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LABORATORY
OUTLINES
IN
PHYSICAL
CHEMISTRY
BY
T. R. BRIGGS
ITHACA, NEW YORK
1920
LABORATORY
OUTLINES
IN
PHYSICAL
CHEMISTRY
BY
T. R. BRIGGS
ITHACA, NEW YORK
1920
LABORATORY OUTLINES
IN PHYSICAL CHEMISTRY '
INTRODUCTION
It is the purpose of this course to acquaint the student with some
of the factors governing physical and chemical change and to enable
him to recognize these factors and to measure their intensity by their
effects. Painstaking accuracy is not required in most of the experi-
ments, which have been designed primarily to illustrate principles
and to encourage intelligent thinking. It is believed that work of
this kind proves more interesting and stimulating to the average
student than do the more tedious and exact measurements carried
out commonly in laboratories of physical chemistry.
Completion of the work of this course entitles the student to three
hours of University credit per term, two of which are given for
experiments performed satisfactorily in the laboratory and one for
written reports based upon these experiments. The following
Laboratory Outlines describe the work to be accomplished during
the year, though certain of the experiments may be omitted at the
discretion of the Professor in charge. No work of a similar nature
done elsewhere at another college or university is required to be
repeated, provided the work be submitted to the Professor in charge
for his approval.
LABORATORY
In performing the majority of these experiments, students are to
work in groups of two. Partners are to be chosen at the first labora-
tory period and this partnership is to be maintained throughout the
year so far as possible. It is absolutely essential, however, that both
partners work in cooperation on the same experiment. Independent
work on different experiments in a given group will not be permitted.
Since this course is introductory in nature, the student is not given
the most delicate instruments or the purest materials. The apparatus
supplied will nevertheless be found quite sufficient for the require-
ments of these experiments. When the student has determined how
closely his calculations must be made, he can readily ascertain the
allowable error, how carefully his measurements must be made and
what degree of delicacy he must look for in his measuring instruments.
All burettes and pipettes should be calibrated according to the
methods of Experimental Group I and should be cleaned in chromic-
1
sulf)lno-40 nyi<l 'mxaMV before using. The use of dirty or poorly
assembled apparatus will not be tolerated.
All special apparatus must be returned clean and dry and should
never be locked away in a desk except by special permission.
Little attention is given in the lectures in Physical Chemistry
(Course 50) to the methods of experimental physical chemistry.
Reference should therefore be made constantly to the Laboratory
Manuals and to other reference books in the Chemical Library.
Before commencing work on any experiment, the directions should
be read and a clear idea of the principle involved should be obtained.
The student should supply himself with a suitable laboratory
notebook in which his own observations are to be neatly recorded at
the time of performing the experiment. Recording observations on
loose sheets of paper will not be permitted. Notebooks are to be
submitted to the Instructor for approval before entering upon work
in this course.
When making measurements, the student is urged to compute the
results so far as possible in the laboratory at the time the work is
being done and, if feasible, to plot rough curves on cross-section
paper. On the completion of each Experimental Group, the labora-
tory notes are to be submitted to the Instructor for inspection and
approval before writing the final report. No report will be accepted
unless this is done. .
REPORTS
Each report should include a description and discussion of all work
completed in the laboratory together with answers to all questions
and problems appearing in the Laboratory Outlines. Reports
should be written in ink and on one side of the paper only, and
should be enclosed in a "Department of Chemistry" cover. Care
should be taken to describe the experiments in the order in which
they appear in the Laboratory Outlines.
In writing the reports, the general outline given below should
be followed:
(1) Purpose of the experiment and theory illustrated.
(2) Apparatus and manipulation.
(3) Experimental data and curves.
(4) Discussion.
At the time of inspecting the laboratory data, the Instructor will
assign a date on which the written report is due. A deduction of 2
per cent per diem will be made for unexcused lateness in submitting
reports. All reports are to be handed in on or before the day of the
final examination in Course 50. After inspection the reports will be
returned to the student. If "double checked" the report is accepted
as written. If "single checked" it is returned for correction and
should be resubmitted with corrections not later than one week after
its return. When a report is received by the Instructor he will make
a note to that effect on the Bulletin of Reports posted in the labora-
tory. Students are requested to consult this bulletin and to notify
the Instructor of any mistakes or omissions.
A term grade of "Incomplete" will be given in Course 51 if at the
end of the term all the reports have not been handed in and accepted.
STANDARD REFERENCES IN PHYSICAL CHEMISTRY
General Texts. Abbreviation
Arrhenius: Theories of Chemistry (1907) Arrhenius
Bigelow: Theoretical and Physical Chemistry (1912) Bigelow
Getman: Outlines of Theoretical Chemistry (2d ed. 1918) Getman
Hildebrand: Principles of Chemistry (1918) Hildebrand
Jones: Elements of Physical Chemistry (4th ed.) 1915 Jones
Kremann (Potts) : Application of Physico-Chemical Theory (1913)
Kremann-Potts
Lehfeldt: A Textbook of Physical Chemistry (1899) Lehfeldt
Lewis: A System of Physical Chemistry, 3 vols. (1916-1918) Lewis
Lincoln: Physical Chemistry (1918) Lincoln
Nernst (Tizard): Theoretical Chemistry (7th ed. 1916) Nernst
Ostwald: Lehrbuch der allgemeinen Chemie (1891-1902) Lehrbuch
Ostwald (Morse): The Fundamental Principles of Chemistry (2d
ed. 1917) OFF
Ostwald (Walker and Taylor) : Outlines of General Chemistry (2d
ed. 1912) OO
Senter: Outlines of Physical Chemistry (1911) Senter
van't Hoff (Lehfeldt): Lectures in Theoretical and Physical
Chemistry (1898) VHL
Walker: Introduction to Physical Chemistry (8th ed. 1920) Walker
Washburn: Principles of Physical Chemistry (1915) Washburn
Laboratory Manuals.
Biltz (Hall, Blanchard) : Laboratory Methods of Inorganic Chemistry
(1909) Biltz
Biltz (Jones, King) : Practical Methods of Determining Molecular
Weights (1899) BJK
Ewell: Physical Chemistry (1909)
Findlay: Practical Physical Chemistry (1917) F
Getman: Laboratory Exercises in Physical Chemistry (1908) G
Gray: Manual of Practical Physical Chemistry (1.914)
Lamb: Laboratory Manual of General Chemistry (1916) Lamb
Ostwald (Walker) : Physico-Chemical Measurements (1894) OW
Stabler: Arbeitsmethoden usw, 3 vols. (1913) Stabler
Traube (Hardin): Physico-Chemical Methods (1898) T
Physical and Chemical Tables.
Biedermann: Chemiker Kalender (annual) Kalender
Castell-Evans: Physico-Chemical Tables (1902)
Landolt-Bornstein-Roth : Tabellen (1912) LBR
Tables Annuelles de Constantes (1910)
Methods of Calculation Problems.
Knox: Physico-Chemical Calculations (1916)
Mellor : Higher Mathematics for Students of Chemistry and Physics
(1902)
Partington: Higher Mathematics for Chemical Students (1911)
Prideaux: Problems in Physical Chemistry (1912) Prideaux
3
journals.
Zeitschrift fur physikalische Chemie (1887) Zeit. Phys. Chem.
Journal of Physical Chemistry (1896) Jour. Phys. Chem.
Journal de Chimie Physique (1903 ) Jour. Chim. Phys.
Journal of the American Chemical Society (1879 )
Jour. Am. Chem. Soc.
Journal of the Chemical Society (London) (1849 )
Jour. Chem. Soc.
Abstract Journals.
Abstract Journal of the American Chemical Society (1907 )
Abstracts of the Journal of the Chemical Society of London
Chemisches Centralblatt (1856)
Science Abstracts (Chemistry and Physics)
PRELIMINARIES
1. Check apparatus in desk.
2. Make wash bottle. Use 1000 cc. flask in desk.
3. Prepare cleaning mixture as follows:
Dissolve 50 grams of powdered commercial Na2Cr2O? in about 200
cc. of warm water. After cooling this solution, add to it, slowly and
with constant stirring, 300 cc. of concentrated H2SO4 (commercial).
Keep in a 500 cc. wide mouth bottle, for cleaning grease from glass
vessels.
EXPERIMENTAL GROUP I
CALIBRATION OF VOLUME MEASURING APPARATUS
The following group of experiments is designed to give practice in
testing and calibrating the volume measuring apparatus, supplied to
you in your equipment. For accurate work apparatus as supplied
by the maker should never be regarded as correctly graduated unless
accompanied by the certificate of the United States Bureau of Stand-
ards or of the German Reichsanstalt.
Discussion.
The best method of procedure is to take a liquid whose specific
volume is known accurately and, completely filling with it the appara-
tus to be tested, to determine the weight of the liquid either contained
or delivered. In most cases the liquids chosen are water and
mercury.
Since bodies usually expand on being heated, it is necessary in
calibrating to make the determinations at the same temperature
as that at which the apparatus is to be used. Instead of doing this,
however, one may calculate the volume changes due to temperature
variations and may introduce the necessary corrections. Such cor-
rections are absolutely essential when the volume of the apparatus
is large (flasks, etc.).
For accurate work the calibrating liquid must be pure and its
surface free from contaminating impurities affecting the surface
tension and hence the shape of the meniscus. All volumes are to be
read from the meniscus, using a suitable background (white or
black).
References.
Read Bulletin U. S. Bureau of Standards, 4, 553 (1908) or an
abstract f this article in Mahin: Qualitative Analysis, 140 (1914).
Note carefully units of capacity; milliliter; Mohr units; parallax
and its avoidance; cleaning apparatus; error due to surface con-
tamination; outflow time and drainage; limit of error for burette;
tables for calculation, etc. Cf. also Foulk: Quantitative Analysis,
79 (1910); OW,82; F, 29.
EXPERIMENT 1
Calibration of Burettes
Calibrate a 50 cc. burette following the procedure recommended
by Richards. The following is quoted from the original article by
Richards: Jour. Am. Chem. Soc., 22, 149 (1900).
"In the original description of this process it is assumed that the
calibrator delivers exactly an integral number of cubic centimeters,
but if a few instruments only are to be calibrated, it is both trouble-
some and expensive to secure such a precise instrument. We have
found it convenient to use a calibrator of any size, and in parallel
columns to compare its multiples with the actual readings of the
burette. The capacity of this calibrator is most conveniently
obtained in the following manner: Suppose that as a, mean of
several comparisons it has been found that sixteen fillings of the
calibrator correspond to 49.53 cc. on a given burette, . . . The
burette is now refilled and exactly this amount of pure water is run
into a weighed flask, with all the precautions which would be used
in an actual titration. The weight of the water gives by appropriate
calculation the true volume of sixteen fillings of the calibrator.
Suppose this was found to be 49.44 cc. ; then the volume of the cali-
brator as it is actually used in a calibration must be
ID
The differences between the successive readings of the burette and
the successive numbers, 3.09, 6.18, 9.27, . . . etc., give at once
the errors of the graduation of the tube at these intervals. These
differences or corrections may be plotted on a diagram in which the
ordinates are volumes and the abscissas corrections. The correction
to be applied for 50 cc. is obviously -0.09 cc."
Notes. Allow the burette to drain for two minutes before making
a reading. See precaution 23 below under Expt. 2. Clean the burette
with cleaning mixture until the "film of water wetting the interior,
will remain continuous for at least five minutes" (Bureau of Stand-
ards requirement). Results are of no value if grease is present.
Reduce all weights to weights in vacuo.
EXPERIMENT 2
Calibration of Pipettes
Calibrate a pipette to deliver 10 cc. at room temperatures
(18-25). Follow the procedure described in laboratory manuals
suchasF,32; OW,84.
Notes.
To secure uniform delivery in case of burettes and flasks see Pro-
ceedings of American Chemical Society, 21 (1904).
"Certain precautions will be taken to secure uniform delivery.
"18. All such -apparatus will be made so clean internally that the
film of water wetting it will remain continuous for at least five
minutes.
"21. Pipettes with one mark will be held vertical with the delivery
orifice touching the side of the receiving vessel during the free outflow
and for fifteen seconds thereafter.
"23. From burettes, after the desired volume shall have been
taken, the suspended drop will be removed with a glass rod and the
reading will be taken at the end of two minutes."
6
Note carefully that the delivery orifice of a pipette must be of such
a size that the free outflow shall last not more than two minutes and
not less than
12 seconds, if capacity is not more than 10 cc.
15 " " " lies between 10 and 50 cc.
20 " " " " 50 and 100 cc.
30 is more than 100 cc.
EXPERIMENT 3
Morse-Blalock Bulb and Flask
Calibrate a Morse-Blalock bulb and flask, the bulb to deliver
exactly 500 cc. at 20, the flask to hold 500 cc. at 20 C.
References.
Morse: Exercises in Quantitative Chemistry, 84 (1905) ; Mahin:
Quantitative Analysis, 155 (1914) ; also article in Am. Chem. Jour.
16, 479 (1894) or in Olsen: Quantitative Analysis, 236 (1910).
The following table will be found very helpful in calibration work.
In it is given the true volume of one apparent gram of water when the
latter is weighed in the air with brass weights.
Volume of one apparent gram of
Temperature (C.) water (cc.)
10 1.0014
11 1.0015
12 1.0016
13 1.0017
14 1.0018
15 1.0019
16 1.0021
17 1.0023
18 1.0024
19 1.0026
20 1.0028
21 1.0030
22 1.0033
23 1.0035
24 1.0037
25 1.0040
EXPERIMENTAL GROUP II
VAPOR DENSITY
The following group of experiments is designed to afford practice
in determining molecular weights by measuring the density of
vapors. The method employed was introduced by Victor Meyer and
makes use of the principle of air displacement. Before commencing
experimental work, study the method carefully, since success requires
skilful and intelligent manipulation.
References.
BJK,6-33; F,49; T,39; OW, 101; G, 30.
Weiser: Jour. Phys. Chem., 20, 532 (1916):
Nernst: 253 (1911) for measurements at high temperatures.
Turner : Molecular Association, 6-21 (1915) .
Young: Stoichiometry (2d ed. 1918).
EXPERIMENT 1
Molecular Weight from Vapor Density
Determine the vapor density and molecular weight of an unknown
liquid. Use either (a) Victor Meyer apparatus or (b) the Weiser
modification. See Instructor. Calculate and report molecular
weight. Check results before reporting.
Notes.
Use water as the heating liquid. Boil rapidly and steadily.
It is a good plan to cork the jacket to insure more even heating.
The cork, of course, must be notched to permit the steam to escape.
Use the earthenware burner guard to protect the flame from drafts.
This will insure steady boiling.
The inner tube must be cleaned and dried after each determination.
Dry by blowing in air from the blast, using a long delivery tube
reaching to the bottom of the inner tube. Pass air through a
CaCh tube or tower. The air in the apparatus must be dry at the
beginning of each run.
The bottom of the inner tube must be covered with mercury, clean
sand, or glass wool to protect it against breaking.
It is essential that vaporization should take place as rapidly as
possible. If it takes place slowly, diffusion and condensation of the
vapor on the upper and cooler parts of the tube may occur. The
volume of air displaced should be read as soon as bubbles cease to pass
over into the collecting eudiometer.
Better results are obtained by protecting the outer jacket from
draughts. Cover the outer cylinder with asbestos paper. The inner
tube should not extend far above the cork at the top of the heating
8
jacket. The air displaced by the vapor of the liquid must be at the
same temperature as the vapor displacing it. Explain why this is
necessary.
Do not attempt to start this experiment until you understand the
operation of the apparatus and know the reasons for the many and
important precautions.
Take the following readings during each determination:
(1) Weight of sample.
(2) Final volume of air displaced.
(3) Barometer reading and barometer temperature.
(4) Temperature of water and temperature of air surrounding
eudiometer tube. These temperatures should be the same.
(5) Height of water column in eudiometer.
Calculate the molecular weight of the unknown substance from the
above data.
EXPERIMENTAL GROUP III
LIQUIDS AND LIQUID MIXTURES
The purpose of this group of experiments is to study some of the
interesting properties of liquids and liquid mixtures, with special
attention to volume changes, refractive indices and viscosity.
References.
Dunstan and Thole: The Viscosity of Liquids (1914).
Kuenen: Verdampfung und Verflussigung (1906).
LeBas: Molecular Volumes of Liquid Chemical Compounds
(1915).
Smiles: Chemical Constitution and Physical Properties (1910).
Turner: Molecular Association (1915).
Young: Stoichiometry (1918).
EXPERIMENT 1
Change of Volume and Temperature on Mixing Liquids
Reference. Kuenen: Verdampfung und Verflussigung, 142.
Part 1. Mix 54 cc. of water and 46 cc. of alcohol. Measure
temperature change and also change in volume. Have water and
alcohol at same temperature before mixing and read temperature to
1/5 degree centigrade. Obtain thermometer from Instructor.
Part 2. Mix equal parts by volume of carbon disulphide and
acetone. Proceed as before.
This experiment illustrates the fact that unexpected and profound
internal changes often accompany the mixing of two liquids.
EXPERIMENT 2
Refractive Index of Liquid Mixtures
The refractive index of ordinary glass is 1.54; that of benzene,- 1 .51 ;
while carbon bisulphide has a refractive index of 1.64 for the same
wave leqgth of light. One- can prepare a mixture of benzene and
carbon bisulphide having the same refractive index as glass for a given
wave length. The glass practically disappears as the refractive index
of the solution approaches that of the glass. Explain. H. G. Wells
has made fantastic use of this principle in his "Invisible Man."
In a test-tube place about 5 cc. of CeHe. Add CS 2 until a clean
glass rod, when dipped into the mixture, becomes invisible.
This experiment illustrates the fact that "the properties of liquid
mixtures are often not widely different from the algebraic sum of the
properties of the constituents."
10
EXPERIMENT 3
Relative Viscosity of Benzene and Water
Procedure. F, 83; OW, 162; etc.
"Having thoroughly cleaned a viscosity tube, introduce into the
larger bulb, by means-of a pipette, a known volume of water, recently
boiled and allowed to cool, sufficient to fill the bend of the tube and
half, or rather more than half, of the large bulb.
"Fix the viscosity tube in the thermostat and after allowing ten to
fifteen minutes for the temperature of the tube and the water to
become constant, attach a piece of rubber tubing to the narrower
limb of the viscosity tube and suck up the water to above the upper
mark. Then allow the water to flow back through the capillary and
determine the time of outflow by starting the stop watch as the
meniscus passes the upper mark. Repeat the measurement four or
five times and take the mean of the determinations. If the time of
outflow is about 100 seconds, the different readings should not deviate
from the mean by more than 0.1 to 0.3 second. Greater deviations
point to a soiled capillary tube.
"The viscosity tube must now be dried and an equal volume of pure
benzene introduced into the tube in place of water. Readings are
made as in the case of water." F, 87.
The density of benzene at 20 and water at 20 are given below. The
viscosity of benzene, relative to that of water at 20 is calculated by
means of the formula:
Relative viscosity of benzene = Time x density of benzene
Time x density of water
Notes.
Use a large beaker (1500 cc.) as a water bath and to insure a
constant temperature keep well stirred. The compressed air
furnishes an excellent means of stirring.
It is important to keep the temperature constant because the
viscosity changes rapidly with the temperature (about 2% per
degree).
Record the temperature frequently. It is not sufficient to work at
room temperature; the temperature must be that specified in the
directions.
See that the viscosimeter is immersed far enough to cover the
upper bulb.
Part 1. Determine the relative viscosity of water and benzene
at 20.
Part 2. Repeat at 40 C.
These two experiments show the influence of temperature on the
viscosity. The densities follow:
Liquid 20 C. 40 C.
Water 0.9982 0.9920
Benzene 0.8790 0.8600
11
EXPERIMENT 4
Viscosity of Mixtures of Ethyl Alcohol and Water
Find the relative viscosity (water as standard) of the following:
Absolute ethyl alcohol and mixtures containing 80, 60, 50, 40,
and 20 parts of alcohol in 100 parts by weight of alcohol and water.
Work at 20 C. 0.1.
The viscosimeter must be clean. It is a good plan to rinse thor-
oughly with the mixture whose viscosity is to be measured. Always
employ the same volume of alcohol- water mixture in the viscosimeter.
Record the temperature during each determination.
Draw a curve with viscosity as ordinates and composition as
abscissas.
The following data will be required:
Parts Alcohol in Mixture Density (20 C.)
0.9983
20 0.9688
40 0.9351
50 0.9140
60 0.8913
80 0.8437
100 0.7895
This experiment is another illustration of the fact that in
many cases the properties of a mixture are unexpectedly different
from the properties of the pure constituents. Compare with Experi-
ment 1 above. Could one use viscosity measurements as a means of
determining the alcohol content of alcohol-water mixtures?
EXPERIMENT 5
Relative Viscosity of Unknown
Determine the relative viscosity of an unknown solution, using
water as standard. The Instructor will supply the unknown solution
and will state the temperature at which to work, and the density of
the solution.
EXPERIMENT 6
Specific Gravity Flotation
If a mixture of dry sawdust and iron filings is thrown into water,
the sawdust will float and the iron filings will sink, the two being
separated by means of a liquid whose specific gravity lies between
those of the mixed solids.
