in addition to the H2S04, 
some alcohol is added and ignited. In the absence of salts of 
copper, the green flame is evidence of this acid. This green 
flame will show best when the acid and alcohol are added to the 
dry salt. 
HH2P02 is readily detected by its odor. ACIDS, 
81 
• h2co3 is recognized by effervescence with the dilute acid in 
the cold without odor. This test should be applied to the 
original solution, using preferably acetic acid. When other 
gases are given off at the same time, the C02 may be detected 
by passing into lime water, which will cause a precipitate of 
CaC03. 
H4Si04 is precipitated as a white, gelatinous mass when the 
H2S04 is added to a soluble silicate. Insoluble silicates will 
be treated of later. 
H4Fe(CN)6 and H6Fe2(CN)12 give the odor of HCN. Their 
presence is confirmed by the use of FeS04 and Fe2Cl6. 
III. The acids of this group not detected by H2S04 are sul- 
phuric and the three phosphoric acids. 
H2S04 is precipitated, on the addition of BaCl2, as BaS04, 
insoluble in HC1 or HN03. 
H3P04 is precipitated yellow on the addition of AgN03, the 
precipitate being soluble in PIN03 (distinction from HC1, HBr 
and HI). (NH4)2Mo04 is the most distinctive test. It causes 
a yellow precipitate when a few drops of the solution are added 
to some of the reagent in a test tube and warmed gently (dis- 
tinction from H4P207 and HP03). Magnesia mixture is also a 
characteristic test for this acid. 
H4P207 gives a white precipitate with AgN03. It is further 
distinguished from the ortho-acid by (NH4)2Mo04, and from 
the meta-acid by its behavior with MgS04, and with albumen. 
HP03 also gives a white precipitate with AgN03. It is best 
distinguished from the other varieties by albumen, by negative 
tests, and by boiling for some time and then applying the tests 
for the ortho-acid. 
H3As04 and H2Cr04 might be classified with this group, but 
a little care on the part of the student will detect them among 
the bases, and a little thought will tell him how to prove 
whether they are present as acids or bases. 82 
ANALYTICAL CHEMISTRY. 
GROUP III. 
ACETIC ACID. 
VALERIANIC ACID. 
STEARIC 
OLEIC 
LACTIC 
OXALIC 
SUCCINIC 
MALIC 
TARTARIC ACID. 
CITRIC “ 
CARBOLIC “ 
BENZOIC 
SALICYLIC “ 
GALLIC 
TANNIC 
REACTIONS OF ACETIC ACID (HC2H302). 
Use a solution of sodium acetate (NaC2H302). 
1. In the free state acetic acid is readily recognized by its 
odor. 
2. H2S04, added to an acetate and the mixture warmed, 
gives the characteristic odor. 
3. H2S04 and C2H5OH in equal volumes added to an 
acetate form acetic ether, readily recognized by its odor. 
4. Fe2Cl6, with a neutral acetate, forms a deep red color, 
due to ferric acetate—Fe2(C2H302)6. 
VALERIANIC ACID (HC5H902). 
The odor is sufficient evidence of the presence of this acid. This odor is 
developed by moisture and heat, and in the case of the salts by the addition of 
H2S04. 
Stearic acid is a white, fatty solid, melting at 69.2° C., giving when combined 
with potassium a soft soap, and with sodium a hard soap; from both it separates 
as an oily liquid on the addition of HC1, becoming solid on cooling. The lead 
salt, lead stearate, Pb(C18H3502)2 is insoluble in ether. 
STEARIC ACID (HQaHsA). 
OLEIC ACID (HC18H3302). 
Oleic acid is an oily liquid at ordinary temperatures, but becomes solid at 40 C. 
and remains so until the temperature rises to 140 C., when it again becomes liquid. 
Lead Oleate—Pb(C18H3302)2, prepared by saturating sodium oleate with acetic 
acid and adding lead acetate, is almost entirely soluble in ether. This is an 
important distinction from stearic acid. 
LACTIC ACID (HC3H503). 
Lactic acid is a colorless, syrupy liquid, of a slight unpleasant odor and a very 
sour taste, soluble in water, alcohol and ether, but insoluble in chloroform. When 
heated with H2SO+, CO is evolved. Lactates are all soluble in water, most of 
them sparingly; they are insoluble in ether. 
Hg2(N03)2 boiled with strong solution of a lactate, deposits crimson mercurous 
lactate—Hg2(CsH508)2. ACIDS. 
83 
REACTIONS OF OXALIC ACID (H2C204.2H20). 
1. K2Mn208, acidulated with H2S04, is decolorized. 
2. BaCl2 produces a white precipitate of barium oxalate— 
BaC204, sparingly soluble in acetic or oxalic acid, and freely 
soluble in HC1, HN03 and NH4C1. 
3. AgN03 precipitates white silver oxalate—Ag2C204, 
readily soluble in hot concentrated HN03, sparingly so in 
dilute acid, soluble in NH4OH. 
4. CaCl2 produces, even in highly dilute solutions, a white 
precipitate of calcium oxalate—CaC204, soluble in HC1 and 
HN03, but insoluble in acetic acid. 
5. FeS04 with dilute solutions causes a yellow color; with 
more concentrated solutions and warming, a yellow precipitate 
ferrous oxalate—FeC204, falls. This precipitate is insoluble 
in acetic acid, but is dissolved by HC1 and HN03. 
Use a solution of ammonium oxalate (NH4)2C204). 
SUCCINIC ACID (H2C4H404). 
1. BaCl2 produces no precipitate, but on the addition of alcohol a white pre- 
cipitate of barium succinate—BaC4H404, falls, soluble in NH4C1. 
2. Fe2Cl6 causes a brownish-red, bulky precipitate of ferric succinate— 
Fe2(C4H404)3. 
3. Pb(C2H302)2 precipitates white amorphous lead succinate—PbC4H404, 
soluble in HNOs. 
Use a solution of sodium succinate (Na2C4H404). 
MALIC ACID (H2C4H405). 
1. CaCl2 fails to give a precipitate with the acid or its salts, until the mixture 
is boiled, when calcium malate—CaC4H405.H20, separates. The addition of 
alcohol will cause this precipitate without boiling. 
2. Pb(C2H302)2 precipitates white lead malate—PbC4H405-3H20. 
REACTIONS OF TARTARIC ACID (H2C4H406). 
Use a solution of sodium tartrate (Na2C4H406). 
1. BaCl2 precipitates white barium tartrate—BaC4H406, 
soluble in ammonium salts, and in HC1. 
2. CaCl2 produces a white precipitate of calcium tartrate 
—CaC4H4064H20, ammonium salts prevent this precipitation. 
On adding KOH to this precipitate, it dissolves; boil this 
solution, calcium tartrate is again precipitated. Acetic acid is 
also a solvent for this precipitate. 84 
ANALYTICAL CHEMISTRY. 
3. AgN03 causes a white precipitate of silver tartrate— 
Ag2C4H4Oe; on adding a few drops of NH4OH to dissolve 
the precipitate, and boiling, a mirror of metallic silver forms 
on the test tube. The test tube must be perfectly cleaned for 
this reaction; to accomplish this, rinse the tube with KOH 
solution and then thoroughly with water. 
' 4. Ca(OH)2 in excess causes a precipitate of calcium tar- 
trate—CaC4H406. This precipitate is flocculent at first, in 
which state it is readily soluble in NH4C1, but after standing 
for some time it becomes crystalline, in which case it is in- 
soluble in NH4C1. 
5. Tartaric acid and tartrates char on heating, and with 
H2S04 the odor of burnt sugar is given off. 
REACTIONS OF CITRIC ACID (H3C6H507.H20). 
Use a solution of sodium citrate (Na3C6H607). 
1. BaCl2 produces a white precipitate of barium citrate— 
Ba3(C6H507)2, soluble in excess of water, in ammonium salts, 
and in acids. 
2. CaCl2 precipitates white calcium citrate—Ca3(C6H507)2 
more insoluble in hot than in cold water, soluble in cold 
NH4CI, but insoluble in KOH. 
3. AgN03 produces a white precipitate of silver citrate— 
Ag3C6H507; on boiling no metallic mirror is formed. 
4. Ca(OH)2, in excess, does not cause a precipitate until 
the mixture is boiled. This is an important distinction from 
tartaric acid. 
5. Citric acid and citrates char on heating, and with H2S04 
give off the odor of burnt sugar. 
CARBOLIC ACID (C6H5OH). 
1. HN03, with an aqueous solution of the acid, forms a yellow color, due to 
formation of picric acid—C6H2 (N02)3OH. 
2. Fe2Cl6 produces a violet blue color. 
3. A piece of pine wood dipped in the acid, and then exposed to the fumes ot 
HC1, becomes after a short time colored blue. 
4. Bromine water causes a white flocculent precipitate. 
5. Albumen is coagulated by the free acid. ACIDS. 
85 
BENZOIC ACID (HC7H502). 
1. Fe2Cl6 precipitates from neutral solutions, flesh-colored 
ferric benzoate—Fe2(C7H502)6, soluble in HC1, with separa- 
tion of benzoic acid. 
2. HC1 causes the separation of benzoic acid from cold so- 
lutions of the benzoates. 
