No- Waste Lab Manual
for Educational Institutions
Prepared by
College of the Redwoods
Under a Grant From
California Department of Toxic Substances Control
WE'VE CHANGED!
On July 17, 1991, the California Environmental Protection Agency
officially came into existence and the Toxic Substances Control Program
became the Department of Toxic Substances Control under that Agency.
The Toxics Program is no longer affiliated with the Department of Health
Services or the Health and Welfare Agency. The wording within this
particular document has not been changed to indicate this new affiliation.
The new mailing address follows:
Department of Toxic Substances Control
Office of Pollution Prevention and Technology Development
400 P Street, 4th Floor
P.O. Box 806
Sacramento, CA 95812-0806
(916) 322-3670
93 83948
NOWASTE LAB MANUAL - A procedure that
eliminates Toxic Waste Production from
Introductory Chemistry Laboratory Courses.
ABSTRACT
The feasibility of producing a laboratory manual for introductory
chemistry courses which incorporates procedures which produce little or
no toxic waste has been demonstrated.
This is accomplished by the use of
consecutive chemical reactions, so that the product of one reaction is used
as the starting material for the next.
In this fashion, all hazardous
chemical products are retained in the cycle, and benign waste, which serves
as a bleed on the system, is converted into educationally useful products.
Another major benefit derived from this procedure is a substantial
reduction in the expense of administering laboratory courses, an expense of
This provides a strong incentive for the adoption of
concern to educators.
the described approach.
Other advantages are accrued.
The school becomes
a role model of waste recycling and thus influences student attitudes
toward our environment.
Student interest in chemistry is strengthened by
creation of a purpose, meaningfulness and connection to their concept of
the "real world".
work as they learn.
The newly acquired chemical knowledge is put to useful
Conservation of dwindling resources is aided.
The goal is to significantly reduce or eliminate the generation of
hazardous wastes by our school laboratories.
Acknowledgements
The author would like to express gratitude to Karen Prentice,
Chris Schomers, and Louise Jensen for their help and encouragement.
This report was submitted in fulfillment of Award Number 87-TO137
by the Department of Health Services, Toxic Substances Control
Division under the sponsorship of the Department. Work was completed
as of April 30th 1989.
Disclaimer
The statements and conclusions of this report are those of the
Grantee and not necessarily those of the State of California.
The
mention of commercial products, their source, or their use in
connection with material reported herein is not meant to be construed
as either an actual or implied endorsement of such products.
_-
TABLE OF CONTENTS
...............................................
Recommendations .......................................................
Introduction ..........................................................
Summary and Conclusions
PAGE
1
3
4
Discussion
.............................................................
Experimental Outline (Schematic) ....................................
Experimental Details ................................................
1 . Preparation of Oxygen .........................................
2 . Preparation of Hydrogen .......................................
3 . Chemistry of Alkaline Earths ..................................
4 . Determination of an Empirical Formula .........................
. 5 . Chemistry of the Halogens .....................................
6 . Reactions of Ions in Solution .................................
7 . Types of Chemical Reactions I .................................
8. Types of Chemical Reactions I1 ................................
9. Activity Series ...............................................
10. Titration .....................................................
General
.
12 .
13 .
14.
15 .
16 .
17 .
18.
19 .
20 .
21 .
22 .
23 .
11
..
.........................................
Fractional Crystallization ....................................
Chemistry of Aluminum .........................................
Chemistry of Lead .............................................
Solubility of Salts and Saturation ............................
Preparation of Lead Nitrate and Iodine from Lead Iodide .......
Chemistry of Chromium
.............
Chemical Separation of Zinc ...................................
Redox Titration ...............................................
Electrolysis of Zinc Sulfate ..................................
Ion Exchange. Recovery of Oxalic Acid .........................
Kinetics. Clock Reaction ......................................
Synthesis of Methyl Benzoate ..................................
24 . Saponification of Methyl Benzoate .............................
25 . Preparation of Iodine .........................................
.
27 .
26
Preparation of Silver Nitrate from Silver Halides
....................
...................................
10
14
51
51
54
58
60
61
66
67
67
71
72
73
73
75
76
77
78
79
84
84
85
85
86
87
87
88
Preparation of Ammonium Iodide from Iodine
89
Synthesis of Sodium Nitrate
91
TABLE OF CONTENTS
PAGE
.
29 .
30 .
31 .
32 .
......................................
Distillation ................................................
Separation of Salts .........................................
Preparation of Alumina ......................................
Preparation of Bleach .......................................
33 . Applications of Industrial Chemicals ........................
34 . Hydrolysis of Salts .........................................
35 . Determination of Purity by Melting Point ....................
Waste Products ....................................................
Cost Reduction ....................................................
Impurity Buildup ..................................................
Storage Space .....................................................
28
Kjeldahl Determination
...............................
Outline of Introductory and Physical Chemistry
Experiments ................
Chemicals and Costs of a Typical Chemistry Manual ......
Costs of Chemicals in NOWASTE Manual ...................
fnnovative Nature of NOWASTE MANUAL
Appendix I
Appendix I1
Appendix 111
Appendix IV
Suggested Names for Solids and Solutions Used in
NOWASTE Manual
.............
92
93
93
94
94
95
95
95
96
99
102
104
105
107
111
113
114
SUMMARY AND CONCLUSIONS
A feasibility study has confirmed the practicality of developing a
laboratory manual for introductory chemistry courses of a novel nature.
This manual is intended to minimize or eliminate the production of toxic
wastes by using the products of one experiment as the starting materials
for the same or other experiments.
In this manner, the toxic chemicals are
retained in the cycle of experiments.
As a valuable dividend, the costs of conducting laboratory classes are
reduced, since the toxic chemicals are not discarded and do not have to be
repurchased in amounts formerly required.
This cost reduction can serve as
a strong incentive for adoption of this methodology by our institutions,
particularly at a time when the rising costs of laboratory classes are of
concern to all administrators.
Other advantages are accrued.
disposal.
The school becomes a role model of waste
The pattern of responsibility for waste is implanted in the
student mind.
Student interest in chemistry in the sense of meaningfulness,
purpose and continuity is expected to be enhanced.
The student's chemical
knowledge is immediately put to use in a real and worthwhile cause.
An
experiment in which iodine is carried through a series of reactions producing remarkably different products, and then is regenerated, has an obvious
educational advantage.
eliminated.
Costs of hazardous waste disposal are essentially
Conservation of irreplaceable resources is achieved.
This study was carried out by a combination of desk research and laboratory work. A concerted review of pertinent literature and published
laboratory manuals was conducted.
This process helped to define the types
of experiments necessary in introductory chemistry courses and recent
tendencies in emphasis of topics.
Having defined the scope of necessary
experiments, procedures were created that could accomplish the stated goals.
These have been tested in the laboratory to the point of establishing their
suitability for intended purpose.
There have been several approaches to design of the experiments.
Sometimes solutions can be used repeatedly as in the electrochemical cell
voltage and crystallization experiments.
In other cases, an element may
be used in a succession of chemical reactions to produce a useful final
product, for example, plaster of paris.
- 1 -
The chemistry of aluminum is
demonstrated by carrying aluminum through a series of reactions and then
converting it in the final experiment to activated alumina, which is then
used as a desiccant in future experiments.
In general, however, the
approach has been to devise a series of reactions which retain the chemicals
within a cycle.
This method represents a general approach to the design of laboratory
experiments for educational purposes.
Since the procedures virtually
eliminate hazardous waste production and at the same time-reduce expenses to
a significant degree, it is felt that this approach may well represent the
method of experimental design for academic purposes in the future.
In conclusion, the feasibility of utilizing a novel approach consisting
of consecutive chemical reactions has been demonstrated.
Further work will
concern perfection of the methods to the point that they may be easily and
safely performed by our beginning students of chemistry.
Beyond that goal,
appIications to general chemistry, organic chemistry and analytical chemistry
courses are envisioned.
- 2 -
RECOMHENDATIONS
A feasibility study has demonstrated that the concept of a basic
chemistry laboratory manual which uses toxic chemical products from experiments as starting materials for others is a practical concept.
following tasks remain to complete the work.
The procedures must be cor-
rected and modified to fit laboratory time requirements.
term reuse of chemicals must be ascertained.
The
Effect of long
Final design of the experiments
must be completed.
Exact quantities of chemicals necessary for the procedures
must be determined.
the manual, must be written, proofread and finally pub-
lished.
Efforts to achieve adoption of the manual must be carried out.
Later, similar manuals should be devised for general, organic and analytical chemistry courses.
- 3 -
INTRODUCTION
The purpose of this project is to determine the feasibility of a novel
approach which eliminates the production of toxic chemicals or waste products from chemistry laboratory classes.
The experiments are designed in
such a manner that toxic chemicals are recycled to other or the same experiments.
Build-up of impurities is eliminated by bleeding the salts formed
and converting them to either chemicals used in the cycles (e.g., bleach,
nitric acid) or other products which are educationally useful (e.g.,
of paris, desiccants).
plaster
A laboratory manual will be written which is intended
for use in our highschools and community colleges.
The ultimate goal is to
extend the use of this method to general (college) chemistry, non-science
major chemistry, organic chemistry, and perhaps analytical chemistry courses.
It is helpful to understand the general background of this project
from the point of view of a chemistry teacher.
All laboratory procedures
used for educational purposes produce chemical wastes.
At present, the means
of disposal of these chemicals is in sanitary sewage systems and landfills,
and ocassionally, particularly in the case of organic wastes, incineration.
The moral and, increasingly legal, responsibilities2 dictate that sanitary
sewage or sanitary landfill disposal are unsatisfactory methods in the long
term.
They are, nevertheless, the only options available at this writing.
For example, a 1988 publication by the American Chemical Society recommends
sanitary sewage disposal for copper salts, alcohols, organic acids, acetone,
barium sulfate, and all filtrates produced after precipitation of the sulfides
of lead, cobalt and silver with sodium sulfide.
It also reccrnnnends ahat the
precipitates be disposed of in sanitary landfills. 3 Versions of this procedure are found in other references.4
1 McKusik, B.C., J Chem. Educ., 61, A152 (1984).
2 for example, U.S. Environmental Protection Agency (EPA) regulations under
the Resource Conservation and Recovery Act (RCRA), 45 Federal Register,
10 May 1980, and subsequent amendments; Title 40 Code of Federal
Regulations, Chapter 1.
3 Summerlin, L.R., Ealy, Jr., J.L. and Borgford, C.L., "Chemical
Demonstrations, A Sourcebook for Teachers", Vol 1, American Chemical
Society, Washington, D.C.,
1988.
Browne, L.M., and Weir, G.L., "Prudent
Practices for Disposal of Chemicals from Laboratories", National
Academy, Washington, D.C.,
1983., Armour, M.A.,
Browne, L.M., and
Weir,. G.L., "Hazardous Chemicals Information and Disposal Guides",
2nd edition, University of Alberta, 1984, p.165.
4 for example, see Armour, M.A.,
-_
- 4 -
There have been a few suggestions at chemical recycle, the theme of
this text.
Walton suggests designing an experiment which "illustrates a
disposal method of a product from a previous experiment".
A multistep
organic synthesis has been devised by Stradling and Gage in which acetophenone is carried through several reactions, the products being used as
starting materials until eventually an end product is produced which must
be discarded.2 Neckers, et al, describe recyle of organic solvents, and a
sequential experiment to reduce the costs of maintaining an organic chemistry
laboratory.3 Schumm has suggested a series of experiments using alum
synthesized from scrap aluminum.4
Other than these limited suggestions, use of microscale experiments,
or disposal by landfill or sanitary sewer, there have been no alternatives
offered to the chemistry instructor.
of solution at the present time.
This presents a problem of great urgency
Failure to provide a solution will permit
undesirable practices to be perpetuated.
This project departs completely from the disposal approach.
All toxic
chemicals are continually reused; losses occur only from accidents or tracequantities not removed from solution.
Even non-toxic wastes are retained
and processed to produce chemicals useable in the experimental cycle, or
useful educationally.
This project has consisted of three phases:
the desk research, and the laboratory research.
The literature search,
The product of this work
is an outline of experiments to be incorporated into a laboratory manual,
after further research has established the practicality and efficacy of the
proposed experiments.
The literature search provided information on previous related work
and familiarization with a number of existing laboratory manuals.
The
search was also invaluable in providing the germ of innovative experiments
in the proposed mode.
J Chem Educ., A69 (1987)
1 Walton, W.A.,
2 Stradling, S. and Gage, C., J Chem. Educ., 1116 (1985)
3 Neckers, D.C., Duncan, M.B., Gamor, J. and Grasse, P.B., J. Chem.
Educ., 54, 690 (1977).
4 Schumm, M.K., "Recycling Materials and Organic Identification Without
Instruments", Montgomery College, Roclnrille, M.D., 20850, Paper
MC4-02, AMW, paper presented at the Seventh Biennial Conference in
Chemical Education, August, 1982.
- 5 -
The following laboratory manuals were investigated by listing each
experiment, its subject and a brief expository paragraph on the subject.
Thirteen manuals were analyzed in this fashion. They were:
1.
"Laboratory Kanual" by Joseph E. Davis, W. Keith McNab, A.L.
McClellan, and Paul R. OConner, 1st Edition, D.C. Heath and Co.,
Lexington, Mass., 1983.
2.
"Laboratory Manual for Chemistry" by Maxine Wagner, 1st Edition,
Allyn and Bacon, Inc., Newton, Mass., 1983.
3.
"Foundations of Chemistry in the Laboratory" by Morris Wein,
Leo R. Best and Robert L. Miner, 5th Edition, Brooks/Cole Pub. Co.,
Monterey, Calif., 1982.
4. "Laboratory Experiments" by Charles M. Wynn and Gary A. Joppich,
2nd Edition, Wadsworth Pub. Co., Belmont, Calif., 1981.
5.
"Chemistry in the Lab" by William L. Masterson, Emil J. Slovinski,
and Edward T. Walford, 1st Edition, Holt, Rinehart and Winston,
New York, N.Y.,
6.
1980.
"Laboratory Chemistry" by L. Neal Carmichael and David F. Haines,
1st Edition, Charles E. Merrill Pub. Co., Columbus, Ohio, 1979.
7.
"Laboratory Manual" by Phyllis M. Dietz, Robert L. Telefsen,
Robert W. Parry and Luke E. Steiner, 2nd Edition, PrenticeHall, Inc., Englewood Cliffs, N.J.,
1975.
8. "Laboratory Experiments", by L. Neil1 Carmichael and David F.
Haines, Charles E. Merrill Pub. Co., Columbus, Ohio, 1979.
The above manuals are intended for high school and community college
introductory chemistry courses.
In addition, the following general chemistry manuals were examined.
9. "Encounters in Experimental Chemistry" by William L. Jolly,
1st Edition, Harcourt Brace Jovanovich, Inc., New York, N.Y.
1972.
10.
"Chemistry in the Laboratory" by Charles W.J. Scaife and O.T.
Beachley, Jr., 1st Edition, Saunders College Pub. Co.,
Philadelphia, Pa., 1987.
-_
11.
"Basic Laboratory Studies in College Chemistry", by Grace R.
Hered, 7th Edition, D.C. Heath and Co., Lexington, Mass., 1980
- 6 -
12.
"Laboratory Manual for General Chemistry" by James E. Brady and
Jo. A. Beran, 2nd Edition, John Wiley and Sons, New York, 1982.
13.
"Fundamentals of Chemistry" by Frank Brescia, John Arents,
Herbert Neislick, Amos Turk and Eugene Weiner, Academic Press, Inc.
1980.
From this survey, a familiarity was gained with the subject matter
covered and with the educational approach. Some specific information was
New York, N.Y.,
For example it was determined that 52% of the experiments in the
introductory manuals involved chemical reactions and that the most frequent
obtained.
topics covered, in order of decreasing number of experiments were stoichiometric
relationships, organic chemistry, electrochemistry, equilibrium, qualitative
analysis, redox, kinetics, various reactions of ions, chemistry of halogens,
types of chemical reactions, empirical formulas, titration, and hydrates.
Egghty per cent of the topics covered were in these categories.
These
analyses were used as guidelines in the design of this manual.
In the general survey of literature pertinent to the subject of introductory chemical courses, easily the most valuable source proved to be the
Journal of Chemical Education, published by the American Chemical Society.
This journal contained many articles specific to this topic, as is demonstrated
by the number of refernces in the text.
helpful.
Two texts proved to be particularly
These were "Qualitative Analysis" by "herald Moeller, 1st Edition,
McGraw-Hill, N.Y.,
1958, and "General Chemistry" by Henry F. Holtzclaw, Jr.,
William R. Robinson and William 11. Nebergall, 7th Edition, D.C. Heath and
Co., Lexington, Mass., 1984.
inorganic chemical reactions.
These are excellent references for general
And lastly, the "Handbook of Chemistry and
Physics", published each year by the Chemical Rubber Co., Cleveland, Ohio,
was invaluable in providing, in particular, data on solubilities of inorganic compounds.
The general approach to the methods developed within this report is
best explained by contrasting conventional experiments with those developed
here. In almost every experiment investigated in conventional procedures.
experiments are performed non-sequentially. For example, if the experiment
is designed to illustrate the activity series of the metals, the instructions
__
would be, for example:
- 7 -
1)
add copper metal to an aqueous solution of silver nitrate and observe
the results (silver is precipitated and the copper goes into solution).
Discard the mixture.
add lead metal to an aqueous solution of copper nitrate and observe the
2)
results (copper precipitates and lead goes into solution).
Discard
the mixture.
3)
add iron metal to a solution of lead nitrate and observe the results
(lead precipitates and iron goes into solution).
Discard the mixture.
........etc.
In this process, waste products of copper, silver nitrate, copper nitrate,
lead, lead nitrate, iron, and iron nitrate have been created.
Contrast this
with the approach used in the method of this report.
1)
add copper metal to an aqueous solution of silver nitrate and observe
.
the results.
After the reaction is complete, remove the unreacted
copper and precipitated silver (for details see Experiment 9.
2)
add lead metal to the copper nitrate solution from 1) and observe
the results.
After the reaction is complete, remove the unreacted
lead and precipitated copper.
3)
add iron metal to the lead nitrate solution from 2) etc.
In other words, by carrying out the reactions sequentially, the products
are minimized.
reactions.
The second phase is to use all products in other chemical
In the case above silver produced eventually returns to the cycle
as silver nitrate, copper as copper nitrate, etc.
So that, in this case,
they create a closed cycle.
There are some cases where the cycle is not closed, and the starting
material ends up as a product useful in some way to the students. . In this
manner aluminum is carried through a series of reactions to finally exit
the cycle as alumina, which is then used as a desiccant in future experiments.
Calcium metal is carried through sequential reactions eventually exiting as
plaster of paris, from which the students fashion an object such as a mold.
These products serve as a bleed to the system to prevent buildup of impurities.
The general scheme, then, is to create a sequence of reactions which
either
retains the species within a cycle, or utilizes it as a useful
-_
product. This is contrasted with the conventional approach, which carries
- 8 -
out reactions separately and in a non-connected fashion, and thus produces
a multitude of waste products.
The educational advantage to the sequential
approach should be noted here; the connection between chemicals and reactions
becomes much more pervasive.
The work carried out here is limited to introductory chemistry courses.
The concept can be expanded to include general (college level), organic
and analytical chemistry courses.
It may be somewhat applicable to certain
phases of the biological sciences, perhaps in cooperation with the chemistry
department.
More generally, it is a way of thinking that should have
applicability to a number of laboratory procedures.
- 9 -
DISCUSSION
Chemistry laboratory instruction is an essential part of a chemistry
education.
But a chemistry laboratory generates toxic chemical waste which
must be eliminated in some manner.
Historically this waste has been discharged
into the sewers but this practice is now regarded as unsatisfactory and, in
some cases, illegal.
Concern for the environment is reflected in the passage
of the Tanner Bill, AB 2948, requiring each county to assess its hazardous
waste streams and adopt alternatives to land disposal, the Safe Drinking and
Toxic Enforcement Act of 1986 which prohibits discharge of carcinogens or
teratogens into water or land, and Farr Assembly 3i11 685, which provides
grants to programs demonstrating hazardous waste source reduction, recycling,
or treatment technologies.
In any case, toxic waste discharge of any kind is undesirable environmentally, economically and conservationally.
This project was conceived to
address this issue, eliminating toxic waste disposal entirely by integrating
the toxic chemicals into a cyclical laboratory instructional process which
uses products of one experiment as a starting point for the same or other
experiment.
Considerable attention has been given recently to the problem
of toxic waste from academic laboratories.
Solutions suggested have been
either reduction of amounts of chemicals employed for instruction to the
microscale range or disposal by commercial haulers,
form of concentration.
usually after some
Neither of these methods is entirely satisfactory.
\
Experiments on a micro scale do not have the visual impact that macro scale
experiments have, in the same way that a 6 inch TV screen is:.less'irhpressive
than a 22 inch screen.
Further, special techniques must be taught to inex-
perienced students in order to achieve results equivalent to the macro scale.
This is why micro scale experiments have been used only to a limited extent
in the past.
Disposal of concentrated chemical wastes adds an additional
cost to the laboratory instruction, and concentration of wastes requires
additional handling.
Further, it is precisely this land contamination which
we are attempting to avoid.
1 "Prudent Practices for Disposal of Chemicals from Laboratories", National
_,
Research Council, National Academy Press, Washington, D.C.,
1983
2 See, for example, A m o u r , M., J. Chem. Educ., 65, A64 (1988); Armour, K.,
Browne, L.M. and Weir, G.L. , J. Chem. Educ. , 6T; A93 (1985).
- 10 -
?
This project was conceived to address the toxic waste issue, and
eliminate the waste disposal problem by integrating these chemicals into a
cyclical process whereby products of one experiment are used as starting
points for the same or another experiment.
In designing these experiments,
a number of considerations were kept in mind.
These considerations were:
1) Types of experiments to be emphasized.
2) Safety of the experiments.
3) Creation of interesting, educational and entertaining experiments.