Employing this principle, separate the mixture of two solids
which is found on the shelf. Use as the liquid the solution formed
when HgI 2 dissolves in an excess of KI.
Specific gravity of the lighter solid is 2.05.
Specific gravity of the heavier solid is 3.60.
12
The specific gravity of the liquid should be about midway between
these values.
The solution is made by adding saturated KI solution to the
saturated HgCl 2 solution on the shelf. Avoid large excess of KI.
Put in a test tube and shake violently, for two or three minutes. See
T, 17; also Stahler 1, 626 (1913). Danger; Mercuric chloride is
extremely poisonous !
Compare flotation of this type with froth flotation now used on so
large a scale for the concentration of sulphide ores. See Mineral
Industry, 24, 807 (1915); Megraw: The Flotation Process (1917).
13
EXPERIMENTAL GROUP IV
VAPOR PRESSURE
The following group of experiments is designed for the purpose of
studying and measuring the pressure exerted by the vapor phase
when in equilibrium with a pure liquid or with a liquid mixture.
References.
Kuenen: Verdampfung und Verflussigung 117, 129 (1906).
Young: Stoichiometry 260 (1908).
Also see references under "Distillation," Experimental Group VII;
Papers by Smith and Menzies in Jour. Am. Chem. Soc. (1910 ).
Apparatus.
One heavy-walled test tube 150 mm. long, 25 mm. ext. diameter,
fitted with two-hole rubber stopper; Chapman water suction pump
(large size) ; Mercury manometer ; Y-tube ; two glass stopcocks ;
pressure tubing, etc. Instead of test tube and rubber stopper a
special glass stoppered test tube may be employed with better results.
Procedure.
Refer to the diagram of apparatus. The heavy test tube A, the
vaporization vessel containing the liquid under investigation, is
immersed in a constant temperature water bath. The two stopcocks
are placed at P r and P 2 . B serves as a trap. The remainder of the
diagram requires no explanation.
First assemble the whole apparatus, connect with manometer and
pump and, closing P t and opening P 2 , test the apparatus for leaks.
If the pump is working properly a "vacuum" of 2-3 cm. should be
obtained. Read the barometer in the balance room and from this
reading subtract the reading obtained on your manometer. The
pressure in A should not exceed 35 mm. and should remain constant
on closing P 2 .
Place in A the liquid whose vapor pressure is to be measured.
Use 10-20 cc. Replace stopper completely, submerge the whole
test tube in the constant temperature bath and proceed with the
measurement. Close P x and open P 2 . Gently agitate the liquid in
A by shaking the test tube back and forth; this will tend to prevent
bumping during vaporization. When the liquid begins to vaporize
or to boil slightly, close P 2 and continuing the shaking to hasten
equilibrium, read the manometer when the latter remains constant.
Again open P 2 and vaporize for an instant. For the second time
close P 2 and read the manometer as before. When repeated vaporiza-
tions of very short duration fail to cause an appreciable change in the
manometer readings, and the difference in the mercury levels in the
two arms of the manometer reaches a maximum and is constant,
subtract this difference from the height of the barometer. The value
so calculated is the vapor pressure of the liquid in A.
14
Notes.
Be sure that the suction pump is clean and is operating properly.
Be on guard against violent bumping when the liquid in A boils.
Shake A to prevent this and to hasten adjustment of thermal equili-
brium between the liquid in A and the water bath. (Experiment:
place some ether in A and connect with the vacuum pump. Note
temperature of ether).
Make certain that all air has been removed from A before taking
the final reading of the manometer. The vaporizing process should
remove the air.
In dealing with solutions vaporize no more than is just necessary
to remove air. Boiling or vaporizing a solution almost always
changes the composition of both liquid and vapor. Explain.
Compare the vapor pressures so determined with those given in the
tables in LBR. The results with pure liquids should not be more
than 2 per cent in error.
EXPERIMENT 1
Vapor Pressure and Composition
Part A. Ascertain the vapor pressure- com position relations in
the system ethyl alcohol and benzene, a pair of consolute liquids.
Temperature 20 C.
Measure the vapor pressure of the following: Pure ethyl alcohol;
benzene; mixtures of alcohol and benzene containing 10, 25, 32, 50,
75, and 90 parts of alcohol by weight in 100 parts of mixture. Den-
sity benzene = 0.88; alcohol (absolute) = 0.78 at 20 C. Mix, using
burettes.
Plot a curve as you proceed with the determinations ; pressures as
ordinates, compositions as abscissae.
Note that the results with the solutions are approximate only,
because the vaporization process, especially if prolonged, causes the
composition of the liquid phase to change and thereby to be different
from the composition of the original mixture. For accurate work the
composition of the liquid at the end of the experiment should be
determined. The method as described approaches sufficiently close
to the more accurate method to enable the student to obtain the
characteristic pressure-composition diagram.
Part B. Ascertain the vapor pressure-composition relations in
the system acetone and chloroform. Temperature 20 C.
Measure the vapor pressure of the following: Acetone; chloro-
form; mixtures of acetone and chloroform containing 15, 25, 40, 50,
60, 75 and 85 parts of acetone in 100 parts of mixture. Density
acetone = 0.80; chloroform = 1.52.
Plot a curve as you proceed. Compare with Part A.
Note. Do either Part A or Part B as assigned.
EXPERIMENT 2
Two Liquid Layers
Determine the vapor pressure of pure ethyl acetate at 20 C.
Look up the vapor pressure of water.
15
Determine the vapor pressure of the following mixtures of ethyl
acetate and water containing 25, 50 and 75 parts of ethyl acetate in 100.
Explain your results. Density ethyl acetate = 0.923.
EXPERIMENT 3
Raoult's Law
Part A. Determine the lowering of vapor pressure when 5 g. of
naphthalene are dissolved in 20 g. of acetone. From this calculate
the molecular weight of naphthalene, using Raoult's formula:
a)
p ' N+n
p = vapor pressure of pure acetone ; p' = vapor pressure of solution ;
n = gram molecules of naphthalene ; N = gram molecules of acetone.
Part B. Determine the molecular weight of nitrobenzene (8g.)
in ether (25 cc.). Density ether = 0.73; nitrobenzene = 1.2.
Determine vapor pressure of ether separately.
Note. Do either (A) or (B) as assigned.
Part C. Optional Experiment; Menzie's Method.
Using Menzie's apparatus (see Instructor) determine the molecular
weights of naphthalene or nitrobenzene. Reference: Bigelow, 320.
Notes.
Look up the vapor pressures of naphthalene and nitrobenzene at
20 C. Are these solutes volatile at this temperature?
How is the lowering of vapor pressure made use of in Burger's
method of determining molecular weights when very small amounts
of substances are available? Jour. Chem. Soc., 85, 286 (1904);
Chamot: Chem. Microscopy, 216 (1915).
EXPERIMENT 4
Vapor Pressure of Aqueous Solutions
For this work connect the apparatus with the rotary vacuum
pump, protecting the latter from moisture by means of a tower or
tube containing anhydrous calcium chloride. Temperature 20 C.
Determine the vapor pressure of (a) water (b) 5 per cent cane
sugar solution (c) 30 per cent cane sugar solution (d) a solution
containing enough calcium chloride to be equimolecular with the
sugar solution in (b). Explain all results.
EXPERIMENT 5
Vapor Pressure (Dissociation Pressure) of Salt Hydrates
A crystalline salt hydrate will effloresce (dissociate) when exposed
to the air if the partial pressure of water vapor in the air is less than
the dissociation pressure of the hydrate. Measure the dissociation
pressure of Glauber's salt (sodium sulphate decahydrate), correspond-
ing to the reaction:
Na 2 S0 4 . 10H 2 = Na 2 S0 4 + 10H 2 O
Reference. Findlay: The Phase Rule, 83 (1904).
Compare with Experiments 1 and 2 of Experimental Group VIII.
16
EXPERIMENTAL GROUP V
ELEVATION OF THE BOILING POINT
This group of experiments deals particularly with the changes
produced in the boiling point when a soluble, non-volatile substance
is 'added to a pure solvent. The differences between electro-
lytes and non-electrolytes are emphasized and explained^ by means
of the theory of electrolytic dissociation. Molecular weights are
determined by the so-called "boiling point" method.
References. BJK, 141-197; F,138; OW,184; G,77; T,97. Read
Bigelow, 317.
EXPERIMENT I
Preliminary Experiment to Illustrate the Difference Between an
Electrolyte and a Non-Electrolyte
Typical non-electrolytes sugar, urea, etc.
Typical electrolytes NaCl, KC1, etc.
This experiment is to be performed by groups of students working
together under the direct supervision of the Instructor.
Place two clean graphite electrodes in 350 cc. of distilled water
contained in a 500 cc. beaker. Connect electrodes to 110-volt
alternating- current circuit in series with a lamp-bank resistance.
Short circuit the current across the electrolyzing cell and observe
the brightness of the lamps. The brightness is roughly a measure
of the current flowing. Then pass the current through the distilled
water and observe again the brightness of the lamps.
Substitute 350 cc. of tap water for distilled water. What do you
observe?
Finally test in order 5 g. of the following substances dissolved in
350 cc. of distilled water: Sodium chloride, mercuric chloride,
cane sugar, and acetic acid. Carefully wash the graphite electrodes
after each solution has been tested.
What can you say regarding the power o'f the above solutions to
conduct the electric current? Is a good electrolyte always an
inorganic salt, and are all inorganic salts good .electrolytes?
Note. A very neat experiment similar to the above is described
in Lamb, 33.
EXPERIMENT 2
Elevation of the Boiling Point
Procedure. Cf. Bigelow, 319, Walker and Lumsden: Jour.
Chem. Soc., 73, 502 (1898).
For this experiment a modified and simple form of the Lands-
berger apparatus for vapor heating is employed. The import-
ant features are three, viz, vapor (steam) generator, boiling chamber
17
(fitted with Beckmann thermometer, outer jacket and exit tube)
and suitable condenser. See the diagram.
The steam generator should be operated at constant speed and
without "bumping". To ensure this, protect the burner with an
earthenware guard and add pumice generously to the water in the
round-bottom flask. Do not change the rate of boiling during a
given run and do not shut off or move the burner under any circum-
stance.
Set a Beckmann thermometer for the boiling point of water.
Make sure that the mercury is low on the scale. Handle with care
the delicate and expensive thermometer.
Start the generator boiling and, when ready, connect to the
boiling chamber containing the solvent. The boiling chamber should
be well insulated thermally. This may be done by using a Dewar
tube (thermos vacuum bottle), by surrounding the tube with the
vapor of the solvent as in the McCoy apparatus (which see), or by
slipping the large test tube serving as boiling chamber into a wide-
mouth bottle, fitting snugly, and closing the annular space at the
neck with felt or cotton wool. The delivery tube for the steam
should reach to the bottom of the boiling chamber and the Beckmann
thermometer should be immersed far enough to submerge the bulb.
Weigh the dry test tube so that the weight of the solution whose boil-
ing point is measured may be determined.
Place pure water in the boiling chamber and boil with steam.
When the mercury reaches a steady position on the scale, take a series
of ten consecutive readings at intervals of ten seconds. Use a read-
ing glass (to be obtained from Instructor). The readings should not
fluctuate by more than one of the smallest divisions (0.01 C.).
Read the barometer before and after.
Then, without interrupting the boiling, disconnect the steam line
from the boiling chamber, lift the cork holding the thermometer,
and drop into the water in the tube a weighed quantity of solute.
Determine the new boiling temperature.
Stop the run, remove the large test-tube, cool and weigh. Cal-
culate the weight of water in the solution. Using the formula
M=520^ (1)
compute the molecular weight of the dissolved solute.
Beckmann Differential Thermometers.
Cf. F, 129-133.
"Some thermometers have scales which allow the adjustment of
the zero point when desired. One kind has a scale which may be
screwed up or down from the top. Another kind permits a change in
the volume of mercury. The Beckmann is of the latter type. This
thermometer has at the upper end of the capillary a mercury reservoir
which allows one to decrease or increase the actual amount of mer-
cury in the bulb and capillary thread. To decrease the mercury in
the bulb, the bulb is heated until the needed amount of mercury
appears in the reservoir as a globule, then a sharp tap with the hand
will separate it, if the thermometer is held in an upright position. It
18
is apparent then that the temperature of the bath should be higher
than the required zero reading by the number of degrees correspond-
ing to the length of thread which is not required."
A good Beckmann thermometer should fulfill these requirements:
(1) The upper and lower mercury reservoirs should branch into
the capillary in a conical fashion.
(2)^ Mercury should be clean.
(3) Thermometer should not be unnecessarily clumsy.
Part 1. Non-Electrolytes.
Determine the molecular weights of urea and cane sugar. Use
1 /20th g. molecule of each substance. From your own data calculate
the elevation you would have observed if the solutions had contained
exactly 500 grams of water. . How do these elevations compare
with each other?
Part 2. Electrolytes.
Proceed as in Part 1 with KC1 and K 2 SO 4 . Compare with Part 1.
Part 3. Mercuric Chloride.
Proceed as in Part 1 with mercuric chloride (poison). Compare
with urea and sugar and with KC1 and K 2 SO4.
Part 4. Unknown.
Determine the molecular weight of an unknown substance.
After obtaining good checks report your results to the Instructor.
Part 5. Ethyl Alcohol as Solvent.
Place absolute alcohol in the outer compartment and about 6 g. of
absolute alcohol in the inner compartment of a McCoy vapor heater.
Guard against fire by connecting a long rubber tube to the side arm.
When the alcohol has boiled for some time close this rubber tube with
a pinchcock and heat the alcohol in the inner compartment with
alcohol vapor. The inner compartment is fitted with a stopper con-
taining an exit tube connected with a condenser and a Beckmann, the
bulb of which is immersed in the alcohol.
Determine the boiling point of pure absolute alcohol.
Then add molecular weight of urea to the alcohol in the inner
compartment and heat the solution with the vapor. When the
boiling point has reached a maximum, pour the contents of the inner
tube into a bottle and determine the weight of the solution.
For this experiment the Beckmann must be set for the boiling
point of absolute alochol.
Boiling constant for ethyl alcohol, 1170. Calculate the molecular
weight.
Notes.
Redetermine the boiling point of the pure solvent before each run.
If this is not done and only one determination is made, the barometric
(atmospheric) pressure may change enough to give very misleading
results. At about 100 C. a change in pressure of only 1 mm. of
mercury produces a temperature difference in the boiling point of
nearly four hundredths of a degree. Bigelow, 317.
19
For a critical discussion of the method and a very elegant apparatus
for determining the elevation of the boiling point read Cottrell;
Jour. Am. Chem. Soc., 41, 721 (1919) and Washburn and Read:
Ibid., 41,737 (1919).
EXPERIMENT 3
Lowering of the Boiling-Point
Discussion.
A non- volatile solute added to a pure liquid always raises the boil-
ing point. When however a non- volatile solute is added to a mixed
solvent containing two volatile liquids a depression of the boiling
point may be produced instead of an elevation.
Let A and B be two volatile substances forming a single homogene-
ous solution. Call A the solvent and B the solute. As B is added to
A the concentration of the solution increases and the partial pressure
of A in the vapor becomes smaller (Raoult's law). At the same time
the partial pressure of B increases in the vapor (Henry's law). When
A is saturated with B the solution is in equilibrium with pure B and
the partial pressure of B in the vapor is practically equal to the
vapor pressure of pure B.
It follows from this that, for a given concentration of B in A, the
greater the solubility of B, the smaller is the partial pressure of B in
the vapor. Anything which decreases the solubility will tend to
increase the partial pressure of the solute in the vapor.
The solubility of B in A may be made less by the addition of a suit-
able third substance. If the latter is non-volatile and soluble both
in A and B, it can affect the total vapor pressure of the solution in
two ways, as follows:
(1) By decreasing the solubility of B in A (or A in B).
(2) By dissolving in A and B.
Influence (1) points in the direction of increased vapor pressure
and may in fact be greater than influence (2) which tends toward a
lower vapor pressure. (Why?) The total vapor pressure, which
is equal to the sum of the partial pressures of A and B, may thereby
be increased and the boiling point depressed. " The experiments
which follow illustrate the point.
Procedure.
Determine the boiling point of a mixture of 50 parts alcohol and
50 parts water. Use a flask and reflux condenser, determining to
tenths of one degree with a special-thermometer (not the Beckmann).
Then add sodium carbonate to the alcohol-water mixture and rede-
termine the boiling point. Do two layers appear as carbonate is
added?
Repeat, using cane sugar instead of Na 2 CO 3 .
It may be found advisable to add the Na 2 CO 3 , or sugar in several
portions, determining the boiling point each time.
20
EXPERIMENTAL GROUP VI
DEPRESSION OF THE FREEZING POINT
The object of the following group of experiments is the study and
use of the freezing point method of determining molecular weights.
References. F, 125-138; T, 81-90; OW, 180-184; etc.
Procedure.
Apparatus: Freezing point apparatus (see diagram) ; Beckmann
thermometer; reading glass; etc.
Set the Beckmann for the freezing point of water.
As solvent use 10-15 cc. of distilled water, i. e. enough to cover the
bulb of the thermometer.
In the battery jar place a freezing mixture of salt and ice. The ice
must be pounded fine and be well mixed with salt. The best
temperature for the freezing bath is about 5 C. A lower tempera-
ture than this is undesirable. Record temperature of freezing mix-
ture. See Findlay on "convergence temperature."
In the freezing mixture, place the outer tube or jacket, and in the
jacket, the inner tube, which must not come in contact with the walls
of the outer. The jacket should be closed by a cork through which
the outer tube passes.
Determine first the freezing point of the solvent, noting the degree
of undercooling (supercooling) and tapping the thermometer fre-
quently to prevent stiction. The water must be stirred constantly
to prevent excessive undercooling. Take the tube out of the jacket
and warm in the hand until the ice melts. Redetermine the freezing
point. Undercooling should not exceed 1 C.
The preliminary cooling may be hastened by placing the inner tube
directly in the freezing mixture. Take care that no salt from the
freezing mixture is introduced into the solution and dry the tube very
carefully before replacing in the outer tube.
The inner tube should be closed by a cork through which the ther-
mometer and stirrer pass.
EXPERIMENT 1
Water as Solvent
Determine the molecular weight of an unknown salt. Use about
15 (weighed to O.lg.) of water and not more than 0.3g. of the un-
known. When your results check satisfactorily, report them to the
Instructor. Constant for water, 1860.
21
EXPERIMENT 2
Benzene as Solvent
Part 1 . Determine the molecular weight of naphthalene or anthra-
cene in benzene. Use about 1 /1000th gram-molecule of solute in 10 g.
of benzene (thiophene free). Set the Beckmann for benzene (5.5 C.)
and use ice alone (no salt) as the freezing agent. Constant for
benzene, 5000.
Part 2. Proceed as above with benzoic acid in benzene. How do
you account for the high value of the molecular weight?
For accurate results the benzene should be anhydrous and should
be protected from moisture in the air.
EXPERIMENT 3
To show how the Freezing Point of a Metal may be affected by other
Metals
The melting point is used as a criterion of purity, especially in
organic chemistry. This experiment shows how one substance affects
the melting point of another.
The fusible alloy is a mixture of Bi, Cd, Pb, Sn.
Place about 0.1 g. in a small glass tube which has been closed at one
end by drawing down and fusing. Find the melting point of the
alloy in a water bath.
Look up the melting points of the metals composing the alloy in
LBR, 190.
EXPERIMENTAL GROUP VII
DISTILLATION OF LIQUID MIXTURES
The following experiments are designed to illustrate the distillation
of mixtures both constituents of which are volatile at the boiling
point. Particular emphasis is laid on the relations existing between
boiling temperature and the composition of residue and distillate.
References.
Kuenen: Verdampfung und Verfliissigung (1906).
Ostwald: Fundamental Principles of Chemistry 123-148.
Rosanoff: Jour. Am. Chem. Soc. (1909-).
Young: Fractional Distillation (1903).
Young: Stoichiometry (1918).
EXPERIMENT 1
Hydrochloric Acid and Water
Solutions.
(a) 1 liter of 10 per cent HC1 i. e. (10 g, HC1, 90 g. water).
(b) 500 cc. 30 per cent HC1.
(c) 2 liters 2NNaOH, standardized against N/l HC1.
(Use rubber stopper for reagent bottle).
Part 1. Distillation of the 10 per cent Mixture
Discussion.
When two miscible liquids are distilled, the composition of residue
and distillate (vapor) will generally differ at any given temperature
of ebullition and the latter will rise as the distillation is continued.
The distillate (vapor) will always be richer in respect to the more vola-
tile constituent or, if the pair of liquids gives a mixture of minimum
boiling point (water and ethyl alcohol), the distillate will be richer
than the residue in respect to this mixture. If, however, the pair of
liquids gives a mixture with a maximum boiling point (HC1 and water
HNO 3 and water; H 2 SO4.and water) the distillate will be richer than
the residue in respect to either one of the pure constitutents, depend-
ing upon conditions. What these conditions are will be shown by the
following experiments.
Procedure.
Place 500 cc. of the 10 per cent solution in a liter distilling flask
connected with condenser and receiver. Place the thermometer in
vapor and use ebullition tubes or pumice to prevent bumping.