3. BaCl2 and CaCl2 produce no precipitates with either the 
free acid or its salts. 
4. Benzoic acid is readily soluble in petroleum ether. 
5. The pure acid melts at I2I.4°C. 
1. Fe2Cl6 produces a deep violet color, which is very char- 
acteristic. 
2. Warmed with H2S04 and methyl alcohol, the fragrant 
odor of methyl salicylate (oil of wintergreen) is developed. 
3. HC1 causes the separation of the free acid from cold so- 
lutions of the salicylates. 
4. Salicylic acid is insoluble in petroleum ether; this fact 
enables one to separate mixtures of benzoic and salicylic 
acids. 
5. The pure acid melts at 155.5°C. 
SALICYLIC ACID (HC7H603). 
GALLIC ACID (HC7H505.H20). 
1. FeS04 produces no change. 
2. Fe2Cl6 produces a bluish-black precipitate, which disap- 
pears on heating. 
3. KOH, if not in excess, develops slowly a deep green 
color, which becomes red on the addition of acids. Alkaline 
carbonates cause the same green color, although more slowly. 
4. No precipitate is produced with either gelatin or the 
alkaloids. With the former, however, a precipitation takes 
place when gum is present. 
TANNIC ACID (C14H10O9) 
I. FeS04, when perfectly pure, causes no change; in'the 
presence of oxygen, however, a dark color rapidly develops, 
which on standing, slowly becomes a precipitate. 86 
ANALYTICAL CHEMISTRY. 
2. Fe2Cl6 produces a bluish-black precipitate. 
3. Normal solution of iodine mixed with a small quantity 
of ammonia, previously diluted with ten times its volume of 
water, produces a brilliant red color. This reaction will take 
place with only traces of tannin. 
4. Gelatin causes a white flocculent precipitate. This re- 
action is more delicate in the presence of small quantities of 
alum. 
5. Alkaloids produce white precipitates, soluble in acetic 
acid and alcohol. 
6. K(Sb0)C4H406 causes a white, gelatinous precipitate. 
Most metallic salts cause precipitates with tannin. 
DIRECTIONS FOR THE DETECTION OF THE ACIDS IN A SOLU- 
TION CONTAINING SOLUBLE SALTS OF GROUP III. 
As many of the acids in this group would indicate their 
presence by odor or physical appearance, a method of sepa- 
rating only the more important and closely related ones will 
be given. The list will therefore be limited to Acetic acid, 
Oxalic acid, Tartaric acid and Citric acid. 
If the solution is Acid to Litmus Paper, neutralize with NaOH. 
I. To a small portion add H2S04 and warm. Acetic acid 
will, if present, be detected by its odor. 
II. To another portion add NH4OH until slightly alkaline, 
and then CaCl2; allow to stand (avoiding heat) for ten minutes, 
and filter. 
Ppt. H2C204, H2C4H4Ob. 
Wash, pour on the filter HC2H302. 
Filt. H3CeH607. 
Boil to remove N H4OH, a white ppt. 
slowly forms on sides of the tube. 
Ppt. H2C204. 
Confirm by testing 
original solution with 
HC2H302 and CaCl2. 
Fill. H,C4H4Oe. 
Add NH40H until slightly 
alkaline, white ppt. 
Confirm by forming mirror with 
AgN03 in original solution. 87 
DETECTION OF BASES AND ACIDS. 
SECTION III. 
DETECTION OF BASES AND ACIDS. 
The bases must alzvays be determined first in a portion of the 
solution or powder, according to chart, page 6j. 
If only the alkali metals are present and the reaction is 
neutral, proceed to search for the acids according to Divi- 
sion I. 
If other than the alkali metals are present and the reaction 
is acid, proceed to search for the acids according to Division 
II. 
When the substance is not entirely dissolved by H20, HC1, 
or a mixture of HC1 and HN03) proceed for both bases and 
acids according to Division III. 
It will be seen from the above that before beginning the 
analysis of a solution it is necessary to try its action on litmus 
paper. If the reaction is alkaline it is necessary to neutralize 
with HC1, bearing in mind the precautions on page 68. 
When the analysis of a solid is to be performed, the action 
on litmus paper of the portion soluble in water is also to be 
noted, and the solution made neutral if necessary. 
All the physical properties of the solid should be carefully 
observed. 
Certain acids, when present, like benzoates and salicylates, 
are precipitated by HC1, in which case it would first be neces- 
sary to evaporate to dryness and heat to redness. 
SPECIAL OBSERVATION. 
When phosphates are present in acid solution, determined 
by adding a few drops of the solution to some HN03 and 
(NH4)2Mo04 and warming gently, the following chart must be 
used in the separation of Group IV. That is, bring the pre- 
cipitate containing Fe, Ce, Al, Cr, from chart, page 67, and 
work according to this. Oxalates also cause an exception 
EXCEPTION IN SEARCHING FOR BASES. 88 
ANALYTICAL CHEMISTRY. 
to the general chart for separation of bases ; when, therefore, 
their presence is suspected, heat the substance to redness 
before examining for bases. 
DIRECTIONS FOR THE ANALYSIS OF GROUP IV WHEN PHOSPHATES 
ARE PRESENT. 
In addition to Fe, Ce, Al, Cr, there may be present the phosphates of Ca, Sr, Ba, Mn, Mg. 
Dissolve the precipitate in HCl, add Na2HP04 in excess, then excess of NH4C2H302; boil, 
filter. 
Ppt. Fe2(P04)2 A12(P04)2 
Ce2(P04)2. 
Wash with hot H2G, then pour 
on the filter hot solution 
KOH. 
Filt. Ba, Sr, Ca, Mn, Mg, Cr. (If green Cr is present.) 
Add K2Cr04, warm, filter. 
Ppt. 
Ba. 
Yellow. 
Filt. Sr, Ca, Mn, Mg. 
Add very dilute H2S()4, allow to stand, filter. 
An excess of H2S04 should be avoided. 
Ppt. Fe, Ce. 
Wash, dissolve in 
HC1 and test a 
portion with 
K4Fe(CN)6 
for Fe. 
Filt. 
A12(P04)2. 
Acidify with 
hc2h3o2 
white 
ppt. = Al. 
Ppt. 
Sr. 
Confirm 
by 
flame 
test. 
Filt. Ca, Mn, Mg. 
Add NH4C2H302 and (NH4)2C204, 
filter. 
Ppt. 
Ca. 
White. 
Filt. Mg, Mn. 
Evaporate a portion to dryness and 
test with borax bead for Mn. 
To the remaining portion add 
Fe2Cl6, filter out the Fe2(P04)2, 
to filtrate add 
NH4C1, NH4OH and NH4HS, 
to separate Fe and Mn, 
filter and test filtrate with 
Na2HP04, white 
ppt. = Mg. 
To another por- 
tion add 
h3c6h5o,, 
then NH4OH in 
excess and 
Na2HP04, 
white 
precipitate = Ce. 
DIVISION I.—When the bases are the alkali metals, and the 
reaction is neutral. 
Rule I.—Evaporate a portion of the solution to dryness, and 
slowly heat to redness. 
If the mass chars, one or more of the following organic 
acids are indicated : Acetates, Tartrates, Citrates, Gal- 
lates, Tannates. 
Note.—Oxalates do not char, although if the heating take place slowly a 
grayish coloration may be noticed, the residue in this case giving off COs on the 
addition of H2S04. 
Fe2Cl6 immediately detects Tannates and Gallates. Gallic 
acid crystallizes in the cold on acidifying the solution. Tannic 
acid precipitates with solution of gelatin. The other organic 
acids are separated and detected according to chart, page 86. DETECTION OF BASES AND ACIDS. 
89 
Rule II.—Add to a second portion of the concentrated solu- 
tion, or the dry salt, strong H2S04, warm gently, and note any 
of the follozvmg effects:— 
Effervescence with dilute acid in the cold, 
no odor, 
► Carbonates. 
Effervescence on heating, no odor, . . 
' Oxalates, confirm by CaCl2 and 
Effervescence with dilute acid on heating, ■ 
odor of H2S, 
HCjHjOj. 
- Sulphides. 
Odor of S02, • . 
Sulphites. 
Odor of S02 with precipitation of S, . . 
Thiosulphates. 
Dark brown color and violet fumes, . . . 
Iodides, confirm by starch and Cl 
Dark red color and reddish fumes,. . . . 
water. 
Bromides, confirm by starch and Cl 
Odor of HCN, 
t water. 
Cyanides. 
OdorofHCN with crystalline deposit, . . 
Ferro- or Ferri-cyanides, confirm by 
Odor of acetic acid, 
Fe2Cl6 and FeS04. 
Acetates. 
With dilute acid, odor of chlorine, . . . 
Hypochlorites. 
With strong acid, odor of chlorine, . . . 
Chlorates. 
Strongly acid suffocating fumes, . . . • 
Chlorides, confirm by AgN03. 
Strongly acid fumes, becoming red when 
Nitrates, confirm by FeS04 and 
metallic Cu is added, 1 
H2S04, also by indigo solution. 
Characteristic odors 
' Benzoates, Succinates, Valerianates, 
l Carbolates, Hypophosphites. 
Rule III.— To a third portion add BaCl2. 
