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45
UC-NRLF
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GIFT OF
Harry East Miller
A LABORATORY MANUAL
OF
General Chemistry
For Use in Colleges
BY
WILLIAM C M BRAY,
Professor of Chemistry in the
University of California
AND
WENDELL M. LATIMER
Instructor of Chemistry in the
University of California
A LABORATORY MANUAL
OF |
General Chemistry
For Use in Colleges
BY
WILLIAM C. BRAY,
Professor of Chemistry in the
University of California
AND
WENDELL M. LATIMER
Instructor of Chemistry in the
University of California
Copyright, 1921, By William C. Bray and Wendell M. Latimer
Price, Fifty Cents
LEDERER, STREET & ZEUS Co., Publishers
BERKELEY, CALIFORNIA
1921
PREFACE
The present laboratory manual has been prepared primarily for the use of
students in general inorganic chemistry in the University of California. These
students have usually had a year in elementary chemistry in high school, and
many of them will take no further work in chemistry. No distinction is made
between students on the basis of the various curricula which they are following,
as we believe that a course in the fundamentals of general chemistry is equally
suitable for all students.
The laboratory and lecture work are correlated as closely as possible. In the
present manual page references are given to Professor Joel H. Hildebrand's
"Principles of Chemistry/' MacMillan, 1918, the reference book written for the
course.
The laboratory work is a study of chemical principles, rather than a presen-
tation of descriptive material. It is hoped that the division of the manual into
Sections, and the statements in the first paragraphs of the various Assignments,
will materially assist the student in recognizing the relation between the
experimental details and the principles involved.
When the course extends over two terms, as at the University of California, a
satisfactory division is to take Sections I to III in the first term, though in some
cases it may be possible also to begin the first Assignment on Qualitative Analysis.
It is recommended that the Assignments in the last two Sections be taken in the
order noted in the text.
The following editions of the manual have been printed : Laboratory Directions
in Chemistry 1A, edited by William C. Bray, 1915; 21 Assignments. A Labora-
tory Manual of General Chemistry, William C. Bray and Ludwig Rosen stein,
1916; 26 Assignments. The, 6 s^rne, . revised, by. William C. Bray, 1917; 31
Assignments; reprinted 1918, \9i9,l92^J frfiS present manual contains 5 Sections
with a total of 30 Assignments, and is an* almost complete revision of the 1917
manual. A : .::: : ".i :Y ': : :/.'"":
In the development of this manual from 1912 to the present time a great deal
has been contributed by the instructors in the course. We wish especially to
acknowledge our indebtedness to Professors G. N. Lewis, J. H. Hildebrand,
Edward Booth and E. D. Eastman and to Doctors Ludwig Rosenstein and
W. L. Argo.
WILLIAM C. BRAY,
WENDELL M. LATIMER.
June, 1921.
*
a
TABLE OF CONTENTS
Pago
Note to Students 5
Special Laboratory Directions 5
List of Apparatus 6
SECTION I. WEIGHT RELATIONS IN
CHEMICAL REACTIONS.
Assignment 1. A Chemical Reaction: The Synthesis of a Sulfide
of Copper 7
Notes on Glass Manipulation 8
Assignment 2. The Relation between the Mass and Volume of
Gases: The Determination of the Volume of a Mol of Oxygen.. 9
Assignment 3. The Reaction between Certain Metals and Hydrochloric
Acid 11
Assignment 4. The Analysis of Copper Oxide 13
Assignment 5. The Reaction between an Acid and a Base in Solution.
Concentration in Terms of Mols per Liter 15
Assignment 6. The Titration of Solutions of Acids and Bases : An
Illustration of Volumetric Analysis 17
Assignment 7. Volumetric Analysis, Continued : The Determination of
the Equivalent Weight of an Unknown Acid 19
SECTION II. IONIC THEORY.
RAPID REVERSIBLE REACTIONS AND EQUILIBRIUM
Assignment 21. Solutions of Strong Electrolytes 21
Assignment 22. Strong and Weak Acids. The Use of Indicators
to Measure Hydrogen Ion Concentration 23
Assignment 23. Strong and Weak Bases. The Use of Indicators
to Measure Hydroxide Ion Concentration 26
Assignment 24. Rapid Reversible Reactions and Equilibrium 27
Assignment 25. The Reversibility of Neutralization Reactions. Hydrolysis.... 30
SECTION III. REACTIONS OF IONS
Assignment 31. The Properties of Sodium, Potassium and Ammonium
Ions. Tests for Chloride, Sulfate and Nitrate Ions 33
Assignment 32. Calcium Ion 36
Assignment 33. Carbonate Ion, Bicarbonate Ion and Carbonic Acid 38
Assignment 34. Sulfates, Chlorides and Nitrates of Copper, Silver and
Zinc 42
Assignment 35. Hydroxides of Copper, Silver and Zinc 44
Assignment 36. Complex Ions of Copper, Silver and Zinc with Ammonia 46
Page
Assignment 37. Carbonates and Sulfides of Copper, Silver and Zinc 48
Assignment 38. Review of the Chemistry of Positive Ions Already
Considered - 50
SECTION IV. REACTIONS OF IONS, CONTINUED
Assignment 41. Oxidation and Reduction. Replacement Reactions.
Electrical Cells 52
Assignment 42. Oxidation of Metals to their Ions. Table of Oxidizing
and Reducing Agents 54
Assignment 43. Ferrous and Ferric Ions 56
Assignment 44. Mercurous and Mercuric Ions 58
Assignment 45. Lead Ion, Chromate Ion 61
Assignment 46. Stannous and Stannic Ions. Amphoteric Sulfides 61
Assignment 47. Ions of Aluminum and of Chromium. Peroxides 62
SECTION V. QUALITATIVE ANALYSIS
Assignment 51. The Development of a Scheme of Analysis for a
Limited Number of Positive Ions 65
(To follow Assignment 38)
Assignment 52. The Standard Scheme of Analysis. Methods of
Dissolving Difficulty Soluble Substances 67
(To follow Assignment 44)
Assignment 53. Lead, Tin, Aluminum, Chromium and Barium in the
Scheme of Analysis 70
(To follow Assignment 47)
NOTE TO STUDENTS
1. Decide what is the real purpose of each Assignment. Before beginning
the experimental work and preferably before coming to the laboratory, read the
first paragraphs of the Assignment, study carefully the References, and review
earlier, related work. In general look for the connection between the lectures
and laboratory work and between each Assignment and the preceding ones.
2. Master each idea before proceeding to the next one. Form the habit
of at once consulting the instructor whenever you are not certain of the cor-
rectness of your answer to questions and of your conclusions from the experi-
ments. The frequent short written examinations will be of great assistance to
you in deciding whether or not you have really understood the work.
3. An average student who has understood the earlier work can complete
an Assignment in the regular laboratory time allotted to it. Students who can-
not finish in the stated time, announced by the instructor, must consider them-
selves behind the class, and should plan immediately to do extra work at home
and in the laboratory. For their convenience the building is open from
8:00 A.M. to 4:30 P.M. (Saturdays 8:00 to 12:00).
4. Your success depends upon your own efforts. If you are in serious
difficulty then there is something wrong with your methods. The instructor
can assist you in finding out what is wrong, but he cannot do your work for
you. No effort on his part can make up for lack of initiative on your part,
failure to assume the responsibility of mastering each idea, or inability to
improve your method of doing the work.
SPECIAL LABORATORY DIRECTIONS
5. A laboratory note book about 6^ inches wide, opening at the side, and
not loose-leaved, is recommended. This book and the laboratory manual will be
needed at each meeting of the laboratory section, including the first one.
6. The recording of experiments, observations and conclusions at once in
the note book is an essential part of the laboratory work. Entries made from
memory or from memoranda on scraps of paper are not records of the experi-
mental work. The value of the original record is improved: by dating each
day's notes; by numbering the pages of the note book; by never erasing an
entry or tearing out a page; by leaving space for additions and corrections;
by adopting a plan of distinguishing between the record of the experiments
actually performed, and the other entries such as answers to questions, solu-
tions of problems, etc. ; by making all calculations neatly at the bottom or side
of the page ; and by writing entries in such a way that they will be easily under-
stood when the work is reviewed. The descriptions of the experiments per-
formed should be very brief when detailed directions are given in the manual,
but must be complete when, as in the later work, experiments are devised by
the student. A passing grade in the laboratory work will not be given unless
the experiments have been completed and the results properly recorded in the
note book.
7. The laboratory desk must be kept neat and dry. An old towel should be
used for cleaning the desk top and another towel should be kept clean for use
on apparatus. When cleaning apparatus use tap water and a brush to remove
all visible dirt and rinse finally with a little distilled water. Before leaving the
laboratory the apparatus should be locked up in the desk.
8. The wash-bottle should only be used to hold distilled water. Before using
sterilize the mouth-piece by boiling in water, and never lend or borrow a
wash-bottle.
9. The contamination of laboratory reagents can be avoided by keeping each
stopper clean and replacing it at once in the proper bottle, and by never pouring
anything back into a reagent bottle.
10. Experiments which give rise to disagreeable or dangerous fumes must
always be performed out of doors or in a fume-closet.
11. At the first meeting of each laboratory section the instructor will dis-
tribute the desk keys together with lists of apparatus similar to the one given
below. Check the apparatus in the locker, exchange damaged articles at the
office, sign the list of apparatus (surname first) and return it to the instructor.
Begin work on Assignment I.
LIST OF APPARATUS
1. Regular equipment of each locker. Additional articles may be obtained
at the office by filling out an "order slip" and signing your name and locker
number. Whenever any article is returned to the office sign a "return slip." At
the end of the term the locker must contain the same amount of apparatus, no
more and no less; the locker must be clean; the apparatus must be clean and
dry, and in good condition; glass stoppers must fit, and be protected by the
insertion of a piece of paper.
1 Key.
5 Beakers, 100 cc., 150 cc., 250 cc.,
400 cc., 600 cc.
5 Reagent Bottles.
2 Sample Bottles, 50 cc.
1 Graduated Cylinder, 50 cc.
4 Flasks, 500 cc., 250 cc., and two
125 cc.
1 Wash-bottle, equipped with glass
tubing and rubber stopper.
2 Funnels.
2 Blue Glasses.
2 Glass Rods, 12 cm. and 18 cm.
30 cm. Glass tubing.*
12 Test-tubes.
1 Watch Glass.
1 Casserole.
1 Crucible, with cover.
2 Evaporating Dishes.
1 Crucible Tongs.
1 Bunsen Burner, with rubber tubing.
1 Iron Wire.*
1 Wire Gauze.*
1 Triangle.
1 Test-tube Brush.*
1 Test-tube Holder.
1 Test-tube Rack.
1 Package Filter Paper.*
1 Rule.
2 Towels.
Litmus Paper* in a bottle.
2. The following additional articles may be obtained at the office :
(a) By signing the regular order slips. Small short-stemmed funnels ; glass
flasks, 50 cc. ; matches * ; corks ; rubber stoppers.
(b) By signing "temporary order ships." Special apparatus for Assignments
2 and 4; burettes, with clamps and pinch-cocks; graduated cylinders, 10 cc.
and 250 cc. ; thermometers ; paraffin. These articles should be returned
when possible during the same laboratory period.
*Not returnable. At the end of the first term students should retain these articles for
use in the second term.
SECTION i v /.\ \ fj \ \- v K - /
WEIGHT RELATIONS IN CHEMICAL REACTIONS
ASSIGNMENT 1
A CHEMICAL REACTION : THE SYNTHESIS OF A
SULPHIDE OF COPPER
References. Hildebrand, Principles of Chemistry, Chapter I, and pages 40-43.
1. In Assignment 1 we shall study quantitatively a chemical reaction in which
two elements, * a metal and a non-metal unite to form a pure compound. The
experiment consists in determining the weight of the compound that is formed
from a weighed amount of copper when heated with excess of sulfur. From
these experimental data, and the atomic weights of the two elements, the relative
number of atoms of copper and sulfur in the compound will be calculated.
Questions. If 3.04 grams of a certain metal, when burned in oxygen, yield 5.04 g.
of a pure compound of the metal and oxygen, what weight of oxygen will combine
with 1.00 g. of this metal? What additional information is necessary before the
relative number of atoms of the two elements in the compound can be calculated ?
2. Experiment. Support a clean porcelain crucible, with a cover, on a triangle
and heat with the colorless flame of a bunsen burner to low r redness. Let the
crucible cool about 10 minutes, and weigh it, with the cover, to 10 milligrams.
Note. Do not make any weighings until instructions in the use of the balance
have been given.
3. While the crucible is cooling obtain from the shelf a clean piece of copper
wire, weighing about 1 gram, and weigh it to 10 mg.
4. Place the copper in the weighed crucible and add enough powdered sulfur
to cover the copper. Place the cover on the crucible and heat gently (with a
small flame) until the sulfur ceases to burn at the edges of the cover, but do not
remove the cover while the crucible is hot. Then heat more strongly until the
bottom of the crucible just becomes dull red. Again allow to cool about 10
minutes and weigh.
5. Carefully remove the cover and note the appearance of the contents of the
crucible, but do not touch the substance. If there is any free sulfur on the cover
or the wall of the crucible, replace the cover, heat the crucible and cover, and
weigh again. Check the accuracy of the final weight by adding a small quantity
of sulfur and repeating the experiment ; continue until two consecutive results
agree within 10 mg. At the end of the experiment remove the substance formed,
break it and describe its properties. Clean the crucible with hot nitric acid in
a porcelain dish, wash with distilled water, dry by heating, and check the
original weight.
6. Questions. What conclusions can you draw from each of the following
observations: (a) the properties of the product are different from those of
either copper or sulfur; (b) the product appears to be homogeneous and its
weight is greater than that of the copper used? What additional evidence is
necessary to prove that the product is a pure substance and not a solid solution?
7. Calculations. Summarize your experimental results and make the calcula-
tions necessary to complete a table similar to the following:
* It is suggested that the student form the habit of writing out the meaning of each
italicized word in the text and of giving an example whenever possible.
[7]
(a) Weight of crucible
(b) Weight of copper
(c) Weight of crucible and product
(d) Weight of product
(e) Difference between (d) and (b)
First Second Value
Weighing Weighing Chosen
Calculate what the increase in weight (e) would have been if one gram-atom
of copper had been used in the experiment, and enter in the table as line (/) ;
show this result to your instructor at once. How does this number compare with
the atomic weight of sulphur? How many gram atoms of sulfur have combined
with one gram atom of copper? What, then, is the simplest formula of the
substance formed? What is the corresponding molecular weight? Write the
equation for the reaction, and write out in words what this equation means, in
terms of (a) atoms and molecules, (b) gram atoms and mols, (c) grams, and
(d) pounds.
8. Problems. (/) In order to determine the effect of a small error in
weighing the copper repeat your calculations, Paragraph 7, using for the weight
of copper a value 10 mg. greater than your experimental value. What per cent
of the weight of copper is 10 mg. ? This would be the percentage error in the
weight of copper if a 10 mg. error in weighing had been made. What is the
corresponding percentage error in your value for the weight of sulfur that
would combine with one gram atom of copper?
(2) Calculate the percentage composition of the copper sulfide formed :
(a) from your experimental data, and (b) from the formula and the atomic
weights of copper and sulfur. Compare the results.
(j) A sulfide of iron contains 53.8% iron. What is the formula? (In solving
this problem consider one gram atom of iron and one gram atom of sulfur as
the fundamental units for iron and sulfur. Calculate first the weight of sulfur
and then the number of gram atoms of sulfur combined with one or more
gram atoms of iron.)
(4) The formulas of hydrogen sulfide and of ferrous sulfide are H L .S and
FeS, respectively. What are their molecular weights. What weight of sulfur is
contained in one mol of hydrogen sulfide? In one mol of ferrous sulfide? What
weight of hydrogen sulfide could be made from one mol of ferrous sulfide?
NOTES ON GLASS MANIPULATION
To bend a piece of ordinary glass tubing, hold it with both hands in a
fan-shaped gas flame and rotate it slowly between the thumb and fingers until
a 2]/2. to 3 inch portion is uniformly heated and is soft enough to be bent to the
proper angle. Set it aside to cool ; glass will remain hot enough to burn the
hand for some time after it no longer appears to be hot.
To cut glass tubing, scratch it with a file at the proper place, grasp it firmly
on each side of this mark (protecting the hands with a cloth), and bend the
tube away from the mark.
Always remove the sharp edges of freshly cut glass at once with a file, or by
heating in a gas flame.
To draw down a piece of tubing to a capillary, heat a portion about 1 inch
long in an ordinary gas flame to a higher temperature than was necessary in
bending the tubing. Hold the tube with both hands and rotate it to ensure
uniform heating and prevent the hot portion from sagging. Withdraw from the
flame and draw apart slowly to obtain a thick-walled capillary.
[81
ASSIGNMENT 2
THE RELATION BETWEEN THE MASS AND VOLUME OF GASES: THE
DETERMINATION OF THE VOLUME OF A MOL OF OXYGEN
References. Hildebrand, Chapter II, and pages 52 and 57
1. It is often necessary to know the volume of a given mass of a substance, or
conversely the mass of a given volume. While in the case of a solid or liquid
the relation between the mass and the volume must be determined for the
particular substance, the problem is simplified when we are dealing with a gas,
since a mol of every gas occupies nearly the same volume under similar condi-
tions. In this Assignment we shall determine, under definite conditions of
temperature and pressure, the volume of a known weight of oxygen and calculate
the volume of one mol at standard conditions. Questions. What information
is needed before you can calculate the weight of 10 cc. of mercury? A given
solution of sodium chloride in water contains 25.0 percent sodium chloride and the
density of the solution is 1.19 g. per cc. ; what volume of solution in cc., and in
liters, contains 100 grams of sodium chloride?
2. When solid potassium chlorate is strongly heated it decomposes with the
evolution of oxygen, and the loss in weight gives the weight of oxygen evolved.
The volume of the oxygen is determined by measuring the volume of water
displaced by an equal volume of oxygen. When pure potassium chlorate is used
it must be heated in a hard glass (difficulty fusible) test-tube. The potassium
chlorate decomposes more readily and at a lower temperature when a small
quantity of manganese dioxide is present. The hard-glass test-tube may then
be replaced by a heavy-walled test-tube of ordinary easily-fusible glass ; but care
must be taken not to heat the latter tube to a higher temperature than is necessary
for the reaction. The manganese dioxide is a catalyst in this reaction, and all
of it may be recovered after the potassium chlorate has been decomposed into
potassium chloride and oxygen.
3. Two students working together obtain from the office a heavy-walled glass
tube, a rubber stopper, rubber tube, pinch-cock, clamp and small tube containing
about 5 grams potassium chlorate. The yellow order slip for "special" apparatus
for Assignment 2 should be signed by both students. The apparatus should be
returned as soon as the experiment is finished. Each student must keep a complete
record of the experiment in his notebook.
4. Experiment. Set up the apparatus according to the accompanying diagram.
Since the apparatus must be gas tight, glass tubing and stoppers must be care-
fully fitted. If your tubes and stoppers do not fit, exchange them at the office.
Do not use the glass tubing of your wash bottle. Directions for bending glass
tubing are given on page 8.
5. Place in the heavy-walled test-tube about 5 grams potassium chlorate. Add
about 50 mg. manganese dioxide, estimating the amount by comparison with the
sample in the laboratory; mix it with the
potassium chlorate, by jarring the tube; and
wipe off any powder that is on the outside of
the tube or on the inside near its mouth.
Assemble the apparatus as before.
6. Heat the tube gently with a moving gas
flame, leaving the pinch-cock open on the
rubber tube outlet. Moisture will appear on
the walls of the test-tube, which shows that
the potassium chlorate and manganese dioxide
were not perfectly dry. Gradually warm the
tube to within about one inch of the
stopper and at the same time heat the potassium chlorate until gas
evolution begins and some water passes over into the beaker. Drive out
the moisture with the oxygen by carefully heating the walls of the tube,
but do not scorch the rubber stopper. When 50 to 100 cc. water have passed into
the beaker, allow the apparatus to cool; close the pinch-cock near the outer end
of the rubber tube, which should now be filled with water. Disconnect the
chlorate tube, and weigh it with the dry material inside to 10 mg.
7. Immediately replace the tube in its proper position, open the pinch-cock
while the end of the delivery tube is under the water, raise the beaker until the
surfaces of the water inside and outside the flask are at the same level, close the
pinch-cock again, place a dry beaker under the delivery tube, and open the pinch-
cock. If the water continues to siphon into the beaker your apparatus is not gas
tight and must be rebuilt. Again heat the tube gradually until gas evolution
begins, and continue to heat the potassium chlorate carefully and not too
strongly until from 250 to 300 cc. water have been forced over into the beaker.
Allow the tube to cool, equalize the level of the water in the beaker and the flask,
and then close the pinch-cock on the siphon tube. By means of a 250 cc. graduated
cylinder measure the amount of water which the oxygen has forced out of the
flask. Finally weigh carefully the hard glass tube containing the partially
decomposed chlorate. Question. Why is it necessary to cool the test-tube and to
have the water in the beaker and flask at the same level before closing the
pinch-cock?
8. Repeat the experiment, Paragraph 7. The three weighings and two volume
measurements give two independent sets of experimental data. Compare the
results of your two experiments by preparing a table which will show for each
experiment : the weight of oxygen, the corresponding volume of water displaced,
the data referred to in 10 to 11 below, and the results of the calculations (12).
9. Clean the test-tube by placing water in it and shaking. The manganese
dioxide is difficultly soluble in water, while both potassium chlorate and potas-
sium chloride dissolve readily. Suggest a method of recovering the manganese
dioxide and obtaining a mixture of dry potassium chloride and potassium
chlorate practically free from manganese dioxide.
10. To make the calculations it will be necessary to know the barometric
pressure at the time you perform each experiment, and the temperature of the
water in the flask. The temperature of the water may be assumed to be that of
the room, and the barometric pressure will be posted on the blackboard. Enter
these data in your notebook before leaving the laboratory. Below is given a table
of the vapor-pressure of water at different temperatures :
Vapor Pressure of Water.
Temp. C. Vapor Pressure Temp. C. Vapor Pressure
14 1.2 cm. mercury 24 2.2 cm. mercury
16 1.3 " ' " 26 2.5 "
18 1.5 " " 28 2.8 "
20 1.7 " " 30 3.2 "
22 2.0 " 32 3.5 "
11. Questions. Assuming that levels of the water in the beaker and the flask
were the same when the pinch-cock was closed, what was the total pressure of
the gas in the flask? What was the partial pressure of the water- vapor? Of
the oxygen?
12. Calculations. From each of your two sets of experimental data, by means
of the Gas Laws, calculate the volume of 1 mol (32 grams) of pure oxygen at
1 atmosphere pressure and O C. Compare your results with the value given in
your text book. Your results should not differ from this value by more than
F 101
5 percent (check your calculations). Show your tabulated results to your
instructor, and repeat the experiment if necessary.
13. Questions. What is the formula of the oxygen molecule? What value is
usually accepted as a close approximation for the volume of 1 mol of gas under
standard conditions? Use this value to calculate (a) the weight of a liter of
hydrogen chloride gas, HC1, under standard conditions; (b) the molecular
weight of a gas whose density at standard conditions is known to be 0.001977
g. per cc.
14. Write an equation to represent the decomposition of potassium chlorate,
KC1O 3 , into potassium chloride, KC1, and oxygen; and write out what this
equation means in terms of (a) molecules; (b) mols, and (c) grams of the
substance involved.
15. Problems, (i) One gram of potassium chlorate is completely decomposed
into potassium chloride and oxygen. Calculate (a) the weight of oxygen that
could be obtained ; (b) the volume of the oxygen (in liters and in cc.) at standard
conditions, and (c) the volume of the oxygen at 27 C. and 750 mm. mercury
pressure.
(2) It is an experimental fact that 2 volumes of carbon monoxide gas react
with 1 volume of oxygen to form 2 volumes of carbon dioxide gas. Give the
reasoning by which, from this result you can conclude that the molecule of
oxygen contains an even number of atoms.
(5) The formulas of carbon monoxide and carbon dioxide are CO and CO 2 ,
respectively. Write the equation for the reaction considered in the preceding
question, and interpret it in terms of (a) mols, (b) liters, and (c) grams.
ASSIGNMENT 3
THE REACTION BETWEEN CERTAIN METALS AND HYDROCHLORIC ACID
References. Hildebrand, pages 84-86, and 47-50.
1. Certain metals, aluminum, zinc, magnesium, etc., react with a solution of
an acid, with evolution of hydrogen gas and formation of a salt in solution. In
this assignment we shall dissolve a definite weight of a metal in excess of hydro-
chloric acid, measure under definite conditions of temperature and pressure the
volume of the hydrogen liberated, and calculate the weight of the metal that
would form one gram-atom of hydrogen. From this result and the atomic
weight of the metal we can then determine : the number of atoms of hydrogen
formed when one atom of the metal reacts with the acid, the number of molecules
of acid (HC1) which react with one atom of metal, and the formula of the
chloride of the metal formed. Questions. What experimental facts and
reasoning have led to the conclusion that the formula of the hydrogen molecule
is H ? If the valence of hydrogen in HC1 is + 1 what is the valence of the
chlorine in this compound?
2. The instructor will supply to each student a sample of a metal as an
"unknown." Experiment. Take a portion of the metal weighing between 0.4 and
0.5 g. Clean it, if necessary. Weigh to 5 mg.
3. Obtain a small short-stemmed funnel at the office, and select a beaker of
such size that the funnel when placed in it can be completely covered with water.
Place the weighed metal in the beaker, place the inverted funnel over it, and
pour freshly distilled water into the beaker until the funnel is completely covered.
Note. Tap water contains a relatively large amount of dissolved air, and should
[in
not be used in this experiment unless it has been heated to boiling to expell the
greater part of dissolved air.
4. Pour distilled water into a half liter flask until the water completely fills the
flask. Moisten a piece of filter paper slightly larger than the mouth of the flask, cover
the mouth of the flask with paper, taking care that no bubble of air remains below
the paper. Invert the flask (over an empty vessel) and lower it into the beaker
in such a manner that the stem of the funnel enters the neck of the flask. If
a bubble of air enters the flask repeat this operation. The apparatus now
consists of a beaker containing a funnel inverted over the metal, and a flask
filled with water and inverted over the funnel. Place this apparatus in a large
beaker or other vessel, to prevent the water from overflowing on the desk during
the remainder of the experiment.