Before starting to distill determine the KC1 content of the solution
by titration with standard NaOH. Withdraw a 5 cc. sample from
the flask with a pipette.
23
Distill and collect the distillate in a measuring cylinder. Wherr
about 30 cc. of distillate have been collected, remove the measuring
cylinder and empty it of its contents as completely as possible. Then
collect between 5 and 8 cc. of fresh distillate, noting the average
temperature at which it comes over. Stop the distillation.
Withdraw a 5 cc. sample of distillate and determine its HC1 con-
tent. Next withdraw rather more than 5 cc. of hot residue in a flask,
cool and titrate a 5 cc. sample.
Again distill; collect another 30 cc.; throw this away as before
and collect a second sample of 5 to 8 cc., observing the tempera-
ture. Continue until nearly all of the acid has been distilled over.
Arrange data as follows:
(a)
(b)
(c)
(d)
(e)
(f)
Number
Nature of
Volume of
NaOH
Grams
Tempera-
of Sample
Sample
of Sample
(cc.)
HC1 per
ture
(cc.)
100 cc.
(average)
No. 1
Original
5
6.60
9.45
99.5
No. 2
Distillate
5
0.15
0.22
110.5
Residue
5
6.80
9.65
100.5
etc.
etc.
etc.
etc.
etc.
etc.
Part 2. Distillation of the 30 per cent Mixture (Hood).
The procedure requires modification, since at the start nearly pure
gaseous HC1 is given off. Do not determine the composition of the
distillate until the distillation is nearly finished. Instead, analyze
samples of the residue at appropriate intervals and observe the
temperature immediately prior to withdrawing the samples.
Connect the condenser to an absorption train for the removal of
HC1 fumes (adapter dipping below the water in beaker) and work
in hoods.
When the temperature has reached a nearly constant value remove
the absorption apparatus and proceed exactly as in the previous case,
analyzing both distillate and residue.
Part 3. Distillation of 10 per cent Mixture with Vigreux Column.
Start with 500 cc. of acid mixture in a round bottom flask to which
a long Vigreu column has been fitted. Place a thermometer at the
head of the column in the usual fashion, also a thermometer in the
vapor in the flask. Take simultaneous reading of both thermometers
throughout.
Proceed with the 10 per cent solution just as in Part 1, analyzing
both distillate and residue. Continue the distillation until residue
and distillate have the same composition.
Computations and Curves.
Calculate results in terms of grams of HC1 in 100 cc. of solution.
On a single sheet, draw curves between temperatures as ordinates and
24
composition as abscissae. Do all the curves approach a common
point? What is the effect of the Vigreux column? Explain.
Compute the percentage of HC1 by weight in the mixture of maxi-
mum boiling point. Consider the specific gravity of the mixture to
be 1 .1 . Use the data as determined by the experimental curves.
From the data derive a formula for the constant boiling mixture,
assuming that it is a definite hydrate of hydrochloric acid. How was
the simple hydrate theory disproved?
EXPERIMENT 2
McCoy Apparatus and Vapor Heating
Part 1. In the outer compartment of a McCoy apparatus place
ethyl alcohol. Connect a condenser to one of the side arms; to the
other a short piece of rubber tubing fitted with a pinch cock. Keep-
ing side arm open, heat the alcohol to boiling. In the inner compart-
ment place 5 or 10 cc. of benzene and close the tube with a cork carry-
ing a thermometer dipping into the benzene. When the alcohol is
boiling very gently and evenly, close the pinchcock and pass alcohol
into the benzene. Read time and temperature at intervals of 15
seconds. Draw a curve with times as abscissae and temperatures as
ordinates.
Precaution. Do not begin heating with vapor until the ther-
mometer in the benzene registers higher than 75 C; then pass in
alcohol vapor as slowly as possible. The rate of heating should be
kept constant throughout.
Part 2. Repeat with acetone in the outer compartment and
chloroform in the inner.
Part 3. Repeat with water in the outer compartment and methyl
alcohol in the inner.
Part 4. Repeat with ethyl acetate in the inner compartment and
water in the outer. Observe carefully the formation of two layers.
Why does the temperature remain constant and how does it compare
with the boiling temperature of pure ethyl acetate and pure water?
Explain.
Part 5. Repeat with water in the inner compartment and ethyl
acetate in the outer.
See Experiment 4 below.
%
EXPERIMENT 3
Steam Distillation
Take two 1000 cc. distilling flasks. In one place distilled water,
beads to prevent bumping, and a thermometer reading to 110
immersed in the liquid. In the other place a concentrated solution
of NaCl and add NaCl in large excess. In this flask place a thermo-
meter reading to at least 125 and immerse in the liquid. See sketch.
25
Boil the water in the first flask and when the water is boiling
gently, connect to the other flask and pass steam into the salt solution .
Note the temperature in each flask, making frequent readings.
When the temperature in the flask containing the solution has
reached a maximum, take the temperature of the vapor in each flask.
Thoroughly wash the thermometer with water after withdrawing
from the solution, and again take the temperature of the vapor.
Explain the results.
Regarding the differences observed when the thermometer is
immersed in the vapor and not in the liquid, see Hite: Am. Chem
Jour. 17, 510 (1895); Sakurai: Jour. Chem. Soc., 61, 495 (1892).
EXPERIMENT 4
To Map out the Boiling Point Composition Diagram for a Binary
Liquid Mixture
Determine the temperature at which the liquid mixture boils
steadily. Use a small round-bottomed flask and not more than 30
grams of liquid in each case. The neck of the flask should be fairly
wide and should be fitted with a cork carrying a thermometer and
connected with a reflux condenser. Place the thermometer in the
liquid mixtures (chloroform-acetone or benzene-alcohol) that you
studied in Experimental Group IV, Experiment 1 A or 1 B. Having
determined the boiling point, plot the values against the composition.
Compare with the pressure-composition diagram.
2t>
* EXPERIMENTAL GROUP VIII
DISSOCIATION
The following experiments are designed to illustrate qualitatively
the dissociation of ehemical compounds, either as the result of an
increase in temperature or as the result of dissolving the substance in
a solvent. Dissociation of the first type is called thermal; dissocia-
tion of the second type is called electrolytic when ions are formed.
We have already studied some of the phenomena due to electrolytic
dissociation, especially in Experimental Groups IV and V. Other
instances of electrolytic dissociation and its effects will be studied in
the Experimental Groups which follow.
References. Solutions and Electrolytic Dissociation.
Abegg: Die Theorie der elektrolytischen Dissociation, Ahren's
Sammlung 8 (1903).
Arrhenius: Theories of Solution (1912).
Findlay: Osmotic Pressure (2nd Ed. 1919).
Jacques: Complex Ions (1914).
Jones : The Nature of Solution (1917) .
Ostwald (Muir) : Solutions (1891).
Rothmund: Die Loslichkeit (1907).
Scxidder : Electrical Conductivity and lonization Constants (1914) .
Seidell: Solubilities of Inorganic and Organic Compounds (1919).
Stieglitz: Qualitative Analysis, Vol. I (1917).
EXPERIMENT 1
Thermal Dissociation of Nitrogen Tetroxide
In a test tube heat a small quantity of Pb(NO 3 ) 2 and pass the
resulting gas through a delivery tube into a test tube which is
surrounded by a freezing mixture of ice and salt.
The NO 2 will condense, under these conditions, as a bluish green
liquid, N 2 O4. On removing from the cooling bath the colorless gas
N 2 O4 will be formed first and on further heating this will dissociate
into NO 2 . Note color changes.
References. Nernst, 453 (1911). Ostwald: Principles of Inor-
ganic Chemistry, 329 (1908).
EXPERIMENT 2
Thermal Dissociation of Limestone
Heat some powdered marble in a hard glass tube. Show that dis-
sociation takes place.
References. Bigelow 4 etc. For study of the reaction used in
lime burning read Kremann- Potts: 107; LBR, 398.
27
Define "dissociation pressure" and draw a curve showing how
dissociation pressure changes with the temperature for the following
reaction: 2NaHCO 3 = Na 2 CO 3 + H 2 O + CO 2 .
Reference. LBR, 398.
*
EXPERIMENT 3
Electrolytic Dissociation and Color
Part 1. Compare the colors of concentrated 'solutions of the fol-
lowing salts: CuSO 4 , CuCl 2 , CuBr 2 . Dilute until they have the
same blue color. Start with about one cc. of solution. Explain.
Part 2. Add concentrated hydrochloric acid to a greenish-blue
solution of CuCl 2 . Note color change. Also heat some of the same
solution. Explain.
References. Ostwald: Prin. Inorg. Chem., 642 (1908); also
Mellor: Inorganic Chemistry, 468 (1914).
Part 3. Color changes with CoCl 2 solutions.
Dissolve a little cobalt chloride in absolute alcohol.
Add two or three drops of water to the solution.
Add ether to the solution.
Add water again.
Also to the pink solution in water add (1) solid magnesium chloride ;
(2) concentrated hydrochloric acid.
References. Donnan and Bassett: Jour. Chem. Soc., 102, 939
(1902). Ostwald: Prin. Inorg. Chem., 623 (1908); Nernst, 389
(1911).
EXPERIMENT 4
Reactions depending upon Degree of Dissociation
Part 1. Pass chlorine gas into AgNO 3 solution. Does a precipi-
tate form at once?
Part 2. Add carbon tetrachloride to AgNO 3 solution. Explain.
Part 3. Add chloroform to AgNO 3 solution. Does a precipitate
form at first?
Let -the mixture stand in the light until the next period. Does a
precipitate form on standing? Explain.
EXPERIMENT 5
Complex Ions
Part 1. To AgNO 3 solution add KCN in excess. Test for silver
with NaCl. Do not use HCU Beware of HCN and remember that
KCN is poisonous. Use hoods.
Part 2. To AgNO 3 solution add sodium thiosulphate'in excess.
Test for silver. Explain results. Cf. Walker, 343.
28
EXPERIMENT 6
Relative Stability of Complex Ions
Part 1. Add KCN in excess to dilute CdSO 4 . Test for cadmium
with H 2 S.
Part 2. Add KCN in excess to dilute CuSO 4 . Test for copper
with H 2 S.
Explain. Cf. Walker, 343.
Caution. Cyanogen is formed in Part 2. Use hoods.
EXPERIMENT 7
Hydrolysis
Part 1. Test KCN and Na 2 CO 3 solutions with litmus paper.
Explain.
Part 2. Test CuSO 4 and A1 2 (SO 4 ) 3 solutions with litmus paper.
Explain.
Part 3. Precipitate PbSO 4 completely from lead acetate solution
by adding A1 2 (SO 4 ) 3 . Then add water and boil. Filter and test
the filtrate for lead and aluminum.
Precaution. It is essential to use very little Al2(SO 4 ) 3 in excess.
At any rate, add plenty of water and boil thoroughly for several
minutes. Explain.
EXPERIMENT 8
Conductivity and Electrolytic Dissociation
Discussion.
The conductivity is the reciprocal of the resistance. From the
resistance of a solution, its conductivity may be calculated. In this
experiment the relative resistance of N/10 HC1 and N/10 CH 3 COOH
is measured by reading the current and voltage across graphite elec-
-p
trodes which dip into the solution. From Ohm's law, I = ' the
R
resistance may be computed.
By maintaining the temperature constant, keeping the electrodes
the same distance apart, and having them immersed to the same
extent, a rough approximation of the conductivity of these two
equivalent acid solutions may be obtained.
The conductivity of a solution depends, among other things, upon
its dissociation. If two solutions are of equivalent concentration and
at the same temperature and if both are placed in the same vessel for
measuring the conductivity, the better conducting solution is either
more completely ionized or else contains the more mobile (the more
rapidly moving) ions. If the difference in conductivity is very great,
as in the present case, the poorly conducting solution is almost certainly
the less strongly dissociated. Since both solutions have the hydrogen
ion in common and since the chlorine and acetate ions are about
29
equally mobile, the relative conductivity is here a very nearly exact
measure of the relative ionization.
Procedure.
Measure the relative resistance of N/10 HC1 and N/10 CH,COOH
solutions.
Follow the procedure used in the experiment which showed the
distinction between an electrolyte and a non-electrolyte Use
alternating current and a-c meters.
Look up the per cent ionization of N/10 HC1 and N/10 CH 3 COOH.
Are your conductivity values proportional?
30
EXPERIMENTAL GROUP IX
SOLUTION AND SOLUBILITY
The experiments of the following group are designed to illustrate
the process of solution, the properties of saturated solutions, the cor-
rosion or solution of metals and the determination of solubility.
References. See under Group VIII Dissociation.
EXPERIMENT 1
Quantitative Determination of Solubility
References. F, 302; OW, 176; G, 234, etc.
The solubility of a salt in water depends chiefly upon the nature of
the salt and the temperature. The rate at which the salt dissolves
depends upon the same factors plus several others besides, such as
size of particles, rate of stirring, presence of catalysts, and so forth.
Solubility may be determined directly, provided the salt is not too
slightly soluble, by saturating a solution with an excess of salt at a
desired temperature, and analyzing a definite weight or volume of the
solution.
Determine the solubility of an assigned salt at 25 C. Place in a
bottle an excess of finely powdered salt, add water and shake in a
thermostat until equilibrium is reached, or until there is no change of
density between successive tests, when measured with a delicate
hydrometer. In a second bottle place finely divided salt and add,
not water, but a solution of the salt saturated at some temperature
(usually a higher one) at which the salt is more soluble than it is at
25 C. Shake as before and determine the density of the saturated
solution. The final densities should be the same in both bottles.
Withdraw samples for analysis using a dry pipette and a small
filtering tube to prevent the entry of solids. Determine the concen-
tration of the saturated solution either by chemical analysis, or by
evaporating a weighed sample to dry ness in an oven or desiccator.
Check results. Determine the density of the solution at 25 C. and
calculate the solubility of the salt in grams per 100 grams of solution;
also in terms of the "molar fraction" of the solute.
EXPERIMENT 2
Cryolite and Water
Add a little finely powdered cryolite to water in a test tube. Does
it dissolve? Explain.
31
EXPERIMENT 3
Solution and Catalysis
Chromic chloride appears in two forms, as the hexahydrate
(CrCl 3 . 6H 2 O) green in color, and as the anhydrous salt (CrCl 3 )
which is violet. The anhydrous form appears to be nearly insoluble
in water while the hydrate dissolves readily. According to Moissan
the violet form dissolves slowly at high temperatures to a green
solution, and Ostwald believes that the apparent insolubility at
ordinary temperatures is due to the extreme slowness with which
solution occurs; in other words, that the violet form is not really in
equilibrium with water. Drucker under Ostwald's direction showed
that the violet modification dissolves readily in the presence of
chromous chloride (CrCl 2 ) in solution, the latter acting as a catalyst.
With these facts in mind perform the following tests:
(1) Try to dissolve violet CrCl 3 in water.
(2) Dissolve some green hexahydrate in water.
(3) To a small quantity of the violet salt add water plus a crystal
of the green hexahydrate. Add a bit of zinc and acidify with HC1.
See whether the violet salt dissolves in time. Explain.
(4) To th'e violet salt add a bit of metallic chromium, then add
dilute HC1. Does the salt dissolve?
(5) Prepare chromous chloride by dissolving metallic chromium
in dilute HC1. Add this solution to a few particles of the violet salt.
Do they dissolve?
(6) Repeat (d) adding zinc instead of chromium. Explain.
References. Ostwald: Prin. Inorg. Chem. 615; Mellor: Inorg.
Chem. 258.
Drucker: Zeit. phys. Chem., 36, 173 (1901).
EXPERIMENT 4
Relative Solubility
Part 1. Precipitate PbSC>4, let it settle, wash once or twice by
decantation, then add KI solution to the residue. Note the color
change. Then warm it. What color change occurs?
Part 2. Precipitate AgCl, repeat procedure in (a) using KBr
solution. What change occurs in the precipitate? Explain.
Part 3. Repeat (2) using KI solution. Walker, 356 (1913).
Part 4. Precipitate AgBr, add KC1 solution. Is there any visible
change? Explain.
Part 5. Prove by simple experiments, which is the more soluble,
CaSO 4 or CaCO 3 .
EXPERIMENT 5
Compound Solvents
Part 1. Add about 20 cc. of impure commercial sulphuric acid
to an equal volume of water. What is the precipitate? Explain.
32
Part 2. Add about 5 cc. of 95 per cent ethyl alcohol to (1) a
saturated solution of Na 2 SO4 (2) a saturated solution of Na 2 CO 3 .
cf. Group V, Experiment 3.
Part 3. Determine by experiment qualitatively the effect of
sodium chloride on the solubility of phenol in water. Repeat with
sodium acetate instead of sodium chloride.
EXPERIMENT 6
Solubility Product
Discussion.
When a salt, dissociating into univalent cations and anions,
is in equilibrium with its saturated solution, the Law of Mass
Action leads to the conclusion that the product of the concen-
trations of cation and anion is a constant for a given temperature,
provided the nature of the solvent undergoes no change. The
product of the ion concentrations when the solution is saturated is
called the solubility product. Thus :
[cation] [anion] = Ks, (1)
where the symbols "[cation]" etc., represent the concentrations.
Reference. Stieglitz, I 141 (read page 142 for criticism of
theory) ; Washburn, 298.
When the salt dissociates into polyvalent ions or into ions of mixed
valence, the relation is more complex. Cf. Washburn, 301.
It is possible to distinguish between two cases, as follows :
(1) When to a solution saturated with a given solid electrolyte
there is added a soluble salt containing a common ion, the product of
the concentrations of cation and anion momentarily becomes greater
than the solubility product. The solution is no longer in equilibrium
with the saturating solid salt and the latter is precipitated, until new
conditions of equilibrium are established. These new conditions
correspond to diminished solubility.
(2) When the concentration of one or both of the ions produced by
the saturating solid is decreased by any kind of physical or chemical
reaction, the product of the concentrations of cation and anion
momentarily becomes less than the solubility product. The
solution is no longer in equilibrium with the solid and fresh solid
dissolves until new conditions of equilibrium are established, the
latter corresponding to increased solubility.
Procedure.
Part 1. To a BaCl 2 solution in a test tube add concentrated HC1,
then add water. Explain.
Part 2. Repeat, using a CaCl 2 solution.
Part 3. To a saturated solution of NaCl, add concentrated HC1.
Part 4. To a saturated solution of HgCl 2 add a saturated solution
of NaCl.
Account for what happened in (1) and (3), by applying the theory
of the solubility product. Explain the very different results of (2)
and (4).
33
How might all these experiments be explained in the light of
Experiment 4?
Reference. Ostwald: Prin. Inorg. Chem., 675 (1908).
Part 5. Treat some freshly precipitated and washed AgCl with
(a) Na 2 S 2 O 3 solution; (b) with KCN solution (poison); (c) with
NH 4 OH.
Part 6. Treat some freshly precipitated calcium oxalate with
HC1.
Part 7. Shake a little HgO with a solution of KI. Note any color
change. Filter and test the nitrate with red litmus.
Part 8. Prepare some Cd(OH) 2 and wash thoroughly with water.
Shake with water and test the supernatant liquid with red litmus.
The solution should be neutral. To one-half of the Cd(OH) 2 add a
small amount of KNO 3 and shake again. Test the supernatant
liquid with red litmus. To the second half of the Cd (OH) 2 add a
little KI, shake and test the supernatant liquid* with red litmus.
Explain.
Reference. Ostwald: Prin. Inorg. Chem., 637 (1908).
EXPERIMENT 7
Solubility of Glass
Part 1. Phenolphthalein Test. Boil some clean, finely-powdered
glass with water in a beaker, then add a drop of phenolphthalein.
Explain.
Part 2. Eosin Test. "If a glass surface is brought into contact
with watery ether, it draws water from the solution and gives up
alkali to it. On the other hand, the orange-yellow solution of iod-
eosin in ether is changed by the alkali into red. Mylius, who had
previously used the color reaction for another purpose, has applied it
to the practical testing of glasses. Commercial ether is shaken up
with water at ordinary temperature until it is saturated with water.
It is then poured from the rest of the water and eosin is added in the
proportion of 0.1 g. to 100 cc. of the liquid. The solution is filtered
"Glass vessels are tested by pouring in the solution. The first
step is to clean the surface from any products of weathering which
may adhere to it, by carefully rinsing with water, with alcohol, and
lastly with ether. Immediately after the cleaning with ether, the
eosin solution is poured in, the vessel is carefully closed and the solu-
tion is allowed some twenty-four hours to do its work. It is then
emptied out and the glass rinsed with pure ether. The surface of the
glass is now seen to be colored red; and the strength of the color
furnishes an indication of the susceptibility of the glass to attack by
cold water."
Reference. Hovestadt (Everett) : Jena Glass and its Scientific
and Industrial Applications.
34
Following these directions make up 100 cc. of eosin solution.
Then test the surface of a new 50 cc. beaker and a new test tube as
described above.
Place a small sample of powdered glass in an 8-dram vial and add
some eosin solution.
Note the color the powdered glass assumes on standing twenty-four
hours. Note also the color of the walls of the vial.
If the powdered glass becomes colored, filter it and wash thoroughly
with water. Does the water remove the color? Pour off the water
and add alcohol. Does the alcohol remove the color?
Eosin as Indicator. Take a few cubic centimeters of the eosin
solution and add a few drops of dilute NaOH.