4 ) Use of economical and easily available materials.
1.
Types of experiments
A survey of a number of modern introductory level laboratory manuals
was conducted to determine those topics most frequently covered.
and authors of these manuals are listed in the introduction.
Titles
In addition,
the chemical education literature was researched to ascertain modern trends
There appears to be a major movement back
.in chemistry laboratory education.
to descriptive chemistry, taught almost exclusively prior to the 1940's.
By descriptive chemistry is meant syntheses, reactions, commercial processes
of chemicals.
In other words, the students need to understand that facts
come before theories.
A tendency away from descriptive chemistry began with
Linus Pauling's classic "Nature of the Chemical Bond" which introduced a
great many theoretical concepts (e.g.,
etc.)
electronegativity, orbital hybridization,
into the field of chemistry and dramatically altered the mode of present-
ation.2 This has continued up to the present time.
There is now a feeling
among educators that the theoretical emphasis must yield somewhat to the
descriptive aspect of the ~ c i e n c e .This
~
tendency has been followed in the
design of the present manual, which is ideally suited to this mode of
presentation.
2.
Safety of the experiments.
The safety of the experiments will be given careful consideration.
Cautionary warnings will be given whenever hazardous chemicals are to be
used.
__
A full section on safety consideration will appear at the beginning
1 Zuckermann, J.J., J. Chem. Educ., 63, 829 (1986)
2 Pauling, L., "The Nature of the Chemical bond", 3rd Ed., Cornel1
University Press, Ithaca, N.Y.,
1960
3. for example, see Whisnant, D.M., J. Chem. Educ., 2,792 (19821,
"What Happened to Descriptive Chemistry?", Editorial, J.Chem. Educ.,
915 (1985) and Zuckermann, J.J., J. Chem. Educ., 63, 829 (1986).
62,
-
11
-
of the manual.
Apparatus set-ups are and will be designed to trap all noxious
gases or vapors evolved.
Substitutions of less harmful chemicals for those
known to be dangerous have been made when possible.
Thus, cyclohexane has
been substituted for carbon tetrachloride in several experiments.
The use
of potassium chlorate, a potentially explosive compound when dry, has been
avoided in the preparation of oxygen.
In some cases where important chemistry
is involved, such as that of chromium, a known carcinogen, appropriate measures
such as wearing of gloves, cautionary statements, and strict supervision will
be emphasized.
A large number of hazardous compounds result from the combin-
ations of the chemical elements, and it is a purpose of chemistry to examine
the nature of these substances.
Since this is unavoidable, the best approach
is the teaching of safe and appropriate procedures, much as we instruct
students in driver's training or woodshop in safe procedures with potentially
hazardous equipment.
3.
*
Creation of interesting, educational and entertaining experiments.
If this manual is to be used extensively, it must not only eliminate
toxic waste and save money, but it must provoke interest of the teachers and
students.
The author has attempted to do this by including experiments with
colorful reactions, such as the chromium series of experiments or the reactions
of copper, striking phenomena such as the crystallization of alum, the formation of iodine from lead iodide, or the clock reaction, and practical descriptive
experiments such as the manufacture of plaster of paris and chemical
recovery techniques such as the recovery of zinc sulfate (beautiful fractal
crystalline growth of zinc is observed from a colorless solution) or fractional
distillation.
In addition, a number of interesting laboratory techniques
and student built pieces of equipment have been introduced or suggested.
4.
Use of economic and easily available materials.
An effort has been made to incorporate the use of readily available
materials in the course in order to reduce expense and to provide student
connections to "real-world" experience.
This is reflected in the use of
bleach, aluminum foil, vinegar, salt, etc.
Using the guidelines outlined above, research was conducted to conceive
and carry out a laboratory program containing sequentially designed experiments.
The feasibility of such an approach has now been demonstrated.
The
possible consequences of teaching an introductory course in chemistry with
a laboratory course designed in this manner are worth analysis.
-
12
-
First, let's consider student interest in chemistry.
An often heard
comment from beginning students is that chemistry has no connection with
the real world.
If they are reminded that the stuff of nature is all chemical,
they remain unconvinced.
The chemicals they ordinarily use in the laboratory
experiments come out of bottles not found in any store, and the products from
their experiments are quickly disposed of to be seen no more.
Contrast that
experience with possible student reaction using the sequential approach.
The chemicals used now are chemicals they have themselves produced.
products of their labors are used for further work.
familiar substances are used:
These
In addition, ordinary
copper wire, baking soda, bleach, aluminum foil.
The connections begin to be made.
Chemistry becomes "real-world".
From an
educators point of view, these connections are invaluable.
Further, students experience their newly acquired skills being used to
improve their world.
They are accomplishing a goal, as well as becoming
.educated in the process. And they are learning an attitude:
responsible for one's actions.
One must be
In the experience of the author, most students
have an environmental awareness and should respond very positively to this
program.
And the school then sets the pattern.
model.
- 13 -
Our schools should be the role
SCHEMATIC OUTLINE OF CYCLICAL EXPERIMENTS
The following symbols are used in the outline of experiments:
denotes an experiment number.
For example,
[SI indicates
Experiment 14.
EtOH
denotes ethyl alcohol
MeOH
denotes methyl alcohol
HOAC
denotes acetic acid
0
denotes the phenyl group
T
denotes a gas
4
A
denotes a solid
0
denotes that the compounds are by-products or unreacted species.
denotes heat
This outline will be helpful in following the discussion, and providing
a quick summary of the experiments.
This will be followed by a narrative.
This summary and discussion represent only the descriptive chemistry experiments that are proposed to be included in the manual.
Other types of experi-
ments, including those of an introductory nature on physical chemistry
experiments, will not be discussed in this report.
of these are included in Appendix I.
-
14
-
A list and brief summary
DISCUSSION:
Note:
Experimental Outline
All chemicals should be assumed to be in aqueous solution unless
specified as solid by ,&
Experiment 1:
H202
.
Preparation of Oxygen
Fe
3
O2t
___)
Fe
f
Purpose :
This experiment illustrates the preparation and properties
of oxygen, and the action of a catalyst.
Description:
Oxygen is generated from 3% hydrogen peroxide and an iron
nail as catalyst. The reaction mixture is heated.
Oxygen is collected over water and tested in the usual
manner.
Comments :
This experiment eliminates the hazards of conventional
preparation of oxygen from potassium chlorate (explosion
potential) or 9% hydrogen peroxide (chemical burn potential),
as well as the necessity of disposing of manganese dioxide,
the conventional catalyst. Costs of the experiment are
reduced.
- 15 -
Experiment 2: Preparation of Hydrogen
Purpose:
This experiment illustrates the preparation and properties
of hydrogen.
Description:
Hydrogen is prepared by the action of clean aluminum,
prepared by pickling in hydrochloric acid, on sulfuric
acid. The hydrogen is collected over water and
conventionally tested.
Comments :
Use of zinc metal in this preparation is avoided completely
by the substitution of aluminum. Aluminum is rendered
reactive to dilute sulfuric acid by pretreatment with
hydrochloric acid which removes the protective oxide
coating. Scrap aluminum from cans, foil, or other source
work well. The necessity of dealing with zinc (and
its inevitable companion cadmidin the comparatively
large quantities required for this experiment is
eliminated.
-
16
-
Experiment 3:
Chemistry of the Alkaline Earths
Purpose:
This experiment acquaints the student with some properties
of the Group I1 metals and their compounds.
Description:
Reactions of calcium and magnesium with water are carried
out and the resulting hydroxides are tested with litmus
and then dissolved in hydrochloric acid. The chlorides
are reacted with sulfuric acid. Calcium sulfate precipitates. Magnesium sulfate does not. The products from
both experiments go to 1301. In set 2, the metals are
shown to combine directly with iodine. Unreacted iodine
is converted to iodide with sodium thiosulfate. The
carbonates are precipitated. The iodide-containing filtrates
- and the precipitates in
are processed in [25]
[a].
-
17
-
Experiment 4 :
Determination of an Empirical Formula
Purpose :
As stated.
Description:
Ammonium iodide is reacted with lead nitrate to produce
lead iodide. The formula of lead iodide can be calculated
from the known percentage of lead in the nitrate and the
mass of lead iodide obtained. The lead iodide is converted
16].
back to lead nitrate and iodine in [-
-
18
-
Experiment 5:
Chemistry of the Halogens
- 19 -
Purpose :
This experiment illustrates some chemical properties of the
halogens as well as techniques that may be employed to
separate halides, and to identify halides.
Description:
In Set 1 the silver halides of iodine and chlorine are
precipitated to illustrate the classic test for halide ion.
The mixtures are heated to coagulate the halides and filtered.
The filtrates go to [=I.
The precipitates are treated with
ammonium hydroxide on'-ihe filter. The silver chloride
dissolves, the silver iodide does not, illustrating a halide
separation technique. The ammoniacal solution of silver
chloride and the silver iodide precipitate are treated with
hydrochloric acid. The silver chloride reforms. The silver
halides go to [-I-
.
In Set 2, ammonium halides of iodine and chlorine are treated
with chlorine generated in situ from acidified bleach. The
free halogens are extracted into cyclohexane, yielding a
yellow and a violet organic layer, respectively. Phase
separation is acheived by passage of the mixture through a
water-wetted filter paper. The organic phase is retained by
the paper. The aqueous phase goes to waste acid. The organic
phases are then treated with aqueous thiosulfate which converts the halogens to the halide ions. The mixture is phase
separated, the or anic layer going to
and the aqueous
layer going to
[z]
-
Comments :
By confining the experiments to chlorine and iodine, the
extremes of the behavior of typical halogens are illustrated.
Fluoride ion is avoided as atypical (see Discussion) and
because of its toxicity. Bromide ion is avoided because of
the complications introduced into the procedures by its
inclusion.
The experiment may prove to be too long for a one hour
period, and may be brobea:ppi.~nto;-.twa
sacperhnts.
-
20
-
Experiment 6:
Reactions of Ions in Solution
ORANGE-RED
BLACK
PALE BLUE
YELLOW
Purpose :
This experiment gives a sense of characteristic precipitates
of the various ionic combinations.
Description:
Salt solutions are mixed and the reactions observed. These
are carried out in quantities of several drops on spot
plates containing eight depressions. In this manner, the
student is able to keep track of the various reactions.
The waste products are minimized.
Comments :
The reactants are in such small amounts that the waste
produced is minimal (see Discussion).
The waste may
be accumulated over a period of years, and processed as
a student project.
- 21 -
Experiment 7:
Types of Chemical Reactions I: Combination, Double Replacement
Purpose:
As stated.
Description:
is ground
Copper obtained the previous year from
in a mortar, placed in an evaporating dish covered with
aluminum foil and heated under an atmosphere of oxygen
to produce copEr oxide. The oxygen is generated as
described in [ 1 ] fron 3% hydrogen peroxide. The oxygen
thus produced is conducted through a hole in the foil and
provides a blanket of oxygen for the copper as it is heated.
The hot copper reacts with the oxygen to produce copper
oxide. This is allowed to cool and dissolved in hydrochloric
acid to produce copper chloride for the next experiment.
C
O
~
t
:S
[E]
The union of the copper with oxygen exemplifies a combination
reaction. The copper oxide is then dissolved in hydrochloric
acid in a double replacement reaction.
- 22 -
Experiment 8 :
Types of Chemical Reactions 11: Double Replacement,
Decomposition, Single Replacement
Description:
Copper (11) chloride from [7]
is carefully neutralized with
sodium hydroxide and then rendered basic to precipitate
copper hydroxide. This is heated and the green hydroxide
is converted to black copper (11) oxide in a decomposition
reaction. The oxide is dissolved in sulfuric acid to
produce the sulfate. This is treated with aluminum metal
preconditioned with hydrochloric acid to yield copper metal
This is removed by filtration and goes to [TI. The filtrate
goes Bo]. The reclamation of the copper m z a l illustrates
a single replacement reaction.
Comments:
This experiment produces a number of striking chemical
change and a final product of copper, the starting material
for [ 7 1 .
-
23
-
Experiment 9:
Activity Series
Purpose :
This experiment demonstrates the Activity Series of metals
in a sequential manner.
Description:
In sequence copper replaces silver, lead replaces copper,
aluminum replaces lead. By heating the solutions, the
replacement reactions can be performed in sequence in
about one hour.
Copper wire is added to silver nitrate solution and heated
to precipitate silver. The wire is removed from the mixture,
washed and set aside. The mixture is filtered to yield
copper nitrate solution and silver metal. This sequence is
repeated with the aqueous copper nitrate and lead (lead in
the form of shot is useful here). The lead shot is washed
free of copper and removed. 'Phe aqueous lead nitrate is
heated with scrap aluminum (soda or beer cans) and the
process repeated. The aluminum is pretreated with hydrochloric
acid. The wire, lead sinkers and scrap aluminumAre reused
the following year. The silver powder goes to 1171. The
copper powder goes to [TI.
- The lead powder g o e s 7 0 [-I.
-
Comments :
The Activity Series is well visualized in this experiment
as each more active metal replaces the less active.
-
24
-
Experiment 10:
Titration
Purpose:
This experiment illustrates the principles of acid-base
titration.
Description:
Sodium hydroxide is standardized with oxalic acid. The
standardized solution is then used to standardize hydrochloric acid and determine the percentage of acetic acid
in vinegar.
Comments:
This experiment is mentioned here because the sodium
oxalate produced is used as raw material for
which produces oxalic acid for reuse in this experiment.
[z]
- 25 -
Experiment 11:
Chemistry of Chromium
K2C1-207
HOAc
H202
K2cr041
(HOAc)
HoAc
Purpose:
,
I-
(HOAc
Cr05
(H202)
( KOAc )
K2Cr207
(HOAc)
Cr(OAc13
( HOAc )
-
1
}
KOH,
KCr02
(KOAC
1
[El
The chemistry of a transition metal is explored. Hydrogen
peroxide is shown to act as both an oxidizing and reducing
agent
.
Description:
~
Potassium dichromate is treated with acidic 3 % hydrogen
peroxide to form the deep blue peroxo complex. Heat
converts this complex to red chromium acetate. Potassium
hydroxide precipitates green chromium hydroxide which
with excess base is converted to potassium chromite.
Basic hydrogen peroxide oxidizes this to yellow potassium
chromate which is converted by acetic acid to orange
potassium dichromate.
Comments :
This is a visually interesting experiment because of the
color changes. The varied chemical reactions are shown to
eventually produce the starting material. Several oxidation
states of chromium are achieved.
-
26
-
Experiment 12:
K2Cr207
(KOAc)
(HOAc)
Fractional Crystallization
1
evaporate
chill
filter
Purpose :
This experiment will recover the potassium dichromate
and illustrate a chemical separation technique.
Description:
Aqueous potassium dichromate solution obtained from
is concentrated by evaporation and chilled to produce
crystals of sparingly soluble potassium dichromate.
The crystals are washed with ethanol and dried. The
mother liquor is retained and added to the next years
dichromate solution (see diagram). Every five years or
so the mother liquor should be processed as follows:
the solution is treated with acidic hydrogen peroxide to
convert all chromium to the trivalent form and this is
then replaced with aluminum, the chromium precipitating
the solid. The filtrate goes to
and the chromium
in the third
is dissolved in acetic acid and used in [A]
step.
[z[
[z]
- 27 -
Experiment 13:
The Chemistry of Aluminum, an Amphoteric Metal
Purpose :
The chemistry and amphoteric properties of aluminum are
demonstrated.
Description:
Aluminum is reacted with sulfuric acid to illustrate its
metallic properties. It is then treated with potassium
hydroxide to illustrate its amphoteric properties.
Sulfuric acid is added to the basic reaction mixture to
re-precipitate aluminum hydroxide. The product is used
to prepare alum.
Comments:
The prepared alum may be used in experiments demonstrating
stoichiometry of a hydrate and melling point determination
or in the preparation of alumina [32].
-
- 28 -
Experiment 1 4 :
Chemistry of Lead
[COLOFUESS]
[WHITE]
[CHOCOLATE]
[WHITE]
filter
Purpose :
As stated.
Description:
Lead powder from
-3 reacted with nitric acid to
produce lead nitrate. these series of reactions are
carried out in a reaction flask equipped with thistle
tube and gas exit. The gases conducted through two
more flasks containing sodium hydroxide which act as
nitrogen dioxide absorbers in this first reaction. The
absorbed nitrogen dioxide goes to [=I.
[z]
-
The lead nitrate is carefully neutralized with sodium
hydroxide to precipitate lead hydroxide, which is oxidized
with bleach to yield lead dioxide. This is converted to
lead chloride. Chlorine produced by lead (IV) oxide
oxidation of chloride ion is reduced to chloride by sodium
thiosulfate. The lead chloride is converted to the less
soluble lead iodide with ammonium iodide. The lead iodide
goes to [E].
-
- 29 -
Experiment 15:
Solubility of Salts and Saturation:
Demonstration
seed
NaOAc
&
--
NaOAc
T
Purpose :
The principles of unsaturated, saturated and supersaturated
solutions are discussed and the striking phenomenon of
rapid crystallization from supersaturated solutions is
demonstrated.
Description:
This experiment represents a closed cycle. Supersaturated
solutions of sodium acetate and sodium thiosulfate are
retained. When used, the solutions are heated to dissolve
all salts and then seeded to produce large crystals of
sodium thiosulfate and sodium acetate. Heats of crystallization are observed. Solutions may be used indefinitely.
- 30 -
Experiment 16:
Preparation of Lead Nitrate and Iodine from Lead Iodide
Purpose :
. Description:
This experiment demonstrates some chemistry of lead and
iodine, and illustrates chemical synthesis. In the process,
lead nitrate and iodine are recovered.
Lead iodide is treated with nitric acid in a flask equipped
with thistle tube and gas exit tube. The evolving gas is
conducted into a flask containing water. The water traps
any iodine which escapes the reaction flask as vapor. Heat
is applied if necessary to initiate reaction. At first
nothing happens, and then the lead iodide begins to disappear,
the solution turns brown and violet vapors of iodine are
observed as well as brown nitrogen dioxide.
The lead nitrate-iodine solution is chilled and filtered
to produce a brown solution of lead nitrate and solid
iodine. The iodine goes to [ g ] . The brown lead nitrate
solution is extracted with cyFiohexane to yield a colorless
nitrate and iodine in cyclohexane. The lead nitrate goes
and the cyclohexane to (291.
to
-
[XI
-
31
-
Experiment 17:
Preparation of Silver Nitrate from Silver Halides
Purpose:
The experiment demonstrates the economically important
reclamation of silver from insoluble silver halides.
Description:
Any mixtures containing silver are treated with zinc
and sulfuric acid to produce silver metal. The silver
is separated by filtration and dissolved in nitric acid
to produce silver nitrate. The filtrate is processed
in [UI.
Comments :
The silver nitrate produced here is used in
This
experiment serves to further purify the silver by chloride
formation and subsequent dissolution in ammonia.
[z].
- 32 -
Experiment 18: Chemical Separation of Zinc
Z ~ C JO ~
Na I
7
-
1191
Purpose:
This experiment illustrates a chemical separation
technique by precipitation of the carbonate of zinc
in order to isolate the zinc from other chemicals.
Description:
Aqueous zinc from all sources is treated with sodium
carbonate to precipitate zinc carbonate which is separated
by filtration. The Zinc carbonate goes to [TO].
- The
25].
iodide-containing filtrate goes to [ -
- 33 -
Experiment 19: Redox Titration
m4=
Purpose :
This experiment will demonstrate a redox titration. It
also serves todetermine the concentration of iodide ion
produced by [ 2 5 ] the previous year to assure that exsss
-I .
ammonium iodide is used in silver precipitation in [i5
Description:
A redox titration is carried out using bleach as the
oxidizing agent and starch indicator. The solution is
back titrated with sodium thiosulfate to reduce the
iodine and from the data the amount of iodide ion
originally presentis determined. The chemical products
are processed in [ZJ.
- 34 -
Experiment 20: Electrolysis of Zinc Sulfate
ZnC03
H2S04
,
ZnSOq
CO21
elec.
v
znh
>- ,
H2S04
wash
_j
[IT]
[Tg
Purpose :
The experiment illustrates the processof electrolysis
and reclaims the zinc necessary for
Description:
Zinc carbonate and sulfuric acid are charged to a beaker
and the zinc sulfate electrolyzed using platinum or
carbon electrodes. Beautiful fractal crystals of zinc
are produced. A battery charger is used as a source
of D.C. current. The zinc is in a form such that a
large amount of surface area/mass is pEsent, which
renders it particularly suitable €or
[LA.
[LA.
- 35 -
Experiment 21:
Na2C204
Ion Exchange:
cation
H2C204
exchange?.
Recovery of Oxalic Acid
.
chilr
evap
H2C204
$.
wash
dry
’
[El
Purpose :
This experiment introduces ion exchange to the student
and utilizes it to recover oxalic acid.
Description:
The sodium oxalate solution from the titration experiment
is passed through a cation exchanger to convert
sodium acetate to oxalic acid. This is then purified by
The
fractional crystallization and reused in
mother liquor is evaporated to dryness and ignited to
carbon dioxide and water. Alternatively, the ion exchange
may be skipped and the sodium oxalate evaporated to dryness
and ignited. The ash would then be dissolved in water and
go to
.*..-:::.
[a
[E].
[,so]:..
-
- 36 -
Experiment 22:
Kinetics:
~ a I 0 3 NaHS03,
starch
Clock Reaction
12.starch
(Na2S04)
boil
'
I2 4
(K2SO4)
chill
f ilte?
124
K2SO4
I2
[ZI
cyclthexane
phase separate
Purpose :
This experiment will demonstrate the effects of temperature
and concentration of reactants on the speed of a chemical
reaction.
Description:
Potassium iodate and sodium bisulfite are reacted in the
presence of starch. A deep blue suddenly appears after
some seconds. This delayed reaction permits the observation
of the effects of time and concentration on the time of
color production. The starch-iodine adduct is destroyed
by boiling which precipitates iodine. This product is
chilled and filtered to remove solid iodine, and the
filtrate is extracted with cyclohexane to r z o v e dissolved
iodine. Solid and dissolved iodine go to &I.
