Robert W. Woodruff
Library
Kelly Miller Library
EMORY UNIVERSITY
Special Collections & Archives
QUANTITATIVE
CHEMICAL ANALYSIS
by
H. G. SGURLOGK, A. M„ M. D„
PROFESSOR OF CHEMISTRY, HOWARD UNIVERSITY
copyright, 1915
WASHINGTON, D. G.
Beresford, Pr., 605 F Street
1915
FOREWORD.
The author, in preparing this book for his pupils, has had in mind
the purpose of supplying an elementary course for those students who
wish to extend their knowledge of general chemistry beyond the
courses in Inorganic Chemistry and Qualitative Analysis, so as to
include Quantitative Analysis, and especially for those who wish to
prepare for further work, or to enter the Medical, Dental and Phar¬
maceutical Colleges.
The exercises contained herein are intended to give practice in the
common examples of determinations by gravimetric and volumetric
methods. The plan of the book is in accordance with the lecture
course which has been designed, and a great deal has been left for
consideration in the class room. The student will be constantly
referred to his text-books on Inorganic Chemistry and Qualitative
Analysis, and the larger text-books on Quantitative Analysis included
in the following list:
Olsen, "Quantitative Analysis,"
Fresenius, "Quantitative Analysis,"
Morse, "Quantitative Chemistry,"
Treadwell and Hall, "Analytical Chemistry, Vol. II,"
Clowes and Coleman, "Quantitative Analysis,"
Sutton, "Handbook of Volumetric Analysis."
Attention is directed to the importance of Quantitative Analysis to
the general subject of Chemistry and its large application in the in¬
dustrial world. It covers an extensive field, and a proper knowledge
of the subject can be gained only by large laboratory experience and
careful reading of the literature.
Cleanliness, order, promptness and accuracy in making notes, and
the recording of things just as they are found, even if they should
vary from expected results, are absolutely essential in any kind of
experimental work.
CONTENTS.
Section I. General Principles.
PAGE.
Chapter I. Introduction. Apparatus i
Chapter II. Stoichiometry 18
Section II. Gravimetric Analysis.
Chapter III. General Operations 25
Chapter IV. Gravimetric Exercises 37
Chapter V. Electrolytic Methods 48
Section III. Volumetric Analysis.
Chapter VI. General Principles 52
Chapter VII. Acidimetry and Alkalimetry 59
Chapter VIII. Oxidizing Reactions 65
Chapter IX. Iodimetry 70
Chapter X. Silver Nitrate and Potassium Sulphocyanate Titrations 73
Chapter XI. Colorimetric Methods 75
Section IV.
Chapter XII. Miscellaneous Analysis 78
Index
87
QUANTITATIVE CHEMICAL ANALYSIS.
SECTION I.
GENERAL PRINCIPLES.
Chapter I.
Introduction. Apparatus.
Quantitative Analysis comprises, in two general meth¬
ods,—gravimetric and volumetric,—the determination of
the quantity of one or more of the constituents of the sub¬
stance analyzed ; as, the amount of an element in a com¬
pound, of a metal in an ore, of a given substance in an
impure sample, of an acid or other substance in solution,
etc.
In the gravimetric method, a weighed portion of the
substance to be analyzed is put into solution and the
components to be determined are precipitated and isolated
in the form of some pure compound of known constitution.
The amount of the new compound is found by direct
weighing, and from this weight the substance sought is
found by calculation.
A volumetric determination is accomplished without the
isolation and weighing of an insoluble product. A known
amount of the substance to be analyzed, in solution, is
treated with a test solution which contains a known amount
of the test reagent. The volume of the test solution
required to complete a reaction with a given volume of the
solution analyzed furnishes the basis for calculation of the
quantity of the substance sought. The completion of the
reaction is shown by an indicator,—a substance added to
the solution tested, which by a color change, indicates t e
2
quantitative chemical analysis.
end-point; e. g., the neutralization of an alkaline solution
by an acid is indicated by the change of color of litmus,
phenolphthalein, etc.
Materials used in quantitative analysis must be weighed
and measured with as great a degree of accuracy as possible ;
hence, there are especially required (i) analytical balances,
sensitive to small differences of weight; (2) weights, ac¬
curately calibrated ; (3) measuring vessels, correctly gradu¬
ated.
THE BALANCE.
Construction and Use.—The analytical balance is illus¬
trated in Fig. 1. Upon examination of a balance the
student will have no difficulty in recognizing the fol¬
lowing parts, and should endeavor to early learn their
functions : beam, post, knife edges, stirrups, hangers, pans,
QUANTITATIVE CHEMICAL ANALYSIS.
3
graduated scale, pointer, rider carrier, beam release and pan
arrests. The balance is enclosed in a glass case to prevent
accumulation of dust and to avoid draughts of air while
weighing. By placing a drying agent within the case the
air may be kept reasonably free from excessive moisture
which might, if present, be condensed on substances while
being weighed.
Errors in analytical work cannot be properly charged
against the balances used, but rather against the improper
use of them and carelessness.
The balance is put in adjustment by an instructor and
should not be tampered with by the student until he is
sufficiently familiar with it to make slight adjustments, and
only then after an instructor has given permission to make
them.
Observe the following directions for using the balance
Sit or stand directly in front of the balance so as to avoid
errors of parallax while observing the movements of the
pointer over the graduated scale.
Dust off the pans with a camel's hair brush, but avoid
giving them a rotary motion.
Keep the balance closed while weighing.
Do not weigh a hot object.
Do not weigh chemicals directly on the scale pans.
Weigh them on a tared watch glass or in a weighing
bottle.
Keep beam and pans in arrest except during the actual
operation of weighing.
Avoid jarring the balances at any time. If they are
placed on a table take care not to lean on it while weighing,
and at no time use the table to rest substances upon while
pounding them.
WEIGHTS.
Construction and Use. — Analytical weights are grad¬
uated in denominations of the metric system, usually
stamped in grams and milligrams. The gram denomina-
4
QUANTITATIVE CHEMICAL ANALYSIS.
tions are made of brass, while the fractional ones (milli¬
grams) are made of platinum. Ordinary weights are
lacquered or nickel-plated ; higher grade weights are plated
with gold or platinum.
Only the highest grade of weights conform closely to
the masses stamped on them, and it is necessary with all
weights to ascertain by comparison with a standard the
amount of error which may be present. The errors should
be recorded and the weights which vary from the correct
mass should be marked with light prick-punch marks for
identification.
Weights of the lower grades show a considerably greater
variation from the correct mass, and every such set should
be calibrated before being used, at least one of them being
compared with a standard, and then the relation of all the
members of the set found. It will be understood from this
why it is necessary that the weights from one box should
not become mixed with those from another.
Handle all weights with pincets ; never with the fingers,
and familiarize yourself at once with the denominations,
and arrangement of the weights in your box.
The Rider.—The rider is a piece of platinum wire so
formed that it may be conveniently lifted by the rider-
carrier and placed at any position along the graduated
beam. The beam is usually divided into 5, 6, 10 or 12
large divisions, each of which is subdivided into 10 small
divisions. Riders whose weights correspond to the number
of large graduations on the beam will indicate at the
extreme end graduation their weight in milligrams, and the
value of each small division will be one-tenth milligram.
Usually the rider is used to obtain equilibrium when a
weight under 10 milligrams is required. In Fig. 2 is
illustrated the graduated beam, divided into 60 equal
divisions, with a six-milligram rider indicating a weight
of 3.7 milligrams.
Zero Point.—The zero point is the point on the graduated
quantitative chemical, analysis.
5
scale over which the pointer would finally come to rest
when the beam is swinging free, and should be very close
to the middle graduation. It seldom coincides exactly
Fig. 2.—Graduated end of beam showing use of rider.
with the middle graduation. It is essential to any weigh¬
ing operation that the zero point be known ; otherwise one
could not know when the balance is in equilibrium. The
first thing to do, then, before weighing, is to find the zero
point.
Exercise I.
Determination of the Zero Point of a Balance.
Release the pans and then the beam with a slow, steady
motion of the hand. If the beam does not start swinging
at once, cause it to do so by gently fanning with the hand
or a piece of paper. Close the door of the balance and
observe the amount of swing of the pointer on each side of
the middle graduation on the ivory scale. After the first
few swings record the readings on each side of the middle
graduation, making the last reading on the same side as
the first. Example,
Left
3.8 divisions
3.6
M "
3)10.8
3.6, mean swing to left.
Right
3.4 divisions
3. ° "
2)6.4
3.2, mean swing to right.
6
QUANTITATIVE CHEMICAL ANALYSIS.
^= .2, gives the zero point as .2 of a division to
the left of the middle graduation. It will often be found
necessary to estimate fractional parts of a scale division by
the eye. This may be facilitated by the use of a magnifying
glass before the ivory scale.
Sensibility.—By the sensibility of a balance is meant the
amount of deviation of the zero point, measured in scale
divisions, caused by an excess weight of one milligram
added to one of the pans. The amount of this deviation
will vary, generally, with varying loads, and should be found
for different masses up to the limit of capacity.
Knowledge of the sensibility of a balance facilitates
"weighing operations. The more sensitive the balance the
more nearly can the exact mass of an object be obtained.
For ordinary analytical purposes a balance which shows a
sensibility of .i milligram with full load is sufficiently
sensitive. Balances are constructed to carry a certain
maximum load which should never be exceeded, for in
such cases the beam is deflected and the weighing incorrect.
Exercise II.
Determination of the Sensibility of a Balance.
With the pans empty, add a weight of one milligram to
the right-hand pan, or place the rider at the one-milligram
division on the beam. Find the zero point. The number
of scale divisions of the displacement from normal zero is
the sensibility with no load. Then find the sensibility
for loads of 5, 10, 15, 20, 30, 40, 50, 60, 75 and 90 grams,
by placing weights of these denominations in each pan and
adding the excess weight of one milligram to the right-
hand pan. If it should be found that two weights of the
same denomination do not have the same mass, put the
lighter into the right-hand pan and bring the balance into
equilibrium with the rider. Record all findings in your
note book.
QUANTITATIVE CHEMICAL ANALYSIS.
7
Weighing.—An object is weighed when it is exactly
counterbalanced by a known mass. When a substance is
weighed in the air its true weight is not found because of
the difference due to the buoyancy of the air, hence it
weighs less in air than it would in a vacuum. Weighings
are reduced to vacuum values as follows : Divide the air
weight of the object by its specific gravity, and the air
weight of the weights by the specific gravity of the
metal of which they are made (brass), to find the cc. of air
displaced by them. Multiply the difference between the
cc. of air displaced by the object and by the weights by
the density of air at the temperature and pressure of the
determination. This will give the amount in grams
to be added to the air weight of the object.
Exercise III.
Determination of the Weight of an Object.
Place a suitable object, e. g., a small beaker, on the left-
hand pan, and on the right-hand pan place weights sup¬
posed to be about equal to it. Slowly release the balance,
pans first and then the beam, until it is seen which way the
pointer swings, and then as slowly arrest it again. Add
or remove weights as may be required until equilibrium is
nearly reached.
Use the rider for differences of weight under ten milli¬
grams, moving it along the beam to right or left as required.
In the same manner as under determination of the zero
point, determine the position at which the pointer would
now come to rest, by observing the amount of swing on
each side of the zero point already established, making, as
before, the last reading on the same side as the first.
As an illustration of how to arrive at the correct weight
of an object, suppose the weights in the pan amount to
16.520 grams, and that the rider indicates 3.7 milligrams ;
also, suppose that the point of equilibrium now is 6 divis-
8
QUANTITATIVE CHEMICAL ANALYSIS.
ions to the left, and that the sensibility of the balance,
previously ascertained, is 5 divisions for 1 milligram.
According to the supposition of the example the zero point
is .2 of a division to the left; therefore, the present devia¬
tion is 5.8 divisions, and since 1 milligram causes a dis¬
placement of 5 divisions it is evident that 5.8 divisions
represent a difference of 1.16 milligrams, and the weight is
too great by this amount. Then from 16.5237 gms. (16.520
+ .0037) there must be subtracted .00116 gm. giving
16.52254 gms., the air-weight of the object.
The proper way to count the weights on the pan is to
first record in the note book the weights missing from the
box and check them when replacing. The record may be
made as follows :
Gms. Mg.
10
5
500
020
0037 (rider)
16
5237
(Zero, .2, left; displacement, 6, left; correction, 5.8, =
— .00116 gm. Correct weight, 16.52254 grams.)
MEASURING VESSELS.
Measuring vessels consist of cylindrical or conical grad¬
uates, volumetric flasks, burettes, and pipettes, graduated
in cubic centimeters ; the latter three forms of vessels are
shown in Fig. 3.
Graduates.—The graduates ordinarily found in the lab¬
oratory are not intended to be used in making accurate
measurements. The conical form is not intended at all to
be used for such purposes. Cylindrical graduates are
sometimes graduated to contain accurate quantities. The
U. S. Bureau of Standards accepts for standardization
QUANTITATIVE CHEMICAL, ANALYSIS.
9
"only cylinders graduated to contain. The inside diameter
of cylinders must not be more than one-fifth the graduated
length."
A
T=f
JtO. /31f
Conta-ins
1000CC
ao°C
M*7
cc
wc
-&■
-33-
/ \
10 CC
2JfC
US'Sec.
\ J
Fig. 3.—Volumetric flasks, burette, transfer pipette and
measuring pipette.
(CC\
ZOt
-m
IO QUANTITATIVE CHEMICAL ANALYSIS.
Flasks.—Volumetric flasks are made to contain or to
deliver accurate quantities of liquids. If only one mark is
found around the neck without other indication it means,
usually, that the flask when filled to the mark will contain
the quantity of liquid engraved on the body of the vessel.
If there are two marks it will deliver the amount engraved
on the flask when filled to the upper mark.
All graduated vessels are calibrated at a given temperature
which is also engraved on them and is the temperature
at which liquids are to be measured for accurate results.
The temperature used by the U. S. Bureau of Standards
and in most laboratories in this country is 20° C.
Transfer Pipettes.—Transfer pipettes are made to deliver
the amount of liquid engraved on the body of the instrument.
They are used by sucking the liquid up into the instrument,
well above the mark; the upper end is closed with the
finger and the excess liquid allowed to run out by slightly
releasing the pressure of the finger so as to admit a small
stream of air. When the liquid has run down to the mark
the upper end of the pipette is again tightly closed with
the finger. The liquid is then run into the vessel to which
it is to be transferred, the lower end of the pipette being
held against the side of the receptacle and allowed to drain
for several seconds after the liquid seems to cease flowing.
There will always be a portion of the liquid remaining in
the pipette. Do not blow this out into the vessel; the
correct amount has been delivered when the instrument
has drained itself as far as it will by gravity.
Measuring Pipettes.—Measuring pipettes are used to
deliver any quantity up to their capacity in cubic centi¬
meters or fractional parts of a cubic centimeter.
Burettes.—The chief use of the burette is for the ad¬
dition of the test solution in titrations in volumetric
analysis; but it may be also used for the accurate meas¬
urement of liquids. For most purposes a burette having
a graduated capacity of 50 cc. is suitable, though instru-
quantitative; chemical analysis.
II
ments of 25 cc. are also much used. The cc. divisions are
divided into fifths and tenths, and by estimation may be
read to hundredths of a cc.
The tip of some burettes is a separate piece of glass
tubing drawn to a proper sized outlet and attached to the
main instrument by a short piece of rubber tubing kept
closed by a pinch-cock. Such instruments will not answer
for use with liquids which attack rubber, and when reagents
of this kind are used the rubber-tipped burette should be
replaced by a glass-stoppered one.
The burette is to be held in a vertical position in a burette
clamp, and filled well above the zero mark with the liquid
to be measured ; then the liquid is slowly and carefully run
out until its level at the lower part of- the meniscus coin¬
cides with the zero mark. Before beginning a measurement
from the burette, or a titration, the drop of liquid adhering
to the tip of the instrument should be removed with a piece
of filter paper or a glass rod. When transferring a measured
portion of liquid, the last adhering drop is removed by
bringing the wet surface of the vessel in contact with the
drop, or it is removed with a glass rod.
Method of Reading Burettes.—It is important in reading
the level of liquids in measuring vessels that the eye be
brought in a direct line of vision with the mark so as to
avoid the errors due to parallax. The following instruc¬
tions are issued by the U. S. Bureau of Standards : " In all
apparatus where the volume is limited by a meniscus the
reading or setting is made on the lowest point of the
meniscus. In order that the lowest point may be observed
it is necessary to place a shade of some dark material
immediately below the meniscus, which renders the profile
of the meniscus dark and clearly visible against a light
background. A convenient device for this purpose is a
collar-shaped section of thick black rubber tubing, cut open
at one side and of such a size as to clasp the tube firmly."
Unless measuring vessels have already been standard-
12
QUANTITATIVE CHEMICAL ANALYSIS.
ized they should be calibrated before use, and the various
pipettes, burettes and flasks tested for agreement with one
another. Exercises in calibration are given below.