Part 3. Tetrachlorgallein Test. Add to a beaker of boiling dis-
tilled water a few drops of alcoholic tetrachlorgallein. Continue the
boiling and observe the color change. Make a blank test with fresh
distilled water.
EXPERIMENT 8
Corrosion of Metals
Discussion.
Many metals dissolve more or less readily in aqueous solutions,
appearing in the solution in the form of cations for at least a limited
time and displacing during this process an equivalent weight of some
other cation, usually hydrogen, from the solution. Thus zinc and
sulphuric acid give zinc sulphate and hydrogen; zinc and copper
sulphate give zinc sulphate and metallic copper, the salts and acids
being in solution. Under these circumstances the zinc is said to
corrode.
It is generally believed that the process of corrosion is electro-
chemical in nature. For-example, when zinc corrodes, two so-called
"electrochemical" reactions take place as follows:
(1) Metallic zinc gives zinc ions plus negative charges, the latter
being retained by the metal.
(2) Hydrogen ions in solution give hydrogen gas plus positive
charges, the latter neutralizing the negative charges on the metal.
Represented by symbols, these reactions may be written:
+26 (1)
+20 (2)
If one adds reactions (1) and (2), the total reaction becomes
Zn + 2H+ ^Zn + + + H 2 (3)
It is interesting to note that the anions appear to play no part
whatsoever.
Applying the Law of Mass Action to the two reactions given above,
it is possible to draw the following conclusions regarding the rate of
corrosion :
(1) A metal tends to corrode more readily in an aqueous solution
the greater its "electrolytic solution pressure," i. e., the greater the
driving force of reaction (1) or the greater the ion-forming tendency
of the metal.
35
(2) The smaller the concentration of the dissolving metal as ion
in the solution, the faster is the corrosion. The' ion concentration
may be kept low by the formation of complex ions, by hydrolysis, etc.
(3) The greater the hydrogen ion concentration in the solution the
faster the corrosion. Other things being equal, metals tend to cor-
rode more readily in acids than they do in alkaline solutions.
(4) Anything that reacts with and removes the discharged hydro-
gen tends to aid corrosion. Oxidizing agents may do this, in which
case they are called "hydrogen depolarizers." Note the part played
by air in the experiments; also the formation of nitrites in Part 2b.
(5) The absence of stable, difficulty soluble protecting films (pas-
sivity) favors corrosion.
(6) Miscellaneous. Metal should have irregularities, etc., in
surface to aid in setting up local "galvanic" couples. Also the
"overvoltage" for hydrogen should be low. These points belong
properly under electrochemistry and cannot be discussed here.
All the conditions favoring corrosion do not have to be fulfilled
simultaneously. Copper for example corrodes in aqueous ammonium
hydroxide in the presence of air. The electrolytic solution pressure
of copper is very small and the hydrogen ion- concentration in ammo-
nium hydroxide solution is very slight, but these conditions which
tend to prevent corrosion are more than offset by the fact that the
copper ion concentration in the solution is practically zero (complex
Cu(NH 3 ) 2 cations) and air oxidizes the discharged hydrogen under
the conditions of the experiment. The reaction as a whole may be
written :
Cu + 2NH 4 OH + O >Cu (NH 3 ) 2 (OH) 2 + H 2 O
Iron corrodes readily in moist air. Moisture is essential inasmuch
as it furnishes the hydrogen ions which are displaced by the iron, the
latter entering the solution in the form of ferrous ions. These are
almost immediately oxidized by air to ferric ions which combine with
the hydroxyl ions of the water to form hydrous ferric oxide. The
iron thus passes from solution and corrosion is thereby accelerated.
Carbon dioxide stimulates corrosion by dissolving in the film of mois-
ture and thus increasing the hydrogen ion concentration by the forma-
tion of H 2 CO 3 . Air increases corrosion by removing the dissolved
iron as explained above and by serving as the hydrogen depolarizer.
Procedure.
Part 1. Solubility of Metals in Acids and Alkalies, (a) Place
a small strip of copper foil in aqueous NH^OH in a test tube. Shake
thoroughly from time to time. Note the color change and explain.
(b) Experiments with concentrated H 2 SO4.
In a few cc. of concentrated H 2 SO4 test the solubility of cast iron,
iron wire, nickel wire, and copper wire. Set aside for an hour.
Dilute the acid five fold with water and repeat, using the same test
pieces. Dilute the acid until the rate of solution is rapid. Caution.
Dilute the acid properly.
36
(c) Experiments with concentrated HNO 3 .
In a few cubic centimeters of concentrated HNO 3 , test the solu-
bility of iron wire and nickel wire. Set aside for an hour. Repeat
with acid diluted twice. Why are metals often more readily attacked
by HNO 3 than they are by HC1?
Test the solubility of aluminum in caustic soda solutions. Explain.
Aluminum forms complex anions in NaOH.
Part 2. Solubility of Metals in Salt Solutions. Clean the metal
thoroughly, and, after weighing, set aside for ten days in a test tube
with 10 cc. of the salt solution. Cover up loosely with filter paper.
Shake from time to time. Clean the test piece and weigh again.
Record the time and note any change in the metal.
(a) Copper in 10 per cent NaCl, test alkalinity of filtered solution
at end.
(b) Cadmium in 10 per cent NH4NO 3 , test alkalinity of filtered
solution at end, and also test for nitrites.
(c) Iron in 10 per cent sodium tartrate. Test as before.
Reference. Chem. News, 90, 142 (1904).
Part 3. Passive Iron.
Discussion.
The passivity of iron is probably due to an adsorbed and stabilized
film of a higher oxide, the formula of which is possibly FeO 2 . The
oxide, which is very difficultly soluble in HNO 3 , is formed by certain
oxidizing agents such as HNO 3 , NO 2 , etc., or when iron is made
anode in an electrolytic cell through which a sufficiently high current
passes. Passivity is removed and activity is restored by destruc-
tion of the oxide film. Reducing agents may destroy the film or the
same thing may be done by making a passive rod cathode with a
sufficiently high current. Consult the Instructor.
Procedure.
(a) Make an iron rod passive in concentrated nitric acid (sp. gr. =
1.4). Wash in water carefully and dip in dilute HNO 3 (sp. gr. 1.2).
What happens?
(b) Having immersed the rod in the dilute acid, touch the rod with
a fresh (active) iron rod. Explain. Repeat, touching passive rod
with zinc.
(c) Immerse an active and a passive rod in dilute (1.2) HNO 3f
taking care to dip the active rod deeply and the passive rod only
slightly beneath the surface of the liquid. Connect the two rods out-
side of the cell with a copper wire. What happens?
(d) Repeat experiment (c), having a large surface of the passive
rod and only a small surface of the active one dipping into the acid.
To understand (c) and (d) see Bennett's paper, p. 220. (Schonbein's
experiments) .
(e) Immerse a passive rod in dilute acid and scratch the surface.
Does the rod become active?
Reference. Bennett and Burnham: Trans. Am. Electrochem.
Soc., 29, 217 (1916).
37
EXPERIMENTAL GROUP X
REACTION VELOCITY AND CATALYSIS
This group of experiments is designed to illustrate in a semi-
quantitative manner the Law of Mass Action and its bearing on the
velocity of chemical change. Simple experiments illustrating cata-
lysis are also included.
Standard References.
Bancroft: Papers in Jour. Phys. Chem. (1917 ).
Henderson: Catalysis and Its Industrial Applications (1918).
Herz: Ahren's Sammlung, 11, 103-145 (1906).
Jobling: Catalysis and Its Industrial Applications (1916).
Mellor: Chemical Statics and Dynamics (1609).
Ostwald: Uber Katalyse (2nd Ed. 1911).
Rideal and Taylor: Catalysis in Theory and Practice (1920).
van't Hoff: Lectures; Vol. 1, Chemical Dynamics (1898).
van't Hoff (Evan) : Studies in Chemical Dynamics (1896).
Woker: Die Katalyse (1915-16).
Law of Mass Action.
The rate at which chemical change occurs is a function of the
concentration of each of the substances taking part in the reaction.
The rate is also a function of the temperature and pressure and it is
affected by catalysts and by various other influences, such as light,
electrical and surface forces.
The law is illustrated by the reaction between bromic and iodic
acids
6 HI + HBrO, -^ HBr + 3 H 2 O + 3 I 2 ,
in which the course of the reaction can be followed color imetrically,
using starch as an indicator.
The rate at which iodine is set free is directly proportional to the
ion concentrations of iodine and bromate and to the square of the
concentration of hydrogen as ion. Clark: Jour. Phys. Chem., 10,
700 (1906). If one keeps the concentration of hydrogen ions con-
stant and does not allow the volume of the solution to vary, the
velocity with which iodine is liberated at any moment is expressed in
terms of the mass law by the equation
_^ = k(a x)(b x) (1)
dt
in which a and b refer respectively to the amount of iodine and bro-
mate present as ions at the beginning of the experiment and are there-
fore proportional to the initial quantity of HI and HBrO 3 , while x
refers to the amount of iodine or bromate ions used up and is accord-
ingly proportional to the quantity of free iodine liberated.
38
If the reaction is allowed to proceed for a relatively short time only
and in such a way that x is small by comparison with a and b, the
velocity equation takes the form
^ X =kab (2)'
whence, on integrating between the limits x = o and x = Xi ; t = o,
and t = t, the following expression results :
t = constant ^L (3)
ab
General Procedure.
In the experiments which follow iodide and bromate are mixed in
acid solution and the reaction is allowed to proceed until a definite
constant quantity of iodine is liberated, as determined by the forma-
tion of a definite "standard" blue color with starch as indicator. The
initial quantities of iodide and bromate are varied and the time
required to reach the standard blue is determined by means of a stop-
watch.
Under these experimental conditions, it is evident from equation
(3) that the time required to reach a standard blue at constant
temperature and volume varies inversely as the product of the initial
quantities of iodide and bromate, as long as the amount of iodine set
free is small. It is also obvious that this statement becomes less
exact as the depth of the standard blue becomes greater.
For comparison times, the relative values of a and b may be sub-
stituted for absolute values.
EXPERIMENT 1
Mass Action
Acid Mixture.
800 cc. distilled water
26 cc. N/2 HC1 (shelf)
20 cc. starch solution
To prepare the starch solution rub one gram of starch with 5 cc. of
cold water in a mortar; pour 150 cc. of boiling water over it, allow
the undissolved part to settle, and decant the supernatant liquid.
Standard Blue.
Take two 100 cc. bottles (glass stoppered) and in one make a
standard blue solution as follows:
80 cc. distilled water
2 cc. starch solution (described above)
3-6 drops "iodine mixture" (shelf)
Procedure.
Part 1 . In the test bottle place 80 cc. acid mixture
1 cc. N/2 KBrO 3 (shelf)
1 cc. N/2 KI (shelf)
in the order named. Add the KI quickly and take the time from the
moment it is added. Shake at the moment of adding KI and note
the time required for the solution to assume the same blue as the
standard. Run a parallel.
39
Notes.
Place the standard and the test bottle against a white background.
Avoid using a standard with too deep a blue. The time taken in
Part 1 should not exceed two minutes nor be less than one minute.
Be careful to work throughout at constant temperature (20 C.).
Record.
Part 2. 80 cc. acid mixture
2 cc. bromate
1 cc. iodide
Shake. Note time as before. Run a parallel.
Part 3. 80 cc. mixture Part 4. 80 cc. mixture
1 cc. bromate 2 cc. bromate
2 cc. iodide 2 cc. iodide
Shake. Note time. Run a Shake. Note time. Run a
parallel. parallel.
EXPERIMENT 2
Catalytic Effect of Acids
The effect of acids in accelerating certain chemical reactions is
roughly proportional to their electrical conductivity. The effect is
dependent primarily on the hydrogen ions.
Prepare a mixture as follows :
400 cc. water
10 cc. bromate (shelf)
10 cc. iodide (shelf)
10 cc. starch solution
Part 1. Take 80 cc. of the above mixture in a 100 cc. bottle, add
2 cc. N/2 HC1. Shake. Note time required. Run a parallel.
Part 2. Take 80 cc. of mixture and 2 cc. of N/2 H 2 SO 4 . Shake.
Note time required. Run a parallel.
Part 3. Take 80 cc. of mixture and 2 cc. of N/2 CH 3 COOH.
Note time required. Shake. Run a parallel. Explain.
EXPERIMENT 3
Catalytic Effect of Ferrous Sulphate
Mixture of 160 cc. H 2 O.
8 cc. KI (shelf)
8 cc. KBrO 3 (shelf)
4 cc. starch solution
Part 1. Take 80 cc. of the above mixture and 10 cc. of N/2 acetic
acid. Shake. Note time required to become blue.
Part 2. Take 80 cc. of the mixture and 10 cc. of N/2 acetic acid
to which is added one drop of neutral saturated FeSO4. Proceed as
before. Explain.
40
Part 3. To 25 cc. of an extremely dilute solution of chromic acid
(CrO 3 ) add a little starch solution.
(a) To 5 cc. of this solution add 2 to 3 drops of KI solution. Note
time as before.
(b) To 5 cc. of the solution add KI as before and a little iron dust.
Note time.
(c) To another 5 cc. portion add KI and a few drops of a ferrous
sulphate solution. Note time.
(d) To another 5 cc. portion add KI and a few drops of ferric
sulphate solution. Note time.
Part 4. (a) Mix in the following order: Dilute CrO 3 solution,
ferrous sulphate solution and starch; shake and wait ten
minutes; then add KI. Note time to reach standard blue
after adding KI.
(b) Mix in the following order: CrO 3 solution, KI solution and
starch; wait ten minutes ; then add ferrous sulphate. Note
time after adding FeSO 4 .
Compare (a) and (b) and explain.
EXPERIMENT 4
Hydrolysis of an Ester Catalysis
Place 50 cc. of distilled water and 5 cc. of ethyl acetate in a clean,
glass stoppered bottle. Shake thoroughly and titrate duplicate
samples (2 cc.) with N/10 NaOH, phenolphthalein as indicator.
In a second bottle place 50 cc. N/2 HC1 plus 5 cc. ethyl acetate.
Shake and titrate as before.
Set both bottles aside for 24 hours (shaking occasionally) and again
titrate duplicate samples (2 cc.). .
Note differences and explain. How is this phenomenon used to
measure the strength of acids?
EXPERIMENT 5
Reactions in Heterogeneous Systems
Part 1. Size of Particles. Whenever one of the reacting sub-
stances is a solid, the speed of the reaction is a function of the surface
area of the solid, or more accurately, of the surface per unit weight of
solid (specific surface). The specific surface, in turn, is a function of
the size of the particles and increases rapidly as the particles become
smaller. Read W9 Ostwald: Grundriss der Kolloidchemie, 30
(1912).
Prepare about 2 grams of finely divided copper by placing some
granulated zinc in a concentrated solution of CuSO 4 . Shake from
time to time to remove the finely divided copper from the zinc.
After most of the copper has been precipitated, remove the zinc,
wash the precipitate with water and dry in an air bath. Mix the
finely divided metal with powdered sulphur and ignite cautiously
with a match. What is formed?
41
Dissolve sulphur in CS 2 and into this solution dip a clean strip of
copper. What is the substance formed on the copper?
Show how this experiment illustrates the principle discussed.
Part 2. Protecting Films, (a) Clean a strip of aluminum foil by
immersing it in 10 per cent NaOH. Rinse and plunge the wet
metal quickly into clean mercury. Hold it there until amalgamated.
Remove and rub off the excess of mercury adhering to the aluminum,
then expose to the air. What happens? Explain.
(b) Place freshly amalgamated aluminum in contact with warm
water. What happens? Compare with sodium.
(c) Dip a rod of metallic magnesium into warm water. What
happens?
(d) Dip a rod of metallic magnesium into warm NH 4 C1 solution.
What happens? Explain.
The passivating films might be regarded as negative catalysts.
References. Ostwald: Prin. Inorg. Chem., 560 (1900); Wis-
licenus: Jour. Praktische Chemie, (2), 54, 41 (1896).
Note. The amalgamated aluminum may be prepared by cleaning
the metal in 10 per cent NaOH, rinsing carefully and then dipping the
wet metal into dilute mercuric chloride.
Part 3. Reactions between Solids. Incompatible Hydrates.
Use small quantities in proportions approximately equivalent.
Weigh out roughly, except in (a), where a few crystals are enough.
(a) Grind together in a mortar HgCl 2 + KI.
(b) " " "" " Na 2 SO 4 -10H 2 O + NH 4 NO 3 .
(c) " " " " " (NH 4 ) 2 SO 4 + NaNO 3 .
<d) " " " " " K 2 S0 4 + NH 4 N0 3 .
(e) " " "" " MgS0 4 -7H 2 O + NH 4 N0 3 .
(f) " '" " " " CuSQ 4 -5H 2 O + NH 4 NO 3 .
References, van't Hoff: Studies in Chem. Dynamics, 173
(1896); Schiff: Liebig's Annalen, 114, 68 (1860).
(g) Grind together 5 g. NHCNS and 10 g. crystalline barium
hydroxide Ba(OH) 2 8H 2 O. What happens? Explain.
Part 4. Halogen Carriers. Support a 250 cc. distilling flask upon
a ring stand and connect its side arm with a funnel the mouth of which
dips just below the surface of a caustic soda solution." Place in the
flask 2 cc. of bromine. Provide a cork stopper for the flask. Now
pour into the flask 15 cc. of benzene. Work at the hoods.
Test for HBr with ammonia fumes. Then add about a quarter of
a gram of iron powder. Again cautiously test for HBr. Be re'ady
to stopper the flask and leave stoppered until the reaction is over.
42
EXPERIMENTAL GROUP XI
SAPONIFICATION OF AN ESTER
The experiment which follows is designed to demonstrate quanti-
tatively the Law of Mass Action as applied to the kinetics of a simple
irreversible reaction. The reaction to be studied is a reaction of the
second order.
References.
Mellor : Chemical Statics and Dynamics (1909) .
Warder; Am. Chem. Jour., 3 340 (1882).
F, 270-272; G, 246-248; OW, 246-252, etc.
EXPERIMENT 1
Saponification of Ethyl Acetate
Solutions Required.
A. N/20 NaOH (free from carbonates) .
A carbonate-free normal solution of NaOH is supplied (shelf).
From this prepare a solution slightly stronger than N/20 being careful
not to waste any of the carbonate-free sodium hydroxide. Make up
two liters of solution and standardize against an acid of known titre
(shelf). Finally, dilute until the solution is exactly N/20 and again
standardize to make sure that the work has been done correctly.
The normal titre of the solution should not differ from the required
value (N/20) by more than 1 per cent. Phenolphthalein as indicator.
Save the residue of this solution for use in Group XIV.
B. N/20 HC1.
Prepare two liters and standardize carefully against N/20 NaOH.
Protect the burette containing the latter by means of a soda-lime tube.
Phenolphthalein as indicator. Save the residue of .this solution for
use in Group XIV.
C. N/20 Ethyl Acetate.
Ethyl acetate being difficult to obtain pure, it is necessary to pre-
pare this solution as follows: To 800 cc. distilled water, contained
in a liter glass stoppered graduated cylinder, add 5 cc. of redistilled
ethyl acetate (special reagent). Stopper quickly to prevent loss of
ester by volatilization, and shake thoroughly to dissolve.
In a 100 cc. glass stoppered bottle place exactly 25 cc. of N/20
sodium hydroxide (burette) and to this add from a pipette (cali-
brated) exactly 10 cc. of the ethyl acetate solution. Replace the
stopper quickly and securely and heat the bottle in a water bath for
thirty minutes, or until the ester is completely saponified. Remove
from bath and cool, add a few drops of phenolphthalein and determine
43
the excess of sodium hydroxide by titration with N/20 HC1. Run in
duplicate. The solution should be more concentrated than N/20
at this point.
Then calculate the volume of water necessary to dilute the ethyl
acetate exactly to N/20, allowing for the amount already withdrawn.
Saponify as before in order to verify the work. The ethyl acetate
solution should now be N/20 +_1 per cent.
Procedure.
Part 1. Adjust the automatic thermostat to 25 C., or if this is not
available use a large pan or beaker of water kept at 25 + 0.1 C.
Measure the temperature with a thermometer graduated to tenths.
Stir with compressed air.
Add exactly 250 cc. of N/20 NaOH to a 500 cc. glass stoppered
bottle. Place in the thermostat and shake occasionally. In a glass
stoppered measuring flask (250 cc.) place an equal amount of N/20
ethyl acetate. Place in a thermostat.
When both solutions have reached 25 C. quickly pour ester into a
bottle containing NaOH, replace the stopper and shake instantly.
Start the stop-watch at the moment of mixing and at the same time
read the hour and minute on a watch, in case the stop-watch should
prove faulty. The reaction begins at the instant of mixing.
Pipette out 10 cc. samples at the following times:
2, 3, 5, 8, 12, 16, 20, 25, 30, 40, 50, 60, 80, 120 minutes.
At the desired moment stop the reaction by emptying the pipette
into an accurately known volume (about 7 cc.) of N/20 HC1 con-
tained in 50 cc. of water -f- 1 drop phenolphthalein in an Erlenmeyer
flask. Add the acid from a burette.