Comments :
Students are always intrigued by the sudden appearance of
color in this reaction. Boiling of the solution then
causes the color to vanish.
- 37 -
Experiment 23:
0-c
r
0
-
Synthesis of Methyl Benzoate
MeOH
+
H2S04
OH
Phase
-te
0-f
-
()Me
Na2C03
(MeOH)
( 2s04 )
MeOH
0-C
.
0-C
**
- OMe
MeOH
Na2S04
04,
- ONa
0
]
-
0-CP- OMe
(Org.)
- ONa
1
I
b
a
wash
0-c"
0-C"'
filter
MeOH
Na2S043
-Me
-OH
=
____)
1241
[301
Purpose :
This experiment permits the student to carry out a simple
organic reaction which generates a pleasant smelling
product.
Description:
Benzoic acid and methyl alcohol are placed in a large
test tube equipped with a glass tube and stopper to act
as an air condenser. Sulfuric acid is carefully added
to the mixture. The tube is placed in a boiling water
bath and the mixture is reflexed for about one-half
hour. The mixture is cooled, washed with sodium carbonate
and phase separated. The organic layer is crude methyl
benzoate. This is washed with water. The aqueous phases
are combined and acidified to produce unreacted benzoic
acid.
- 38 -
Experiment 24:
Saponification of Methyl Benzoate
0
0-C*
- OMe
NaoH
,
0
$-C
4
0
-
ONa
(MeOH)
H2S04
~
0-c4 - OH
cool
filter MeOH
[
2
3
1
Na2C03+ MeO€I]+
to
neutral ~iiiSO4
Purpose:
Another simple organic chemistry experiment that can be
performed by introductory students. The chemistry of
ester formation is punctuated by its reverse process
in this experiment.
Description:
The ester is placed in a flask equipped with an air
condenser with sodium hydroxide (2M) and refluxed for
30 minutes or until the mixture is homogeneous. The
mixture is then cooled with running water and carefully
acidified and heated to boiling to drive off methanol.
It is then again cooled in running w s e r and filtered.
The recovered benzoic acid goes to [231.
-
- 39 -
ral,
Experiment 25:
(NH4Cl)
(NaC1)
Preparation of Iodine
I
filter
Purpose :
Description:
(NaCl)
cyclohexan
phase sepa?ate
.This experiment-demonstrates a-commercial process for
the preparabion of iodine from seaweed and brine, and
serves to separate iodine from-salts.
Iodine and iodide ion from all sources are combined and
treated with nitrite ion to preferentially oxidize iodide
to iodine. The iodine is removed by filtration and goes
The filtrate is extracted with cyclohexane
to
and phase separated to remove dissolvg iodine. The
cyclohexane-iodinesolution goes to [La].The extracted
filtrate goes to
[E].
[z].
- 40 -
Experiment 26:
Preparation of Ammonium Iodide from Iodine
Purpose:
This experiment illustrates a chemical synthesis and
serves to convert the iodine generated in previous
experiments to ammonium iodide, and iodide salt useful
in our experimental scheme.
Description:
Iodine is converted directly to ammonium iodide by
reaction with hydrogen peroxide and ammonium hydroxide
by mixing at ambient temperature. The cyclohexane layer
is removed by phase separation and goes to [ 2 9 ] . The
ammonium iodide is freed of ammonia by boiling and then
is rendered slightly acidic to neutralize any ammonia
remaining. This anrmmnium iodide solution is used the
4 I.
following year in [-
- 41 -
Experiment 27:
Synthesis of Sodium Nitrate from Nitrogen (11) Oxide:
Demonstration
Purpose :
This experiment illustrates some chemistry of nitrogen
and isolates nitrates produced in previous experiments,
Description:
A l l nitrate solutions from previous experiments are
combined and after iodide ion is removed are concentrated.
Reaction with copper is initiated by heat, as evidenced by
evolution of brown nitrogen dioxide. The reaction is
carried out in a flask equipped with thistle tube and gas
exit, which is conducted through two other traps containing
cold aqueous sodium hydroxide. The nitrogen dioxide is
absorbed, producing sodium nitrate and nitrite. Unreacted
copper is removed with forceps and the reacted copper is
precipitated by the addition of iron o r aluminum. The
solution is filtered to remove the copper, and the filtrate
goes to
[z].
Comments :
This experiment tends to concentrate the nitrates and
nitrites so that they may be more easily handled in
This experiment may best be a demonstration experimenL
28].
lasting several periods and done in connection with [ -
[E].
-
42
-
Experiment 28: Kje.fdahl Determination:
Demonstration
Purpose :
This experiment illustrates a classical analytical
procedure, and converts nitrates to harmless ammonium
salts.
Description:
The nitrates and nitrites are heated in strong base with
an aluminum catalyst to convert them to ammonia. The
ammonia is t r a m e d in waste base using the apparatus
described in [ 7 1 . Students then titrate the unreacted
sulfuric acid and determine the amount of nitrogen
originally present in the solution. The digestion
portion of the experiment is performed by the instructor
as it involves concentrated base.
- 43 -
Experiment 29:
Distillation
Cyclohexane
( impurities)
distill
methanol
ethanol
(impurities)
distill
cyclohexane
(impurities)
methano1
ethanol
(impurities)
___)
A
>-
l-5- 1
bo-]
k51
-
112.1
bo-1
Purpose :
This experiment illustrates a physical method of
purification of liquids.
Description:
Using student-built condensers and distillation columns,
cyclohexane is separated from any impurities. Methanol
and ethanol are separated by fractional distillation.
Connnents :
The bottoms from these distillations may be added to the
distillation charge for the following year. They are
expected to be minimal. The separation of methanol from
ethanol is not critical. The methanol is used to synthesize methyl benzoate and any ethyl benzoate formed
will not interfere with the experiment in any way. The
ethanol is used in conjunction with the fractional
crystallization of potassium dichromate. The presence
of methanol in the ethanol used will not affect the
results of that experiment.
-
44
-
Experiment 30: Separation of Salts
WASTE ACID
WASTE BASE
to neutral’
filter
K+
Mg2+
K+
Description:
reduce ,Na2S044
All waste acids and waste bases are combined carefully and
neutralized with sulfuric acid or sodium hydroxide as needed.
Calcium sulfate preckitates, is removed by filtration and
goes to Experiment b 3 ] . The solution is basicified with
sodium hydroxide andTeated. Ammonia is expelled and is
absorbed in dilute sulfuric acid with phenolpthalein added to
indicate exhaustion. Ferric hydroxide and magnesium hydroxide
coprecipitate, are removed by filtration, and dissolved in
sulfuric acid. The sulfates are treated with activated (that
is, pretreated with hydrochloric acid) aluminum. The undissolved aluminum is removed and washed, and the solution heated
t o allow any remaining aluminum to dissolve. The precipitated
and the
iron is removed by filtration and goes to
sulfates of aluminum and magnesium go to [AI.
m,
- 45 -
Experiment 31:
Preparation of Alumina, an Abrasive, Desiccant and Catalyst
Purpose:
This experiment illustrates the industrial preparation of
alumina and serves to process aluminum from various
experiments.
Description:
Aluminum compounds from various experiments are neutralized
with sulfuric acid to the point at which aluminum
hydroxide precipitates. After filtration of a small amount
of the mixture, the filtrate is tested for completeness of
precipitation. When precipitation is complete the aluminum
hydroxide is removed by filtration and ignited to alumina.
A small amount of magnesium oxide will also be present.
- 46 -
Experiment 32:
NaCl
Preparation of Bleach:
1
Demonstration
NaClO
Purpose:
As stated
Description:
The preparation of bleach is demonstrated by electrolyzing
the two salts to yield chlorine and alkali metal hydroxides
which react together to yield bleach. This is used to bleach
cloth and flowers. The destruction of microorganisms by
bleach is demonstrated with a culture of protozoans and a
microscope. The bleach is then combined with stock bleach.
- 47 -
Experiment 33: Applications of Industrial Chemicals
CAS04
. 2H20
( W 4 )2SO4
ALUMINA
Fe
02
I__)
A
,
, 2CaS04 . H20
Plant
Plaster of Paris
[CELLS ]
Fiber
[PAPER
H20
Fez03
[CAST]
1
[ALUMINA-H70 ]
Linseed
Oil
0
~
ALUMINA
[PAINT ]
Purpose:
This experiment will illustrate some practical applications
of various chemicals produced from Experiment
Description:
Gypsum is heated to produce plaster of paris from which
a mold is constructed. Ammonium sulfate is utilized
as a fertilizer for a laboratory plant. Sodium and
potassium sulfates are used to size paper. Alumina is
used as a desiccant, and iron powder is oxidized to iron
oxide from which rouge or paint is manufactured.
Comments:
The amounts of these products produced in a year is small.
These may be collected over a period of several years,
and one of the above projects carried out per term.
- 48 -
Experiment 3 4 :
Hydrolysis of Salts
Purpose :
Determination of the pH of salts
Description:
Dilute (0.1M) solution of various salts are tested for
pH. General rules are developed to predict basicity,
neutrality or acidity. Suggested salts are ammonium
sulfate, calcium nitrate, sodium chloride, potassium
acetate, zinc sulfate, sodium iodide, ammonium chloride,
potassium aluminum sulfate.
Comments :
These salts are not used up in the experiment, and may
be used repeatedly each term.
- 49 -
Experiment 35:
Determination of Purity by Melting Point
Oxalic Acid(C0OH)z
from
[El
(COOH)2 dissolve)
melting
point
in H20
determination
evaporate
cool, filter
(COOHI2
co2
ignire’H20
Purpose:
The purpose of this experiment is to illustrate how the
purity of an organic solid may be determined by melting
point. This serves to identify the purity of the oxalic
acid produced from [XI
- and to repurify it as necessary.
Description:
The melting point of the oxalic acid produced in
is determined by placing it in a capillary tube along
with another tube filled with oxalic acid of the desired
purity. The tubes are attached to a thermometer, placed
in an oil bath and heated until they melt. If the oxalic
acid is not of sufficient purity for
it is
recrystallized and the purity established by repeating
the described procedure.
[u]
[E],
- 50 -
DISCUSSION:
EXPERIMENTAL DETAILS
In this discussion, the various experiments proposed for the
laboratory manual will be discussed and examined. along with the reasons
and experimental results that led to their tentative adoption.
Letters
and numbers in parentheses refer to the laboratory notebook reference, for
example (G-1-77-11) refers to the second experiment in laboratory notebook 1,
page 77.
Experiment 1:
Preparation of Oxygen
The reaction initially considered for the preparation of oxygen was the
classic manganese dioxide catalyzed thermal decomposition of potassium
chlorate.
The compound is mixed with a small amount of manganese dioxide,
placed in a test tube and heated to yield oxygen, which is collected over
water and tested in various ways to determine its chemical properties.
procedure initially proposed was as follows:
The
after carrying out the above
described procedure, heating would be continued to quantitatively convert
all of the potassium chlorate to its thermal decomposition product, potassium chloride.
Evidence of completion of reaction would be the cessation
of oxygen evolution.
The residue would then be extracted with water to
remove the potassium chloride, and the remaining manganese dioxide would be
removed by filtration and reused the next year for the same reaction.
This thermal decomposition of potassium chlorate has been in common
useage for decades as the preferred method of oxygen preparation.
however, a serious drawback.
It has,
Potassium chlorate can react explosively if
heated in contact with reducing agents, such as dust, cork or rubber, all
possible contaminants in an introductory class.
It has been suggested that
the largest contributor to chemical accidents in academic laboratories is
the decomposition of a-chloraEe!
- SinaeisafetyLis: of :.primary-cnncern.
other
methods of oxygen preparation were investigated.
A laboratory preparation of oxygen from hydrogen peroxide is described
by Hein, Best, and Miner.2 Thirty percent hydrogen peroxide is diluted
to 9Xw and its decomposition is catalyzed by manganese dioxide. The method,
howevern still possesses a safety risk, and that is the possibility of
--chemical burns resulting from skin contact of peroxide in this concentration.
1
2
Burns, C.J., J.Chem. Educ.,Z, 508 (19561, Winderlich, R.J., J.Chem.
Educ., 27, 670 (1950).
Hein, Best, and Miner, "Foundation of Chemistry in the Laboratory",
6th Edition, Brooks-Cole Publishing Co., p.21
51
-
-
Consequently, an investigation was undertaken to determine if nonhazardous commercially available 3 % hydrogen
~
peroxide could be used to
generate oxygen.
A preparation was attempted (6-1-77-1),
following the
procedure outlined by Hein et a1 but substituting 3% peroxide for the 9%.
Into a 250cc flask equipped with thistle tube and gas exit tube was placed
25 mls H20 and approximately 0.5g of manganese dioxide.
The mixture was
heated (here the procedure departed from Hein) and 3% hydrogen peroxide
was added in 15-30 ml portions as needed to keep approximately one hour.
Evolution was somewhat sluggish, even with the heat.
An additional pro-
blem was that manganese dioxide was occasionally forced up the thistle tube
a few centimeters due to surges in pressure, and there was sone danger of
plugging the tube, which acts as a pressure relief device.
The procedure was repeated (G-1-77-11] but without preliminary addition
of the 25mls of water.
This improved the oxygen evolution rate.
An addi-
tional flask was added with an inlet tube reaching near the bottom of the
flask. to act as a trap for any water that distilled from the reaction
mixture.
This second flask was charged with 50cc of water.
It also served
as a monitor of the rate of oxygen evolution, as the bubbles of oxygen were
easily observed.
With the elimination of the initial 25mls of water charged
to the first experiment, oxygen evolution was more rapid and four 150 mls
bottles were filled in approximately 45 minutes.
Four 15Oml bottles is
This procedure still had the
.
disadvantage of the necessity of reclaiming the catalyst and the attendant
sufficient quantity for the usual experiments.
problem of potential plugging of the thistle tube.
Wikoff and Brown3 report a preparation of oxygen using yeast as a
catalyst.
Their preparation called for yeast cake.
This method was explored.
substituting dry yeast cake macerated with water to make a slurry (G-1-78-1).
The reaction flask was charged with 15Occ of hydrogen peroxide, and the
macerated yeast slurry added to the peroxide.
to avoid denaturing the yeast.
The mixture was not heated
The reaction produced several bottles of
oxygen in approximately 30 minutes time but the reaction was sluggish and
other methods were explored.
Heavy metals will catalyze the decomposition of hydrogen peroxide.
With
_this in mind, a homogeneous catalytic reaction was investigated using
copper chloride as a catalyst.
3
Using the apparatus previously described,
Wikoff, H.L. and Brown, J.B.,
J. Chem Educ.,
- 52 -
2, 1.434 (1926)
approximately 0.5g of the chemical was dissolved in 150cc of 3% hydrogen
peroxide, and the mixture gently heated.
Evolution of oxygen was not
vigorous, and the preparation (G-1-78-11] was discontinued.
Also investigated was the acid catalyzed decomposition of hydrogen
peroxide using iodide ion as ~ a t a l y s t .Approximately
~
0.5g of potassium
iodide was charged to 150cc of 3% hydrogen peroxide and 20cc of 3m H2SO4
added by means of the thistle tube (G-1-80-11).
No visible reaction occurred.
The mixture was warmed to about 30C but there was still little evidence of
reaction.
The feasibility of decomposing 3% hydrogen peroxide using Fe/Fe304 as
catalyst was investigated by charging 15Omls of hydrogen peroxide to the
reaction flask previously described and adding a rusty iron nail.
action mixture was gently heated.
The re-
Oxygen expelled was again passed through
a second flask containing 5Oml of water to act as a trap for any steam
generated and to serve as an indication of oxygen flow.
As soon as heat was
applied, bubbles of oxygen were observed to form on the surface of the nail,
and evolution of oxygen was vigorous.
Four 150cc bottles of oxygen were
collected in about 15 minutes.
The reaction was repeated by a laboratory assistant with the same
results (61-79-11).
This preparation appears satisfactory and is an inno-
vative and improved procedure over the classical method for the following
reasons :
Safety hazards are considerably reduced over the thermal decomposition
of potassium chlorate or generation from concentrated peroxide solutions.
The expense of the catalyst is essentially eliminated.
The 3% hydrogen peroxide is easily available.
The catalyst is easily removed from the solution without filtration.
The reaction product is water.
The oxygen formation on the catalyst is easily observed and is
convincing evidence to the student of catalytic activity of the iron.
4 KOtz, J.C. and Purcell, K.F., "Chemistry and Chemical Reactivity"
1st Edition, Saunders College Publishing CO., Philadelphia, PA.
1987, p.475
- 53 -
Experiment 2:
Preparation of Hydrogen
Hydrogen may be conveniently prepared in the laboratory by displacement
from a dilute acid by a metal, electrolysis of water, or reaction of an
active metal with water.
The amounts of hydrogen produced by the latter
methods are generally insufficient in the usual allotted time to be conveniently tested by the introductory student and the method of choice, for
this reason,is usually the reaction of zinc with acid, and that is the
method that is considered here.
A typical preparation5 is as follows:
Zinc
metal is placed in a flask equipped with a thistle tube and gas outlet tube,
covered with water, and concentrated sulfuric acid added slowly until generation of hydrogen ensues.
The hydrogen is collected by water displacement
in a pneumatic trough.
The major work carried out on this experimental method concerned not
the hydrogen preparation itself, but the reclamation of the zinc from the
zinc salts produced in the reaction.
Two methods were considered, the
electrolysis of zinc salt solutions and the precipitation of the hydroxide
followed by the reduction of the hydroxide to zinc metal with carbon at
elevated temperature.
Of the two methods the former is the simpler and was
the method of choice.
For an introductory chemistry laboratory there are two simple sources
of DC current for electrolysis, an automobile battery or an inexpensive
battery charger ($30
-
$50) available from auto supply houses.
All experi-
ments conducted here used a Viking Model VA 7612 Solid State Battery Charger
on the 12V setting.
A typical hydrogen preparation was carried out by charging 10.7g
mossy zinc to a flask, covering the zinc with water (approximately 50cc)
and adding lOcc of concentrated sulfuric acid.
vigorous.
Evolution of hydrogen was
The reaction of the zinc was completed by the addition of 5cc
more sulfuric acid.
The reaction mixture was filtered to remove suspended
particulate matter to yield a colorless filtrate.
The following series of experiments dealt with finding a practical
simple method of recovering zinc from the zinc sulfate solution by electrol-
yeis.
In the first experiment (G1-34-11),
the zinc sulfate was electro-
€yzed using the battery charger and two electrodes consisting of pencil
, "Laboratory Manual for Chemistry". 1st Edition,
Allyn and Bacon, Inc., Newton, Mass. , 1983, p. 89
5 Wagner, M.
- 54 -
graphite.
No effort was made to hold current density or concentration
constant.
Hydrogen was evolved at the anode, and a dark zinc deposit
formed at the cathode, but rather slowly.
The procedure was repeated (G-1-34-1111 but continued for a period of
approximately 5 hours.
At the end of this period the pencil graphite
electrodes had disintegrated to the point that one of the electrodes was
no longer immersed in the solution.
A search was made for more durable
electrodes.
Two strips of 304 stainless steel were obtained and the procedure
described above was repeated (G-1-68-11).
The strip used as the anode
slowly dissolved and the solution became colored, presumably due to cobalt
or nickel present in the steel.
Since these metals are undesirable con-
taminants, this procedure was discontinued.
However, use of stainless for
the cathode was promising.
Several carbon rods used in arc welding were obtained.
These consist
of a carbon rod plated with copper with approximately a centimeter of bare
carbon at the tip.
Using these rods as electrodes, the anode again dis-
integrated as the electrolysis was continued (Gl-68-11).
A platinum elec-
trode was then used as an anode, and the carbon rod as the cathode (G-1-69-1).
This arrangement proved satisfactory.
When two platinum electrodes were
used (G1-69-111), constant useage and removal of zinc from the cathode
caused mechanical damage to the electrode, with the platinum eventually
becoming separated from the copper wire attached to it.
As the zinc is removed from solution by electrolysis, it deposits in
beautiful fractal crystals resembling ferns.
As this deposit grows and
branches, it tends to grow toward the anode, which could cause eventual
ehort circuiting of the charger.
was sought to prevent this.
zinc cathode from the anode.
Consequently, a simple inexpensive method
A porous cup (6-1-70-1) was used to isolate the
Since a quantity of electrolytic zinc produced
by this method was needed for work on zinc reduction of silver halides
(see Experiment 17), electrolysis of zinc solutions was carried out intermittently over a period of several weeks.
The porous cup was frequently
removed from the solution and set in a beaker.
Zinc sulfate crystals
frequently formed on the surface of the cup and presumably within the pores.
In any case, after about two weeks useage the porcelain cup began to
- 55 -
disintegrate.
This may have been caused by mechanical strain due to
internal crystal formation in the pores.
If this were the case, immersion
in water when not in use would prevent this.
In any case, isolation of
the electrodes were sought.
A simple salt bridge between cells has been suggested by Howell, et al. 6
A length of 1/8 in. ID (1/32 wall thickness) amber latex tubing was filled
with a hot solution of 3% agar in 0.1M zinc sulfate and the solution allowed
to cool as gel formation occurred (G-1-71-11].
This cell bridge was used
to span two beakers containing zinc sulfate, an electrode in each.
Inadequate
electrical current was obtained between the cells as judged by the rate of
hydrogen evolution from the anode.
In another experiment (6-1-71-1) a 6 inch length of tubing was filled
with the zinc sulfate solution and plugged with cotton at both ends and
placed in two beakers with electrodes.
Once again, current density was
severely diminished.
An experiment (G-1-68-1 and G-1-70-11) was conducted in which a beaker
was roughly divided into two cell compartments with a piece of polycarbonate
plastic cut loosely to shape.
The fit permitted solution circulation but
isolated the anode from the cathode directly.
An aquarium aerator provided
agitation of the solution to minimize electrode polarization.