5. Insert a thistle tube or long-stemmed funnel into the water so that the lower
end touches the bottom of the beaker at the rim of the inverted funnel, and
through it pour 25 cc. concentrated hydrochloric acid. If the liquid is not stirred
the concentrated acid, which is 1.18 times as dense as water, will remain for some
time as a layer at the bottom of the beaker, and the metal will be dissolved
rapidly. If all the water in the inverted flask is displaced by the hydrogen you
have used too much metal or too small a flask and must begin the experiment
over again.
6. When the metal has all dissolved (except a few dark-colored flakes of
impurities of negligible weight), place the apparatus in a large basin of tap
water and carefully remove the beaker and funnel without allowing any air to
enter the inverted flask. Keep the flask in the water for several minutes in order
that it may be at the same temperature as the water. Then raise or lower the
flask until the level inside and outside the flask is the same. (What is now the
pressure of the gas inside the flask?) While the flask is in this position, cover
the mouth of the flask with the palm of the hand, remove the flask from the water
and invert it. While the gas is escaping, test to prove that it is hydrogen.
7. Measure the volume of the gas which was contained in the flask by filling
the flask completely with water and observing the volume needed. Record in
your notebook the barometric pressure (written on the blackboard) and the
temperature of the water in which the flask was immersed.
8. You now have the weight of metal taken, and the volume, at a definite
temperature and pressure, of a corresponding amount of hydrogen saturated
with water vapor. What is the partial pressure of the water vapor at the
temperature of the experiment? What was the partial pressure of the hydrogen
in the flask?
9. Calculate from these data :
The volume at standard conditions that the hydrogen would occupy if it
were dry.
The weight of the hydrogen. (Use the molecular weight 2.016 and the volume
of 1 mol of gas, Assignment 2.)
The weight of metal that would have liberated 1 gram-atom of hydrogen.
Report this value to your instructor, who will tell you the name of the metal if
your result is correct to within about 5%.
By means of the atomic weight of the metal calculate the number of atoms of
hydrogen formed when 1 atom of the metal dissolves in acid.
10. Questions. How many molecules of HC1 react with 1 atom of the metal ?
Assuming that the hydrogen of the acid is replaced by the metal, what is the
formula of the chloride formed? What is the valence of the metal in this
compound? Write the equation for the reaction.
11. Problems. (i) From the density of hydrogen at standard conditions,
[12.]
0.00008987 g. per cc., calculate the actual volume of 1 mol of hydrogen. What
percentage error did you make in the calculations in Paragraph 9 by assuming
the value 22.40 liters.
(2) If sulfuric acid, H 2 SO 4 , had been used in the above experiment instead of
hydrochloric acid, the same result would have been obtained and the final solution
would have contained a sulfate of the metal. Write the equation for the reaction.
(j) The student will have noted that in the first three Assignments we have
assumed a knowledge of atomic weights. The arbitrary choice of the unit
0=16.00 should present no difficulty, but it is often not clear why a particular
value is chosen for the atomic weight of an element rather than some fraction or
multiple of this value, e. g., in the case of chlorine why 35.46 is chosen instead
of say 17.73 or 70.92. To illustrate how this choice is made on the basis of
the experimentally obtainable quantities, molecular weight and percentage
composition, the following data may be used. (The molecular weights given in
the second column of the table are the accurate values, and not approximate
values such as would be obtained directly from the weight of 22.40 liters of gas
reduced to standard conditions.)
No. Grams Chlo-
Substance Molecular Weight % Chlorine rine in i Mol
Chlorine 70.92 100
Hydrogen chloride 36.47 97.2
Chlorine oxide (/) 86.92 81.6
Chlorine oxide (?) 67.46 52.6
Phosphorus chloride 137.42 77.4
Carbon chloride 153.84 92.3
Calculate the values required for the fourth column of the table. What value
would you choose for the atomic weight of chlorine? No compound of chlorine
has even been made which contains in 1 mol less than 35.46 grams of chlorine.
How many atoms of chlorine are contained in a molecule of each of the six
substances listed in the table?
ASSIGNMENT 4
THE ANALYSIS OF COPPER OXIDE
Reference. Hildebrand, Chapter III.
1. In this Assignment, as an example of chemical analysis, we shall determine
the composition of an oxide of copper. The analysis will be made by heating a
weighed portion of the oxide in a current of hydrogen and weighing the metallic
copper which remains. The oxygen of the oxide unites with the hydrogen to
form steam. As in Assignment 1, we shall determine the formula of the
compound by assuming the atomic weights of copper and oxygen. It is to be
noted, however, that the results could be used to determine the relative atomic
weights of copper and oxygen if the formula of the compound were known.
Our experimental data, of course, will not be sufficiently accurate to make worth
while the calculation of the atomic weight of copper.
2. Experiment. Two students may work together ; both should sign the order
slip for "special apparatus for Assignment 4," which consists of a thick-walled
hard glass test tube with a rubber stopper and glass tubes, a thistle tube, and
two-holed rubber stopper, a clamp, and a calcium chloride tube (with 2 rubber
stoppers, 2 glass tubes, 2 rubber tubes and 2 short glass rods). The apparatus
should be returned as soon as the experiment is finished.
3. Set up a ''hydrogen generator" by fitting your half liter flask with a thistle
tube extending through the rubber stopper nearly to the bottom of the flask,
[13]
and an outlet tube bent at right angles (see note on glass manipulation). Place
in the flask about 10 grams of zinc and cover it with about 100 cc. water. To the
outlet tube attach a "drying tube" (containing solid calcium chloride, which has
the property of absorbing moisture). Make sure that the apparatus is air-tight
and wrap the flask in a towel.
4. Set up the remainder of the apparatus according to the directions of the
instructor. Dry the thick-walled glass test-tube that is to contain the copper
oxide by heating it gently. When it is cool weigh it carefully, together with
any portion of the apparatus that may come in contact with the copper oxide.
Place in the tube about 1 gram of copper oxide, wipe off any particles that are
not in the portion of the tube that is to be heated. Weigh again carefully to
obtain the weight of copper oxide used. Attach the apparatus to the hydrogen
generator, pour about 40 cc. concentrated hydrochloric acid down the thistle tube,
and allow the hydrogen to pass through the apparatus until it has swept out
the air. (Caution. Do not place a flame near the outlet nor heat the oxide
while the apparatus contains a mixture of oxygen and hydrogen. A dangerous
explosion might result.) Collect the gas in small test tubes by displacement
of water and ignite it. Explain how this test may be used to determine when the
hydrogen is no longer mixed with oxygen.
5. When pure hydrogen is passing over the copper oxide, begin to heat the oxide
very gently with a small flame and continue to heat cautiously until all the oxide
changes color. If moisture collects in the farther end of the tube, drive it out by
heating the tube carefully. Question. Where does this moisture come from?
6. Allow the tube to cool in the current of hydrogen, and weigh it. If you
have time, check this result at once by repeating the heating in the current of
hydrogen and then weighing; if not, set the tube aside in order that you may do
so if the results of the following calculations are unsatisfactory.
7. Calculate the number of (a) grams; (b) gram atoms of copper that are
combined with 1 gram atom of oxygen. What, then, is the formula of this oxide
of copper? Repeat the experiment if your results are inconclusive.
8. Write the equation for the reduction of copper oxide by hydrogen, and
interpret in terms of (a) atoms and molecules; (b) gram atoms and mols, and
(c) grams.
9. Calculate the percents of copper and of oxygen in this oxide of copper
(a) from your experimental results; (b) from the formula and the atomic
weights of copper and oxygen.
10. Problems. ( i) The formulas of cuprous oxide and cupric oxide are Cu 2 O
and CuO, respectively. Write equations for the reactions between the heated
oxides and hydrogen to form copper and steam. What weight of copper would
be obtained from one gram of each oxide? What is the percentage composition
of each oxide?
(2) What weight of water could be obtained from 1 gram of cupric oxide?
What volume would this amount of water occupy at 1 atmosphere pressure and
(a) 4 C; (b) 273 C? What volume of hydrogen at 273 C and 1 atmosphere
pressure is required to form this amount of water? Is this the amount that
would be used in an experiment similar to the one actually performed?
(3) What are the formulas of cuprous sulfide, cupric sulfide, and hydrogen
sulfide? Cuprous and cupric sulfides are also reduced to copper when they are
heated in a current of hydrogen; hydrogen sulfide is formed. Write equations
for the reactions.
14
ASSIGNMENT 5
THE REACTION BETWEEN AN ACID AND A BASE IN SOLUTION
CONCENTRATION IN TERMS OF MOLS PER LITER
References. Hildebrand, Chapter V, pages 76-80, Chapter VIII, pages 105
and 106.
1. In this Assignment we shall study the reaction between sodium hydroxide
and hydrochloric acid in solution and shall determine the amount of salt that is
formed from a measured volume of sodium hydroxide solution. In the preceding
Assignments the amount of a substance was determined by weighing or by
measuring the volume of the pure substance. It is often more convenient to
determine the quantity of a substance by measuring the volume of a solution
which contains a known amount of the substance in a unit volume of the solution.
The amount of the substance in a unit volume of solution is called the
concentration. Question. If the concentration of a salt solution is known to be
10 g. per liter, what volume would you measure out in order to have 0.5 g. of salt?
2. It is necessary first to examine separately the properties of the three solutions,
the base, acid and salt. Experiment. Prepare dilute solutions of sodium
hydroxide, NaOH, and of hydrochloric acid, HC1, by diluting 5 cc. of the
laboratory solution of each with 50 cc. of distilled water, and also a dilute solution
of NaCl by dissolving between one and two grams of the salt in 50 cc. of distilled
water.
To 10 cc. portions of each of the three solutions add a few drops of litmus
solution. Repeat using phenolphthalein.
Taste each solution by dipping a glass rod into the liquid and touching it to
the tongue. (Caution. Do not taste any substance in the laboratory unless
directed to do so.)
Test a drop of each solution in a colorless gas flame by means of an iron (or
platinum) wire. A yellow flame proves the presence of sodium.
Evaporate to dryness in a casserole 1 cc. of HC1 solution, and of NaCl solution.
In each case examine if there is a residue. Question. What conclusion can you
draw in regard to the volatility of water and hydrogen chloride as compared to
sodium chloride? What result would you predict if a solution containing both
NaCl and HC1 were evaporated? Pure sodium hydroxide is a stable non-volatile
substance which would be left as a solid when a solution containing it is
evaporated to dryness. (Caution. Do not evaporate alkaline solutions to dryness.
Glass and porcelain are slowly attacked by hot concentrated alkaline and a porcelain
dish is spoiled if an alkali residue is heated strongly in it.)
Summarize in a table the properties of the three solutions examined above.
3. Experiment. Take 40 cc. of your laboratory NaOH solution, place it in a
clean half liter flask and dilute with 440 cc. of distilled water. Shake the flask in
order that the solution shall be uniform throughout. Cork the flask and label it
"NaOH solution for Assignments 5, 6 and 7." Question. What approximately
is the ratio of the initial volume of the sodium hydroxide solution to the final
volume?
4. Dry a porcelain dish and watch-glass large enough to cover it. Weigh the
dish and watch-glass to 10 mg. Measure out in your graduated cylinder as
accurately as possible 50 cc. of the sodium hydroxide solution prepared above.
Pour the solution into the weighed evaporating dish, and sufficient phenolph-
thalein to give a pink color and then add small portions of your laboratory
hydrochloric acid, stirring after each addition, until the solution is colorless.
Approximately 5 cc. of the acid will be required to give the colorless solution.
If the addition of the last portion of acid is made, a few drops at a time, it will
f 151
be observed that the color changes abruptly. An excess of 1 cc. of HC1 may
now be added.
5. Place the dish containing the solution on your wire gauze and heat until
the solution begins to boil. Reduce the size of the flame and allow the solution
to boil gently, or to evaporate slowly just below the boiling point, until the
bottom of the dish is covered with solid material. Then, to avoid loss from
bumping and spattering during the evaporation to dryness, cover the dish loosely
with the watch-glass, leaving an open space at one side, and continue the heating,
first with a small flame, then more strongly until no further trace of water-vapor
is expelled. Allow the dish and residue to cool for 10 minutes while covered
with the watch-glass and again weigh to 10 mg. Heat the dish and residue gently
for five minutes longer, let cool, and weigh again. If the two weights are not the
same within 20 mg. repeat this process until two weights are obtained which
check to 20 mg.
6. Dissolve a portion of the residue in a small amount of water and test as
in Paragraph 2. State what evidence you have that a reaction has taken place
between the acid and base, and that the solid residue obtained is sodium chloride.
Write the equation for the neutralisation reaction between NaOH and HC1 and
interpret it in terms of (a) molecules, (b) mols, (c) grams of the substances
involved.
7. Calculations. From the weight of NaCl found in Paragraph 5, calculate
(a) the weight of NaOH which must have been present in the 50 cc. of NaOH
solution, (b) the number of grams of NaOH in 1 cc. of the solution, (c) the
concentration of the NaOH in grams per liter? As we shall see in the next
assignment this value may not be very accurate. In addition to errors such as
that made in measuring out the volume of the NaOH solution, the laboratory
NaOH contains small amounts of impurities, as NaCl.
8. The reaction just considered is typical of the reaction between any acid and
any base. In every case H of the acid unites with OH of the base to form water.
Write the equations for the reactions between the following bases and acids, and
interpret each equation in terms of mols of the substances involved :
Sodium hydroxide and nitric acid
Sodium hydroxide and sulfuric acid (to form two molecules of water)
Barium hydroxide and hydrochloric acid
Barium hydroxide and sulfuric acid.
9. Since the mol is a convenient unit of weight to use in studying chemical
reactions, concentration is often expressed in terms of the number of mols of
substances in a liter of solution. {Definition. A solution which contains in one
liter one mol of dissolved substance is called a molal solution; one which
contains in one liter one-tenth mol of dissolved substance is called a tenth molal
solution, etc. Molal hydrochloric acid is designated thus: M HC1; tenth molal
sulfuric acid would be written 0.1 M H 2 SO 4 , etc. Questions. Calculate the
number of mols of NaCl obtained in Paragraph 5. How many mols of NaOH
were there in 50 cc. of solution ? What then is the concentration of your NaOH
in mols per liter? If the laboratory solution were exactly 6 M (which probably
is not the case) what would this concentration be, as a result of the twelve fold
dilution in Paragraph 3?
10. Problems, (i) How many (a) mols, (b) grams of sulfuric acid are in
50 cc. of 0.2 M H 2 SO 4 ?
(2) What is the concentration in mols per liter of a solution which contains
5.8 g of NaCl in 125 cc. of solution.
16
ASSIGNMENT 6
TITRATION OF SOLUTIONS OF ACIDS AND BASES : AN ILLUSTRATION
OF VOLUMETIC ANALYSIS
Reference. Hildebrand, Chapter VIII, pages 106-107.
1. In Assignment 5 we learned that an acid and a base in solution will neutralize
each other. In Assignment 6 we shall see how this reaction can be used in
determining the concentration of one of these solutions when that of the other is
known. The operation is called a titration. It is evident that a pure solution
of a salt may be prepared by mixing the corresponding acid and base in exactly
the right proportion; this end-point in the titration is determined by means of a
suitable indicator .The relative concentration of the two solutions can be calcu-
lated when the relative volumes of the two solutions used in the titration are
known. Questions. How many mols of NaOH are required to neutralize
exactly 0.01 mol of (a) HC1, (b) H,SO 4 ? How many cc. of 0.50 M NaOH
solution will exactly neutralize 0.01 mol of (a) HC1, (b) H 2 SO,.?
2. Experiment. Prepare 300 cc. approximately 0.5 M HC1 from your labora-
tory 6 M solution, place it in a flask and shake it. Cork the flask and label it
"HC1 solution for Assignments 6 and 7." Clean a small flask and label it
"known H 2 SO 4 solution/' rinse the flask with distilled water and set it aside to
drain in order to have it ready for use later in this Assignment.
3. Experiment on the determination of an end-point and the choice of the
indicator to be used in the titration. Dissolve approximately 0.5 g. NaCl in about
50 cc. water, add 2 drops phenolphthalein and stir the solution. Add NaOH
solution (approximately 0.5 M) drop by drop, stirring after each drop is added,
and note how many drops are needed to give a distinct color. Then determine
how many drops of your 0.5 M HC1 solution are required to decolorize the
solution. Note. Small drops are conveniently added from a glass tube drawn
out to a point, the solution being held in the tube by placing the finger over the
upper end of the tube. Be sure that the tube is clean ; before using, rinse it once
with the solution. Repeat the experiment with litmus solution instead of phenol-
phthalein, and determine whether the change in color gives a satisfactory end-
point for the titration of HC1 and NaOH solutions. Finally determine the
nature of the end-point with each indicator when about 1 g. sodium sulfate,
Na 2 SO 4 - 10 H 2 O is used instead of NaCl.
4. The volumes of the solutions used in a titration are measured by means of
burettes; these are uniform glass tubes graduated in cubic centimeters and tenths
(or fifths) of cubic centimeters. Two burettes will be placed on your desk before
the beginning of the Assignment. At the end of the period rinse them with
distilled water and leave them on your desk.
5. Experiment. Fill each burette with distilled water. Air may be removed
from the small tube below the pinch-cock by tilting the tip upward and allowing
the liquid to flow through the pinch-cock. Question. Why is it necessary to
remove any bubble of air trapped in this small tube? Practice reading a burette:
bring your eye to the same level as the liquid and note the reading of the burette
corresponding to the bottom of the meniscus; repeat until consecutive readings
check to better than 0.05 cc. Question. Why is it important to have the eye at
the same level as the liquid before making a reading? Never attempt to adjust
the volume of the solution in a burette so that the reading will be some exact
amount. Allow the water to flow slowly out of the burette. If drops remain on
the inner surface of a burette, exchange it at the office for a clean one.
6. Rinse one burette with a little of your approximately 0.5 molal HC1 solution,
and fill the burette with this solution. Rinse and fill the other burette with the
0.5 molal NaOH solution.
[ 171
7. Record the readings of the burettes side by side in your note-book ; run
about 15 cc. of the acid solution into a clean beaker or flask standing on white
paper, and record the final burette reading under the initial reading. Add two
drops of phenolphthalein, and about 20 cc. distilled water. Then run in the
sodium hydroxide solution from the other burette, a little at a time, and towards
the end, very carefully, a drop or two at a time, stirring the mixture constantly,
until the faintest perceptible permanent pink color is obtained. Wash down the
inside of the beaker by means of a jet of water from the wash bottle. If two
much of the basic solution is added, decolorize the solution by adding a little of
the acid and determine the end-point again. Record the final readings of each
burette, and the actual volumes of each solution used in titration. Calculate the
volume of sodium hydroxide necessary to neutralize one cubic centimeter of the
hydrochloric acid.
8. Repeat this experiment, using about 20 cc. of the acid solution, and in each
case make the same calculations. Do not fill up the burette each time unless there
is not enough solution in it for the titration.
9. Questions. If the error in measuring out a volume of solution by means of a
burette is 0.10 cc. what is the percentage error if 1 cc. of solution is measured?
If 20 cc. of solution are measured? Why is it important not to use less than
10 cc. in any titration?
10. Compare the volume ratios calculated from two titrations. If the result
differs from the average by more than 1%, perform additional titrations until
you are satisfied that you have determined the volume ratio with an accuracy
better than 1%.
11. Questions, (a) From an examination of the equation for the neutraliza-
tion of sodium hydroxide by hydrochloric acid state the ratio of the number of
mols of acid and base added to the beaker when exact neutrality was reached.
(b) From your average volume ratio state which solution, acid or base, is the
more concentrated, and what is the ratio of the concentrations.
12. Take your clean, dry, labelled flask, Paragraph 2, to the office to obtain a
sulfuric acid solution of known concentration.
13. Empty the HC1 out of the burette, rinse it with about 5 cc. of the sulfuric
acid solution of known concentration, and fill it with the sulfuric acid
solution. Determine the volume ratio as before from three (or more) titrations.
From the volume ratio and the reaction between sodium hydroxide and sulfuric
acid calculate the concentration of the sodium hydroxide solution in mols per
liter. Calculate also the concentration of your hydrochloric acid solution.
14. Make a list of the sources of error. (Many of them have been mentioned
in the above directions.)
15. Save the remainder of the NaOH and HC1 solutions, whose concentration
you have determined, in corked flasks for use in Assignment 7.
16. Problems, (i) How many cc. of 0.01 M Ba(OH) 2 will be required to
neutralize 10 cc. of 0.5 M HC1?
(2) What is the concentration in mols per liter of a sulfuric acid solution,
25 cc. of which neutralizes 20 cc. of 0.20 M NaOH?
[18
ASSIGNMENT 7
VOLUMETIC ANALYSIS, CONTINUED: THE DETERMINATION OF THE
EQUIVALENT WEIGHT OF AN UNKNOWN ACID
Reference. Hildebrand, Chapter VIII, pages 108-111.
1. In Assignment 7 there will be introduced another unit of quantity, the
equivalent. This unit and the corresponding unit of concentration, equivalents
per liter, are frequently more convenient than the units, mol and mols per liter
The mol and equivalent are identical for HC1, HNO 3 , NaOH, NaCl, etc., but a
mol of H 2 SO 4 , Ba(OH) 2 or Na 2 SO 4 , etc., contains two equivalents. A solution
which contains one equivalent in a liter is called a normal solution and is desig-
nated i N. The convenience of this unit of concentration depends upon the fact
that when one equivalent of any acid reacts with one equivalent of any base
the resulting solution contains one equivalent of the corresponding salt.
Question. What is the normal concentration of a molal solution of H 2 SO 4 ,
NaOH and Na 2 SO 4 respectively? How many equivalents of acid can be
neutralized by 10 cc. of 0.1 N NaOH? Calculate the normal concentrations of
your solutions of NaOH, HC1 and H 2 SO 4 used in Assignment 6.
2. Weighed .portions of an unknown solid acid will be titrated with your NaOH
solution. By means of the concentration of the NaOH solution (determined in
Assignment 6) the number of grams in one equivalent of the acid will be calcu-
lated. The correctness of this result depends of course, on the accuracy of your
work in Assignment 6. It is to be noted that if you had started with a solid and
of known equivalent weight, you could have used the results obtained in this
Assignment to determine the concentration of the sodium hydroxide solution.
3. Experiment. Obtain from the office a sample bottle containing approxi-
mately 2 grams of a crystalline acid of unknown composition. Clean two % liter
flasks and label them No. 1 and No. 2. Weigh the sample bottle with its cork and
contents. Remove the cork, taking care that none of the solid which may be
sticking to it drops off, and shake about one gram into flask No. 1. Replace the
cork and weigh again. Shake the remainder of the sample into flask No. 2, and
weigh the empty sample bottle and -cork. All weighings should be made to 5 mg.
4. Dissolve the contents of each flask in about 50 cc. of distilled water, add
two drops of phenolphthalein and titrate to the appearance of the first pink color
with the sodium hydroxide solution prepared in Assignment 8. If you should
pass the end-point in a titration add from a second burette an acid solution of
known concentration until the pink color is discharged, and again titrate to the
end-point with the sodium hydroxide.
5. For each sample calculate the number of cubic centimeters of NaOH
solution needed to neutralize 1 gram of the acid. If the results differ by more
than 2% repeat the experiment.
6. From your two (or more) measurements obtain an average value of the
number of cubic centimeters of sodium hydroxide solution per gram of acid.
From this average value and the concentration of the NaOH solution calculate
the number of mols of NaOH needed to neutralize 1 gram of the acid.
7. Questions. How many equivalents are in one gram of the acid? How
many grams are in one equivalent of the acid?
Report this value to your instructor at once. If your result is unsatisfactory
check your calculations in Assignments 6 and 7 and repeat as much of the work
as is necessary.
8. Problems, (i) Among the acids suitable for this Assignment are: oxalic
H 2 C 2 O 4 2H 2 O ; citric, H 3 C 6 H 5 O 7 H 2 O ; tartaric, HX 4 H 4 O e and potassium acid
sulfate, KHSO 4 . Write the equations for the reaction between NaOH and (a)
[19]
oxalic acid to form Na 2 C 2 O 4 , and (b) KHSO 4 ; and calculate the equivalent
weight of the acid in each of these reactions.
(2) What is the normal concentration of a sulfuric acid solution, 25 cc. of
which neutralizes 20 cc. of 0.20 AT NaOH?
(j) Chemically pure ("C. P.") sulfuric acid, nitric acid, hydrochloric acid
and ammonia, as supplied by the manufacturers, are concentrated aqueous
solutions of these substances. The concentration of each solution is guaranteed
not to be less than a certain minimum value, and this is tested by measuring the
density (or the specific gravity). The following table contains the density and
the percentage composition by weight of the concentrated laboratory reagents.
Density Concentration,
g. per cc. % by weight Equivalents per liter
H 2 SO 4 1.84 95.6% H 2 SO 4
HNO 3 1.42 69.8% HNO 3
HC1 ' 1.19 37.2% HC1
NH 4 OH 0.90 57.5% NH 4 OH
For each solution calculate the normal concentration and record the results in
the fourth column of the table.
Caution! The concentrated acids, especially sulfuric and nitric, produce
dangerous burns and should not be used carelessly. When you are directed to
use one of these acids carefully pour just enough of it for your experiment into
a dry beaker.
20
SECTION II
IONIC THEORY
RAPID REVERSIBLE REACTIONS AND EQUILIBRIUM
ASSIGNMENT 21
SOLUTIONS OF STRONG ELECTROLYTES. IONIC EQUATIONS
Reference. Hildebrand, Chapter X, pages 124-137.
1. This Assignment, which contains no experimental work, is introduced in
order that the student may become familiar with the fundamental ideas under-
lying the Ionic Theory before proceeding to use these ideas in the following
Assignments.
2. The use of the terms acids and bases in designating two distinct groups of
substances implies that the members of each group have a set of properties in
common. The properties characteristic of all acid solutions, which were observed
in Assignment 5, are ascribed to a substance called hydrogen ion, represented
by the symbol H + ; and those of basic solutions to the substance hydroxide ion,
written OH~. In addition to the properties of the hydrogen ion, each acid in
solution has a group of properties different from those of any other acid but
common to solutions of all salts of that acid. Thus, hydrochloric acid has a set
of properties which is characteristic of solutions of all chlorides and is ascribed
to a substance called chloride ion, Cl~. Likewise, a solution of sodium hydroxide
has a group of properties which is characteristic of solutions of all sodium salts
and which is attributed to the sodium ion, Na + . Questions. What is the evidence
from freezing point data that there are approximately 2 mols of substance pres-
ent when one mol of hydrogen chloride is dissolved in water? How does the
electrical conductivity of hydrochloric acid solution support the idea that the
molecule of hydrogen chloride in solution is broken up into two new substances?