A white precipitate insoluble in HC1 indicates Sulphates. 
Rule IV.— To a fourth portion add CaCl2. 
A white precipitate soluble in excess of H20 indicates 
Sulphates ; if insoluble in excess of H20 and in acetic acid, 
Oxalates are indicated. A white precipitate soluble in KOH, 
reprecipitated on boiling, indicates Tartrates ; confirm by 
boiling with NH4OH and AgN03, forming a mirror of silver 
on the test tube. 
A white precipitate soluble in NH4C1 and reprecipitated on 
boiling, indicates Citrates. Citrates and Tartrates are also 
detected by heating some of the dry salt with H2S04, when 
the odor of burnt sugar is developed. 
When all four of these acids are suspected to be present, 
treat the solution with HC1 and BaCl2 to remove sulphates ; 90 
ANALYTICAL CHEMISTRY. 
then neutralize the excess of HC1 with NaOH, and proceed 
according to chart, page 86. 
Rule V.— To a fifth portion, acidified with H N 03, add AgN08. 
A white, curdy precipitate, immediately and completely 
soluble in NH4OH, indicates Chlorides. If slowly soluble 
in NH4OH, Bromides may be present; confirm by chlorine 
water and starch. 
A yellowish precipitate insoluble in NH4OH = Iodides. 
A white precipitate soluble in strong, hot HN03 = Cyanides. 
When bromides and iodides are present, they may be separated 
by adding to the original solution chlorine water and starch ; 
continue the addition of chlorine water until the blue color of 
the starch iodide is discharged, then shake with chloroform ; 
when a yellowish-red color is imparted to the separated chlo- 
roform, Br is indicated. 
For further instruction regarding the separation of these 
acids, consult chart, page 72. 
Rule VI.—Add to a sixth portion of the solution Magnesia 
Mixture. 
A white precipitate indicates Phosphates and Arsenates. 
In another portion separate the As by HC1 and H2S, and 
repeat the test for H3P04 in the filtrate by neutralizing and 
adding magnesia mixture. 
Rule VII.— To a seventh portion add Fe2Cl6. 
This reagent readily indicates Tannates, Gallates, Ferro- 
and Ferri-cyanides. 
Rule VIII.— To an eighth portion apply the special test with 
H2S04 and FeS04 for Nitrates. 
Rule IX.— To a ninth portion of the dry salt, or the solution 
evaporated to dryness, add a few drops of strong H2S04 and iy 
to 20 drops of alcohol, mix well and apply flame. Borates, if 
present, will impart a green color to the flame. DETECTION OF BASES AND ACIDS. 
91 
DIVISION II.—When the solution is acid, or, if a solid, requires the 
use of HC1 to dissolve it. 
After determining the bases, bearing in mind the exception in 
regard to phosphates, a portion of the substance is tested accord- 
ing to Rules I and II, DIVISION 1. 
A second portion is treated as follozvs:— 
Boil with strong solution Na2COs ; filter. 
Ppt. consists of the 
carbonates, hydrates, 
and oxides of the 
bases present, and 
is disregarded. 
Filt. contains the acids in combination with Na. 
Divide in two parts. 
First portion. 
Add HC1 until slightly acid, 
boil to remove C02, neu- 
tralize with NaOH, and 
test according to Rules 
III, IV, VI, VII, VIII and IX, 
DIVISION I. 
Second portion. 
Add HNOj until slightly acid, 
boil to remove C02, neutral- 
ize with NaOH, and test 
according to Rule V, 
and if necessary, 
III, IV, VI, VII and IX, 
DIVISION I. 
DIVISION III.—When the substance is not entirely dissolved by 
H20, HC1, ora mixture of HC1 and HNOs. 
Add water and filter; the filtrate, if it contains solid matter, 
is examined for bases by chart, page 67, and for acids by 
Division I. The residue, after treatment with water, is treated 
with HCl,or if necessary, with a mixture of HC1 and HN03, 
and, if it contain any solid matter, examined for bases by 
chart, page 67, and for acids according to Division II. The 
insoluble residue consists of one or more of the following 
substances : Sulphates of Ba, Sr, Ca, Pb ; Chlorides, Bromides 
and Iodides of Ag, Pb ; certain Oxides which have been highly 
heated, as Fe203, A1203, Cr203 and Sn02; Silica and Silicates ; 
Carbon and Sulphur (the last two are readily detected in the 
preliminary examination by heat). This insoluble residue is 
mixed with about four times its weight of dry sodium car- 
bonate and fused on a platinum foil; the mass is boiled with 
water, and the soluble portion examined for bases by chart, 
page 67, and for acids according to Division I. The in- 
soluble portion is treated with HC1, evaporated to dryness, 
redissolved in water with a little HC1, and examined, after 
determining the bases, according to Division II. Any insolu- 
ble substance remaining is probably silica, which may be 
determined by a bead of salt of phosphorus before the blow- 92 
ANALYTICAL CHEMISTRY. 
pipe, showing the skeleton of silica. If the metals lead and 
silver are present in this insoluble residue, they are best deter- 
mined by treating the residue insoluble in HC1, by HN03 and 
examining according to Division II. Many of these insoluble 
substances are detected in a preliminary examination by the 
use of the blowpipe, as explained in the tests given under the 
individual elements. 
The following chart, taken largely from Dr. Muter’s Ana- 
lytical Chemistry, will be found useful to consult after deter- 
mining the bases and before commencing the examination for 
acids. DETECTION OF BASES AND ACIDS. 
Bases 
found. 
If soluble in water, test for the 
following acids. 
If insoluble in water, but 
soluble in acids, test for 
the following acids. 
If insoluble 
in acids, fuse 
with Na2C03 
and test for, 
Ag. 
hno3,hno2,h2so4, 
hc2h3o2, 
Oxide, Sulphide, H2C03, 
h3po4,hcn,h2c2o4, 
H2C4H4Ofi, H3C6H507. 
Cl, I, Br. 
Hg(ous), 
hno3,hc2h3o2,ii2so4, 
Oxide, Sulphide, Cl, I, 
Oxysulphate. 
Cl, I, Br. 
Hg(ic), 
Cl, HNOs, H2S04, HC2H302, 
Oxide, Sulphide, I, 
Oxysulphate. 
Sulphide, 
Iodide. 
Pb, 
HC21I302, hno3, 
0xide,Sulphide,H2C03, 
H3P04, H2C204. 
h2so4,ci, 
1, H2Cr04. 
Bi, 
HN03, Cl, H2S04, HC2H302, 
h3c6h5o7, 
Oxynitrate, Oxychloride, 
Oxide, Sulphide, H2C03, 
h3po4. 
None. 
Cu(ic), 
Cl, hno3, h2so4, hc2h3o2, 
Oxide, Sulphide, II2C03, 
H3P04, Oxyacetate. 
(6 
Cu(ous), 
h2so4, 
Oxide, I. 
(i 
Cd, 
Cl, hno3,i,h2so4, 
Oxide, Sulphide, H2C03, 
h3po4. 
u 
Sb, 
Cl, h2c4h4o6, 
Oxide, Sulphide, 
Oxychloride. 
u 
Sn(ic), 
Cl, 
Oxide. 
a 
Sn(ous), 
Cl, h2so4, 
Oxide, Sulphide, H3P04, 
H2Cr04. 
« 
Au, 
Cl, 
Sulphide. 
a 
Pi, 
Cl, 
Sulphide. 
it 
Fe(ic), 
ci, hno3, h2so4, iic2h3o2, 
Oxide, Sulphide, 1, 
h3po4. 
a 
Fe(ous), 
ci,h2so4,i, 
Oxide, Sulphide, H2COs, 
h3po4. 
t< 
Al, 
Like Iron, 
Oxide, H3P04. 
u 
Ce, 
Cl, 
Oxide, H2C204. 
Oxide. 
Cr, 
ci, h2so4, hc2h3o2, hno3, 
Oxide, H3P04. 
Oxide. 
Mn, 
Like Chromium, 
Oxide, Sulphide, H2C03. 
h3po4. 
None. 
Zn, 
Like Chromium, 
Like Manganese. 
“ 
Co, 
ci,hno3,h2so4, 
Oxide, Sulphide, H2C03, 
h3po4. 
a 
Ni, 
Like Cobalt, 
Like Cobalt. 
“ 
Ba, 
ci, hno3,hc2ii3o2, 
h2co3,h3po4,h2c2o4, 
H2Cr04. 
Sulphate. 
Sr, 
Like Barium, 
Like Barium. 
i< 
Ca, 
Like Barium, 
Like Barium. 
Sulphate. 
Mg, 
C1,H2S04, 
Oxide, H2C03, H3P04. 
None. 
Li, 
All Radicals, 
h3po4. 
“ 
K, 
All Radicals, 
K2PtCl6, KHC4II406. 
“ 
Na, 
All Radicals, 
None. 
<< 
nh4, 
All Radicals, 
None. 
“ 94 
ANALYTICAL CHEMISTRY. 
SECTION IV. 
SOME OF THE REACTIONS AND TESTS OF PURITY OF THE 
MORE IMPORTANT ORGANIC COMPOUNDS. 