Determine as soon as possible the excess of N/20 HC1 by titrating
with N/20 NaOH. cf. Group X Experiment 5.
Shake the bottle in the thermostat every two minutes.
Precautions.
This is an experiment requiring accurate manipulation. Burettes
and pipettes should be calibrated and placed in cleaning mixture
for at least twenty-four hours before use. While in use, see that they
are kept filled with solution or distilled water, because drying in air
causes glassware to acquire a grease-like film. When reading bur-
rettes try to estimate to hundredths of a cubic centimeter.
Computations.
From the data recorded during the experiment compute the num-
ber of cc. of N/20 NaOH consumed by the ester during each of the
time intervals. If this experiment is carried out properly these
values should rise from zero at the beginning to nearly 5 cc. at the end.
Draw a curve between cc. of NaOH consumed as ordinates and
time in minutes as abscissae.
44
Using the equation of a second order reaction
k = _* (1)
at (a x)
compute values of the velocity coefficient k corresponding to the
different times.
Part 2. Dilute exactly 250 cc. of the ethyl acetate solution to 500
cc. making it N/40. Do the same with 250 cc. of the NaOH. Then
repeat Part 1 and draw a curve between cc. of N/40 NaOH used up
and time in minutes.
Compare the curves obtained in Parts 1 and 2. Determine in
each case the time required for one-half of the original NaOH to
disappear. How do these times compare and how are they related
to the initial concentrations of ester and base?
From these results determine the order of the reaction.
45
EXPERIMENTAL GROUP XII
THE STUDY OF A REACTION
The experiments which comprise this group constitute a detailed
study of the reaction between oxalic acid, potassium permanganate
and sulphuric acid, in aqueous solution:
5 (COOH) 2 + 2 KMn0 4 + 3 H 2 SO 4 = 2 MnSO 4 + K 2 SO 4 +
10 CO 2 + 8 H 2 O
The reaction is familiar to every chemist because of the important
part it plays in volumetric analysis.
Reference. Harcourt and Esson: Jour. Chem. Soc., 20, 460
(1867).
Discussion.
The reaction as written above is really the result of a series of
simpler reactions. The reaction will be studied by means of velocity
determinations made by ascertaining how much permanganate is used
up under definite experimental conditions, and by systematically
varying these conditions.
The important factors constituting the experimental conditions
are the following:
(1) Amount of KMnO 4 .
(2) Amount of oxalic acid.
(3) Amount of H 2 SO 4 .
(4) Amount of MnSO 4 .
(5) Amount of K 2 SO 4 .
(6) Amount of CO 2 in solution.
(7) Volume of solution (amount of water).
(8) Temperature.
(9) Pressure.
(10) Illumination.
(11) Time.
Factors 1 to 7 inclusive are concentration factors. In the present
study temperature, pressure, and illumination are kept as nearly
unchanged as possible without special precautions and the experi-
ments are carried out at constant volume. The work is done in
open vessels so that factor (6) is practically constant throughout.
Arbitrary values are assigned to four of the first five factors while the
fifth is being varied in a systematic manner. The time during which
the reaction takes place (factor 11) is four minutes.
46
Solutions.
Prepare the following solutions:
(1) Potassium permanganate 1. 58 g. per liter.
(2) Oxalic acid 3.15 " " "
(3) Sulphuric acid 1.47 " " "
(4) Manganous sulphate 2.23 " " "
(5) Potassium sulphate (500 cc.) 0.87 " " "
(6) Potassium iodide (500 cc.) 25.00 " " "
(7) Sodium thiosulphate 2.48 " " "
The first five of the above solutions are of such strength in every
case that one liter contains - of the number of molecules taking
i\J\J
part in the reaction. Thus, in the case of the sulphuric acid:
3 x M. W. = 3 x 98 = 294, and
?^ = 1.47 (grams per liter)
200
Comparison of Solutions.
(a) Titrate the oxalic acid against the permanganate in
strongly acid solution. These solutions must be equivalent i. e. 10 cc.
KMn0 4 = 10 cc. (CO 2 H) 2 .
(b) Decompose 5 cc. of the KMnO 4 solution by adding 15 cc. of
KI solution. Determine the amount of iodine liberated, by titrating
with the thiosulphate. Five cc. of the KMnO 4 should require about
25 cc. of the thiosulphate.
In titrating, the iodine with thiosulphate, do not add the starch
indicator until most of the iodine has been reduced. When the
solution has acquired a pale straw color, add the starch. A blue
color should appear. Practice this titration until satisfactory end-
points are obtained.
The starch indicator may be prepared by rubbing a gram of arrow-
root starch into a paste with cold water, and to this paste adding
about 200 cc. of boiling water.
Experimental Procedure.
The required amounts of all of the reacting substances except the
permanganate, are mixed, diluted to 100 cc. and placed in Erlen-
meyer flasks. These are then allowed to come to the same tempera-
ture. Take the temperature of each mixture and record it. The
permanganate is then added quickly, the flask shaken immediately
and the time taken with a watch. The reaction commences with the
addition of the permanganate.
After exactly four minutes have elapsed, the excess of undecom-
posed KMnO 4 is destroyed by adding an excess of the KI solution
(10 to 15 cc.) and the iodine set free is determined with thiosulphate
solution, using starch as indicator.
47
The reactions are as follows :
KMnO 4 + 5 KI + 4 H 2 SO 4 = 3 K 2 SO 4 + 51+4 H 2 O + MnSO 4
and
2 Na 2 S 2 3 + 2 I = 2 Nal + Na 2 S 4 O fi
The amount of decomposed permanganate is proportional to the
volume of thiosulphate used in reducing the iodine, and we may thus
determine the permanganate used up in the reaction, the thiosulphate
titre of the permanganate solution being known.
Part 1. Variation of Sulphuric Acid.
KMnO 4
H 2 C 2 4
H 2 S0 4
(cc)
(cc)
(cc)
(a) 5
5
5
(b) 5
5
10
(c) 5
5
15
(d) 5
5
25
(e) 5
5
40
(f) 5
5
60
Part 2.
Variation of Manganous
Sulphate.
KMn0 4
H 2 C 2 4
H 2 SO 4
(cc)
(cc)
(cc)
(a) 5
5
15
(b) 5
5
15
(c) 5
5
15
(d) 5
5
15
(e) 5
5
15
(f) 5
5
15
(g) 5
5
15
Part 3.
Variation of Oxalic Acid.
KMnO 4
H 2 C 2 4
H 2 S0 4
(cc)
(cc)
(cc)
(a) 5
1
25
(b) 5
2
25
(c) 5
3
25
(d) 5
4
25
(e) 5
5
25
(f) 5
6
25
(g) 5
7
25
(h) 5
8
25
(i) 5
9
25
(j) 5
10
25
(k) 5
11
25
(D 5
12
25
(m) 5
15
25
(n) 5
20
25
(o) 5
25
25
(p) 5
30
25
(q) 5
50
25
MnSO 4
(cc)
5
5
5
5
5
5
MnSO 4
(cc)
1
3
5
8
10
15
MnSO 4
(cc)
10
10
10
10
10
10
10
10
10
10
10
10
10
10
10
10
10
48
Part 4. Variation of Potassium Sulphate .
KMnO 4 H 2 C 2 O 4 H 2 SO 4 MnSO 4 K 2 SO 4
(cc) (cc) (cc) (cc) (cc)
(a) 5 5 15 10 5
(b) 5 5 15 10 25
(c) 5 5 15 10 75
Part 5.
To 5 cc. of the KMnO 4 solution in a test tube add 10 cc. of the
MnSO 4 solution. Shake and allow to settle. What is the precipi-
tate? What color is the supernatant liquid? Test it with litmus.
Computations and Curves.
From your data, compute (in cc.) the amount of potassium per-
manganate used up in the reaction after four minutes. Call these
numbers "y." Tabulate the values of y along with the correspond-
ing values of the substance undergoing variation, called "x." With
values of x as abscissas and of y as ordinates, draw four curves pictur-
ing the results obtained in each of the four parts.
49
EXPERIMENTAL GROUP XIII
REVERSIBLE REACTIONS AND CHEMICAL EQUILIBRIUM
The following experiments deal particularly with chemical reac-
tions which occur readily in both directions and are therefore dis-
tinctly reversible, tending to reach a condition of equilibrium.
Several examples of reactions of this type have already been studied,
notably in Groups VIII and IX.
References. See under Group X.
EXPERIMENT 1
Homogeneous Chemical Equilibrium
This is illustrated very simply by the equilibrium between the
reciprocal pairs, ammonium thiocyanate-ferric chloride and ferric
thiocyanate-ammonium chloride. The amount of ferric thiocyanate
formed in solution may be estimated by the intense red-brown color
that the undissociated salt imparts.
"If the reaction is represented by
3 NH 4 CNS + FeCl 3 = Fe (CNS) 3 + 3NH 4 C1
and the amount of ferric sulphocyanate is judged by the depth of
color of the solutions, the reaction between equivalent quantities
must be regarded as incomplete."
Procedure.
The following solutions will be found as shelf reagents :
Solution A. Ammonium thiocyanate 38 g. NH 4 CNS per liter.
Solution B . A mixture of the following :
Ferric chloride 30 g.
Concentrated hydrochloric acid 115 cc.
Water, 1000 cc.
Take equal volumes of solutions A and B, 5 cc. of each. Dilute to
2 liters. Stir thoroughly.
The solution should be a definite orange in color.
Divide into five 400 cc. portions.
To the is added Color becomes
First portion 5 cc. NH 4 CNS solution ?
Second portion 5 cc. FeCl 3 solution ?
Third portion 50 cc. sat. NH 4 C1 solution ?
Fourth portion solution of HgCl 2 ?
The first portion is kept for comparison. Explain all results.
Reference. Miller and Kenrick: Jour. Am. Chem. Soc. 22, 292
(1900).
50
EXPERIMENT 2
Heterogeneous Chemical Equilibrium
BaSO 4 + Na 2 CO 3 = BaCO 3 + Na 2 SO 4
Part 1. In an evaporating dish over a water bath heat together
1/100 molecular weight of BaSO 4 , 1/6 molecular weight of Na 2 CO 3 ,
and 100 cc. of water. Stir constantly and replace water that
evaporates.
After heating for one hour, test the supernatant liquid for sul-
phates, taking care to decompose the carbonates before testing.
Wash the residue by decantation and finally on a filter until the
wash water gives no test for carbonates.
After washing to free from soluble carbonates test the residue.
What is it? Explain.
Part 2. Take 1/100 molecular weight of BaCO 3 , add 100 cc. of
water, then add 1/6 molecular weight of Na 2 SO 4 . 10 H 2 O.
Follow the same procedure as in Part 1.
Test the supernatant liquid for carbonates.
Test the residue, after washing, with HC1. Is there a residue after
treating with HC1? What is this residue?
Explain.
Reference. Walker, 290 (1913); Mellor: Chem. Statics and
Dynamics, 243 (1909).
EXPERIMENT 3
Heterogeneous Chemical Equilibrium
. 2 NaCl + H 2 SO 4 = 2 HC1 + Na 2 SO 4
Part 1. Concentrated H 2 SO 4 is poured into its own volume of a
saturated solution of sodium chloride in a small evaporating dish.
Warm very gently in the hood. Set aside until crystallization begins,
then pour the liquid off and dry the crystals on a piece of unglazed
porcelain.
The product is sodium sulphate containing hydrochloric acid. To
test for the former it is necessary to get rid of the latter. Dissolve in
the least possible amount of water and precipitate the sulphate by
adding absolute alcohol.
Filter, and after drying the residue, test for sodium chloride and
sulphate.
Reference. Miller and Kenrick: loc. cit. (Cf. Expt. 1).
Part 2. Cover a crystal of Na 2 SO 4 . 10 H 2 O on a watch glass with
concentrated HC1. After the reaction is complete pour off the acid
on an unglazed porcelain plate, as before.
To analyze the resulting product warm gently in a test tube and
remove any HC1 fumes from the tube by blowing out with air. Then
dissolve in water and test for sodium chloride and sulphate.
51
EXPERIMENT 4
Distribution of a Base between Two Acids
Weigh out 5 grams of Ba(OH) 2 8 H 2 O and dissolve in 50 cc. of
water. Make up a mixed solution of H 2 SO 4 and HC1 (obtained by
calculation and reference to acid tables) containing just enough of
each acid to neutralize all the Ba(OH) 2 in the first solution. Dilute
this solution to 50 cc. Then mix the two solutions. Shake well.
After settling, test the supernatant liquid for barium. How does the
base distribute itself between the competing acids? Why is H 2 SO 4 ,
the "weaker" acid, more active in this case? Define the term
"weaker" acid.
Explain your results.
EXPERIMENT 5
Addition of a Common Ion
Discussion.
The dissociation of a weakly ionized acid or base is greatly reduced
by the addition of one of its neutral salts. According to the Mass
Law, the product of the concentrations of the two ions of the acid is
proportional to the concentration of its undissociated portion and
since the concentration of the anion is greatly increased by the addi-
tion of the neutral salt, the ratio of the concentration of the H ion to
that of the undissociated acid must decrease in the same proportion.
In the following experiment, in order to show the difference between
the concentration of the hydrogen ion in the two cases, use is made of
the relative effect of the acid, in the absence and presence of its
neutral salt, in accelerating the bromate-iodide reaction.
Procedure.
Make a standard blue solution. See X, Experiment 1.
Make a solution as follows: 175 cc. of water, 5 cc. of N/2 KBrO 3 ,
5 cc. of N/2 KI, and 3 cc. of starch solution.
Part 1. Solution (1) Take 90 cc. of the above mixture.
Solution (2) Then 25 cc. of N/2 acetic acid and mix with 25 cc.
water.
Mix (1) and (2) and note time to reach the standard blue.
Part 2. Solution (1) Take 90 cc. of the above mixture.
Solution (2) Then 25 cc. of N/2 acetic acid and 25 cc. of N/2
sodium acetate mixed.
Mix (1) and (2) and note time as before.
52
EXPERIMENTAL GROUP XIV
INDICATORS
This group of experiments is divided into two parts, the first com-
prising a study of several of the more common indicators employed
in acidimetry and alkalimetry, particularly methyl orange and phe-
nolphthalein. The second part comprises the rough determination of
hydrogen ion concentration by the use of a set of indicators.
References.
Glaser: Die Indikatoren (1901).
Noyes: Jour. Am. Chem. Soc., 32, 816 (1910).
Prideaux: Theory and Use of Indicators (1917).
Thiel: Der Stand der Indikatorenfrage, Ahren's Sammlung 16,
307-422 (1911).
SUB-GROUP I
STUDY OF INDICATORS
Procedure.
Prepare the following solutions :
N/20 HC1 1 liter.
N/20 Acetic acid 500 cc.
N/20 NaOH 1 liter.
N/20 NH 4 OH 500 cc.
The bases should be free from carbonates. Cf. Group XI.
Use calibrated burettes and make sure that these are absolutely
clean. Cf. Group I. Before taking readings allow burettes to drain
exactly two minutes and use every precaution in titrating. Protect
NaOH from CO 2 in the air. Never leave the burettes standing
partly empty exposed to the air, but keep them filled with distilled
water" when not in constant use.
Always use the same amount of indicator each time.
Prepare a standard comparison end-point for use with each indica-
tor, and match this shade and color carefully each time.
Keep the temperature as constant as possible.
EXPERIMENT 1
Comparison of Indicators
Part 1. Titrate 10 cc. HC1 with NaOH. Dilute acid to 50 cc.
each time.
(a) Phenolphthalein (Ppn) in acid.
(b) Methyl orange (MO) in acid.
(c) Purified litmus (special reagent) in acid.
53
Note. Acid and base should be very closely equivalent with
litmus. Explain different readings obtained in (a), (b), and (c).
Cf . Experiment 4, this group.
Part 2. Titrate 10 cc. of HC1 with NH 4 OH. Dilute to 50 cc.
(a) MO in acid.
(b) Ppn in acid.
Part 3. Titrate 10 cc. acetic acid with NaOH. Dilute to 50 cc.
(a) Ppn in acid.
(b) MO in acid.
Part 4. Titrate 10 cc. acetic acid with NH 4 OH. Dilute to 50 cc.
(a) MO in acid.
(b) Ppn in acid.
From your results draw conclusions as to the proper indicator to
use under the various conditions. MO as indicator seems to behave
as a weak base; Ppn, as a weak acid. Cf. Waddell: Jour. Phys.
Chem.,2,171 (1898).
EXPERIMENT 2
Indicators as Acids or Bases
Indicators are -weak acids or weak bases. Is there therefore any
difference in the amount of acid or base required for neutralization,
depending on whether the indicator is placed in the acid or in the
base?
Part 1. Titrate the number of cc. of NaOH required to neutralize
the acid in Experiment 1, Part la, with HC1, adding the indicator to
the base. Dilute to 50 cc.
Part 2. Titrate the number of cc. of NaOH required to neutralize
the acid in Experiment 1, Part Ib, with HC1, adding the indicator to
the base. Dilute to 50 cc.
EXPERIMENT 3
Effect of Heat on Indicators
Part 1. Take 10 cc. of HC1, dilute to 50 cc., add Ppn and nearly
neutralize with NaOH. Then heat to 80-90 and complete the
titration at this temperature.
Part 2. Take 10 cc. of HC1, dilute to 50 cc., add MO and nearly
neutralize with NaOH. Heat to 70-80 and complete the titration
at this temperature.
EXPERIMENT 4
Effect of Volume
The neutral (end-point) color of an indicator occurs at a definite
concentration of hydrogen ions in the solution. Study the table in
Washburn 333 and posted in the laboratory. The hydrogen ion
concentration of the end-point is different for the different indicators.
With this in mind and remembering that concentration is defined as
mass divided by volume, perform the following :
54
Part 1. Take 10 cc. of HC1, dilute to 250 cc., add Ppn and titrate
with NaOH.
Part 2. Take 10 cc..of HC1, dilute to 500 cc., add Ppn and titrate
with NaOH.
Part 3. Take 10 cc. of HC1, dilute to 250 cc., add MO and titrate
with NaOH.
Part 4. Take 10 cc. of HC1, dilute to 500 cc., add MO and titrate
with NaOH.
EXPERIMENT 5
Phosphoric Acid
Discussion.
Phosphoric acid dissociates in three stages:
(1) H,PO 4 =H+ + H 2 P04
(2) H 2 PO = H+ + HPO (somewhat)
(3) HPO7 = H+ + PO 4 = (very slightly)
Accordingly phosphoric acid is really a fairly strong monobasic
acid, but as a dibasic acid it is weak.
On adding NaOH, the reaction first takes the following course:
H 3 PO 4 + NaOH = NaH 2 PO 4 + H 2 O
The ions are Na + and H 2 PO 4 . Referring to stage 2 in the ioniza-
tion of phosphoric acid it is seen that H 2 PO 4 also ionizes somewhat
into H+ and HPO 4 = . The hydrogen ions are so few, however, that
their concentration is not sufficient to turn MO red, but is sufficient to
render Ppn colorless. On adding a second molecule of NaOH, the
reaction becomes:
Na 2 HPO 4 + NaOH = Na 2 HPO 4 + H 2 O
The ions are now Na+ and HPO 4 =. Since HPO 4 = gives scarcely
any H + and PO 3 = ions (stage 3) and does not react readily
with NaOH, a very small quantity of base in excess of two equivalents
will give a solution sufficiently alkaline to turn Ppn pink. Read
Stieglitz, I 103. Do your results check with the theory?
Procedure.
Part 1. Titrate 10 cc. of M/20 phosphoric acid (shelf) with
NaOH, as indicator.
Part 2. Titrate 10 cc. of M/20 phosphoric acid with NaOH.
Ppn as indicator. Walker, 359 (1913).
EXPERIMENT 6
Titration of Carbonates. Effect of CO 2
Dissolve 0.3 gram of NaiCO 3 in 120 cc. of H 2 O
Part 1. Take 20 cc. of this solution and titrate with N/20 HC1
(MO) as indicator. When the end-point is reached, heat to boiling.
55
Part 2. Take 20 cc. of this solution, heat to boiling, and titrate
with N/20 HC1 (MO) as indicator.
Part 3. Take 20 cc. of this solution and titrate with N/20 HC1
with Ppn as indicator. When the end-point is reached, heat to
boiling.
Part 4. Take 20 cc. of this solution, heat to boiling, and titrate
with N/20 HC1 with Ppn as indicator.
Part 5. To a solution of Na 2 CO 3 add Ppn.
To a solution of Na 2 CO 3 add MO.
To a solution of NaHCO 3 add Ppn.
To a solution of NaHCO 3 add MO.
Part 6. To a dilute solution of NaOH add Ppn. Pass CO 2 into
the solution. Does the pink color disappear? Does it reappear on
passing in more CO 2 ?
Repeat, using MO. Explain all results.
Part 7. To a solution of Na 2 CO 3 add Ppn, then pass CO 2 into the
solution.
Repeat, using MO. Explain.
Hint. Consider the reaction as occurring in two stages:
2 NaOH + C0 2 = Na 2 CO 3 + H 2 O
Na 2 C0 3 + C0 2 + H 2 = 2 NaHCO 3
Compare with phosphoric acid in Experiment 5 above. Explain.