In what way does a chloride ion differ from an atom of chlorine, and a hydrogen
ion differ from an atom of hydrogen? List briefly differences in properties of
the substances hydrogen ion and hydrogen gas.
3. Many substances in dilute solution may be considered as completely ionized.
These substances are called strong electrolytes and include :
Practically all salts
A few acids as HC1, HNO 3 and H 2 SO 4 , and
A few bases as NaOH, KOH and Ba(OH) 2 .
The student should memorize this list and should form the habit of thinking
of solutions of strong electrolytes in terms of the ions present and not merely in
terms of the specific solid, liquid or gas used in making the solution. It is impor-
tant to realize, however that strong electrolytes are not ionized in the solid or
gaseous state ; thus, while hydrochloric acid and sodium chloride solutions consist
of H + and Cl~ and Na + and Cl~, respectively, gaseous HC1 and solid NaCl are
not ionized, and each has its own specific properties. Questions. What are the
principal substances present in each of the following solutions and what is the
approximate concentration of each substance in mols per liter :
(/) A solution which contains 0.1 mol of H 2 SO 4 in 1 liter. (Answer. H + and
SO 4 ~" at concentrations 0.2 M and 0.1 M, respectively).
(-2) A solution which contains 0.2 mol of NaOH in 1 liter.
(5) A solution which contains 0.1 mol of Na SO 4 and 0.1 mol of NaCl in
1 liter.
[21]
(4) A solution which is made by mixing equal volumes of (i) and (2).
4. Ionic equations. Having realized what substances are present in solutions
of strong electrolytes we are now in a position to interpret reactions involving
such solutions. We shall first consider what is the ionic reaction when a strong
acid neutralizes a strong base, and shall take as an example the reaction between
sulfuric acid and sodium hydroxide solutions. The equation
H,SO 4 + 2NaOH = 2H 2 O + Na 2 SO 4
is a statement of the reaction that takes place whenever H 2 SO 4 and NaOH
solutions are mixed, and we have seen that it may be interpreted in terms of
molecules, and of mols, equivalents, grams, or any other weight units. It records
no. experimental details, such as volumes or concentrations of solutions, or the
use of either reagent in excess; and additional notes are necessary when such a
record is desired.
This equation can be interpreted according to the ionic theory (a) by adding
a note that H 2 SO 4 , NaOH and Na 2 SO 4 are strong electrolytes and that H 2 O is a
weak electrolyte, or (b) by rewriting the equation to give the same information:
(2H + + SO 4 ~) + (2Na + + 2OH~) = 2H 2 O + 2Na + + SO 4 ~.
It is evident that the Na + and SO 4 ~~ shown to be in the final solution were present
in the two initial solutions, and that these substances have undergone no change
during the reaction. It is incorrect to say that "sodium ion and sulfate ion have
combined/' The reaction that has taken place is simply the formation of the weak
electrolyte water:
5. The statement that a substance is a weak electrolyte is equivalent to saying
that the ions of that substance cannot exist in the presence of each other except
at low concentrations and it is obvious that any reaction which involves the
formation of a weak electrolyte from its ions may be expected to take place.
Accordingly, we shall be able to predict a number of reactions when we have
classified the substances involved as strong or weak electrolytes in solution.
6. Another reaction which involves the ions of a strong electrolyte is the
precipitation of a sparingly soluble salt. Again the statement that a salt, such as
silver chloride, AgCl ; is sparingly soluble is equivalent to saying that its ions
cannot exist together in solution except at the low concentrations corresponding
to the solubility of the salt. Write the ionic equation for the reaction that takes
place when dilute solutions containing equivalent amounts of NaCl and AgNO 3
are mixed. What substances are present at high concentrations in each of the
initial solutions and in the final solution? Write the ionic equation for the
reaction. Consider next the case in which silver nitrate solution in excess is
added to a sodium chloride solution, answer the same question and write the
ionic equation. It should be obvious that the reaction is the same in both cases,
and that, therefore, the two equations should be identical.
7. Does anything happen in the following experiments ? A dilute solution of
sodium chloride is mixed with a dilute solution of (i) potassium nitrate,
(2) nitric acid?
8. Problems, (i) Write ionic equations for the following reactions: (a) A
precipitate of barium sulfate, BaSO 4 , is formed by mixing solutions of barium
chloride and sulfuric acid, (b) A sodium sulfate solution is evaporated to dry-
ness. (c) Hydrogen chloride gas is dissolved in water.
(2) The solubility of lead iodide, PbI 2 , is 0.002 mol per liter at 18. What
is the concentration of the ions present in a saturated solution?
(j) If the freezing point of water is lowered approximately 1.86 per mol of
substance in solution in 1000 grams of water, what is the freezing point of a 0.1
molal solution of (a-) a non-electrolyte, (b) hydrochloric acid, (c) barium
chloride, BaQ 2 ?
[221
ASSIGNMENT 22
STRONG AND WEAK ACIDS. THE USE OF INDICATORS TO MEASURE
HYDROGEN ION CONCENTRATION
Reference. Hildebrand, Chapter X, pages 138-142.
1. Indicators can be used to measure the concentration of hydrogen ion, or of
hydroxide ion, in a solution. The color change for each indicator occurs in a
definite range of concentrations of hydrogen ion (or hydroxide ion), which is
characteristic of the indicator. In the present Assignment, by using solutions of
known concentration of the strong acids listed in Assignment 21, we shall develop
a method of determining approximately the concentration of hydrogen ion in
any acid solution. This indicator method will then be used to measure the concen-
tration of hydrogen ion in solutions of a typical weak acid. Two reactions
involving this acid will be studied.
2. Experiment. Prepare 60 cc. N HC1 by adding distilled water to the proper
volume of the 6 N laboratory reagent and shaking the mixture. From this
solution, or from your known HC1 solution, Assignment 6, prepare between 50
and 100 cc. of 0.10 N HC1, and from this solution prepare 50 to 100 cc. of 0.001
N HC1. On account of the error in measuring volumes by means of a graduated
cylinder, and the probable variation of the concentration of the laboratory
solution from 6 N, the concentrations of these solutions are known only
approximately.
3. Pour into marked test tubes 10 cc. of each HC1 solution (N, 0.1 N, 0.01 N,
0.001 N) and pour 10 cc. of water into a fifth test-tube. Add to each solution
from a glass tube a single drop of methyl violet solution. Hold the tubes in a verti-
cal position over a piece of filter paper, look down through the surface, record the
color of each solution and note the smallest concentration of hydrochloric acid
that shows with this indicator a color different from that of water. If the
indicator solution is so dilute that one drop does not produce a distinct color add
1 or 2 more drops, but record the number of drops, and use the same number in
all the tubes. State how the indicator, methyl violet, may be used to determine
the approximate concentration of a hydrochloric acid solution. (The color in the
more concentrated solution will fade on standing. It may be restored by adding
another drop of the indicator.) Note. Set aside the remainder of the 0.01 N
and 0.001 TV" HC1 for use later in this Assignment, Paragraphs 5 and 8.
4. Experiment. Repeat the experiment with nitric acid or with sulfuric acid,
using the same concentrations as before (N, 0.10 N, 0.01 N, 0.001 N). Compare
the colors obtained with the different acids. Questions. Are the colors charac-
teristic for each acid? If not, what substance determines the color? If
hydrochloric acid is completely ionized what conclusion can you draw with respect
to the ionization of nitric acid and sulfuric acid? Note. Save the 0.001 N
solution for later use, Paragraph 5.
5. If you have performed the above experiments correctly and understood
them you will realize that the indicator methyl violet can be used to determine,
approximately, concentrations of hydrogen ion between N and 0.001 N.
However, while the concentration of H + in the dilute solution, 0.001 N, (which is
often written 10~ 3 N), is small compared with the normal solution, it is 10,000
times as great as in pure water. We shall now make use of the indicator
methyl orange to examine solutions in which the concentration of H + is between
10~ 3 N and that of pure water, 10~ 7 N. Experiment. By a 10 fold dilution of
your 10- 3 A 7 HC1 solution, prepare 50 or 100 cc. 10~ 4 N HC1; and from this
prepare a 10~ 5 N solution. Test 10 cc. portions of these three solutions and of
water with 1 drop of methyl orange. Repeat the experiment with nitric or sulfuric
[23]
acid, starting with your 10~ 3 A 7 " solution. Note. Save one of the lO' 1 N solutions
for later use, Paragraph 8.
6. Summarize your results with the two indicators in a table which shows the
color obtained at various concentrations of H + . Show your table to the
instructor.
7 '. The weak acid which w r e shall now proceed to study is acetic acid,
HC 2 H 3 O 2 . It is a white, crystalline, rather volatile solid which melts near room
temperature and is very soluble in water. The formula of its sodium salt, sodium
acetate, is NaC 2 H 3 O 2 . We shall abbreviate these formulas to HAc and NaAc,
respectively. Questions. How many mols of acetic acid are contained in 1 liter
of 6 N acid? How many mols of sodium acetate could be prepared from
this quantity of acetic acid?
8. Experiment. From the laboratory 6 N HAc prepare solutions which are
approximately N, 0.5 N, 0.05 N and 0.01 N. Place a 10 cc. portion of each of
the acetic acid solutions (N to .01 N) in a labelled test-tube, place 10 cc. distilled
water in another test-tube, and test each solution with methyl violet as in Para-
graph 3. For comparison, repeat the test with HC1 solutions of suitable
concentrations. (Note. In color comparisons it is not safe to trust the memory,
or even written descriptions.) Determine the lowest concentration of acetic acid
at which the color with methyl violet is distinctly different from that with water,
and estimate approximately the concentration of hydrogen ion in two of the acetic
acid solutions, say in the N and O.I N solutions. Repeat the experiment, but
use methyl orange instead of methyl violet, and estimate the concentration of H +
ion in the O.I N solution and in a more dilute solution.
9. The acetic acid in the solution must be present either in the form of ions,
H + and Ac~, or in the un-ionized form, HAc. From your estimate of the concen-
tration of H* in the 0.1 normal solution, calculate the fraction of the acetic acid
which is ionized, and the fraction which is un-ionized. The fraction of the acid
which is in the form of ions is called the degree of ionization. State also the
concentrations of acetate ion, Ac~, and of the un-ionized acid, HAC, in the 0.1
normal acetic acid solution. Is acetic acid a weak or a strong acid?
10. Calculations. The concentration of the ions in acetic acid solutions have
been determined by other methods more accurately than is possible by these color
experiments. The concentrations of hydrogen ion in these solutions at room
temperature are given in the following table:
Concentration Concentration Concentration Concentration of Degree of
acid. of H + . of Ac~. un-ionized HAc. Ionization.
1 N .004 N
0.5 AT .003 N
r\N .0013 N
O.OIN .0004 AT
Fill in the remaining columns of the table and show your table to the instructor
at once.
11. Memorize the fact that acetic acid and water are weak electrolytes.
Questions. Refer to Assignment 21, Paragraph 5, and state what you would
expect to happen if a solution containing Ac~ at high concentration were mixed
with a solution containing H + at high concentration. What solutions would you
mix to try this experiment, and how could you prove with an indicator that a
reaction had taken place?
12. The reaction between sodium acetate and hydrochloric acid solutions.
Experiment. Prepare some approximately half normal hydrochloric acid.
Measure out two 15 cc. portions in test-tubes, and add two or three drops of
methyl violet to each. Measure 10 cc. 1 N sodium acetate and add the solution,
[24]
a few drops at a time, to one of the 0.5 N hydrochloric acid solutions. After
each addition shake the mixture, record the color and note the volume of the
sodium acetate solution added. Write the equation for the reaction that has
taken place between H + and Ac~. Questions. If solutions containing 0.10 mol
HC1 and 0.05 mol NaAc were mixed what substances would be present in the
resulting solution? What would be the concentration of each if the final volume
were (a) 1 liter, (b) 500 cc.?
13. The neutralization of NaOH solution by acetic acid. The concentration
of H + in an acetic acid solution is small in comparison with the total concentration
of acid in the solution, but the student must not jump to the conclusion that the
results of Assignments 5 and 6 on neutralization would have been materially
different if acetic acid had been used throughout instead of hydrochloric acid.
In the following experiment we shall study qualitatively the reaction between
NaOH solution and acetic acid; cf. the quantitative experiment in Assignment 5,
Paragraphs 4 and 5. Experiment. To 10 cc. 6 TV NaOH in a porcelain dish
add about 8 cc. 6 A r HAc. (Does the mixture become warm?) Add 1 drop of
phenolphthalein and continue to add acetic acid slowly until the solution becomes
colorless; test again with the indicator (since the phenolphthalein color fades
in a concentrated NaOH solution) and finally add about 2 cc. acid in excess.
Evaporate the solution until crystals of salt begin to separate. (Caution. Do
not evaporate the solution to dryness, since the NaAc may decompose.) Allow
the mixture to cool, collect some of the moist salt, NaAc 3 H 2 O, and dry it
between filter papers. Prove that the salt contains sodium by the flame test, and
acetate by warming with 6 TV sulfuric acid and noting the odor of acetic acid.
Question. What evidence is furnished by this experiment that the neutralization
reaction
HAc + NaOH == H 2 O + NaAc
takes place when solutions of HAc and NaOH are mixed? The concentration
of an acetic acid solution can be determined by titration with a known sodium
hydroxide solution when a suitable indicator is used, phenolphthalein in this case
(and if time permits the student may perform this titration).
14. We shall now examine the above reaction in the same way as we have
already done in Assignment 21, Paragraph 4, in the case of the reaction between
a strong acid and a strong base. Questions. Which of the four substances are
strong and which weak electrolytes? What substance is present at high concen-
tration in an acetic acid solution which is not present at high concentration in a
sodium acetate solution? What substance is present in a NaOH solution which
is not present at high concentration in sodium acetate solution? What substance
is present at high concentration in a sodium acetate solution which is not present
at high concentration in either of the initial solutions? Write the equation for
the main reaction, and show it to your instructor. The study of this reaction will
be continued in Assignment 25. V <
15. Problems, (i) 20.0 cc. acetic acid solution were found to neutralize the
same volume of NaOH solution as 16.0 cc. 0.50 TV H 2 SO 4 ; phenolphthalein was
the indicator in both titrations. What is the concentration of the acetic acid
solution (a) in mols per liter, (b) in equivalents per liter, and (c) in grams
per liter?
(2) Outline experiments to distinguish between
(a) 1.0 N HNO 3 and 0.10 N HNO 3
(b) A solution of nitric acid and one of acetic acid which give the same
bluish color with methyl violet.
[25]
ASSIGNMENT 23
STRONG AND WEAK BASES. THE USE OF INDICATORS TO MEASURE
HYDROXIDE ION CONCENTRATION
1. In Assignment 23 we shall develop a method of measuring approximately, by
means of indicators, concentrations of hydroxide ion between normal and 10~ 7
normal, the concentration in pure water. This method will be used in studying
solutions of a typical weak base. Since there is throughout a close relation with
the preceding Assignment only brief direction will now be given. The student
is expected to make use of the discussion and Questions in Assignment 22 in
correlating the results of the two Assignments.
2. Experiment. Prepare solutions of sodium hydroxide which are approxi-
mately normal, 0.1 normal, 0.01 normal, and 0.001 normal. (State in your
note-book how you prepared these solutions.) To 10 cc. of each solution in a
test-tube, add 1 drop of a solution of the indicator, trinitrobenzol. Record the
color obtained in each case, and observe especially the most dilute solution that
gives a color with the indicator. Make a second series of observations using 1 a
larger amount of the indicator, say 6 drops, in each case. Note that by usirig
the different amounts of indicator you can determine approximately concentra-
tions of OH~ in one case between N and .01 N, and in the other between O.I N
and 0.001 N.
3. Repeat the experiment with potassium hydroxide solution, and compare the
colors obtained at each concentration. If sodium hydroxide in solution is
completely ionized, what conclusion can you draw with respect to potassium
hydroxide? What concentrations of hydroxide ion can be measured by means
of this indicator, trinitrobenzol?
4. Experiment. Prepare solutions of NaOH or KOH which are approximately
10- 4 N and 1O 5 N. Test 10 cc. portions of the 10~ 3 N, 10~ 4 N and 1Q- 5 N
solutions, and water with the indicator phenolphthalein. Repeat, using litmus
instead of phenolphthalein. Note. Since the quantity of alkali in a given volume
of these dilute solutions is extremely small it is evident that large errors in
concentration may result if the test-tubes and flasks are not thoroughly washed
with .distilled water before use. Check your results by preparing fresh portions
of the 10~ 3 , 10~ 4 and 10~ 5 N solutions and repeating the experiment.
5. Summarize your results in a table which shows the color obtained with
each indicator at various concentrations of hydroxide ion. Compare this table,
and the corresponding tph^e in Assignment 22, Paragraph 6, with the table given
by Hildebrand on page 181.
6. In the next experiment we shall determine the concentration of hydroxide
ion in solutions of the weak base, ammonium hydroxide, NH^OH. The concen-
trated laboratory reagent, cf. Assignment 7, Problem (j), is prepared by dissolving
the gas ammonia, NH 3 , in water until the solution is nearly saturated with NH 3
at room temperature. Write the equation for the formation of ammonium
hydroxide from ammonia and water. Ammonium sulfate, (NH 4 ).,SO4, is an
example of a salt of this base. Question. What is the concentration in mols per
liter and in equivalents per liter of ammonium ion, NH 4 + , and of sulfate ion in a
0.1 molal solution of ammonium sulfate?
7. Experiment. From the laboratory 6 TV NH 4 OH solution prepare solutions
which are approximately N, 0.1 N, 0.01 N and 0.001 N. By experiments with the
indicators trinitrobenzol and phenolphthalein (write out the details of these
experiments in your note book) estimate the concentration of OH~ in at least two
of these solutions. In each case give also the concentration of NH 4 + and of
un-ionized NH 4 OH, and calculate the degree of ionization. Note. Accurate
[261
determinations show that at each concentration the degree of ionization of
ammonium hydroxide is almost the same as that of acetic acid at the same
concentration; see Assignment 22, Paragraph 10.
8. The reaction between ammonium chloride and sodium hydroxide solutions.
Predict what will happen when a solution of ammonium chloride is added to a
solution of sodium hydroxide; cf. Assignment 22, Paragraphs 11 and 12. Plan
an experiment to demonstrate this result and try the experiment. Write the
ionic equation for the reaction.
9. The neutralization of H^SO^ solution by ammonium hydroxide . Experiment.
To 10 cc. 6 N H 2 SO 4 in a" flask add 6 N NH 4 OH from a graduate until the
solution, after shaking, has a distinct odor of NH 3 . Evaporate in a porcelain
dish until a considerable quantity of salt has separated, cool the mixture, collect
some of the salt and dry it between filter papers. Prove that the salt contains
(a) sulfate, by dissolving some of it in water and adding barium chloride (to
precipitate BaSO 4 ), and (b) the ammonium radical, by warming with sodium
hydroxide solution and noting the odor of NH 3 . Give the experimental evidence
and reasoning in favor of the conclusion that, when solutions of a strong acid
and ammonium hydroxide solutions are mixed, the main reaction is:
H + + NH 4 OH = H 2 O + NH 4 *
The study of this reaction will be continued in Assignment 25.
10. Problems. (/) State how you would determine whether an unknown
solution is more acidic or more basic than water. (2) Suggest experiments to
distinguish between (a) 0.1 N KOH and 0.01 N KOH. (b) A solution of a
strong base and one of a weak base which have the same hydroxide ion
concentration.
ASSIGNMENT 24
RAPID REVERSIBLE REACTIONS AND EQUILIBRIUM
References. Hildebrand, Chapter XII, pages 155-171; Chapter XI, pages 145
and 148.
1. While certain reactions proceed to completion, as the transformation of
metallic copper into cuprous sulfide studied in Assignment 1, many reactions do
not. For example, when solutions containing equivalent amounts of a strong acid
and sodium acetate are mixed, cf. Assignment 22, Paragraph 12, the resulting
solution still contains about 1% of the reacting substances H + and Ac~, i. e., the
reaction H + -f- Ac" = HAc, although it takes place very rapidly, stops when about
99% of the possible amount of HAc has been formed. In the final solution, the
concentration of the three substances involved in the reaction are the same as in
the corresponding acetic acid solution. Similarly, if we had started with pure
acetic acid, which is un-ionized in the solid, liquid or gaseous state, and dissolved
it in water, the reaction HAc = H + + Ac~ would have taken place rapidly, but
only until about 1% of the acetic acid had been ionized. Question. From your
results in Assignment 22, what would be the concentrations of H + , Ac~ and HAc
in a solution made (a) by dissolving 1 mol of HAc to give a liter of solution, and
(b) by dissolving 1 mol of HC1 and 1 mol of NaAc to give a liter of solution?
2. The reaction just considered is an example of a rapid, reversible reaction,
and the three substances involved in this reaction are in equilibrium with each
other in the final solution. In general whenever it has been shown experimentally
that a reaction can be made to take place in both directions, i. e., is reversible,
then it may be concluded that under suitable experimental conditions a state of
equilibrium can be realized in which all the substances involved in the reaction
[27]
are present together. For each set of experimental conditions there is a definite
state of equilibrium ; and when the experimental conditions are altered, e. g., by
changing the concentration of one or more of the substances involved or by
changing the temperature the reaction takes place in one direction or the other
until equilibrium is again established. The problem is to learn to predict what
will happen in any given case when the experimental conditions are altered.
3. The effect of changing the concentration of one of the substances involved
in an equilibrium. Experiment. Place two 10 cc. portions of N acetic acid in
two test-tubes, a 10 cc. portion of 0.1 N acetic acid in a third test-tube, and a
10 cc. portion of water in a fourth test-tube. To each solution add the same
number of drops of methyl violet solution. To one of the normal solutions add a
small amount of solid sodium acetate (or of 4 AT" solution) ; compare the colors
of the four solutions. Repeat the experiment with more dilute solutions of
acetic acid, using methyl orange instead of methyl violet. Question. What
conclusion can you draw with regard to the change in the concentration of
hydrogen ion in this experiment? What reaction must have taken place? When
equilibrium has again been established after the addition of the sodium acetate,
is the concentration of each of the substances H + , Ac" and HAc greater or less
than its concentration in the original acetic acid solution? State briefly how this
experiment illustrates the general statement : the effect of changing the concen-
tration of one of the substances involved in an equilibrium is to cause that reaction
to take place which tends to neutralize the change. The effect of increasing the
concentration of acetate ion could also have been predicted from the quantitative
statement of the Mars Law : (Concentration of H + ) (Concentration of Ac~) /^Con-
centration of HAc) = constant, when equilibrium has been established at a
definite temperature.
4. Outline an experiment to demonstrate that the reaction NH 4 + -|- OH~ =
NH 4 OH takes place when a solid ammonium salt is added to a solution of
ammonium hydroxide ; cf . your experiments in Assignment 23. Perform this
experiment, and explain how it illustrates the italicized statement in the preceding
Paragraph.
5. As another example of the effect of change of concentration upon equilibrium
we shall study the equilibrium between solid silver acetate and its ions. The
reversible reaction is:
AgAc ( solid )= Ag+ + Ac~
Note. The solubility of silver acetate is 0.06 mol per liter at room temperature;
while it is much more soluble than a sparingly soluble salt, as silver chloride,
it is much less soluble than such salts as silver nitrate, sodium nitrate or sodium
acetate. Experiment. Prepare some solid silver acetate by adding to 25 cc.
4 N NaAc solution (free from chloride) * about 20 cc. 0.1 N AgNCX solution,
and shaking the mixture several times. Collect the solid on a filter paper, and dry
it by pressing between dry filter papers. Prepare a saturated solution by shaking
the solid with 10 cc. water at intervals for about 10 minutes. (The saturated
solution can be prepared somewhat more conveniently by warming the mixture to
40 or 50, but it must be cooled to room temperature before continuing the
experiment). Allow the solid to settle. Pour half of the clear saturated
solution into another test-tube and set it aside for later use, Paragraph 7. To the
remaining mixture of solid and saturated solution add 6 N HNO 3 drop by drop,
shaking the mixture after each drop is added. When the silver acetate has
dissolved, heat the solution nearly to boiling and note the odor.
Note. Place all silver residues, including any solution which contain an appre-
ciable amount of silver, in the jar maked "silver waste."
* If chloride is present it may be removed from NaAc solution by adding first some
AgNOs solution, shaking the mixture and filtering off the precipitated AgCl. The resulting
solution contains some NaNOs and should be used only in this Assignment.
[28]
6. The disappearance of solid silver acetate on the addition of a strong acid
may be considered to be due to a shifting of the equilibrium
AgAc (solid) = Ag + + A<r
as a result of the establishment of the second equilibrium
H + + Ac- = HAc
When a drop of nitric acid is added the H + uses up Ac~ to form HAc; and
in order to again build up the concentration of Ac', and thus restore equilibrium,
some solid AgAc goes into solution. When more HNO 3 is added the same
processes are repeated, until finally all the silver acetate is dissolved. From this
discussion of the mechanism of the reaction it is evident that the result of the
experiment could have been predicted from the fact that acetic acid is a weak
acid. The principal substances that have disappeared are solid AgAc and H + ,
and those that have formed are Ag + and HAc; and we may, therefore, write
for the main reaction the equation :
AgAc (solid) + H + == HAc + Ac-.
The same equation is obtained by adding together the two equations considered
' in discussing the mechanism of the reaction. It is important, however, to realize
that in such a case as this no single equation can represent all that has happened
in the reaction. Thus, the concentration of Ac~, although small throughout the
experiment, is much smaller in the final solution than in the initial AgAc solution ;
this small decrease is due to the reaction H+ + Ac~ = HAc, but is not taken into
consideration in the equation which we have written for the main reaction.
7. Predict what will happen when solid sodium acetate is added to a saturated
solution of silver acetate. Experiment. Test your answer by adding a small
amount of 4 AT NaAc to the clear saturated solution of silver acetate which you
have just prepared. Question. How does this experiment illustrate the italicized
statement in Paragraph 3? Re-word that statement to make it apply specifically
to the change of solubility of a salt when the concentration of one of the ions of the
salt is changed. What will happen when solid silver nitrate is added to a saturated
solution of silver acetate?
8. Problems. (/) For each of the following reactions:
(a) H 2 O (liquid) =H 2 O (gas)
(b) H 2 O (solid) =H 2 O (liquid)
(c) NH 3 (in solution) = NH 3 (gas)
(d) Ag + + Cl- AgCl (solid). (Note the solubility of AgCl is 1.0 X 10' 5
mols per liter) ;
outline an experiment to show that the reaction can take place as written and a
second experiment to show that the reverse reaction can also take place.