It is absolutely essential that measuring vessels shall be
clean. Sometimes, in the absence of visible accumulations,
a vessel may be very dirty, as when drops of water adhere
to the surface of a burette.
Cleaning Solution.—An efficient cleaning solution for
use when there is plenty of time consists of a mixture of a
few crystals of potassium dichromate and concentrated
sulphuric acid. The vessel to be cleaned is filled with the
cleaning mixture and set aside for several hours, after which
the solution is returned to the bottle, and the vessel washed
thoroughly with distilled water. A more rapid cleansing
may be effected by washing with sodium hydroxide solution,
then with dilute acid, and, finally, with distilled water.
Vessels of suitable size may be dried by placing in the
oven. Sometimes it is convenient to rinse the vessel
with alcohol, and then with ether, after which it can
be made quite dry in a short time in the oven. Great
caution is always necessary when using ether near an open
flame. The drying of large or long pieces of apparatus,
such as a burette, may be assisted by passing through them
a blast of dry air.
Care of Stop-cocks.—Stop-cocks and glass stoppers should
be lubricated slightly with paraffin. When the vessel is to
be set aside for some time or be returned to the storeroom, it
should be thoroughly cleaned and the stop-cock properly
lubricated, or all lubrication removed and a piece of paper
placed between the stop-cock and its seat at one side with
one end projecting outward. This is to prevent the stick¬
ing of the stop-cock, with the probable loss of the instrument.
Units of Capacity.—The units of capacity employed by
the U. S. Bureau of Standards are (i) the liter, defined as
the volume occupied by a quantity of pure water at 40
Centigrade having a mass of 1 kilogram ; and (2) the cubic
QUANTITATIVE CHEMICAL ANALYSIS.
13
centimeter or milliliter, which is the one-thousandth part
of a liter. One milliliter is 1.000027 cubic centimeter,
hence the cubic centimeter is not exactly one-thousandth
part of a liter, but the difference is so small that in volu¬
metric analysis it makes no material difference.
Although the standards of capacity are defined at 40 C.,
760 mm. pressure, vessels are graduated at 20° C., which
is close to the average temperature of most laboratories.
Pure water at 20° C. has a density of 0.9982343 ; i. e., one
liter of pure water, measured at 20° C., would weigh in a
vacuum 998.2343 gms. Weighing is affected by the buoy¬
ancy of the air, and to find the true weight, reduction
must be made to weight in a vacuum.
One liter of air at 760 mm. pressure weighs 1.2 gms.
and exerts this amonnt of buoyancy upon a liter of water.
The density of brass is 8.4 ; therefore, a weight of 998.2343
gms. of brass has a volume of 998.2343 h- 8.4 = 118.8527
cc., which is also the volume of air displaced, having a
weight of 0.1426 gm.; 1.2—0.1426 = 1.0574 = loss of
weight of a liter of water when weighed in air at 20°. The
weight of a liter of water at 20° in air is 998.2343 — 1-0574
= 997.1769 gms. Calculate the weight of the fractional
parts of a liter from this factor.
Exercise IV.
Calibration of a Transfer Pipette.
Bring the water to be used to the temperature etched
upon the pipette. Measure accurately into a carefully-tared
weighing bottle the full contents of the pipette, touching
the tip against the side of the bottle to remove the last ad¬
hering drop, but do not blow into the bottle the portion of
water remaining within the tip. Stopper the bottle and
carefully take its weight. Subtract the weight of the
bottle alone from the weight of bottle and contents to find
the weight of the water delivered. Record this weight in
14
QUANTITATIVE CHEMICAL, ANALYSIS.
the note-book. By reference to the table of water volumes
calculate the volume of water delivered and record in the
note-book.
Exercise V.
Calibration of a Burette.
Fill the burette with distilled water at the required
temperature; run out successive portions of 5 cc. into a tared
flask whose capacity is as great as that of the burette.
Weigh after the addition of each successive portion of 5 cc.
and from the results calculate the correct readings and
record in the note-book. The first calibration should be
checked by a duplicate. It is important that in measuring
the water from the burette that the reading should not be
taken earlier than after three minutes, in order to allow
the film of liquid adhering to the walls of the instrument
to run down and indicate the correct height in the instru¬
ment.
Exercise VI.
Calibration of a Measuring Pipette.
Calibrate a measuring pipette in the same manner as a
burette.
Exercise VII.
Calibration of a Flask.
(a) By Weight.—Placed the counterpoised flask on the
left-hand pan of the balance and add to the right pan the
requisite number of grams, on the basis of 997.177 grams
for one liter. Fill the flask with an amount of distilled
water at 20° C. slightly less than the required amount.
Then, carefully, drop in water until equilibrium is reached.
If the flask is ungraduated a film of melted paraffin should
be smeared around the neck before beginning the operation,
and when the height of the liquid is noted a thin line is
drawn around the neck of the flask through the paraffin, and
QUANTITATIVE CHEMICAL ANALYSIS.
a little hydrofluoric acid placed in the groove thus made to
etch the line upon the glass. If the flask is already grad¬
uated it may be filled to the mark and weighed. From
the weight obtained correction, if necessary, is to be made
for the contents of the flask. Or, if the existing mark is
erroneous a new mark may be made and so indicated.
{b) By Morse-Blalock Bulbs.—Where several vessels are
to be calibrated to the same capacity it may conveniently
be done with the Morse-Blalock bulbs, figured in Fig. 4.
U
Fig 4.—Morse-Blalock bulbs for calibrating flasks, burettes
and pipettes.
In the figure the large bulb has a capacity of 500 cc. from
the upper mark to the first mark (zero) on the graduated
stem. The stem is graduated in millimeter divisions. The
n
U
j5 quantitative chemical analysis.
combined capacity of bulb and stem at o° C. is not less
than 500 cc. A delivery tube with stop cock is attached
by rubber tubing to the lower end.
To calibrate the instrument, it is filled with distilled
water (including the delivery tube) to the upper mark.
The water is then run out into a tared flask to the first
mark on the stem, and weighed. Then run the water in
the stem into a weighing bottle to the last mark. Weigh
the bottle. The temperature of the water must be taken
and the weight divided by the weight of 1 cc. of water at
the observed temperature, making correction for air dis¬
placement. In this way the capacity of the bulb and the
capacity of the stem at the observed temperature are found.
Divide the number of cc. in the stem by the total number
of graduations on the stem to find the value in cc. of each
division, and select the mark on the stem to which water
must be run out in order, when added to the bulb, to measure
just 500 cc. at the temperature of the calibration.
If practicable bring the water to 20° C. Then the fac¬
tor for dividing the weight of the water will be .997177
(.9982343 — .001574). The apparatus is then ready for
calibrating any number of flasks at that temperature.
Procedure.—Smear the necks of the flasks to to be cali¬
brated with a thin coating of melted paraffin. Fill the
standard bulb with distilled water, as before, at the temper¬
ature of the experiment. Run the water into the flask
down to the mark on the stem necessary to deliver 500 cc.
Then make a thin groove through the paraffin down to the
glass at the lower edge of the meniscus and etch it in with
hydrofluoric acid. Also etch the capacity of the flask, the
temperature of the calibration, and your initials on the
body of the flask.
These directions call for a 500 cc. flask, but flasks of any
size may be calibrated in like manner by using the proper
standard bulbs.
QUANTITATIVE CHEMICAL ANALYSIS.
17
Questions.
1. What is the chief difference between gravimetric and
volumetric methods ?
2. What special equipment is required for quantitative
analysis ?
3. What is the zero-point of a balance, and how is it
found ?
4. What is meant by the sensibility of a balance, and
how is it determined ?
5. Does the sensibility of a balance vary with the load ?
6. State how weights which do not conform exactly to
the masses stamped on them may be used in quantitative
analysis.
7. A 10-milligram rider is used on a beam having 12
large divisions, each divided into 10 small divisions. What
is the value of the rider for each small division ?
8. Advance a reason why objects should not be weighed
hot.
9. Why should weights not be handled with the fingers ?
10. What does the term " weight" mean when used for
the findings with the analytical balance ?
11. Describe how a transfer pipette is used for measuring
liquids.
12. Describe how a burette is used and the manner of
reading a burette.
13. How may glass vessels be made clean?
14. What gives evidence that a clean-appearing vessel is
dirty ?
15. Define the units of capacity used in graduating
vessels.
16. At what standard temperature are vessels graduated
in this country ?
17. Is the weight of an object as ordinarily taken the
true weight ?
18. Discuss the factors bearing on the correct answer
to 17.
2
18 QUANTITATIVE CHEMICAL ANALYSIS.
19. Describe the calibration of a burette.
20. Describe the calibration of a flask by the use of
Morse-Blalock bulbs.
Chapter II.
STOICHIOMETRY.
The calculations of analytical chemistry depend upon
the quantities of substances used and found and the atomic
or molecular proportions entering into their reactions.
Atomic and molecular weights are made use of here in the
same manner as in the problems worked in a first course in
chemistry. Usually a simple proportional statement is all
that is required for most calculations ; but in order to select
the proper quantities for the proportional statement it is
necessary to understand the reactions which take place and
to be able to write the equations.
BALANCING EQUATIONS.
Metathetical Reactions.—The reactions most commonly
met with in a quantitative analysis are double decomposi¬
tions between (1) two salts ; (2) a salt and an acid ; (3) an
acid and a base ; resulting respectively, in an exchange of
radicals, the formation of a salt of the acid used, and neu¬
tralization with the formation of a salt.
In writing equations it is essential that the equivalent
quantities of the reacting substances be set down. Silver
nitrate (AgN03) and sodium chloride (NaCl) form silver
chloride (AgCl) and sodium nitrate (NaNOs). Inspection
of the formulas of these substances shows that both sodium
chloride and the silver chloride formed, each contain one
equivalent of chlorine, are univalent combinations, and
the reaction, a simple exchange of radicals ;
AgN03 + NaCl = AgCl + NaNOa.
quantitative; chemical analysis. 19
Silver also forms silver chloride with barium chloride
(BaCl2). Barium chloride contains two atomic equivalents
of chlorine, and since the product of the reaction with silver
is silver chloride, which contains but one atomic quantity
of chlorine, two molecular quantities of silver nitrate are
required to furnish the silver necessary for the complete
reaction ;
2AgN03 + BaCl2 = 2AgCl + Ba(N03)2.
It is evident that in attempting to balance an equation
one must have knowledge of the products of the reaction,
and a review of the following illustrations will be of assist¬
ance :
1. Two salts exchange radicals, as in examples above.
2. Strong acids decompose the salts of weaker acids with
the formation of a salt of the stronger acid ;
CaC03 + H2S04 = CaS04 + H20 + C02.
3. An acid and a base form a salt and water ;
HC1 + NaOH = NaCl + H20.
4. A metalliooxide and a non-metallic oxide form a salt \
BaO + C02 = BaC03.
5. A metallic oxide and an acid form a salt and water ;
BaO + H2S04 = BaS04 + H20.
6. Carbonates yield carbon dioxide, as in 2 ; when roasted
they give also a metallic oxide ;
BaCOs = BaO + C02.
Other cases of the formation of oxides by heating are
given under Ignition.
Oxidizing Reactions.—When a substance gains in positive
valences it is said to be oxidized ; the opposite is reduction ;
20
QUANTITATIVE CHEMICAL ANALYSIS.
thus, when ferrous chloride (FeCl2) is oxidized to ferric
chloride (FeCl3), Fe gains a positive valence. In the light
of the theory of electrolytic dissociation oxidation and re¬
duction are looked upon as the assumption of positive and
negative charges, respectively. Since oxidizing reactions
involve a change of valence they are not so easy to balance
as the metathetical equations just considered.
A common example of oxidation is that of iron by pot¬
assium dichromate in the presence of hydrochloric acid.
First, consider the reaction of potassium dichromate
(K20.2Cr03) when reduced, thereby furnishing the oxidizing
charge; and, second, the oxidation of ferrous chloride,
finally combining the two equations,
K2Cr207 + 8HC1 = 2KC1 + 2CrCl3 + 4H20 + 3O. (1)
6FeCl2 + 6HC1 + 30 = 6FeCl3 + 3HA (a)
6FeCl2+K2Cr2Or+i4HCl=6FeCl3+2KCl+2CrCl3 +
7H20. (3)
The molecular equivalent of potassium dichromate may
be found from equation 1, in which it is seen that three
atoms of oxygen are liberated : K2Cr207 = 30 = 6H ;
therefore, one-sixth the formula-weight of potassium di¬
chromate is equivalent to one atomic weight of hydrogen.
Iron, by oxidation, gains one valence, and K2Cr207 = 30 =
6Fe.
Oxidation by nitric acid also furnishes an illustrative
example. When heated, concentrated nitric acid decom¬
poses as shown in equation (4).
2HNO3 = H20 -f- 2NO2 -f~ O. (4)
2FeCl2 + 2HCI + O = 2FeCl3 + H20. (5)
2FeCl2 -f 2HN03 + 2HCI = 2FeCl3 + 2H20 + 2N02. (6)
Two formula-weights of nitric acid as a unit yield one
atom of oxygen ; hence, 2HN03 = O = 2H = 2Fe.
QUANTITATIVE CHEMICAL ANALYSIS.
21
Calculations based on equations.—Write the equation for
the reaction considered and set down under each member
of the equation the equivalent quantities involved, thus :
H2S04 + BaCl2 = BaS04 + 2HCI
98.09 208.27 233.43 72.924.
1. The amount of a substance required to combine with
a given amount of another substance.
How much barium chloride required to combine with 18
grams of sulphuric acid ?
98.09 : 208.27 : : 18 : x = 38.21 gins.
2. The amount of a substance required to form a given
amount of another substance.
How much sulphuric acid required to form 115 milligrams
of barium sulphate ?
233-43 : 98.09 : : 115 : x = 48.324 mgs.
3. The amount of a substance formed by a given amount
of another substance.
How much barium sulphate is formed by 20 grams of
barium chloride ?
208.27 : 233.43 :: 20 : x = 22.949 gms.
4. The amount of a substance equivalent to a given
amount of another substance.
Twenty cc. of a sulphuric acid solution by analysis give
175 milligrams of barium sulphate. How much sulphuric
acid in a liter of the solution ?
233-43 : 98-°9 T75 : x =73-54 tng.=H2S04 in 20 cc., and
73.54 Xi.ooo cc. (1 liter) ,
1 = 3>677 or 3-677 gms-,
amount of H2S04 in one liter.
22
QUANTITATIVE CHEMICAL ANALYSIS.
Analytical Factors.—An analytical factor is found by
dividing the equivalent weight of the substance sought by the
•equivalent weight of the substance found ; thus, the factor
for sulphuric acid from barium sulphate: 98.09 -h 233.43 =
-.42021. To calculate the quantity of sulphuric acid found
by precipitation as barium sulphate, multiply the amount
of barium sulphate obtained by the factor .42021. (See
also volumetric factors.)
Use of Logarithms.—Arithmetical calculations are short¬
ened by the use of logarithms, and errors are usually de¬
tected more easily. It is advisable that the student acquire
facility in using logarithmic tables.
The slide rule, based on logarithms, is a mechanical
means of solving problems and may be used to advantage
in many instances.
Stating Results.—If the sample analyzed is a liquid, the
results may be stated in grams per liter, or if the specific
gravity of the liquid is known the percentage strength may
be given. When solids are examined the report is usually
made in f>er cents.
The ratio of the amount of the calculated substance to
the amount of the sample used, the latter taken as 100,
the base of percentage, will give the per cent of the sub¬
stance as found by analysis.
Substance calculated x , . ,
—_ . . . :: — : or, put into a proportional
Original sample 100 r r r
statement with x as the last member, A : B :: 100 : x (A =
sample, B = calculated substance).
Chemical Tables.—As a result of very careful analyses
tables have been prepared for the use of the chemist which
give important information. Among them will be found
lists of the elements and the compounds and radicals com¬
monly encountered in analytical practice, together with
their atomic and molecular weights and their logarithms ;
also, lists of analytical factors and their logarithms.
Tables which show the per cent of substances in solution
QUANTITATIVE CHEMICAL, ANALYSIS. 23
at different specific gravities are exceedingly useful. For
example, suppose it were desired to make a solution of very
nearly 10 per cent of hydrochloric acid from the reagent in
the stock bottle. Referring to a table of the properties of
hydrochloric acid it is found that 10 per cent hydrochloric
acid has a specific gravity of 1.050. It would be a simple
matter, using a hydrometer, to add water to a portion of
the stronger acid until the specific gravity fell to the desired
point. (See "Making Volumetric Solutions.") Or, suppose
again that the acid in the stock bottle is concentrated, i. e
its specific gravity is 1.2 ; it would then contain 39.11 per
cent hydrogen chloride (HC1). The proportions of acid
and water to be mixed may be easily calculated by alligation.
Subtract the S. G. of water (1) from the desired S. G. of
acid (1.050) and place the result under "acid;" subtract
desired S. G. of acid from S. G. of strong acid and place
under " water," as shown below.