“ A heavy, clear, colorless, diffusive liquid, of a character- 
istic, pleasant, ethereal odor, a burning, sweet taste, and a 
neutral reaction.” 
CHLOROFORM (CHC13). 
1. Agitate about 5 c. c. with twice its volume of distilled 
water, separate and test the aqueous portion with litmus 
paper; no change should be produced, showing the absence 
of acids. To a small portion of the water add AgN03; a 
white precipitate would indicate HC1, the result of decom- 
position. To another portion add solution of KI; a reddish 
color would indicate Chlorine. Another portion warmed with 
solution KOH should not become colored, showing the ab- 
sence of aldehyde. 
2. 10 c. c. of chloroform mixed with half its volume of 
strong H2S04, and, after agitation, set aside, should not cause 
a color in either liquid for twenty-four hours. 
3. H2S04 and K2Cr207, mixed with an equal volume of 
chloroform and warmed, will become green, indicating the 
presence of a small quantity of alcohol. 
IODOFORM (CHI3). 
Iodoform occurs in small, lemon-yellow, lustrous crystals, 
having a saffron-like odor, and an unpleasant, slightly sweetish, 
iodine-like taste. Its solutions have a neutral reaction. 
I. Digest a small quantity of iodoform with alcoholic solu- 
tion of KOH, acidify with dilute HN03, and add starch paste, 
when the blue color of starch iodide will be developed. 
2. Agitate a small quantity of iodide with distilled water 
and filter; the filtrate should not affect litmus paper. Add to 
a portion of the filtrate AgN03; a yellowish precipitate would 
indicate iodides as an impurity. Evaporate another small REACTIONS AND TESTS OF ORGANIC COMPOUNDS. 
95 
portion to dryness; no residue should remain, indicating the 
absence of soluble impurities. 
ALCOHOL (C2H5OH). 
“ A transparent, colorless, mobile and volatile liquid, of a 
characteristic, pungent and agreeable odor, and a burning 
taste. It should not change the color of blue litmus paper, 
previously moistened with water. It boils at 78° C. and is 
readily inflammable, giving a blue flame without smoke.” 
1. Add to 5 c. c. of alcohol in a test tube an equal volume 
of H2S04to which has been added a small quantity of K2Cr207, 
and warm; the odor of aldehyde will be developed, the liquid 
in the tube becoming green. 
2. On heating 10 c. c. of alcohol with about one-fourth its 
volume of H2S04 the odor of ether will be evolved. 
3. Equal volumes of H2S04, acetic acid and alcohol, on 
warming, will give the odor of acetic ether. If much amyl 
alcohol be present the odor of amyl acetate may also be 
recognized. 
4. Evaporate 10 c. c. of alcohol to about one-fifth its bulk, 
and add an equal volume of strong sulphuric acid; a reddish 
color will indicate amyl alcohol. 
5. Mix in a test tube 5 c. c. of alcohol with an equal volume 
of solution KOH ; an immediate darkening will indicate the 
presence of methyl alcohol, aldehyde or oak tannin. 
CHLORAL HYDRATE (C2C130H H20). 
{Chloral, U. S. P.) 
Chloral hydrate occurs in colorless, transparent crystals, 
slowly vaporizing on exposure to the air, having an aromatic, 
penetrating and slightly acrid odor, bitterish, caustic taste 
and a neutral reaction. It liquefies when mixed with carbolic 
acid or with camphor. 
I. In a test tube heat to boiling a small quantity of chloral 
hydrate and water and add solution KOH ; a vaporous milky 
mixture of chloroform results, readily recognized by its odor, 
while formate of potassium remains in solution. 
2. To a hot aqueous solution of chloral hydrate in a test 96 
ANALYTICAL CHEMISTRY. 
tube, add AgN03, then NH4OH and boil; a silver mirror is 
deposited on the tube. 
3. To another portion of an aqueous solution add a few 
drops of HNO3 and AgN03; no precipitate should be pro- 
duced, indicating the absence of HC1. 
4. Add NH4HS to an aqueous solution; a reddish-brown 
coloration results, which on standing deposits a reddish-brown 
compound mixed with sulphur. 
5. A crystal of chloral hydrate heated on platinum foil 
should volatilize without residue. 
“A clear, colorless liquid, of a syrupy consistence, oily to 
the touch, hygroscopic, without odor, very sweet and slightly 
warm to the taste, and neutral in reaction.” 
GLYCERIN (C3H5(OH)3). 
1. Make a borax bead in the loop of a platinum wire, 
moisten it with a small quantity of glycerin (previously ren- 
dered alkaline with a dilute solution of soda), and hold in the 
colorless flame of a Bunsen burner; the flame is tinged green, 
owing to the liberation of boric oxide. 
2. Warm 5 c.c.’of glycerin with an equal volume of H2S04; 
no coloration should result, showing the absence of cane sugar. 
3. Heat a few drops on platinum foil; no residue should be 
left. 
4. A portion heated nearly to the boiling point with an 
equal volume of Fehling’s solution, should not deposit a red 
precipitate of cuprous oxide, Cu20. The same result should 
be obtained if the glycerin be previously boiled with a small 
quantity of HC1, showing the absence of sugars, starch and 
dextrin. 
GLUCOSE (C6H1206). 
(Dextrose, Grape Sugar.) 
i. To a dilute solution of glucose, add an equal volume of 
Fehling’s solution and boil; a red precipitate of cuprous 
oxide, Cu20, is deposited. 
2. A strong solution mixed with H2S04, in the cold; no 
change should occur; if, however, organic impurities are 
present, a dark coloration will result. REACTIONS AND TESTS OF ORGANIC COMPOUNDS. 
97 
3. To a strong solution of glucose add a solution of KOH 
and warm ; the solution becomes yellow and, finally, brown. 
SACCHAROSE (C12H22Ou). 
[Cane Sugar.) 
1. Boil a moderately dilute solution of sugar with a few 
drops of HC1, neutralize and add Fehling’s solution; red 
cuprous oxide, Cu20, will be precipitated. If the sugar be 
pure this reaction will not take place unless it be previously 
boiled with the acid. 
2. To a cold saturated solution add an equal volume of 
concentrated H2S04; there is an immediate blackening and 
swelling until the tube is filled with a dry coke. 
AMYLOSE (C6H10O5)x. 
(,Starch.) 
Starch usually occurs as a fine, white powder, insoluble in 
cold water, soluble in boiling water forming a thick paste 
when cold. 
1. To a cold dilute solution add potassium iodide and a 
few drops of chlorine water; the deep blue color of starch 
iodide results. This is decolorized by heating and reappears 
on cooling. It is bleached by excess of chlorine water. 
2. Heat a small quantity of starch on platinum foil; it 
should char and finally disappear without residue. 
3. Triturate a small quantity of starch with cold water, 
filter, and test the filtrate with litmus ; no change should take 
place. 
MORPHINE (C17H19N03.H20). 
I. To some crystals of morphine sulphate on a crucible 
lid, or other white porcelain surface, add a small quantity of 
neutral solution of ferric chloride; a blue color is produced, 
which rapidly fades. The presence of free sulphuric acid 
must be avoided. 
2. Similarly, to another portion add a drop of strong HN03; 
a red color at first forms, which rapidly becomes yellow. 
3. To a solution of the sulphate add carefully a dilute 98 
ANALYTICAL CHEMISTRY. 
solution of KOH ; a white precipitate forms, which is readily 
soluble in excess. NH4OH and Na2C03 produce the same 
precipitate insoluble in excess. 
4. A few crystals on a white porcelain surface should entirely 
dissolve in strong H2S04 without color, becoming reddish on 
standing, and on the addition of a crystal of K2Cr207 a green- 
ish color is produced. A purple or violet color would indi- 
cate the presence of strychnine or brucine. 
STRYCHNINE (C^H^N-A). 
1. A small quantity of the alkaloid dissolved in a drop or 
two of H2S04, and a small crystal of K2Cr207 added will cause 
a deep blue color, rapidly changing to violet, then cherry-red, 
and finally fading. 
2. Strong HN03 added to strychnine should not cause 
more than a very faint red color, showing the absence of 
brucine. 
3. KOH added to a solution of the sulphate causes a white 
precipitate, insoluble in excess. 
1. To an aqueous solution of the sulphate add fresh chlorine 
water and then NH4OH in slight excess, when a green color 
is formed, due to Thalleioquin. 
QUININE (UH^NAdH.O). 
2. To a very dilute, slightly acid solution, add chlorine 
water, then a small quantity of K4Fe(CN)6, and finally a few 
drops of NH4OH ; a red color is produced, which rapidly 
fades. 
3. Add KOH to an aqueous solution of the sulphate; a 
white precipitate forms, insoluble in excess. Under the same 
circumstances NH4OH produces a white precipitate, soluble in 
excess. 
CINCHONIDINE (QoH^O). 
i. To an aqueous solution of the sulphate add NH4OH ; a 
white precipitate forms almost insoluble in excess. 
2. KNaC4H406 added to a neutral solution produces a white 
precipitate. REACTIONS AND TESTS OF ORGANIC COMPOUNDS. 