(2) In each case in Problem (/) point out the conditions under which equi-
librium can be realized. A satisfactory answer in (a) is the following: whenever
water and water vapor are present together in a closed vessel equilibrium is
quickly established, and remains unchanged as long as the temperature remains
constant; when the closed vessel contains only water and water vapor the
pressure at equilibrium is 1 atmosphere when the temperature is 100 ; cf .
Assignment 2, Paragraph 10, for the pressure of water vapor at room temperature.
(j) Give an example of a reaction which can be shown to be reversible at high
temperature, but which proceeds so slowly at room temperature that equilibrium
will not be realized in several months. Note that all the substances involved in
this reaction can exist together at room temperature but that such a fact does not
constitute a proof that equilibrium has been established.
(4) From your table in Assignment 22, Paragraph 10, state what fraction of
[29]
the acetic acid is ionized in the 0.1 N and 0.01 N solutions. What reaction takes
place when 1 liter of 0.1 N solution is diluted to 10 liters of 0.01 N solution?
(5) State how the experimental result in Problem (4) could have been pre-
dicted from kinetic considerations, on the assumption that the two opposing
reactions, HAc = H + -j- Ac~ and H + -f- Ac~ = HAc, are actually taking place at
equilibrium.
(6) When one of the substances involved in an equilibrium is a solid, what
is the effect of adding more of this solid to an equilibrium mixture? Does the
solubility of a salt at room temperature depend on the amount of the salt in
contact with the saturated solution? Give another example of a reversible
reaction involving a solid substance.
ASSIGNMENT 25
THE REVERSIBILITY OF NEUTRALIZATION REACTIONS. HYDROLYSIS
References. Hildebrand, Chapter XIII, pages 185-192; as a review, read pages
173 and 180-183.
1. In earlier Assignments we have studied various examples of neutralization
reactions and have shown that the main reactions that take place are represented
by the following equations :
(a) H + + OH" = H 2 O when a strong acid reacts with a strong base.
(b) HA + OH~ = H 2 O + A~ when a weak acid reacts with a strong base.
(c) H + + BOH = H 2 O + B + when a strong acid reacts with a weak base.
In the present Assignment we shall find that these reactions are reversible and
we shall study the equilibrium that is established in each case. When the
Assignment is finished the student should be able, from a knowledge of the
relative strengths of various acids, and of various bases, to predict the approxi-
mate concentrations of the various substances at equilibrium in any given
neutralization reaction.
2. We wish to examine experimentally the properties of the solution that is
obtained when equivalent amounts of a given acid and base react. To do this we
shall take advantage of the fact that a solution of a pure salt is identical with the
solution prepared by mixing exactly equivalent amounts of the corresponding
acid and base. In other words we shall make use of the reaction which is the
reverse of neutralization in establishing the equilibrium state that we desire to
study. The experimental work will consist in determining by means of indicators
the approximate concentration of H + or OH~ in solutions of various salts.
Note. In the following experiments it is important that all test-tubes should be
clean; rinse each several times with distilled water before using it.
3. Experiment. Test 10 cc. portions of N NaCl solution and distilled water by
adding to each the same number of drops of litmus solution. Repeat using
phenolphthalein and methyl orange. Question. What conclusions can you draw
with regard to the relative concentration of H + or OH~ in sodium chloride solution
and in water? Write the ionic equation for the neutralization of NaOH by HC1
solutions, and state whether the Na + and Cl" are involved in the reaction. In
pure water concentrations of H + and OH~ are the same and equal to 10~ 7
mols per liter, what are the concentrations of Na + , Cl~, H + and OH~ in the
solution made (a) by dissolving 1 mol NaCl in water to give a liter of solution,
and (b) by dissolving exactly 1 mol of HC1 and 1 mol of NaOH to give a liter
of solution?
4. Experiment. Test 10 cc. portions of 4 N NaAc and distilled water first
[301
with litmus solution and then with phenolphthalein. Estimate approximately the
concentration of OH" in the 4 N NaAc solution. Questions. Is the concentration
of OH" in 4 N NaAc solution greater or less than in pure water? From the fact
that the solution must be identical with that obtained by mixing exactly equivalent
amounts of acetic acid and sodium hydroxide, what conclusions can you draw
with regard to the completeness of the neutralization reaction between HAc and
OH"? To account for the formation of OH" when NaAc is dissolved in water,
what reaction must take place? What are the approximate concentrations of
Na + , Ac", HAc and OH- in the 4 N NaAc solution?
5. It will be observed that the main reaction which accounts for the neutrali-
zation of HAc by NaOH,
HAc + OH- + (Na + ) = Ac- + H 2 O + (Na + )
or the reverse of this reaction considered in the preceding paragraph, involves
the disappearance of one weak electrolyte and the formation of another. Both
the OH" and the Ac" are competing for the H + . From the fact that water is a
far weaker electrolyte than acetic acid, it could have been predicted that when
equilibrium is reached the OH" will have taken nearly all the H + away from
the Ac" ; i. e., the concentration of the HAc and OH" will be very small and that
of the Ac" large. Question. From your table of the concentrations of H + in
acetic acid solutions, Assignment 22, what is the ratio of the concentration of H +
in N acetic acid to that in pure water?
6. As indicated in the preceding paragraph, the mechanism of this reaction
may be considered to be parallel to that of the dissolving of silver acetate in a
solution of a strong acid, Assignment 24. When an acetic acid solution is
neutralized by NaOH solution, the OH" unites with H + which is present at small
concentration in the acetic acid solution :
H + + OH- = H 2 O
and, in order to again restore equilibrium, the HAc dissociates,
HAc = H + + Ac-
These two reactions continue to take place simultaneously until equilibrium is
established. The sum of the two gives the equation which we wrote for the main
reaction.
7. Discuss in a similar way the mechanism of the reaction between Ac" and
H 2 O which takes place when sodium acetate is dissolved in water. This, of
course, involves the reverse of the two reactions written in Paragraph 6. Write
the equation for the main reaction in this case.
8. The reaction which is the reverse of neutralization is called hydrolysis.
In the reaction which we have just considered a small amount of the acetate ion
is said to be hydrolyzed. Question. Why would you expect the Ac" to be
hydrolyzed slightly while the Cl~ is not hydrolyzed at all?
9. The dissociation of HCXV, bicarbonate ion, into H + and carbonate ion,
CO 3 ", is far less than the dissociation of acetic acid. Questions. What reaction
would you expect to take place between CO 3 ~~ and H 2 O when Na 2 CO 3 is
dissolved in water? Which would you expect to be the more basic, a solution of
Na 2 CO 3 or a solution of NaAc of the same molal concentration? Test your
answer by the following experiment
10. Experiment. Estimate the approximate concentration of OH" in a 0.5 M
Na 2 CO 3 solution by testing 10 cc. portions with litmus, phenolphthalein and
trinitrobenzol. The main reaction for the hydrolysis of CO 3 ~ " is
C0 3 -- + H 2 = HCO 8 - + OH-
and the concentration of OH~ in a 0.5 M Na 2 CO 3 solution is about 0.01 M.
Questions. What would be the concentration of OH- if the CO 3 - " were completely
[31]
hydrolyzed? What fraction of the CO 3 ~~ has been hydrolyzed? What would be
the approximate concentrations in mols per liter of Na + , CO 3 ~~, HCO 3 ~ and OH"
in a solution prepared by dissolving 0.05 mol NaHCO 3 and 0.5 mol NaOH to
give 1 liter of solution?
11. Experiment. Test a 10 cc. portion of 4 AT NH 4 C1 with litmus solution.
Compare the color with that obtained by adding the same amount of litmus
solution to (a) 10 cc. water and (b) 10 cc. water containing 1 drop of 6 N HC1.
Questions. How do you account for the fact that a solution of NH 4 C1 is slightly
acid while that of NaAc is slightly basic. Write the main reaction for the
hydrolysis of the NH 4 + , and the two reactions which may be used to explain the
mechanism of the reaction. The concentration of H + in 4 N NH 4 C1 is about
4 X 10~ 5 N. What is the concentration of each of the principal substances
present in a solution made by dissolving 4 mols of HC1 and 4 mols of NH 4 OH
to form a liter of solution?
12. Problems. (1) What is the concentration of H + in N NaOH?
Note. Consult Hildebrand, page 180.
(2) From the results of the experiments in this Assignment and the table of
indicators (Hildebrand, page 181) select one or more indicators which might be
used to determine the end-point (i e., to determine when equivalent amounts of
acid and base are present) in each of the following titrations :
(a) NaOH solution with HC1 or HNO 3 solution.
(b) HAc solution with NaOH solution.
(c) NaHCO 3 solution with NaOH solution.
(d) NH 4 OH solution with HC1 solution.
Note. Compare Assignment 6, Paragraph 3.
(5) An ammonium acetate solution gives the same colors with indicators as
does water : What hydrolysis reactions must occur when solid NH 4 Ac is dissolved
in water? State what substances are present in the solution.
(4) Point out parallelisms between the strengths of acids and bases and the
hydrolysis of the corresponding ions.
32
SECTION III
REACTIONS OF IONS
ASSIGNMENT 31
THE PROPERTIES OF SODIUM, POTASSIUM AND AMMONIUM IONS.
TESTS FOR CHLORIDE, SULFATE AND NITRATE IONS
Reference, for Assignments 31 and all later Assignments: A Standard Text on
Inorganic Chemistry.
1. In this Assignment we shall study the properties of the ions of the common
acids and bases.
2. Aside from the physical properties, such as color and taste, the properties
of an ion depend on its behavior toward other substances. The study of the
chemistry of an ion thus consists in determining whether or not various substances
react with the ion, and in studying the reactions that do take place. For each
reaction it is necessary to know :
(1) the products of the reaction,
(2) whether the reaction takes place rapidly or slowly,
(j) whether the reverse reaction can be made to take place, and
(4) the nature of the equilibrium if the reaction is rapid and reversible.
In this course the task is greatly simplified on account of the following consid-
erations : Nearly all the reactions studied take place very rapidly ; this is
characteristic of ionic reactions, such as ionization, precipitation, neutralization,
etc. Also the number of different types of reactions is not large ; and, when for
one type a single example is understood, additional examples should present no
real difficulty. Finally an equilibrium state is reached in many cases, and a
knowledge of the principles of equilibrium enables the student to predict what
will happen when there is a given change in the experimental conditions. While
the number of facts to be remembered is thus greatly reduced, it is necssary,
however, to memorize such facts as; the formulas of substances, the relative
solubility of salts, the relative volatility of various solid or liquid substances and
the relative strength of weak electrolytes.
3. Properties of sodium ion. The chemistry of Na + is extremely simple and
may be summarized by saying that, at ordinary temperatures, with a few
uncommon exceptions, all of its compounds are non-volatile, readily soluble and
strong electrolytes.
4. Flame test for sodium. Experiment. Clean an iron wire by dipping it into
dilute HC1 solution in a test-tube and holding the wire in the flame until it gives
only a faint yellow color. (Note. Do not contaminate your HC1 solution hy
dipping the iron wire into the bottle.) Try the flame test with a solid sodium salt,
the laboratory solution of a sodium salt, a very dilute solution and distilled
water. The test is extremely delicate and substances which do not contain
measurable amounts of sodium usually give a slight yellow color for a short
time. Therefore, before concluding that sodium is present in an "unknown"
it is well to make comparative tests with known solutions (/) free from sodium
and (2) the same solution containing a small amount of sodium salt. Such
experiments are called blank tests.
5. Properties of potassium ion. What is the relation of K and Na in the
Periodic System? Potassium ion resembles .sodium ion very closely. One
difference, however, is the precipitation of potassium cobaltinitrite when cobalti-
nitrite ion, Co(NO 2 ) fi is added to a solution containing K + . Experiments. To
[ 33 ]
5 cc. of a dilute, neutral solution of a potassium salt add a few drops of acetic
acid and 2 to 5 cc. of the sodium cobaltinitrite reagent ; let the mixture stand
10 minutes.
6. Flame test for potassium. Try the flame test with solid KC1, KC1 solution,
KC1 solution containing a small amount of Nad, Nad solution, and NaCl solution
containing a small amount of KC1. Use the blue glasses to shut off the yellow
color of the sodium flame and practice until the presence or absence of both
potassium and sodium can be determined. Note. Since the flame test depends
upon the amount of solid present it is well to concentrate a dilute solution by
evaporation before making the test. Chlorides give more satisfactory flame tests
than sulfates or oxides, which is due to the relatively greater volatility of the
chlorides.
7. Properties of ammonium Ion. The ammonium ion resembles sodium ion and
potassium ion in that with few exceptions its compounds are soluble, strong
electrolytes. Like the potassium ion it forms a precipitate with Co(NO 2 ) 6 .
Repeat the experiment in Paragraph 4, using NH 4 + instead of K + . Ammonium
compounds give no flame test. The chemistry of NH 4 + is made somewhat more
complicated by the fact that NH 4 OH not only is a weak base, but is decomposed
easily into NH 3 and H 2 O. NH 3 , ammonia, enters into many reactions which
will be considered in later Assignments. Experiment. Note the odor of the
6 N NH 4 OH laboratory reagent; prepare some dilute solutions of NH 4 OH, e. g.,
0.1 A 7 " etc., heat gently about 10 cc. of each solution (moving the test-tube back
and forth through the flame) and note the odor from time to time; decide what
is the most dilute solution you can detect by means of the odor of NH 3 gas.
Determine by experiment whether you can detect the odor of NH 3 gas when a
solution of NH 4 C1 or (NH 4 ) 2 SO 4 is boiled. Suggest a method of testing for
ammonium ion based on the formation first of NH 4 OH in the solution and then
of NH 3 gas. Write equations. Experiment. Try your method with a dilute
solution of an ammonium salt.
8. Volatility of ammonium compounds. Ammonium compounds are much
more volatile than those of sodium and potassium. Experiment. Evaporate a
small amount of (a) NH 4 C1 solution and (b) a mixture of NH 4 C1 and KC1
solutions to dryness in a porcelain dish and heat until fumes are no longer given
off. Suggest a method other than the flame test of determining the presence of
K + in a solution which contains both K + and -NH 4 + .
9. Properties of H + and OH~. Aside from the tendency of H + and OH" to
unite with each other to form water each of these ions reacts with many other
ions to form weak electrolytes. We have already noted the reaction between H +
and Ac~ and OH~ and NH 4 + . Acids are more volatile than their salts : the relative
volatility of the common acids will be considered in Assignment 34. The OH~
also forms precipitates with many metallic ions. Question. How would you
test for (and estimate the concentration of) H + or OH" in an unknown solution?
10. Although the ions NO 8 ~, Cl~ and SO 4 " will not be considered in detail in
this Assignment, certain of their reactions will be studied in order that their
presence in solutions may be detected. All nitrates are soluble, but a few
chlorides and sulfates are sparingly soluble.
11. Test for chloride ion. Experiment. To 1 cc. portions of solutions of NaCl,
Ba(OH) 2 , Na 2 CO 3 and Na 2 SO 4 add 10 cc. water, shake and add a few drops
of AgNO 3 solution. In each case, record your observations, noting the color of
the precipitates. Now add HNO 3 until each solution is acid. From your results
state how you would test an unknown solution for chloride ion. Note. Ba(OH) 2
was used instead of NaOH because Cl~ is usually present in the laboratory NaOH
solutions. Experiment. Test a dilute solution of your NaOH for Cl~.
[34]
12. Test for Sulfate ion. Experiment. To dilute solutions of NaCl, NaOH,
NaXO 3 and Na 2 SO 4 add a few drops of Ba(NO 3 ) 2 solution. Now add HNO 3
until each solution is distinctly acid. From your results state how you would
test an unknown solution for SO 4 ~~.
13. Test for nitrate ion. Try the following experiment with a solution which
contains nitrate ion and with one free from nitrate ion, first acidifying with
6 N H 2 SO 4 if the solution is alkaline. To about 2 cc. of the solution to be
tested for nitrate ion add in excess a solution of ferrous sulfate, FeSO 4 (or of
ferrous ammonium sulfate, a double salt of ferrous sulfate and ammonium
sulfate), filter if there is a precipitate, hold the test tube in a slanting position,
and pour carefully down the side of the test-tube (from a small beaker) 2 or
3 cc. concentrated sulfuric acid. The concentrated acid sinks to the bottom
of the tube and a dark brown ring forms on its surface when nitrate is present.
Note. The brown substance decomposes rapidly when the mixture is hot.
If the result of the test is negative and the test-tube feels hot to the hand (owing
to the heat liberated when the concentrated H 2 SO 4 mixes with the solution),
repeat the test more carefully.
14. Make up solutions of various concentrations of Cl~ and try experiments
to determine if your test can be used to decide whether the concentration of
Cl~ is large or small. Repeat for the nitrate and sulfate tests.
15. Analyses Nos. A and B. Analyze the unknowns for the four positive and
four negative constituents considered in this Assignment. Try to distinguish
between large amounts, small amounts, and traces.
16. Problems. (/) What is the valence of each of the 4 positive and 4 negative
ions considered in this Assignment? Write the formulas of the sixteen possible
compounds, each of which contains one positive and one negative constituent
(cf. Hildebrand, pages 90-91). Name each of the sixteen substances, state
whether it is a solid, liquid or gas at room temperature, whether it is readily
soluble in water or not, and whether it is a strong or weak electrolyte.
(2) What is the valence of sodium in (a) solid sodium sulfate, (b) metallic
sodium?
(j) A solid unknown was dissolved in wa f er and the solution was found to give
distinct tests for Na + , K + , Cl~ and SO 4 ~ ~. What conclusions can you draw with
regard to the nature of the solid salts in the original mixture?
(4) Outline experiments to decide whether or not each of the following
contains the impurity named :
(a) Barium chloride impurity in the barium hydroxide solution.
(b) Potassium chloride in solid sodium chloride.
(c) Potassium chloride in solid ammonium chloride.
(d) Ammonium chloride in solid potassium chloride.
(e) Sodium nitrate in solid sodium sulfate.
(5) In each of the following it is assumed that two solutions are mixed each
of which contains one of the substances at moderate concentration, say 0.1 M .
Mark the cases in which no reaction takes place. In the others write the equation
for the reaction. In (a), (b), (c) state what solutions you would use in the
experiment.
(a) Ac-andCl- (d) NH 4 OH and K* (g) K + and Ac~
(b) NH t + and Na + (e) H + and Ac- (h) Ag + and Cl-
(c) NH 4 OHandH + (/) K + and OH- (i) Ba ++ and Cl~
F35 1
ASSIGNMENT 32
CALCIUM ION
References. Hildebrand, pages 175-179, 192-193 and 207.
1. In Assignment 32 we shall study the chemistry of metallic calcium and of
calcium ion, including its formation from metallic calcium. We shall find that
certain compounds of this element differ from the corresponding compounds of
sodium, potassium, and ammonium in that they dissolve to a much smaller extent
in water; and we shall study the equilibrium between these solids and their
saturated solutions.
2. Note. In future write equations for all reactions.
3. Obtain from the, office a piece of metallic calcium. Describe its properties as
far as you can observe them by physical examination. In what respects does
calcium show the physical properties of a metal?
4. Reaction between calcium and water. Experiment. Drop the calcium
into about 20 cc. of water. Stir the mixture and warm it gently until the metal
has disappeared. Test the solution for OH-. What reaction has occurred ? The
white solid formed, calcium hydroxide, Ca(OH) 2 , is a strong base but only
moderately soluble in water. Cork the test-tube and save the mixture for use
in a later experiment. Question. Name two other metals that react readily with
water. What reactions occur?
5. Calcium hydroxide. If solid calcium hydroxide is shaken with water until
a saturated solution is obtained an equilibrium between solid Ca(OH) 2 and the
ions Ca + * and OH~ is established. The equation for the reaction that has taken
place is
Ca (OH ) 2 (solid) = Ca ++ + 2OH-
The solubility of Ca(OH) 2 at room temperature is 0.02 mols per liter. Questions
What is the concentration of calcium ion, of hydroxide ion, (/) in mols per
liter, and (2) in equivalents per liter? Is the reaction just considered reversible?
What solutions would you mix in order to prepare solid Ca(OH) 2 ? Try the
experiment.
6. Which contains the larger concentration of OH~, a saturated solution of
Ca(OH) 2 or 1 AT NH 4 OH (see Assignment 23)? State what you think will
happen when a dilute solution of calcium chloride is made alkaline with ammonium
hydroxide. Try the experiment. Explain the result if your prediction was
incorrect.
7. State what would happen if a solution of calcium hydroxide were treated
with (i) hydrochloric acid, (2) ammonium chloride. Predict what would
happen if solid Ca(OH) 2 were also present in each case. Experiment. Test
your answers by experiments with portions of the mixture prepared in Par-
agraph 4. Questions. How would you use calcium hydroxide to test a solution for
ammonium ion?
8. Calcium carbonate. The solubility of CaCO 3 in water Lt room temperature
is very small, viz., 0.00013 mols per liter. Question. What is the concentration
of calcium ion, of carbonate ion, in this saturated solution, (i) in mols per
liter, (2) in equivalents per liter? What solutions would you mix in order to
form a precipitate of calcium carbonate? Try the experiment. Write the
simplest ionic equation for the reaction. Continue the experiment with very
dilute solutions of a calcium salt in order to determine if this reaction can be
used as a delicate test for calcium ion. In the more dilute solutions, if a pre-
cipitate does not form at once, heat to boiling and let stand 10 minutes. If the
liquid is turbid compared with water a precipitate has formed.
[361
9. Experiment. Dilute 5 cc. N CaCL with about 100 cc. water, add about 8cc.
N Na 2 CO 3 solution, heat the mixture to boiling, filter, wash the precipitate and
reject the wash water. Consider what substances may be present in the filtrate.
Test for chloride ion and for calcium ion. Which reagent was present in excess,
calcium chloride, or sodium carbonate?
10. Experiment. Treat a portion of the CaCO 3 precipitate with hydrochloric
acid solution. Repeat the experiment with nitric acid. The gas given
off is carbon dioxide, CO 2 . What substances are present in the final solution in
each case after the solution has been boiled to expel the CO 2 ? Note. CO 2 is
moderately soluble in cold water and the solution contains carbonic acid,
H 2 CO 3 , which is a weak acid. CO 2 gas will not be evolved until the concentra-
tion of carbonic acid in the solution becomes sufficiently high. Questions. How
is the equilibrium between solid CaCO 3 and its ions disturbed by the addition
of a strong acid? What is the analogy between this example and the action
of an acid on solid Ca(OH) 2 , and on solid silver acetate? What is the reason
for the difference between these results . and that obtained when dilute nitric
acid is added to silver chloride or barium sulfate?
11. Test for carbonate. Experiment. Fit a test-tube with stopper and delivery
tube. Place a small amount of any carbonate in the tube and add a few cc. of
6 N H 2 SO 4 . Insert the stopper and pass the gas given off into a clear solution of
Ca(OH) 2 .* If only a small amount of carbonate is present the CXX will not pass
over into the Ca(OH) 2 solution. A satisfactory test may be obtained in this case
by placing the test-tube in a beaker of boiling water, or by adding a small piece of
zinc to the cold mixture in the test-tube. The CO 2 is carried over with the
dissolved air and steam in the first method, and with the hydrogen gas in the
second.
12. Calcium oxalate. Experiment. Prepare some CaC 2 O 4 by treating a
solution containing calcium ion with ammonium oxalate solution. Test the
action of excess of a strong acid, as HC1 or HNO 3 on CaC 2 O 4 . Question. Is
oxalic acid, H 2 C 2 O 4 , a weak or a strong acid?
13. Experiment. Determine if the precipitation of CaC 2 O 4 is a delicate
test for calcium by treating very dilute solutions of Ca ++ with ammonium oxalate
solution and a little ammonium hydroxide (to make sure that the mixture is not
acid). Observe the same precautions as in the test for calcium by means of
CO 3 . Note. The solubility of calcium oxalate in water is even less than that
of calcium carbonate.
14. Calcium sulfate. Experiment. To 10 cc. normal CaCl 2 solution add
2 cc. 6 normal H 2 SO 4 . If no precipitate appears at once, heat the solution
gently, and let it stand. Filter. Test a portion of the filtrate for Ca + + by adding
NH 4 OH until the solution is no longer acid, and then (NH 4 ) 2 CO 3 solution, and
warming the mixture. Test another portion for SO 4 ~~. What conclusion do you
draw with regard to the solubility of CaSO 4 in water? Which is the less
soluble, (i) CaSO 4 or CaCO 3 , (2) CaSO 4 or BaSO 4 ?
15. Predict what will take place when solid CaSO 4 is heated with excess
Na 2 CO 3 solution? Experiment. Test your prediction by boiling some solid
CaSO 4 with normal Na 2 CO s solution. Test the filtrate for SO 4 ~. Wash the
precipitate with water, and test it for carbonate.
16. Experiment. Test whether the reverse reaction will take place, by heating
solid CaCO 3 with sodium sulfate solution, filtering and testing the precipitate and
filtrate.
^Instead of Ca(OH) 2 solution, Ba(OH)s. solution may be used; BaCO 3 , like CaCOs is
a difficultly soluble substance.
[37]
17. Flame test for calcium. Experiment. Determine whether calcium salts
give a characteristic flame test. Can you distinguish the flame with a calcium
salt from that with sodium or potassium salts?
18. Note. The solubilities of many common salts at 18 are listed on the inside
cover page of Alexander Smith's text-book. Do not try to remember the
actual solubilities, but make lists of the readily soluble and difficultly soluble
salts of each metal studied and memorize these lists. Arrange the compounds ot
calcium according to their solubilities in water, distinguishing readily soluble,
moderately soluble, and difficultly soluble substances. Point out the compounds
which are much more soluble in dilute hydrochloric or nitric acids than in water.
19. Give as many methods as you can of testing for Ca ++ in a solution
which is known to contain H + , Na + , K + and NH 4 + .
20. Problems, (i) What is the minimum volume of water needed at 18 to
dissolve i gram of calcium carbonate?
(2) Can the following substances be present at moderate concentrations in
the same solution? If not, what is formed?
(a) H + andNO 3 - (e) H + and C 2 O 4 - -
(0) H + and OH- (/) Ca ++ and CO 3 - -
(c) H + and SO 4 - - (g) Ca ++ and NO 3 ~
(rf) H + and CO 3 - - (h) Ca ++ and NH 4 OH
(j) How would you prepare:
(a) Solid CaCO 3 from solid Ca(OH) ?
(b) Solid CaSO 4 from solid CaCO 3 ?~
(c) Solid CaC 2 O t from solid CaSO 4 ?