Desired S. G. of acid
1.050
t * N,
S. G. water S. G. acid
1 1.2
.150 .050
It is apparent that the ratio of the two numbers obtained
by the above subtraction is 3 : 1 ; hence, three volumes of
water to one of acid will give an acid solution of the desired
specific gravity. One point to be considered is whether
contraction of volume takes place on making the mixture.
If so, it is to be allowed to stand until contraction has
ceased and then bring up to the correct volume. The
percentage strengths may be used in the same way as the
specific gravities with the same results, but it is not to be
forgotten that the acid strength of water is- zero.
24
QUANTITATIVE CHEMICAL ANALYSIS.
Questions.
1. Point out the difference between metathetical and ox¬
idizing reactions.
2. What is meant by a molecular equivalent ?
3. Define the " analytical factor."
4. Name some ways in which analytical results are
stated.
5. Balance the following equations and state for each
whether it is metathetical, oxidizing, or reducing :
(a) KI + H202 = KOH + I.
(^) MnS04 -f- Na2C03 -f- O = C02 H- Na2S04 -)- Na2 Mn04.
(c) Co(NO,)2 + HN02 = H20 + NO + Co(N02)3.
{d) KC103 + H2S04 + FeS04 = Fe2(SOA + H20 + KC1.
(e) HgCl + NH3 = NH4C1 + HgNH2Cl + Hg.
(/) K2Cr207 + HI +H2S04=K2S04+Cr2(S04)3+H20 + I.
Or) HgCl2 + SnCl2 = SnCl4 + HgCl.
6. Solve the following problems :
(a) What weight of ferrous sulphate corresponds to 1.5
gms. of Fe203?
(,b) How much Mg2P207 will be obtained by the ignition of
•75 gm- of MgHP04?
(c) How many cc. of sulphuric acid, S. G. 1.210, must be
measured to get 3 gms. of H2S04 ?
(d) How many cc. of hydrochloric acid, S. G. 1.12, will be
required to dissolve 1 gm. of Fe203 ?
(e) Calculate the factor for P205 from Mg2P207*
(/) Reduce the weight of 750 gms. of water at 20° C. in air
to weight in a vacuum.
(j?) Ten gms. of a sample of sodium carbonate yield 1890 cc.
of C02; what is the degree of purity of the sample
in per cent of Na2C03?
QUANTITATIVE CHEMICAL ANALYSIS.
25
SECTION II.
GRAVIMETRIC ANALYSIS.
Chapter III.
General Operations.
The usual steps of gravimetric analysis are (1) preparation
and weighing of the sample; (2) solution ; (3) precipitation;
(4) collection of precipitate on filter and washing ; (5) drying
of precipitate and filter; (6) ignition; (7) cooling and
weighing crucible .and contents.
Preparation and Weighing of Sample.—The sample is
reduced to the state of a fine powder, and the mechanically-
held water is removed by drying in the oven at ioo° C. to
constant weight.
It is not always necessary to weigh out a certain exact
amount of the sample, but the exact weight of the portion
taken must be found. The amount of sample weighed out
should be such that the precipitate to be formed will not
exceed 1 gm. A convenient and satisfactory method of
weighing out the required amount is to place the powdered
substance in a weighing bottle and weigh the whole accur¬
ately. Then shake into a beaker an amount of the substance
judged to be about what is required. Again carefully weigh
the bottle and contents ; subtract the last weight from the
first and the difference will be the amount of the sample
transferred to the beaker. During the operation of weigh¬
ing, as just described, the bottle should not be handled
with the bare fingers, but with a piece of cheese-cloth or
filter paper.
Should it be necessary to weigh out a certain exact
quantity of the sample, it must be weighed in a tared watch-
glass or other suitable vessel, and never directly on the
26
QUANTITATIVE CHEMICAL ANALYSIS.
scale pans. When the desired amount has been nearly
reached, small quantities may be put on or taken off as
required with the point of a thin blade of a pocket knife
or a spatula. Keep the weighed substance covered with a
clean watch-glass to prevent loss from air draughts. Loss
must also be avoided in transferring a substance. The last
adhering particles are brushed off with a small camel's hair
brush.
Solution.—Under each exercise the method of procedure
includes the solution of the sample. The solvents com¬
monly employed for inorganic substances are: (i) water;
(2) hydrochloric acid ; (3) aqua regia,—a mixture of hydro¬
chloric and nitric acids, three parts and one part by volume,
respectively; (4) fusion with sodium and potassium car¬
bonates, and solution in water,—applicable to silicates.
Water alone is used whenever possible.
If heating is required, the beaker must be covered with
a clean watch-glass, and the mixture stirred with a glass
rod, the cover being raised just high enough on one side to
permit it. If the mixture is to be boiled, or if there is a
reaction in which a gas is evolved which may cause me¬
chanical loss of liquid, the substance should be dissolved
in an Erlenmeyer flask with a small funnel placed in its
mouth.
Precipitation.—A suitable precipitant, in solution, is
added in slight excess to the solution of the sample. Pre¬
cipitation may be effected in most cases at ordinary room
temperature, but it is generally best to heat the mixture.
The directions for heating during solution apply to heating
during precipitation. When the operation is concluded
the cover of the vessel is to be washed with a stream of
water from the wash bottle, the stream flowing into the
beaker containing the mixture.
The precipitant chosen is the one which will form the most nearly
insoluble product. All substances are more or less soluble; those whose
solubility is so slight as to be almost negligible are usually said to be
insoluble.
quantitative; chemical analysis.
27
The theory of electrolytic dissociation holds that substances in solution
dissociate into their constituent groups to a greater or less degree de¬
pending on certain conditions, notably temperature and concentration.
For example, sodium chloride, which is freely soluble in water, dissociates
into sodium ions, bearing positive electrical charges, and chlorine ions
bearing negative charges. The word ion is derived from a Greek verb
meaning "go," and ions are named according to the electrical pole to
which they are attracted in electrolysis; thus, sodium is a cathion and
chlorine is an anion.
The degree of ionization means the per cent, of the total dissolved
molecules which become dissociated into ions. For any given compound
there' is always a constant ratio of dissociated molecules to undissociated
molecules, represented by the formula a ^ ^ = k, where ab is the dis¬
solved undissociated molecules, a and b the ions of dissociated
molecules, and k a constaht.
... Na+XCl~
For sodium chloride, this ionic equation would read —— = ■ *
now another soluble chloride, e. g., hydrogen chloride were added to the
solution it would also ionize, giving hydrogen ions and chlorine ions.
This increase of chlorine ions would make the product a X b momentarily
greater, but the equilibrium between sodium and chlorine ions and undis¬
sociated sodium chloride, as shown in the equation above, would be
maintained. Hence, the effect of the addition of more chlorine ions by
hydrogen chloride would be to drive back the ionization of the sodium
chloride and there would be fewer dissociated molecules of this salt. If
the sodium chloride solution were saturated, the effect would be to pre¬
cipitate the molecules whose ionization had been checked. This fact is
sometimes made use of in the laboratory to prepare pure sodium chloride.
In the case of difficultly soluble substances the amount which passes
into solution is so small that it may be considered totally ionized, and
the ionic product becomes the solubility product. For example, the solu¬
bility of barium sulphate at 180 C. is 0.00001 mole per liter (1 mole =
molecular weight in grams). The ionic equation for barium sulphate is
BaSQ4~7~'>Ba + + +SO4 ; the ionic product, 0.00001 X 0.00001 = .091
(solubility product). If a soluble barium salt is added to a solution of
barium sulphate a larger number of barium ions enters into the solution
so that the ionic product Ba + + X SO4 exceeds the solubility product
of the barium sulphate with the consequent precipitation of the latter salt.
In this is the reason for the addition of the precipitant in slight excess.
Also, it will be apparent from the above paragraph why large quantities
of water are to be avoided in washing precipitates. A large excess of
precipitant is to be avoided because of the increased difficulty in com¬
pletely washing the precipitate.
Filtration and Washing.—The precipitate is separated
from the supernatant liquid by filtration, ordinarily by use
of filter paper in a funnel whose sides include an angle of
28
QUANTITATIVE CHEMICAL ANALYSIS.
6o°. The paper is folded into a cone which fits the funnel
closely, and the latter must be of such size that the paper
will not fill it completely to the rim, and it should have a
fairly long stem. Papers of 9 and 11 cm. diameter are
commonly used, but the size will be determined by the
character and bulk of the precipitate to be filtered.
Quantitative filter papers are described as " ashless," be¬
cause they have been thoroughly washed with suitable
reagents to remove almost completely all inorganic mate¬
rials. The ash content of the smaller papers is negligible
in the majority of gravimetric determinations.
Some papers filter rapidly, others more slowly, and still
others are of such close texture that it is advantageous to
use the filter pump with them. The filter flask and pump
are illustrated in Fig. 5. Papers of rather close texture are
Fig. 5.—Filter flasks and filter pump. At the left, conical funnel
with platinum cone ; right, Gooch crucible in holder.
required for certain precipitates, which, like barium sul¬
phate, contain particles so fine that they readily pass
through the more rapidly filtering papers.
The apex of the folded filter should be supported during
filtration by suction to prevent rupture of the paper. For
this purpose a platinum cone, conforming to the shape of
quantitative; chemical analysis.
29
the funnel, or a very small, folded, filter paper of tough
texture may be used, care being taken that the support
does not interfere with the filter lying close against the
sides of the funnel.
Some precipitates are advantageously filtered upon a felt
of asbestos in a Gooch crucible supported in a suitable
holder connected to a suction flask. A rubber band or
hollow rubber stopper to receive the crucible forms an air¬
tight joint around it. Asbestos of suitable form is prepared
by washing in hydrochloric acid, afterwards with distilled
water and then ignited. It is shaken up with distilled
water and poured into the crucible, suction being applied
to draw off the water and to imbed the asbestos upon the
perforated bottom of the crucible. The felt formed in this
way should be at least one milliter thick. After the water
is drawn out as much as possible, the crucible is dried in
the oven to a constant weight. It is then ready to be used
in filtering off precipitates, in which operation the suction
pump is used as in the preparation of the crucible.
Crucibles made of alundum for filtering purposes are
sufficiently porous to permit of filtration by suction, and
require no preparation other than cleaning and drying to
constant weight.
In using the filter pump do not turn on the full force of
the pump, but proceed with gentle suction.
The method of transferring a liquid from the beaker to
the filter is illustrated in Fig. 6. Be careful to direct the
stream against the side of the filter, and not upon the apex.
The point on the outer rim of the beaker where the glass
rod is held should be smeared with a little melted paraffin.
After nearly filling the filter the liquid is allowed to run
through until nearly all has passed, when it is again nearly
filled, and so on until all has been introduced.
Adhering portions of the precipitate are detached from
the sides of the beaker by rubbing off with a rubber-tipped
glass rod (" policeman"), made by passing a short piece of
QUANTITATIVE CHEMICAL ANALYSIS.
QUANTITATIVE CHEMICAL ANALYSIS. 31
rubber tubing partly over one end of a stirring rod, the
edges of the free, projecting end being pressed together and
sealed with a little rubber cement. The particles are
finally washed onto the filter as shown in Fig. 7.
Precipitates should be well washed by decantation and
all washings passed through the filter. After the precip¬
itate has been removed to the filter, washing is continued
with a fine stream from the wash-bottle until the filtrate,
upon test, shows an entire absence of soluble materials.
The manner of washing a precipitate is important. Use a
fine stream and direct it over all portions of the precipitate
until the whole of the latter is well covered. Allow this
water to pass through the filter and repeat the process
until the precipitate is completely washed.
The character of the soluble materials is known and a
suitable test must be applied to the last portions of the
filtrate to determine when they are completely removed.
For example, in washing a precipitate of barium sulphate,
precipitated from magnesium sulphate by barium chloride,
magnesium chloride and barium chloride will be present
in the solution and mixed with the precipitate. Complete
removal of these chlorides will be indicated at that point
in the washing, when the filtrate, tested with silver nitrate
solution, gives no cloudiness (absence of chlorides).
Drying Precipitate and Filter.—The precipitate and
filter are put into the oven where a temperature ioo°-iio°
is maintained, until dried to a constant weight, if they are
to be weighed together in the form in which dried. If the
precipitate is to be ignited, and this is usual, the drying
need not be carried to constant weight. In fact, they need
not be weighed at all, but are dried only until all sensible
moisture is removed. No attempt should be made to re¬
move a wet filter from the funnel. Press a moistened piece
of filter paper over the rim of the funnel for a cover and
place in the oven. When a precipitate is dried in a crucible,
the crucible may be put into a small beaker and the latter
covered with a watch-glass.
32
QUANTITATIVE CHEMICAL ANALYSIS.
Ignition.—The object of ignition may be (i) to burn off
the filter paper from the precipitate; (2) to remove all
traces of organic matter ; (3) to drive off volatile impurities ;
and (4) chiefly, in most cases, to convert the precipitate
into a substance more stable and better suited for taking
the exact weight than the original precipitate. Chlorides^
sulphates and oxides are the forms in which most metallic
substances are weighed. If a metal is precipitated as
hydroxide, carbonate or oxalate, it will be changed during
ignition to oxide. Many nitrates also change by ignition
to oxides. The following equations illustrate these changes:
2Fe(OH)3 + A = Fe2Os + 3H20.
0aCO3 -f- A = CaO -(- C02.
CaC204 -f- A = CaO -)- C02 -I- CO.
2Pb(N03)2 + A= 2PbO -f 02 + 2N2Oi.
The precipitate is ignited in a crucible which has been
previously made thoroughly clean, carefully ignited, cooled
in the desiccator, and weighed. It is better to weigh the
crucible and its cover separately on account of the possibility
of breaking the cover during subsequent ignition. Should
this happen it would not necessarily interfere with further
carrying out of the determination. The weights are recorded
in the note-book. This preparation of the crucible is to be
done every time the crucible is to be used, but time may
be saved in the weighing by reference to the former weight.
Porcelain crucibles are much used for general purposes,
but are not suitable for alkaline fusions, or where the sub¬
stance is to be treated with hydrofluoric acid. They should
always be gently heated at first.
Silica crucibles are used for much the same purposes as
porcelain crucibles, and have the advantage that they may
be subjected to sudden changes of temperature without
danger of fracture.
Platinum crucibles are best for ignitions, on account of
the well-known properties of the metal, such as its resistance
QUANTITATIVE CHEMICAL ANALYSIS.
33
to oxidation, and the effect of most acids, except a mixture
of hydrochloric and nitric acids. Also, platinum crucibles
cannot be used in the following cases : (i) fusion with the
hydroxides and nitrates of the alkalis ; (2) fusions in which
bromine, fluorine and iodine are set free ; (3) fusions in
which lead, silver, zinc, tin, bismuth, arsenic and antimony
are reduced to the metallic state, on account of their fusion
with the metal; (4) it is better not to ignite phosphates on
a filter paper in platinum crucibles because of the possibility
of reduction, whereby phosphorus may be set free and
combine with the platinum.
Platinum ware is cleaned by scouring with sea sand.
Firmly adhering stains may be removed by fusing potassium
bisulphate in the crucible and then removing the fused
mass with boiling water.
Crucibles are supported on triangles made of clay, iron
wire, silica, nichrome' wire, and platinum. Platinum
crucibles must not be supported on iron triangles.
If the precipitate has been dried before ignition it should
be removed as far as possible from the paper and collected
on a clean watch-glass or a piece of black glazed paper, and
covered with a watch-glass to prevent loss. The filter
paper is folded and a piece of platinum wire coiled about
it. It is then held over the weighed crucible and burned,
the ash dropping into the crucible. The remainder of the
precipitate is transferred to the crucible, with a camel's
hair brush, and the whole ignited. The paper is burned
separately for the reasons that many substances are reduced
when burned in contact with paper, and the burning
of the latter is accomplished easier on account of exposure
to a larger volume of air than it would meet within the
crucible.
When the precipitate and the paper are ignited together,
the crucible lid should be in place and heat gently applied
until the paper is charred, and then the whole brought to
a dull red heat. In case the paper burns with difficulty, or
3
34 QUANTITATIVE CHEMICAL ANALYSIS.
the precipitate is moist, the crucible must be slightly tilted
on the triangle as shown in Fig. 8, with the lid touching
it near the upper edge, but making a larger opening at the
lower edge, so as to direct the flame into the crucible. In
these operations every care must be exercised to prevent
loss of material. Sometimes a deposit of carbon accumu¬
lates on the surface of the crucible and the lid and must be
removed by applying the flame to such portions and burning
it off.
The Bunsen burner of suitable size answers very well
for ordinary heating operations, but the Tirrill burner illus¬
trated at A, Fig. 9, is, perhaps, the best burner for general
use in the quantitative laboratory, since it permits the regu¬
lation of both gas and air, and the flame may be made very
small or quite large. Gas is regulated by the milled screw
near the bottom ; air regulation is obtained by screwing the
body of the burner up or down as required. The air inlets
quantitative; ch^micai, analysis.
35
should not be opened to the extent that the flame is pro¬
duced with a roaring or decided blowing noise.