99 
3. To a cold saturated solution of the sulphate add KNa- 
C4H4Os in slight excess, allow to stand a short time at about 
150 C., filter and to the filtrate add a drop of NH4OH ; not 
more than a slight turbidity should result, showing the absence 
of more than small quantities of the other cinchona alkaloids. 
CINCHONINE (C20H24N2O). 
1. To a solution of the sulphate add NH4OH ; a white pre- 
cipitate results, insoluble in excess. 
2. To a saturated solution add chlorine water and then 
NH4OH in slight excess; a white precipitate is formed, free 
from green color. 
ATROPINE (C17H22N03). 
1. Add to a small quantity of the alkaloid a drop of H2S04, 
no change occurs ; add to the mixture HN03, no color is pro- 
duced, showing the absence of and difference from morphine; 
similarly when a crystal of K2Cr207 is added no change occurs, 
indicating the absence of and difference from strychnine. 
2. To a small quantity of atropine in a test tube add H2S04 
and heat until it turns brown ; dilute with water and boil, when 
the characteristic odor of orange flowers will be developed. 
3. Boil with K2Cr207 and dilute H2S04, and add KOH in 
excess, when a herring-like odor will be given off. 
CAFFEINE (C8H10N4O2.H2O). 
i. To an aqueous solution of the alkaloid add a solution of 
potassio-mercuric iodide; no precipitation occurs, indicating 
the absence of and distinction from most other alkaloids. 
2. Treat a small quantity with chlorine water and evaporate 
to dryness, when a yellow mass will remain, which on moist- 
ening with NH4OH will become purple. 
VERATRINE. 
I. To a small quantity on a crucible lid, add a drop of 
HN03; it dissolves with a yellow color, which soon passes to 
a reddish-yellow, then to an intense scarlet and, finally, to a 
violet-red. 100 
ANALYTICAL CHEMISTRY. 
2. Similarly, to another portion add H2S04; very little 
change takes place, but on warming a resinous mass is formed, 
which dissolves with a deep red color. 
3. Repeat the above reaction, using HC1, when a blood-red 
color results. 
SALICIN (C13H1807). 
1. Add cold, concentrated H2S04 when a red color is pro- 
duced, which disappears on adding water, a dark red powder 
insoluble in water and in alcohol being deposited. 
2. The aqueous solution of salicin should not be precipi- 
tated by tannic or picric acids, nor potassio-mercuric iodide, 
indicating the absence of and difference from most alkaloids. 
SANTONIN (C15H1808). 
1. To an alcoholic solution of KOH add santonin; it will 
dissolve with a red color, gradually becoming colorless. 
2. To another portion of santonin add cold concentrated 
H2S04; there is gradually formed a yellow color, becoming 
red and, finally, brown. 
3. Santonin should give no precipitates with the ordinary- 
reagents for the alkaloids. 
Note.—All the preceding organic compounds may be tested 
for mineral impurities by heating on platinum foil; no residue 
should be left. • PART THIRD. 
QUANTITATIVE ANALYSIS. PART III.—QUANTITATIVE CHEMICAL 
ANALYSIS. 
SECTION I. 
GRAVIMETRIC ESTIMATION. 
PRELIMINARY DIRECTIONS. 
The following course supposes the student to be familiar 
with the operations in qualitative analysis, as well as with the 
process of weighing, or in a position to be instructed by some 
one familiar with the use of analytical balances and weights. 
The following rules, however, should be kept constantly before 
those using a fine balance :— 
1. Never put any chemical substance directly on the pan of a 
balance, but always on a clean, dry watch crystal. 
2. Never put on or take off a weight, or anything else, from a 
balance when it is resting on the knife edges. 
3. All volatile acids and other corrosive substances should be 
weighed in stoppered weighing tubes. 
4. Do not keep the balance open longer than is absohitely 
necessary. 
PRECIPITATION. 
This should be conducted in beaker glasses. Avoid a large 
excess of the reagent, except in a few cases where it is 
directed. The filtrate should always be tested with a few 
drops more of the reagent, to determine that the precipitation 
has been complete. Distilled water is used in all quantitative 
determinations. 
FILTRATION. 
A filter paper having a very small ash should be used; the 
Swedish is probably the best. Plain filters are always used, 
made by folding a circular piece of the paper twice, and open- 
ing so as to form a cone, which will fit exactly into a funnel, 
and allow all the liquid to pass through the precipitate and 104 
QUANTITATIVE ANALYSIS. 
escape at the lower point into the neck of the funnel. The 
precipitate should be allowed to settle in the beaker after pre- 
cipitation; the clear supernatant liquid is then poured on the 
filter, using a glass rod against the edge of the beaker, to 
conduct it into the funnel without loss. The precipitate is 
then washed into the filter with the aid of the wash bottle, 
and the washing continued, until by the appropriate test it is 
found to be complete. The washing, in many cases, is to be 
done with hot water. 
DRYING. 
The filter and its contents are usually dried in the funnel 
when intended for ignition, but if to be weighed without burn- 
ing, it is transferred to two watch crystals clamped together, 
and dried in an air bath, at from 120° to 130° C., until it 
ceases to lose weight. 
IGNITION. 
This is usually done in a platinum or porcelain crucible, the 
precipitate being separated as much as possible from the filter, 
and the latter burned on the lid, then placed in the crucible 
with the precipitate and the whole brought to a low or bright 
red heat, as the case requires. It is then cooled in a desicca- 
tor, until nearly the temperature of the balance room, when it 
is ready for weighing. 
Examples for Practice by Gravimetric Estimation. 
BARIUM CHLORIDE (BaCl2.2H20). 
Estimation of the Barium. 
Weigh about .500 gram of the pure salt, dissolve in 100 c. c. 
of water in a beaker of about 200 c. c. capacity. Add dilute 
sulphuric acid, drop by drop, as long as any precipitate is pro- 
duced, and then boil the mixture. Allow the precipitate to 
settle, pour the clear supernatant liquid into a filter, boil with 
a fresh quantity of water, and after pouring in the clear liquid, 
transfer the precipitate to the filter. Remove the small part 
of the precipitate which adheres to the sides of the beaker 
with a glass rod having a short section of rubber tubing 
slipped over the end. Wash the precipitate with hot water, 105 
GRAVIMETRIC ESTIMATION. 
until the washings cease to cause a turbidity with solution of 
barium chloride, then dry, ignite and weigh as barium sul- 
phate, BaS04. 
The following calculation will serve to illustrate the general 
method:— 
Ba = 136.8 = 56.16 per cent. 
Cl2 = 70.8 ■= 29.06 “ 
2H20 = 36. = 14.78 “ 
243.6 100.00 
Supposing we take .365 gram of barium chloride and find 
•349 gram of barium sulphate, we have as— 
Molecular wt. 
of BaS04 
232.8 
Atomic wt. 
of Ba. 
136.8 
Wt. of BaS04 
found. 
•349 
Wt. of Ba. 
found 
•205 + 
Wt. of BaCl2.2H20 
taken. 
As .365 
Wt. of Ba 
found. 
.205 + 
Per cent, of Ba 
found. 
56.16 
IOO 
ESTIMATION OF THE CHLORINE. 
To accomplish this weigh another portion of the salt, 
dissolve in about 100 c.c of water, acidify with HN03, add 
AgN03 in slight excess and heat to the boiling point. Stir 
the mixture while hot, until the precipitate coagulates, allow 
to settle, decant the clear liquid to a filter, then rapidly transfer 
the precipitate, and wash it well with hot water, dry and trans- 
fer the precipitate to a crucible, burning the filter on the cru- 
cible lid. If an appreciable amount of the precipitate adheres 
to the filter it is best to moisten the residue after burning with 
a drop or two of nitrohydrochloric acid, to reconvert any me- 
tallic silver which may have formed into chloride. The lid, 
crucible and contents are heated in an air bath to 130° and 
weighed. The whole operation should be conducted as rap- 
idly as possible, as silver chloride is acted on by the light. 
COPPER SULPHATE (CuS04.5H20). 
Weigh from .500 to 1 gram, dissolve in a small quantity 
of water in a porcelain dish ; add potassium hydrate in slight 
excess and boil, to convert the cupric hydrate first formed into 
Estimation of the Copper. 106 
QUANTITATIVE ANALYSIS. 
cupric oxide. Wash two or three times by decantation, then 
transfer to the filter and again wash with hot water until the 
filtrate ceases to change red litmus paper. The precipitate is 
separated from the filter, and the latter after burning is moist- 
ened with a drop of nitric acid and again heated, to reconvert 
any metallic copper which may have been formed ; this residue 
is then transferred to the crucible with the precipitate and the 
whole ignited and weighed as cupric oxide CuO. 
ESTIMATION OF SULPHURIC ACID. 
Weigh another portion of the copper sulphate, dissolve and 
precipitate the sulphuric acid with barium chloride. Treat the 
precipitate exactly as in the estimation of barium. 
POTASSIUM NITRATE (KNO,). 
Estimation of the Potassium. 
To a weighed portion of the salt dissolved in a small 
quantity of water, add hydrochloric acid, and evaporate to 
dryness, dissolve in a small quantity of water, add platinic 
chloride in excess and evaporate nearly to dryness on a water 
bath, keeping the water in the bath just below the boiling 
point. Add 80 per cent, alcohol to the residue, allow to stand 
some time, then transfer to a small weighed filter, and wash 
repeatedly with small quantities of the same strength alcohol. 