Write equations for all reactions, and point out what equilibria are involved.
(4) What is the position of calcium in the Periodic System? Name the
elements in the alkaline earth group. Write a brief note illustrating gradations
of properties in this group, after examining the table given by Hildebrand on
page 268.
ASSIGNMENT 33
CARBONATE ION, BICARBONATE ION, AND CARBONIC ACID
1. This assignment, in which we shall study the chemistry of carbonic acid and
its ions, is introduced, not only because this subject is an extremely important
one, but also because it serves to illustrate how a large number of facts can be
correlated when the equilibria involved in a few simple, reversible reactions are
understood. Carbonic acid, H 2 CO 3 , is a weak, dibasic acid; and, like other weak
acids which contain more than one acidic hydrogen in the molecule, it ionizes in
steps, to form (/) bicarbonate ion, HCO 3 ~, and (2) carbonate ion, CO 3 ~~, thus:
H.CO, == H + + HCCV (1)
HCO 3 - = H + + CO 3 -- (2)
In order to understand what happens in the various reactions considered, it is
necessary to know the relative concentrations, in the initial and final solutions,
of the substances involved in these two rapid reversible reactions. Our problem,
then, is to determine how the concentrations of H + (or OH"), H 2 COo, HCO 3 ~
and CO 3 ~~ at equilibrium in a solution vary as the experimental conditions are
changed.
2. Obtain at the office a large and small two-holed rubber stopper, two pieces
of rubber tubing and a thistle tube for use in the experiment in paragraph 3 arid
a hard glass test-tube for use in the experiment in Paragraph 12.
[381
3. Preparation and properties of carbon dioxide. Experiment. Make a carbon
dioxide generator by fitting your half-liter flask with the thistle tube
(passing through the stopper nearly to the bottom of the flask), and an outlet
tube bent at right angles. Make a wash bottle for washing the CO 2 gas by
equipping a small flask with an entry tube extending nearly to the bottom of the
flask, and a delivery tube. Fill the small flask about half full of water. Place
in the generator a few lumps (about 5 grams) of limestone, cover with water,
and add hydrochloric acid, a little at a time, through the thistle tube.
Note the color and the odor of the gas. Calculate from the molecular weights
the relative densities of carbon dioxide and oxygen, and of carbon dioxide and,
nitrogen, at the same temperature and pressure. Question. Is carbon dioxide
denser or lighter than air? Experiment. Collect some carbon dioxide in a
test-tube by displacement of air; and determine if it is inflammable, and if it
supports combustion.
4. Suggest an experiment to prove that carbon dioxide is moderately soluble in
water at room temperature. Try the experiment. When CO 2 dissolves in water
the reaction is
CO 2 (gas) + H 2 O = H 2 CO 3 (in solution).
Prove by an experiment that the reverse reaction can take place. Under what
conditions is there an equilibrium? Note. At room temperature a solution in
equilibrium with CO 2 gas at 1 atmosphere pressure contains about 0.04 mols
H 2 COo in 1 liter. Question. How is this equilibrium altered by an increase of
temperature ?
5. We shall now consider what the principal substances are in a saturated
carbonic acid solution. Experiment. Prepare a small quantity of nearly
saturated solution of CO 2 and test it with the indicators, phenolphthalein. litmus,
methyl orange, and methyl violet; for comparison repeat indicator tests with
water and with an acid solution of known H + concentration. What conclusion
can you draw with regard to the concentration of H + in the saturated CO 2
solution? Note. The concentration of H + in this solution has been determined
by other methods to be about 0.0001 N (W~ 4 N). From this value for the
concentration of H + , the concentration of OH" in the same solution can be
calculated to be 10 ~ 10 N; explain how this calculation is made, cf . Hildebrand,
page 180.
6. Experiment. Test the action of carbonic acid on Ca ++ by passing CO 2
gas into a dilute solution of CaCl 2 . What conclusion can you draw with regard
to the concentration of CO 3 ~~ in carbonic acid solution? The concentration of
CO 3 " in the 0.04 M H 2 CO 3 solution has been found by other methods to be
of the same order of magnitude as that of OH~. Questions. Making use of the
fact that the only ions in this solution are produced by the ionization of H 2 CO 3
(and H 2 O), what conclusion can you draw with regard to the extent to which
each of the reactions (/) and (^), Paragraph 1, has taken place in a solution
of carbonic acid? What is the approximate concentration of HCO 3 " in the 0.04
M H 2 CO 3 solution? List the substances present in this solution (a) at moderate
concentration, (b) at small concentration ( two substances), and (c) at extremely
small concentration (two substances).
7. The neutralization of carbonic acid. When NaOH solution is added gradu-
ally to a carbonic acid solution the neutralization of H 2 CO 3 takes place in steps
corresponding to the two steps in the dissociation of H 2 CO 3 . The main reaction
for the first step in the neutralization is
HoCO s + OH- + (Na + ) =* H 2 O + HCCV + (Na + ),
and the solution made from eaual molal quantities of H 2 CO 3 and NaOH is a
sodium bicarbonate soultion. The second step in the neutralization of H 2 CO 3 is
HCO 3 - + OH- == H 2 O + CO 3 - -
[39]
This is the reaction for the neutralization of HCO 3 ~ by a strong base, and was
considered in Assignment 25.
8. Substances present in a solution of sodium carbonate. Review your notes,
Assignment 25, on the relative concentrations of CO 3 ~ ~, HCO 3 ~ and OH~ in a
0.5 M Na 2 CO 3 solution; and calculate the concentration of H + in this solution.
The concentration of H 2 CO 3 in this solution is also extremely small, but of course
is greater than the concentration of H + . Questions. How would you prepare 1
mol of solid Na 2 CO 3 from 1 mol of NaHCO 3 , and also from 1 mol of H 2 CO 3 ?
What quantity of NaOH would be required in each case? List the substances
which are present in a solution of Na 2 CO 3 , (a) at moderate (or high) concentra-
tion, (b) at relatively small concentration (two substances), and (c) at extremely
small concentration (two substances).
9. Substances present in a solution of sodium bicarbonate. From the fact that
HCO 3 ~ is a weak acid, what would you expect to be the principal substances in
a solution of NaHCO 3 ? Noting that the substance HCO 3 ~ is intermediate in
composition between H 2 CO 3 and CO 3 ~~, and considering the equilibria involved,
would you expect the concentration of CO 3 ~" in a solution of NaHCO 3 to be
greater or less than in a solution of (a) H 2 CO 3 , (b) Na,CO 3 ? The following
experiment, on the precipitation of CaCO 3 by NaHCO s solution at room temp-
erature, furnishes evidence with regard to the concentration of CO 3 ~~ in this
solution, (but it should be remembered that more CaCO 3 is precipitated than
corresponds to the actual concentration of the CO 3 ~~ since on account of the
displacement of equilibrium some more CO 3 ~~ forms as the CaCO 3 is precipitated.
Experiment. To 10 cc. M NaHCO 3 in a flask add about 50 cc. water, and 10 cc.
N CaCl 2 ; shake the mixture. Filter and note the amount of the CaCO 3 preci-
pitate. Compare the result with that obtained in the experiment with Ca ++ and
H,CO 3 , Paragraph 6. Repeat the experiment using 10 cc. TV (0.5 M) Na 2 CO 3
instead of 10 cc M NaHCO 3 . Heat the two nitrates to boiling, and test the
gas evolved in the NaHCO 3 experiment by passing it into a clear solution of
Ca(OH) 2 or Ba(OH) 2 (cf. Assignment 32, Paragraph 11). Compare the
amount of the precipitate now obtained with that formed in the cold solution.
Complete the following equation :
Ca ++ + 2HCO 3 - = CaCO 3 (solid)
and suggest an explanation for the shifting of the equilibrium when the mixture
is heated to boiling. The reaction that takes place when a solution of NaHCO 3
is heated to boiling is
2HC0 3 - = H 2 C0 3 + C0 3 " = C0 2 (gas) + H 2 O + CO 3 "
Questions. How will the concentrations of H 2 CO 3 and CO 3 ~~ in a NaHCO 3
solution be changed (/) when the solution is boiled, and (2) when CaCL
solution is added at room temperature? What conclusions can you draw with
regard to the relative concentrations of HCO 3 ~, and H 2 CO 3 and CO 3 " ~ in a freshly
prepared, cold solution of NaHCCX?
10. Experiment. Test a freshly prepared, approximately molal solution of
NaHCO 3 with litmus and with phenolphthalein and estimate approximately the
concentration of OH~ and H + in the solution. Repeat the experiment with a
solution which has been heated to boiling, or which has been allowed to stand
in the laboratory for several days. From your results in Paragraphs 9 and 10
state which substances are present in a pure NaHCO..- solution, (a) at moderate
(or high) concentration, (ft) at relatively small concentration, (two substances)
and (c) at still smaller concentration (two substances).
11. Summarize in tabular form the lists referred to in the last sentence of
each of the three Paragraphs, 6, 10 and 8, arranging- the columns in the order
H,CO 3 , NaHCO 3 and Na 2 CO 3 . Note the regular change in the concentration
of "each substance, e. g. H 2 CO 3 , when the three solutions are considered in this
[40]
order. If the arrangement of the various substances in your table is not sym-
metrical you have probably made some mistake. Question. What is the equation
for the main reaction when 0.1 mol of strong acid is added to (a) 0.1 mol
Na,CO 3 , and (b) 0.05 mol Na 2 CO 3 ?
12. The decomposition of solid sodium bicarbonate. Experiment. Place about
1 gram NaHCO 3 in a hard-glass test-tube and lead the delivery tube into a
solution of Ba(OH) 2 , or of CaCL and NH 4 OH. Heat the tube until no more
gas is given off, remove the delivery tube from the solution, and allow the hard-
glass tube to cool. What is the evidence that CO 2 was formed? That water
was formed ? What must the residue be if H 2 O and CO 2 are formed in equimolal
quantities? (This experiment may be performed quantitatively by weighing the
hard-glass tube (/) empty, (2) with NaHCO 3 , and (5) with the residue.)
Experiment. Continued. Clean the delivery tube and lead it into a fresh
solution of Ba(OH) 2 . Pour water carefully into the hard-glass tube until it is
half full, but do not stir the mixture. Add 2 cc. 6 N HC1, and at once insert the
stopper. What is the gas evolved? Warm the mixture slowly by placing the
hard-glass tube in a beaker of water and heating the latter. Questions. If the gas
volumes had been measured, what would have been the ratio of the volumes in
the two parts of the experiment? If all the CO 2 evolved had been converted into
BaCO 3 , what would have been the relative amounts obtained in the two parts
of the experiment?
13. Reaction between an acid and the negative ion of a weaker acid. Give
examples of the action of a strong acid on a solution of the salt of a weak acid.
Experiment. Determine the action of dilute acetic acid on NaHCO 3 .
Which is the stronger acid, HAc or H 2 CO 3 ? Taking into consideration the fact
that HCO 3 ~, the negative ion of H 2 CO 3 , is also a weak acid (whose ions are H*
and CO 3 ), and the additional fact that H 2 CO 3 is a stronger acid than HCO 3 ',
predict what will happen when H 2 CO 3 is introduced into a solution of a car-
bonate.
14. Experiment. Test your answer to the last question by placing a small
amount of freshly precipitated CaCO 3 in about 100 cc. water and saturating the
mixture with CO 2 . Filter off any CaCO 3 that remains and heat the filtrate to
boiling.
15. Summarize all the evidence, presented in this Assignment that the reaction
2HC0 3 - = H 2 C0 3 + C0 3 "
is reversible, noting especially the experimental conditions under which the re-
action will proceed almost completely, (a) as written, and (b) in the reverse
direction. Question. Which solution would be more alkaline at room tempera-
ture (on account of hydrolysis), a molal solution of a salt of a weak monobasic
acid of the same strength as H 2 CO 3 , or a molal solution of NaHCO 3 ? Give your
reasoning.
16. Problems. (/) Can the following substances be present at moderate con-
centrations in the same solution? If not, what is formed?
(a) H 2 CO 3 and Ca ++ (/) HCO 3 - and H +
(b) H 2 CO 3 and OH' (g) HCO 3 - and OH-
(c) H 2 CO 3 and CO 3 " (h) HCCV and CO,"
(<0 H 2 CO 3 and HCO,- (i) CO 8 and OH-
(e) H,CCL and H H
(2) What weight of Na 2 CO 3 can be obtained from 8.4 g. NaHCO 3 , (a) by
heating, and (b) by treating the bicarbonate with NaOH solution? What is
the least volume of N NaOH that can be used in (b) ?
[41] .
(j) What volume of CO 2 at standard conditions can be obtained from 8.4 g.
NaHCO 3 , (a) by heating, and (b) by treating the bicarbonate with HC1 solution?
What is the least volume of N HC1 that can be used in (b) ?
(4) When a current of CO 2 is passed into a solution of Ca(OH) 2 a precipitate
is observed to form, and then dissolve. If the final solution is now heated to
boiling, a precipitate again appears. State what has happened in this experiment,
and write an equation for the main reaction in each of these stages.
ASSIGNMENT 34
SULPHATES, CHLORIDES AND NITRATES OF COPPER AND ZINC
1. In Assignment 34 and the four following Assignments we shall study the
chemistry of copper ion, Cu ++ , silver ion, Ag + , and zinc ion, Zn ++ . We shall
first devote our attention to the common soluble salts, and to the method of
transforming one salt into another ; and then we shall study the sparingly soluble
compounds, the methods of dissolving them, and the equilibria involved. The
same method of treatment will be used later in studying the chemistry of ions
of other metals, and many of the results now obtained in studying copper, silver
and zinc are also true for other metals. Read again Paragraph 2, Assignment 31.
2. Solubility of nitrates, acetates, chlorides and sulfates of metals. The
nitrates and acetates of all metals are soluble in water and this is also true foi
nearly all the chlorides and sulfates. Question. Give an example of a diffi-
cultly soluble chloride and of a difficultly soluble sulfate. Give the formulas of
the nitrates, acetates, chlorides and sulfates of copper, silver and zinc.
3. Preparation of sulfates from nitrates and chlorides. Experiment. Test
the relative volatility of HC1, HNO 3 and H 2 SO 4 by evaporating a few drops of
a concentrated solution of each of these acid's in a casserole (out of doors or in
a fume closet). Cf. Hildebrand, pages 173-175. Suggest a method of preparing
solid copper sulfate from copper nitrate based on the difference in volatility
of HNO 3 and H 2 SO 4 . Experiment. Try your method with 15 cc. of the labora-
tory solution of copper nitrate. Recrystallize the copper sulfate by mixing the
residue, after it is cold, with 2 or 3 cc. water, heating the mixture to boiling
and letting it cool slowly. Wash the crystals with a very little water, test a
small portion for NO 3 ", and save the remainder.
4. Suggest a similar method for the preparation of solid zinc sulfate from
zinc chloride. How would you test whether the final product is free from
chloride ?
5. The problem of preparing a soluble chloride or nitrate from a soluble
sulfate will be considered in the following Assignment.
6. Conversion of soluble chlorides into nitrates, and nitrates into chlorides.
Experiment. Mix 2 cc. concentrated HNO 3 solution and 6 cc. concentrated
HC1 solution, and let the mixture stand. Evaporate a few drops of the solution
to dryness (out of doors or in a fume closet) and note if there is a residue.
The gradual deepening of the color of the solution at room temperature proves
not only that a reaction is taking place but that this reaction is not a rapid one.
Questions. How would the speed of the reaction be altered (a) by raising the
temperature, and (b) by using less concentrated acids? Note. Although this
reaction does not take place in a dilute solution which contains H + , Cl~ and
NO 3 ", it can be made to do so by concentrating the solution by evaporation to
a small volume. Since both Cl~, and NO 3 ~ are used up in this reaction and the
products are volatile, we can make use of this reaction (/) to remove Cl~ from
a solution by heating with excess of concentrated HNO 3 , or (2) to remove
[42]
NO 3 ~ from a solution by heating with excess of concentrated HC1. Question.
If excess of concentrated HNO 3 is added to a small amount of NaCl and the
mixture evaporated just to dryness, what would you expect the solid residue
to be. Note. The mixture of concentrated HNO 3 and HC1 is called aqua regia,
and the reaction that takes place is NO 8 - + 3 Cl~ + 4H + = NOC1 + CL, + H 2 O.
7. Experiment. To about 0.2 g. copper chloride (or 2 cc, N ZnCl 2 solution)
in a casserole, add 3 cc. concentrated HNO 3 and stir; add one or two drops of
this mixture to 10 cc. water, test this solution for Cl~ and set it aside for com-
parison with later results. Evaporate the mixture of the chloride and con-
centrated HNO 3 just to dryness (out of doors or in a fume closet)., add 5 cc.
concentrated HNO 3 being careful to dissolve any solid on the side of the
dish, and test one or two drops of the solution for Cl~ as before. Again
evaporate the mixture just to dryness, and test a portion of the residue for
chloride. Questions. What is the final solid residue? What conclusion can
your draw from a comparison of the amounts of the AgCl precipitates in your
three tests? Does this experiment furnish any evidence that the reaction is
slow, even at a temperature in the neighborhood of 100, when the concentration
of one of the reacting substances is small?
8. How would you convert copper nitrate, or zinc nitrate, into the chloride?
9. Treatment of silver chloride with hot concentrated nitric acid and sul-
furic acid. Experiment. Prepare some silver chloride, and wash it with dilute
HC1 and then with water. Place a small amount of the silver chloride in a
porcelain dish, and 3 or 4 cc. concentrated HNO 3 , evaporate the mixture to
dryness. Add a little water and test the solution for Ag + . Repeat the experi-
ment with the residue. Compare the result with that in the experiment in
Paragraph 7, and state how the speed of the reaction between Cl~ and NO 3 ~ in
acid solution depends on the concentration of Cl~.
10. Test the action of hot concentrated H 2 SO 4 on silver chloride by heating
for several minutes a small amount of the salt with 3 or 4 cc. of the acid in a
porcelain dish covered with a watch glass. Set the dish aside until it is cool,
pour the mixture into water, and test the solution for Ag + . Note. The chloride
is expelled as HC1 gas in this experiment, but since the temperature may have
exceeded 300 it is not surprising that the result is different from that obtained
in the experiment with AgCl and concentrated HNO 3 .
11. Problems. (/) What will be the solid residue when each of the follow-
ing solutions is evaporated to dryness?
(a) A solution which contains Ag + , H + and NO 3 ~.
(b) A solution which contains Cu ++ and NO 3 ~ at small concentrations and
H + and SO 4 '~ at large concentrations.
(c) A solution which contains Cu ++ and NO 3 ~ at large concentrations and
H + and SO 4 ~~ at small concentrations.
(d) A solution of zinc sulfate to which nitric acid has been added.
(2) How would you prepare solid silver sulfate from (a) solid silver nitrate,
(b) solid silver acetate?
(j) When solid copper sulfate is prepared by crystallization from an aaueous
solution it contains water of crystallisation and its formula is CuSO 4 5H 2 O.
When this blue substance is heated, white anhydrous CuSO 4 is finally obtained.
Write the equation. The reverse reaction takes place when air saturated with
water vapor is passed over the anhydrous salt at room temperature. If a vessel
containing CuSO 4 5H 2 O were evacuated for a few -minutes and then closed,
what substances would be present in the closed vessel? What reason do you
have for the conclusion that the pressure of water vapor in the vessel at room
temperature must be less than the partial pressure of water vapor .at the same
F431
temperature? How would the pressure in the vessel be altered by raising the
temperature? Give two other examples of a reversible reaction in which one
solid substance dissociates into another solid and a gas, and in each case give
an approximate value for the equilibrium pressure at some definite temperature.
(Refer to the lectures on the dissociation of calcium carbonate and calcium
hydroxide when heated.)
ASSIGNMENT 35
HYDROXIDES OF COPPER, SILVER AND ZINC
References. Hildebrand, Chapter V pages 72-73. 80-82; Chapter XIII pages
194-195.
1. We have learned from the laboratory work and the lectures that, the
hydroxides of sodium, potassium and the other alkali metals, and ammonium
hydroxide are readily soluble in water, and that the hydroxides of barium,
strontium and calcium are moderately soluble (with the solubility decreasing
rather rapidly in the order named). The hydroxides of all other metals are
difficulty soluble in water.
2. Cuprlc hydroxide. State what solutions you would mix to prepare
Cu(OH) 2 . Point out the substances involved in the equilibrium between the
solid and its saturated solution and predict what will happen when the solid is
treated with HC1, HNO 3 or H 2 SO 4 solution. Experiment. Test you*
answer by preparing some copper hydi oxide and treating it with acids. Also
determine if it dissolves in 0.5 'N NaOH.
3. Cuprlc oxide. Experiment. Collect some Cu(OH) 2 on a filter and
wash it once with water. Mix some of the Cu(OH) 2 with water and heat the
mixture to boiling. Place the remainder of the Cu(OH) 2 in a procelain dish
and heat gently. Cupric oxide has formed. Does the reverse reaction between
CuO and H 2 O take place at room temperature? Predict what will happen when
NaOH solution is added to a solution of Cu(NO 3 ) 2 at 100, and test your
answer. Test the action of HC1, HNO 3 or H 2 SO 4 solution on CuO. Question.
How would you prepare solid CuCl 2 from Cu(OH) 2 or CuO?
4. Sliver oxide. Experiment. To 10 cc. 0.1 TV AgNO 3 solution add
NaOH solution slowly until, after shaking, the solution reacts strongly alkaline
to litmus. Silver oxide, Ag 2 O, is formed. Filter, acidify the filtrate with
HNO 3 and test it for Ag + . Write the equation for the reaction between Ag +
and OH~ to form Ag 2 O. Point out the substances involved in the equilibrium
between the solid substance and its saturated solution, and predict what will
happen when the solid substance is treated with HNO 3 solution. Test your
answer by an experiment.
5. Relation between oxides and hydroxides of metals. Contrast the action of
water or water vapor at room temperature on calcium oxide, and on copper
oxide or silver oxide. Review Problem 3, Assignment 34; and note that the
reaction
Oxide of a metal (solid) + H 2 O (gas) = Hydroxide of a metal (solid)
takes place completely in the case of oxides of the alkali metals, as Na 2 O, does
not have any tendency to take place in the case of oxides of the more noble
metals, as Ag 2 O, and can be shown to be reversible for many oxides which
are intermediate between these two extremes. (In many cases equilibrium is
reached only very slowly,' especially at low temperatures.) On account of this
relation between oxides and hydroxides of metals, it is not surprising that oxides
of metals in general are capable of neutralizing acids. Questions. What is
anhydride of a base? Give examples. Is it always correct to say than an
[441
anhydride of a base will react with water to form the base? What is an acid
anhydride!' Give examples. Also give examples of reactions between an acid
anhydride and a base, and between an acid anhydride and a basic anhydride.
6. Acids which contain oxygen may be considered to be related to the hydrox-
ides of non-metals. Thus H 2 SO 4 may be regarded as H 6 SO 6 or S(OH) from
which 2 molecules of water have been withdrawn
S(OH) 6 = H 2 SO 4 + 2 H 2 O;
and the formula of phosphorous acid, H 3 PO 3 , may be written P(OH) 3 . The
characteristic properties of these acids, to yield negative ions containing oxygen
on dissociation or neutralization, is then to be attributed to the existence of
strong bonds between the non-metal and the oxygen in the compound; which is
really due to the great tendency of non-metals to hold electrons firmly. On the
other hand metals in compounds do not hold electrons firmly, and positive ions
of metals are formed when bases or salts dissociate. However, certain elements
which form a positive ion also can form a negative ion containing oxygen; the
hydroxide of such an element has the properties of a base in that it can neutralize
an acid, and also has the properties of an acid since it can neutralize a base: it
is said to be amphoteric.
7. Zinc hydroxide. Experiment. Prepare some Zn(OH) 2 by treating a
solution containing Zn ++ with a very small amount of NaOH. Collect the
Zn(OH) 2 on a filter, treat a portion with a strong acid, and another portion with
NaOH or KOH solution. Question. What evidence is furnished by this
experiment that zinc hydroxide is an amphoteric substance? The reaction be-
tween zinc hydroxide and excess OH~ is analogous to the first stage in the neu-
tralization of carbonic acid :
H ZnOo (solid) + OH- = H,O + HZnCX-
H 2 CO S + OH- H 2 O + HCO S -
The negative ion, HZnO 3 ", is usually zincate ion, though strictly speaking this
name should be reserved for the ion ZnO 2 ~~. The latter substance probably
forms to some extent when the concentration of OH" is very large. Solid
NaHZnO, has not been prepared, but Na^ZnOo can be made by fusing zinc oxide
with NaOH.
8. Predict what will happen when a solution of a strong acid is added slowly
to sodium zincate solution. Test your answer by an experiment.
9. Preparation of a nitrate or chloride from a soluble sulfate. Experiment.
Prepare a solution of copper nitrate from copper sulfate solution by precipitat-
ing copper hydroxide from a very dilute CuSO 4 solution, washing the Cu(OH) 2
precipitate, and transforming it into the nitrate. Test the product for sulfate.
10. Suggest a second method based on the removal of the sulfate ion by
means of barium ion (cf. Hildebrand. pages 207-208). Try your method, and
test the copper nitrate solution for SO 4 ~~ and for Ba ++ . Which is the better
method to use on a small scale in the laboratory?
11. Suggest two methods of preparing solid zinc chloride from zinc sulfate.
12. Problems. (/) Write equations for the action of HC1 solution in excess
on calcium oxide, on ferrous and ferric oxides (FeO and Fe.X) 3 ), on ferrous and
ferric hydroxides, and on sodium zincate solution.
(2) A solution containing 0.2 mol Na 2 SO 4 is mixed w r ith a solution contain-
ing 0.1 mols Ba(NO 3 ) 2 , the precipitate is removed by filtration, the wash water
is run into the filtrate until the final volume of the latter is 1 liter. What sub-
stances are present in the final solution and what is the concentration of each?
Suggest a method of preparing pure sodium nitrate from sodium sulfate.
( j) Give two methods of preparing CuCL from CuSO 4 .
[45]
(4) Give an example of an amphoteric hydroxide other than Zn(OH) 2 , and
write the equation for its reaction with (a) a strong acid, (b) a strong base.
ASSIGNMENT 36
COMPLEX IONS OF COPPER, SILVER AND ZINC WITH AMMONIA
References. Hildebrand, pages 184 and 202-205.