The M£ker burner, illustrated at B, Fig. 9, differs from
the Bunsen burner in that the volume of air admitted is
considerably larger for the same amount of gas, affording
better combustion and a hotter flame. The expanded top
contains a nickel grid which serves to break what would
otherwise be one solid flame into several small flames which
coalesce into one large flame without the large inner cone
of cold gases. Unless the burner is suited to the gas pres¬
sure at hand there will be much annoyance from " striking
back" of the flame. The Mdker burner may be used in¬
stead of the blast lamp, in many instances, for igniting pre¬
cipitates.
The blast lamp, illustrated at C, Fig. 9, is used for ob¬
taining a high-temperature flame for ignitions. Gas is
led in at the side tube and lighted at the mouth of the
burner. Air under pressure is led in at the straight con-
36
QUANTITATIVE CHEMICAL ANALYSIS.
nection and should be so regulated with reference to the
gas flow that the flame does not roar, but has only a slight
blowing sound. The size of the full flame is determined
by the opening in the tip. Blast burners are usually fur¬
nished with interchangeable tips having different sized
openings.
Cooling and Weighing.—Before a precipitate can be
weighed after ignition it must be cooled to the temperature
of the room. If the crucible and contents are weighed hot,
air currents will be set up within the balance case and un¬
equal expansion caused of the balance parts, both of which
render the weighing incorrect. Cooling, is allowed to take
place in a desiccator, Fig. 10, so that moisture may not
condense on the crucible. The air within the desiccator is
kept dry by a layer of sulphuric acid about i cm. deep, or
the bottom compartment may be fairly filled with lumps of
anhydrous calcium chloride. The ground rims of the
desiccator and its lid should be smeared lightly with vase¬
line, and the lid must be kept in place except when it is
necessary to remove it to introduce or remove something
from the desiccator.
As soon as the crucible and contents are cold they should
be weighed. Do not allow substances to remain for a great
while in the desiccator before weighing.
Fig. 10. —Desiccator.
QUANTITATIVE CHEMICAL ANALYSIS.
37
Questions.
1. What are the steps of a gravimetric determination?
2. State how a substance is prepared for analysis and
how weighed.
3. Discuss the solution of the sample.
4. What governs the choice of the precipitating agent ?
5. Why is the precipitant added in " slight excess ?"
6. What is meant by ionization, and what is the ioniza¬
tion constant?
7. What is the effect of adding to a solution a substance
giving ions common to existing ions ?
8. Describe u quantitative" filter papers.
9. Describe the preparation of a Gooch crucible.
10. Why are large quantities of water to be avoided in
washing most precipitates?
11. How is the completion of the washing of a pre¬
cipitate determined?
12. What is the object of ignition of a precipitate?
13. Why is the filter paper burned apart from the pre¬
cipitate?
14. Describe the ignition of moist precipitates.
15. Should an object be allowed to stand indefinitely in
the desiccator before weighing ?
Chapter IV.
GRAVIMETRIC EXERCISES.
Exercise VIII.
Determination of Iron.
Iron is determined gravimetrically by precipitation in the
ferric state by ammonium hydroxide as ferric hydroxide>
and ignited to ferric oxide, in which form it is weighed.
FeCl3 + 3NH4OH = Fe(OH)3 + 3NH4C1.
2Fe(OH)3 = FeA + 3HA
33
QUANTITATIVE CHEMICAL ANALYSIS.
The instructor will furnish you with a solution of an
iron salt. You are to report the amount of iron in the
portion of solution given you.
Transfer the solution to a 250 cc. beaker and rinse the
smaller vessel originally containing the solution with dis¬
tilled water and pour into the larger beaker. Add distilled
water enough to bring the volume up to at least 50 cc.
Add about 3 cc. of hydrochloric acid, and boil for a few
minutes.
Then add 2 cc. of nitric acid and continue to boil for
fifteen or twenty minutes. Set the solution aside to cool,
and then pour into 200 cc. of water containing 10 cc. of
ammonium hydroxide and again boil for two or three
minutes, stirring the mixture constantly.
Allow the precipitate to settle, and then pour off the
supernatant liquid through a filter, leaving the precipitate
in the beaker. Wash the' precipitate with hot water by
decantation, passing all the washings through the filter-
Finally, transfer all of the precipitate to the filter and wash
with hot water from the wash bottle until the absence of
chlorides in the filtrate is shown by the test with silver
nitrate solution.
Dry the precipitate in the oven and ignite over a Tirrill
burner; at the last, use the blast burner for a few minutes.
Cool the crucible and contents in the desiccator and weigh.
Calculate the amount of iron in the solution.
Exercise IX.
Determination of Lead.
In this exercise the amount of lead in lead nitrate is
determined by ignition of the salt to lead oxide. Any salt
of lead capable of being changed into the form of nitrate
may be analyzed for its lead content in this way. Also,
the soluble salts of lead may be determined by precipitation
with hydrogen sulphide and weighed as lead sulphide.
quantitative; chemical analysis.
39
Weigh out from the weighing bottle about 0.5 gram of
lead nitrate and put into a crucible which has been previ¬
ously cleaned, ignited and weighed. Place the lid on the
crucible and heat gently until all crepitation has ceased,
and a large amount of nitrous fumes is no longer given
off; then, at dull red heat, continue the heating until no
fumes at all are evolved. Cool the crucible and contents,
and weigh. From the amount of lead oxide found calculate
the per cent of lead in the sample.
Pb(N03)2 = PbO + 2N02 + NO.
Exercise X.
Determination of Mercury.
a. Precipitation as Mercurous Chloride.—Mercury must
be in the mercurous form to be precipitated as mercurous
chloride. If a mercuric salt is to be determined it must first
be reduced.
For this exercise you are given a solution of a mercuric
salt. Add to the solution 3 cc. of hydrochloric acid and
an excess of phosphorous acid and heat for 15 minutes at
550 to 6o° C.; not higher than 6o°. Allow the precipitate
to settle, and filter through a Gooch crucible which has
been previously properly prepared. Report the amount of
mercury.
b. Precipitation as Sulphide.—Dilute the mercury solu¬
tion given you to about 100 cc. Saturate the solution with
hydrogen sulphide, and filter through a Gooch crucible.
Wash thoroughly with cold water, dry at ioo° C. and weigh.
It may happen that free sulphur is mixed with the precipi¬
tate, and for this reason it would be well to again wash
with carbon disulphide, and dry again at no0 C. Report
the amount of mercury.
4°
QUANTITATIVE CHEMICAL ANALYSIS.
Exercise XL
Determination of Magnesium.
Magnesium is precipitated in alkaline solution as the
acid phosphate (MgHP04), which is afterward ignited to
magnesium pyrophosphate, Mg2P207, in which form it is
weighed.
The solution given you is to be diluted to about 50 cc.,
and treated with 5 cc. of hydrochloric acid and 10 cc. of di-
sodium phosphate solution. Then, with constant stirring,
add dilute ammonia until the mixture is distinctly alkaline.
Be careful not to rub the sides or bottom of the beaker
with the glass rod, on account of the formation at such
points of a strongly adhering precipitate. Cover the beaker
and set aside for twelve hours. Filter through a 9 cm.
filter paper and wash with very dilute ammonia until the
washings, tested with silver nitrate, show only the slight
turbidity given by the ammonia.
Cover the funnel and contents, convey to the oven and
dry at ioo°. By means of a small brush of camel's hair
remove as much as possible of the dry precipitate from the
filter to a piece of black glazed paper. Then fold the filter
paper for burning on a platinum wire, allowing the ash to
drop into a weighed crucible.
If difficulty is found in completely burning portions of
the paper within the crucible, it may be moistened with
nitric acid after cooling and the Bunsen flame again applied.
When the paper has been completely burned to a white
ash, add the precipitate to the crucible. Heat the crucible
and contents gently for a few minutes and then turn on the
full flame of the burner; finally, ignite over the blast
burner. Cool in the desiccator and weigh. From the
quantity of magnesium pyrophosphate obtained calculate
the amount of magnesium in the solution given you.
2MgHP04 = Mg2P207 + H20.
quantitative; chemical analysis.
41
Exercise XII.
Determination of Sulphuric Acid.
Sulphuric acid and soluble sulphates are determined by
precipitation with barium chloride. The resulting barium
sulphate is very insoluble, but the particles, unless formed
in hot solution, are very fine and easily pass through most
filters. For this reason precipitation is carried out in boil¬
ing solution, and the mixture afterward boiled for several
minutes, which causes still greater enlargement of the
particles. A large excess of barium chloride is to be avoided
in the precipitation on account of the tendency which
barium sulphate has to carry down other substances with it
(absorption) which renders complete washing of the precipi¬
tate difficult. The method to be described for the deter¬
mination of sulphates may be used for the determination of
barium by using a soluble sulphate as precipitant.
Dilute the solution of a neutral sulphate given you to
about 100 cc., add 3 cc. of hydrochloric acid, cover and
heat to boiling. Also, heat 50 cc. of 5 per cent barium
chloride solution to boiling and pour cautiously into the
boiling sulphate solution. Boil for at least five minutes,
stirring constantly. Should the precipitate not settle
readily, boil again for a few minutes.
Decant the supernatant liquid through a small filter pre¬
viously wetted with hot water, retaining the precipitate in
the beaker. Add 100 cc. of hot water and 1 cc. of dilute
hydrochloric acid to the precipitate in the beaker, stir well,
allow the precipitate to settle, and pour the liquid through the
filter. Again wash the precipitate by decantation with hot
water alone, and then transfer the precipitate to the filter.
Continue the washing with hot water from the wash bottle
until the final washings show no trace of chlorides with
silver nitrate.
The filter and contents may be dried in the oven before
ignition, or they may at once be transferred to a weighed
42
quantitative; chemical analysis.
crucible. If the latter is done the crucible should be ar¬
ranged on the triangle as described under ignition and
illustrated at Fig. 8. Heat should be applied gently at first
to the cover so as to dry the filter, and when the water has
been expelled the crucible is gently heated to drive off volatile
portions of the paper. Afterward, apply the flame to the
bottom of the crucible, using a large flame, until the paper
is all charred. Allow the crucible to cool, moisten the
residue with concentrated nitric acid and again ignite until
the residue is of a uniform, grayish-white appearance. Cool
in the desiccator, weigh, and calculate the quantity of sul¬
phuric anhydride, S03, in the solution.
Exercise XIII.
Determination of Hydrochloric Acid.
Hydrochloric acid and soluble chlorides are determined
by precipitation with silver nitrate and weighed as silver
chloride. When a soluble chloride is used as the precipitant
the same method may be used for the determination of
silver.
Silver nitrate is also used for the 'precipitation of bro¬
mides, iodides and cyanides. Chlorates, after reduction to
chloride, may be determined by precipitation with silver
nitrate.
Dilute the chloride solution given you to about 150 cc.,
add ten drops of concentrated nitric acid and 25 cc. of silver
nitrate solution. Carefully heat the mixture to boiling,
stirring occasionally, until the precipitate forms in curds and
promptly settles. Add a drop or two more of silver nitrate
solution to the clear supernatant liquid to observe if further
precipitations occurs. During the boiling protect the
solution from direct sunlight and as far as possible from
strong diffused light.
When precipitation is complete, pour off the supernatant
liquid through a prepared Gooch crucible, leaving as much
QUANTITATIVE CHEMICAL ANALYSIS.
43
as possible of the precipitate in the beaker. Wash the
precipitate two or three times by decantation with hot water,
then transfer to the Gooch crucible, and wash , with hot
water from the wash bottle until the filtrate gives no cloud¬
iness with a drop of hydrochloric £cid.
Dry the crucible and contents at 130° to constant weight,
and calculate the chloride as hydrogen chloride.
Exercise XIV.
Determination of Phosphoric Acid.
{a) By Precipitation as Phosphomolybdate.—Prepare a
concentrated solution of ammonium nitrate. Prepare also a
molybdate solution as follows: Dissolve 37.5 gms. of am¬
monium molybdate in 250 cc. of water to which a little
ammonia is added. Filter if turbid.
Make a mixture of 125 cc. each of concentrated nitric
acid and water, and while constantly stirring, pour the
solution of ammonium molybdate into it. Set the solution
aside in a warm place for a few days and then filter before
use.
Heat the phosphate solution given you nearly to boiling
and, while constantly stirring, add 50 cc. of molybdate
solution. Stir for a while, maintaining the temperature on
a water bath for about ten minutes. Allow the mixture to
stand for an hour; then, filter through a Gooch crucible.
Test the filtrate for the presence of phosphoric acid by the
addition of molybdate solution ; if present, return the pre¬
cipitate to the crucible.
Wash the precipitate thoroughly with a solution contain¬
ing 50 gms. of potassium nitrate and 40 cc. of concentrated
nitric acid per liter.
Place the Gooch crucible inside of a larger crucible and
heat gently until there is no further evolution of ammonia ;
then ignite, cool, and weigh as phosphomolybdic anhydride.
44
QUANTITATIVE CHEMICAL ANALYSIS.
H3P04 + i2(NH4)2Mo04 + 2iHN03=
(NH4)3P04i 2M0O3 + 21NH4NO3 + i2H20.
2(NH4)3P04i2M0O3 = 24M0O3P2O5 + 6NH3 -j- 3H20.
Report the amount of P2Os in the portion of solution given
you.
(b) Precipitation as Ammonium-Magnesium Phosphate.
—The phosphate may be at once precipitated by magnesia
mixture (see determination of magnesium), or, the phos-
phomolybdic precipitate may be redissolved and precipitated
by magnesia mixture. If the latter method is employed,
after washing the molybdate precipitate a beaker is placed
under the funnel and the filter pierced ; wash the precipi¬
tate into the beaker with a fine stream of water. Moisten
the paper with a little ammonia and wash with a small
quantity of hot water. Add sufficient ammonia to dissolve
the precipitate. Add hydrochloric acid to neutralize the
solution, but if phosphomolybdate should be again pre¬
cipitated add more ammonia to dissolve it. Again make
the solution acid with hydrochloric acid, and add sufficient
magnesia mixture to precipitate all the phosphate present.
With constant stirring, add ammonia to the mixture until
it is neutral and then add a slight excess of ammonia.
Set the mixture aside over night, then filter off the pre¬
cipitate, wash with dilute ammonia, ignite and weigh. The
residue is magnesium pyrophosphate.
If the sample of phosphate to be analyzed contains only
the alkali metals and if ammonium salts are present in
moderate quantity only, the determination by magnesia
mixture is satisfactory; but if other metals are present,
precipitation by phosphomolybdate solution is advantage¬
ous, but hydrochloric, sulphuric and boric acids must be
absent. The quantity of nitric acid present should not
exceed 1.6 cc. of the concentrated variety for every 100
mg. of phosphoric acid. In case the latter method is em¬
ployed better results are secured by transforming the
quantitative: chemical analysis.
45
molybdate precipitate into ammonium-magnesium phos¬
phate, and igniting to magnesium pyrophosphate.
The determination of phosphorus constitutes one of the
important features of fertilizer analysis, and the method of
the Association of Official Agricultural Chemists should be
read (U. S. Dept. Agr., Bur. Chem., Bull. 107).
(a) By Loss.—Use the apparatus- shown in Fig. 11,
Mohr's alkalimeter. Weigh out 2-3 grams of the powdered
carbonate material and put into the flask a; nearly fill
bulb c with dilute hydrochloric acid, partly fill bulb f with
concentrated sulphuric acid. Insert the bulbs into their
respective places in the flask. Weigh the whole apparatus,
Exercise XV.
Determination of Carbon Dioxide.
Fig. 11.—Mohr's alkalimeter.
46
QUANTITATIVE CHEMICAL ANALYSIS.
then remove from the balances and turn stop-cock d so as to
allow the hydrochloric acid solution to flow into the flask
and decompose the carbonate. When the acid has all run
out close the stop-cock again, allow the apparatus to set
until the evolution of C02 has ceased, and then warm gently
on the hot plate to expel the gas. Bring the apparatus to
the temperature of the laboratory and replace the carbon-
dioxide in the apparatus with air by careful suction at e,
having opened stop-cock d. Weigh the apparatus; the loss
in weight will be the carbon dioxide.
This method of the determination of carbonates is not
accurate, but will serve very well for such work as the
comparison of available C02 in different baking powders.
In such cases water would be used in bulb c instead of
quantitative; chemical analysis.
47
hydrochloric acid, unless the object were to determine total
C02.
(ib) By Absorption.—This method depends on the absorp¬
tion of C02 by caustic potash and the determination of the
quantity involved by the increase of weight of the bulb.
A satisfactory way of doing this is by the method of Rose,
illustrated at Pig. 12. Tube g, which is about 70 cm.
long, is half-filled with calcium chloride, as shown, a plug
of glass wool being placed near its middle. A stream of
dry carbon dioxide is blown through the tube to saturate
the calcium chloride, after which dry air is blown through.
The potash bulb h is partly filled with caustic potash solu¬
tion, and the U-tube k is filled with calcium chloride. The
whole apparatus must be tested for air-tightness and the
rubber connections are to be coated with shellac.