Dry the precipitate first in the air, then in the air bath at 130° 
C., and weigh between two watch crystals. From the potas- 
sium platinic chloride (K2PtCl6) the potassium is readily calcu- 
lated. 
Fuse in a platinum crucible, at a low temperature, some 
potassium nitrate, and pour out on a warm porcelain slab. 
Ignite two or three grams of silica in the crucible and mix 
with it about one-fourth its weight of the fused and powdered 
nitre. Bring the mixture to a low red heat and keep at that 
temperature until it ceases to lose weight. The loss indicates 
the amount of N205 from which the HN03 can be calculated. 
Sulphates and chlorides are not decomposed at this tempera- 
ture. 
ESTIMATION OF THE NITRIC ACID. GRAVIMETRIC ESTIMATION. 
107 
MAGNESIUM SULPHATE (MgS04.7H20). 
To a weighed quantity of the salt dissolved in water, add so- 
lution of ammonium chloride in slight excess, then water of 
ammonia until after stirring the mixture smells distinctly of it. 
If the ammonium hydrate produce a precipitate, it is evidence 
that sufficient ammonium chloride has not been added. When 
the solution remains clear after the addition of the hydrate, 
add sodium phosphate in excess, stir the mixture, without al- 
lowing the rod to touch the sides of the beaker, and set aside 
for twenty-four hours without warming. 
The clear liquid is poured on a filter and the precipitate 
then transferred to the same, using a little of the clear filtrate 
to wash the last portions on if necessary. 
The precipitate is then washed with a mixture of three 
parts water and one part ammonia water until a few drops of 
the washings acidified with nitric acid give no precipitate with 
silver nitrate. 
The precipitate is then dried and transferred to a crucible, 
the filter burned separately and the whole ignited strongly until 
the weight is constant. The residue consists of magnesium 
pyrophosphate—Mg2P207. 
ESTIMATION OF THE SULPHURIC ACID. 
The sulphuric acid is estimated in another portion of the 
salt by precipitation with barium chloride in the presence of 
hydrochloric acid. The resulting barium sulphate is collected, 
washed and ignited as under barium. 
ESTIMATION OF THE WATER. 
The water of crystallization may be estimated in magnesium 
sulphate by exposing a weighed quantity of the salt in a 
crucible to a low red heat until the weight becomes constant. 
CALCIUM CARBONATE (CaC03). 
Estimation of the Calcium. 
Dissolve a weighed quantity of the carbonate in dilute 
hydrochloric acid, avoiding excess of acid, add ammonium 
hydrate until slightly alkaline, then ammonium oxalate, and 108 
QUANTITATIVE ANALYSIS. 
set aside for twelve hours. Decant the clear supernatant 
liquid to a filter, wash once or twice by decantation, transfer 
the precipitate to the filter and wash with small portions of 
water, allowing each portion to completely run through before 
adding more. Dry and ignite the precipitate, using a strong 
Bunsen burner with a chimney, so as to bring the crucible to 
a white heat, until it ceases to lose weight; if this cannot be 
effected with a Bunsen burner, use a blast lamp. The residue 
is weighed as calcium oxide, CaO. 
ESTIMATION OF THE CARBONIC ACID. 
This in some cases is done by ignition, but the method 
adapted to all cases is to treat the carbonate with an acid, as 
HC1 or H2S04, and estimate the loss of C02. This is best 
accomplished by using a carbonic acid apparatus, that known 
as Schroetter’s is probably the best. 
An apparatus may be constructed, however, by taking a 
small, wide-mouth flask, and adapting to it a small tube con- 
taining calcium chloride. A weighed quantity of the material 
is then placed in the flask with some water, a test tube con- 
taining sulphuric acid is put in so as to retain the acid until 
the cork with the drying tube is fixed in place, and the whole 
apparatus weighed. 
The flask is then turned over so as to allow some of the 
acid to flow out, and this is repeated, allowing only a small 
quantity to run out at a time, until the carbonate has been 
completely decomposed. The flask is then heated, to drive 
out all the C02, and after cooling the apparatus is weighed. 
The loss indicates the amount of C02 in the material used. 
FERROUS SULPHATE (FeS04.7H20). 
Iron is usually estimated gravimetrically in the ferric state ; 
for this purpose the above salt when pure will answer the 
purpose. For the estimation of ferrous iron see the next 
section. 
A weighed quantity of the salt is dissolved in water, a few 
drops of sulphuric acid added and the mixture heated to the 
boiling point; nitric acid is then dropped in until oxidation GRAVIMETRIC ESTIMATION. 
to ferric iron is effected, and then water of ammonia is added 
until the solution smells strongly of it after stirring. The 
precipitate of ferric hydrate is allowed to settle and the clear 
supernatant liquid poured on a filter, more water is added to 
the precipitate in the beaker and after pouring the clear liquid 
on the filter the precipitate is transferred to the same and 
washed thoroughly with hot water, until the washings cease 
to have an alkaline reaction. The precipitate is dried, ignited 
and weighed, the result being ferric oxide, Fe203. 110 
QUANTITATIVE ANALYSIS. 
SECTION II. 
V.O LUMETRIC ESTIMATION. 
Volumetric analysis is the process of determining the 
amount of a substance by the use of a suitable reagent in 
solution of a known strength. The requisite apparatus con- 
sists of a liter flask and a 1000 cubic centimeter cylinder, 
in which to prepare the solutions. A 100 c. c. burette (Geissler’s 
is the most desirable) with a suitable holder. A few pipettes, 
assorted sizes, some of which are graduated to -fo c. c. The 
simplest process is that by neutralization, where an alkali is 
determined by neutralizing it with standard solution of oxalic 
acid. In such a case it is necessary to have a means of know- 
ing when the alkali is exactly neutralized. Such a substance 
is known as an indicator. Solution of litmus maybe used, but 
a solution of phenolphtalein gives most satisfactory results ; 
it is colorless in neutral or acid and bright red in alkaline 
solution. It is made by dissolving 1 part phenolphtalein in 
25 parts alcohol and adding sufficient water to make 100 
parts. 
A normal solution contains in every liter the molecular 
weight of the compound in grams. This is sometimes 
made semi-normal, when the compound, an acid, for instance, 
is dibasic. 
A deci-normal solution is one-tenth the strength of the 
normal one. While in some of the following examples the 
normal solution is used, that of deci-normal strength is recom- 
mended for the use of students. 
VOLUMETRIC SOLUTION OF OXALIC ACID. 
(HaC204.2H20‘= 126). 
The semi-normal solution is used, containing 6j grams in 1 liter. 
Take of oxalic acid, carefully purified by crystallization, 
63 grams, transfer it to a liter flask, add about 800 c. c. of VOLUMETRIC ESTIMATION. 
111 
water, agitate until dissolved and bring the measure, with 
water, to 1000 c. c. 
ESTIMATION OF POTASSIUM HYDRATE. 
Dissolve a convenient quantity of potassium hydrate, for 
instance .500 gram, in a beaker containing 100 c. c. of water; 
add three or four drops of the phenolphtalein solution; set 
the beaker on a piece of white paper and bring over it a 
burette containing some of the oxalic acid solution. Allow 
the acid solution to run in slowly, constantly stirring with a 
glass rod, until the red color just ceases to be visible. The 
proof of this is that one drop of a similar alkali solution will 
restore the red color. If the potassium hydrate were pure, it 
would be found that 8.94 c. c. were used, but if we suppose 
that 8.50 c.c. were necessary, we have every c. c. of the oxalic 
solution = .056 gram of KOH, therefore 8.50 X .056 = .476 
gram—the amount of KOH in the .500 gram = 95.20 per 
cent. 
One c. c. containing 0.063 gram of oxalic acid is the equiva- 
lent of— 
Gram. 
Ammonia, NH3, 0.0170 
Ammonium Carbonate, NH4HC03NH4NH2C02, 0.0523 
Potassium Acetate, KC2H302 (alter ignition), 0.0980 
Potassium Bicarbonate, KHCOg, 0.1000 
Potassium Carbonate (dry), K2C03, 0.0690 
Potassium Citrate, K3C6H307.H20 (after ignition), 0.1080 
Potassium Hydrate, KOH, 0.0560 
Potassium Permanganate, K2Mn2Os, 0.0314 
Potassium Sodium Tartrate, KNaC4tI406.4H20 (after ignition), 0.1410 
Potassium Tartrate, (K2C4H406)2.H20 (after ignition), . . . 0.1175 
Lead Acetate, Pb(C2H302)2.3H20, 0.1892 
Lead Subacetate, Pb20(C2H302)2, 0.1367 
Sodium Bicarbonate, NaHC03, 0.0840 
Sodium Borate, Na2B407.ioH20, 0.1910 
Sodium Carbonate, crystallized, Na2C03.ioH20, 0.1430 
Sodium Carbonate, anhydrous, Na2C03, 0.0530 
Sodium Hydrate, NaOH, 0.0400 
When carbonates are estimated it is necessary to boil, 
toward the end of the reaction, after each addition of the 
acid, in order to drive off the C02, which affects the indica- 
tor. 112 
QUANTITATIVE ANALYSIS. 