EXPERIMENT 7
Miscellaneous
Part 1. To 20 cc. of alcohol plus a few drops of phenolphthalein
add several drops of aqueous ammonia, and shake the solution.
Water is added slowly up to 5 cc. Then add 25 cc. of alcohol.
Explain.
References. Elements of Phys. Chem., 295 (1907).
Hildebrand's explanation, Jour. Am. Chem. Soc., 30, 1914 (1908).
Jones' explanation, Am. Chem. Jour., 18,377 (1896).
Part 2. Add a little Ppn to concentrated H 2 SO 4 . Add a little
Ppn to concentrated NaOH.
Reference. McCoy: Am. Chem. Jour., 31, 516 (1904).
Part 3. Divide a dilute acetic acid solution into two portions, and
add MO to each. To one add sodium acetate. Show that this solu-
tion is still acid to litmus. Explain. Cf. Stieglitz, I, 113.
SUB-GROUP 2
HYDROGEN ION CONCENTRATION
Discussion.
Read Prideaux on Indicators, or the more recent general texts, such
as Washburn or Lewis. All aqueous solutions, whether acid, neutral
56
or alkaline, contain both hydrogen and hydroxyl ions, the product of
their concentrations being roughly 1.0 x 10-14 at 25 C. In neutral
solutions these concentrations are equal and lie close to 10~ 7 gram ions
per liter. A solution normal with respect to hydrogen ions would
represent a hydrogen ion concentration of 10 or unity; a tenth-
normal solution a hydrogen ion concentration of 10" 1 or 1/10 and so
on ; a solution normal with respect to hydroxyl ions would represent
a hydrogen ion concentration of 10~ 14 . It follows therefore that
The degree of acidity or alkalinity of any solution may be expressed
in terms of its hydrogen ion concentration.
Sorensen has suggested that the hydrogen ion concentration be
represented in terms of an index represented by the symbol PJJ
This "Index" is the common logarithm of the hydrogen ion concentration
with the minus sign omitted. Thus if PH = 1, the solution has a
hydrogen ion concentration of 10 -1 and is tenth-normal; PH = 7
would represent a neutral solution, and so on.
When PH is greater than 7, the solution is alkaline; when it is less
than 7, the solution is acid, provided one is dealing with so-called
"room" temperatures (18-25 C.).
The most accurate and reliable method of measuring the hydrogen
ion concentration of a certain solution is an electrical one employing
a hydrogen electrode. This electrometric method is studied in the
laboratory course in electrochemistry, Course 56b. Since, however,
indicators undergo their characteristic color changes and show their
neutral colors at very definite hydrogen ion concentrations, a set of
indicators may be used to measure hydrogen ion concentration, pro-
vided the critical or neutral color concentration is known for each
indicator and the range covered is sufficiently great.
Cf. Clark: The Determination of Hydrogen Ions (1920).
The indicator method may be carried out by comparing the
unknown solution with a set of standard solutions of known hydrogen
ion concentration and determining with which of these standard solu-
tions the unknown is most nearly identical. The following standard
solutions are available (special reagents) :
Standard Solutions of Known Hydrogen Ion Concentration
Reference. Noyes: Jour. Am. Chem. Soc., 32, 822 (1910).
(1) PH = 3. Mix 570 cc. of N/10 acetic acid with 430 cc. of
water. Acetic acid contains 6 g. per liter.
(2) P H = 4. Dissolve 2.7 g. CH 3 COONa-3H 2 O in 1 liter of N/10
acetic acid. CH 3 COONa.3H 2 O is crystalline sodium acetate.
(3) PH = 5. Dissolve 15.0 g. of CH 3 COONa-3H,O in 500 cc. of
water, and add 500 cc. N/10 acetic acid.
(4) PH = 6 to 11. Make up a tenth molecular solution of Na 2
HPO 4 '12H 2 O. Prepare also N/10 HC1 and N/10 NaOH (free from
carbonate). Mix the solutions as follows:
57
PH
Mix
6
7
8
9
10
11
600 cc.
700 cc.
1000 cc.
1000 cc.
1000 cc.
1000 cc.
N/10 NaH
N/10
N/10
N/10
N/10
N/10
:P0 4 + 500 cc.
+ 3,50 cc.
+ 47 cc.
+ 5 cc.
+ 3.6 cc.
+ 36 cc.
N/10 HC1.
N/10 "
N/10 "
N/10 "
N/10 NaOH.
N/10 NaOH.
These solutions will be found on the reagent shelf.
For other mixtures giving solutions of known hydrogen ion concen-
tration consult Walpole: Biochemical Journal 5, 207 (1911).
Indicator Solutions.
0.5 per cent thymolphthalein (Tpn) in alcohol.
0.5 per cent phenol phthalein (Ppn) in alcohol.
0.5 per cent rosolic acid (RA) in 50 per cent alcohol.
0.1 per cent methyl red (MR) in water.
0.1 per cent methyl orange (MO) in water.
0.1 per cent purified litmus (L) in water.
Extract of cochineal (Coc) in water.
EXPERIMENT 1
To Determine Hydrogen Ion Concentration Corresponding to Neutral
or Critical Color of Indicators
Noyes: Jour. Am. Chem. Soc., 32, 824 (1910).
Obtain twenty-seven test tubes, clean and dry, then place in nine
groups of three. To each test tube add 10 cc. of the various standard
solutions of known hydrogen ion concentration and to these 0.1 cc.
(two drops) of the various indicators, according to the following
scheme :
Standard Solutions PH =3 45 67 8 9 10 11
(1) MR MR MR MR Tpn Tpn Tpn Tpn
(2) Coc Coc Coc Coc Ppn Ppn Ppn Ppn
(3) MO MO MO RA RA RA RA
Determine where the critical color change occurs. How do your
results agree with what Noyes found, or with the table given in
Washburn?
Note that the experiment, performed as outlined above, is only
roughly quantitative. For accurate work the color changes should
be observed in a colorimeter. Washburn, 332.
EXPERIMENT 2
To Determine Hydrogen Ion Concentration of an Unknown
Determine the approximate hydrogen ion concentration of N/10
methylamine hydrochloride. Employ the set of indicators listed
above.
From your results compute the per cent hydrolysis of N/10
methylamine hydrochloride.
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Sodium Alizarin-
Sulphonate
Rosolic Acid
Guiac Tincture
Phenolphthalein
Thymolphthalein
Methyl Red
Litmus
Cochineal
EXPERIMENTAL GROUP XV
EQUILIBRIUM AND THE PHASE RULE
The series of experiments outlined in this group constitutes a study
of physical and chemical equilibrium from the point of view of the
Phase Rule. It includes phase equilibria in systems of one, two and
three components. Compare the experiments on distillation (Group
VIII) and vapor pressure (Group IV). In carrying out the experi-
mental work keep the Phase Rule in mind.
References :
Bancroft : The Phase Rule (1897) .
Desch: Metallography (1913) ; Intermetallic Compounds (1914).
Duhem (Burgess) : Thermodynamics and Chemistry (1903).
Findlay: The Phase Rule (1917). (FPR).
Roozeboom : Die heterogenen Gleichgewichte (1901-1913) .
Tammann: Kristallizieren.und Schmslzen (1903).
SUB-GROUP 1
INVERSION POINTS
Discussion.
Read carefully the appendix in FPR, 335 (1911) or F, 307-315.
Do not begin experimental work until you are thoroughly familiar
with the principles involved. Note "suspended transformation,"
FPR, 69 (1911).
"It frequently happens that in place of determining the complete
concentration-temperature curve and from the break determining
both the concentration and temperature at the inversion point, one
prefers to' measure the temperature at which such changes occur.
Since a change in the solid phase brings a change in practically all the
physical properties, the close observation of the variations of any
one of these with the temperature will decide at which temperature
the inversion takes place. The different properties whose variations
are accessible to easy measurement are crystal form, volume, color,
vapor pressure, conductivity, and electromotive force. The
variation of the physical properties is accompanied by a variation of
the energy content so that by measurement of the variation of some
energy quantity with the temperature, the inversion point may
readily be found by all the methods; as in analysis, every particular
case shows one method which ought to be employed in preference to
the others, because of its sharpness in detecting the change.
"In practically all cases where phase changes (inversions) occur,
there is a lag or reluctance to change, which may be more marked in
one direction than in the other. This reluctance to change gives rise
to metastable phases and to metastable equilibria. Even when the
60
change of phase (inversion) is actually occurring, time is required for
the change and this may, and usually will, introduce a complicating
factor in the experimental determination of inversion temperature."
EXPERIMENT 1
Optical Method
Part 1. Determine by means of color change the inversion tem-
perature of mercuric iodide. One component.
Carry out this determination with the aid of a Thiele bulb, as you
would make a determination of the melting point. Use H 2 SO 4 and
heat very slowly. Note the point at which the color change occurs
with both rising and falling temperature. What is the cause of the
difference? What is this phenomenon called?
In a test tube heat the red HgI 2 until it becomes yellow. Pour
melted vaseline over some of the yellow HgI 2 and cool quickly. Like-
wise cool the remainder of the yellow iodide exposed to the air. Is .
the stability of the yellow form affected by the presence of vaseline?
Reference. FPR, 75 (1917).
Part 2. Following the same procedure, determine the inversion
point of copper potassium chloride. Pick out blue crystals of the
hydrated double salt in preference to the green ones. Three com-
ponents: CuCl a 2KC1 2H 2 O = CuCl 2 KC1 + KC1 + 2H 2 O.
How many phases are in equilibrium at the inversion point?
EXPERIMENT 2
Thermometric Method (Cooling Curves)
Discussion.
If a system of phases is at a temperature different from the sur-
roundings it will either absorb or give off heat according to its tem-
perature. If at any temperature there occurs in the system some
change where heat is evolved or absorbed there must necessarily be
a break in the curve of heating or cooling. Since the appearance or
disappearance of a phase is always accompanied by a heat change, one
may easily and rapidly make the determination by observing the
temperature-time curve indicating the rapidity of heating or cooling
of the system.
Procedure.
Part 1. Determine the inversion temperature of sodium sulphate
decahydrate (Glauber's salt) by the thermometric method.
In a test tube place sufficient powdered salt to cover completely the
bulb of a large thermometer graduated in tenths. The test tube
should be half full. Place the test tube in a water bath and beginning
at 28 heat slowly to 36, stirring the contents" of the test tube con-
stantly with the thermometer. Raise the temperature of the bath
at a uniform rate, not exceeding one degree in five minutes.
Read the temperature on the thermometer immersed in the salt
at regular intervals of two minutes. At the same time record any
changes which may be visible in the contents of the tube. Draw a
curve between temperature and time and note the "break" at the
inversion temperature.
61
Next cool the test tube and contents from 36 to 28 proceeding as
you did before. Draw a cooling curve between temperature and
time.
If undercooling becomes excessive and persists, add a crystal
of decahydrate and stir vigorously. Account for the sudden rise of
temperature.
How many components and phases are there at the inversion
point? How does the inversion point differ in this case from a melt-
ing point? Has Glauber's salt a melting point?
Part 2. Determine the inversion temperature of mercuric chloride
methylalcoholate, HgCl 2 CH 3 OH. Saturate methylalcohol at 45 C.
with HgCl 2 . Cool and determine the temperature at which HgCl 2
ceases to be deposited and the alcoholate makes its appearance.
The reaction may be written
HgCl 2 + CH 3 OH = HgCl 2 CH 3 OH.
Reference. Jour. Phys. Chem., 1, 298 (1896).
Caution. Work at the hoods.
EXPERIMENT 3
Dilatometric Method (Volume Changes)
The powdered solid is introduced into the bulb of a glass dilato-
meter through the larger tube below the bulb. The capillary tube
is closed by means of a small piece of glass to prevent the solid sub-
stance from clogging the capillary. This piece of glass may best be
made by drawing out a glass rod, then forming a bead at one end
by holding it in the flame for an instant. The bulb is then nearly
filled with the solid and the larger tube sealed off.
The dilatometer must now be filled with some measuring liquid,
e. g., petroleum or xylene. This is best done by attaching an adapter
to the end of the capillary tube by means of a rubber stopper fitting
the wide end of the adapter and then connecting the latter to a suc-
tion pump after filling with xylene. The air from the dilatometer
bubbles through the oil, which, when the pressure is released, is
drawn back into the dilatometer, Cf. F, 312 (1917).
This operation is repeated until all the air is withdrawn from the
dilatometer and replaced by xylene. This capillary tube of the dila-
tometer should be tapped frequently to loosen any adhering air
bubbles. Any excess of xylene may be removed from the capillary
by means of a long finely drawn out capillary tube, so that when the
dilatometer is placed in the water bath the xylene meniscus may
remain on the scale. The capillary tube is not sealed. A suitable
millimeter scale is used for reading the change in volume. This
method is especially useful for determining inversion points when the
amount of substance obtainable is relatively small.
By means of the method described find the inversion temperature
of sodium thiosulphate pentahydrate, Na 2 S 2 O 3 5 H 2 O.
After the dilatometer has been filled, place it in a large beaker of
water and starting at 46, heat to 52 at the rate of 1 every five
minutes, noting the change in volume. Then allow the dilatometer
to cool very slowly, taking readings of temperature and volume.
62
Finally, start at a temperature about two degrees below the inver-
sion temperature and heat to a temperature of about two degrees
above, at a rate of 1 every ten minutes. Again allow dilatometer to
cool, taking readings of temperature and volume.
Does suspended transformation cause trouble?
How does the inversion point of sodium thiosulphate differ from
the inversion point with Glauber's salt?
SUB-GROUP 2
EUTECTIC POINTS
EXPERIMENT 1
Cryohydric Points.
In this case the problem is to determine the conditions under which
solid solvent (ice), solid solute (K 2 SO 4 ), solution and vapor may
co-exist. Under the conditions of the experiment, using vessels open
to the air, the system may not really be in equilibrium with the vapor
and may be under a pressure different from that of the invariant
system, ice, salt, solution and vapor. Actually, however, the slight
and slowly acting readjustments due to these causes do not have
much influence upon the temperature at which ice, salt and solution
are in equilibrium; and the eutectic temperature of a system com-
posed of non- volatile or slightly volatile salt, ice, solution and vapor,
determined at atmospheric pressure in open vessels, does not differ
appreciably from the temperature of the system, salt, ice, solution
and vapor in complete equilibrium.
Part 1. Prepare a saturated solution of K 2 SO 4 in water and place
this solution in a test tube immersed in an ice-salt freezing mixture.
Note the temperature at one minute intervals, immersing the
thermometer in the solution. Draw the usual curve between time
and temperature.
Part 2. Prepare a dilute solution of K 2 SO 4 and repeat the pro-
cedure of Part 1 using 5 g. K 2 SO 4 in 93 cc. water.
The concentration of the solution at the cryohydric temperature
may be ascertained by removing a sample with a pipette, being care-
ful to prevent the introduction of any solid material into the pipette.
This sample may be analyzed and its sulphate content determined by
precipitating with barium chloride.
EXPERIMENT 2
Eutectic Points by Cooling Curves
By the thermometric method determine the eutectic point of one
of the following pairs: naphthalene-anthracene, naphthalene-
phenol, naphthalene-diphenylamine.
Compare with data and phase diagrams in LBR.
63
SUB-GROUP 3
TWO LIQUID LAYERS
EXPERIMENT 1
Melting under the Solvent. Add an excess of para-toluidine to
water in a test tube. Heat on a steam bath to 45 C. and note what
happens. At what temperature does the para-toluidine melt?
What is the melting point of pure paratoluidine? FPR, 129 (1917).
EXPERIMENT 2
Phenol and Water. Make up mixtures of phenol and water of the
following composition in parts of phenol in 100 parts of mixture:
5, 8, 10, 20, 30, 40, 50, 60, 70, 80, 90.
Weigh the required amounts of phenol out as quickly as possible
to prevent absorption of moisture from the air. Let the combined
weight of phenol and water in each mixture be 15 or 20 g. Add the
required amount of water from a burette and immediately close the
mouth of the test tube with a cork.
Beginning with the mixture containing 10 per cent of phenol, heat
each succeeding mixture (up to the 90 per cent one) by immersing the
test tube in a water bath (i. e. a beaker). Place a thermometer in
the test tube and stir thoroughly. Stirring by means of a slow stream
of air is very effective. When the two layers disappear, and the
liquid becomes homogeneous, observe the temperature.
Next remove the test tube from the bath, and with constant stirring
and slow cooling, observe the temperature at which the two layers
reappear, i. e. when the solution becomes milky.
Next place the test-tube in a freezing mixture and determine the
temperature at which the phenol solidifies under the solution. Is this
the same temperature as the eutectic point? Explain. Determine
the eutectic point.
Draw a curve with concentrations as abscissae and temperatures as
ordinates.
The 8 (and perhaps the 70) per cent solutions should be homogen-
eous at ordinary temperatures. On immersion in cold water, how-
ever, the liquid layers will be formed just as in the other cases.
Determine at what temperature this occurs.
The 5, 80 and 90 per cent mixtures should also be homogeneous at
room temperature. On cooling in a freezing mixture, these solutions
do not separate into two liquid layers but deposit a solid phase.
Determine the temperature at which solid first begins to appear and
ascertain the nature of the solid phase. Ice or phenol?
Note. Do not throw away the phenol-water mixtures but return
them to the bottle marked "phenol residues."
Note the following:
Take a 30 per cent mixture of phenol in water and heat to about
75 C. At this temperature, add 5 to 10 grams of solid phenol.
Do two liquid layers form? Allow the solution to cool down until
the layers appear, noting the temperature. Represent what you did
graphically on the curve obtained in Experiment 2.
64
Precaution. Phenol is very corrosive. Do not let it remain in
contact with the skin.
References. LBR, 592; Lehfeldt, 228; Rothmund: Die
Loslichkeit (1907).
EXPERIMENT 3
Sulphur and Aniline. (Optional)
Proceeding exactly as you did in the case of phenol and water make
up the following mixtures of sulphur and aniline:
25, 40, 50, 60, 70, 80, 85, 90, 93 per cent sulphur.
Determine the temperatures at which the layers appear (i. e. the
clear liquid becomes turbid) on cooling the clear solutions from a
temperature of 140-160 C. The turbidity will be noticed between
the temperature limits of 102 and 140 C. Stir vigorously. Also
ascertain at what temperature the pure sulphur melts and at what
temperature it melts under the solvent. To do this note the tem-
perature at which the lower layer of aniline in sulphur in one of the
above mixtures solidifies to a crystalline yellow mass.
Draw a curve between temperature and composition.
Precaution. Use roll sulphur (not flowers of sulphur).
Reference. LBR, 595.
EXPERIMENT 4
Three Components Chloroform, Acetic Acid, and Water
Make up mixtures of chloroform and water of the following compo-
sition (by weight) : 98, 95, 90, 80, 70, 60, 50, 40, 30, 20, 10, 5, 2 parts
of chloroform in 100 of mixture. Total weight of each mixture to
be 40 grams.
Mix in 100 cc. glass stoppered bottles, shake vigorously, heat to
about 40 in a water bath, cool and allow to come to equilibrium by
standing a week.
When this has been done and the bottles are at the same tempera-
ture (record) add glacial acetic acid from a burette until a homo-
geneous (non-cloudy) solution is obtained. Shake constantly during
the addition of the acid. Calculate the weight of acetic acid neces-
sary to produce a homogeneous solution and plot your results upon a
triangular diagram.
Reference. FPR, 249 (1917).
SUB-GROUP 4
PREPARATION OF COMPOUNDS
The object of this set of experiments is to give practice in applying
phase rule methods to the preparation of compounds by the system-
atic use of temperature-composition diagrams.
65
EXPERIMENT 1
Hexahydrate of Calcium Chloride
Diagram in FPR, 155 (1917). Plan your procedure carefully and
report it to the Instructor before doing this experiment. Follow this
plan throughout.
Note. Filter the CaCl 2 solution, as it may be turbid on account of
basic chlorides if made from desiccated CaCl 2 . Show the compound
to the Instructor.
EXPERIMENT 2
Hydrates of Potassium Hydroxide
Part 1. Prepare the monohydrate of KOH.
Part 2. Prepare the dihydrate of KOH.
Show the crystals to the Instructor.
References. Pickering: Jour. Chem. Soc., 63, 899 (1893).
Note properties of crystals, 898. Complete data and diagram in
LBR, 477.
EXPERIMENT 3
Monohydrate of Sulphuric Acid
Prepare the monohydrate of H 2 SO 4 .
References. Pickering: Jour. Chem. Soc., 57, 338 (1890).
Complete data (SO 3 and water) and diagram in LBR, 493.
Hint. Since the solubility curve for the compound H 2 SO 4 ' H 2 O
passes through a very sharp maximum in respect to temperature,
unless the concentration of the solution is very accurately adjusted
to be equal to that of the maximum point, one is very apt to meet
with failure unless the solution is cooled to a very low temperature.
Prepare the solution and divide it into two equal parts. _ Try to
crystallize out the monohydrate. If you fail, the solution is either
too concentrated or too dilute (unless supersaturation has caused the
trouble). To one of the tubes add a drop of water, to the other a
drop of concentrated acid and again attempt to crystallize the mono-
hydrate. Continue this procedure until you succeed. Show the
crystals to the Instructor and record the temperature at . which the
last crystals disappear on warming.