Weigh out about 1 gram of the carbonate material and
put into flask a ; put about 50 cc. of dilute hydrochloric
acid into dropping funnel c, with which is connected the
tube b filled with calcium chloride. Carefully weigh the
potash bulb and record its weight. When the potash bulb
has been again attached and the apparatus made air-tight,
open the stop-cock of the dropping funnel and allow the
dilute acid to run into the flask at such a rate that no more
than two bubbles of gas per second pass through the potash
bulb. After the acid has all run in, close the stop-cock and
gently heat the flask, gradually bringing it to the boiling
point. It must not be boiled to the extent that the gases
pass through the caustic potash in a solid stream, but in a
series of bubbles which may be counted. When the opera¬
tion is completed, open the stop-cock of the dropping funnel
and draw air through the apparatus to thoroughly remove
all carbon dioxide. Remove the caustic potash bulb and
again weigh. The increase in weight will be the weight
of the carbon dioxide.
This method has been slightly modified from the usual
manner of conducting the determination and some details
48
QUANTITATIVE CHEMICAL ANALYSIS.
omitted, but careful attention to carrying out the experiment
as given will result in a fair degree of accuracy.
Chapter V.
ELECTROLYTIC METHODS.
Electrolytic analysis depends on the fact that many
metals may be separated from solutions of their salts by
deposition on a suitable electrode, dried and weighed. For
the purpose of studying electrolytic effects, consider first
two platinum electrodes immersed in pure water and con¬
nected with a source of electricity. It is easily demonstrated
that no current will flow through the water from electrode
to electrode. Now, dissolve in the water some copper
sulphate ; a current passes.
The explanation offered is that copper sulphate in solution
dissociates into Cu cathions and S04 anions, and the current
is conveyed through the solution by these ions. The
copper ions, bearing a positive charge, go to the negative
electrode and there yield up their electrical charge, the
ionic copper being deposited as metallic copper.
The amount of electrical charge borne by each ion is
proportional to the valence of the substance, hence, all ions
of equal valence carry the same amount of electrical charge.
It also follows that the amount of substance deposited by a
given current must be directly proportional to its atomic
mass and inversely proportional to its valence.
For each metal there is a minimum of potential difference
at which it will be deposited, which makes it possible to
separate metals in solution by electrolyzing at increasing
voltages as each metal is successively removed.
Metals tend to redissolve in the media in which they are
deposited and the amount of current in amperes must be
sufficient to deposit the metal faster than it dissolves. The
surface area of the electrode will determine the amount of
QUANTITATIVE CHEMICAL ANALYSIS.
49
current, since the whole surface of the electrode presents
deposited metal to the action of the solvent.
Current Density.—By current density is meant the
amount of amperage per unit area, usually 100 sq. cm. One
ampere per 100 sq. cm. is expressed by the formula,
CD100; 1.25 amperes per 100 sq. cm. by CD100= 1.25 amperes.
Source of Current.—The commercial no-volt direct-cur¬
rent lighting circuit may be used if means are employed to
reduce the voltage and control the amperage. Amperage
may be satisfactorily controlled by passing the current
through lamps of suitable resistance connected in parallel.
Lamps are rated in candle-power or in watts. The watt is
an electrical unit found by multiplying the voltage of the
circuit by the amperes used. A lamp of 25 watts rating
will draw from the no-volt circuit, .2277 amperes (25 -r-
110 = .2277); 50 watts, .4554 amperes; 100 watts, .910
amperes, etc. Carbon filament lamps rated as 16 candle-
power, pass .4 ampere; 32 candle-power, .8 ampere ; 50
candle-power, 1.2 amperes.
When lamps are connected in parallel the value of the
current is the product of the amperage of one lamp by the
total number of lamps ; if in series, it is the quotient of the
amperage of one lamp by the total number of lamps.
The voltage employed must be sufficient to carry the de¬
sired amount of current through the solution being elec-
trolyzed, since the current in a circuit is directly proportional
to the voltage and inversely proportional • to the resistance
(Ohm's law). When the no-volt circuit is used the elec¬
trolytic solution is placed in shunt with an adjustable re¬
sistance and in series with the parallel lamp bank. The
voltage is brought to the proper value by sliding the shunt
contact along the resistance member until the voltmeter or
ammeter indicates the proper passage of current.
The storage battery is also a convenient source of cur¬
rent, the amperage being controlled by suitable resistance
coils. The single storage cell gives a voltage of 2.2 when
4
5°
QUANTITATIVE CHEMICAL ANALYSIS.
charged, and when connected in series the voltage is the
product of a single cell by the total number of cells. Do
not attempt to measure the current of the storage battery
with an ammeter; there is danger of damaging the outfit
through the short circuiting of the battery. An adjustable
resistance is to be used in series with the solution.
Fig. 13.—Open cylindrical and helical electrodes.
Electrodes.—The electrodes for electrolysis consist of
cylinders of platinum, one smaller than the other, so
as to be placed within it, or the smaller electrode may
be a spiral or helix of platinum wire. Sometimes the op¬
eration is carried on in a platinum dish which forms the
electrode for receiving the deposit.
The time devoted to electro-deposition is shortened by
warming the solution to 6o°-y^° C.; also, deposition takes
place in better form if the electrode is rotated.
Exercise XVI.
Determination of Silver in Solution.
Pour the silver solution furnished you into 25 cc. of
potassium cyanide solution containing 3 grams of the salt.
Connect the electrolysis apparatus, the cylinder electrode
quantitative chemical analysis.
51
being attached to the negative wire, the helix to the
positive. Warm the solution to 65° C. Pass a current of
.4 ampere at a pressure of 2.5 volts. At the end of 3 hours
remove a drop of the solution on a clean glass rod and pass
into hydrochloric acid in a small test tube and observe if
there is any cloudiness produced; if so, continue the
electrolysis until the test shows tke absence of silver in the
solution.
Siphon off the solution, at the same time adding distilled
water. Caution : Do not inhale the vapors arising during
electrolysis, nor Jill the siphon by suction.
Rinse off the electrode with alcohol and dry for ten.
minutes in the oven at ioo° C. Cool, weigh and calculate,
the amount of silver nitrate in the solution given you.
Exercise XVII.
Determination of Copper in Solution.
Weigh out about 1.2 gms. of crystallized copper sulphate,
dissolve in 100 cc. of water and add 1 cc. of concentrated
nitric acid. Warm the solution to 70° C. Connect the
cylinder electrode to the negative wire, as in the exercise
above, and pass the current at .1 ampere at 2 volts.
When there is no longer any blue color left in the
solution remove a few drops and treat with concentrated
ammonium hydroxide to test whether the copper has all
been removed. If not, neutralize the portion tested for
copper by nitric acid, return to the beaker containing the
solution and continue the electrolysis for some time longer,
and until the test shows that all copper has been removed
from the solution. Siphon off the liquid, adding water,
and then rinse with alcohol, dry and weigh as in the
exercise above. Calculate the per cent of copper in crys¬
tallized copper sulphate.
52
QUANTITATIVE CHEMICAL, ANALYSIS.
SECTION III.
VOLUMETRIC ANALYSIS.
Chapter V.
General Principles.
It has already been stated that volumetric determinations
require the use of test solutions of definite strength, and the
amount of substance to be determined is calculated from
the volume of the test solution required to complete a
reaction.
Solutions for chemical operations are made up in terms
of molar concentration, percentage strength, or normality;
and, in some instances, an arbitrary standard of strength is
selected.
Volumetric test solutions are usually made up on the
basis of normality.
A Normal Solution is one which contains in one liter of
solution an amount of the test reagent equivalent to one
gram of hydrogen. On this basis solutions may be multiples
or submultiples of a normal solution. Usually for analytical
purposes submultiples of the normal solution are used, as,
half-normal (N/2), fifth-normal (N/5), tenth-normal (N/10),
centi-normal (N/100), etc. These solutions contain respect¬
ively, one-half, one-fifth, one-tenth, one-hundredth, etc., as
much of the reagent in one liter of solution as contained in
a normal solution.
It has already been indicated that volumetric estimations
are based on the complete reaction between the reagent
and the substance tested. Consider the formula weights of
hydrochloric acid, sulphuric acid and sodium carbonate,
and the reactions between these two acids and the carbon¬
ate : the formula weight of hydrogen chloride is 36.462 ; of
hydrogen sulphate, 98.09 ; of sodium carbonate, 106. The
reactions are as follows :
quantitative: chemical analysis.
53
2HC1 + Na2C03 = 2NaCl + H20 + C02.
H2S04 + Na2C03 = Na2S04 + H20 + C02.
It is plain that hydrochloric acid contains one equivalent
of hydrogen, sulphuric acid two equivalents, while sodium-
carbonate, though not containing hydrogen, is equivalent
to one hydrogen sulphate, therefore, equivalent to two-
hydrogen grams. Normal solutions of these would contains
in one liter one formula-weight of hydrogen chloride, one--
half formula-weight of hydrogen sulphate and of sodium
carbonate.
Rule.—The amount of substance required to make a liter
of normal solution is found by dividing the formula-weight
by its hydrogen equivalent.
For ordinary reactions the hydrogen equivalent may be
determined by the valence of the compound. If the sub¬
stance is a solid containing water of crystallization, the
water is added to the formula-weight of the substance.
Making Volumetric Solutions.—If the reagent is a pure,
solid substance the required amount is carefully weighed
out and transferred to a volumetric flask of suitable capa¬
city. A portion of distilled water is added, the reagent
thoroughly dissolved, and more water added exactly to the
mark on the neck of the flask, and at the temperature at
which the flask was calibrated. Then the solution should
be thoroughly mixed.
Instead of trying to get an exact, fixed weight of reagent
it will be found more convenient to weigh out a quantity
of material as near the amount wanted as convenient, and
calculate the volume of solution which this quantity of
reagent will make. Example: Suppose it is desired to
make a normal solution of sodium carbonate, and that
53* I5° grams have been weighed out (53 grams are required
to make a liter of normal solution). By the proportion 53 :
1,000 cc. : : 53.150 : x = 1,002.8 cc., the volume (1,002.8
cc.) is found to which the solution of the weighed material
54
QUANTITATIVE CHEMICAL ANALYSIS.
must be made up. If the weighing and measuring are
correctly done the solution will be normal.
In the example cited above it would be convenient to
choose a flask marked to deliver 500 cc. Placed the
weighed material in the flask, add water enough to dissolve
the salt and then fill exactly to the mark. Now transfer
the solution to a bottle large enough to hold the entire
amount of solution to be made. Add more water to the
500 cc. flask, giving the flask a rotary motion so as to
thoroughly incorporate the solution from the first mixing
which still adheres to the flask, and finally again fill exactly
to the deliver mark and transfer to the bottle ; 2.8 cc. still
-remain to be added and may be measured with a pipette or
with a burette, after which the solution should again be
Ihoroughly mixed.
If it is desired to make a solution of a reagent which is
already in solution, or which is a liquid, its strength may
be ascertained by a gravimetric analysis ; or, if its specific
gravity is known reference to tables will show the strength.
Then a suitable portion may be diluted to a concentration
somewhat greater than that desired, and afterward standard¬
ized against another solution of known concentration, or
a gravimetric determination may be made.
Solutions of the mineral acids are conveniently made in
this way. Sulphuric acid may be standardized by precipi¬
tation with barium chloride; hydrochloric acid, by silver
nitrate, etc.
There is to be found in the market a variety of apparatus
for finding the specific gravity of liquids by direct weighing.
Of specific-gravity bottles, one of the simplest forms is that
figured at C, Fig. 14.
At B, Fig. 14, is illustrated a hydrometer graduated for
densities of 1.000 to 1.200. The instrument is immersed
in the liquid to be tested and the specific gravity is read off
directly when the liquid is at the temperature of stand¬
ardization of the instrument.
QUANTITATIVE) CHEMICAL ANALYSIS.
c
Fig. 14.—A, Westphal balance; B, hydrometer ; C, specific
gravity bottle.
The Westphal balance is shown at A, Fig. 14. The
beam of the balance is divided into 10 equal parts, the
tenth division coinciding with the point of suspension of
56
quantitative; chemical analysis.
the hanger to which the plummet is attached. The
plummet embodies a thermometer, and when immersed in
pure water the pointer on the beam is brought to coincide
with the fixed pointer of the balance by raising or lowering
the apparatus by means of the screw shown in the base of
the balance. In the illustration the largest weight is
shown at 7, the next larger at 9, and the smallest at 5.
The specific gravity would be read .795« For specific
gravities greater than water a weight equal to the largest
is swung from the hanger. The other weights are read as
already indicated, e. g., 1.795.
Titration.—Titration means the addition of the test
reagent to the solution tested in measured quantities until
the reaction is complete. The addition of the test reagent
is accomplished by means of the burette, described under
measuring vessels.
The solution to be tested may be measured from a burette
or a pipette ; but it is important that the measuring vessels
used should agree, or the proper correction made for any
errors which may exist.
Always, in using a burette, allow a few minutes to elapse
before making the reading in order that the film of liquid
adhering to the walls of the instrument may drain ^nd
show the correct height of reagent.
When the end-point is reached the flow of reagent is
stopped immediately, and if the color change is very deep
the titration should be repeated, adding less reagent than
at first and completing the titration by the addition of single
small drops. In any case titrations should be done in du¬
plicate and should agree within one-tenth of a cc.
Correction Factor.—When a test reagent is greater or less
than the desired concentration the increment of error is
added to or subtracted from that strength as unity, making
a factor for correction of the amount of reagent added, thus,
"1.057 N/10 hydrochloric acid," or, "0.996 N/10 hydro¬
chloric acid." The number of cc. used, multiplied by the
quantitative; chemical analysis.
57
correction factor gives the correct number of cc. in the
basic normality.
Indicators.—The end-point, or completion, of a reaction
is shown by the color change of a substance added to the
solution tested, called an indicator. Some indicators pos¬
sess a dual character of acid and alkaline properties with
one or the other predominating. For this reason some are
more sensitive to acids than to bases, and vice versa, while
others are moderately sensitive to both acids and bases.
Some indicators commonly used with acids and bases
are : i, phenolphthalein ; 2, rosolic acid ; 3, methyl orange ;
4, cochineal; 5, litmus. One and 2 are very sensitive to
acids ; 3 and 4, to bases ; 5 is moderately sensitive to both
acids and bases.
PhenolphthalEin is a yellowish-white powder, almost
insoluble in water. Solutions for use in volumetric analyses
are prepared by making a 1 per cent solution in 50 per cent
alcohol. It is colorless with acids; with bases it gives a
bright pink to deep red color.
Rosolic Acid occurs in fairly large crystals of an irri-
descent red color. It gives a red color with bases and a
yellow with acids. It is made up in 1 per cent solution in
60 per cent alcohol.
Methyl Orange occurs as a yellow powder, yielding a
yellow solution in water. Solutions of about .1 per cent are
generally used. It is yellow with bases and red with acids.
Cochineal owes its coloring properties to carminic acid,
and is derived from the bodies of the dried female insect,
coccus cacti, Linn£. The indicator solution is made by
digesting the dried insects with alcohol. It is violet with
bases and red with acids.
Litmus is derived from certain lichens. Watery solu¬
tions are used, giving red with acids and blue with bases.
Its moderate sensitiveness render it of little use in close
work in analysis.
Care must be taken not to use too much indicator solu¬
tion. Usually one drop, or two, will be sufficient, and in
QUANTITATIVE CHEMICAL ANALYSIS.
any case not more than three drops of solution will be re¬
quired.
Volumetric Equivalents and Calculations.—A volumetric
equivalent is the amount in grams of the substance tested
equivalent to the number of grams of the test reagent con¬
tained in i liter of the test solution.
Volumetric Factors.—By dividing the quantity per liter
of normal strength by 1,000, the equivalent per cc.,—volu¬
metric factor,—is obtained ; or, if the test solution is tenth¬
normal, divide by 10,000. When the manner of arriving
at the volumetric factor is understood it will be seen that,
expressed in milligrams, it is equal to the number of grams
of the substance required to make a solution of concentration
equal to the test solution. Thus :
(a) What is the normal volumetric factor for sodium
carbonate ?
The amount of sodium carbonate required to make a liter
of normal solution is 53 grams; therefore, 53 grams h-
1,000 = 0.053 gram = 53 milligrams.
(3) Sodium carbonate solution is titrated with 1.02-nor-
mal sulphuric acid solution ; what is the volumetric equiva¬
lent or factor?
Multiply the normal volumetric factor for sodium car¬
bonate by 1.02 X 0.053 §"m- = 0.05406 gm.
The following examples of volumetric calculations are
added by way of illustration :
1. 25 cc. of a solution of hydrochloric acid require 60 cc.
of N/10 sodium hydroxide solution for saturation; how
much hydrochloric acid in the solution ?
The N/10 factor for hydrochloric acid is 0.003646 gm.
Then, 60 X 0.003646 = 0.21876 gm. HC1 in the 25 cc. of
solution.