Fractal
growth of the zinc in this arrangement was virtually eliminated and the zinc
growth took place directly on the cathode.
This technique appeared to
work well.
Having achieved a simple method of zinc recovery by electrolysis, an
experiment was conducted to determine the rate of zinc recovery.
At this point, consideration was given to avoiding the entire problem
of zinc reclamation by employing aluminum instead of zinc.
Zinc has pre-
sumably been the metal of choice in the past because of its high rate of
reactivity, which is necessary to produce the quantity of hydrogen required
in the short amount of laboratory time generally allotted to beginning classes.
Aluminum has a lower reduction potential (-1.66V) than zinc (-0.76) and
consequently should have a high reactivity with dilute acid, as well as
producing a relatively non-toxic product, alumina.
Aluminum is ubiquitous
and scrap may be obtained at essentially no cost (scrap cans, foil, etc.).
It does have an oxide coating which renders it unreactive to most dilute
6
Howell, B . S . ,
Cobb, v.S.
and Hacksima, R.A.,
- 56 -
J Chem Educ.,
60, 273
(1983)
acids.
This coating is easily removed by a quick pretreatment in hydro-
It is then quite reactive toward dilute sulfuric acid. These
experiments demonstrate that fact.
A sample of aluminum foil was placed in about lOOcc of lOXw sulfuric
acid. No reaction was apparent. In another beaker was placed another lOOcc
chloric acid.
of sulfuric acid, but about 3ml of hydrochloric acid were added.
Hydrogen
evolution began almost immediately.
In a 250cc flask was placed approximately lg of aluminum foil and sufficient water was added to cover the foil (G-1-67-1).
Two ml of conc. sulfuric
acid was added to the flask.
No reaction was observed.
hydrochloric acid was added.
No reaction was visible.
Five drops of 6M
Eventually after
addition of a total of 12cc of conc. H2SO4 and 10 drops of 6M HC1 was added,
evolution of hydrogen began and built in intensity.
The evolution became
quite satisfactory for a laboratory preparation.
In another experiment (G-1-67-111), aluminum foil was immersed in
6M HC1 until hydrogen evolution commenced (approximately 10 seconds).
It
was then placed in 10% w sulfuric acid at which point hydrogen evolution
became quite vigorous.
These experiments have led to the tentative design of Experiment 2
in which aluminum metal is pretreated with hydrochloric acid for a few
seconds to remove the oxide coating and then reacted with dilute sulfuric
The reason that hydrochloric acid is not
acid in the conventional manner.
used for the reaction is that fumes from this acid are noxious, and since
hydrogen is prepared early in the beginning chemistry course, good laboratory
skills have not yet been developed, and danger of hydrogen chloride fumes
escaping into the room is greater.
In any event, it is proposed that aluminum be substituted for zinc in
this preparation because of two distinct advantages. It is much less expensive (at todays prices mossy zinc sells for about $30/lb., aluminum for
$.30/lb.)
and produces a relatively non-toxic product, aluminum sulfate,
which can be converted t o alumina, a major component of earth.
The procedures developed for zinc recovery by electrolysis will not
--
be for nought, however, but will be utilized in Experiment 20 for recovery
of zinc produced from silver halide reduction.
- 57 -
Experiment 3:
Chemistry of the Alkaline Earths
This experiment has been developed to utilize a number of reactions of
the alkaline earth metals and their compounds in a sequential fashion so
that the reactions illustrate the various properties.
been divided into two sets.
The experiment has
In Set 1, calcium and magnesium are reacted
with water to produce the hydroxides, which are shown to be bases by litmus
paper, and insoluble in water.
These are dissolved in hydrochloric acid to
demonstrate the solubilities of the chlorides, and calcium sulfate is precipitated with sulfuric acid, whereas magnesium sulfate remains in solution,
illustrating a means of separation of the two metals.
The resulting mixtures
go to Experiment 30.
In Set 2 , calcium and magnesium metals are reacted directly with iodine
to illustrate the direct combination of the alkaline earths with a halogen.
The reaction mixtures are filtered to remove unreacted iodine and the iodides
are treated with sodium carbonate to precipitate the carbonates and the
sodium iodide is removed by filtration.
The carbonates are then dissolved in
acid to illustrate their solubility in that medium, and the product goes to
waste acid.
A l l of the reactions described are straight-forward with the exception
of the reactions of the alkali metals with iodine.
To investigate this
reaction 0.24g magnesium and 3.00g iodine were placed in an evaporating dish
with the intention of adding 15cc of water.
As the water was being added to
the mixture the reaction was initiated, resulting in spectacular emission
of purple iodine vapor resulting from the heat of reaction. This would make
a good demonstration reaction (61-107-1).
The reaction was repeated with the following modifications (6-1-107-11).
Iodine (3.03g10.015 mole) was added with agitation.
The mixture immediately
turned dark brown, and a vigorous reaction ensued with evolution of heat.
The dark reaction mixture was filtered to yield a brown filtrate containing
the magnesium iodide and a trace of dissolved iodine.
The solution was
converted to colorless with several drops of 1M sodium thiosulfate and the
colorless magnesium iodide solution was ready for treatment with sodium
garbonate.
- 58 -
The reaction of calcium with iodine was carried out as follows (61-108-1).
Iodine (3.05g10.012 moles) was placed in lOml of water in a 5Oml flask and
calcium turnings (0.60g=0.01 moles) added.
An exothermic reaction ensued.
The resulting mixture was yellow with a small amount of precipitate calcium
hydroxide (white) evident.
The mixture was filtered to yield a yellow
filtrate which was decolorized with several drops of 0.1M sodium thiosulfate
solution.
Presence of iodide ion in the filtrate was confirmed by precip-
itation of lead iodide by reaction with 1M lead nitrate.
The two methods described appear satisfactory for the demonstration
of the reactions of calcium and magnesium with iodine, Set 2.
- 59 -
Experiment 4: Determination of An Empirical Formula
In this experiment, lead nitrate is reacted with ammonium iodide obtained
from Experiment 26 to yield lead iodide, which is removed by filtration, dried,
and weighed.
From the data obtained and the amount of lead known to be in the
lead nitrate, the empirical formula of lead iodide may be calculated.
reaction produces as a by-product potassium nitrate.
The
In the early stages of
this project, an effort was made to eliminate nitrates from the experimental
schemes.
Consequently, consideration was given to reacting lead chloride instead
of lead nitrate with potassium iodide.
chloride and lead nitrate.
The products would then be potassium
Although relatively insoluble in cold water
(O.O95g/cc at 20°C), lead chloride increases in solubility considerably in
hot water (3.34g/cc at 100°C).
rhus, if the concentration is held to this
level or below at 100°C, the lead will be in solution.
* T o investigate this reaction, 0.75% of lead chloride was added to 5Omls
of water and 5Omls of 3% sodium iodide, and the mixture brought to a boil
(61-41-1).
Addition of 25mls of 1M sodium iodide immediately produced the deep
yellow lead iodide.
The mixture was decanted and the precipitate washed with
three 5Oml portions of hot water and the suspension filtered.
was dried and weighed,
A recovery of 90.3% was obtained, losses apparently
due to solubility of lead iodide in hot water.
using 25ml boiling water.
losses (61-41-11).
the lead iodide
The experiment was repeated
A recovery of 100% was achieved, including solubility
This approach to eliminate nitrate ion, was eventually
rejected as cumbersome, and the original experiment using lead nitrate was
adopted, subsequent to methods being developed for the handling of nitrates
(Experiment 27) and iodides (Experiment 25).
7 "Handbook of Chemistry and Physics", Weast, R.C. and Selby, S.M. Editors,
47th Edition, Chemical Rubber Co., Cleveland, 1964, P.B-187.
-_
-
60
-
Experiment 5:
Chemistry of the Halogens
The reader is referred to the Experimental Outline to aid in following
this discussion.
This type of experiment, which surveys a number of chemical
reactions and produces many products, proved to be the most difficult sort of
experiment to adopt to the methods used in this report.
The author will
attempt to guide the reader through the processes which eventually led to
the proposed scheme.
The original intention was to divide the experiment into four sets of
reactions:
1) reaction of halides with calcium ion; this set would contrast
the insolubility in water of calcium fluoride to the solubility of the other
calcium halides, 2) reaction of halide ion with silver nitrate; this set
would contrast the solubility of silver fluoride with the insolubility of the
other silver halides, 3) reaction of insoluble silver halides with ammonia to
aillustrate the varying solubilities, 4) oxidation of halides to the free
halogen with chlorine.
Methods used here were those used for each experiment discussed in this
report.
These were A), create a sequence of reactions instead of the
conventional approach of a number of one-step reactions to minimize the
number of products necessary to deal with and B), incorporate the reaction
products in the rest of the cycle of experiments.
Set 11, above was first considered.
In the original scheme, sodium
fluoride, chloride, bromide and iodide would be reacted with calcium nitrate
in aqueous solution.
These reactions would yield only one insoluble compound,
calcium fluoride, and thus illustrate the marked difference in behavior of the
fluoride ion.
The problem with this is that calcium fluoride is very difficult
to incorporate into further reactions because of its very low solubility and
lack of reactivity.
The following plans were considered.
First, react
calcium fluoride with sulfuric acid to produce hydrogen fluoride as a demonstration reaction.
This was dismissed as too hazardous.
Second, convert
calcium fluoride to silicon tetrafluoride with silicon dioxide and sulfuric
acid, dismissed for the same reason as the first.
Third, bottle the calcium
fluoride for some future experiment, rejected as antithetical to the purposes
__ of this report.
-61
-
Since none of these alternatives were desirable, the decision was made to
avoid use of fluoride ion entirely.
This is justifiable on the basis that the
chemistry of fluoride ion is so different from the other halides that it is
placed in a different group for purposes of qualitative analysis.
reason Set 1) was rejected.
Next considered was Set 2.
For this
In this series of reactions the original
intention was to react silver nitrate with sodium fluoride, chloride, bromide,
and iodide.
The silver salts of the latter three halides precipitate, the
fluoride does not.
experiment.
Elimination of fluoride from condsideration-simplifiesthe
Further, if the precipitated silver halides can be easily removed
by filtration, then Set 2 could be combined with Set 3 , the reactions of the
silver halides with ammonium hydroxide.
To investigate the filterability of the silver halides, 5mls of 0.1M
silver nitrate was combined with 5mls 1.OM sodium chloride, and the reaction
repe'ated with the bromide and iodide.
The three resulting mixtures were
placed in a boiling water bath for about 5 minutes and the mixtures filtered.
The ageing of the precipitates in the water bath was effective and the resulting curdy precipitates were satisfactorily separated from the mixtures by
filtration (61-23-1).
The satisfactory isolation of the precipitates by filtration meant that
they could be treated with ammonia on the filter, and the solubility of the
silver halides observed.
Silver chloride is completely soluble, silver bromide
much less so and the iodide is insoluble.
The partial solubility of silver
bromide in the amounts of ammonium hydroxide which are practical to use here
presents a problem in that some of the bromide will appear in the filtrate
and some will remain on the filter.
Bromine poses another problem; it com-
plicates procedures later used (see Ehperiments 25 and261 for halogen
recovery.
The decision was made to restrict the halogens investigated to
chlorine and iodine.
This seems reasonable, the two representing extremes of
halogen behavior of the representative halogens, fluorine being atypical.
Separation and recovery schemes are considerably simplified.
In this fashion the sequence of reactions for Set 1 were conceived.
These are the following.
The silver halides of chlorine and iodine are
precipitated, removed by filtration, and treated with ammonium hydroxide.
- 62 -
Silver chloride dissolves, silver iodide does not.
The dissolved silver
chloride is regenerated from ammoniacal solution by acidification with
hydrochloric acid.
The undissolved silver iodide is washed from the filter
paper into hydrochloric acid to demonstrate its insolubility in that medium.
The resulting silver halides are processed in Experiment 17.
Consideration was then given to the oxidation of halides by chlorine
(originally Set 4).
This is an important reaction in halogen chemistry,
since it is used to prepare iodine and bromine industrially. 8Bleach
(aqueous solution of sodium hypochlorite) was chosen as an inexpensive source
of chlorine.
Reaction of bleach with an acid will produce chlorine in situ,
and thus eliminate the hazards of handling chlorine water.
The intent of
the experiment is to demonstrate the colors of chlorine and iodine in an
organic solvent.
This can be used as a test for halogens.
For example,
in carbon tetrachloride, chlorine is yellow, bromine is red, and iodine is
.violet.
The experiment also demonstrates the oxidizing power of chlorine
with respect to halogens of greater atomic number.
To carry out this experiment, bleach and an organic solvent (cyclohexane)
are added to an aqueous solution of the halide.
acid causes the liberation of chlorine.
Addition of dilute sulfuric
In the case of the chloride, the
liberated chlorine preferentially extracts into the organic phase, yielding
a yellow solution.
In the case of the iodide, the chlorine oxidizes the
halogen to free iodine, which extracts into the cyclohexane layer as a violet
solution.
The organic phases may be separated from the aqueous phase using
a simple technique developed.
on the funnel.
A filter paper is prewetted with water while
When a mixture of a water phase and a water-insoluble phase
is poured onto the paper, the aqueous phase easily passes through and the
organic phase is retained by the paper.
The organic phase may then be re-
covered by punching a hole in the filter paper and allowing the phase to
drain into a container or by simply pouring it from the filter.
This method
is used in this report whenever there is a need for phase separation.
The cyclohexane solutions of halides thus prepared are shaken with a
dilute sodium thiosulfate solution to reduce the halogens to halides.
The
halides and cyclohexane solutions of halogens are processed in Experiment 25.
This procedure becomes Set 2.
8 Holzclaw, Jr., H.F.,
Robinson, W.R. and Nebergall, W.H.,
Chemistry, 7th Edition, 1984, p.569
-
63
-
General
A number of experiments were conducted concerning Set 2.
In (61-25-1)
the suitability of phase separation using filter paper was investigated.
Two mls of 0.1M solutions of fluoride, bromide and iodide were treated with
20 drops of chlorine water and 2mls of carbon tetrachloride.
phases became, respectively, clear, red, and violet.
The organic
When the water-carbon
tetrachloride mixture was poured on the filter some phase separation took
place, but the carbon tetrachloride somewhat hindered the passage of the
aqueous phase since it is denser than water and settles to the bottom of the
cone.
Nevertheless, it worked well enough to be promising.
Carbon tetrachloride is now considered undesirable as a reagent for the
introductory laboratory, due to its carcinogenicity and particularly because
of its immediate toxic effect on the kidneys, and the possibility of absorbtion through the skin in toxic q~antities.~
Consequently substitutes for
carbon tetrachloride were investigated.
. Experiment
(61-24-11) repeated the procedure described above (61-25-1)
but using kerosene as the organic solvent.
This worked well as the extractant
but bromine and chlorine solutions were observed to fade in color with time
indicating possible halogenation of the hydrocarbons, an undesirable side
reaction producing toxic compounds.
As a result, kerosene was not used in
later experiments.
In (G1-28-I),
mineral oil, a petroleium fraction, was used in a
repetition of (Gl-25-11.
This proved to have the disadvantage of kerosene
in that halogen solutions decolorized on standing, and in addition it was
difficult to remove from glassware using conventional cleaning procedures.
On (61-24-IV), cyclohexane was used as extractant, and this appears to
be the most acceptable of the carbon tetrachloride altematives tested.
Use of bleach as a chlorine generator was investigated as follows
(Gl-24-111).
Two
m l s of a 0.1M sodium iodide solution was mixed with 2mls of
bleach and 2mls of kerosene, acidified with several drops of dilute sulfuric
acid and shaken.
colorless.
The kerosene became violet in color and the lower layer
The procedure was repeated, using 0.1M sodium bromide.
kerosene layer became red.
The
The mixture was phase separated with pre-
moistened filter paper.
The experiments described were carried out to develop the procedures
-_
used in Set 2 and described earlier. The final proposed procedure was
9 MeinLloow,J., Keeffe, J.R., and Bemstein, R.L., J.Chem. Educ.,
5 8 , All (1981).
-
64
-
carried out for Sets 1 and 2 (61-291).
The total elapsed time was about 4 5
minutes, so it was decided to add another reaction set.
selective oxidation of iodide ion by acidic nitrite ion.
(Gl-261) demonstrates this reaction.
This involved
An experiment
To 2Omls of 0.1M sodium iodide was
added 2mls of 6M nitric acid, 8mls of 0.1M sodium nitrite, and 40 mls of
cyclohexane.
The mixture was agitated and then allowed to phase separate.
Addition of further nitrite produced no more iodine, as judged by the color
of the mixture.
reaction.
Repetition of the procedure with chloride ion produced no
This selective oxidation was chosen as Set 3 of the halogen exper-
iment.
In (61-31-1) the complete experiment was carried out in about 50 minutes,
a practical duration.
A later modification to the experiment was the sub-
stitution of the ammonium salts of the halides for the sodium salts, since
ammonium iodide is generated from Experiment 26.
- 65 -
Experiment 6 :
Reactions of Ions in Solution
This experiment is conducted to illustrate a variety of insoluble substances produced in solution.
The proposed reactions are:
1. reaction of silver nitrate with sodium chloride to produce silver chloride.
2. reaction of mercury (11) chloride with sodium carbonate to produce mercury
(11) carbonate.
3 . reaction of barium chloride with sodium sulfate to produce barium sulfate.
4 . reaction of lead nitrate with potassium chromate to produce lead chromate.
5. reaction of mercury (11) chloride with potassium iodide to produce mercury
-
(11) iodide.
6 . reaction of copper (11) chloride with sodium sulfite to produce copper (11)
sulfide.
7 . reaction of calcium chloride with sodium carbonate to produce copper
carbonate.
8:reaction
of barium chloride with potassium chromate to produce barium
chromate.
A large number of diverse products are generated in an experiment of this
type, making it difficult to fit into the proposed method of this report.
The
final solution to this dilemma was to reduce the experiment to the semimicro
scale so that the amounts of chemicals produced as waste are miniscule.
The
reactions are carried out on a spot plate, which helps the students to keep
track of the reactions.
The total amount of chemicals produced per year were
estimated as follows.
There are eight reactions.
Assuming 0.1M solutions and 2 drops of each
chemical, an average formula weight of 200g/mole, a class of 30 consisting
of 15 students pairslexperiment, the amount of chemicals produced would be:
-
8 reactions x 15 experiments x 2 chemicals x 2 drops x ml
experiment
term
reaction chemical 20 drops
x Liter
lOOOml
x 0.1 mole
liter
x 200g
mole
-
0.48 chemicals
term
The amount of waste generated in 200 years of a chemistry class is
thus 192g.
- 66 -
Experiments 7 and 8: Types of Chemical Reactions I and I1
An
experiment often referred to as the copper cycle has been widely
employedlOas an experiment in introductory chemistry.
This procedure consists
of a sequence of steps beginning and ending with metallic copper, and generally
consisting of the following steps: copper is dissolved in nitric acid, the
hydroxide precipitated by addition of sodium hydroxide, the hydroxide converted to copper oxide by heat, the copper oxide converted to copper sulfate
with dilute sulfuric acid, the copper precipitated by addition of zinc, the
excess zinc dissolved by heat, the copper removed by filtration, dried and
weighed.
The experiment is generally used to demonstrate the principle of
conservation of mass, as near quantitative yields of copper are obtained.
"his is the only procedure found by the author in the literature, a
portion of which was adaptable to the methods of this report directly, due
to its sequential nature.
Here it is used to illustrate the basic types of
ionic reactions; combination, decomposition, single replacement and double
replacement.
The first reaction, an oxidation-reduction reaction of copper with
nitric acid, does not fit easily into the four basic catagories described.
Consequently another reaction of copper metal was sought which would be
relatively simple and which would illustrate a combination reaction.
Copper is known to combine directly with aqueous ammonia to produce the
deep blue tetramine copper (11) ion.
This reaction is catalyzed by the
presence of air. l1To investigate this reaction approximately 5g of copper
foil was immersed in lOOml of concentrated ammonium hydroxide (Gl-8-11). Air
was sparged into the solution, and evolved ammonia was trapped in an adjacent
flask containing dilute sulfuric acid.
overnight.
The reaction was allowed to proceed
The next morning a deep blue solution resulted containing some
precipitated copper hydroxide.
The reaction makes an excellent demonstration
reaction because of the beautiful color generated.
But the conversion of
copper (27.4%) was not complete enough to warrant using this reaction as a
first step.
10 Condike, G.F., J.Cjem Educ.,X, 615 (1975); Weiner, E., "Fundamentals of
Chemistry; Laboratory Studies", 4th Edition, Academic Press, New York,
1980 pp141-151; Davis, Jr.,J.B., MacNab,K.W., McClellan, L.L., and
__ O'Conner, P.R., "Laboratory Manual for Chemistry", 1st Edition, D.C.
Heath and Co., Lexington, Mass., 1982, pps.58-60, Todd, D. and Hobey, W.D.,
J Chem Educ.,%,
177, (1985).
11 Holzclaw, Jr., H.F., Robinson, W.R, and Nebergail, W.H., "General Chemistry"
7th Edition, 1984, p860.
- 67 -
Copper reacts with hydrogen peroxide and acetic acid to produce copper
A procedure is described by Umans and deVos.12Two to 3mls of 20%
acetate.
acetic acid is mixed with 2 to 3mls of 5I hydrogen peroxide and warmed gently
without boiling.
The reaction is initiated by the addition of copper.
The
procedure was carried out as described, using copper powder obtained by a
single replacement reaction of copper sulfate with iron.
This results in a
fine copper powder which has a large amount of surface area, and is typical
of the copper that will be produced in the experiments in this report.
copper was further reduced with a mortar and pestle.
The
Carrying out the reaction
as described, a bluish-green solution was obtained overnight.
The unreacted
copper was removed from solution by filtration, washed with acetone, dried,
A conversion of 3.4% was obtained, too low to be practicable here.
and weighed.
Copper will react directly with sulfur to produce copper (I) and copper (11)
sulfides. 13A reaction was carried out using copper powder prepared as described
above (61-83-1).