EXPERIMENT 4
Carnallite KC1 -MgCl 2 '6H 2 O
Discussion.
Cf . FPR, 280-298 (1917) ; see isothermal diagram for 25 C. in
Whetham: Solutions, 404 (1902); excellent discussion by Hilde-
brand: Jour. Ind. Eng. Chem., 10, 97 (1918).
It is obvious that if one prepares a solution containing equimole-
cular quantities of KC1 and MgCl 2 '6H 2 O and evaporates until the
66
solution phase just disappears, carnallite will be formed, since this
salt is stable above 21. This method however is not elegant and
if the evaporation is discontinued at any point short of complete
disappearance of the liquid phase a mixture of carnallite and KC1 will
be obtained. It is important to remember that carnallite cannot be
in equilibrium with a solution containing MgCl 2 and KC1 in the ratio
of 1:1. When carnallite is dissolved in water the solution soon
becomes saturated with KC1 and this salt is precipitated while
carnallite continues to dissolve. It is not until the MgCl 2 content
of the solution rises to a high value by the precipitation of KC1, that
carnallite can exist as stable phase in contact with solution.
Procedure.
Prepare a solution of MgCl 2 and KC1 in the proper molecular ratio
to insure the separation of carnallite as the first solid phase
on cooling or dehydrating. Show the crystals to the Instructor.
Prove that they really are carnallite.
For data regarding the composition of the solution cf. FPR, 298,
296 (1917).
Suggest a simple method of obtaining KC1 from Stassfurt carnallite.
EXPERIMENT 5
Copper-potassium Chloride
Following the procedure used in preparing carnallite, make the
blue double salt. Test for purity by determining the inversion point
for the breakdown, 2 KC1 CuCl 2 2H 2 O -+ KC1 ; CuCl 2 + KC1 + 2H 2 O
If the salt is green the result is not entirely satisfactory.
Note. 2 KCl-CuCl 2 -2H 2 O, like carnallite, is unstable in contact
with solution containing KC1 and CuCl 2 in the ratio 2:1, but is stable
in contact with a solution containing these salts in the ratio 1 :1 or 1 :2.
Bancroft: Phase Rule, 176 (1897).
EXPERIMENT 6 (OPTIONAL)
Lead Potassium Iodide
Prepare lead potassium iodide, PbI 2 'KI'2H 2 O. Bancroft, 179
(1897); Abegg: Handbuch, III (2) 667; Schreinemakers: Zeit. phys.
Chem. 10, 467 (1892).
Note. Schreinemakers' diagram on page 471 indicates that the
double salt is stable only in a solution containing KC1 in excess.
EXPERIMENT 7 (OPTIONAL)
Astracanite, Na 2 SO 4 - MgSO 4 - 4H 2 O
Reference. FPR, 264 (1911).
Report results. Write the reaction.
67
SUB-GROUP 5
INDIRECT ANALYSIS
Discussion.
Under some circumstances solid separates out from a liquid phase
in a form which renders direct analysis very difficult and uncertain.
The solid may be unstable and it may be impossible to remove adher-
ing mother-liquor. Indirect analysis is then resorted to. Many
methods of indirect analysis have been proposed; the following
experiment illustrates one of the most satisfactory.
References. Bancroft: Jour. Phys. Chem., 6, 178 (1902).
Browne: Ibid, 6, 281 (1902).
FPR, 236, 310 (1917).
EXPERIMENT 1
Determination of Solid Phases
Discussion.
Let us suppose a system to be composed of three components A, B,
and C, all of them miscible in the liquid phase. Starting with a
system composed of the homogeneous (unsaturated) solution in con-
tact with vapor, let the composition of the solution be a per cent
of A, b per cent of B, and c percent of C.
Next, without changing the total amount of A, B, and C in the
system (no loss by evaporation, etc.) cool until a single solid phase
separates out and the system solid-liquid is produced. Suppose that
a qualitative analysis of the solid phase indicates that C is not present
in the solid. There are three possibilities, as follows:
(1) Solid is pure A or pure B.
(2) Solid is a compound of A and B.
(3) Solid is a solid solution of A and B or an absorption compound.
Without, removing the solid, pipette out some of the clear mother-
liquor and analyze it. Let the composition now be (in per cent) a',
b', and c'. The following relations hold true for the two solutions:
a + b +c = 100 (1)
a' -f b' + c' = 100 (2)
Next divide (2) by ^-, whence
vSince C has not separated out in the solid phase and the total
amount of C in the liquid phase therefore remains unchanged, the
composition of the solid phase must be proportional to (FPR, 232) :
68
If M A and M_, are the respective molecular weights, then the
A. 13
molecular composition of the solid phase is given by the expression
ao'-a'A /bc-jyc
C 'M A ; ^ , M B
From (5) the number of molecules of B combined with one mole-
cule of A becomes
_^A (bc> - b'A
M ^ac' - a'c/
B
Procedure.
Prepare a solution of 50 g ; sodium sulphate decahydrate (Glauber's
salt) and 10 g. sodium chloride in 100 cc. of distilled water. Filter
the hot solution. Cool to 45 C. and analyze the solution for sodium
chloride and sodium sulphate. See below for procedure. Keep the
solution in a stoppered flask or Erlenmeyer. Run in duplicate.
Cool the solution until solid crystallizes out in considerable amount,
then carefully pipette two samples of the solution for analysis. It
may be found advisable to fit to the end of the pipette a bit of glass
tubing containing glass wool or cotton to serve as a filter. Separate
some of the solid and wash with a very little water. Has any sodium
chloride been precipitated?
Analysis. Determine NaCl in one sample (1 g.) with standard
silver nitrate (shelf) using K 2 CrO 4 as indicator. Evaporate a
second sample to dryness (being careful to avoid spattering) and
determine total chloride and sulphate. Determine water by
difference. Using equation (6) determine the chemical formula of
the solid phase, assuming that no solid solutions are formed in this
experiment.
How else might one determine approximately the composition of
the solid in the above experiment, using, of course, an indirect
method?
Outline the procedure in case component C also separates out in
the solid phase. See references (Triangular Diagrams).
How could one distinguish between compound and solid solution?
Why must the two salts have an ion in common?
How can one tell whether the number of solid phases precipitated
from the solution is one or two?
EXPERIMENTAL GROUP XVI
COLLOID CHEMISTRY
This comprehensive group of experiments serves to illustrate some
of the more important and interesting properties of colloidal systems.
Typical colloids are prepared and studied, particularly from the
point of view of Bancroft: Jour. Phys. Chem., 18, 549 (1914).
Read the article before beginning experimental work in this group.
General Texts in Colloid Chemistry.
Alexander: Colloid Chemistry (1919).
Bancroft: Applied Colloid Chemistry (1920).
Burton: Physical Properties of Colloidal Solutions (1916).
Cassuto: Der Kolloide Zustand der Materie (1911).
Freundlich: Kapillarchemie (1909).
Hatschek: An Introd. to the Physics and Chemistry of Colloids
(1919).
Miiller: Chemie der Kolloide (1907).
Ostwald (w) : Grundriss der Kolloidchemie (1911-12).
Ostwald (w) (Fischer): Theoretical and Applied Colloid
Chemistry (1915).
Ostwald (w) (Fischer) : Handbook of Colloid Chemistry (1915)
Svedberg: Die Methoden zur Herstellung kolloider Losungen
usw. (1909).
Taylor: The Chemistry of Colloids (1915).
Willows and Hatschek: Surface Tension (1915).
Zsigmondy (Alexander) : Colloids and the Ultramicroscope (1909) .
Zsigmondy: Kolloidchemie (1912).
Zsigmondy (Spear): Colloid chemistry (1917).
Journals
Journal of Physical Chemistry, (special articles).
Kolloidchemische Beihefte (special articles).
Kolloid-Zeitschrift. (1906).
SUB-GROUP 1
DIFFUSION, DIALYSIS AND MEMBRANES
EXPERIMENT 1
Diffusion of Solutions.
Obtain six test tubes, fitting each with a rubber stopper (one hole),
and prepare six 15 cm. lengths of narrow-bore (2. 5-3 mm. internal
diam.) glass tubing. Seal one end of each length of tubing and fill
70
completely with distilled water. Place 10 cc. of solution to be tested
in each test tube, insert a water-filled diffusion tube in the stopper
and place it in the test tube, immersing open end of the diffusion
tube just below the surface of the solution. Work carefully. Set aside
the test tubes in a safe place and make observations at regular inter-
vals, recording the time. Test the following solutions:
KMnO 4 solution N/50.
KMnO 4 solution N/5.
Congo red 1/5 of one per cent.
Methyl violet or safranine 1/5 of one per cent.
Arsenious sulphide sol. (See Part 3 below).
Ferric oxide sol. (Sse Part 3 below).
Optional Method. The following experiments are similar to those
of Graham. A small, two-dram vial is fastened to the bottom of a
tall, narrow beaker (250 cc. capacity) by means of paraffin.
Fill the vial carefully with the solution containing the solute whose
rate of diffusion is to be measured and cover it securely with a small
cover-glass (20 millimeters). Be sure that no solution is spilled
from the vial during the process of filling and covering. Pour dis-
tilled water into the beaker until it is nearly full and the vial is well
covered, taking care to have the water level at the same height in
each beaker. Finally, slide the cover glass carefully off the mouth of
the vial by means of a clean glass rod.
A two cc. test-sample is then pipetted from the liquid in the beaker
at a point about three centimeters above the open mouth of the vial.
Mark this position by means of a label placed on the wall of the
beaker. Be careful not to stir the liquid. Test for chlorine as ion
with silver nitrate making a rough nephelometric estimation of the
relative amounts of silver chloride formed in each sample. Test for
organic matter by evaporating a test sample to dryness in a clean
porcelain dish and carbonizing the residue.
It is essential that the water levels be the same in each beaker, that
the sample be pipetted from equal distances above the mouth of the
vial and that the beakers and solution remain absolutely undisturbed.
Withdraw test samples at the beginning and after 1, 2, 4 and 7 days,
noting the exact time.
The following solutions are to be tested:
(1) One per cent solution of gelatine.
(2) Five per cent solution of sodium chloride.
(3) Twenty-five per cent solution of sodium chloride.
Note. Prepare a 5 per cent solution of gelatine for this and subse-
quent work as follows: Soak 2 g. of gelatine in cold water until soft,
pour off. the water and to the softened gelatine add enough warm
water to make about 40 cc. of solution. On cooling, a jelly will form
which readily melts when the beaker with the jelly is warmed on the
steam bath. Do not warm over a flame as the beaker will almost
certainly crack. Dilute the gelatine solution as required.
71
EXPERIMENT 2
Diffusion Through a Jelly
Obtain eight small test tubes and fill each half full of liquid 5 per
cent gelatine and allow this to solidify. Pour into the tubes, on top
of the gelatine, the solutions or sols specified below, being careful
that the latter are cold so that they do not liquefy the jelly.
If they diffuse, the substances in solution will tend to pass from the
upper aqueous layer into the lower portion occupied by the gelatine
and the process may be observed by means of the coloration produced
in the jelly. If the colored substance forms a true solution, the
diffusion of the solute through a jelly occurs almost as rapidly as
through pure water itself. On the other hand, colloidal solutions
show practically no evidence of diffusion. We may, therefore, dis-
tinguish between the two classes of solution by means of this method,
provided the jelly is not "semi-permeable" to the dissolved solute.
Observe the condition of each tube after twenty-four hours and
again after a week. Keep the tubes in a cool place. Use the follow-
ing solutions (shelf) :
(1) Eosine
(2) Congo red
(3) Safranine
(4) Picric acid
(5) Methylene blue
(6) Arsenious sulphide sol
1/5 of one per cent.
1/5 of one per cent.
1 /5 of one per cent.
1/5 of one per cent.
1/5 of one per cent,
(see below).
-(7) Ferric 'oxide sol (see- below).
(8) Mixture Congo red and picric acid, picric acid in excess.
From the data obtained in these experiments what conclusion do
you draw regarding the nature of the above solutions?
EXPERIMENT 3
Dialysis v/ith Collodion. Instead of using parchment, prepare
collodion dialyzing tubes as follows: Take one of the inner test
tubes of heavy glass used in the free2ing point determinations and
wet the inner walls completely with a fairly thick film of collodion
solution (soluble cotton in a mixture of ether and alcohol) . Do this
quickly while spinning the tube to make the collodion film uniform.
As soon as the collodion "sets" blow air into the tube to remove the
ether. This process should take about five minutes. Then pour
water into the test tube and gradually loosen the collodion from the
glass. With moderately careful manipulation, a transparent, tough
dialyzing tube can be obtained which is more convenient and less
expensive than the parchment dialyzers ordinarily used. Having
prepared the tube, test for leaks by filling with water and if intact,
immerse completely in a large beaker of water to remove the alcohol.
Soak until the next period, changing the water from tirn.e to time.
Make three dialyzing tubes.
Fill one nearly full with a mixsd solution containing 1 per cent
gelatine plus 5 per cent of sodium chloride. Place this in a beaker of
distilled water and test the water at stated interval for NaCl and
.gelatine.
72
Fill the second tube with a solution of safranine. Place this in a
second beaker of water and observe diffusion. In the third tube place
a solution of Congo red. Does this diffuse?
EXPERIMENT 4
Semipermeable Membranes
Into a small bottle pour, very carefully and in the order given, the
following liquids: Chloroform, water, and ether. Three layers
should be present. Note the thickness in mm. of each layer.
Let the bottle stand undisturbed for a week and again measure
the thickness of the layers. Continue the experiment until one of the
three original layers disappears. Explain.
Reference. Kahlenberg: Jour. Phys. Chem., 10, 146 (1906).
EXPERIMENT 5
Osmosis and Semipermeable Membranes
Part 1. Fill a test tube with a M/2 CuSO 4 , then, by means of a
pipette placed in this solution add slowly and carefully a small
amount of M/2 potassium ferrocyanide. A globule should form,
consisting of the solution of ferrocyanide surrounded by a gelatinous
membrane of brown copper ferrocyanide. Carefully detach the
globule from the end of the pipette and it will sink, owing to the
greater density of the ferrocyanide solution.
Observe carefully any changes that may occur in the copper
sulphate solution surrounding the globule. Set aside the test tube
and keep it constantly under observation. What happens?
Explain.
Part 2. Plant-like Growths. Fill a small beaker with dilute
sodium silicate (water glass) solution and drop into the liquid one or
two crystals each of CuSO 4 , MnSO 4 , CoSO 4 , etc. What happens?
Explain.
SUB-GROUP 2
ADSORPTION
The following experiments are designed to illustrate adsorption
phenomena. Adsorption is the basis of colloid chemistry. All the
experiments of Sub-groups 3 and 4 illustrate this point.
EXPERIMENT 1
Adsorption by Bone Black
Part 1. Boil a dilute solution of litmus with bone black. Filter.
Part 2. Repeat, using dilute solution of indigo. Are the colors
removed? Explain.
Part 3. Prepare a dilute solution of silver nitrate. Divide this
into two portions. To one portion add about one-tenth its volume of
bone black and shake vigorously for at least three minutes. Then
73
filter and add NaCl to both portions. Compare the amounts of
precipitated silver chloride.
Bone black or animal charcoal contains 85 per cent. of calcium
phosphate and about 15 per cent of carbon.
EXPERIMENT 2
Selective Adsorption
Part 1. Prepare about 250 cc. of indicator solution as follows:
To 250 cc. of distilled water add a little phenolphthalein and a trace
of NaOH, just enough to color the liquid pink.
Part 2. Ina test tube shake fuller's earth with distilled water and
add some of this muddy suspension to one of the test tubes containing
the indicator. Is the color removed?
Part 3. Allow this muddy suspension to settle and then add the
supernatant clear liquid to a second test tube colored with indicator.
Filter the supernatant liquid to remove all the fuller's earth. Is this
filtered liquid acid?
Part 4. Moisten a little fuller's earth with boiled water and test
with blue litmus by pressing the latter down on the earth.
Reference. Cameron: Jour. Phys. Chem., 14, 400 (1910).
Part 5. Add blue litmus solution to fuller's earth suspended in
water. Notice the change.
Part 6. Add some fuller's earth to a dilute solution of methyl
violet and shake. Filter, noting color of filtrate and of earth. Is
the color removed from the earth by water or alcohol?
Part 7. Repeat the last experiment, using eosin instead of methyl
violet. Note any differences in behavior.
Part 8. Moisten some absorbent cotton with freshly boiled water
(free from CO 2 ) and wrap it around a strip of blue litmus paper.
For comparison of the original and the final color, let about half an
inch of the paper protrude beyond the cotton. Explain your results.
Compare Part 4, above.
EXPERIMENT 3
Adsorption by Iron Oxide. The Antidote for Arsenic Poisoning
Hydrous ferric oxide is precipitated from a solution of ferric
sulphate or chloride by adding an excess of magnesia. Shake
vigorously. Then prepare a dilute solution of As 2 O 3 and filter, and
test the filtrate for arsenic with H 2 S.
Be sure that the As 2 O 3 solution is very dilute. Test half the
original solution with H 2 S for arsenic. Only a slight test should be
obtained, if the experiment is to work well. Then test the second
half of the As 2 O 3 solution after treatment with the ferric hydroxide
mixture. Has the arsenic been adsorbed? Should the arsenic be
completely adsorbed? Explain.
74
EXPERIMENT 4
Adsorption Compounds. Carey Lea's "Photohalides"
Reference. Carey Lea: Am. Jour. Science, (3) 34, 349, 480,
489 (1887).
Method suggested by Luppo-Cramer : Kolloid-Zeit., 2, 360
(1908).
To 3.5 cc. of 10 per cent KBr add 5.5 cc. of 10 per cent AgNO 3 .
To this mixture containing AgBr plus AgNO 3 in excess add the
following solution:
7.5 cc. Rochelle salts (1:3) plus 2.5 cc. of ferrous sulphate (1:3).
Do not add the Rochelle salts and ferrous sulphate solutions sepa-
rately.
Wash the dark colored precipitate several times by decantation and
finally with a mixture of equal parts concentrated HNO 3 (1.4 sp.gr.)
and water. An intense blue-violet color should develop.
The photohalides of silver are adsorption compounds of silver with
silver chloride and are similar to the "subsalts" of silver composing
the "latent image" in an exposed photographic plate.
EXPERIMENT 5
Selective Adsorption and Capillary Diffusion
Part 1. Place several drops of a mixed solution of CuSO 4 and
CdSO 4 (shelf) on the center of a square of blotting paper (6 by 6 in.).
Allow the drops to diffuse until a large round spot has formed, then
hold the paper in a stream of H 2 S gas. Which "diffuses" farthest,
water, CuSO 4 , or CdSO 4 ? Cf. Gordon: Jour. Phys. Chem., 18,
337 (1914).
Part 2. Suspend strips of blotting paper (1 cm. broad and 20 cm.
long) in water solutions of the following substances: Congo red;
picric acid; cosin; methylene blue; methylene blue plus cosin.
Note the height to which the water and dye rise.
Reference. Goppelsroeder: Kapillaranalyse (1906).
EXPERIMENT 6
Adsorbed Air in Charcoal
Fit a cylinder (100 cc.) with a three-hole rubber stopper. Into one
hole introduce the delivery tube of a burette filled with water. In
the second place a thermometer. In the third place a glass tube lead-
ing to a water- filled graduated cylinder (capacity 250 cc.) inverted
over water in a trough.
Place a volume of 50 apparent cc. of granular cocoanut charcoal in
the cylinder. Then add water slowly from the burette, recording the
volume added. Continue to add water until its level rises to the
surface of the charcoal. Measure the volume of air displaced.
Take the temperature before and after adding the water. Have
the water in the burette and the charcoal at the same temperature in
the beginning.
75
SUB-GROUP 3
PEPTIZATION
EXPERIMENT 1
Peptization by Adsorbed Ions
Lottermoser: Jour. Praktische Chemie, [2] 72, 39 (1905); 73,
374(1906); Zsigmpndy (Spear) 179; Ostwald (Fischer) : Theoretical
and Applied Colloidchemistry, 115.
Part 1. Prepare a small quantity of silver bromide and wash the
precipitate thoroughly by decantation. Place approximately equal
amounts of the freshly prepared silver bromide in each of five
stoppered test tubes. In the first test tube place distilled water (10
cc.); in the second, N/100 KBr; in the third, N/30 KBr; in the
fourth, N/10 KBr, and in the fifth, N/5 KBr. Shake thoroughly and
after allowing the test tubes to remain standing several minutes,
describe the appearance of each tube. In which is the supernatant
liquid most turbid?
The process constitutes a dispersion method of preparing colloidal
silver bromide.
Part 2. Fill two burettes with N/20 AgNO 3 and N/20 NH 4 CNS
(shelf). Fit a small Erlenmeyer flask with a solid rubber st.opper.
Perform the following experiments:
(a) To 10 cc. AgNO 3 in flask add quickly 10 cc. NH 4 CNS, stopper
and shake.
(b) To 10 cc. AgNO 3 in flask add quickly 12 cc. NH 4 CNS, stopper
and shake.
(c) To 10 cc. NH 4 CNS in flask add quickly 10 cc. AgNO 3 , stopper
and shake.
(d) To 10 cc. NH 4 CNS in flask add quickly 12 cc. AgNO 3 , stopper
and shake.