2. How much HC1 in a liter of the solution ?
Since 25 cc. contain 0.21876 gm., 1,000 cc., or 1 liter,
,, ^ . 1,000 X 0.21876
would contain — = 8.7504 gms.
^5
3. What is the per cent strength of the acid solution ?
QUANTITATIVE CHEMICAL ANALYSIS.
59
To calculate the per cent of acid the specific gravity of
the solution must be known. Then,
(25 X S. G. of solution) : (HC1 in 25 cc.) : : 100 : x.
In the above statement (25 X S. G.) gives the weight of
25 cc. of the solution.
4. What is the normality of the solution ?
A normal solution of hydrochloric acid contains in 1 liter
36.464 gms. HC1; therefore, the solution is normal.
36.464
Or, since a N/10 solution contains 3.464 gms. HC1 per liter,
8.7504 -7- 3.464 = 2.4 i. e., the solution is 2.4 X N/10.
5. How should the solution be diluted to bring it to
decinormal concentration ?
Since the solution is 2.4 X N/10, one volume of the
strong solution must be diluted with distilled water to
make exactly 2.4 volumes.
6. The solution just tested was made by taking 10 cc. of
acid from the stock bottle and diluting to 400 cc,; what is
the strength of the acid in the stock bottle ?
Since 10 cc. of the stock acid were used and diluted to
400 cc. the solution is 1/40 the concentration of the stock
acid ; or, in other words, the stock acid is 40 times the
concentration of the solution, which has been shown to be
2.4 N/10, or .24N. Then 40 X .24N = 9.6N, the concen¬
tration of the stock acid.
Chapter VI.
ACIDIMETRY AND ALKALIMETRY.
Acidimetry is the determination of the amount of an acid
in solution by volumetric methods. Alkalimetry is like¬
wise the determination of the amount of a base in solution
by volumetric methods. Standard solutions of bases and
acids are required in the respective cases.
Standard Oxalic Acid Solution.—Crystallized oxalic acid
(C2H204.2H20) may be obtained of sufficient purity to be
used for making a standard solution directly from the
6o
QUANTITATIVE CHEMICAL ANALYSIS.
weighed portion. It has the disadvantage that its solutions
are not permanent, but the advantage of a pure, solid acid
substance for a starting solution is obvious. The recently-
prepared oxalic acid may be used for standardizing an
alkaline solution which may then be likewise employed for
more stable acids.
Carefully weigh out from a weighing bottle what is
judged to be about 6.3 grams of oxalic acid (H2C204.2H20).
Calculate the volume of N/10 solution to be made from the
amount weighed out. Since 6.3 grams of oxalic acid are
required to make a liter of N/10 solution, the following
proportion will give the volume of solution to be made :
6.3 : 1,000 cc. : : weight found : required volume of solution.
Standard Hydrochloric Acid Solution.—Make a N/10
solution of hydrochloric acid.
First Method.—Take the S. G. of a sample hydrochloric
acid from the stock bottles and, by reference to a table of
properties, calculate the amount of dilution necessary to
bring a portion of the acid to a strength somewhat greater
than N/10. Prepare such a solution, and determine its
strength by precipitation with silver nitrate and a gravi¬
metric determination. Standardize according to the findings
of the gravimetric determination, and again determine its
strength in the same way.
Density and Concentration of Hydrochloric Acid at 15°.
(Normal HC1 =36*47 grams per litre.)
Density
16°/4°.
100 grams
contain
grams
HC1.
1 litre
contains
grams
HC1.
Nor¬
mality.
Density
100 grams,
contain
-grams
HC1.
1 litre
contains
grams
HC1.
Nor¬
mality.
1*010
2*14
22
0*6
1*110
21*9
243
6*7
1*020
4*13
42
1*2
1*120
23*8
267
7-3
1*080
6*15
64
1*8
1*180
25*7
291
8*0
1*040
8*16
85
2*3
1*140
27*7
315
8*6
1*050
10*17
107
2*9
1*150
29*6
340
9*3
1*060
12*19
129
3*5
1*160
31*5
366
10*0
1*070
14*17
152
4*2
1*170
33*5
392
10*8
1*080
16*15
174
4*8
1*180
35*4
418
11*7
1*090
18*1
197
5*4
1*190
37*2
443
12*1
1*100
20*0
220
6*0
1*200
39*1
469
12*9
QUANTITATIVE CHEMICAL ANALYSIS. 61
Second Method.—Prepare a N/io hydrochloric acid solu¬
tion from constant-boiling acid. Hydrochloric acid by dis¬
tillation reaches a concentration at which the boiling point
remains constant. This is known as constant-boiling acid,
which, at 760 mm. pressure, in 180.170 grams contains one
mole, or 36.464 grams of hydrogen chloride, the quantity
required to make one liter of normal solution.
To make a N/10 solution, measure into a tared 100 cc.
flask 17.3 cc. of constant-boiling hydrochloric acid; weigh
carefully, and calculate the volume of N/10 solution to be
made;
18.017 : 1,000 cc. : : gms. found : x= number of cc. of
total volume of solution. Pour the weighed acid into a
flask of suitable size, as already described, and rinse the 100
cc. flask several times with distilled water, pouring each
rinsing into the flask with the acid. Finally bring the
volume up exactly to that which was calculated. This
solution may be standardized against a standard solution of
an alkali or by a gravimetric determination. If, however,
due caution has been taken that everything should be just"
right in the making, the solution maybe accepted as N/10.
Standard Sulphuric Acid Solution.—A convenient way
of making a standard solution of sulphuric acid is to dilute
the concentrated acid to about 30 per cent, at which strength
it possesses the advantage of greater change in specific
gravity with smaller changes in concentration, and is not
hygroscopic. Take the specific gravity and find the per
cent of H2S04 by reference to tables. The requisite amount
of acid in cubic centimeters or grams is taken and diluted
to make the solution of desired strength. It is preferable
to weigh the acid, since errors by this method will be
smaller than by volume measurement.
The amount of acid required in grams is found by the
proportion ; per cent of acid : 100 : : gms. H2S04 required
: x = gms. of acid to weigh out.
The amount of acid to be measured in cc. is found by
6z QUANTITATIVE CHEMICAL AN"AI,Y"SFS.
dividing the number of grams of H2S04 required, multiplied
by 100, by the product of the per cent of the acid and the
specific gravity. Example: 4.9 gms. of H2S04 are required
to make one liter of N/10 solution. If it is to be made from
30 per cent acid, S. G., 1.2228, then, *°°8 = I3-35 cc-
Density and Concentration of Sulphuric Acid at 15°.
(Normal H2S04 = 49*04 grams j>er litre!)
Density
150/4°.
100 grams
contain
grams
H2SO4.
1 litre
contains
grams
H2S04.
Nor¬
mality.
Density
100 grams
contain
grams
H2SO4.
1 litre
contains
grams
H2SO4.
Nor¬
mality.
1-006
1
10*1
0*20
1*449
55
797
16*25
1013
2
20*3
0*41
1*502
60
901
18-38
1*020
3
30*4
0*62
1-558
65
1013
20*65
1*026
4
41*0
0*84
1*615
70
1130
23*05
1*033
5
51*6
1*05
1*674
75
1248
25*60
1*040
6
62*4
1*27
1*732
80
1386
28*20
1*047
7
73 *"3
1*49
1*784
85
1520
30-92
1*054
8
84*3
1*72
1-820
90
1540
33*40
1*061
9
95*5
1*95
1*825
91
1660
33*86
1*068
10
106*8
2*18
1*829
52
1680
34*32
1*104
15
160*6
3*38
1*833
93
1710
34*76
1*142
20
228
4*66
1*886
94
1730
35*20
1*182
25
296
6*02
1*839
95
1750
35*62
1*222
30
367
7*48
1-841
96
1770
36*03
1*264
35
444
9*02
1-841
97
1790
36-42
1*306
40
522
10*66"
1*841
98
1800
36*79
1*851
45
608
12*40
1-839
99
1820
37*13
1*399
50
700
14*26
(1-836)
100
(1836)
(37-4)
Standard Sodium Carbonate Solution.—Sodium carbon¬
ate (Na2C03 = 106) is obtainable in its pure, anhydrous
form, or may.be prepared by heating pure bicarbonate of
sodium to a temperature of 300 0 C. for about 1 hour. Tests
should be applied to the carbonate, or to the bicarbonate if
the latter is to be the source of the carbonate, for the
presence of chlorides and sulphates.
Normal solutions of sodium carbonate contain 53 grams
per liter, and if the solution is carefully made it will not
be necessary to standardize it, and it may be used, if desired,
as a starting solution in acidimetry and alkalimetry.
QUANTITATIVE CHEMICAL, ANALYSIS.
63
Standard Sodium Hydroxide Solution.—With pincets
remove a " stick" of sodium hydroxide from the bottle and
quickly, between pieces of filter paper, break off a length of
about 5 cm. Transfer immediately to a suitable bottle and
add about 1,000 cc. of distilled water, stopper the bottle,
and shake to dissolve the hydroxide. Set aside to cool;
again shake well to mix thoroughly, and titrate against a
standard acid as follows :
Fill a burette with N/io hydrochloric acid and bring the
level of the acid to the zero graduation. Measure from a
second burette, or carefully, with a transfer pipette, 25 cc. of
the sodium hydroxide solution into a 150 cc. Erlenmeyer
flask, being careful that none of it lodges in the neck of
the flask. Titrate immediately with the N/10 acid, using
methyl orange as the indicator.
Solutions of sodium hydroxide must be carefully protected
from the atmosphere of the laboratory, on account of
absorption of carbon dioxide, and should frequently be
checked against a standard acid.
Exercise XVIII.
Determination of A mount of a Base in Solution.
An instructor will furnish you with a solution of a base.
Do not remove from the beaker in which it is given you,
but dilute with distilled water, add a drop of phenol phtha-
lein solution, and titrate with N/10 hydrochloric acid solu¬
tion. Report the amount of potassium hydroxide in the
solution.
Exercise XIX.
Determinatio7i of a Carbonate.
Proceed in general as in exercise XVIII. Titrate with
a N/10 acid, using methyl orange as indicator. Report
sodium carbonate.
64
QUANTITATIVE CHEMICAL ANALYSIS.
Exercise XX.
Determination of a Bicarbonate.
Add to the solution of bicarbonate given you an excess
of sodium hydroxide solution, after which a slight excess of
barium chloride solution is added to precipitate the carbon
dioxide. Add a drop of phenolphthalein solution and titrate
the excess alkali with N/io acid, constantly stirring the
mixture. From the difference between the amount of
alkali added and that found by titration, calculate the amount
of sodium bicarbonate.
Exercise XXI.
Determination of Carbonate and Bicarbonate in Mixed
Solution.
Determine total alkalinity by titration as in Exercise XIX.
Treat a second portion of the solution as in the determina¬
tion of bicarbonate. Calculate the amount of sodium car¬
bonate and sodium bicarbonate in the solution.
Exercise XXII.
Determination of Caustic Soda and Sodium Carbonate in
Mixed Solution.
Add a drop of phenolphthalein to the solution and
titrate with N/io acid until the pink color just disappears.
Record the number of cc. used. Add a drop of methyl
orange to the same solution and continue the titration
until a faint permanent pink color is produced.
In the first titration, using phenolphthalein as indicator,
the NaOH and half the Na2C03 were neutralized according
the equations;
NaOH + HC1 = NaCl + H20 (i)
Na2C03 + HC1 = NaHCOs + NaCl (2).
QUANTITATIVE CHEMICAL ANALYSIS.
65
In the second titration, with methyl orange as indicator,
the bicarbonate formed by the reaction in the first part was
neutralized ;
NaHC03 + HC1 = NaCl + H20 + C02.
The sodium carbonate titration is found by taking twice
the number of cc. used in the second part of the titration,
and this number subtracted from the total number used,
gives the titer of the caustic soda.
Exercise XXIII.
Determination of Acid Strength.
Withdraw 10 cc. of acid from one of the stock bottles
designated by an instructor, dilute to 250 cc., and use por¬
tions of 25 cc. for titration with N/10 caustic soda solution.
Find the specific gravity of the acid in the stock bottle and
calculate its per cent strength.
Chapter VII.
OXIDATION REACTIONS.
N/10 Potassium Permanganate Solution. — Potassium
permanganate, as an oxidizing reagent in the presence of
sulphuric acid, yields free oxygen according to the follow¬
ing equation :
2KMn04 + 3H2S04 = K2S04 + MnSO, + 3H20 + 5O.
5O = 10H, hence, 1/10 the double formula weight of
potassium permanganate (316 X 1/10= 31-6) is the hydro¬
gen equivalent, or the amount in grams which dissolved
in one liter will make a normal solution.
To prepare a N/10 solution, weigh out slightly more
than 3.16 grams of the salt, transfer to a perfectly clean
liter flask, add about 500 cc. of distilled water and dissolve.
5
66
QUANTITATIVE CHEMICAL ANALYSIS.
When solution is complete, add water to the mark, and
mix thoroughly. Allow to stand for about 24 hours and
filter through asbestos. Put into a clean, glass-stoppered
bottle and protect from light.
The solution may be standardized against a solution of
iron wire, ferrous ammonium sulphate, oxalic acid, etc.
Ferrous ammonium sulphate is a very convenient sub¬
stance for the purpose. Weigh out, as nearly as possible,
1 gm. of ferrous ammonium sulphate, dissolve in 25 cc. of
dilute sulphuric acid, and titrate with the potassium
permanganate solution. The end-point of the reaction
will be shown by the pink color imparted by the un¬
changed potassium permanganate.
If exactly 1 gm. of ferrous ammonium sulphate were
used, and if the permanganate solution were exactly N/10,
25.5 cc. of the permanganate solution would have been
required to oxidize the salt. But suppose that 24.7 cc.
were used ; then, 24.7 : 25.5 : 11 : x = 1.0323 N/10 (1 in the
equation standing for the unit strength of the solution, i. e.,
N/10). If the concentration is close to N/10 it is better to
make no attempt to dilute it to exact N/10 concentration,
but put on the label the strength as found. Potassium
permanganate may be used for the determination of several
substances, such as iron, calcium, oxalic acid, oxalates,
peroxides, nitrites, nitrates, etc.
N/10 Potassium Dichromate Solution.—Potassium di-
chromate liberates oxygen according to the equation,
K2Cr207 -f- 8HC1 = 2KCI -f- CrCl3 -f- 4H20 -f- 3O.
3O = 6 H, hence, 1/6 the formula weight of potassium
dichromate (294 X 1/6 = 49) is the hydrogen equivalent,
or the quantity required to make one liter of normal solution.
Make a N/10 solution by dissolving 4.9 grams in a liter
of water in the same general manner as already described
under other solutions.
QUANTITATIVE CHEMICAL ANALYSIS.
67
Potassium dichromate may be purchased of such a high
degree of purity that solutions carefully made, need not be
standardized. However, it may be standardized, if desired,
against ferrous ammonium sulphate as follows : Prepare a
fresh solution of potassium ferricyanide, by washing off a
small crystal to remove superficial ferrocyanide, and dissolve
to a concentration of about 0.1 per cent. Prepare a ferrous
ammonium sulphate solution as in standardizing potassium
permanganate.
Run into the ferrous ammonium sulphate solution, from
a burette, an amount of potassium dichromate solution
which is calculated to be less than required to react with the
iron salt. Remove a drop of the mixture in the beaker with
a clean glass rod to a piece of porcelain. With another clean
glass rod bring a drop of the potassium ferricyanide
indicator into contact with the drop of the mixture; a pre¬
cipitate is formed, showing the presence of ferrous iron.
Continue the titration with the dichromate solution, adding
only very small quantities at the time, and after every
addition make the test with a small drop of the mixture as
before. As the quantity of ferrous iron in the solution
diminishes, the color of the precipitate will pass from blue
to light green, and finally, when the end point has been
reached, there will be no color at all. The first titration
will likely be incorrect and several titrations should be made
to get the correct titer.
Potassium dichromate solutions may be used for the
determination of iron and for the standardization of sodium
thiosulphate solutions. It possesses the following advan¬
tages over potassium permanganate: (1) when the pure
salt is used it is not necessary to standardize the solution;
(2) the solution is stable; (3) it does not attack rubber and
may be used with any type of burette; (4) it may be used
in the presence of hydrochloric acid. Its chief disadvantage
lies in the fact that titrations are somewhat tedious on
account of the necessity for using an indicator outside of
the solution.
68
QUANTITATIVE CHEMICAL, ANALYSIS.
Exercise XXIV.
Determination of Iron in an Ore.
(a) By Potassium Permanganate.—Weigh out from a
weighing bottle about 0.3 gram of the ore into a casserole.
Add 1—3 cc. of 2.5 per cent stannous chloride solution and
20 cc. of hydrochloric acid, S. G. 1.12. Cover the casserole
and heat the mixture on the hot plate until the iron is all
dissolved and the residue is white.
The solution should be yellow ; if it is colorless, add
hydrogen peroxide to destroy the excess of stannous chloride
and to cause the solution to become yellow.
Drops of the solution adhering to the sides of the casserole
and the cover must be washed down into the casserole with
a stream from the wash-bottle.