Copper (2.93g) and sulfur (3.01g) were intimately mixed by
grinding the two together in a mortar and pestle.
The mixture was placed in
an evaporating dish, and an attempt was made to initiate reaction by striking
a match and plunging it into the mixture, but without success.
The mixture was
placed on a hot plate and heated for about 3 minutes on high heat.
A struck
match plunged into the mixture immediately initiated the reaction.
The initial
incendiary point of black copper sulfide spread to the entire dish in about 30
The reaction was rapid and the yield essentially quantitative (5.90g
seconds.
copper sulfide).
Unfortunately, the reaction does have disadvantages.
Sulfur
dioxide is evolved, and sulfides are created which can generate poisonous
hydrogen sulfide in the presence of acid.
Sulfides can be oxidized in acidic media to sulfur. 14This can be affected
by permanganate, iodide, dichromate, nitrate ion, etc. in the presence of
hydrogen ion.
(61-84-1).
warmed.
A reaction was investigated using 0.5g of copper sulfide
Approximately 5cc of 6M nitric acid was added and the mixture
Evolution of nitrogen dioxide was observed.
sulfide was detected.
No odor of hydrogen
The reaction mixture was filtered, and the residue,
which had the appearance of sulfur was ignited and burned with evolution
of sulfur dioxide,
12 Umans, T. and deVos, W., J.Chem. Educ., 2, 52 (1982).
13 Mahan, B . H . , "University Chemistry", 3rd Edition, Addison, Wesley
Pub. Co., Menlo Park, Ca. (1975) p710
14 Moeller, T., "Qualitative Analysis", 1st Edition, McGraw-Hill, New York
1958, p221
- 68 -
The above procedure is cumbersome.
sought.
A more satisfactory reaction was
Copper will react with the oxygen of the air or pure oxygen to yield
copper oxide.
It was felt that since the processes in this manual produce
a finely divided copper, it might present enough surface to yield good conversions to copper oxide in a reasonable amount of time.
If, instead of air,
pure oxygen were used as a reactant, the reaction time should be shortened.
To acheive this, copper powder (2.07g) was placed in the center of a 2cm.
diameter glass tube 20cm. long and stoppered at either end (61-84-11). An
oxygen generator described in Experiment
supplied oxygen. The copper in
the tube was heated with a Meker burner, and the oxygen generator, charged
with 3% hydrogen peroxide and a rusty nail, was heated to generate oxygen.
The copper in the tube immediately blackened on the surface. The reaction
was carried out in 45 minutes. The tube was allowed to cool and the product
removed, placed in a beaker, and 5Omls 6M hydrochloric acid added. The
mixture was examined 2 days later. At that time, the solid was completely
dissolved and a green copper (11) chloride solution was obtained.
The above reaction was repeated with the following modifications (61-81 I).
The equipment was simplified by replacement of the glass tube with an evaporating dish to contain the copper.
Aluminum foil was placed over the dish,
form-fitted around the top, and a small hole made in the center to permit
entry of the tube containing oxygen.
A flask acting as steam trap and bubbler
was installed between the oxygen generation flask and the dish containing the
copper.
Oxygen from the generator was conducted through a tube in this flask
which terminated close to the bottom of the flask, which was charged with 5Omls
of water.
A gas exit tube permitted the oxygen to exit the trap and enter
the vapor space above the copper in the dish.
excess oxygen escaped to the atmosphere.
Since the foil was a loose fit,
The dish was charged with 2.21g
copper powder, heated, and oxygen passed over the copper for 30 minutes.
The dish was allowed to cool, and the contents transferred to a 250cc flask,
and 5Oml of 6M hydrochloric acid was added.
t o yield a
The residue dissolved completely
gum solution of copper (11) chloride.
This conversion of copper to copper oxide (combination) and then to
copper chloride (double replacement) is quite satisfactory for the purposes
Intended.
The two reactions require about 50 minutes for completion.
I
69
-
It
now remained to employ the standard sequence of reactions of the copper cycle
to complete the experiment.
1.
These are
Conversion of the chloride to copper hydroxide with sodium hydroxide
(double displacement).
2.
Conversion of the hydroxide to the oxide by heat (decomposition).
3.
Conversion of the oxide to the sulfate with sulfuric acid
(double replacement).
4.
Conversion of the sulfate to copper metal with aluminum
(single replacement).
This series of reactions was carried out in toto.
and physical changes is quite striking (61-87-1).
made to the procedure.
The sequence of chemical
A further modification was
An experiment (61-87 11) was carried out to determine
if aluminum (oxide coating removed by treatment with hydrochloric acid) could
be ‘substituted for zinc (a more toxic and expensive element).
Pretreated
aluminum (log) was added to 1 5 h l of a solution of copper sulfate and heated.
Conversion of the copper sulfate (5Xw aqueous) was complete in about 5 minutes,
as evidenced by the change of color from blue to colorless.
The excess
aluminum was removed by forceps and 25mls of 1M H2SO4 was added to dissolve
remaining unreacted aluminum.
The mixture was heated for about 5 minutes
to complete reaction, and the reaction mixture filtered to yield 3.0g copper.
A summary of the reaction sequence follows: 1) copper metal is treated
with oxygen to yield the oxide (combination). 2) the oxide is converted to
the chloride with hydrochloric acid (double replacement).
3) the chloride is
converted to the hydroxide with sodium hydroxide (double replacement).
4) the hydroxide is heated to yield the oxide (decomposition).
5) the oxide
treated with sulfuric acid yields copper sulfate (double replacement).
6) the copper sulfate reacts with aluminum to give copper (single replacement).
The reaction sequence occupies about 2 50 minute lab periods. It illustrates
the four types of simple ionic reactions, conservation of mass, and some
chemical properties of copper and its compounds.
-
70
-
Experiment 9: Activity Series
A typical activity series experiment is found in Hein, et all5 In six
separate experiments, copper is reacted with silver nitrate, lead with copper
nitrate, zinc with lead nitrate, zinc with magnesium sulfate, copper with
sulfuric acid, and zinc with sulfuric acid.
The approach used in the work
reported here is the creation of a sequence of reactions so that products
are minimized.
on paper.
A number of different reaction schemes were initially proposed
Most were quite complicated.
It was then realized that reduction of
this experiment to a sequence of reactions in which a lesser metal is replaced
by the more active would not only greatly simplify the scheme, but make the
point of the experiment much more accessible to the student.
The most promising sequence was:
.
1)
silver nitrate plus copper yields copper nitrate plus silver.
2)
copper nitrate plus lead yields lead nitrate plus copper.
3)
lead nitrate plus zinc yields zinc nitrate plus lead.
The zinc would ultimately be recovered from the nitrate salt by electrolysis.
The silver is eventually converted to silver nitrate, the copper to copper
nitrate, and the lead to lead nitrate.
The cycle is complete, the increasing
activity of the metals easily comprehended.
There are several problems with the scheme.
Can the reactions be completed
in a 55 minute period? Can the unreacted more active metals used in the single
replacement be separated from the less active metal produced from the reacting
salt? For example, suppose an excess of copper powder is added to silver
nitrate to produce a finely divided silver, the copper being in excess to ensure
completeness of reaction.
and copper.
Then the mixture contains both finely divided silver
These are not easily separable.
An excess of the more active
metal i s necessary for completeness of reaction, otherwise the reaction mixture
is made even more complicated by containing unreacted salt.
problems to be addressed.
These were the
Replacement reactions are usually carried out at ambient temperature in
An experiment was carried out to explore the effect of
temperature increase on the reaction rate (61-48-1). Five ml of 0.1M silver
introductory courses.
nitrate were placed in a test tube and copper in the form of rolled up copper
screen was added (2g).
15
The mixture was placed in a water bath at 100°C for
Hein, M., Best, L.R., and Miner, R.L., "Foundations of Chemistry in the
Laboratory", 5th Edition, Brooks/Cole Publishing Co., Monterey, Ca.
1982, p.75.
- 71 -
approx,mately
complete.
10 minutes.
At the end of this period, the reaction appeared
The unreacted copper was removed with forceps and washed free of
silver with water.
The copper was set aside to dry on a paper towel.
The
reaction mixture containing the silver and copper nitrate was filtered, and
the recovered silver treated on the filter with 6M nitric acid, yielding
silver nitrate.
The filtrate, containing copper nitrate, was treated with
lead shot and placed in a boiling water bath and the tube occasionally removed
and shaken.
The mixture became colorless in about 10 minutes, indicating
completeness of reaction.
The lead was removed with forceps and washed free
of copper with water into the mixture.
retrieve the copper powder.
The mixture was then filtered to
To the lead nitrate solution was added a strip of
zinc metal and the sequence repeated.
The lead was obtained as a powder.
The products resulting from the reaction sequence are then zinc nitrate,
silver powder, copper powder, and lead powder, as well as unreacted pieces of
copper, lead, and zinc, which are reused.
The zinc from the zinc nitrate may
be recovered by electrolysis (Experiment20).
verted to silver nitrate in Experimentl7.
The silver powder will be con-
The copper powder will be used as
starting material for the copper cycle, Experiment 7.
The lead powder will be
used as starting material for Experiment 1 4 , The Chemistry of Lead.
The problem of separation of the excess of active metal from the precipitated less active metal is resolved by using the more active metal in
bulk(i.e.,
not powdered) form.
The former may then be easily retrieved with
forceps, and washed free of the less active metal.
laboratory time (maxi”
The problem of limited
time is assumed to be 55 minutes) is overcome by
conducting the reactions near the boiling point of water.
Another modification
made to the reaction scheme is the use of aluminum or iron to substitute for
zinc, which reduces the amount of zinc necessary to recover by electrolysis
(Experiment 20 1.
Experiment 10:
Titration
In this experiment, oxalic acid is used as a primary standard to determine
the normality of sodium hydroxide, used as a titrant in the experiment. This
experiment is mentioned here since oxalic acid is part of a cycle represented
by--hperiments10, 21, and 36.
- 72 -
Experiments 11 and 12: Chemistry of Chromium and Fractional Crystallization
A sequence of reactions involving chromium was sought, which could be
incorporated into a cycle.
These reactions should be visually interesting
to the beginning student, and demonstrate solubilities, chemical properties,
and, ideally, several oxidation states of the element.
Solutions of chromate ion are bright yellow.
acidified, the orange dichromate ion is formed.
When these solutions are
This is a powerful oxidizing
agent and is reduced readily to chromium (III)!6 Reaction in acidic solution
of chromate or dichromate with hydrogen peroxide yields a peroxo compound
(contains 0-0 bonds).
atures but
This compound is comparatively stable at low temper-
decomposes in water slowly to yield chromium (111) ion and oxygen.
Excess acid favors this. '17Chromium is amphoteric and with excess base chromium
hydroxide is converted to sodium chromite. 18Alkaline peroxide will oxidize
chromite ion to chromate ion.
These reactions suggest a sequence of experiments that might be performed.
The first sequence attempted was as follows:
1)
sodium chromate is converted to sodium dichromate with sulfuric acid.
2)
sodium dichromate is converted to the peroxo complex with acidic
hydrogen peroxide.
3)
chromium peroxo complex is converted to chromium sulfate with heat.
4) chromium sulfate is reacted with sodium hydroxide to produce first
chromium hydroxide, and then with further base sodium chromite.
5)
sodium chromite is converted to sodium chromate with alkaline hydrogen
peroxide.
This sequence of reactions was attempted in (61-36 I).
However, it was
soon realized that potassium dichromate is a much less soluble salt than
either sodium or potassium dichromate. Solubilities are sodium dichromate:
238g/lOOcc at OOC, potassium dichromate:
4.9g/lOOcc at O o , sodium chromate:
50g/lOOcc at 10°C.19Step 1) was then modified to 1) potassium dichromate is
converted to potassium chromate with potassium hydroxide.
The resulting
chromate ion will then be in basic solution, and this must be made acidic for
the peroxide reduction to chromium (111) ion.
If the acid chosen here is
and Humiston, G.E., "General Chemistry", 3rd Edition, John
Wiley and Sons, New York, 1975, p671
17 Jolly, W.L., "Encounters in Experimental Chemistry", 1st Edition, Harcourt,
Brace Jovanovich, Inc., New York, 1972, p134.
18 Moeller, T. , "Qualitative Analysis", 1st Edition, McGraw Hill, New York,
1958, p.221
19 "Handbook of Chemistry and Physics", Weast, R.C. and Selby, S.M., Editors,
47th Edition, Chemical rubber Co., Cleveland, Ohio, 1964 , p.B-203 and B-223.
73
16 Brady, J.E.,
-
-
acetic instead of sulfuric as used in the previous scheme, then the potassium
acetate (solubility 253g/lOOcc at 20°C) will be much less likely to cocrystallize in the final step with the potassium dichromate than would potassium
sulfate (solubility 12g/lOOcc at 25OC).
Therefore acetic acid was substituted
for sulfuric acid, formerly introduced in Step 1).
Finally, a Step 6) is
added in which potassium chromate is converted to dichromate with acetic acid,
and the potassium dichromate is isolated by crystallization.
The modified
reaction scheme then becomes:
1)
potassium dichromate is converted to potassium chromate with
potassium hydroxide.
2)
potassium chromate is converted to the chromium peroxo complex with
acetic acid and hydrogen peroxide.
3)
chromium peroxo complex is converted to chromium acetate with heat.
4)
chromium acetate is reacted with excess sodium hydroxide to eventually
produce potassium chromite.
5)
potassium chromite is converted to potassium chromate with alkaline
hydrogen peroxide.
After several modifications the experiment described below (61-111 I) was
carried out, and represents a successful application of the sequential reaction
descussed.
Potassium dichromate (2.82g) was placed in 5Oml of water in a
50Oml beaker.
Five cc of glacial acetic acid was added and then 50cc of 3%
hydrogen peroxide.
then reddish black.
The solution turned dark blue immediately, then black,
Oxygen evolved.
of oxygen was evolved.
The mixture was heated and a small amount
Fifty mls of 6M KOH were added and the solution turned
an emerald greem (chromium (111) ). Fifty m l s of hydrogen peroxide were added.
The solution, over about a three minute period, turned from emerald green to
light green to gold in color.
The mixture was heated to reduce the solution
to about fifty mls in volume.
It was cooled and 27mls of glacial acetic acid
added at which point the solution turned from gold to orange indicating
conversion of the chromate to dichromate.
added.
Three m l s more of acetic acid were
The volume of the solution was now 70 mls.
This was chilled and filtered.
The dichromate was washed with two 25ml portions of ethanol and dried on low
heat to constant weight.
seating at 118% yield.
The yield of potassium dichromate was 3.338,
repre-
The obvious contaminant was potassium acetate, which
will not interfere with the cyclical use of the potassium dichromate.
This
contaminant can undoubtedly be reduced by reduction in the amount of potassium
hydroxide used.
-
74
-
The mother liquor from the experiment is added to the next year's crude
potassium dichromate.
Ever so often, the mother liquor should be treated
with hydrogen peroxide to reduce the chromium to chromium (111), the hydroxide
precipitated and removed by filtration and added to the chromium acetate of
Experiment11
hydroxide.
, making
sure sufficient acetic acid is present to dissolve the
This procedure will rid the cycle of excess potassium acetate.
Alternately the chromium can be reduced to chromium (111) and precipitated
with aluminum and the potassium acetate removed by filtration.
The complete
experiment occupies about 2 hours and may be separated into 2 experiments,
the first being the series of chemical reactions and the second the fractional
crystallization of the potassium dichromate.
The reaction sequence exhibits chromium in two oxidation states and
colored compounds of yellow, orange, deep blue, green and red.
The sequence
of the two experiments is a complete cycle.
Experiment 13:
The Chemistry of Aluminum, an Amphoteric Wetal
This procedure was developed by a combination of a sequence of reactions
of aluminum and its compounds, partly based on a procedure described by
Summerlin, et a1.20A sequential scheme was developed for aluminum in which
the metal is reacted with potassium hydroxide to produce hydrogen and
potassium aluminate, this converted to hydroxide with sulfuric acid and with
excess to alum, potassium aluminum sulfate, which precipitates in large
crystals.
-_
This procedure illustrates an industrial process.
20 Summerlin, L.R., Borgford, C.L. and Ealy, J.B., "Chemical Demonstrations",
1987, p54.
Vol 2, American Chemical Society, Washington, D.C.,
- 75 -
Experiment 1 4 :
The Chemistry of Lead
Lead is easily brought into solution with nitric acid in a reaction
impressive to students, in which nitrogen dioxide is evolved.
is one of the few soluble lead compounds, lead (11) nitrate.
The product
Lead (IV) oxide
can be synthesized by oxidation with bleach of lead (11) hydroxide.
compound is a chocolate brown color.
This
This compound in acidic solution is a
strong oxidizing agent, capable of oxidizing chloride ion to chlorine.fO
Thus lead (IV) oxide is converted to lead (11) chloride with hydrochloric acid.
Lead chloride in the presence of iodide ion will form the less soluble lead
iodide.21The lead iodide is the link that can complete the cycle.
The lead
iodide is converted to lead nitrate in Experiment 16, and then to lead metal
in Experiment 9
, completing the cycle. The complete cycle is then lead
(silver metal) to lead nitrate (colorless solution) to lead hydroxide (white
solid.) to lead (IV) oxide (chocolate-brown solid) to lead chloride (white
solid) to lead iodide (yellow solid) to lead nitrate (Experiment 1 4 ) to
lead (Experiment 9 ).
The set of reactions would illustrate a number of
compounds of lead, interesting physical and chemical changes, the two
oxidation states of lead, and the oxidizing power of lead dioxide.
To investigate this series of reactions, 0.5g of lead foil was covered
in a beaker with concentrated nitric acid and heated to initiate reaction.
A total of 5 d of nitric acid was added to dissolve all of the lead. Fifty m l s
of water was added and the solution heated until all of the nitrogen dioxide
was expelled, as indicated by the color of the solution. Twenty ml of 6M
potassium hydroxide was added to precipitate lead hydroxide, which appeared
as a creamy white precipitate.
When heated this formed curds and settled.
Ten m l s of bleach were added and the precipitate immediately darkened and
upon further heating turned chocolate brown, indicating formation of lead
dioxide.
acid.
To this was carefully added, in the hood, 3Omls of 6M hydrochloric
Considerable chlorine was evolved and upon heating the mixture slowly
turned from brown to off-white (lead chloride).
was added.
One gram of sodium iodide
At first a yellow precipitate formed but this turned t o white and
brown iodine was apparent, formed by oxidation of iodide by the chlorine present.
20 Brady, J.E. and Humiston, G.E., "General Chemistry", 3rd Edition, John
Wiley and Sons, New York, 1975, p.
21 Moeller, T., "Qualitative Analysis", 1st Edition, McGraw-Hill, New York,
1958, p.277
-
76
-
A small amount (0.5g) of sodium thiosulfate was added which immediately caused
a beautiful yellow precipitate to form, and the supernatent liquid to become
clear.
The lead iodide was removed by filtration (61-109 I).
Because of the evolution of nitrogen (IV) oxide and chlorine from the
reaction mixture, the following modifications were made to the above procedure.
The reaction was conducted in a 2.5cm ID test tube equipped with thistle tube
and gas exit tube.
The gas exit was conducted to a trap consisting of a 250cc
flask equipped with an inlet tube terminating near the bottom, and a gas exit
The trap was initially charged with 10% sodium hydroxide to trap
tube.
oxides.
nitrogen
One half gram of lead was charged to the tube and 5ml of concentrated
hydrochloric acid added to the flask.
oxides and complete the reaction.
Heat was applied to expel the nitrogen
The sodium hydroxide was removed from the
flask and replaced with 1M sodium thiosulfate to absorb chlorine.
Six molar
hydrochloric acid (3Omls) was added and the tube heated until no more chlorine
whs evolved.
Approximately 2mls of 1M sodium thiosulfate was added via the
thistle tube to reduce all dissolved chlorine.
the tube and 1.0 g sodium iodide added.
The stopper was removed from
The mixture was heated.
The yellow
precipitate of lead iodide formed.
This reaction was further modified to essentially eliminate chlorine
evolution as follows:
the above procedure was repeated, reducing the amount
of lead charged to 0.25g.
The sequence of reactions was carried out, halving
the amounts of reagents added.
After the lead (IV) oxide was formed, 0.5g
of sodium thiosulfate was added to the mixture to reduce chlorine as formed.
Hydrochloric acid was then added to yield lead chloride without evolution of
chlorine.
The rest of the reaction was carried as described above.
Experiment 15:
Solubility of Salts and Saturation:
Demonstration
This is an instance in which the solutions may be used indefinitely.
In
this experiment, solutions of sodium thiosulfate and of sodium acetate are
heated unci1 all the solid is dissolved.
They are then cooled to room temp-
erature to supersaturate them, and they are ready for demonstration of
saturation, supersaturation, supercooling, and heat of crystallization.
A
thermometer is placed in the solution and equilibrated.
The solutions are
then seeded with well-formed crystals whereupon a shower of crystal forms and
a riee in temperature is observed.
The solutions are then stored for the
following year.
- 77 -
Experiment 16:
Preparation of Lead Nitrate
This conversion of lead iodide to lead nitrate and iodine was considered
one of the key experiments to the success of this method.
Lead iodide is
produced in Experiments 4 and 14. Lead nitrate is used in Experiments 4
and 9 .
The iodine produced from the process here is converted to ammonium
iodide, which is used in Experiments 4, 5
and 19.
The first method to be explored was the attempted oxidation of lead iodide
with nitric acid.
One gram of lead nitrate was dissolved in 5Omls of water.
To this was added
2.20g of potassium iodide in 5Omls of water.
The mixture
was stirred and lead iodide was allowed to settle, removed by filtration, and
washed with two 25ml portions of water, two 25ml portions of ethanol and dried
with a heat gun (61-42 I).
Lead iodide thus prepared was placed in a 250 ml flask with 25mls of water.
Twenty five mls of concentrated nitric acid was added to the mixture.
there was no visible sign of reaction.