1. The ions of certain metals have the power of forming compounds with
NH 3 , which are examples of "complex ions". The ammonia is supplied by
adding NH 4 OH solution. There is an equilibrium between the complex ion,
NH 3 or NH 4 OH, and the ion of the metal; and when the NH 4 OH is present
in excess the concentration of the ion of the metal is often very small. Some
examples will be considered in the present Assignment.
2. Experiment. Dilute 5 cc. N CuSO 4 with 20 cc. water and add two or
three drops 6 N NH 4 OH. The precipitate is Cu(OH) 2 . Question. What
evidence with regard to the solubility of Cu(OH) 2 in water is furnished by this
experiment? (Consider the low concentration of OH~ in NH 4 OH solution and
the result of the experiment in Paragraph 6, Assignment 32.
3. Continue to add the NH 4 OH solution until, after shaking, the solution is
clear. Observe the volume of NH 4 OH used. What is the evidence that a new
substance is formed? Repeat the experiment with the same volume oi 6 N
NaOH instead oi 6 N NH 4 OH. State the reasoning by which you may con-
clude that the substance obtained in the experiment with NH 4 OH solution is
not formed by a reaction between OH~ and Cu(OH) 2 similar to that between
OH~ and Zn(OH) 2 . Try an experiment with CuSO 4 solution and NH 4 C1
in order to determine if Cu ++ reacts with NH 4 + in the same way as with NH 3
or NH 4 OH.
4. The new substance formed when Cu(OH) 2 dissolves in NH 4 OH solution
is Cu(NH 3 ) 4 ++ . The final solution contains NH 4 OH, and OH- must be present
at an appreciable concentration ; it may be concluded that the concentration of
Cu ++ is extremely small, since, otherwise, Cu(OH) 2 would precipitate. Questions.
Which solution has the greater concentration of Cu ++ , a saturated solution of
Cu(OH) 2 or a solution of the complex ion containing NH 4 OH? How is the
equilibrium between Cu(OH) 2 and its saturated solution affected when NH.,OH
is added?
5. The equation for the main reaction when a solution of a copper salt is
treated with excess NH 4 OH is
Cu ++ +4NH 4 OH = Cu(NH 3 ) 4 ++ + 4H 2 O.
A simple method of demonstrating that the reverse reaction will take place is to
remove the NH 4 OH. Suggest two methods of doing this and try your methods.
6. Zinc and silver also form complex ions with ammonia,
Zn(NH 8 ) 4 " and Ag(NH 3 ) 2 + .
Experiment. To a dilute solution of ZnSO 4 add NH 4 OH solution drop by
drop, shaking the mixture after each addition of NH 4 OH; when some but not
all of the Zn(OH) 2 has dissolved test the solution with trinitrobenzol to determine
approximately the OH~ concentration. Repeat the experiment, using NaOH
instead of NH 4 OH, and note again the OH- concentration in the solution when
Zn(OH) 2 is partly but not completely dissolved. Question. From a comparison
of the OH~ concentration at equilibrium in the two cases what conclusion can
you draw with regard to the presence of zinc as a different substance in the
NH 4 OH and NaOH solutions?
[46]
7. Experiment. Prepare some silver oxide and some silver chloride, and
treat each substance with NH 4 OH solution. Predict what will happen when
the resulting solutions are acidified with nitric acid, and test your answers by
experiments.
8. Which solution would contain a greater concentration of Ag + , a saturated
solution of silver acetate (see Assignment 24, Paragraph 5) or a solution which
contains the complex ion and NH 4 OH? What then must be the action of
NH 4 OH solution on solid AgAc?
9. Questions. When AgCl is dissolved in excess NH 4 OH solution, what is the
principal positive ion, and the principal negative ion, in the resulting solution?
When o.i mols AgCl are dissolved and the final volume is 100 cc. what is the
approximate concentration of each of these ions? Is the concentration of Ag*
smaller or larger than in a saturated solution of AgCl in water?
10. When Cu(OH) 2 is dissolved in excess NH 4 OH solution what is the
principal negative ion in the final solution? , It is important to realize that
Cu(NH 3 ) 4 (OH) 2 , Zn(NH 3 ) 4 (OH) 2 and Ag(NH 3 ) 2 OH are soluble, strong
bases.
11. We shall next consider the action of NH 4 OH solution on silver iodide, a
salt which is much less soluble in water than either silver oxide or silver chloride.
Solid Agl, suspended in water, is in equilibrium with Ag + and I~ in the
saturated solution. When NH 4 OH is added some Ag + must react with it to
form Ag(NH 3 ) 2 + , and some solid Agl must dissolve to establish again the
solubility equilibrium. We can therefore predict that Agl must be more soluble
in NH 4 OH solution than in water, but it is not safe to predict that a large
amount of it will dissolve. Experiment. Prepare a solution of silver oxide
or silver chloride in NH 4 OH and add to this solution a few drops of 0.1 N potas-
sium iodide solution. In another experiment collect some solid silver iodide on
a filter paper, wash it with water, and treat it with dilute NH 4 OH solution.
Devise and try an experiment with the solution thus prepared to determine
whether an appreciable amount of Agl has dissolved.
12. Predict what will happen when silver chloride is treated with a solution
of potassium iodide. Try the experiment.
13. The ions of the alkali metals, as Na + , K + , etc.; those of the alkaline earth
metals, as Ca ++ , Sr TT and Ba ++ ; and many other positive ions do not combine
with NH 3 to form complex ions. Question. Will a difficultly soluble salt of
calcium, as CaCO, or CaC 2 O 4 , dissolve in NH 4 OH solution to a greater extent
than in water?
14. Problems, (i) Write equations for the main reactions in the following
experiments :
(a) A few drops of NH 4 OH solution are added to a solution contain-
ing Zn ++ .
(b) Excess NH 4 OH solution is added to a solution containing Zn ++ .
(c) An ammoniacal solution of silver chloride is acidified with nitric acid.
(2) The solubilities of AgCl and Agl in water are 10~ 4 and 10~ 7 mols per
liter, respectively. In two experiments with NH 4 OH solution each substance
is found to be one thousand times more soluble than in water. How many grams
of each has dissolved in i liter of the NH 4 OH solutions?
(j) Give two examples of complex ions other than those with ammonia.
(4) From the positions of Cu, Ag, and Zn in the Periodic System (Hildebrand
page 257), state- what other ions might be expected to form ammonia complexes.
Check your answer by referring to Hildebrand, page 269.
[47]
ASSIGNMENT 37
CARBONATES AND SULFIDES OF COPPER, SILVER AND ZINC
References. Hildebrand, pages 199-201 and 195-197.
1. In Assignment 33 we learned that when carbon dioxide dissolves in
water the solution contains the weak acid H 2 CO 3 , and in the saturated solution
we recognized an equilibrium between H 2 CO 3 in the solution and CO 2 in the
gas space. We found that carbonic acid ionizes in two steps and can be neu-
tralized in stages to give solutions of the two corresponding types of salts, and
that when these salts are treated with a strong acid carbon dioxide is evolved.
In the present Assignment we shall continue the study of difficultly soluble car-
bonates, and we shall deal with the salts of another dibasic acid, H 2 S, which is
in many respects analogous to carbonic acid.
2. Hydrogen sulfide, like carbon dioxide, is a gas which is soluble in water,
and when the solution is saturated an equilibrium is established between gaseous
and dissolved H 2 S. The solution has the properties of a weak acid which ionizes
in two stages as does carbonic acid. Sulfides, in general, are less soluble than
carbonates ; in the case of the alkali metals both the sulfides and carbonates are
readily soluble, while in the case of the alkali earth metals the sulfides are
readily soluble and the carbonates are not. Questions. Write the two equations
which show the ionization of H 2 S (dissolved in water) in two stages. What
are the names of the negative ions formed ? In analogy with the ions of car-
bonic acid, which is the stronger acid, H 2 S or HS"? What are the principal
substances present in a solution of sodium sulfide, Na 2 S? In a solution of
sodium hydrogen sulfide, NaHS? What reactions take place when a strong acid
is added to any solution containing sulfide ion? How, then, will the solubility of
difficultly soluble sulfides be affected by the addition of a strong acid?
3. Experiment. Try the action of Na 2 CO 3 solution on 0.1 N solutions of
Cu+ + , Ag + , and Zn ++ . In each case collect the precipitate on a filter, wash it
with water, and test a portion for carbonate. Predict the action of nitric acid
solution and of ammonium hydroxide on these precipitates, and test your pre-
diction by experiments.
4. In the above experiment with Cu ++ and in many other cases, the precipitate
obtained by the addition of Na 2 CO 3 to a solution of a salt is not a pure car-
bonate, but a substance which contains both the carbonate and hydroxide
radicals. Such substances, called basic carbonates, are difficult to obtain pure,
the actual precipitates consisting of variable mixtures of these basic carbonates,
carbonates, and hydroxides.
5. Experiment. Prepare 20 cc. of 0.1 N solutions of Cu ++ , Ag% and
Zn + +. To each add I cc. 6 N H 2 SO 4 and pass in H 2 S gas* until the liquid is
saturated. To determine this, close the end of the test-tube or flask, shake
thoroughly and test the odor cautiously. Collect the precipitates on separate
filters. In the experiment with Zn ++ and H 2 S, to the filtrate (or to the clear
solution if there was no precipitate), add 10 cc normal NaAc solution, and
again saturate with H 2 S gas. Question. What effect has the addition of Ac" on
the concentration of H + ?
6. H 2 S and HS~ are weaker acids than H 2 CO 3 and HCO 3 ~, respectively. What
can you conclude about the relative concentrations of S~~ and CO. { ~~ in solutions
of H 2 S and H 2 CO 3 ? It will be recalled that 'H 2 CO 3 does not precipitate CaCO 3
* Caution. H 2 S is poisonous. Work out of doors if possible and do not breathe the gas.
[481
from a solution of CaCL (Assignment 33, Paragraph 6). Question. What
conclusion can you draw with regard to the relative solubility of CaCO 3 and
the sulfides prepared in the preceding Paragraph?
7. From the equation for the main reaction
M ++ + H 2 S = MS (solid) + 2H+
it is evident that both M ++ and H* are competing for S~~ and that an increase
in the H + concentration tends to reverse the reaction, as we have seen in the
case of ZnS. However in the case of a very insoluble sulfide, CuS for example,
even though H + concentration is fairly large the concentration of M ++ will be
negligible if the solution is saturated with H 2 S. Determine by experiment which
of these sulfides will dissolve completely in 2 N H 2 SO 4 . Which then has the
greater solubility in water? Write the two equations which show the mechanism
of the main reaction. State how you would precipitate a sulfide which is spar-
ingly soluble in water, but many times more soluble than ZnS (cf. the precipita-
tion of CaCO 3 ).
8. Experiment. To solutions containing Cu(NH 3 ) 4 ++ , Ag(NH 3 )o + , and
Zn ( NH 3 ) 4 + +, in separate experiments, add (NH 4 ) 2 S solution. Repeat the
experiments, using H 2 S gas instead of (NH 4 ) 2 S solution. State in each case
which has the smaller concentration of the ion of the metal, a solution con-
taining the complex ion, or a solution saturated with the sulfide. State also
which is the less soluble in water, Cu(OH) 2 or CuS, Ag 2 O or Ag 2 S, Zn(OH) 2
or ZnS. From the fact that the sulfides are less soluble than the carbonates,
predict the action of H 2 S on mixtures of basic copper carbonate and water,
silver carbonate and water, and zinc carbonate and water. Test your answers
by experiment.
9. Experiment. Test the action of H 2 S on solutions of nitric acid of
different concentrations, 0.1 N, 2.0 N, and 6 N. Saturate each solution with
H 2 S gas and heat the mixture almost to boiling. The white precipitate formed
is solid sulfur and the principal reaction is
3H 2 S + 2NO 3 - + 2H + = 3S (solid) + 2NO + 4ELO.
Questions. What evidence is furnished by your experiments that this is a slow
reaction? How is the speed affected by the concentration of the nitric acid, and
by the temperature? (Note. The products of the reaction, S, NO, and H 2 O,
do not react with each other under any ordinary conditions of temperature and
pressure, and it is practically impossible to realize an equilibrium involving the
substances shown in the equation.) Because of this reaction nitric acid in a
strongly acid solution will remove sulfide ion from the solution more completely
than will H + alone since it destroys the H 2 S which is in equilibrium with sulfide
ion. We may therefore expect that nitric acid will be a better solvent for
difficulty soluble sulfides than sulfuric or hydrochloric acid.
10. Experiment. Collect some CuS on a filter, transfer it to a casserole,
add 2 normal HNO 3 and boil the mixture. Filter, and test the filtrate for
Cu ++ . The residue collected on the filter is sulfur; the dark color is due to a
little CuS enclosed in the sulfur. Complete the equation
3CuS (solid) + 2NO 3 - = 3S (solid) + 2NO
Problems. (/) Hydrogen sulfide is passed into 1 N CuSO 4 until the precipita-
tion of copper sulfide is complete. The liquid is filtered and H 2 S is removed by
boiling. What concentration of H + is left in the solution.
(2) Write the equations for the neutralization of (a) one mol H.,S with one mol
OH-, (b) H 2 S with excess OH~, (c) HS~ with OH~.
(j) Predict the effect of NH 4 OH on the solubilities of CuS, ZnS and
T491
Ag 2 S. Give your reasoning. (If you try the experiments use freshly prepared
sulfides moistened with water which contains H 2 S.)
(4) Give a method for the preparation ot
(a) CuSO 4 solution from CuS.
(b) Ag 2 S(solid) from AgCl.
(c) ZnCO 3 (solid) from ZnS.
(d) Na 2 CO 3 solution from Na 2 S solution; see experiment at end of
Paragraph 8, and Hildebrand, pages 207-208.
ASSIGNMENT 38
REVIEW OF THE CHEMISTRY OF THE POSITIVE IONS ALREADY CONSIDERED
1. This Assignment is inserted in order that the student may summarize and
study the chemistry of the positive ions considered in the last seven Assign-
ments. If the student makes an accurate summary, and notes carefully similar-
ities and differences in the results for the various ions, he will find that it is not
difficult to remember the large number of facts involved. In later work, as
soon as the chemistry of another ion has been investigated, the results should
be summarized, and compared with the results for the ions already studied.
2. Summary of results. Tabulate acording to the following plan the results
which you have obtained in earlier Assignments ; refer to your laboratory notes
whenever necessary. Write in a horizontal line across a double page of your
notebook the names of the positive ions, silver, cupric copper, zinc, calcium,
sodium, potassium and ammonium ; in the vertical columns write the following :
(/) the formulas of the ions, including complex ions;
(2) the formulas of the compounds which are readily soluble in water,
(j) the formulas of the compounds which are only moderately soluble, as
Ca(OH) 2 , CaS0 4 , Ag 2 SO 4 .
(4) the formulas of the insoluble substances (the term difficultly soluble
or sparingly soluble is preferable). Indicate which is the least soluble com-
pound of each metal. When any ion or difficultly soluble compound is colored,
give the color in the table. Next make as general statements as you can about
the solubilities in water of the nitrates, chlorides, sulfates, carbonates, hydroxides
(or oxides), and sulfides of these and other metals.
3. Show how you can apply your knowledge of the equilibria involved, in
choosing reagents to dissolve the various difficultly soluble substances listed in
the preceding Paragraph. In this connection, for each complex ion compare
the concentration of the ion of the metal in a solution which contains the complex
ion with that in a saturated solution of the various difficultly soluble salts of
the metal; cf. Hildebrand's discussion of the reactions of silver ion, pages
204-205, and the following example. Since ZnCO 3 dissolves readily in excess
NH 4 OH solution, and ZnS does not, it may be concluded that in a solution
which contains Zn (NH 3 ) 4 ++ and NH 4 OH the concentration of Zn ++ is less than
in a saturated solution of ZnCO 3 but greater than in a saturated solution of ZnS.
4. The Summary in Paragraph 2 was obtained from the results of experiments
on the reactions between the positive ions and various reagents. Conversely
the results of such experiments can be predicted by means of the Summary and
a knowledge of the various equilibria involved.
[50]
5. Table of Reactions. Prepare for future reference a table of the reactions
between dilute solutions of the positive ions considered above and the following
reagents :
(a) HC1 or a soluble chloride,
(b) NaOH in small amount,
(c) NaOH in excess,
(d) NH 4 OH in small amount,
(e) NH 4 OH in excess,
(/) H 2 C0 3 or CO, gas,
(g] a soluble carbonate,
(A) H,S in 0.3 N H + solution,
(*) a soluble sulfide as Na 2 S, or (NH 4 ) 2 S in NH. t OH solution,
(/ ) sodium cobaltinitrite in the presence of a few drops of HAc.
When no reaction takes place mark X at the proper place in the table. In
other cases write the formulas of the substances formed, note the colors, and
mark the precipitates by underlining the formulas. Be sure that you can
write the equation for each reaction. Try experiments whenever it is neces-
sary, e. g. when your laboratory notes are incomplete or seem to be incorrect,
and when there is a blank space in your table.
6. Test your knowledge of the reactions just considered by writing out from
memory portions of the Table of Reactions, Paragraph 5 ; write equations for the
reactions that take place, and discuss briefly the equilibria involved.
Note. It is recommended that the student proceed with Assignment 51, the
first in the Section entitled Systematic Qualitative Analysis. It is important that
the student become thoroughly familiar with the reactions already considered
before proceeding to study other reactions, and this Assignment deals only with
the ions already considered.
51
SECTION IV
REACTIONS OF IONS, CONTINUED
ASSIGNMENT 41
OXIDATION AND REDUCTION. REPLACEMENT REACTIONS. ELECTRICAL CELLS
References. Hildebrand, Chapter 'VI pages 93-94, Chapter V page 70, Chapter
XV pages 234-237.
1. In Assignment 41 and the three following Assignments we shall consider
examples of oxidation and reduction reactions. In these reactions changes of
valence occur, and they are therefore easily distinguished from other reactions
such as the formation of weak electrolytes or the precipitation of difficultly soluble
salts. The following reactions can now be recognized as oxidation and reduc-
tion reactions :
Zn (solid) + 2H + = Zn ++ + Ho (Assignment 3)
H, + CuO (solid) = Cu (solid) + HoO (Assignment 4)
Ca (solid) + 2H 2 O = Ca +V + 2OH- + H, (Assignment 32)
3H 2 S + 2NO 3 - + 2H + == 3S + 2NO + 4EUO (Assignment 37)
3d- + NO 3 - + 4H + = 3/2 C1 2 + NO + 2H.,O
= C1 2 + NOC1 + 2H 2 6 (Assignment 34)
Point out the changes of valence in each reaction.
2. In the first of the above reactions the oxidation of Zn to Zn ++ and the
reduction of 2H + to H 2 consist simply in the transfer of two electrical charges
from two hydrogen atoms to one zinc atom (which in reality is the transfer
of two electrons from the zinc atom to the two hydrogen atoms). In the pres-
ent Assignment we shall study a number of examples which are characterized
by such an interchange of electrical charges. The reaction between copper and
silver ion will first be studied quantitatively and then a number of other cases
will be taken up qualitatively.
3. Experiment. Obtain at the office 100 cc. 0.10 N silver nitrate, and a
special filter paper which burns to give an ash of negligible weight. Clean a
straight piece of thin copper wire and weight about 0.3 g. to 5 mg. Cover the
copper wire with 100 cc. of 0.10 N silver nitrate solution and let it stand until
the copper has disappeared. The precipitate formed is metallic silver. Questions.
What substance is responsible for the color of the final solution? What is
the color of silver ion? Heat a porcelain crucible, let it cool, and. weight to 5
mg. Collect the silver by filtering through the special filter paper and wash
the precipitate until the washings are colorless. (Save the filtrate, and place in
it another piece of copper in order to recover the rest of the silver.) Fold the
filter paper so that it will fit into the crucible, heat with a small flame until it is
dry, and then with a larger flame until the paper is completely consumed. Let
the crucible cool and again weigh to 5 mg. Place all the silver in the jar marked
"Silver Waste."
4. Calculations. Calculate (a) the weight of silver formed in the experi-
ment; (b) the weight of silver precipitated by one gram atom of copper. Com-
pare the latter number with the atomic weight of silver. (The formula of
silver ion is Ag + .) What is the formula of copper ion?
5. Questions. Give the formulas of silver nitrate, silver sulfate, copper
nitrate and copper sulfate. If one gram molecule of silver ion has associated
with it 96,500 coulombs of positive electricity what quantity of electricity is
[521
associated with one gram molecule of copper ion? With a gram molecule of
nitrate ion? Write the equation for the action of silver ion on copper. How
would the result of this experiment have been influenced by replacing the silver
nitrate with silver sulfate solution?
6. Experiment. Into 10 cc. N Cu(NO 3 ) 2 place a piece of lead.
7. Repeat the experiment with a solution of lead nitrate, Pb(NO 3 ) 2 , and
zinc.
8. Write the equations for all the reactions between metals and positive ions
which you have studied (Assignments 3 and 41). Questions. The negative
ion present in each experiment is not shown in the equation; what experi-
mental fact is thereby implied? If in one of the above experiments the dry
salt had been used instead of its solution, would the reaction occur sufficiently
rapidly at room temperature to be detected? Try an exepriment, if necessary,
with lead and solid copper sulfate.
9. Experiment. Determine by experiment which of the metals used in
this Assignment will dissolve readily in dilute hydrochloric acid (or sulfuric
acid).*
10. Arrange the metals used in these experiments in the order in which
they will replace each other from solution. After referring to your lecture
notes state where hydrogen belongs in this replacement series. From this
replacement series predict what will happen when,
Silver nitrate solution is treated with lead.
Copper nitrate zinc.
Copper nitrate silver.
Experiment. Test the correctness of your prediction.
11. Replacement series for the halogens. Experiment. Make a small gas
generator (see Assignment 33, Paragraph 3) and charge the flask with about
2 g. of manganese dioxide. Pour 5 cc. concentrated HC1 into the flask
and warm gently until a continuous stream of chlorine is evolved. Pass a few
bubbles of the gas through a very dilute solution of potassium bromide, pre-
pared by diluting a few drops of the laboratory solution with 5 cc. water.
Shake the solution with I cc. of carbon bisulfide. The brown color in the car-
bon bisulfide layer is characteristic of bromine.
12. Experiment. To 1 cc. potassium iodide solution add a drop of bromine
water. Shake with carbon bisulfide. The violet color is characteristic of iodine.
13. Questions. In each of the above experiments write the equation for the
reaction which has taken place. Point out the changes of valence and the inter-
change of electrical charges. Predict what reaction will occur when (a) chlorine
gas is passed through KI solution, (b) bromine is added to KC1 solution, and
(c) iodine is added to KBr solution. Experiment. Test the correctness
of your predictions.
14. The Conversion of Chemical into Electrical Energy. The reactions studied
in the present Assignment may be utilized to convert chemical into electrical
energy. In each of the reactions an interchange of electrical charges has taken
place; and the problem in the construction of an electrical cell is to arrange the
substances involved in the reaction in such a way that this interchange of
electricity can take place only through an external conductor such as a wire
joining the electrodes. Thus, in the reaction between zinc and copper ion two
changes of an electrical nature are involved: on the one hand zinc is changed
to zinc ion, and, on the other, copper ion is changed to copper. When the zinc
*The action of nitric acid on metals will be studied in Assignment 42.
[531
is placed directly in a solution of a copper salt the electrical changes both occur
at the surface of the zinc and the transfer of electricity cannot be directly
detected; but when we arrange the materials so that the zinc electrode dips into
a zinc sulfate solution contained in a porous cup, and this is placed in a vessel
containing copper sulfate solution and the copper electrode, then no reaction
takes place. If now we join the two electrodes by a wire, electricity passes
through the wire and zinc will change to zinc ion at the surface of the zinc,
while copper ion changes to copper at the surface of the copper. The total
change in the cell is of course the same as when zinc acts directly on copper
ion.
15. Draw a diagram of the elecrical cell described above. Question. What
substances carry the electricity from one electrode to the other through the
solution ?
16. In a similar way the reaction between chlorine and bromide ion may
be used to generate electrical energy; but, to lead the current into and out of
the solution, we must use some substance, as graphite, which conducts elec-
tricity and is not attacked by chlorine or bromine. In one part of this elec-
trical cell there would thus be a graphite rod dipping into a solution which
contains bromide ion and bromine, and in the other part a graphite rod in a
solution of chloride ion and chlorine. The two solutions must be separated,
e. g., by a porous cup, in order to prevent the chlorine from reacting directly
with the bromide ion. When the graphite electrodes are joined by a wire a
current will flow, and at the one electrode bromide ion will change to bromine
while at the other electrode chlorine will change to chloride ion.
17. Problems. (/) What quantity of electricity (in coulombs) flows
through the circuit of the cell described in Paragraph 13 when one mol of zinc
disappears? What weight of copper will at the same time be deposited on the
copper electrode?
(2) Sketch the arrangement of a cell in which the reaction between
metallic copper and silver ion is used to produce an electric current. When a
current is taken from this cell what changes take place at each electrode?
(j) How many coulombs of electricity could be obtained by the transforma-
tion of 1 liter chlorine gas (measured at C and 760 mm.) to chloride ion?
ASSIGNMENT 42
OXIDATION OF METALS TO THEIR IONS. TABLE OF OXIDIZING
AND REDUCING AGENTS
Reference. Hildebrand, Chapter XV pages 226-245.
1. In Assignment 41 we found that certain metals can be oxidized by
hydrogen ion and that other metals cannot. We may say, therefore, that certain
metals are sufficiently strong reducing agents to reduce hydrogen ion to hydro-
gen and that others are not, and it follows that any reagent which will change
one of the latter metals to its ion is a more powerful oxidizing agent than is
hydrogen ion. In the present Assignment we shall study the action on zinc,
copper, and silver of two such oxidizing agents, nitric acid and bromine.
2. Arrange copper, silver, zinc and hydrogen, in a vertical column, in a
replacement series. Beside each symbol write the formula of the ion to which
the element can be oxidized. Your table now is a table of oxidizing
and reducing agents arranged according to their strengths, and this table
[541
may be extended to include other substances than the metals and their
ions. Questions. Which is the more powerful reducing agent, Zn or Cu? Cu
or Ag? Which is the more powerful oxidizing agent, Ag+ or Cu ++ ? Ag+ or H + ?
(State the experimental facts upon which you based your answers to the pre-
ceding questions.) In describing the reaction between zinc and dilute hydro-
chloric or sulfuric acid solution is it correct to say that metallic zinc has com-
bined with chloride ion or with sulfate ion? How can the solid salts be pre-
pared from the solutions obtained in such experiments?