What striking differences do you observe and how do you account
for them?
(e) Refill the burettes and, placing 10 cc. AgNO 3 in a flask run in
NH 4 CNS from a burette (not too rapidly) until floccula-
tion occurs. Shake and note the volume of NH 4 CNS
added. Repeat adding NH 4 CNS more slowly as end-
point is reached. The end-point represents the isoelectric
point (define).
(f) Place 10 cc. NH 4 CNS in a flask and add AgNO 3 following
the procedure of (b) 5 preceding. Explain.
EXPERIMENT 2
Peptization by Adsorbed Colloid
Prepare 5 per cent solutions of chromic and ferric chlorides. Mix
in the proportions specified below. Then add 10 per cent NaOH in
excess. Note the color and appearance of the precipitate (if any)
and of the supernatant liquid. Use test-tubes and shake.
76
Ferric Chloride Chromic Chloride Remarks
(cc). (cc.)
10
8 2
5 5
3 7
2 8
1 9
10
Cf. Nagel: Jour. Phys. Chem., 19, 331, 569 (1915).
EXPERIMENT 3
Peptization by Adsorbed Colloid (Protective Colloids)
Solution A: 5 cc. N/2 AgNO 3 + 5 cc. of 5 per cent gelatine.
Solution B: 5 cc. N/2 KBr + 5 cc. of 5 per cent gelatine.
Part 1. After thoroughly mixing each solution, pour B into A,
shake and note any changes. Place the mixture in the sunlight and
note results. Repeat the above experiment, replacing the gelatine
solution by an equal volume of pure water. Was AgBr formed in the
first experiment with gelatine. How might one prove this?
Part 2. Prepare some silver bromide, wash by decantation and
remove to a filter paper. Divide into two portions. Place one por-
tion in an air bath and dry for an hour at 120, being careful not to
exceed this temperature.
To the freshly prepared moist silver bromide add a few cubic centi-
meters of hot 5 per cent gelatine and shake vigorously. Is a suspen-
sion formed? Do the same thing with the dried silver bromide and
note any differences in its behavior compared with that of the freshly
prepared substance. What is the effect of "ageing?"
Part 3. Grind a little roll-sulphur with a 5 per cent gelatine solu-
tion in a mortar until a milky suspension is formed. Pour some of
this suspension into water and note the color.
SUB-GROUP 4
PREPARATION AND FLOCCULATION OF SUSPENSIONS
EXPERIMENT 1
Colloidal Arsenious Sulphide (Condensation Method)
In a clean beaker, boil about 6 grams of As 2 O 3 with 100 cc. of
distilled water for fifteen minutes. Cool, filter and dilute to 100 cc.
Pass clean hydrogen sulphide gas into the solution of arsenious acid
until no further action takes place. Remove excess of H 2 S by blow-
ing a slow stream of air through the suspension and then filter.
Describe the appearance of the suspension as to color, turbidity,
etc., and perform the following tests. (See also diffusion experi-
ments) .
77
(a) To 10 cc. colloidal As 2 S 3 add 2 cc. M/20 HC1.
(b) To 10 cc. colloidal As a S 3 add 2 cc. M/20 NaCl.
(c) To 10 cc. colloidal As 2 S 3 add 2 cc. M/20 MgCl 2 .
(d) To 10 cc. colloidal As 2 S 3 add 2 cc. M/20 Al (NO 3 ) 3 .
(e) To 10 cc. colloidal As 2 S 3 add 2 cc. M/20 Na 2 SO 4 .
Which produces flocculation most quickly? Explain.
Colloidal As 2 S 3 thus prepared is a negative suspension. That is,
the particles of the disperse phase carry a negative charge due to
preferential adsorption of anions from H 2 S present in solution.
Place a test tube containing 10 cc. of As 2 S 3 suspension in an ice salt
freezing mixture until frozen solid. Warm the test tube gently until
the ice is melted. What effect upon the suspension is noticed?
EXPERIMENT 2
Colloidal Ferric Oxide (Condensation Method)
Add about 0.5 gram of crystallized ferric chloride to 100 cc. of
boiling distilled water. Then boil the solution gently for about ten
minutes, replacing the water boiled away. Note the color and appear-
ance of the hot solution, and compare with the color of a solution
made by adding FeCl 3 to cold water. Explain the change. What
is this process called?
(a) To 10 cc. of the iron oxide suspension add 2 cc. M/20 NaCl.
(b) To 10 cc. of the iron oxide suspension add 2 cc. M/20 MgCl 2 .
(c) To 10 cc. of the iron oxide suspension add 2 cc. M/20 Na 2 SO 4 .
(d) To 10 cc. of the iron oxide suspension add 2 cc. M/20 citric acid
(e) To 10 cc. of the iron oxide suspension add trace of H 2 SO 4 .
Which causes the most rapid flocculation? Explain. What is
the precipitate?
The ferric oxide suspension as prepared above is positive.
Optional Experiment. Colloidal Ferric Oxide (Dispersion Method)
Reference. Kratz: Jour. Phys. Chem., 16, 126 (1912).
Prepare Fe 2 O 3 suspension by the method of washing out the coagu-
lating salt, following Kratz's procedure.
% EXPERIMENT 3
Mutual Flocculation of Two Suspensions
Study the mutual flocculation of colloidal As 2 S 2 and Fe 2 O 3/ two
oppositely charged suspensions. Plan your own experiments.
EXPERIMENT 4
Colloidal Silica. (Condensation Method)
To 10 cc. of syrupy sodium silicate solution add 30 cc. of water and
pour the resulting solution into a mixture of 25 cc. of concentrated
hydrochloric acid previously diluted with an equal volume of water.
A limpid mixture will result, consisting of a suspension of hydrated
silica.
78
Warm some of this solution nearly to boiling and allow it to stand
undisturbed for a few minutes. What has occurred? Can the sus-
pension be restored? Study the jelly obtained. How does it differ
from gelatine or agar agar?
EXPERIMENT 5
Colloidal Metals (Condensation Methods)
Part 1. Colloidal Silver. Gelatine as Protecting Colloid.
To 5 cc. of water in a test tube add about 1 cc. of M/10 AgNO 3
solution, mix well and treat with NaOH in slight excess. What is
formed?
To 5 cc. of a 5 per cent gelatine solution in a test tube add about
1 cc. M/10 AgNO 3 , mix well and treat with NaOH in slight excess.
Note any unusual action. Then heat the test tube until contents are
about to boil. What color changes occur? Dilute some of the
silver sol so formed with water and describe its color. What reduces
the silver oxide?
Repeat the above experiment, using a drop or two of hydrazine
hydrate as the reducing agent, besides gelatine.
If unsatisfactory results are obtained, repeat the experiment, using
smaller amounts of AgNO 3 solution and varying other conditions
until successful.
Part 2. Colloidal Silver. Method of Carey Lea.
Prepare two solutions as follows :
Solution A. Mix: 10 per cent silver nitrate solution 20 cc.
20 per cent Rochelle salts solution 20 cc.
distilled water 80 cc.
Solution B. Mix: 30 per cent ferrous sulphate solution . . 10.7cc.
20 per cent Rochelle salts solution 20 cc.
distilled water 80 cc.
Pour B slowly into A, stirring rapidly. The solutions must be
freshly prepared and the work should be done in light as weak as
possible.
Throw out the precipitated silver by means of a centrifuge, wash
with 2 per cent Rochelle salts solution and again separate in a cen-
trifuge.
Obtain a camels-hair brush and paint some of the silver on a watch
glass. Dry slowly (without heating above 50 C.) and note the color
of film obtained.
Place a crystal of iodine in the center of the yellow silver film.
Record all that happens. Explain.
References.
Carey Lea: Am. Jour. Science, (3) 37, 476 (1889); 38, 47, 129,
237(1889); 41,179,259,482(1891); Blake: Zeit, anorg. chem.,
37,243(1903); also Svedberg: Herstellung (1909).
Part 3. Colloidal Copper (Gelatine as Protecting Colloid.)
Mix equal volumes (5 cc.) of 10 per cent gelatine solution (freshly
prepared and warm) and 5 per cent copper acetate. To this solution
79
add, with shaking, a very slight excess of sodium hydroxide (20 per
cent). A purplish-blue, clear solution should result. If a persistent
precipitate remains, repeat the experiment, using a more concentrated
gelatine solution. Perform the same experiment, using 5 cc. of
water in place of the gelatine. What is the precipitate? Does it
dissolve in an excess of sodium hydroxide?
Heat some of the purplish-blue copper oxide-gelatine solution to
boiling and add a few drops of hydrazine hydrate. The latter is a
very powerful reducing agent and will reduce the oxide to metallic
copper in alkaline solution. Continue gently to heat the reaction
mixture until a dark, blood-red liquid is produced. The red color
is due to finely divided copper. Pour some of the liquid into water,
noting its beautiful color. In this connection cf. Paal: Ber. 35,
2206,2219 (1902).
EXPERIMENT 6
Colloidal Sulphur (Condensation Method)
Reference. Raff 6: Kolloid-Zeit., 2, 358 (1908); 8, 58 (1911).
Place a cylinder containing 70 grams of concentrated sulphuric
acid (sp. gr. 1.84) in ice water or in a freezing mixture and into it pour,
drop by drop and with constant stirring a cold solution of 50 grams
of pure crystallized sodium thiosulphate in 30 cc. of distilled water.
Work at the hoods, as H 2 S and SO 2 are given off. When the reaction
is complete, transfer the mixture to a beaker, add 30 cc. of distilled
water and warm to 80 on a water bath until SO 2 and H 2 S cease to be
given off. Then prepare a Buchner funnel and filter, connect with
the suction and pour in hot water until the funnel and filter-flask are
warm. Pour out this wash water and filter the hot sulphur hydrosol.
Cool the warm filtrate in ice water and decant the supernatant acid
liquid. To some of the precipitated sulphur add water. Is it
peptized?
To 10 cc. of this suspension add a little saturated K 2 SO 4 . What
happens? To 10 cc. add some Na 2 SO 4 solution. Is flocculation so
easy? Note difference between K 2 SO 4 and Na 2 SO 4 .
Flocculate some of the sulphur suspension by adding a soluble salt
of potassium and allow the sulphur to settle. Decant the super-
natant liquid and wash once by decantation. Then add water to
the precipitate of sulphur and shake until a coarse yellow suspension
of sulphur is formed. To this add a tiny crystal of Na 2 SO 4 . Con-
tinue to add salt cautiously until a clear yellow suspension of sulphur
is formed. What is this process called? When a large excess of
sodium sulphate is added, what happens?
SUB-GROUP 5
EMULSIONS
References. Bancroft: Jour. Phys. Chem., (1912-1918); Briggs:
Ibid., 19, 210, 478 (1915); 24, 147 (1920).
80
EXPERIMENT 1
Oil-in-Water Emulsions
Part 1. In a 150 cc. glass stoppered bottle place 45 cc. of benzene
plus 5 cc. of 1 per cent sodium oleate solution. Then shake the
bottle and contents steadily and without interruption until the ben-
zene is completely reduced to a milk-white emulsion and no free
benzene remains floating at the surface. Note the time required and
the approximate number of shakes.
Part 2. Discard the emulsion by emptying into the bottle marked
"benzene residues" and repeat the experiment making a change,
however, in the method of shaking. Give the bottle two violent
up and down shakes and then let it stand on the desk for a "rest
interval" of about thirty seconds. Continue the intermittent shak-
ing until emulsion is completed. Note the time and approximate
number of shakes. Compare with (1). Explain.
Part 3. Again discard and make the emulsion in still another way,
as follows :
In glass stoppered bottle, place 2 cc. of sodium oleate solution and
to this add 1 cc. of benzene from a burette. Shake thoroughly until
all the benzene is emulsified. Then add another cc. of benzene and
again shake. Repeat this process until about 100 cc. of benzene
have been emulsified. An emulsion having the consistency and
appearance of blanc-mange should result. As the volume of emulsion
increases, more benzene may be added each time before shaking, but
if too much is added the emulsion may "break" and a fresh start
become necessary.
Add a drop of HC1 to some of this emulsion. What happens?
Explain.
In this emulsion the oil (benzene) exists in drops (disperse phase)
and the soap solution is the dispersion medium.
EXPERIMENT 2
Water-in-Oil Emulsions
In a 200 cc. bottle, as in the previous experiment, place 10 cc. of a
benzene solution of magnesium oleate. Add water from a burette
slowly and with shaking, following a procedure similar to that of the
preceding experiment, until 40 cc. of water have been added. How
does this emulsion compare with the benzene-in-water one? In this
case the water forms the drops (disperse phase) and the soap solution
is the dispersion medium. This may be proved as follows:
Proof. On a glass plate place a drop of water and with a glass rod
stir in some of the emulsion formed in Experiment 1 . Does it mix
easily? On another portion of the plate place a drop of benzene and
stir in some of the emulsion. Does it mix easily?
Do the same thing with some of the emulsion obtained in Ex'-eri-
ment 2, that is, stir it into water and into benzene.
If the aqueous liquid is the outside phase the emulsion will mix
easily with water, but not with benzene. The reverse is true when
benzene forms the outside phase. Newman: Jour. Phys. Chem.
13,35(1914).
81
EXPERIMENTAL GROUP XVII
THERMOCHEMISTRY
It is the purpose of the following group of experiments to study the
thermal effects accompanying chemical action, change of state and
similar phenomena. Many instances of such thermal effects have
been met with in previous experiments.
References. Thomsen (Burke): Thermochemistry (1908).
Thomsen: Thermochemische Untersuchungen (1882-1886).
Sackur (Gibson) : Thermochemistry and Thermodynamics (1917) .
Journal articles.
Mathews and Germann: Jour. Phys. Chem., 15, 73 (1911);
Richards and Rowe: Proc. Amer. Acad., 43, 475 (1908) ; Richards:
Jour. Am. Chem. Soc., 31, 1275 (1909).
Procedure in Laboratory. F, 273-293 (1917); OW, 119-138;
T, 132-152.
General Directions.
For this work a simple, home-made calorimeter may be obtained
from the Instructor.
Two special thermometers are also supplied. These must be com-
pared with each other in the usual way by heating in a well-stirred
water-bath between 10 and 30 C. Number each thermometer and
reduce all subsequent readings of temperature to readings on one of
these thermometers.
Having assembled the calorimeter, determine the water equivalent
by experiment several times. How does this compare with the
calculated water equivalent?
Note. Mix weighed and approximately equal amounts of cold and
warm water so that the final temperature of the mixture is about equal
to that of the room. Weigh out water to grams only on the large
balance.
Report the water equivalent before proceeding with the experiments
which follow.
EXPERIMENT 1
Heat of Solution
Part 1. Qualitative. Half fill a test tube with finely powdered
dry NH 4 NO 3 and close tube with a rubber stopper. Then add
quickly an equal volume of cold water and mix violently to produce
instantaneous solution. Then observe the temperature of the solu-
tion. Explain the extraordinary drop in temperature. How does
82
this method of making a freezing mixture compare with the usual one
(ice-salt) ? Explain.
Read the quaint old paper on this subject by Robert Boyle, re-
printed in the Philosophical Transactions of the Royal Society (Lon-
don), 1, 86 (1666).
Part 2. Quantitative.
Procedure. T, 137.
The weighed solute is introduced into a known amount of water
contained in the calorimeter. A convenient method is to make a thin
walled glass bulb, fill it with the solute and place it in the calorimeter.
When bulb and water are at the same temperature, break the glass
and allow the solute to dissolve as quickly as possible. See that {he
solute is very finely pulverized.
Take the substance assigned from the following:
(1) NH 4 NO 3 in 200 gram molecules of water.
(2) KNO 3 in 200 gram molecules of water.
(3) NH 4 C1 in 200 gram molecules of water.
(4) KC1 in 200 gram molecules of water.
Measure the heat of the solution and derive equation (1) below,
Computations.
S = p(a + w) (t a ti) (1)
S = heat of solution in small calories; t t = initial temperature of
water and bulb in calorimeter; t 2 = final temperature when solution
is complete; a= grams of water;, w = water equivalent; 1/p =
fraction of required molecular quantities actually used experimen-
tally. For further explanation refer to Experiment 3 following.
EXPERIMENT 2
Heat of Dilution
Procedure. T, 139.
In this experiment the solution to be diluted is placed in the upper
vessel and the water is placed in the calorimeter. The solution and
water are then mixed and the thermal affect measured.
Determine the heat of dilution when a solution represented by
NaCl -f 10H 2 O is diluted with 40 gram molecules of water.
Derive equation (2) below.
Computations.
D = p { (tf tb) [(a + b) c + w] (t a tb) (a + w) } (2)
D = heat of dilution in small calories ; ta = initial temperature of
water; tb = initial temperature of solution; tf = corrected final
temperature of mixture whose specific heat = c ; w = water equiva-
lent; a = grams of water; b = grams of solution to be diluted;
1/p = fraction of required molecular quantities actually used
experimentally.
83
EXPERIMENT 3
Heat of Neutralization of 'Acids and Bases
Procedure. T, 133.
Place the acid in the calorimeter and" the base in the upper vessel.
Mix and measure the heat change.
Computations.
N = p [b (tf tb) + (a + w) (tf t a )] (3)
N = heat of neutralization in small calories; t a = temperature of
acid; tb = temperature of base ; tf = temperature of mixture ; a =
grams of water contained in solution of acid; b = grams of water
contained in solution of base; w = water equivalent. . 1/p = frac-
tion of required molecular quantities used experimentally. Here the
specific heat of the mixture is assumed to be unity.
Derive equation (3).
Part 1. Sulphuric Acid and Sodium Hydroxide.
Measure the heat of neutralization for each of the following cases:
(a) (2 NaOH + 200 H 2 O) + (1/2 H 2 SO 4 + 200 H 2 O).
(b) (2 NaOH + 200 H 2 O) + (H 2 SO 4 + 200 H 2 O).
(c) (2 NaOH + 200 H 2 O) + (2 H 2 SO 4 + 200 H 2 O).
Part 2. Phosphoric Acid and Sodium Hydroxide.
(a) (H 3 P0 4 + 200 H 2 0) + (NaOH + 200 H 2 O).
(b) (H 3 PO 4 + 200 H 2 O) + (2 NaOH + 200 H 2 O).
(c) (H 3 PO 4 + 200 H 2 0) + (6 NaOH + 200 H 2 O).
In this work one is dealing with molecular quantities of the sub-
stances involved. For instance (2NaOH + 200H 2 O) means 80
grams of NaOH dissolved in 3600 grams of H 2 O. Again, (1/2H 2 SO 4
+ 200 grams H 2 O) means 49 grams of H 2 SO 4 in 3600 grams of H 2 O.
Obviously such volumes of acid and base cannot be handled con-
veniently, so one chooses some convenient fractional part of the acid
and base solution, for example, 1/16 whence 1/p = 1/16. The quan-
tity of the solutions to use in the case of H 2 SO 4 and NaOH (Part 1)
would be found thus:
1/16 (80 + 3600) = 230 grams of the NaOH solution.
1/16 (49 + 3600) = 228 grams of the H 2 SO 4 solution.
To make up this acid solution mix 3.06 grams of H 2 SO 4 with .225
grams of H 2 O.
H 2 SO 4 and H 3 PO 4 tables may be found in the Kalendar, Vol. 1,
and elsewhere.
EXPERIMENT 4
Thermoneutrality of Salt Solutions
Measure the heat change when solutions of the following pairs are
mixed:
1 . NaCl + 200H 2 O and KNO 3 + 200 H 2 O.
2. NH 4 C1 + 200H 2 O and KNO 3 + 200 H 2 O.
Take the pair assigned, placing one solution in the upper vessel
and the other in the calorimeter.
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EXPERIMENTAL GROUP XVIII
PHOTOCHEMISTRY
The purpose of the following experiments is to study qualitatively
the action of light in producing and accelerating chemical change.
References.
Bancroft: Electrochemistry of Light, Jour. Phys. Chem.
(1908-1912).
Bancroft: Orig. Comm. 8th Int. Cong. App. Chem., 20, 31 (1912).
Sheppard: Photochemistry (1914).
EXPERIMENT 1
Soluble and Insoluble Sulphur
Saturate 10 cc. of CS 2 with roll sulphur. Work in the hood.
Then divide into three portions and place in loosely stoppered test
tubes.
(a) Expose one test tube to direct sunlight. After precipitation
of amorphous sulphur has occurred, set aside in a dark place. The
amorphous sulphur will dissolve. It may be necessary to wrap the
test tube in dark paper to protect it from the light.
(b) Place another portion in a test tube which is immersed in a
solution of CuSO 4 .
(c) Place the third portion in a test tube which is immersed in a
solution of K 2 Cr 2 O7.
Reference. Rankin: Jour.*Phys. Chem , 11, 1 (1907).
Note. Do not stopper the test tubes too tightly when exposing
to the sunlight.
Discussion.
This experiment shows in a very satisfactory way how light dis-
places the equilibrium:
Q " C
soluble insoluble.
It also shows that the activity of light differs for different wave
lengths. In a certain sense "light is a mixture of reagents." Light
of a particular wave length is active if it is absorbed, and absorbed
light tends to shift the equilibrium in such a way as to favor the
production of the substance which absorbs the particular light less
readily.
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