VOLUMETRIC SOLUTION OF SODIUM HYDRATE. 
(NaOH = 40). 
The normal solution is used, containing 40 grams in 1 liter. 
As sodium hydrate is never, or very rarely absolutely pure, 
it is necessary to standardize this solution. For this take 
something more than the theoretical amount (about 50 grams) 
and dissolve in a liter of water. Place 100 c. c. of the standard 
oxalic acid solution in a beaker, and, having added the indi- 
cator, bring it under a burette containing some of the soda 
solution, and note the number of c. c. necessary to exactly 
neutralize the acid solution. Take of the alkaline solution 
ten times the number of c. c. necessary to neutralize the 100 
c. c. of the acid solution, and add sufficient water to bring the 
measure to 1000 c. c. For instance, if it required 95 c. c. of 
the alkaline solution to neutralize the 100 c.c. of the oxalic 
acid solution, then 95 c.c. X 10 = 950 c.c., which amount 
diluted to 1000 c. c. would make it exactly equal in strength 
to the oxalic solution. 
One c. c. containing 0.040 of sodium hydrate NaOH, is the 
equivalent of— 
Gram. 
Acetic Acid, absolute, HC2H302, 0.0600 
Citric Acid, crystallized, H3CfiH507.H20, 0.0700 
Hydrobromic Acid, absolute, HBr, 0.0808 
Hydrochloric Acid, absolute, HC1, 0.0364 
Hydriodic Acid, absolute, HI, 0.1276 
Lactic Acid, absolute, HC3H503, 0.0900 
Nitric acid, absolute, HNOs, 0.0630 
Oxalic Acid, crystallized, H2C204.2H20, 0.0630 
Sulphuric Acid, absolute, H2S04 0.0490 
Tartaric Acid, crystallized, H2C4H4Oe, 0.0750 
VOLUMETRIC SOLUTION OF POTASSIUM BI- 
CHROMATE (K2Cr207 = 294.8). 
A viginti-normal solution is used, containing 14-.74- grams to 
1 liter. 
Place the necessary amount of the salt (14.74 grams) in a 
liter flask, add about 800 c. c. of water, agitate until dissolved, 
and bring the measure with water to 1000 c. c. This solution, 
when acidified with H2S04, is used in the estimation of iron in VOLUMETRIC ESTIMATION. 
113 
the ferrous condition. The end of the reaction is determined 
by taking out a drop of the iron solution, and testing on a 
white porcelain surface with a drop of potassium ferricyanide 
solution ; when this ceases to give a blue color the reaction is 
at an end. 
One c. c. containing 0.01474 gram of potassium bichromate, 
K2Cr207, is the equivalent of— 
Gram. 
Iron in ferrous condition, Fe, 0.01677 
Ferrous Carbonate, FeC03 0.03477 
Ferrous Sulphate, FeS04.7H20, 0.08337 
Ferrous Sulphate (dry), FeS04.HzO, 0.05097 
VOLUMETRIC SOLUTION OF IODINE (I = 126.6). 
A deci-normal solution is used, containing 12.66 grams in 1 liter. 
Weigh the necessary amount of iodine (12.66 grams) in a 
stoppered tube, to prevent loss, as well as the corrosive action 
of the fumes on the balance. Also weigh 18 grams of potas- 
sium iodide and place in a liter flask with the iodine. Add 
about 100 c. c. of water. The iodine dissolves more readily 
in this strength of potassium iodide solution, besides it admits 
of more thorough agitation. When the solution is complete 
add water until the liquid measures 1000 c. c. Starch solution 
is used as an indicator in the determinations with the iodine 
solution. 
One c. c. containing 0.01266 iodine is the equivalent of— 
Gram. 
Arsenious Oxide, As2Os, 0.00494 
Potassium Sulphite, crystallized, K2S03.2H20, 0.0097 
Sodium Bisulphite, NaHS03, 0.0052 
Sodium Hyposulphite,crystallized, Na2S203.5HaO, 0.0248 
Sodium Sulphite, crystallized, Na2S03.7H20, 0.0126 
Sulphurous Oxide, S02, 0.0032 
VOLUMETRIC SOLUTION OF SODIUM HYPO- 
SULPHITE (Na2S203.5H20 = 248). 
A deci-normal solution is used, containing 24.8 grams in 1 liter. 
Sodium hyposulphite cannot be prepared sufficiently pure 
to be relied on, consequently this solution must be standard- 
ized, therefore more than the theoretical amount is taken. 114 
QUANTITATIVE ANALYSIS. 
The U. S. Pharmacopoeia recommends 32 grams. Dissolve 
this amount in 1000 c. c. of water; place 100 c. c. of the stand- 
ard solution of iodine in a beaker and run in the hyposulphite 
solution until the color of iodine nearly disappears, then add 
a small quantity of starch solution and continue until the blue 
color is discharged. Multiply the number of cubic centi- 
meters of the hyposulphite solution used by 10, and to that 
amount add sufficient water to bring the measure to 1000 c. c. 
The substances estimated by this solution either contain free 
iodine, or develop it on the addition of potassium iodide, so 
that starch solution may be used as an indicator. 
One c. c. containing 0.0248 gram of sodium hyposulphite is 
the equivalent of— 
Gram. 
Bromine, Br, 0.00798 
Chlorine, Cl 0.00354 
Iodine, I, 0.01266 
VOLUMETRIC SOLUTION OF SILVER NITRATE 
(AgN03 = 169.7). 
A deci-normal solution is used, containing 16.9J grams in 1 liter. 
As silver nitrate can be obtained or prepared perfectly pure, 
the necessary amount (16.97 grains) is dissolved in sufficient 
distilled water to make 1000 c. c. 
In testing one of the following compounds, ammonium 
chloride for instance, a weighed amount is taken, dissolved 
in water, and a few drops of potassium bichromate solution 
added. The silver nitrate solution is then run in until a red 
precipitate remains permanently. The silver combines with 
the chlorine until the latter is all used, when it forms with 
the chromic acid red silver chromate, so that its formation 
indicates the end of the reaction. 
One c. c. containing 0.01697 gram of silver nitrate is the 
equivalent of— 
Gram. 
Ammonium Bromide, NH4Br, 0.00978 
Ammonium Chloride, NH4C1, 0.00534 
Ferrous Bromide, FeBr2, 0.01077 
Ferrous Iodide, Fel2, . . • 0.01545 
Hydrocyanic Acid, absolute HCN, 0.00270 VOLUMETRIC ESTIMATION. 
115 
Gram. 
Hydriodic Acid, HI, 0.01276 
Potassium Bromide, KBr, 0.01188 
Potassium Chloride, KC1, 0.00744 
Potassium Cyanide, KCN, 0.01300 
Potassium Iodide, KI, 0.01656 
Sodium Bromide, NaBr, 0.01028 
Sodium Chloride, NaCl, 0.00584 
ELEMENTS. 
SYMBOLS. 
ATOMIC WEIGHTS. 
Aluminium 
A1 
27 
Antimony 
Sb 
120 
Arsenic 
As 
74-9 
Barium 
Ba 
136-8 
Beryllium 
Be 
9-i 
Bismuth 
Bi 
207.5 
Boron 
B 
11 
Bromine 
Br . . . 
79-8 
Cadmium 
Cd 
hi.8 
Caesium 
Cs 
132-6 
Calcium 
Ca 
40 
Carbon 
C 
12 
Cerium 
Ce 
140.2 
Chlorine 
Cl 
35-4 
Chromium 
52 
Cobalt 
Co 
58-9 
Copper 
Cu 
63.2 
Didymium 
Di 
147 
Erbium 
E . . . . 
165.9 
Fluorine 
F 
19 
Gallium 
G 
68.8 
Germanium 
Ge 
72-3 
Gold 
Au 
196.2 
Hydrogen 
Ii 
1 
Indium 
In 
II3-4 
Iodine 
I 
126.6 
Iridium 
Ir 
192-7 
Iron 
Fe 
55-9 
Lanthanum 
La 
138-5 
Lead 
Pb 
206.5 
Lithium 
Li 
7 
Magnesium 
Mg 
24 
Manganese 
Mn 
54 
Mercury 
Hg 
199-7 
Molybdenum . . . . 
Mo 
95-5 
Nickel 
Ni 
58 
Niobium 
Nb 
94 
TABLE OF ELEMENTS. 116 
QUANTITATIVE ANALYSIS. 
ELEMENTS. 
SYMBOLS. 
ATOMIC WEIGHTS. 
Nitrogen 
N 
14 
Norwegium .... 
Ng 
214 
Osmium 
Os 
198.5 
Oxygen 
O 
16 
Palladium 
Pd 
105.7 
Phosphorus .... 
P 
31 
Platinum 
Pt 
194 
Potassium 
K 
39 
Rhodium 
Rh 
Rubidium 
Rb 
85.3 
Ruthenium .... 
Ru 
104.2 
Samarium 
Sm 
150 
Scandium . ... 
Sc 
44 
Selenium 
Se 
78.8 
Silicon 
Si 
28.2 
Silver 
107.7 
Sodium 
Na 
23 
Strontium 
Sr 
87.4 
Sulphur 
S 
32 
Tantalum 
Ta 
182.1 
Tellurium 
Te 
128 
Thallium 
T1 
203.7 
Thorium 
Th 
233.4 
Tin 
Sn . ... 