While the solution is still hot, add drop by drop, very
dilute stannous chloride solution until the iron is reduced
as shown by the solution becoming colorless or greenish.
Dilute to about 50 cc. and add 10 cc. of saturated mercuric
chloride solution, which should give a silky precipitate of
mercurous chloride ; if the precipitate is black, an excess
of stannous chloride is present and the solution cannot be
used.
The titration is to be carried out with potassium per¬
manganate, and it is important that the quantity of hydro¬
chloric acid present be as small as possible, and sulphuric
acid depended upon to give acidity to the solution.
For this purpose, and other advantages, use is made of a
solution made by dissolving 160 grams of manganese sul¬
phate in 1,750 cc. of water, to which are added 330 cc. of
syrupy phosphoric acid, S. G. 1.7, and 320 cc. of sulphuric
acid, S. G., 1.82.
Pour the iron solution into 500 cc. of cold water containing
10 cc. of the manganese solution just described, and titrate
immediately with N/10 potassium permanganate solution.
The end point is reached when the first pink color is pro-
QUANTITATIVE CHEMICAL ANALYSIS.
69
duced which diffuses throughout the whole mixture and
then fades away. Calculate the per cent of iron in the ore.
(b) By Potassium Dichromate.—Prepare the solution of
iron ore as given above as far as the addition of the mercuric
chloride solution. Dilute the solution to 500 cc. with cold
water, add 20 cc. of concentrated hydrochloric acid and
titrate with N/10 potassium dichromate solution, using
potassium ferricyanide solution as indicator as described
under standardization of a dichromate solution.
Questions.
1. In what terms is the concentration of chemical solu¬
tions stated, and which of these is generally used for vol¬
umetric test solutions ?
2. What is a normal solution ?
3. Describe the general procedure in making a solution,
of definite strength of a solid reagent.
4. Describe three ways of finding the specific gravity of
a liquid.
5. How may the specific gravity of liquid reagents be
made use of in making solutions of them ?
6. What is meant by titration ?
7. What is an indicator ? Name and describe five in¬
dicators.
8. Define volumetric equivalent and volumetric factor.
9. Define acidimetry and alkalimetry.
10. Describe a method of making and standardizing a
decinormal solution of hydrochloric acid.
11. What would be the correction factor in terms of nor¬
mality for a hydrochloric acid solution which contains 36
gms. per liter ?
12. Describe the making of an approximate 18 per cent
solution of hydrochloric acid from the concentrated form of
specific gravity 1.2.
13. What precautions are to be taken with solutions of
the caustic alkalies ?
7°
QUANTITATIVE CHEMICAL ANALYSIS.
14. Discuss the determination of sodium carbonate and
bicarbonate in mixed solution.
15. Discuss the advantages and disadvantages of potas¬
sium permanganate and' potassium dichromate solutions.
16. What is the ferrous ammonium sulphate equivalent
of 1 cc. of a potassium permanganate solution containing 3
gins, per liter ?
17. One cc. of potassium permanganate solution is equiv¬
alent to .005 gms. of iron ; what is its calcium equivalent ?
18. How much hydrogen peroxide will be oxidized by 50
cc. of N/15 potassium permanganate?
Chapter VIII.
IODIMETRY.
Iodine is an oxidizer both in acid and neutral solutions,
and with sodium thiosulphate forms a useful means of
making several determinations. Substances which are
oxidized by iodine or reduced by sodium thiosulphate may
be determined by the use of these two solutions. For ex¬
ample : (1) copper acetate, treated with potassium iodide
solution, will precipitate cuprous iodide and liberate iodine
according to the equation,
2 Cu(C2H203)2 + 4KI = Cu2I2 + 4KC2H203 + I2.
The mixture is titrated with sodium thiosulphate solu¬
tion, from which the amount of liberated iodine is found
and the equivalent amount of copper is calculated.
(2) Sulphites are oxidized to sulphates by iodine accord¬
ing to the equation,
Na2S03 + H20 + 2l = Na2S04 + 2HI.
The sulphite solution is run into an excess of iodine solu¬
tion in which oxidation takes place. The unused iodine
remains in solution, and is titrated with sodium thiosulphate
quantitative; chemical analysis.
71
solution. Calculation of the amount of sulphite is made
from the difference between the known amount of iodine
originally present and the residual iodine found by titration.
N/10 Iodine Solution.—Although iodine is very little
soluble in pure water, it is readily soluble in watery
solutions of potassium iodide. Weigh out about 12.8
grams of iodine, put into a liter flask, and add about 25
grams of potassium iodide and 50 cc. of water. Shake the
flask until the iodine is all dissolved and then add water to
the mark. Standardize the solution against a solution of
sodium thiosulphate.
Sodium thiosulphate is oxidized to sodium tetrathionate
by iodine:
2Na2S203 + I2 = Na2S406 + 2NaI.
Starch Indicator.—The indicator for use with iodine
solutions is starch, which gives a blue color in the presence
of iodine. Rub up about 1 gram of starch in a mortar
with sufficient cold water to make a thin paste. Pour this
thin paste into about 200 cc. of boiling water, stirring
meanwhile, and allow to boil for a few minutes; after
which, cool and filter.
N/10 Sodium Thiosulphate Solution.—Sodium thiosul¬
phate, Na2S203.5H20, (molecular weight = 248.2) is ob¬
tainable in such a pure state that an accurate solution may
be made without the necessity of standardization. Dissolve
24.82 grams of the pure crystals in about 300 cc. of water
in a liter flask, and then fill to the mark. The water used
should be boiled before use to drive off any possible carbon
dioxide. Transfer to a suitable bottle and keep well
stoppered.
Standardization of the Iodine Solution.—Measure 25 cc.
of the iodine- solution into an Erlenmeyer flask and dilute
with a moderate amount of distilled water. Run in the
sodium thiosulphate solution from a burette until the
brown color of the iodine has become a pale yellow. Then
72
QUANTITATIVE CHEMICAL ANALYSIS.
add i cc. of the starch solution. The contents of the
flask will now assume a blue color. Continue the addition
of the thiosulphate solution carefully until the blue color
just disappears. The method of calculating the relative
strengths of the two solutions is the same as in other cases
of standardization already given.
Exercise XXV.
Determination of Antimony.
The mineral stibnite, Sb2S3, is analyzed for antimony by
bringing into solution with hydrochloric acid in the presence
of tartaric acid, afterward diluted cautiously, and nearly
neutralized with ammonium hydroxide. It is then treated
with an excess of cold saturated solution of sodium bicar¬
bonate and titrated with iodine solution, with starch solution
as indicator.
The present exercise deals with the determination of
antimony in tartar emetic, K(Sb0)C4H406. Add to the
solution given you about 20 cc. of cold saturated solution
of bicarbonate of sodium and a little starch indicator
solution. Titrate with N/10 iodine solution until a blue
color is produced.
Report the amount of tartar emetic in the solution given
you.
Exercise XXVI.
Determination of Iodine in Tincture of Iodine.
Dilute the portion of the tincture given you with about
20 cc. of potassium iodide solution, add a few drops of starch
indicator solution, and titrate with N/10 sodium thiosulphate
solution until the blue color just disappears. Calculate the
amount of iodine in the amount of the sample examined
by you.
QUANTITATIVE CHEMICAL ANALYSIS.
73
Exercise XXVII.
Determination of Copper.
When a solution of copper is treated with potassium
iodide solution, cuprous iodide is precipitated and free
iodine liberated according to the equation : 2CuS04 + 4KI
= Cu2I2 + I2 -f 2K2S04, and the amount of copper is cal¬
culated from the amount of iodine liberated as indicated by
the number of cc. of sodium thiosulphate solution used in
titration. The titration is carried out in a solution in which
the only free acid permissible is acetic acid.
Add to the solution given you sodium carbonate solution
until a small permanent precipitate is formed, then carefully
add acetic acid until the precipitate is dissolved. Titrate
the mixture with N/10 sodium thiosulphate solution,"using
the starch indicator. Calculate the amount of copper in
the solution given you.
Chapter IX.
SILVER NITRATE AND POTASSIUM SULPHO-
CYANATE TITRATIONS.
N/10 Silver Nitrate Solution.—Silver nitrate (AgN03 =
169.9) decinormal solution contains in 1 liter 16.99 grams of
the pure salt. If it is desirable to standardize the solution it
may be done by titration against a standard solution of
pure sodium or potassium chloride of N/10 concentration.
The indicator is potassium chromate solution. When the
reaction with the chloride is complete the silver immediately
combines with the chromate to form silver chromate which
is of a red color. It cannot be used in acid solution on
account of the immediate decomposition of the silver
chromate.
Silver nitrate solution may be used for the determina¬
tion of the chlorides and bromides of sodium, potassium,
calcium, ammonium and magnesium in neutral solution
74
quantitative; chemical analysis.
with potassium chromate as the indicator. It may also be
used for the determination of cyanides in alkaline solution.
The completion of the reaction in the latter case is shown
by the production of a precipitate.
N/10 Potassium Sulphocyanate Solution.—Potassium or
ammonium sulphocyanate solution may be used in conjunc¬
tion with silver nitrate solution for several determinations in
acid solution, in which precipitates are formed, the indicator
being ferric sulphocyanate.
A measured excess of silver nitrate solution is added to
the solution analyzed. The precipitate which is formed is
filtered off and well washed. The filtrate, which contains
the excess silver solution, is treated with 5 cc. of ferric
nitrate solution (indicator solution) and then titrated with
potassium sulphocyanate solution to a red coloration. The
difference between the amount of silver nitrate solution
added and that found by titration is used to calculate the
amount of the substance analyzed.
Exercise XXVIII.
Determination of Mercury in Mercuric Oxide.
Weigh out accurately about 1 gram of mercuric oxide.
Dissolve in concentrated nitric acid, transfer to a 100 cc.
volumetric flask, and add distilled water to make 100 cc.
Pipette out 25 cc. of the solution, add 10 cc. of concentrated
nitric acid and 5 cc. of the ferric indicator, dilute to 75 cc.,
and titrate with N/10 potassium sulphocyanate solution.
Calculate the per cent of mercury in the sample,
Hg(N 03)2 + 2KSCN = Hg(SCN)2 + 2KNO3.
Exercise XXIX.
Determination of Bromide in a Solution.
Acidify the solution given you with nitric acid and add
a measured quantity of silver nitrate in excess. Stir the
mixture until the precipitate settles. Add the ferric in-
QUANTITATIVE CHEMICAL ANALYSIS.
75
dicator and titrate the filtrate which contains the excess
silver nitrate with N/io potassium sulphocyanate solution.
Calculate the amount of potassium bromide in the solution
given you.
Chlorides in acid solution may be determined in the
same manner as bromides, but it would be necessary to
filter off the precipitate of silver chloride and to wash with
cold water before titrating the filtrate.
Iodides are also estimated by the same method, with
the modification that the silver nitrate solution is added in
small quantities at the time, 1-2 cc., and the mixture
shaken vigorously after each addition. It is not necessary
to filter off the precipitate.
Chapter X.
COLORIMETRIC METHODS.
Some substances in solution produce very strong colora¬
tion, and even in very dilute solutions where there may be
only a trace of the substance present, the coloration may be
still discernible. The red coloration produced by ammonium
sulphocyanate with ferric iron is perceptible when the iron
in solution amounts to no more than .0001 gm. per liter.
In carrying out a colorimetric determination the solution
to be examined, if not already very dilute, is made so, and
a known volume of it treated with the color-producing
reagent in solution. A standard solution of a salt of the
kind being estimated is brought to the same volume as the
solution tested, in a series of tubes, each containing a dif¬
ferent amount of the salt. These tubes are treated with
the reagent in the same manner as the solution tested, and
are used for comparison of the color produced in the solu¬
tion tested. Since the amount of salt in each tube is
known, when a coloration is found which matches that of
the tested solution, it is taken that they contain equal
amounts of the salt.
76
QUANTITATIVE CHEMICAL, ANALYSIS.
Exercise XXX.
Colorimetric Determination of Ammonia.
Ammonia in water usually results from contamination
by sewage. Therefore, the examination of water for ab¬
normal quantities of nitrogen is one of the important points
in the sanitary analysis of drinking water. Ammonia pre¬
sents itself in two forms, (a) free ammonia, which is made
evident by a brown color with Nessler's solution ; (b)
" albumenoid ammonia" which is liberated by the action of
potassium permanganate in basic solution.
The reagents required for the determination of ammonia
in water are:
1. Ammonia-free Water.—Redistil distilled water acid¬
ified with sulphuric acid, rejecting the first portions of the
distillate.
2. Nessler's Solution of Mercurio-Potassium Iodide.—
(a) Prepare a saturated solution of mercuric chloride by
boiling an excess of the salt in water; (b) dissolve 25 gms.
of potassium iodide in a little water; (c) make a 50 per
cent solution of potassium hydroxide. Add a to b until a
slight but permanent precipitate remains; add 200 cc. of
the potassium hydroxide solution, and dilute the whole to
500 cc. Allow to stand so that the liquid may become
clear before using.
3. Alkaline Potassium Permanganate Solution. — Dis¬
solve 100 gms. of potassium hydroxide and 4 gms. of
potassium permanganate in 500 cc. of distilled water, and
remove 125 cc. by distillation, then dilute to 500 cc. with
ammonia-free water.
4. Standard Ammonium Chloride Solutions.—(a) Dis¬
solve 1.572 gms. of pure ammonium chloride in 500 cc. of
ammonia-free water. This solution will contain 1 milli¬
gram of NH3 per cc. (b) dilute 5 cc. of a to make 500 cc. ;
this solution will contain .01 milligram of NH3. These
two solutions are to be used for color comparisons.
quantitative; chemical analysis.
77
Connect a distilling flask or a tubulated retort with a
Iyiebig condenser and remove all traces of ammonia by
boiling distilled water in the apparatus until the distillate
shows no traces of color with Nessler's solution. Then
put into the distilling flask 500 cc. of the water to be
examined and distil at the average rate of about 8 cc. per
minute, collecting 200 cc. of the distillate in a 200 cc.
volumetric flask.
Remove the 200 cc. flask and put in its place a 150 cc.
volumetric flask. Introduce into the distilling flask 50 cc.
of alkaline potassium permanganate solution, and continue
the distillation until the flask is filled to the mark.
The 200 cc. of distillate first collected contains the free
ammonia ; the second distillate contains the " albumenoid"
ammonia.
Prepare a series of 12 Nessler tubes containing in cc. the
following amounts of standard ammonium chloride solution :
0.2, 0.4, 0.6, 0.8, 1.2, 1.4, 1.7, 2.0, 2.5, 3.0, 4.0 and 5.0.
Bach cc. of standard ammonium chloride solution represents
.00001 gm. of ammonia.
Dilute the contents of each tube to 50 cc., add 2 cc. of
Nessler's solution and place in a Leeds comparator.
Divide the 200 cc. of distillate into 4 tubes of 50 cc. each
and dilute with ammonia-free water; treat with 2 cc. of
Nessler's solution, and estimate the amount of ammonia in
each by comparison with the standard tubes just prepared.
Record the readings of each, and finally add together all
the readings to find the total ammonia.
Divide the distillate from the second distillation into
three tubes of 50 cc. each and determine ammonia as above.
Report " albumenoid ammonia."
78
QUANTITATIVE CHEMICAL ANALYSIS.
SECTION IV.
Chapter X.
MISCELLANEOUS ANALYSES.
Exercise XXXI.
The Proximate Analysis of Coal.
Coal is a carbonaceous material containing certain com¬
bustible compounds, volatile at and somewhat below red heat,
" fixed" carbon which does not ignite below red heat, inor¬
ganic salts, and some moisture. The common commercial
forms are the anthracite and the bituminous, to which the
terms " hard" and " soft" are respectively applied. Anthra¬
cite coal differs from bituminous chiefly in that the former
contains much less of volatile constituents than does the
latter, and is chiefly "fixed" carbon.
The determination of the components set forth above
constitutes the proximate analysis of coal as distinguished
from the ultimate analysis, in which determination is made
of the elementary constituents. The analysis of coal fur¬
nishes a basis for ascertaining the suitability of any par¬
ticular lot of coal for a given purpose, or for the determina¬
tion of its fuel value. The fuel value as calculated from
the results of a proximate analysis is usually found by the
use of empirical formulas which, while not always express¬
ing the true value, serve very well for ordinary purposes.
The real value of any coal as a fuel is found by burning a
weighed portion of the sample in a calorimeter, from which
the number of heat units evolved may be found.
This exercise is concerned with the determination of (i)
moisture; (2) volatile combustible matter; (3) coke; (4)
fixed carbon and ash.
The coal to be analyzed has been sampled and prepared
quantitative; chemical analysis.
79
for analysis. Reference should be made to larger text¬
books or to special works on coal analysis for the details of
sampling, etc.
Moisture.—Weigh out ten grams of the sample, transfer
to a weighed watch-glass, and dry in the oven for one hour
at no0 C. Cool in the desiccator and weigh. Record the
weight, and subtract from first weight to find moisture.