3. Action of nitric acid on zinc, copper, and silver. Experiment. Treat
about l / grams of zinc with 2 .V HNO 3 in a small beaker, covered with a
watch-glass. If the reaction is slow warm the mixture. The main reaction
under these experimental conditions is
3Zn (solid) + 2NO 3 - + 8H+ = 3Zn + * + 2NO (gas) + 4H 2 O.
4. Repeat the experiment with about 0.5 g. copper, and with a small
piece of silver. The reactions are similar to the one between zinc and
nitric acid, and the change from Zn to Zn + + corresponds exactly to the change
from Cu to Cu ++ and from 2 Ag to 2 Ag + . Write the equations for the reactions.
Question. How is the speed of the reaction between nitric acid and the metals
affected by a change in temperature and by a change in the concentration of the
acid?
5. Action of bromine on zinc, copper and silver. The halogens, chlorine,
bromine and iodine, all have a similar action on metals. To illustrate this
action we use bromine because its aqueous solution is a convenient reagent.
The solution of chlorine in water does not keep well, while iodine is a solid
which is only very slightly soluble in water. Experiment. Treat a small piece
of zinc with bromine water. After a few minutes pour the solution into a
porcelain dish, and heat it to boiling to expel any bromine that is present. To
a portion of the solution add a few drops of silver nitrate: AgBr, like AgCl
and Agl, is a difficulty soluble salt. Test another portion of the solution for
zinc ion. What evidence have you that a reaction takes place between zinc
and bromine? Repeat the experiment with copper and bromine water. Silver,
also, reacts with bromine.
6. The equation for the reaction between zinc and bromine solution is
Zn (solid) + Br 2 = Zn ++ +2Br
Since in this reaction zinc is oxidized to zinc ion we must conclude that bromine
is reduced when it changes to bromide ion. Write the equation for the action
of bromine on copper and on silver.
7. Review the replacement series for the halogens (Assignment 41) and
arrange chloride ion chlorine, bromide ion bromine, iodide ion iodine in a
table so as to give the relative strengths of the halogens as oxidizing agents and
of the halide ions as reducing agents. Add this table, in its proper place, to the
one previously prepared for the metals and their ions. Place Ca Ca + * and
Na Na + in this table. Show your completed table at once to the instructor.
8. We can determine the position of nitric acid in this table of oxidizing and
reducing agents by reference to the results of the experiment in Paragraph 6,
Assignment 34, when hot concentrated nitric acid was used to remove chloride ion
from a solution. Chlorine was formed and we can therefore decide that hot con-
centrated nitric acid is a stronger oxidizing agent than chlorine. Question. Is con-
centrated nitric acid a stronger oxidizing agent than bromine or iodine?
9. Study the method of balancing oxidizing and reducing reactions. Hilde-
brand, pages 226-233.
10. Problems, (i) Summarize the behavior of Na, K, Ca, Zn, Cu,. and other
[55]
metals discussed in the lectures by dividing the metals into the three following
classes :
I. Those which react readily with water. Hydrogen is evolved.
II. Those which react readily with HC1 or H 2 SO 4 solution, but not
with water. Hydrogen is evolved.
III. Those which dissolve readily in HNO 3 or Br, solution, but not in
HC1, H 2 SO 4 or water. H 2 is not evolved.
Write the equation for the action of HC1 or H 2 SO 4 solution on metallic sodium.
(2) In the experiment in Paragraph 11, Assignment 41, manganese dioxide,
MnO 2 , reacted with chloride ion in acid solution to form chlorine. Manganous
ion, Mn ++ , was formed. Write the equation for the reaction. Predict what will
happen when copper is treated with MnO 2 in the presence of dilute sulfuric
acid.
(j) Sketch the arrangement of a cell in which the reaction between zinc
and bromine is used to produce an electric current. What changes take place
at the electrodes? State how you could bring about the reversal of all the reac-
tions in this cell.
ASSIGNMENT 43
FERROUS AND FERRIC IONS
Reference, in this and later Assignments : A standard text on Inorganic Chem-
istry. Note. Read again Asignment 31, Paragraph 2, and Assignment 38.
1. In Assignment 43 we shall study the chemistry of iron as an example
of a metal which forms two series of compounds in which the valence
of the metal is different. You will find on the laboratory shelves solutions and
solid salts of ferrous and of ferric iron. Give the valences of the following:
sulfate ion; sulfate in ferrous sulfate, FeSO 4 ; Fe (metal); ferrous ion, Fe ++ ;
ferric ion, Fe +++ ; iron in a ferrous salt, and in a ferric salt. Write the formulas
of the chlorides, nitrates, hydroxides, and oxides of ferrous and of ferric iron.
2. Many double salts are found among iron compounds.
FeSO 4 '(NH 4 ) 2 SO 4 ' 6H 2 O, ferrous ammonium sulfate, and
Fe 2 (SO 4 ) 3 '(NH 4 ) 2 SO 4 ' 24H 2 O, ferric ammonium sulfate (ammonium
iron alum) are examples of double salts whose solutions dissociate into the ions
of the simple salts. These two salts may be used in studying the properties of
Fe ++ and Fe +++ , respectively. Give the formulas of : alum, potassium iron
alum, sodium iron alum.
3. The ferrocyanides and ferricyanides are examples of double salts
whose solutions contain complex ions at high concentrations. The formula,
Fe(CN) 2 4NaCN, for sodium ferrocyanide shows its relation to the simple
salts. It can be prepared by dissolving the difficultly soluble ferrous cyanide,
Fe(CN) 2 , in excess of NaC'N, but in a solution of the pure salt the concentra-
tions of Fe ++ and CN~ are extremely small. Accordingly the formula is usually
written Na 4 Fe(CN) 6 to show that the ferrocyanide radical Fe(CN) is the
negative constituent of the salt. Similarly Na 3 Fe(CN) 6 is the formula
usually written for sodium ferricyanide. Questions. What is the valence
of the ferrocyanide radical? Of the ferricyanide radical? What is the
formula of potassium silver cyanide? How is it prepared? What ions are
present in its solution? Caution. Soluble cyanides and hydrocyanic acid are
deadly poisons : do not try unnecessary experiments with them. Ferro and
ferricyanides are not dangerous.
[561
4. Color tests for Fe ++ and Fe + + + . Experiment. To a dilute solution of
FeCl 3 add a drop or two of K 4 Fe(CN) 6 solution. Repeat with K 3 Fe(CN) 8 .
Prepare a solution of a ferrous salt by dissolving a small amount of FeSO 4
or of FeSO 4 *(NH 4 ) 2 SO 4 in a little dilute sulfuric acid, and test portions of this
solution at once with K 4 Fe(CN) c and with K 3 Fe(CN) 6 . The character-
istic dark blue precipitates, or colloidal solutions, obtained in two of these
four experiments are probably identical. Equations may be omitted for these
reactions.
5. Oxidation of Fe ++ to Fe +++ . Experiment. Test for Fe +++ an acid solu-
tion of FeSO 4 which has been standing In the laboratory for several days.
The oxygen of the air slowly oxidizes Fe ++ to Fe +++
4Fe ++ + O 2 + 4H + = 4Fe +++ + 2H 2 O.
Experiment. To a few drops of a ferrous salt solution in a porcelain dish add
a little HNO 3 , heat the mixture, dilute with water, and test whether the solu-
tion now contains Fe" 1 " 1 " or Fe +++ . Try an experiment with another oxidizing
agent and determine whether or not it oxidizes Fe ++ to Fe ++ +.
6. Note. Write equations for all reactions.
7. Reaction between Fe and Fe +++ . Experiment. Place some iron wire
or iron filings in a very dilute solution of FeCl 3 , stir the mixture, and test
small portions of the solution from time to time for Fe ++ and Fe +++ .
8. Reactions of iron ivith H + , and ivith Cu^. Experiment. Try the action
of dilute HC1 or H 2 SO 4 on a small piece of iron; heat the mixture, if neces-
sary. Predict whether ferrous or ferric ion is formed, and test the correctness
of your answer. Experiment. Cover a small piece of iron with a solution of
a copper salt.
9. From the results of the above experiments, or from any other source of
information, determine approximately the position of Fe Fe ++ and of Fe ++ Fe +++
in your table of oxidizing and reducing agents, Assignment 42, Paragraphs
2 and 7. It is not advisable to attempt to assign a definite place in one large
table to each of the oxidizing and reducing agents which you will study. Thus
in the case of Fe ++ Fe +++ it is sufficient to note that Fe ++ Fe +++ , I~ 1 2 and
Ag Ag + are in the same part of the table: Fe ++ , I~ and metallic silver are all
readily oxidized by powerful oxidizing agents, while Fe + + + , I 2 and Ag + are
easily reduced by powerful reducing agents. Questions. Which is the stronger
reducing agent, Fe" 1 " 1 " or Fe? Which is the stronger oxidizing agent, Fe* ++
or Fe ++ ?
10. Reduction of Fe +++ by H 2 S in acid solution. Experiment. To a few
drops of FeCL solution add 1 cc. 6 N HC1 and about 20 cc. water, and saturate
the solution with H 2 S gas. Boil the mixture in a porcelain dish to coagulate
the precipitated sulfur and expel the H 2 S, filter and at once test a portion of the
clear solution for Fe +++ . Treat the remainder of the solution again with H 2 S if
Fe +++ seems to be present. Name another reducing agent which might be
expected to reduce Fe +++ to Fe ++ and try the experiment.
11. Other reactions of Fe +++ . Experiments. Treat small quantities of dilute
solutions of FeCl 3 with the following reagents:
(a) NH 4 OH in small amount and in excess. The precipitate is Fe(OH) s .
(b) NaOH in small amount and in excess.
(c) Na 2 CO 3 or (NH 4 ),CO 3 . Filter; acidify the filtrate and test it for Fe +++
with K 4 Fe(CN) 6 ; wash with water the precipitate (obtained by adding
carbonate), and test whether it is a carbonate or hydroxide.
(rf) NH 4 OH or NaOH and then a sulfide as H 2 S or (NH 4 ) 2 S. The black
precipitate is mainly Fe 2 S 3 , but may contain some FeS and S.
[57]
12. Other reactions of Fe ++ . Repeat the experiments in (11) with small
quantities of dilute solutions of ferrous salt. In (a) and (b) note what
happens when moist ferrous hydroxide is exposed to the air; the oxygen
reacts much more rapidly with moist Fe(OH) 2 than with Fe ++ in acid solution.
The brown compound formed in the test for nitrate ion, Assignment 31, Par-
agraph 13, is an addition compound of NO and ferrous ion, probably Fe(NO) ++ .
This compound is formed when NO gas is passed into a cold ferrous salt solution,
and when nitric acid, nitrous acid, or NO 2 is reduced by excess Fe ++ . Questions.
What happens when a solution which contains this addition compound is heated?
Is the rate of reduction of nitrate ion in dilute acid solution by Fe ++ a rapid or a
slow reaction? Experiment. Determine whether the reduction of nitrous acid,
HNO 2 , is a rapid or a slow reaction by adding to 10 cc. water, 2 cc. FeSO 4
solution, a few drops oi 6 N H 2 SO 4 , and a drop of a nitrite solution.
13. Summarize in a table for future reference the reactions of Fe ++ and
Fe +++ with the following reagents: a strong oxidizing agent, as hot moderately
concentrated HNO 3 ; a strong reducing agent, as zinc in acid solution; H 2 S
in 0.3 N acid solution, and the reagents used in Paragraphs 11 and 12. Below the
table give methods of dissolving the difficultly soluble salts of iron which you
have studied.
14. Problems. Try experiments when necessary.
(/) How would you transform
(a) ferric chloride to ferric sulfate;
(b) ferric sulfate to ferric chloride;
(c) ferrous sulfate to ferric sulfate;
(d) ferric sulfate to ferrous sulfate?
(2) Fe(OH) 2 is moderately soluble in a solution of an ammonium salt,
as NH 4 C1, and Fe(OH) 3 is not. Outline an experiment to demonstrate this.
Which is the less soluble in water, Fe(OH) 2 or Fe(OH) 3 ?
(j) Suggest a method of preparing Fe(OH) 3 and ZnS from a single portion
of a solution which contains Fe ++ , Zn ++ , H + and SO 4 ~~.
ASSIGNMENT 44
MERCUROUS AND MERCURIC IONS
1. In the present Assignment we shall study the chemistry of another metal
which forms two series of compounds. You will find on the laboratory shelves
solutions of mercurous nitrate, HgNO 8 ; and of mercuric nitrate, Hg(NO 3 ) 2 .
What is the valence of mercury in a mercurous salt; in a mercuric salt? Write
the formulas of the chlorides, sulfates and oxides of mercurous and of mer-
curic mercury.
2. Mercurous chloride and mercuric chloride. Experiment. To 1 cc.
N HgNO 3 add 20 cc. water and 1 cc. 6 N HC1. Repeat the experiment with
N Hg(NO 3 ) 2 , and give a method of distinguishing between soluble mercurous
and mercuric salts. Note. Mercuric chloride is one of the few examples of a
weak salt. The concentrations of Hg ++ and of Cl~ in its solution are very small.
3. Reduction of mercuric ion in stages. Tin forms two series of compounds
stannous and stannic, in which the valence of tin is -\- 2 and + 4 respectively.
Stannous ion, Sn ++ , is a good reducing agent since it is easily oxidized to Sn +++ *.
The chlorides of tin are both soluble in water; and the common stannous salt
laboratory reagent is SnCl 2 solution to which HC1 has been added to prevent
[581
the precipitation of a basic salt. Experiment. To 1 cc. N Hg(NO 3 ) 2 and
20 cc. water add one drop SnCL solution. Divide the mixture into two portions
and to one portion add SnCl 2 until it is present in excess. Observe the differ-
ence in appearance of finely divided mercury and a large drop of the metal.
Suggest a method of testing a solution for the presence of mercuric mercury.
Try your method with a dilute solution of HgCl 2 .
4. Reaction between mercuric ion and mercury. Experiment. To 1 cc.
N Hg(NO 3 ) 2 add a drop of mercury, shake the mixture gently, and after
several minutes add 20 cc. water. Pour off the solution and test it for Hg +
with HC1. Filter off the mercurous chloride and test the filtrate for mercuric
mercury. If there is more than a mere trace repeat the experiment and leave
the mercury in contact with the solution for a long time. The reaction be-
tween Hg and Hg ++ is similar to that between Fe and Fe +++ . In each case, in
the presence of excess of the metal, the reaction will proceed until at equi-
librium the concentration of the ion of higher valence, Hg ++ or Fe +++ , is ex-
tremely small. Questions. Which is the stronger oxidizing agent, Kg** or
Hg + ? Which is the stronger reducing agent, Hg + or Hg?
5. Reduction of Hg + and oxidation of Hg. By experiments with some of
the metals you have studied and small amounts of HgNO 3 solution determine
approximately the position of Hg Hg* in your table of oxidizing and reduc-
ing agents. In these experiments any metallic mercury formed can usually be
seen on the surface of the metal, but the solution may also be tested for the
dissolved metal after precipitating the mercurous mercury with HC1. From
your result predict whether metallic mercury will dissolve in dilute sulfuric
acid , and test your prediction. Experiment. Treat a very small drop of mer-
cury with about 30 cc. 2 N HNO 3 in a porcelain dish or beaker. If the reac-
tion is slow heat the mixture. While the reaction is in progress pour a small
amount of the solution into 10 cc. water and test for Hg + and Hg ++ . When
a moderate amount of the metal has dissolved, pour off the solution for use
in the next experiment.
6. Oxidation of Hg + and of HgCl. Experiment. Continue the preceding
experiment to determine how rapidly hot 2 N HNCX oxidizes Hg + . Repeat
the experiment with hot concentrated HNO 3 and a small quantity of HgNO 3 .
Experiment. Collect some HgCl on a filter. Treat a small portion in a porce-
lain dish with hot 2 N HNO 3 , and another portion with hot 2 N HNO 3 , to
which some HC1 has been added. Repeat the experiments, if necessary, with
more concentrated HNO 3 , and note whether the speed of oxidation is increased
by the presence of HC1. Experiment. Treat a small portion of HgCl with
bromine water, and test the solution as before for mercuric mercury.
7. Difficulty soluble salts of mercuric mercury. Experiment. Test the
action of each of the following reagents on a dilute solution of Hg(NO 3 ) 2 . Save
the precipitate for use in (8).
(a) H 2 SO 4 or a soluble sulfate.
(b) Na^COg. The precipitate is a basic carbonate.
(c) NaOH. The precipitate is HgO.
(d) NH 4 OH. When a chloride is present the precipitate is Hg(NH 2 )Cl,
ammonobasic mercuric chloride, which is closely related to Hg(OH)Cl,
basic mercuric chloride. Analogous but more complex substances are pre-
cipitated in the absence of chloride. What is the valence of the OH radical?
Of the NH 2 radical?
(e) H 2 S in- acid solution.
8. Predict the action of dilute HC1 or HNO 3 on each of the precipitates just
obtained (except HgS), and test your answers by experiments.
[59]
9. Experiment. Collect some HgS on a filter and wash it with water.
Transfer part of the precipitate to a porcelain dish, add 2 TV HNO 3 boil the
mixture for about a minute, and filter. Test the filtrate for Hg*" 1 ". Repeat the
experiment with hot 2 N HNO 3 to which some HC1 has been added. From the
fact that HgQ 2 is a weak salt suggest an explanation for the result of this
experiment. Repeat the experiment with bromine solution instead of HNO 3 .
10. Other reactions of mercurous ion. Experiment. Treat a dilute solution
of HgNO 3 with each of the following reagents :
(a) H 2 SO 4 or a soluble sulfate.
(b) Na 2 C0 3 .
(c) NaOH. The precipitate is Hg 2 O, which, on heating, decomposes into Hg
and HgO.
(d) NH 4 OH. The precipitate is a mixture of metallic mercury and an am-
monobasic mercuric salt; see 7 (d).
(e) H 2 S in acid solution. The precipitate is a mixture of Hg and HgS.
11. From the results of your experiments in Paragraphs 6, 8, and 9, give
methods of dissolving each of the precipitates obtained in 10. Try experiments
whenever you are doubtful of your method.
12. Table of reactions for mercurous ion and mercuric ion. Summarize in
a table for future reference the reactions of Hg* and Hg ++ with the following
reagents: HC1; a strong oxidizing agent as Br 2 a strong reducing agent as
SnCl 2 (in small amount and in excess), and the reagents used in Paragraphs 7
and 10. Below this table classify the precipitates according to reagents used in
dissolving them and point out the considerations that led to the choice of each
reagent. Be sure that you can write the equation for each reaction.
13. Problems. Try experiments when necessary.
(/) Complete the following equations, and indicate what you would observe
in the corresponding experiments:
Sn+ + + HgCl 2 (excess) =
Sn + + (excess) + HgCl 2 =
(2) What reactions are possible between a sulfide and hydrochloric acid?
Between a sulfide and nitric acid? In connection with your answers discuss
the following experimental facts; CuS dissolves readily in hot 2 N HNO 2 but
not in hot 2 N HC1.
(j) Compare the action of hot 2 A r HNO 3 on CuS (Assignment 37 Paragraph
10) and HgS, and state what conclusion can be drawn with regard to the relative
solubilities of these two sulfides in water. Outline a method of preparing pure
HgCl 2 and pure CuSO 4 from a mixture of the freshly precipitated sulfides.
(4) Suggest methods of making the following preparations:
(a) HgO fromHg(N0 3 ) 2 ;
(b) Hg(NO 3 ) 2 from HgSO 4 ;
(c) Hg(N0 3 ) 2 fromHgN0 3 ;
(d) HgNO 3 from Hg(NO 3 ) 2 .
Note. It is recommended that Assignment 52 on Qualitative Analysis, follow
Assignment 44.
60
ASSIGNMENT 45
LEAD ION, CHROMATE ION
1. In the present Assignment and the following ones we shall continue the
study of important common elements. Only very brief laboratory directions
will be given; each student is expected to plan the details of the experiments,
and is recommended to try additional experiments which are suggested by the
lectures or laboratory work.
2. Chromate ion, CrO 4 ~~, and bichromate ion, Cr,O 7 ~~. Note the colors
. of solutions of sodium (or potassium) chromate and bichromate. What is
the color of CrO 4 ~~? Of Cr 2 O r ~~? What is the valence of chromium in
these ions? Experiment. Add dilute HNO 3 or H 2 SO 4 to a small quantity
of a chromate solution. Add dilute NaOH or NH 4 OH to a bichromate so-
lution; acidify the solution obtained. Note. Many chromates are difficultly
soluble in water, but dissolve in HNO 3 .
3. Reactions of lead ion, Pb ++ . Experiment. Treat small portions of
0.1 A 7 Pb(NO 3 ) 2 with each of the reagents listed in Assignment 38 (Par-
agraph 5), with H 2 SO 4 , and with a soluble chromate.
4. Methods of. dissolving difficulty soluble lead salts. State which of the pre-
cipitates obtained in (3) you would expect to dissolve in 0.2 N HNO 3 . Experi-
ment. Prepare some PbS, collect it on a filter, and test portions of it with hot
2 N HC1, and with hot 2 N HNO 3 . Experiment. Collect some PbCl, on a
filter, pour cold water over it and test the filtrate for Pb* + , e. g. with H 2 S;
then pour hot water over PbCl 2 on a filter and test the filtrate for Pb ++ .
Experiment. Collect some PbCrO 4 on a filter, wash thoroughly with water
to remove any soluble chromate, and determine how readily it dissolves in hot
dilute HNO 3 . What ions are present in the solutions obtained? Experiment.
Collect some PbSO 4 on a filter, transfer some of it to a test tube, and pass in
H 2 S. Which is the less soluble, PbSO 4 or PbS? Experiment. Treat some
PbSO 4 with Na 2 CO 3 solution, as directed in Assignment 52, 11 (5).
5. Relative delicacy of tests for Pb + +. Experiment. Prepare 50 or 100 cc.
portions of 0.01 N, 0.001 N, 0.0001 N Pb(NO 3 ) 2 . Test 5 or 10 cc. of each
solution with the reagents used in (3). Specify the two most delicate tests
Name the two least soluble salts of lead.
6. Lead dioxide, PbO 2 . Experiment. Treat a small portion of PbO 2 with
hot 6 N HC1. What gas is evolved? Repeat the experiment with HNO 3 instead
of HC1.
7. Problems. (/) Outline methods for making the following transformations:
(a) PbSO 4 to PbS. (c) PbCO 3 to PbCrO 4 .
(b) PbS to PbSO 4 . (d) PbO 2 to Pb(NO 3 ) 2 .
(2) Name three amphoteric hydroxides, (cf. Assignment 35). What would
be observed on gradually acidifying a solution of sodium plumbite?
ASSIGNMENT 46
STANNOUS AND STANNIC IONS. AMPHOTERIC SULFIDES
Reference. Hildebrand, pages 197-199.
1. Ammonium polysulfide solution. Experiment. Prepare some colorless
ammonium sulfide solution by passing H 2 S into NH. t OH solution. Divide the
solution into three parts. To one portion add a very small amount of powdered
[61 1
sulfur. Place another portion on a watch-glass to expose it to the action of the
oxygen of the air. Acidify the third portion with dilute HC1. Experiment.
Acidify small quantities of the laboratory solutions of ammonium polysulfide
(yellow ammonium sulfide), and of ammonium sulfide. Note. On account of
hydrolysis the (NH 4 ) 2 S solution is really NH 4 OH + NH 4 SH.
2. Reactions of Sn ++ and Sn ++++ . . Experiment. Prepare 0.1 N solutions
of stannous chloride and of stannic chloride, and determine the action of the
following reagents on each of these solutions :
(a) H 2 S in cold 0.3 N HC1 solution. These tests are delicate and characteristic.
(b) H 2 S in hot 2 N HC1. If there is a precipitate, filter, add water to the filtrate
and test it for tin.
(c) NH 4 OH in small amount and in excess.
(d) NaOH in small amount and in excess. If there is a precipitate after adding
NaOH in excess, filter, acidify the filtrate with HC1 or H 2 SO 4 and test it for
tin.
0) (NH 4 ) 2 S after NH 4 OH has been added.
(/) Ammonium polysulfide after NH 4 OH has been added. Pour through a filter
to remove any precipitate that may remain undissolved, and acidify with
dilute HC1.
(g) A soluble carbonate. Filter, wash the precipitate and test it for carbonate.
(h) A good oxidizing agent in dilute acid solution, such as bromine.
(i ) A good reducing agent in dilute acid solution, as zinc.
(/ ) HgCl 2 solution in small amount and in excess.
3. Action of concentrated HNO^ on Sn. Experiment. Treat a small piece
of tin with concentrated HNO 3 in a casserole. Evaporate the mixture just
to dryness: to avoid heating the residue strongly, evaporate the last portions
of the liquid by moving the casserole back and forth through a small flame.
Warm the residue with dilute HNO 3 , filter, and test the filtrate for tin. Note.
Stannic acid, or hydrated stannic oxide, when prepared in this way by the
action of concentrated HNO 3 on Sn or on a compound of tin, is difficultly
soluble in HNO 3 . Experiment. Suggest and try methods of dissolving this
substance.
4. Problems, (i) If a solution were known to be either SnCL or SnCl t how
would you identify it?
(2) How would you test for SnQ 2 in a solution of SnCl 4 ?
(j) What happens when HC1 is added gradually to a solution of (a)
sodium stannite, (b) sodium stannate, (c) sodium stannite after it has been
treated with a powerful oxidizing agent, as sodium peroxide, Na 2 O 2 , or hypo-
chlorite, (d) sodium sulfostannate?
(4) How would you prepare (a) ammonium polysulfide, (b) sodium sulfo-
stannate from SnS 2 , (c) sodium sulfostannate from SnS?
ASSIGNMENT 47
IONS OF ALUMINUM AND OF CHROMIUM. PEROXIDES
1. Write the formulas of aluminum chloride, nitrate and sulfate; chromium
chloride, nitrate and sulfate; alum and chrom alum. These salts are all readily
soluble in water and their solutions react acid to litmus.
2. Hydroxides. Experiment. To very dilute solutions of an aluminum
salt and of a chromium salt add NH 4 OH a few drops at a time until, after
shaking, the odor of ammonia is just perceptible. Warm the mixture. These
are delicate tests for Al +++ and Cr^*.
T621
3. When a solution of an aluminum, or chromium, salt is treated with a
soluble carbonate or sulfide, the hydroxide of aluminum, or chromium, is pre-
cipitated. Explain.
4. Aluminate ion and chromite ion. A1O 2 ~ and CrCX" (or H.AICX') and
H,CrO 3 ~). Experiment. To dilute solutions of an aluminum salt and of a
chromium salt add a few drops of NaOH solution. To each mixture add
more NaOH and shake after each addition. Filter if a precipitate remains,
add some solid ammonium salt, and heat the mixture.