At first
The flask was placed in a boiling water
bath and in about 3 minutes the violet vapor of iodine was observed in the vapor
space.
The lead iodide dissolved with evolution of iodine vapor.
minutes the reaction mixture subsided.
mixture.
In several
Crystals of iodine were suspended in the
In the first attempts, the iodine was removed from the reaction
mixture by extraction with kerosene, mineral oil, cyclohexane, followed by
phase separation.
This had the disadvantage of requiring considerable volume
of the organic extractant and this procedure was modified.
The reaction mixture
was chilled to ice temperature, the iodine removed by filtration, and the filtrate
extracted with the organic solvent.
Extraction in this fashion produced a clear
colorless solution of lead nitrate, a precipitate of solid iodine, and an
violet organic phase containing dissolved iodine.
A quantitative experiment was conducted to determine the yield of lead
nitrate with several concentrations of nitric acid, 15M, 6M, and 3M.
was prepared by dissolving 2.03g lead nitrate in 25ml water.
was removed by filtration, washed, and dried.
Lead iodide
The precipitate
Three portions, 0.7Og each, were
treated with two ml each of the three concentrations of nitric acid.
tative yields ( 0 . 5 4 g ) of iodine were obtained in each case.
Quanti-
It was necessary
to heat the 3M nitric acid a little longer to complete reaction, but all reactions were complete in several minutes.
-
78
-
These reactions were carried out
in a 250 ml flask equipped with thistle tube and gas outlet.
The gas was
conducted through two flasks equipped with inlet tube leading to near the
bottom of each flask, and a gas exit tube. These flasks were each filled
with lOOmls of cold water.
This reaction produces lead nitrate which is evaporated to dryness and
used in Experiments 4 a n d 9, and iodine, which is converted to ammonium iodide
used in Experiments4
Experiment 17:
,
5 and19
.
Preparation of Silver Nitrate from Silver Halides
Because of the high cost of silver, a number of articles have appeared in
the literature concerning reclamation of silver from its various compounds.
any
of these methods use high temperature reactions (500-10OO0C), or hazardous
material (cyanides, silver-ammonia solutions, concentrated base or aqua regia). 25-27
Experiment 5 produces both silver chloride and silver iodide, and a method
was.sought to easily convert these compounds to silver, from which silver nitrate
map be easily obtained.
iodide to iodine.
MoellerZ8reports that copper (11) selectively oxidizes
Copper (I) iodide, a rather unreactive solid is unfortunately
also produced in the process and would present problems of incorporation into
the experimental cycle.
An industrial method of iodine preparation consists of the following scheme:
A silver nitrate solution is added in sufficient quantity to brine to precipitate
only silver iodide, which is removed by filtration and treated with clean steel
scrap to form metallic silver and a solution of iron (11) iodide. ,The iron (XI)
iodide is then treated with chlorine to yield ferric chloride and iodine.29
This approach seemed promising, that is, treatment of silver iodide(and
silver chloride) with iron to yield silver and iron (11) iodide.
The iron (11)
iodide could be used instead of other iodide salts (sodium or potassium iodide
were originally considered) for the reaction in kperiments 4, 5, and 19.
The
first step of the reaction investigated was the conversion of the silver chloride
to the less soluble silver iodide, which was known to react with iron to yield
silver.
22 Wilbanks, B.L., J Chem. Educ. 2,347 (1953)
23 Steed, S . P . and Hayes, J.M., J. Chem. Educ. 49, 156 (1972)
24 Bush, J.B., and Diehl, H. , J. Chem. Educ., 5
c 54 (1979)
25 A m o u r , M/A., J. Chem. Educ., 65, A66 (1988)
J. Chem. Educ., 63, 537 (1986)
26 Rawait, J.P. and Kamoonpuri, S.I.M.,
27 Hill, J.W. and Bellows, L., J. Chem. Educ., 63, 357 (1986)
28 Moeller, t., "Qualitative Analysis", McGraw-Hill, N.Y., 1958, p231
29 Kotz, J.C. and Purcell, K . F . , "Chemistry and Chemical Reactivity", 1st Ed.
Saunders College Publixhing Co., Philadelphia, Pa., 1987 p851
79
-
-
Fifty mls of 0 . 1 M silver nitrate were treated with 70 mls of 0.1M hydrochloric acid to yield a white precipitate of silver chloride.
of 0.1M silver iodide was added to the light grey residue.
heated on the hot plate for 10 minutes and then filtered.
obtained.
One hundred mls
The mixture was
A dark filtrate was
To determine if this color was due to iodine, a portion of the filtrate
was extracted with mineral oil to yield a clear aqueous layer and a violet upper
layer, confirming iodine, probably formed by nitrate ion oxidation (61-9 I).
To the precipitate of silver chloride was
The experiment was repeated.
added 10Omls of 0.1M sodium iodide.
of conversion to silver iodide.
A yellow-gray mass was obtained indicative
The mixture was brought to a boil to complete
conversion of the silver chloride to silver iodide.
removed from the filter and placed in 25mls of water.
The precipitate was washed,
Two ungalvanized iron
nails, cleansed of rust by pickling in hydrochloric acid, were added to the
precipitate and the mixture heated on a hot plate for 10 minutes and then filtered.
The dark filtrate was extracted with mineral oil to yield a clear aqueous layer
and again a violet upper layer, indicating the presence of a small amount of
free iodine.
It was felt that the iodine produced in the above reaction was due to nitrate
ion oxidation.
The procedure was altered to include this step (61-13-A2)
and
the darkening due to iodide oxidation was eliminated.
Silver iodide conversion was determined in several experiments and the
Silver nitrate (0.77g) was dissolved in 5Omls
following is typi,cal (61-10-1).
of distilled water and a solution of 0.99g of sodium iodide in 5Omls of water
was added with stirring. Light yellow silver iodide precipitated immediately.
The silver iodide precipitate was placed
The mixture was heated and filtered.
in 50mls of water.
Three steel screws were cleaned by brief immersion in
hydrochloric acid, rinsed, and added to the silver iodide.
the mixture was
A black deposit of silver formed immediately
agitated with a magnetic stirrer.
around the metal.
The mixture stood for several days. At the end of this
period a grey-black precipitate remained. The steel screws were removed with
forceps and washed with water.
cold.
The precipitate was treated with 15mls 6M HNO3
There was no visible reaction.
nitrogen (IV) oxide was evident.
The mixture was warmed and evolution of
Within about 2 minutes the black precipitate
haii disappeared and in its place a small amount of yellow-white precipitate
remained (unreacted silver iodide).
The mixture was cooled and filtered.
- 80 -
The weight of unreacted silver iodide remaining on the filter was 0.09g
representing a 91.5% conversion of silver iodide.
orating the silver nitrate solution to dryness.
An
This was checked by evapamount of 0.77g of silver
nitrate representing an 80.5% conversion was obtained.
The lower yield was
attributable to some loss due to spattering of the silver nitrate.
The procedure described above was repeated with steel wool, cleansed by
pickling with hydrochloric acid, rinsed, and immersed in 5Omls of water containing silver iodide prepared from 0.75g of silver nitrate.
Silver chloride
recovered amounted to 0.49g, representing a 75.4% conversion of silver iodide
(Gl-74 I).
Faust30reports a method of recovery of silver chloride residues by reaction
with zinc in acid solution, based on procedures previously reported in the
literature. 3l,32Meyer reports that zinc and iron both reduce silver iodide
to silver, but that its reduction is less easy and less complete than in the
case with silver chloride.
The procedure described by Faust f o r silver chloride is as follows:
A
suspension of the compound in water is vigorously stirred with granular zinc
and this is added until only a flocculent gray solid remains in the reaction
vessel.
Once gas evolution ceases, the mixture is filtered and the grey silver
residue is washed with an aqueous solution of 25% sulfuric acid followed by
water.
The reaction is reported to be complete in 5 minutes.
The procedure was carried out, using electrolytic zinc produced from
Experiment 20.
This zinc is in the form of metallic threads of zinc, or
fractals and presents considerable surface area as does granular zinc, recom-
mended in the Faust procedure.
In a typical experiment, 0.47g of silver nitrate
was dissolved in 5Omls of water and added to 1.02g of sodium iodide in 5 h l s
The precipitate silver iodide was heated to coagulate the precipitate,
cooled, and lomls concentrated sulfuric acid. The solution turned an ochre
water.
color, probably caused by nitrate oxidation of a small amount of iodide ion.
30 Faust, D.F., J. Chem. Educ., 61, 924 (1984)
Chem Abst., 19, 1827 (1925)
31 Meyer, A . , Chem. Ztg., 49, 22?-(1925),
32 Mellor, J.W., "A Comprehensive Treatise on Organic and Theoretical Chemistry",
V01.111, Longmans, Green and Co., N.Y., 1923, p.314
See Reference
- 81 -
The mixture was agitated by air sparging.
(2.09g) was added.
Electrolytic zinc from Experiment
Immediately a black scum formed on the surface of the
In approximately 30 seconds, chunks of black substance were evident.
mixture.
No yellow-white precipitate of silver iodide was evident (61-15 I).
The
reaction mixture was decanted and the supernatent filtered. A golden mother
liquor and grey precipitate were obtained.
To the precipitate was added lOml
of 6M nitric acid.
The mixture was warmed on the hot
plate.
No reaction was noted.
Evolution of nitrogen (IV) oxide was observed and the dark grey pre-
cipitate disappeared leaving a small amount of light grey precipitate.
The
mixture was filtered, washed with 25ml of water, then 25 ml of acetone and
dried with a heat gun.
The filtrate was tested with 6M hydrochloric acid
and confirmed the presence of silver nitrate.
The unreacted silver iodide-
silver (light grey precipitate) was dried to constant weight to obtain 0.07g
unregcted residue.
Assuming the residue is silver iodide, this represents an
861 conversion of silver iodide.
This reaction was immediately put to practical application (61-17 11).
Filter paper containing silver residues contaminated with other substances
(zinc and chromium for instance) were ashed to obtain 38.81g of black ash.
This was placed in a 40Oml beaker, 15Oml of water and 3Oml of concentrated
sulfuric acid were added.
The mixture was placed on a steam bath and heated
for one hour to dissolve any soluble metals.
The mixture then stood overnight
and the next day was washed with two lOOml portions of water.
was added to a chromium/zinc waste container.
The filtrate
Four hundred m l s of water were
used to transfer the solids from the filter to a 60Oml beaker.
Nitric acid
(25ml) was added and a vigorous evolution of nitrogen (IV) oxide ensued.
The
mixture was heated for about 30 minutes and filtered to yield a green (chromium)
filtrate.
Ten grams of sodium chloride in lOOcc of water was added to form
large curds of silver chloride.
The mixture was allowed to settle. A small
amount of dilute hydrochloric acid confirmed completeness of precipitation
of the silver.
The precipitate was filtered and washed with several 5Oml
portions of water.
Enough 10% sulfuric acid was added to precipitate all
the silver as the sulfate.
The sulfate was removed by filtration, placed in
a 250 ml beaker, and 100 mls of water and then lOmls of concentrated sulfuric
acid added, followed by log of zinc metal.
- 82 -
The mixture was agitated and a
black residue forned.
The agitation was stopped because of excessive foaming.
After approximately one hour the evolution of hydrogen subsided. An additional
ml of concentrated sulfuric acid was added to insure completeness of reaction.
The grey silver residue was removed by filtration.
Nitric acid (6M) was
added dropwise on the filter to convert the silver metal to the nitrate
salt.
The silver nitrate solution was evaporated to dryness to yield 2.64g
of silver nitrate.
The described method of zinc reduction seems quite satisfactory for the
reclamation of silver from silver residues.
It uses zinc produced from the
experimental sequences, the conversions of silver iodide are good and the proAn experiment was carried out using a mixture of halides
cedure is simple.
to determine the conversion of the mixture.
Three halides, sodium chloride, sodium bromide, and sodium iodide were
weighed out in the amounts of 0.50g each.
Silver nitrate in the amounts of
l.OOg, 0.64g, and 0.44g were dissolved in 5Omls of water and the three halides
added to each beaker respectively.
The halides were combined and 3.10g sodium
iodide was added to convert the yellowish white to the less soluble silver
iodide.
The mixture was digested over low heat for about 2 hours.
period the precipitate was isolated and washed by decantation.
After this
One hundred
fifty mls of water were added followed by 2.03g zinc produced by electrolysis
(Experiment 20).
Ten mls of concentrated sulfuric acid were added and the
mixture stirred vigorously.
After the bulk of the hydrogen was evolved, 3mls
of sulfuric acid were added and the mixture heated until all evolution of
hydrogen ceased. The mixture was then decanted and washed with 5Omls of water.
The filtrate was saved for G-2-11-1.
nitric acid.
The precipitate wae treated with 33%
When evolution of nitrogen (IV) oxide ceased the mixture was
heated for about 10 minutes and then filtered.
The precipitate (unreacted
silver halides) was washed with acetone to remove water, and then air-dried
to constant weight.
The amount of unreacted halide was 0.45g giving a 79.2%
conversion of silver halide (6-2-9-1).
Recovery of silver as silver nitrate was carried out by adding sufficient
nitric acid (12.5M) to dissolve the silver, and simply evaporating the nitric
acid solution to dryness.
When no more fumes were evolved, the solution was
cooled, the crystals removed from the beaker and bottled.
- 83 ..
In summary, silver may be recovered from a mixture of silver halides by
reduction of silver with zinc in the presence of sulfuric acid.
The silver is
dissolved in nitric acid and the silver nitrate obtained by heating the mixture
until fumes of nitric acid are no longer expelled.
be expected.
Conversions of 75Xw may
Unreacted silver halides are added.
Experiment 18: Chemical Separation of Zinc
Solutions of zinc sulfate, zinc iodide and zinc chloride combined with
sulfuric and hydrochloric acids, and ammonium chloride are produced in Experiment 17.
A convenient method for separation of the zinc appeared to be precipi-
tation of an insoluble zinc compound, perhaps the carbonate or hydroxide.
To
explore this possibility, 15Omls of the filtrate from a silver halide-zinc
reduction (Experiment 17) was treated with approximately 7Omls of 1 0 % ~sodium
carbonate until foraming cleared.
Twenty mls additional sodium carbonate
solution was added to precipitate the zinc as zinc carbonate.
ous mixture was heated for 10 minutes
considerabie settling of the solid.
t3
The white gelatin-
coagulate the precipitate.
This caused
The mixture was filtered through Whatmann
54 filter paper, yielding a white precipitate resembling mashed potatoes.
The
zinc carbonate was dissolved in 1.5M sulfuric acid (17mls) to give a colorless
solution of zinc sulfate ready for electrolysis (G-2-11-1).
This procedure represents a satisfactory method for isolation of zinc from
its contaminants and the experiment consumes about 45 minutes.
Experiment 19:
Redox Titration
This experiment accomplishes the very practical purpose of determining
the percentage of ammonium iodide in the solution produced from Experiment 26,
so that an excess of iodide ion may be added to lead nitrate in Experiment
This is critical to the quantiAt the same time, the student is introduced
t o precipitate all of the lead as the iodide.
tative nature of Experiment 4.
to an oxidation-reduction titration.
In this titration, an excess of sodium hypochlorite is added to the iodide
solution and acidified.
Starch indicator immediately reacts with free iodine
to yield the deep blue starch-iodine guest-host compound.
The iodine is back
titrated with thiosulfate and the percentage iodide is determined.
Alternately,
alkaline hydrogen peroxide may be used as the oxidant with the remainder of
the procedure being identical.
The iodide present in the titration mixture is reclaimed in Experiment 25.
- 84 -
Experiment 20: Electrolysis of Zinc Sulfate
The methods developed in Experiment 2 are utilized here to reclaim zinc
from zinc sulfate.
The procedure consists of electrolyzing the zinc sulfate
solution with a 12V source (a battery charger is convenient) using a carbon
welder's rod as the cathode and a platinum electrode as the anode.
Zinc in
the form of delicate fern-like shapes is produced, which presents a large
amount of surface area and works well in Experiment 17, in which silver halides
ara reduced by zinc to silver metal.
Experiment 21:
Ion Exchange, Recovery of Oxalic Acid
The experiment involves reclamation of oxalic acid converted to the
sodium salt in Experiment 10 by titration with sodium hydroxide.
In this
scheme, the sodium oxalate is passed through a cation exchange resin in the
hydrogen form to produce the oxalic acid.
The column is then regenerated
with acid and the sodium expelled as sodium chloride.
To exmaine the effectiveness of this procedure DOWEX 50W-X8, H+ form
was charged to a column by slow addition of 2ml/min. of a slurry of the resin.
Several column volumes (25ml) of 3M HC1 were passed through the column,
followed by sufficient distilled water to remove completely all the free chloride
as tested by silver nitrate.
A 1.06g sample of oxalic acid was placed in a
250 cc flask, 5Omls water added, and the mixture taken to the phenolpthalein
end point with 0.5M sodium hydroxide.
A flask containing a drop of 0.5M
sodium hydroxide and phenolpthalein indicator was placed under the column.
Change in color of the indicator signals the presence of acid.
oxalate was charged to the columns.
The sodium
Thirty-five m l s of effluent remained pink
but after addition of 4Omls the mixture in the receiver was colorless.
A
different receiver was placed under the column and approximately lOOmls of
effluent was collected.
Titration with 0.5M sodium hydroxide showed 1.Og of
oxalic acid to be present in the effluent for a yield of 94.5%.
- 85 -
Experiment 22:
Kinetics:
Clock Reaction
In this classical reaction3? iodate ion is reduced to iodine which is
reduced to iodide ion until the reductant is exhausted.
reacts with starch to form a colored complex.
The iodine then
The concentration of reactants
may be adjusted so that the color change takes place only after a number
of minutes have expired.
The sudden color change is impressive (colorless
to deep blue) and permits the kinetics of the reaction to be explored.
Effect of time and temperature may be determined, and a pseudo rate constant
and activation energy may be calculated from the data.
Destruction of the starch-iodine inclusion adduct is necessary to isolate
the iodine from the reaction mixture at the conclusion of the experiment.
For example, iodine does not extract into an organic phase in which it is
soluble, when bound to the starch.
An experiment was carried out to determine
if hydrolysis of the starch by sodium hydroxide would destroy the complex
(61-99-1).
Five g of sodium hydroxide was added to 15Omls of the reaction
product from the clock reaction,prepared by addition of 5Oml of a solution
containing 4.3g potassium iodate per liter to 50 ml of a solution containing
4g of soluble starch, 0.2g of sodium hydrogen sulfate and 5nl of 1M sulfuric
acid diluted to 1 liter.
faded to colorless.
The solution was allowed to stand, and the blue
Further literature research revealed that the hydroxide
ion reacts with starch-iodine complex to yield iodide and iodate ion, both
colorless. 34
In another experiment, the starch-iodine complex was heated to determine
if this would destroy the complex(G1-89-1).
in (61-99-1).
became
The complex was prepared as
The mixture after boiling for about 30 seconds suddenly
colorless.
Either of these two methods is satisfactory for complex destruction.
The
iodine may then be recovered from the solutions by chilling to ice temperatures and filtering.
This removes the bulk of the iodine.
The brown filtrate
is then extracted with cyclohexane to remove the last traces of iodine
(61-99-111).
Summerlin, L.R. and Ealy, Jr., J.L., "Chemical
Demonstrations", Vol 1, 2nd Edition, American Chemical Society
__ Washington D.C., 1988, p.107
34 Ophandt, C.E., J. Chem. Educ., 64, 9, 808, 1987.
33 See for example:
- 86 -
Experiment 23:
Synthesis of Methyl Benzoate
Preparation of an ester is a standard experiment for an introductory
class in chemistry, since it is easily accomplished and produces a goodsmelling product, sometimes identifiable.
of a 2
-
This preparation becomes part
experiment cycle in which the benzoic acid produced in Experiment
becomes the raw material for this experiment.
Methyl benzoate is prepared
according to standard procedures .35
Modifying the procedure used in Cason and Rapoport, 2.5g of benzoic
acid and 5.Oml methanol were placed in a-large test tube and 8 drops of
concentrated sulfuric acid were added with constant swirling.
was fitted with stopper and 20cm x 5
The test tube
I D glass tube as reflux condenser.
By suspending the tube above a boiling water bath, a gentle reflux was
sustained for 30 minutes.
The mixture was washed with 25mls of 10% sodium
carbonate and phase separated.
The crude methyl benzoate was washed with
a 25ml portion of water and again phase separated.
The sodium carbonate
layer was carefully neutralized with 6M hydrochloric acid to produce 0.70
of unreacted benzoic acid.
obtained ( G l - 1 0 2 - 1 ) .
A 72% yield ( 2 . 0 0 g ) of methyl benzoate was
This crude ester was then used in the hydralysfs
experiment.
Experiment 24:
Hydrolysis of an Ester
The ester prepared in Experiment 23 is saponified here to yield the
original alcohol and acid.
by the following experiment.
The procedure is standard and is exemplified
The amyl acetate ( 2 . 0 g ) was placed in a
125ml Erlenmeyer flask with 2Oml of 2M sodium hydroxide and the flask
fitted with a reflux tube.
The sample was heated for 30 minutes at which
time the sample was homogeneous.
The solution was cooled and acidified
with 6M hydrochloric acid and heated to boiling.
It was then cooled and
filtered and the benzoic acid (1.62g) washed with cold water (6-1-103-1).
The benzoic acid is used in Experiment 23.
-.
See, for example, Cason, J. and Rapoport, H., "Basic Experimental
Organic Chemistry", Prentice-Hall, Inc., Englewood Cliffs, N.J.,
1964 p.90
- 87 -
Experiment 25:
Preparation of Iodine
The filtrates from Experiments 3, 4, 18 and 19 contain iodides. The
primary purpose of this experiment is to recover these water soluble
iodides as water insoluble iodine, and thus afford a means of separation
from the other components of the mixtures, which include ammonium chloride,
sodium chloride, sodium sulfate, and sulfuric acid.
The oxidation of iodide
may be affected with either sodium nitrite and nitric acid, or with acidic
hydrogen peroxide.