”7-7 
Titanium ..... 
Ti 
48 
Tungsten 
W 
183.6 
Uranium 
U 
238.5 
Vanadium 
V 
5i-3 
Ytterbium 
Yb 
172.7 
Yttrium 
Y 
89.8 
Zinc 
Zn 
64.9 
Zirconium 
Zr 
84.4 
TABLE OF ELEMENTS.—Continued. INDEX. 
PAGE 
Acetates, detection of. 86, 88 
reactions of. 82 
Acetic acid 82 
Acid, acetic 82 
benzoic 85 
boric 78 
carbolic 84 
carbonic 79 
estimation of. 108 
chloric 73 
citric 84 
gallic.... 85 
hydriodic 71 
hydrobromic 70 
hydrochloric 21, 70 
hydrocyanic 72 
hydroferricyanic 79 
hydroferrocyanic 79 
hydrofluoric 71 
hydrosulphuric 74 
hypochlorous 73 
hypophosphorous 77 
hyposulphurous 76 
lactic 82 
malic 83 
metaphosphoric 78 
nitric 76 
estimation of. 106 
oleic 82 
orthophosphoric 77 
oxalic 83 
phosphoric 77 
pyrophosphoric 77 
salicylic 85 
silicic 79 
stearic 82 
succinic 83 
sulphuric 75 
estimation of. 106 
sulphurous 75 
tannic 85 
tartaric 83 
thiosulphuric 76 
valerianic 82 
Acids, analytical reactions of 70 
detection of 70-93 
Acidum hydrochloricum 21 
Alcohol, reactions and tests of. 95 
amylic, detection of 95 
Alkaloids, reactions and tests of. 97 
Aluminium, analytical reactions of. 49 
hydrate 31 
Ammonia, preparation of. 25 
properties of 25 
Ammonium, analytical reactions of. 36 
nitrate 29 
oxalate 29 
Analysis of Acids— 
Group 1 72 
“ II 80 
“ III 86 
Bases— 
Group 1 37 
“ II 40 
Analysis of Bases— page 
Group III 45 
“ IV 50 
“ V 57 
“ VI 62 
“ VII 66 
Analysis of soluble salts 87 
insoluble salts 91 
Antimony, analytical reactions of. 55 
Marsh’s test for 55 
Reinsch’s test for 55 
Aqua chlori 19 
Arsenates, analytical reactions of 53 
detection of. 90 
Arsenic, analytical reactions of. 52 
Fleitman’s test for 54 
Marsh’s test for 53 
Reinsch’s test for 54 
Arsenious oxide 52 
Atomic weights, list of. 115 
Atropine 99 
Balance, rules for using 103 
Barium, analytical reactions of. 38 
chloride, estimation of. 104 
Benzoic acid 85 
Bismuth, analytical reactions of 60 
Bleaching 20 
Boric acid, analytical reactions of. 78 
Bromides, detection of 70, 90 
Bromine, preparation of. 22 
properties of 22 
Cadmium, analytical reactions of. 61 
Caffeine 99 
Calcium, analytical reactions of. 39 
carbonate, estimation of 107 
phosphate, preparation of.. 29 
Cane sugar 97 
Carbolic acid 84 
Carbon dioxide, preparation of 26 
properties of 27 
Carbonic acid, analytical reactions of 79 
estimation of 108 
Cerium, analytical reactions of 48 
Chloral hydrate, reactions and tests of. 95 
Chloric acid 73 
Chlorides, detection of. 72, 90 
Chlorine, estimation of. 105 
preparation of. 19 
water 19 
Chloroform, reactions and tests of. 94 
Chromium, analytical reactions of. 49 
Cinchonidine, reactions and tests of. 98 
Cinchonine, reactions and tests of. 99 
Citric acid, analytical reactions of. 84 
detection of 86, 89 
Cobalt, analytical reactions of. 43 
Copper, analytical reactions of. 61 
estimation of 105 
sulphate, estimation of. 105 
preparation of. 32 
Cyanides, detection of. 72, 90 118 
INDEX. 
PAGE 
Deci-normal solution no 
Dextrose 96 
Drying 104 
Elements, table of. 115 
Ferric hydrate 32 
sulphate 31 
Ferrous sulphate 31, 108 
Filtration 103 
Gallic acid, analytical reactions of. 85 
detection of. 88 
Glauber salt 22 
Glucose, reactions and tests of. 96 
Glycerin, reactions and tests of 96 
Gold, analytical reactions of. 56 
Grape sugar 96 
Gravimetric estimation 103 
Group I, Acids 70 
“ II, “ 73 
“ III, “ 82 
“ I, Bases 35 
“ II, “ 38 
“ III, “ 42 
“ IV, “ 47 
“ V, « 52 
“ VI, « 59 
“ VII, " 64 
Group reagents 69 
Hydrates, analytical reactions of 74 
Hydriodic acid, analytical reactions of.... 71 
Hydrobromic acid, analytical reactions of 70 
Hydrochloric acid, analytical reactions of 70 
preparation of 21 
properties of. 21 
Hydrocyanic acid, analytical reactions of 72 
Hydroferricyanic acid 79 
Hydroferrocyanic acid 79 
Hydrofluoric acid 71 
Hydrogen 17 
preparation of 17 
properties of 18 
arsenide 53 
Hydrosulphuric acid 74 
Hypochlorous acid 73 
Hypophosphorous acid 77 
Hyposulphurous acid 76 
Ignition 104 
Indicator no 
Iodides, detection of 72, 90 
Iodine, preparation of 22 
properties of 22 
Iodoform, reactions and tests of 94 
Iron, analytical reactions of. 47 
estimation of. 108 
Lactic acid 82 
Lead, acetate of 32 
analytical reactions of 65 
Lithium, analytical reactions of 36 
Magnesium, analytical reactions of 39 
carbonate 30 
estimation of. 107 
oxide 32 
sulphate 30 
Malic acid 83 
Manganese, analytical reactions of. 42 
Marsh’s test for arsenic 53 
PAGE 
Mercuric compounds 59 
analytical reactions of 59 
Mercurous compounds 65 
analytical reactions 65 
Metaphosphoric acid 78 
Morphine, reactions and tests of... 97 
Neutralization no 
Nickel, analytical reactions of. 44 
Nitrates, detection of. 90 
Nitric acid, analytical reactions of 76 
preparation of. 26 
properties of. 26 
Nitrogen, preparation of. 24 
properties of. 25 
Normal solution no 
Oleic acid 82 
Oxidation 20 
Oxides 74 
Oxygen, preparation of. 23 
properties of. 23 
Phenolphtalein no 
Phosphates, detection of 90 
Phosphoric acid, analytical reaction of.... 77 
Platinum, analytical reactions of. 56 
Potassium, analytical reactions of. 35 
chloride, preparation of. 28 
estimation of 106 
hydrate, estimation of in 
and sodium tartrate 28 
Precautions on Group II, Bases 42 
“ IV, “ 52 
“ V, “ 59 
“ VI, " 64 
“ VII, “ 68 
Precipitation 103 
Purple of Cassius 56 
Quantitative analysis, preliminary direc- 
tions 103 
Quinine, reactions and tests of. 98 
Reinsch’s test for arsenic 54 
Rochelle salt 28 
Saccharose, reactions and tests of 97 
Salicin 100 
Salicylic acid 85 
Santonin, reactions and tests of. 100 
Seminormal solution no 
Silicic acid — 79 
Silver, analytical reactions of 64 
Sodium, analytical reactions of. 36 
sulphate 22 
Solubilities, chart of. 93 
Stannic compounds 56 
Stannous compounds 55 
Starch 97 
Stearic acid 82 
Strontium, analytical reactions of. 38 
Strychnine, reactions and tests of 98 
Succinic acid 83 
Sulphates, detection of 81, 89 
Sulphuretted hydrogen 74 
Sulphuric acid, analytical reactions of 75 
Sulphurous acid,analytical reactions of.... 75 
Sulphurous oxide, preparation of. 24 
properties of. 24 
Summary Group I, Acids 72 
“ II, “ 80 
“ III, “ 86 
properties of.. 119 
INDEX. 
PAGE 
Summary Group I, Bases, 37 
“ II, “ 40 
“ III, “ 45 
“ IV, “ 50 
“ V, “ 57 
“ VI, “ 62 
“VII, “ 66 
Symbols, table of. 115 
Tannates, detection of. 88 
Tannic acid, analytical reactions of. 85 
Tartaric acid, analytical reactions of. 83 
Tartrates, detection of 86, 89 
Thalleioquin 98 
Thiosulphuric acid 76 
Tin, analytical reactions of 55 
PAGE 
Valerianic acid 82 
Veratrine, reactions and tests of. 99 
Volumetric analysis no 
estimation of alkalies m 
solutions . no 
iodine 113 
oxalic acid no 
potassium bichromate 112 
silver nitrate 114 
sodium hydrate 112 
sodium hyposulphite 113 
Wash bottle 30 
Water 74 
Zinc, analytical reactions of 43 
sulphate, preparation of. 19