5. Ions of Chromium. State which ions of chromium can exist only in an
acid solution and which in an alkaline solution. What is the valence of
chromium in each of these ions? Suggest a method of changing bichromate
ion to chromic ion.
6. Reduction of Cr 2 O 7 ~~ and of CrO.,~~. Experiment. Test the action
of H 2 S on very dilute solutions of a bichromate in the presence of sulfuric
acid, heat the mixture to boiling, filter, and wash the precipitate. Make the
filtrate alkaline with NH 4 OH. Repeat the experiment with H 2 S and a very
dilute solution of a chromate in the presence of NaOH, heat, filter and wash
the precipitate. Acidify the filtrate. Experiment. Determine the action of hot
concentrated HC1 on Cr 2 O 7 ~~, and of Cl~ on CrO 4 ~~ in alkaline solution.
7. Oxidation of chromite ion. Experiment. Prepare a small quantity
of a chromite solution. To a portion of this add bromine solution; boil the
solution to expel any bromine that may remain. Question. How would you
prepare a hypochlorite solution? Write the equation for the oxidation of
CrO 2 ~ by C1O~. Experiment. Boil a small quantity of a chromic salt solution
with concentrated HNO 3 , or with aqua regia.
8. Peroxides. Sodium peroxide decomposes rapidly, even in a cold solution.
NiO HO = 2NaOH
Hydrogen peroxide in a cold, dilute acid solution decomposes only very
slowly. Na 2 O 2 is a very powerful and useful oxidizing agent. It reacts with
chromic hydroxide suspended in water, as follows :
2Cr(OH) 3 + 3Na 2 O 2 = 2CrO 4 " + 6Na + + 2OH~ +2H,O.
Experiment. Test the action of Na 2 O 2 on CrO 2 ~. The evolution of oxygen
observed in this experiment is due to the decomposition of some of the
peroxide.
9. In acid solution H 2 O 2 may be reduced by a powerful reducing agent, as
HoS; but it may also be oxidized by a powerful oxidizing agent, as Cr 2 O 7 ~~.
Cr 2 O 7 " + 3H 2 O 2 -f 8H + = 2 Cr +++ +3O 2 + 7H 2 O.
10. Problems. (/) Complete the following equations and state what would
be observed in the corresponding experiments:
Cr(OH) 3 + H+ (excess) = ....................................................
CrO 2 - + H + (excess) = ....................................................
CrO 4 - + H + (excess) = ....................................................
Cr 2 r ~ + H 2 S0 3 + H+ = Cr+" ............................................
H 2 O 2 + MnO 2 + H+ = Mn" ............................................
H 2 2 + I- + H+ = I, .................................................
(2) State how you would prepare Cr 2 O r ~~ from Cr +++ , Cr(OH) 3 from PbCrO 4 .
(j) Concentrated NH 4 OH solution has a slight solvent action on Al(OH)g
and Cr(OH) 3 , but these hydroxides can be reprecipitated by adding an ammon-
ium salt and warming the mixture. Explain.
(4) List in a vertical column all the amphoteric hydroxides you have studied,
and write in two more columns the formulas of the corresponding ions which
[63]
are formed when the hydroxide is dissolved in. excess of a solution of (a) a
strong acid, and (b) a strong base.
(5) Examine the position in the Periodic System (Hildebrand, page 257) of
each of the elements considered in the preceding problem, and write a brief
note on the gradations of properties in each family in which one of these
elements occurs. Name additional elements which might be expected to have
amphoteric hydroxides.
64
SECTION V
QUALITATIVE ANALYSIS
ASSIGNMENT 51
THE DEVELOPMENT OF A SCHEME OF ANALYSIS FOR A LIMITED NUMBER OF
POSITIVE IONS
(To follow Assignment 38)
1. In Assignment 38 the chemistry of the positive ions already considered
was summarized, and the reactions of these ions with various reagents tabulated.
We shall now consider how this information may be used in analyzing an
unknown solution for these ions.
2. A possible method of analysis would be to devise a separate procedure
for each ion, then, if seven ions were to be considered, seven separate portions
of the original solution (or substance) would be taken, and each treated by an
experiment or series of experiments devised to show whether the particular
substance is present or absent in the unknown. Such tests were used in Assign-
ment 31 for Cl~, SO 4 ~~, and NO 3 ~; in these cases provision was made for
neutralizing OH~, because the presence of OH" would have interfered with or
spoiled the tests. It is to be noted also that the test used for Cl~ is not specific
for Cl~, since any other silver salt which is insoluble in dilute HNO 3 , as silver
iodide, would also be precipitated. It is evident that the application of this
method (of testing separate portions of the unknown for a particular substance)
becomes increasingly difficult as the number of substances which may be present
is increased. Also a lengthy procedure is often necessary, especially in testing
for certain positive ions when ions of several other metals are present.
3. Accordingly, in systematic qualitative analysis the method has been adopted
of using only one portion of the unknown solution in testing for metals, and
of preparing from it a pure compound of each metal. Such a process serves
to identify the positive constituents and also enables a comparison of the amounts
of the various metals to be easily made. The reagents ordinarily used in this
process are acids, NH 4 OH and NH 4 salts; and therefore separate portions of
the unknown are used in testing for H + , NH 4 + , and the various negative ions.
(The development of a systematic procedure for the detection of negative con-
stituents will not be considered in this course.)
4: As an illustration we shall discuss the preparation of pure salts of zinc and
calcium from an acid solution containing Zn ++ and Ca ++ . The Tables in Assign-
ment 38 show the following differences which may be utilized in devising methods
of separation :
(/) CaS is soluble in water and ZnS is difficultly soluble. The latter can
be precipitated, after the solution has been made alkaline with NH 4 OH, by add-
ing (NH 4 ) 2 S solution drop by drop, or better by passing in H 2 S gas. When
the ZnS has been completely precipitated, and removed by filtration, calcium
can be tested for in the filtrate and any desired salt of calcium can be pre-
pared. Experiment. Test this method by the Procedure given below, with an
acid solution which contains both Zn ++ and Ca ++ , prepared for example by adding
to 100 cc. water in a flask 2 or 3 drops of ZnSO 4 , CaCL and HC1 solutions.
Repeat, omitting the calcium salt, in order to determine if a satisfactory blank
is obtained in the calcium test. Procedure. (Be sure that you understand the
reason for each detail in the directions.) Add NH 4 OH, a few drops at a time,
until the mixture after shaking has a distinct odor of NH 3 , pass in H 2 S, shake
[651
the mixture thoroughly, and cautiously note the odor to determine if H 2 S has
been added in excess; pass in more H 2 S if necessary. (If a large precipitate
is obtained, test with litmus and if the solution has become acid add more
NH 4 OH; again pass in H 2 S until it is present in excess.) Warm the mixture
to 50 or 60, and filter. Test the clear filtrate for calcium by adding (NH 4 ) 2 CO 3
solution, and heating the mixture almost to boiling. Note. If in the analysis
of an unknown solution no precipitate is observed in the test for Zn ++ or Ca ++ ,
let the warm solution stand for several minutes before concluding that the ion
is absent.
(2) Zn ++ forms a complex ion with ammonia and Ca ++ does not. Both
carbonates are difficultly soluble in water, but ZnCO 3 , unlike ZnS, dissolves in
NH 4 OH. Accordingly CaCO 3 can be precipitated by treating the mixture with
NH 4 OH in excess and (NH 4 ) 2 CO 3 . Write out a Procedure (cf. preceding Par-
agraph) and note the reason for each of the experimental details.
(5) Zn ++ is converted into zincate ion by excess of a strong base. Pure
NaOH solution would precipitate some of the calcium as Ca(OH) 2 but
Na 2 CO 3 must also be added to make a satisfactory separation. The laboratory
NaOH solution contains some CO 3 ~~ but it is better to add Na 2 CO 3 .
5. Plan three methods of separating silver and calcium, i. e. of preparing
pure salts of silver and of calcium from a mixture of their nitrates. Write
out each method, stating why you think it will work; and show your note-
book to the instructor.
6. By combining your methods of separation, Paragraphs 4 and 5 devise
a "scheme of analysis" for these three positive constituents, silver, zinc and
calcium. Test your scheme by experiments with very small quantities of solu-
tions of (a) AgNO 3 alone, (b) ZnSO 4 alone, and (c) CaQ 2 alone. In each
experiment with a pure salt a satisfactory test should be obtained for one metal
and blanks in the tests for the other two.
7. Experiment. Predict what will happen when a small quantity of CuSO 4
solution is treated by your scheme of analysis, Paragraph 6, and test your answer
by an experiment. Devise a method of including copper in your scheme, and
test your method by experiments.
8. Assume that a sodium salt is present in solution with salts of silver,
copper, zinc and calcium, and consider how all the sodium originally present,
and no more, could be recovered as a pure salt. In other words, include sodium
in your scheme. Show your scheme of analysis for the five positive constit-
uents to the instructor. Question. With which of the five metals would
potassium have been found if it had been present in the original solution? - .
9. When an unknown solution has been analyzed in this way and a result
is inconclusive, owing for example to the small size of the precipitate, a charac-
teristic confirmatory test is made whenever possible, such as the dissolv-
ing of AgCl in NH 4 OH and the reprecipitation with excess HNO 3 , the flame
test for calcium, etc.
10. When the unknowns are solid salts an important part of the work is
the preparation of the solution for analysis. Preliminary experiments with small
quantities are performed to determine a method of dissolving the unknown, and
conclusions with regard to its nature can often be drawn from the results of
these experiments. The salts given out for analysis at this time will dissolve
either in water, in dilute acids or in concentrated acids. . A discussion of
methods of dissolving difficultly soluble salts will be given in Assignment 52.
11. Review your notes in the previous Assignments on the tests for
Na + , K + , NH 4 + , C1-, SO 4 ", CCX" and NO 3 -. Devise a method of testing for
[661
S" based on the formation of H 2 S gas, and determine whether your test will
apply in the case of a very difficultly soluble sulfide, as CuS.
12. Analyses 1, 2 and 3. (Note. Do not begin these analyses until the
instructor has approved your scheme of analysis for the positive constituents,
Paragraphs 6 to 8. ) Test for the positive and negative constituents :
H, Ag, Cu, Zn, Ca, Na, K, NH 4
OH, Cl, N0 3 , S0 4 , C0 3 , S.
Record at once each experiment and observation and write down any con-
clusion that has bearing on the identification. Try to distinguish between large
amounts, small amounts, and traces of the different constituents. When you
report the result of an analysis to your laboratory instructor show him your
notebook record.
13. Problems. (/) What negative ions need not be tested for in solutions
which contain the. following positive ions: (a) H + . (b) Ca + *, (c) Zn ++ , (d)
Cu ++ , (e) Ag*?
(2) Point out the errors in the following reports:
(a) A solution known to contain Ag + was reported to contain SO 4 ~ '
because BaCL gave a white precipitate which did not dissolve in
HN0 3 .
(b) A pure sodium salt which dissolved in water to give an alkaline
solution was reported to be a sulfate because a white precipitate
was obtained on the addition of Ba(NO 3 ) 2 -
(c) A pure salt which did not dissolve appreciably in water was reported
to be zinc nitrate.
ASSIGNMENT 52.
THE STANDARD SCHEME OF ANALYSIS. METHODS OF DISSOLVING DIFFICULTLY
SOLUBLE SUBSTANCES
(To follow Assignment 44)
1. In Assignment 51 a number of schemes were considered for the systematic
analysis of solutions containing Ag + , Cu" + , Zn ++ ,Ca ++ and Na + (and K + ). One of
the most satisfactory schemes, and the one used in practically all text books,
is the following:
To the solution add HC1
Precipitate
AgCl
Filtrate (concentration of H + , 0.3 A r ) : Saturate with H 2 S gas
Precipitate
CuS
Filtrate : Boil to expel H S, * add NH 4 OH and
(NH 4 ) 2 S.
Precipitate
ZnS
Filtrate: Add (NH 4 ),CO 3 and warm
Precipitate Filtrate : Evaporate to dry-
CaCO 2 ness,
[eat to expel NH 4 salts.
Residue: Na (and K)
salts.
* Note. It is really not necessary to expel the H 2 S ; but when* additional positive con-
stituents are included in the scheme, and several substances may be precipitated at this
point, the observation of the effect of adding NH 4 OH in the absence of sulfide often assists
in identifying the positive constituents.
[67]
2. Iron and mercury in the scheme of analysis. From the results of your
experiments on the reactions of the ions of iron and mercury, Assignments 43
and 44, predict what would be observed if a solution containing each of the
following ions was treated by the scheme of analysis, Paragraph 1 :
Fe ++ , Fe +++ , Hg + , Hg ++ . Experiment. Verify your answers by experiments with
dilute solutions of each of these ions. Be sure that the blank tests for the other
positive constituents, e. g. Ca ++ , are negative. Question. What mistake has been
made if in the experiment with a mercuric salt a black precipitate is obtained
with sulfide in the NH 4 OH solution?
3. Separation of ZnS and FeS. Plan a method of dissolving a mixture of
ZnS and FeS and of obtaining pure Fe(OH) 3 and ZnS from the resulting solution
(which contains Zn ++ and Fe ++ ). Experiment. Test your method with small
amounts of (a) ZnS, (b) FeS, and (c) a mixture of ZnS and FeS. See Assign-
ment 51 Paragraph 4 (/) on the method of percipitating ZnS.
Separation of CuS and HgS. Experiment. Prepare a small amount, of (a)
CuS, and (b) a mixture of CuS and HgS. Treat each of these by the procedure
devised in Problem (j), Assignment 44. Questions. Why is the sulfur residue
black when CuS alone is used, (cf: Assignment 37, Paragraph 10)? Is a
confirmatory test for mercuric mercury necessary? For such a test see Assign-
ment 44, Paragraph 3.
Separation of AgCl and HgCl. Devise and confirm by experiment a method
of determining whether a chloride precipitate is either AgCl or HgCl. If a
large amount of HgCl is found and only a small amount of AgCl is present,
the method based on the use of NH 4 OH solution fails on account of the reduction
of Ag + by the metallic mercury. Devise a method of detecting silver in the
presence of a large amount of HgCl, based on the treatment of the mixture
obtained on the addition of NH 4 OH, with an oxidizing agent such as bromine
solution. Experiment. Try your method.
4. The other common positive elements are grouped as follows with the
elements already considered.
Division of Metals into Groups for Qualitative Analysis.
Characteristic Reagents for Precipita- Substances Precipitated.
Metal tion.
Name of Group.
I. Silver
Silver Group.
II. Copper
Copper Group.
[II. Zinc
Zinc Group.
IV. Calcium
Calcium Group.
HC1 or NH 4 C1 and an AgCl, HgCl, PbCl 2 par-
tially, BiOCl partially.
acid.
HS in 0.3 N acid.
NH 4 OH and (NH 4 ) 2 S
or NH 4 OH in excess
and H,S.
CuS, PbS, HgS, CdS,
Bi 2 S 3 , Sulfides of As,
Sb, Sn.
ZnS, A1(OH) 3 , Cr(OH) s
FeS, Fe 2 S 3 , MnS,NiS,
CoS.
(NH 4 ) 2 CO 3 andNH 4 OH. CaCO 3 , BaCO 3 , SrCO 3 , a
compound of rnag-
nesium when alcohol is
present,
none.
V. Sodium
Sodium Group. ,
5. The further separation of the elements has already been illustrated in
Paragraphs 2 and 3, but only for two elements in each of three groups. The task
F681
of developing a scheme of analysis becomes increasingly difficult as the number
of elements is increased, mainly on account of the large number of time-con-
suming experiments which must be made with known mixtures in order to
prove that the proposed methods are satisfactory. Therefore if you ever have
to make an accurate qualitative analysis of a complicated mixture it will be
necessary to consult a standard text-book, such as A. A. Noyes' Qualitative
Chemical Analysis, and to follow carefully the directions there given. Before
attempting to make an analysis with the aid of a book, work through the pro-
cedures with known solutions and study the chemistry involved.
6. The present course, however, is not primarily a course in Qualitative
Analysis, and a text-book will not be necessary. In the time which remains at
our disposal only a limited number of additional elements can be studied in the
laboratory. These will be investigated by the same methods as before, and
from the tables of reactions it will be possible to develop satisfactory methods
for the analysis of simple mixtures.
7 . Analyses Nos. 4, 5 and 6. In addition to the positive and negative con-
stituents considered in Assignment 51, test for iron, mercurous mercury and
mercuric mercury. Do not begin these analyses until the instructor has
approved your methods of analysis, Paragraphs 2 and 3. Also study carefully
the following paragraphs.
8. Analysis of an unknown solid. Record the appearance of the solid. Treat
a small portion with water and heat the mixture in a porcelain dish or beaker
if the solid does not dissolve readily. Always test the solution with litmus paper.
If the solution is neutral, determine if anything has dissolved, either by slowly
evaporating a few drops of the clear solution on a watch-glass, or by quickly
evaporating a larger quantity of the solution in a porcelain dish.
9. If the solid does not dissolve readily in water, in order to find a satis-
factory method of dissolving it, try preliminary experiments, such as are sug-
gested below, with small quantities of the powdered substance. The problem
is much simplified if the solid is known to be a single pure substance; for in
this case, when a reagent is found to react with some of the solid, the re-
mainder can be made to react in the same way by continued treatment with
this reagent.
10. When a method of dissolving the solid has been found, dissolve about
1 gram, and set aside about two-thirds of the solution in order that portions
of it may later be used in testing for NH 4 + and the negative constituents, and in
making confirmatory tests. Treat the remainder by the scheme of analysis for
the positive constituents ; use beakers or flasks in this work, and do not throw
away a portion of a solution whenever you have a large volume of it.
11. Methods of dissolving a difficultly soluble, pure substance. Observations
in these preliminary experiments with small quantities of the solid may be of
great assistance in identifying the substance. (In the following, some substances
are mentioned which will not be encountered in Analyses 4, 5 and 6).
(/) Treatment with dilute acid. Use N HNO 3 , boil if necessary. Note
the odor of any gas evolved. The use of HC1 is avoided because AgCl or
HgCl might be formed. Similarly BaSO 4 or PbSO 4 might form if H 2 SO 4
were used.
(2) Treatment with hot concentrated HC1. If chlorine is evolved a very
powerful oxidizing agent is present, such as MnO 2 or PbO 2 . Concentrated HC1
dissolves AgCl and HgCl readily, with formation of complex negative ions, but
the chlorides precipitate again when water is added to dilute the acid.
(3) Treatment with hot concentrated HNO 3 , a powerful oxidizing agent.
The rapid formation of brown gases indicates the presence of a reducing
[691
agent. There is a residue of sulfur after the action of nitric acid on a sulfide,
and the solution may contain SO 4 ~~.
(4) Treatment with either (a) a mixture of concentrated HC1 and HNO 3
(aqua regia) or (b) bromine water. A few substances, as HgS, are dissolved
more readily by these reagents than by concentrated HNO 3 (which is really
a more powerful oxidizing agent) on account of the formation of weak electrolytes
or complex ions.
Note. Before using for analysis a solution containing concentrated acid
evaporate it almost to dryness to remove the excess of acid, and add water.
Question. What happens when H 2 S is passed into concentrated HNO 3 or
aqua regia?
(5) Treatment with Na 2 CO 3 solution to form a carbonate. Use a con-
centrated solution of Na 2 CO 3 in a porcelain dish, cover the dish with a watch-
glass and boil for several minutes. Filter and retain the nitrate to test for the
negative constituent. Treat the residue with dilute HNO 3 , filter if necessary,
and analyze the solution for the positive constituent. Review the action of
Na.,CO 3 solution on CaSO 4 , Assignment 32; and, experiment, try the above
procedure with BaSO 4 .
(6) Special tests. For example, AgCl dissolves in NH 4 OH solution, while
HgCl turns black; AgCl, when treated with water and H 2 S, is transformed into
the black Ag 2 S, which then may be collected on a filter and dissolved in con-
centrated HI\ f O 3 .
12. When the unknown is a mixture of solid salts, the same method of pro-
cedure may be followed in order to determine a method of dissolving the solid
completely. However, it is often convenient to analyze separately the portions
dissolved in the different operations. For example the mixture might consist of
one salt which dissolves readily in water or dilute HNO 3 , and a second salt
which does not. If this method is followed, care should be taken to dissolve the
first salt completely before continuing the experiments with the residue. Note.
Frequently when a mixture of two salts, as AgNO 3 and NaCl, is treated with
water a reaction takes place.
L3. Problems, (i) Briefly outline experiments by which you could identify
each of the following: AgCl, HgNO 3 , CuO, HgS, ZnCl 2 , Fe(OH) 3 , CaCO 3 .
(2 In order to be able to interpret quickly observations made in preliminary
experiments on difficultly soluble substances, Paragraph 11, prepare a table,
listing in a vertical row all the difficultly soluble compounds, studied thus far, and
in a horizontal row the reagents, (a) dilute HNO 3 , (b) hot concentrated HC1,
(c) hot concentrated HNO 8 , (d) bromine water (or aqua regia). (c) NH 4 OH
solution : If the compound is insoluble in a given reagent mark its position in
the table with a cross, if soluble write in the formulas of the substances formed
when it goes into solution.
ASSIGNMENT 53
LEAD, TIN, ALUMINUM, CHROMIUM, AND BARIUM IN THE SCHEME OF ANALYSIS
(To follow Assignment 47)
Reference. Assignment 52.
1. Lead in the scheme of analysis. State what would be observed if a solution
of pure Pb(NO 3 ) 2 were treated by the scheme of analysis which you used in
Assignment 52. Sugest a method of determining whether a white precipitate
obtained with HC1 contains any PbCL. What .would be the result of treating
[70]
a mixture of HgS, CuS and PbS with hot 2 N HNO 3 (cf. Assignment 52, Par-
agraph 3) ? Give two reactions which are characteristic of Pb ++ but not of Cu ++ ,
and one reaction which is characteristic of Cu ++ but not of Pb ++ . Suggest a
method of separating lead from copper and mercury, i. e. of including lead in the
scheme of analysis. Experiment. Test your method with known solutions.
2. Tin in the scheme of analysis. State what would be observed if solutions
of (a) stannous tin, (b) stannic tin were treated by the scheme of analysis. In
which analytical group would stannous and stannic tin be precipitated? Suggest
two methods of separating tin from copper, mercury and lead. Experiment.
Try one of your methods with known solutions. Repeat the experiment in
Assignment 46, Paragraph 2 (/), using small amounts of dilute solutions of
(a) Cu(NO 3 ) 2 , (b) SnCL, and (c) a mixture of Cu(N(X) 2 and SnCU.
3. Aluminum and Chromium in the scheme of analysis. State what would
happen if separate solutions of Al 2 (SO 4 ) a , Cr 2 (SO 4 ) 8 and K 2 Cr.,O 7 were
treated by the scheme of analysis. In which analytical group are aluminum and
chromium precipitated, and what are the formulas of the substances obtained?
How would you prepare from a mixture of ZnS, FeS, A1(OH) ; < and Cr(OH) 3
an alkaline solution containing HZnO.r, CrO 2 - and AlO.r? What would be
the effect of adding Na 2 O 2 to this alkaline solution? Outline a method of
showing the presence of aluminum, chromium and zinc in the resulting solution.
Note. The alkaline solution which contains CrO 4 ~~ must not be acidfied until
the peroxide present has been decomposed by boiling the solution. Experiment.
Test your method with known solutions containing aluminum and chromium,
and chromium alone. Experiment. Try the procedure with the reagents. NaOH
and Na 2 O 2 , and note whether a satisfactory blank test is obtained for aluminum
and for zinc.
4. Barium in the scheme of analysis. Explain why you might expect barium
to be found with calcium in the scheme of analysis. Look up in a table of
solubilities, such as is given on the front cover page of Smith's text-book, the
solubilities in mols per liter of the hydroxides, carbonates, sulfates and chromates
of barium and calcium. Arrange these salts in one column in the order of their
solubilities in water. Devise two methods for the separation of calcium and
barium.
5. Procedure for the separation of barium and calcium, after precipitation as
carbonates. Experiment. Test the following procedure with a freshly prepared
precipitate of (a) BaCCX, and (b) CaCO. v To dissolve the carbonate pre-
cipitate pour repeatedly through the filter a 10 cc. portion of hot 3 N acetic
acid, and use a little more acetic acid if the precipitate does not dissolve com-
pletely. Add 20 cc. N NH 4 Ac, and heat the solution to boiling in a flask.
Measure out 15 cc. A 7 " Na 2 CrO 4 and add it, a little at a time, heating and shak-
ing the mixture after each addition. Finally heat the mixture at 90 - 100
for one or two minutes, and shake it at the same time. Filter, even if you
cannot see a precipitate, remove the filtrate and wash the precipitate with cold
water. (Pale yellow precipitate, presence of barium.) To confirm the presence
of barium, dissolve the precipitate by pouring hot HC1 through the filter, evap-
orating to a small volume, and applying the flame test. To test for calcium in
the filtrate,- make the solution alkaline with NH 4 OH, add (NH 4 ) 2 CO 3 , or
ammonium oxalate, and heat the mixture just to boiling. Note. The separation
depends on the fact that BaCrO 4 is much less soluble in water than is CaCrO 4 .
If the BaCrO 4 is not precipitated slowly in a hot acid solution, as directed, the
precipitate will be so finely divided that it will run through the filter.
6. Analyses Nos. 7, 8, 9, etc. Test for all the positive and negative con-
stituents you have studied. Verify your results by experiments with known
substances. !
[71 1
7. Problems. (/) Complete the table of the behavior of difficultly soluble
substances towards various reagents, Assignment 52 Problem (<?), by the
addition of the -difficultly soluble substances studied in Assignments 45-47 and 53.
(2) Point out the errors in the following reports:
(a) Mercury is reported to be present because a black residue remained
when the sulfides of the copper group were boiled with 2 N HNO 3 .
(b) Lead is reported to have been found in the silver group but not in
the copper group.
(c) In the analysis referred to in (b) a black precipitate was observed
in testing for the zinc group, although no test for iron was otbained.
What errors in the method of analysis would produce this result?
(d) Aluminum is reported to be present in the zinc group, but the note-
book shows the following records : on precipitating the zinc group
a greenish precipitate was obtained on the addition of NH 4 OH,
later a yellow solution resulted when Na^CX was added ; when this
solution was acidified there was evolution of gas and the yellow
color disappeared; on the addition of NH t OH a precipitate was
formed.
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