To investigate the usefulness of sodium nitrite, 15Omls of the filtrates
obtained from Experiments 3, 4, 18, and 19 were treated with 2 g of sodium
nitrite and 25mls of 6M nitric acid.
and iodine precipitated.
filtered.
The mixture immediately turned brown
The mixture was cooled to ice temperature and
A brown filtrate was obtained which was extracted with cyclohexane
to yield a colorless filtrate and a violet cyclohexane solution (G2-21-1).
Iodine (0.82 g) was obtained as a solid.
The reaction was repeated using 5Omls of 3% sodium hypochlorite and 25mls
of 3M H2SO4.
solution
Iodine (0.79 g) was obtained after chilling and filtering the
.
- a0 -
Experiment 26:
Preparation of Ammonium Iodide from Iodine
This is a key experiment in the cyclical schenes explored in this report.
Iodide ion and iodine are produced in a number of procedures and these species
must be converted to a suitable soluble iodide that may be conveniently used
in the cycle.
The final procedure adopted is the reaction of ammonium
hydroxide and hydrogen peroxide with iodine to yield ammonium iodide.
The
experiments that led to this final adoption will be described.
The first reaction explored was the reaction of calcium hydroxide with
iodine.
The expected reaction products were calcium iodide and a precipitate
of calcium iodate.
The solubility of calcium iodate is 0.20g/lOOcc at 15OC
whereas calcium iodide is soluble to the extent of 209g/lOOcc at 20OC.
The
precipitation of the iodate from the solution was expected to drive the
reaction to the right.
Filtration of the mixture should yield calcium iodate
as a solid and a solution of calcium iodide.
.
At this point in the experimental program, it was intended to extract all
iodine into an organic phase as the means of isolation of the element.
For
this reason, the calcium hydroxide reaction was carried out heterogeneously
with a kerosene solution of iodine (G-1-53-1).
Ten g of calcium oxide was
suspended in water and agitated for approximately one hour to prepare a
saturated calcium hydroxide solution.
the clear limewater decanted.
The solution was allowed to settle and
Fifty mls of a 3% solution of iodine in
kerosene was mixed with 5Omls of limewater.
minutes.
The mixture was agitated for 30
At the end of this period the solution was partially emulsified.
It was placed in a boiling water bath which partially broke the emulsion.
The experiment was modified.
Twenty-five mls of 3% iodine in cyclohexane
was briefly shaken with 25mls of limewater (6-1-53-11).
of the iodine disappeared almost immediately.
upper layer and a clear lower layer.
The violet color
In this case a white turbid
The upper layer was heated in a water
bath to demulsify it, but it remained turbid and this approach was abandoned.
Iodine is reduced by sulfite ion in the presence of acid.36The reaction
products of the direct reaction with sodium sulfite in the presence of
sulfuric acid would be sodium iodide, sodium sulfate and any sulfuric acid
and sodium sulfite in excess. The sodium iodide could be separated by the
following scheme. Excess sodium sulfite would be destroyed with acid to
36
Hoeller, T., "Qualitative Analysis", 1st Edition, McGraw-Hill,
New York, 1958, p.216
- a9 -
yield sodium sulfate and sulfur dioxide.
The excess acid could then be
neutralized with sodium hydroxide to yield sodium sulfate.
The two com-
ponents of the mixture at this stage would be sodium sulfate and sodium
iodide.
Solubilities of these two compounds are greatly different in ethanol.
Sodium iodide is reportedly soluble to the extent of 43g/lOOcc whereas
sodium sulfate is reportedly insoluble .37Extraction of this mixture with
ethanol after evaporation to dryness should then produce an alcoholic solution of sodium iodide.
This entire procedure was explored.
Iodine (0.5g) was placed in a test
tube and l h l s of 1M sodium sulfite added.
Addition of 5mls of 3M sulfuric
acid with agitation brought about solution of the iodine, indicating reduction.
The resulting solution was yellow, probably due to the presence of a small
amount of triodide (6-1-57-1).
To determine the efficacy of separation of sodium iodide from sodium
sulfate, 2.01g of sodium iodide and 1.52g of sodium sulfate were placed in
a 25Oml beaker and extracted with l 0 h l s ethanol by agitation for 45
minutes.
The solution was allowed to settle and the ethanol solution re-
moved by decantation.
dried.
The precipitate was washed with 25mls of ethanol and
The yield of remaining sodium sulfate was 98.0% (1.49g), illustrating
a nearly perfect separation of the iodide from the sulfate.
Because the sulfite reduction and selective extraction with ethanol for
preparation of sodium iodide is rather involved, other procedures for pre38
paration of iodide salts from iodine were investigated. Moeller reports
sulfide ion will reduce iodine to iodide.
Sulfur and sodium hydroxide are
To examine this reaction, 5Omls of 8.1M sodium sulfide
solution was added to 5Omls of 1% iodine in kerosene and the reaction mixture
the by-products.
sparged to agitate the solution.
The kerosene lost color almost immediately.
The lower layer was yellow and the upper kerosene layer colorless.
After
phase separation, the solution was boiled to coagulate the sulfur, but this
was not successful.
of it.
Passage through charcoal helped to remove the bulk
At this point the procedure was abandoned.
37 "Eandbook of Chemistry and Physics", Weast., R.C and Selby, S.M.,
Editors, 7th Edition, Chemical Rubber Co., Cleveland, 1964,
P.B224 and B-226.
38 Moeller, T., Qualitative Analysis", 1st Edition, McGraw-Hill, New York
1958, P.221
- 90 -
Iodine will react directly with zinc dust to produce zinc iodide.39
Iodine (1.Og) was dissolved in 5Oml
This reaction was explored as follows:
kerosene and lOOcc of water in a 25Oml flask.
Electrolytic zinc produced
from Experiment 20 was added and the mixture sparged with air overnight to
agitate the mixture.
The next morning the iodine had vanished and a brownish
kerosene layer and clear lower aqueous layer remained.
complete reaction.
This indicated
This reaction looked quite promising since the zinc iodide
contained no by-products, and therefore no further separation scheme was
necessary (6-1-58-1).
Another promising reaction was the conversion of iodine to ammonium
iodide with ammonium hydroxide and hydrogen peroxide?'
carried out in the following manner.
This procedure was
One gram of iodine was reacted with
28mls of 10% ammonium hydroxide (double the stoichiometric quantity) and
6Omls of 3% hydrogen peroxide.
oxygen was evolved.
Most of the iodine dissolved readily and
Another 2Oml of 3% hydrogen peroxide was added.
iodine completely disappeared and the reaction mixture became yellow.
All
The
solution was evaporated on the steam bath to yield colorless crystals in
near quantitative yield (2.25g, 98,6% yield) (6-1-56-I).
The ammonium hydroxide/hydrogen peroxide reduction of iodine to ammonium
iodide procedure has been adapted because of the ease and rapidity of the
reaction, and the lack of by-products other than water.
This reaction is
used to convert iodine from all sources to a conveniently used iodide salt,
which is used in Experiments 4, 5, 14 and 19.
Experiment 27:
Synthesis of Sodium Nitrate from Nitrogen (IV) Oxide
This experiment is designed to isolate the nitrates and nitrites from
admixtures by converting them to nitrogen (IV) oxides by means of copper
and heat.
They are reabsorbed in sodium hydroxide to produce a solution of
sodium nitrate in sodium hydroxide.
This resulting basic solution of nitrate
ion may then be used in the classical quantitative analysis of nitrate ion
by reduction to ammonia using aluminum as the catalyst.
This destroys the
nitrate and nitrite ion and converts them to the much less toxic ammonium
ion.
The procedure illustrates some nitrogen oxide chemistry and an indus-
trial method for the preparation of nitrates.
39
The evolved gases may also
"Handbook of Preparative Inorganic Chemistry", ed. G. Brauer,
2nd Edition, Academic Press, New York, I, 1963.
- 91 -
be absorbed in water to illustrate the preparation of nitric acid.
The
nitric acid is then reacted with an excess of base to prepare the solution
for Experiment 27, or the prepared nitric acid may be reused in the various
experiments after appropriate concentration.
An experiment was carried out to explore to nitrogen (IV) oxide reaction.
Sodium nitrate (Z.OOg), nitric acid (1.OOg) and sodium hydroxide (1.1Og) were
placed in a 25Oml flask, dissolved in lOOmls of water, and 5.00g of copper
foil added to the mixture.
A stopper equipped with gas outlet tube was
placed in in the flask, and the evolved gas conducted through a tube near the
bottom of the second flask containing lOOmls of 0.5M sodium hydroxide.
second flask contained a vent to the atmosphere.
This
Fifteen nls of concentrated
sulfuric acid was added to the first flask containing the nitrates, and
evolution of nitrogen (IV) oxide began immediately.
The reaction was allowed
to proceed overnight, at which time the reaction had ceased and a blue solution
of copper sulfate was obtained.
No odor of escaping nitrogen (IV) oxide was
ever detected during the procedure.
To confirm the presence of nitrate in the sodium hydroxide absorber,
the solution was made strongly acidic by addition of 3Omls of concentrated
sulfuric acid and 5.0g of copper foil was added. The mixture was heated and
the copper partially dissolved to yield a solution of copper sulfate with
evolution of nitrogen dioxide.
-
This appears to be a satisfactory method for nitrate removal from other
salts.
Experiment 28:
Kjeldahl Determination of Nitrogen
This standard method of nitrogen determination may now be applied to
the product of Experiment 27.
necessary.
The usual digestion procedures are un-
In this procedure, nitrogen compounds are reacted with strong
sodium hydroxide and a catalyst, in this case aluminum, to yield ammonia.
The ammonia is usually absorbed in standardized acid, and back titrated
with standard base to yield the nitrogen content.
be expected here.
No difficulties are to
This procedure serves to convert the toxic nitrates to
the lees toxic ammonia, which ends up as ammonium sulfate, a well known
fertilizer.
- 92 -
Experiment 29:
Fractional Distillation
In this experiment students use a self-constructed distillation column
and condenser to purify cyclohexane and ethanol contaminated with water.
The principles of distillation are discussed.
Clean separations are not a
prerequisite and the major contaminant, water, will not interfere in moderate
proportions with subsequent use of the chemicals.
Experiment 30:
Separation of Salts
In this experiment, the waste acids, bases, and salts are combined,
brought to neutrality with either sodium hydroxide or sulfuric acid as needed
and processed by this and Experiments 32 and 33 to yield chemicals which may
be utilized in several projects (see Experiment 34).
This serves to bleed
the system.
.
The various ions in the combined mixture are sodium, calcium, potassium,
ammonium, iron (III), magnesium, chloride, and sulfate.
precipitates and is removed by filtration.
Calcium sulfate
Addition of sodium hydroxide
precipitates iron and magnesium hydroxides which, after heating the mixture
to coagulate the precipitates, are removed by filtration.
In addition,
heating of the basic solution removes ammonium ion as ammonia.
This is absorbed
in dilute sulfuric acid to produce ammonium sulfate, used in Experiment 33.
The precipitated ferric and magnesium hydroxides are dissolved in acid and
heated with activated aluminum to precipitate iron.
After digestion to remove
unreacted aluminum, the mixture is filtered to isolate the iron, which is
washed with water and dried.
to Experiment 31.
This goes to Experiment 33.
The filtrate goes
The filtrate from the precipitated iron and magnesium
hydroxide is reduced in volume, chilled and the crystallized sodium and potassium sulfates are removed by filtration.
Solubilities of potassium chloride
and sodium chloride are listed as 34.7g/lOOcc water at 20°C and 35.7g/lOOcc
water at O°C respectively, and for potassium sulfate and sodium sulfate
decahydrate, 12g/lOOcc water at 25OC and llg/lOOcc water at ll°C respectively.
If the separation proves difficult, the chlorides may be separated from the
sulfates by use of a cation exchange resin.
This experiment illustrates some principles of chemical separations of
common cations.
No experimental work has yet been done on this set of reactions.
They are all well known and no major problems in the separations are expected.
- 93 -
Experiment 31:
Preparation of Alumina, an Abrasive, Desiccant, and Catalyst
A l l products containing aluminum are conbined, which includes aluminum
sulfate, alum, and sodium aluminate.
Magnesium sulfate is also present, and
is immediately precipitated as magnesium hydroxide.
The mixture is quite
basic, and the pH is adjusted with sulfuric acid to the point at which
aluminum hydroxide also precipitates.
A small portion of the mixture is
filtered and the filtrate tested for completeness of precipitation.
When
enough sulfuric acid has been added to precipitate all of the aluminum, the
mixture is filtered, the precipitate washed and the solid evaporated to
dryness and ignited to alumina and magnesia.
This goes to Experiment 33.
The filtrate contains potassium sulfate which also goes to Experiment 33.
The point of the experiment is to prepare a mixture consisting largely of
alumina which will be employed as a desiccant in Experiment 33.
Experiment 32:
Preparation of Bleach
In this experiment sodium and potassium chloride are electrolyzed with
12V DC current to produce chlorine and sodium and potassium hydroxides. This
procedure simulates the industrial preparation of bleach. The generated
chlorine reacts with the hydroxides to yield potassium and sodium hypochlorite,
the active ingredient in commercial bleach. Since a fair amount of chlorine
is generated, this reaction should be demonstrated by the instructor. The
bleach thus generated may be used to destroy microorganisms, and'its action
observed by aid of a microscope. Objects such as cloth may be bleached. The
apparatus used in this experiment is identical to that used in the electrolysis
of zinc sulfate, Experiment 20.
- 94 -
Experiment 33:
Applications of Industrial Chemicals
This experiment utilizes the non-recycled chemicals produced in the
entire series.
Plaster of Paris is created from calcium sulfate dihydrate
by ignition to dicalcium monohydrate.
The students then use this to form
a mold, and perhaps cast a reproduction in lead.
Ammonium sulfate is used
to fertilize a plant which has been grown from seed using ammonium sulfate
as the sole nitrogen source.
sizing to make paper.
Sodium and potassium sulfates are used as
The mixture of alumina and magnesium oxide is used
as a dessiccant to demonstrate that moist ammonium sulfate becomes dry in its
presence, and then used as a general purpose desiccant f o r future laboratory
work.
Iron powder is used to demonstrate magnetic lines of force, and then
donated to the physics department or used as a source of iron for the plant
fertilized with ammonium sulfate;
Another choice is to oxidize it to iron
(111) oxide and use it to make red paint or rouge.
It should be pointed out
that the quantities of these chemicals that are produced by each class are
rather small.
For example, the amount of calcium sulfate dihydrate produced
in a year is the order of magnitude of 5g/yr.
As a consequence, the chemicals
can be allowed to accumulate for several years, and one or two utilized per
year year thereafter.
Experiment 3 4 :
Hydrolysis of Salts
Solutions of salts are reused each term for pH determination.
See the
Experimental Outline for more details.
Experiment 35:
Determination of Purity by Melting Point
This experiment serves to determine the purity of the oxalic acid produced
in Experiment 21, and thus, whether a fractional crystallization is necessary.
- 95 -
DISCUSSION:
Waste Products
A number of waste products were generated from this work.
These were
isolated into the following categories.
Zinc wasteIIodide Waste
Iodide Waste
Chromium Waste
Iodate/Iodide Waste
Copper Waste
Lead Waste/Iodide Waste
Cyclohexane Waste, Acetone Waste, Ethanol Waste, Kerosene Waste
Silver Waste
Acid Waste
Base Waste
The waste was collected in one gallon plastic jugs.
Processes were
developed for the recycling of these wastes, in general using procedures
developed for the various experiments.
These will be discussed in the order
listed, and an example of the process used will be included.
*
a)
Zinc Waste/Iodide Waste:
Zinc iodide was produced in a number of
cases, hence the iodide was a necessary contaminant.
One gallon of zinc waste
containing zinc sulfate, zinc chloride and zinc iodide was concentrated to about
5OOml.
Three hundred mls of 3% hydrogen peroxide was added to produce a pre-
cipitated iodine.
The mixture was chilled with ice water and filtered into
another one gallon plastic jug through an ordinary coffee filter.
filtrate was obtained.
acid to yield 4.218.
A brown
The iodine was dried in a desiccator over sulfuric
The filtrate was extracted with 2 15Oml portions of
kerosene to yield an essentially colorless filtrate and a violet kerosene solution of iodine.
The kerosene solution (ca 30Omls) was treated with 40mls of
10% ammonium hydroxide and l 2 M s of 3% hydrogen peroxide.
The phases were
agitated with air supplied by an aquarium aerator overnight.
The next morning
a colorless (slightly yellow) cyclohexane layer was removed by phase separation, leaving a colorless aqueous phase containing ammonium iodide, unreacted
hydrogen peroxide and ammonium hydroxide.
dryness to yield 1.30g of ammonium iodide.
This was boiled and evaporated to
The clear zinc-containing aqueous
layer, after extraction with cyclohexane, was electrolyzed with a 12 volt
battery charger, a carbon cathode and platinum anode to yield 12.lg electrolytic
zinc
- a feathery product with considerable surface area.
The cyclohexane
was placed in a separate container to be recovered by distillation.(6-2-11-11).
b)
Iodide Wastes:
Approximately one gallon of iodide waste containing
ammonium iodide, sodium iodide, potassium iodide, iron iodide and free iodine
- 96 -
was distilled and the distillate trapped in a flask surrounded by an ice bath
to trap the vapours.
The volume of the solution was reduced to 50Omls.
To
this mixture was added 25g sodium nitrite and 25mls concentrated nitric acid.
Iodine was immediately precipitated, the mixture chilled to ice temperature
and filtered to remove the iodine.
To the mixture was then added lOOg copper
scrap and 25 mls concentrated sulfuric acid and the mixture heated to decompose
nitrates.
Nitrogen dioxide was evolved and absorbed in cold water to yield
A yield of 12.55g iodine was obtained
a mixture of nitric and nitrous acids.
including distilled iodine and iodine formed after the nitriteinitrate oxidation (6-1-34-1).
c)
Chromium Waste:
Three liters of this waste was concentrated by heat
to a volume of about 1 liter.
This was made strongly basic by the addition
of 30g of potassium hydroxide.
A flacculent precipitate was obtained which
did not dissolve in base, presumably iron hydroxide.
settle by decantation over several days.
This was allowed to
The bulk of the supernatent liquid
was removed by decantation, and the precipitate filtered after the addition
of 10 grams of filter aid.
The solution obtained was treated with 30Omls
of hydrogen peroxide in increments of 5Omls.
The mixture was carefully heated
to destroy excess peroxide and then made acidic with acetic acid as indicated
by the change in color of solution from yellow to orange.
acetic acid was added.
Fifteen m l s excess
Some potassium dichromate immediately precipitated.
The mixture was reduced to about lOOmls of volume, chilled to ice temperature,
and filtered to remove crystals of potassium dichromate.
The crystals were
washed with two 25ml portions of ethanol and oven dried.
A yield of 23.7g
of potassium dichromate was obtained.
The mother liquor was saved and added
to the waste dichromate container (6.2-36-1).
d)
Iodate/Iodide Waste:
All solutions containing iodide/iodate wastes
were treated with an excess of sodium thiosulfate to reduce all iodate to
iodide.
This solution was then combined with b) Iodide Waste and treated as
described earlier.
e)
Copper Waste:
Copper was removed from these solutions by single
replacement with aluminum activated with hydrochloric acid.
For example,
to approximately 1 liter of copper waste was added 2Og of aluminum foil which
had
-_ been activated by previous immersion in concentrated hydrochloric acid.
- 97 -
The mixture was heated until all colot was gone.
The foil was washed with
water to remove precipitated copper and removed from the solution.
Twenty-
five mls of concentrated sulfuric acid was added to dissolve remaining aluminum,
and the solution heated f o r approximately 1 hour.
The mixture was then filtered
to remove the copper, the copper washed on the filter with acetone to remove
the water, and air dried to yield 24.6g red copper powder.
the filtrate is
at present being reduced to dryness by solar evaporation.
f)
Lead WasteIIodide Waste:
This mixture contained solid lead iodide.
Four liters of this mixture was treated with lOOmls of concentrated nitric acid
and heated on the steam bath.
The mixture was chilled to ice temperatures
and the iodine removed by filtration, washed with water and dried over sulfuric
acid to yield 11.6g iodine.
to precipitate lead.
The filtrate was treated with 15g aluminum foil
Unreacted aluminum foil was washed to remove lead and
removed from the solution.
One hundred m l s of concentrated sulfuric acid
was6added to the mixture to dissolve unreacted aluminum.
heated for approximately one hour and filtered.
The mixture was
The silvery lead powder was
washed with acetone and air dried to yield 13.3g of lead.
The filtrate is
being reduced to dryness by evaporation.
g) Organic Wastes: All organic wastes were distilled separately to yield,
respectively, purified cyclohexane, acetone, ethanol and kerosene. The distil-
lations were carried out in a 5OOml round bottom boiling flask equipped with
a 50cm Vigreaux column and condenser.
h)
Silver Waste:
The processing of this waste has been reported in
the Experimental Discussion of Experiment 17.
i), j)
Acid Waste/Base Waste:
These were used to neutralize each other
to pH 7 and are in the process of being evaporated to dryness.
- 98 -
DISCUSSION:
Cost Reduction
Reduction in costs of operating an introductory chemistry laboratory
using the methods outlined in the NOWASTE manual were calculated by the
following procedure.
The steps in the calculation were as follows:
The manual selected for comparison was "Laboratory Manual for
Chemistry'' by Maxine Wagner, Allyn and Bacon, INc., Newton, Mass,
1983.
1.
This manual was selected for three reasons:
The author considers it to be an excellent manual, covering
all major topics in the course.
2.
The manual is used by our local high school
3.
The manual contains a list of all chemicals and their amounts
used.
This is included as Appendix 111.
The costs of the chemicals used for experiments which produce waste
chemicals was calculated.
Cost for the other items listed represent
one-time only costs, since these items are used repeatedly (for
example, copper, lead or aluminum pellets for density determinations).