Volatile Combustible Matter and Coke.—Weigh out one
gram of the sample 'into a weighed crucible and heat for
seven minutes over a Tirrill burner regulated to give a full,
free flame, the bottom of the crucible being 5 or 6 cm. above
the flame. Protect the apparatus from draughts of air.
Cool in the desiccator and weigh. Subtract the weight of
the residue from the first weight of the coal. This will
give the total volatile matter (including moisture) the per
cent of which must be subtracted from the per cent of total
volatile matter to find the combustible volatile matter. The
residue in the crucible is coke, i. e., the fixed carbon and
ash.
Fixed Carbon and Ash.—Place the crucible containing
the coke on a triangle and proceed to burn off the carbon
until no trace of it remains; the ash, however, will not be
white on account of the presence of iron oxide. Cool in
the desiccator and weigh to find the ash content, which,
subtracted from the weight of coke previously found, will
give the amount of fixed carbon. Report all values in per
cents.
Calculation of Fuel Value.—Calculate the fuel value
from the formula of Goutal, 82C + AM = calories per
gram. C = per cent of fixed carbon ; M = volatile com¬
bustible matter; A = coefficient whose value is fixed by
the volatile combustible matter as follows :
When M = 2 — 15, A = 130
15 — 3° 100
30 — 35 95
35 — 4o 9°-
80 QUANTITATIVE CHEMICAL ANALYSIS.
Exercise XXXII.
Examination of Burning Oils.—Kerosene.
Common burning oils are derived from petroleum by
fractional distillation and consist of mixtures of the liquid
hydrocarbons of the methane series, and are usually defined
by their boiling points, or their specific gravities. The
examination of fuel oils has to do more with certain physical
properties rather than with their chemical constituents, the
purpose being to determine their fitness for the use intended,
and usually includes, (i) specific gravity; (2) flash-point;
(3) burning point.
Specific Gravity.—Take the specific gravity of the sam¬
ple at 6o° F. (15.50 C.) with a Westphal balance or
with a hydrometer. In commercial practice the specific
gravity of oils is given in Baume degrees, to which, if
desired, the true specific gravity may be translated. To
convert lighter-than-water densities to Baume degrees, and
vice versa, use the following formulas: B = — 136 ;
o. (J.
S. G. = ^ = -^aum^ ' S. G.= specific gravity).
Flash Point.—The flash point of a burning oil is the
temperature at which it will give off an inflammable vapor.
The safety with which an oil may be used in the ordinary
oil-burning utensils of the household is determined princi¬
pally by its flash-point. The matter is of such importance
that the lower limit of the flash point of oils allowed to be
sold in a commonwealth is fixed by state law. This lower
limit of flash point varies in different places.
The form of oil tester known as the New York State
Board of Health pattern is conveniently used. The outer
shell supports a cup which contains water, into which is
placed an inner cup to contain the oil to be tested. The
oil cup is provided with a glass cover in which is an opening
through which a thermometer may be suspended in the oil
QUANTITATIVE CHEMICAL ANALYSIS. 81
and a larger opening through which the flame is to be
applied, and which is normally kept closed, and is un¬
covered when the flame is to be passed over the heated oil.
Set up apparatus as indicated in the paragraph above.
Light the burner and bring the temperature of the oil up
to 90° F. (32.2° C.), then turn burner very low and slowly
bring the temperature up to 950 F. (350 C.). Remove the
cover from the aperture in the glass cover and pass a lighted
taper over the opening. If there is no flash, carefully
bring the temperature up to ioo° F. (37.70 C.) and make
the test with the flame as before. Continue to make the
test in this way until the temperature at which the oil
flashes is found. The test should be repeated after allowing
the oil to cool somewhat.
Burning Point.—The burning point is the temperature
at which the vapor is given off so rapidly that the flame
will be continuous rather than a mere flash. Find the
burning point by continuing the experiment as for flash
point, until the temperature is reached at which there is a
continuous flame.
Exercise XXXIII.
Determination of Nitrogen.—Kjeldahl Method.
This exercise deals with the determination of nitrogen
by the Kjeldahl method. It depends on the decomposition
of the nitrogenous substance by sulphuric acid with the
formation of ammonia compounds, which are afterward
decomposed by caustic soda, distilled into a measured volume
of standard acid solution, and the excess of standard acid
titrated with standard alkali solution.
1 cc. of N/10 acid = 0.0017 gm. of NH3 = 0.0014 gm.
of N. The number of cc. of acid used by the ammonia is
found by subtracting from the original volume added the
excess number as shown by the titration with the standard
alkali solution. This number multiplied by the factor for
nitrogen, 0.0014, gives the amount of nitrogen in the sam-
6
82
QUANTITATIVE CHEMICAL ANALYSIS.
pie analyzed. If the result is to be expressed as protein
substance the nitrogen content is to be multiplied by the
factor 6.25. The factor 6.25 is derived from the amount
of nitrogen contained in protein substances as found by
analysis. Sometimes the factors 6.33, 6.37 and 6.38 are
used.
Use for this exercise the sample of milk employed for fat
determination. The reagents required are (1) concentrated
sulphuric acid ; (2) potassium sulphate; (3) mercury or
mercuric oxide ; (4) 4 per cent potassium sulphide solution ;
(5) saturated solution of caustic soda ; (6) cochineal indica¬
tor solution ; (7) N/10 acid solution; (8) N/10 alkali solu¬
tion.
Pour 5-10 cc. of the milk into a beaker and weigh, then
introduce as much as possible of the milk into the Kjeldahl
flask. Again weigh the beaker, the difference being the
weight of the milk transferred to the flask. Put into the
flask 25 cc. of concentrated sulphuric acid and 0.7 gm. of
mercuric oxide. Heat gently until there is no longer any
frothing, then heat more strongly for about ten minutes.
The flask should be inclined on its support, and a small
funnel placed in its mouth. Remove the flame and allow
the flask and contents to cool, and then add 10-12 gms. of
potassium sulphate. Again apply heat and keep the mix¬
ture boiling briskly until the solution becomes colorless.
Continue boiling for some time after the solution becomes
colorless or a very light straw color. Allow to cool, add
200 cc. water, and then precipitate the mercury by the
addition of 25 cc. of the potassium sulphide solution. Put
into the flask a few pieces of pumice stone and a piece of
paraffin about the size of a pea.
Set up the apparatus for distillation as shown in Fig. 15,
measuring 50 cc. of N/10 acid into the beaker. Now, pour
50 cc. of the saturated caustic soda solution into the flask
in such a manner that it does not at once mix with the
solution, by inclining the flask and pouring in the alkali
QUANTITATIVE CHEMICAL ANALYSIS.
83
carefully. Immediately connect with the condenser and
gently mix the whole. Distil the contents of the flask into
the beaker until 150 cc. of distillate have passed over. Then
disconnect condenser and rinse with a little distilled water,
and also wash off the adapter inside and outside, the rinsings
and washings being allowed to flow into the beaker. Ti¬
trate the contents of the beaker with N/10 alkali solution,
using cochineal as the indicator. Calculate nitrogen and
protein substance as indicated above.
Exercise XXXIV.
Determination of Fat in a Food.
The food product to be analyzed is first dried and its fat
then extracted with anhydrous ethyl ether or petroleum
ether into a weighed flask. The solvent is distilled off, the
84
QUANTITATIVE CHEMICAL ANALYSIS.
fat remaining in the flask which is again weighed. The
difference between the last weight and the weight of the
flask alone is the weight of the extracted fat.
For this exercise use milk. Make a coil of a strip of fat-
free filter paper about 55 cm. long and 6.25 cm. wide, and
fasten with a thin wire.
Pour into a beaker about 5 cc. of milk and weigh quickly.
Then place one end of the paper coil in the milk and ab¬
sorb as much as possible of it; it should be practically all
absorbed. Remove the paper coil and again weigh the
QUANTITATIVE CHEMICAL ANALYSIS.
85
beaker. The difference between the first weight and the
last weight is the weight of milk absorbed. Place the paper
coil, with the wet end uppermost, in a clean beaker and dry-
in the oven for two or three hours.
When the coil has thoroughly dried, place it in a Soxhlet
extraction apparatus, illustrated in Fig. 16, the end of the
coil which was used to absorb the milk being placed down¬
ward. Pour anhydrous ether into the extractor until it
siphons over ; then connect up the apparatus for the deter¬
mination. Be sure that the apparatus is tight and that
water is flowing through the condenser before applying
heat. It is better to use an electric hot plate for heating,
though, with caution, a safety water bath may be used.
Isolate all open flames from the vicinity of any experiment
in which ether is used. Immerse the extraction flask in a
beaker of water and carefully regulate the temperature so
that the extractor fills and siphons about every 10 to 12
minutes. Continue the extraction for about three hours.
Then remove the paper coil, but connect the apparatus again
and continue heating until nearly all the ether has been
distilled from the flask and collected in the extractor. The
ether is to be returned to the ether can, and the flask placed
in the oven and dried to a constant weight. The difference
between the weight of the flask after and before the extrac¬
tion is the weight of fat. Calculate the per cent of fat in
the sample of milk examined.
Exercise XXXV.
Determination of Alcohol.—By S. G. Method.
The per cent of alcohol in a liquid may be found by
distilling off from a measured quantity of the sample, and
after diluting with water to a volume equal to the volume
of the sample taken, the specific gravity of the water-alcohol
mixture is determined, and by reference to " alcohol tables"
the per cent of alcohol is found.
86
QUANTITATIVE CHEMICAL ANALYSIS.
Alcoholic liquids which may contain volatile acids should
be treated with caustic soda solution before being distilled.
Measure out 100 cc. of the sample to be analyzed, introduce
into a distilling flask and add 25 cc. of 5 per cent caustic
soda solution. Distil over 75 cc. and transfer to a 100 cc.
flask. Fill to the mark with distilled water, mix well and
find the specific gravity as follows :
Weigh a perfectly clean and dry specific-gravity bottle
like that shown at C, Fig. 14. Fill the bottle at i,5°C with
the alcohol-water mixture and insert the stopper. With a
piece of filter paper remove all drops of liquid adhering to
the outside of the bottle and wipe across the top of the
stopper. Do not hold the body of the bottle in the hand
as the warmth will cause the liquid to expand and issue
from the capillary bore in the stopper. Weigh rapidly.
Record the weight of bottle and contents, and then in the
same way and at the same temperature weigh the bottle
full of distilled water. From each of the two weights
obtained subtract the weight of the empty bottle. Divide
the weight of the alcohol by the weight of the water to
find the specific gravity of the alcohol-water mixture. Con¬
sult the alcohol tables to find the per cent of alcohol cor¬
responding to the specific gravity found. Calculate the
per cent of alcohol in the sample.
INDEX.
Acid, determination of strength of, 65.
Alcohol, determination of, 85.
Alkalimetry, 59.
Ammonia, albumenoid, determination of, in water, 77.
free, determination of, in water, 77.
Ammonia-free water, 76.
Ammonium chloride, standard solution, 76.
Analytical equations, balancing, 21.
Analytical factors, 22.
Anthracite coal, 78.
Antimony, determination of, 72.
Ash, determination of, in coal, 79.
Balance, construction and use of, 2.
sensibility of, 6.
Westphal, 55.
Base, in solution, determination of, 63.
Baum6 degrees, to convert to specific gravity, 80.
Bicarbonate, determination of, 64.
Bituminous coal, 78.
Blast burner, 75.
Bottle, specific gravity, 54.
Bromides, volumetric determination of, 73, 74.
Bunsen burner, 35.
Burettes, 10.
calibration of, 14.
reading, 11.
Burning oils, examination of, 80.
Burning point of oils, 81.
Calculations, based on equations, 21.
volumetric, 58.
Capacity, units of, 12.
Carbon dioxide, determination of, 45, 47.
Carbonate, determination of, 63.
and bicarbonate in mixed solution, 64.
and caustic soda in mixed solution, 64.
Care of stop-cocks, 12.
Caustic soda solution, 63.
Chlorides, determination of, 42.
Chromate, silver, as indicator, 73.
88
INDEX.
Cleaning solution, 12.
Coal, proximate analysis of, 78.
Cochineal, 57.
Coke, determination of, in coal, 79.
Colorimetric methods, 75.
Combustible matter, volatile, in coal, 79.
Cooling precipitates, 36.
Copper, determination by electrolysis, 51.
volumetric determination of, 73.
Crucibles, 29, 32.
Current, electric, density, 49.
regulation of, 49.
source of, 49.
Desiccator, 36.
Drying precipitates, 31.
Electric current, control of, 49.
source of, 49.
Electrodes, 50.
Electrolytic methods, 48.
Equations, analytical, balancing, 18.
metathetical, 18.
oxidizing, 19.
Equivalents, volumetric, 58.
Ether, caution in use of, 83.
Factor, correction, 56.
Factors, analytical, 22.
volumetric, 58.
Fat, determination of, in a food, 83.
Ferrous ammonium sulphate, 66.
Filter paper, 28.
Filtration, 27.
by suction, 28.
with Gooch crucible, 29.
Flash point of burning oils, 80.
Flasks, volumetric, 10.
calibration of, 14, 15.
Fuel value of coal, calculation of, 79.
General operations, 25.
Gooch crucible, preparation of, 29.
Graduates, 8.
Gravimetric analysis, general operations of, 25.
exercises, 37.
Hydrochloric acid, gravimetric determination of, 42.
standard solution, 60.
table of properties, 60.
Hydrogen equivalent, 53.
Hydrometer, 50.
INDEX.
Ignition, 32.
of moist precipitates, 34.
Indicator, definition, 1, 57.
Indicators, 57.
Iodides, determination of, 75.
Iodine, determination of, in tincture of iodine, 72.
standard solution, 71.
standardization of, 71.
Iodimetry, 70.
Ionization, 27.
constant, 27.
Iron, determination of, by gravimetric method, 37.
potassium dichromate, 69.
potassium permanganate, 68.
Kjeldahl method of nitrogen determination, 81.
Lead, determination of, 38.
Litmus, indicator, 57.
Logarithms, use of, 22.
Magnesium, determination of, 40.
Measuring vessels, 8.
M£ker burner, 35.
Mercury, determination of, gravimetric, 39.
volumetric, 74.
Metathetical equations, balancing, 18.
Methyl orange, indicator, 57.
Milk, determination of fat in, 83.
determination of nitrogen in, 81.
Miscellaneous analyses, 78.
Moisture, determination of, in coal, 79.
Molybdate solution, 43.
Morse-Blalock bulbs, 15.
calibration with, 16.
standardization of, 16.
Nessler's solution, preparation of, 76.
tubes, 77.
Nitrogen, determination of, by Kjeldahl method, 81.
Ohm's law, 49.
Operations, general, 25.
Oxalic acid, standard solution of, 59.
Oxidation, defined, 19.
Oxidizing reactions, 19.
Phenolphthalein, indicator, 57.
Phosphomolybdate solution, 43.
Phosphoric acid, determination of, 43, 44.
Pipettes, 10.
calibration of, 13, 14.
9°
INDEX.
Potassium dichromate, indicator for, 67.
standard solution of, 66.
Potassium ferricyanide, indicator, 67.
permanganate, alkaline solution, 76.
standard solution, 65.
Precipitates, drying of, 31.
transferring to filter, 29.
washing of, 27.
weighing, 36.
Precipitation, 26.
Preparation and weighing of sample, 25.
Questions, 17, 24, 37, 69.
Results, stating, 22.
Rider, use of, 4.
Rose's method for carbon dioxide, 47.
Rosolic acid, indicator, 57.
Rules for using the balance, 3.
Sensibility of a balance, defined, 6.
determination of, 6.
Silver, determination by electrolysis, 50.
chromate as indicator, 73.
nitrate, standard solution, 73.
and potassium sulphocyanate titrations, 73.
Sodium bicarbonate, determination of, 64.
carbonate, determination of, 63.
hydroxide, standard solution, 63.
thiosulphate, standard solution, 71.
Solubility product, 27.
Solution, 26.
Solutions, normal, 52.
rule for making, 52.
volumetric, making, 52.
Specific gravity, formula for converting to Baum£ degrees, 80.
use of, in making solutions, 23, 54.
Starch indicator, 71.
Stating results, 22.
Stoichiometry, 18.
Suction, filtration by, 28.
Sulphuric acid, determination of, 41.
specific gravity, table for, 62.
standard solution of, 61.
Tables, chemical, 22.
Temperature, standard for volumetric analysis, 13.
Tirrill burner, 34.
Units of capacity, 12.
Valance, change of, in oxidation and reduction, 20.
INDEX.
Volatile combustible matter in coal, determination of,
Voltage of current, control of, 49.
Volumetric analysis, 52.
general principles of, 52.
Volumetric equivalent, 58.
Washing precipitates, 27.
Water, free and albumenoid ammonia in, 76.
ammonia-free, 76.
Weighing, 7.
Weight of an object, determination of, 7.
Weights, construction and use, 3.
Westphal balance, 55.
Zero point of a balance, definition, 4.
determination of, 5.