The total costs for one class of thirty students (15 student-pair
teams) to purchase chemicals which are used each term came to $333
(see Appendix 11).
The amounts of chemicals used in the NOWASTE manual were estimated
and the costs of all chemicals calculated assuming a 10% and 20%
loss in those chemicals which are recycled.
These costs totalled
$3.00 and $3.45 per student team (see Appendix 111).
These costs were multiplied by 15 student teams to give $45.00 and
$52.00 per class for the NOWASTE manual.
Since the Wagner manual contains 39 experiments which produce chemical
waste and the NOWASTE manual contains 34 experiments which produce
chemical waste, the costs of the NOWASTE manual were prorated by
39 = 1.147 to yield a cost comparable to the Wagner manual. These
3F
costs came to $S2/class for a 10% loss and $60/class for a 20% loss
in recycled Chemicals.
Savings in costs were then calculated as follows:
10% loss NOWASTE manual ($333 $52) x 100 84% reduction
- 1
($333 - $60) x
( $333
1
(
20% loss NOWASTE manual
$333
- 99 -
-
in cost.
100 = 82% reduction
in cost.
G)
If costs of disposal by sanitary landfill are included in the estimates,
savings are even greater.
The costs of packing, transporting, and
disposing of "lab p a c k ~ ' * ~ ~ai n"secure" EPA-regulated landfill has
been estimated to be $23/gal in a 1985 publication.
waste from school labs end up as solutions.
Normally, chemical
To estimate the costs
of chemicals disposed as wastes produced from the Wagner manual, the
following procedure was used:
1)
The total waste in grams was calculated = 7300 g
2)
This was prorated from 39 to 34 experiments for reasons described
7300 x 34 = 6350 g
39
3) Aqueous solutions of these chemicals of 25Xw were assumed.
6350
= 25,400 g solution.
0.25
4) A density of 1 . 5 ~was assumed. Volume = 25,400 x 15 = 38,lOOml
ml
or 3.8 L. or 38
qts = 40 qts. or 10 gallons.
0.946
5) 10 gallons x 23 = 230
gal
Thus, disposal costs are close to 2/3 the costs of chemicals.
above.
6)
If disposal costs are included in teh costs of a conventional
chemical lab then savings by use of the NOWASTE manual are
increased to:
-
10% less NOWASTE manual (333+230)-(52)
333+2 30
x 100
20% less NOWASTE manual (333+230)-(60)
333230
x 100 = 89%
90%
In s u ~ ~ p p ~savings
ry,
in costs by use of the NOWASTE manual are estimated
to be between 80%-90% of the chemical costs for the term.
excludes costs of new equipment due to breakage, etc.,
This estimate
encountered in all labs.
There may be some minor capital costs incurred upon adoption of the NOWASTE
A battery charger makes a good source of 12V DC current for the
electrolysis experiment (Experiment 20). For example, a Viking Model VA 7612
manual.
Solid State Battery charger is available at the local auto supply house for
about $40.
Twelve and 6V automobile batteries may also be used.
A pair of
platinum electrodes is also needed for Experiment 20 and are available, for
40 In a "lab pack", 1 gallon bottles of solutions are placed in 55 gallon
drums with vermiculite filler.
41 Walton, W.A., J.Chem. Educ., 64, A669, (1987); "Less is Better"; ACS
Task Force on RCRA:
Washington D.C.,
- 100 -
1985
example, from Central Scientific Co., catalog number 81275 for $17.50.
A
set of 48 500 ml capacity plastic bottles will provide enough storage for
all the solids and solutions generated in the NOWASTE procedure (see DiscussionStorage Space).
These bottles are available, f o r example, from Van Waters
and Rogers, catalog number 16126-278, at 6 for $16.45 or $131.60 for 48.
Total costs for these purchases is approximately $200. Payout for this expense
is approximately one semester.
No special training for instructors is anticipated.
are standard.
All techniques used
Familiarization with the manual is necessary, as with any new
laboratory manual.
Also, no special approval by any state agency is required
for adoption of a new laboratory manual in high schools or community colleges
in the State of C a l i f ~ r n i a . ~ ~
42
-.
Telephone conversation with James Eurst, Board of Directors, Fort
Bragg Unified School District, 12/16/88 and with L. Kavanaugh,
Dean of College, College of the Redwoods; Mendocino Branch, Fort
Bragg, CA 12/17/88.
-
101
-
DISCUSSION:
Impurity Buildup
The possibility that impurities will build up in some of the cyclical
It is extremely difficult to answer this question
without actual experimental evidence, and this is one of the major problems
that will be addressed in future work on this project. However, in general,
some educated guesses may be put forth.
The first point to be made is that _ _ the experiments in this manual
which are cyclical in nature involve a sequence of reactions. These reactions
systems must be examined.
are characteristic for that particular element, and thus tend to be purification
schemes.
For example, let us follow lead through the cycle of reactions
starting with Experiment 14. Outline of Experiments will help the reader
follow this discussion.
Lead might be accompanied by impurities of copper,
antimony, arsenic, bismuth, gold and silver, since these elements are present
as it is obtained from the blast furnace.
The physical separation that occurs
in this experiment is a filtration after addition of ammonium iodide in acidic
medium (HC1).
Assuming all elements mentioned above are present, this reaction
would probably produce soluble copper since copper (11) iodide does not exist46
and copper (I) iodide is soluble in a~id,~~soluble
antimony and bismuth since
the iodides of these elements decompose in hot water,48and soluble arsenic,
probably present at this stage as arsenic acid.
Species accompanying the lead
iodide as precipitates might be gold metal and silver iodide.
The filtered
lead iodide is next reacted with nitric acid in Experiment 16 to yield soluble
lead nitrate and iodine as a precipitate.
Here the gold would join the iodine
as an insoluble residue (if present in sufficient quantity) as would the silver
iodide.
The lead nitrate has thus been freed of impurities arising from the
lead. (The iodine which contains the gold and silver iodide is reacted in
Experiment 26 with hydrogen peroxide and 'ammonium hydroxide to yield water
soluble ammonium iodide.)
The gold and silver iodide are insoluble in water
and would appear as a precipitate at this point, easily removable by filtration.
46 Holtzclaw, Jr., H . F . ,
Robinson, W.R. and Nebergall, W.H., "General
Chemistry", 7th Edition, D.C. Heath andCo., Lexington, Mass. (1984)
47 Ibid., p. 8 6 3 , p. 896
48 "Handbook of Chemistry and Physics", Weast, R.C. and Selby, S. M.,
Editors, 47th Zdition, Chemical Rubber C o . , Cleveland Ohio ( 1 9 6 4 )
p. B-174
-
102
-
Consider chromium, present in Experiments 1 1 and 12.
A common impurity
is iron, originally present in the chromite ore in which chromium is usually
found in nature.
In Experiment 11, the chromium ion is treated with strong
base to convert it to the hydroxide and then to the soluble chromite salt.
Under these conditions, iron would appear as the insoluble iron hydroxide,
which would not dissolve in excess base.
This could be easily removed by
filtration.
Because of many possibilities of impurities introduced in the various
reagents, and the variety of reactions which might occur under varying conditions,
these sort of arguments must be viewed as speculative, and best settled by
experimentation.
Nevertheless, the general conclusion may be drawn that the
series of reactions specific to the elements will tend to purify them.
A second point to be made is that most chemicals used are of reagent grade
purity, and as a consequence, impurity buildup to the point of interference
should be quite slow.
For example, the analysis of "Bakers Analyzed" copper
to have maximum analyzed impurities of 0.01%.
In order for these impurities
to build up to a level of detectable interference, assuming they all interfere
and a level of say lX, would require 100 cycles, assuming replacement of all
the copper each cycle and no loss of impurities.
This represents 100 terms of
chemistry classes.
A third point is that if it is determined that there is a buildup of some
by-product or impurity over a period of years, then there is a good chance
that a simple procedure can be formulated which can remove this impurity on
occasion.
In conclusion, only a series of repetitive runs on each cyclical experiment will determine f o r a certainty if this will be a problem in a particular
case, and it is very likely that some preventative measure can be taken, once
the problem is defined.
-
103
-
DISCUSSION:
Storage Space
All of the chemicals and solutions produced in the suggested procedures
may be stored in 48 wide-mouthed bottles of 500 ml capacity.
These have a
2.5 inch diameter and are conveniently stored in a 15" x 20" space (2 ft.2).
These may be conveniently labelled on the top of the large cap so that they
may be identified easily.
It might be pointed out here that a great deal of
the preparation normally required for a typical course is eliminated by this
type of procedure in which the product of one experiment is the raw material
for the next.
Suggested names for the solutions are given in Appendix IV.
During each exper'bent,
the appropriate product bottles should be made
available to the students by the instructor.
The instructor should check
each solution before it is placed in the product container.
- 104 -
DISCUSSION:
Innovative Nature of NOWASTE Manual
The innovative nature of the NOWASTE manual is essentially an innovation
in method.
The essence of this method is the arrangement of chemical reactions
into a sequence,
so that the product of one reaction produces the starttng
material for the next reaction.
By connecting reactions and entire experiments
in this fashion, the number of chemicals used is reduced greatly.
When entire
experiments can be connected in cycles, then final products can be eliminated.
This is the general approach of this manual.
The result of this approach to the design of academic chemical experiments
is the elimination of chemical waste, a substantial reduction of costs and, an
improvement in meaning and meaningfulness of the experiments.
The intent of
this work is the creation of a method of academic experiment design, which can
To
be applied to numerous chemistry courses involving laboratory experiments.
the.author’s knowledge, no such laboratory manual has ever been published.
A number of innovations have been made with particular experiments.
It
should be realized that the purpose of an introductory chemistry manual is not
to introduce new and unusual chemical reactions, but to demonstrate well-known
and typical properties of the elements.
The innovations of the individual
experiments therefore lie in the details of the experiments, for example, the
choice of a particularly colorful reaction to demonstrate chemical property.
In this respect, the author considers the following experiments innovative
individually for the reasons to be cited.
Experiment 1 uses a novel catalyst, an iron nail, to generate oxygen
from hydrogen peroxide.
This has the advantage of demonstrating to the students
that the nail catalyzes the reaction, as oxygen can be observed forming on the
surface of the nail.
Further, this experiment utilizes 3% hydrogen peroxide
avoiding the hazards of the usual 9% hydrogen peroxide or potassium chlorate.
Experiment 2 is novel in that it utilizes aluminum, an inexpensive metal,
Instead of the conventionally employed zinc to produce hydrogen.
Normally,
aluminum is unreactive to sulfuric acid but,by a pretreatment in hydrochloric
acid to remove the oxide coating, it is rendered reactive.
Experiment 3 demonstrates conventional reactions of the alkaline earths
in addition to an innovative direct reaction of alkaline earths with halogens.
The consecutive nature of the reactions is novel.
- 105 -
Experiment 4 is novel in that it uses ammonium iodide instead of the
conventional potassium iodide.
Experiment 5 is novel in its use of conventional reactions of halogens
in a consecutive fashion.
Experiment 7 employs a reaction novel to the beginning course, oxidation
of copper to its oxide.
This reaction is possible because of the finely
divided nature of the copper produced in Experiment 8, providing the necessary
large amount of surface area, and because a simple apparatus has been developed
to carry out the reaction.
Experiment 8 is novel in that reactions of the well-known copper cycle
reaction have been arranged to illustrate three basic chemical reactions,
decomposition, single and double replacement.
Experiment 9 is novel in that the activity series of the metals is demonstrated in a clear fashion due to the consecutive nature of the reaction.
Experiment 11 is novel in its consecutive nature of the reactions.
Experiment 14 illustrates some chemistry of lead effectively by utilizing
consecutive reactions to produce a series of colorful chemical changes.
Experiment 16 demonstrates a reaction novel to beginning chemistry courses,
yet of a fundamental nature, and indispensable to the recovery of the lead
and iodine.
Experiment 20 demonstrates a technique of recovery of a rather active metal
from solution.
This experiment is novel to beginning chemistry courses and
indispensable to the recovery of zinc.
The nature of the zinc crystals thus
formed permit their use in Experiment 17, the recovery of silver.
Experiment 22 is novel in the destruction of the iodine-starch adduct
and the consequent recovery of iodine.
Experiment 25 is novel in that the nitrite oxidation is employed for
the actual isolation of iodine from other components of the iodide mixtures.
Experiment 26 is novel to introductory courses and enables waste iodides
to eventually be converted to useable ammonium iodide.
Experiment 27 is novel in the sense that it is used to separate all the
nitrogen in the reaction mixture form the other compounds present.
Experiment 30 is novel and employs various techniques to bring about
--separationof salts.
-
106
-
APPENDIX I
OUTLINE OF INTRODUCTORY AND PHYSICAL CHEMISTRY EXPERIMENTS
Experiment A:
Laboratory Equipment Construction I
Purpose :
Glass working and cork boring.
Description:
In this experiment students familiarize themselves with
the bunsen burner, and are taught to bend glass and
construct some simple chemical apparatus.
Schematic :
None
Comments :
Some major construction projects are referred to here.
These are intended to evoke student interest, and can be
used as special projects for single students, groups, or
the entire class, or can be built by the teacher. Some
will have been built by previous classes.
Experiment €3:
Laboratory Equipment Construction I1
More of the sane.
Experiment C:
Alcohol Burner, Scientific Observations
Purpose :
Use of lab equipment and scientific observation.
Description:
Observations on combustion are made. Water production
is proven with cobalt chloride paper. Carbon dioxide
production is demonstrated by its reaction with lime
water. Use of the balance is introduced. Ethyl alcohol
consumed by the burner is related to the temperature
rise in a beaker of water. Conservation of matter and
energy are introduced.
Experiment D:
Density Determination
Purpose :
Determination of a Physical Property
Description:
The balance and graduate are employed to determine the
densities of several substances including several liquids
and solids. An unknown may be introduced. Values obtained
are compared with reported values in the literature.
Comments :
Might use different metals painted the same color and
several colorless liquids to be identified by density.
-
107
-
APPENDIX I (cont.)
Experiment E:
Acids, Bases and Salts.
Organic and Inorganic Compounds
To demonstrate that some common substances can be easily
Purpose :
classified as organic and inorganic. The compounds are
also classified as acids, bases, and salts.
Description:
Tests using litmus paper; tests are performed on table
salt, vinegar, aspirin, ammonia, bleach, wood ashes,
soap, etc. A simple ignition test is performed to
determine if substances are organic or inorganic.
Comments :
Gives students a feeling for acids, bases and salts,
organic and inorganic substances early in the course.
Experiment F:
Percent Oxygen in Air
Purpose :
To determine X v oxygen in air
Description:
A test tube with metric rule attached by a rubber band
contains steel wool cleaned by short immersion in hydrochloric acid. this tube is inverted over water, and from
initial and final volumes of gas the X v oxygen is
determined.
02
+
+
WASTE
Schematic :
Fe
Comment :
The experiment is simple, inexpensive and rapid.
Experiment G:
Fe304
SALT
Flame Tests
To illustrate a rapid method of identification of certain
Purpose :
elements and principles of fireworks.
Description:
Using a platinum wire or cotton swabs (*tips) the flame
colors of salts of calcium, strontium, lithium, copper
and potassium are observed. Fireworks are discussed.
Comments :
Solutions of the salts are retained and reused.
Experiment H:
Characterization of Alum:
After Dehydration.
Melting Point and Mass Remaining
Purpose :
To illustrate a melting point determination, and
determination of water or hydration.
Description:
Alum is one of a few inorganic compounds which have a low
melting point for characterization. A simple melting point
apparatus is assembled and the melting point is determined.
The alum is ignited to remove water of hydration and
weighed in a closed container to prevent rehydration.
- 108 -
APPENDIX I (cont.)
Experiment I:
Acid-Base Indicators
Purpose :
Indicate color change with Ph
Description:
ids
Using a spot plate, differnet con entrations of
and bases are tested with litnus, phenopthalein,
and natural indicators of grape juice, cabbage juice
and wisteria flower extract (or other colored flower).
The ph ranges at which the color changes occur are
established with universal indicator.
Experiment J:
Molar Volume of a Gas
Purpose :
As stated
Description:
Carbon dioxide is generated from dry ice, dried by passage
through drierite and conducted to a tared plastic bag.
When the bag is inflated it is weighed, and the molar
volume at STP calculated from p,T,V and mass.
Experiment K:
Molecular Weight Determination from Molar Volume
Purpose :
As stated
Description:
Unknowns of isopropanol or ethanol are presented to the
students as an unknown and are placed in a flask covered
with aluminum foil with a pinhole. The tared flask is
placed in boiling water and equilibrated at 100" C. The
flask is allowed to cool and is weighed. The molecular
mass is determined from p,V,T and mass of condensed vapor.
Experiment L:
Atomic Weight Determination from Specific Heat
Purpose :
Determination of specific heat and atomic weight estimation.
Description:
The specific heat of a metal is determined using double
Styrofoam cups, a plastic lid, thermometer, water and the
metal. The atomic weight of the metal is then calculated
by the Law of Dulong and Petit.
Experiment M:
Conservation of Mass in Chemical Reactions
Purpose:
Illustrates conservation of Mass in two consecutive chemical
reactions, one involving gas formation and the other
precipita tion.
Description:
Dry ice produces carbon dioxide which is trapped in a plastic
bag. The mass of the apparatus is determined before and
after addition of the sulfuric acid. Aqueous calcium chloride
is then added to precipitate calcium sulfate. All of the
chemicals and apparatus are on the balance so that it may
be established that no change in mass is observable.
-
109
-
APPENDIX I (cont.)
Experiment N:
Electrochemical Cells I
Purpose :
Self-manufactured batteries
Description:
A battery is made from a lemon using zinc and copper
strips. The voltage is tested with a volt-odometer
(if available). The student batteries are combined and
used to set off a flash cube (a capacitor may be necessary).
Using dialysis tubing, magnesium, copper and sodium sulfate,
cells are constructed and used to run a radio and electrolyze zinc iodide (Experiment
).
A few drops of Na2S203 restores the solution to colorless.
A n impressive and fun experiment.
Comments :
The cells are disassembled
and reused next time.
Chemical Demonstrations, Vol 2, ACS,
Washington, D.C., 1987, p.107.
Reference:
Experiment 0:
Enthalpy
Purpose :
Determination of a Heat of Combustion
Description:
A calorimeter is assembled from an Erlenmeyer flask,
thermometer and a coffee can to shield the flask from
drafts. The calorimeter is calibrated by burning a
weighed amount of methanol. The heating value of different
types of wood may be determined. The heat of combustion
of ethanol may be determined.
Experiment P:
Entropy (Demonstration)
Purpose :
To demonstrate the 2nd Law of Thermodynamics.
Description:
Stretched rubbers possess negative thermal expansion
contraction is observed on heating,
coefficients
expansion on cooling. In the apparatus illustrated, the
length of the rubber band is measured at two different
temperatures. Rubber bands used for model airplane
engines work well. The effect of having a rather heavy
load slowly being raised nearly to its breaking point is
a convincing demonstration of a change in entropy.
Reference:
Bader, M.,
-
J.Chem.
Educ.,
- 110 -
z, 285,
(1981)
APPmIX 11
CHEMICALS AND .COSTS OF iTYPICAL CHEHISTRY MANUAL, "LABORATORY MANUAL FOR CHEMISTSY~~,
by MAXINE WAGNER, 1 s t EDITION, ALLYN AND BACON, INC., NEWTON, MASS., 19R3
Ibt includa approximate qUraUUU for O m C b S of
900 mL
1% m L
MO
c acid. glacial
M
CIcnonc
r-
rrcu
. .
h
l
b
w
C
U
Ammonia water (NHI.H.0)
Ammonium acetate*
Ammonium sulfate-.A m y l alcohol
M u m chloride dihydrate
Buium hydroxide octohydnte
M u m nitrate
Barium Sulfate e b
Calcium turnings
Cdt5um carbonate chips
W u m chloride JCISdum hydroxi&
CIildum nitrate t e u a h y d n t c - ~
a k i u m sulfated P
t u b o n disulfide
-s
-
IS grams
15 grams
15 grains
30 grams
20 grams
1s grams
210
I
8
APPENDIX 111
Costs of Chemicals in NOWASTE Manual
Chemical
*
Mass/Vol VWR Catalog No.$/g or ml
cost
loss
1 0% T O %
Hydrogen Peroxide 3%
345 mls
-
0.0020/ml
0.69
Aluminum Foil
13.0s
-
o.o022/g
0.03
Calcium turnings
0.4g
321262-4
0.2192/g
0.09
Hydrochloric Acid
9.8g
JT9535-4
0.0065 /ml
0.06
Sulfuric Acid
20.2g
JT968 1-4
O.O046/g
0.09
Magnesium
0.4g
EM MxOO10-1
0.3445/g
0.14
Sodium Carbonate,
5.5g
JT3604- 1
O.O486/g
0.25
Iodine
2.6g
EM XXO125-1
0.2290/g
(0.60) 0.06 0.12
Lead Nitrate
JT 2322-1
0.1301/g
(0.43) 0.04 0.08
Potassium Iodide
3.3g
0.4g
JT 3165-11
O.O207/g
Silver ‘Nitrate
0.2g
JT 3426-6
1.541 7/g
Ammonium Chloride
0.03
JT 0660-11
O.O435/g
Bleach
18mls
Sodium Thiosulfate
1.5g
Sodium Nitrite
1.2g
Cyclohexane
1 3 . 0 ~ ~ EX CX2290-7
Nitric Acid
5.3cc
Copper Wire
.
-
(0.29) 0.03 0.06
O.O007/ml
0.01
JT 3954-1
O.O442/g
0.06
EM SXO665-1
0.0573/g
0.07
0.0479 /ml
(0.62) 0.06 0.12
EM NXO409-14
o.o112/g
0.06
4.2g
JT 1736-1
O.O930/g
(0.39) 0.08 0.16
Sodium Hydroxide
13.9g
VN 6720-5
O.O214/g
0.30
Lead Foil
0.6g
36996-807
o.o220/g
(0.01)
oxalic Acid
0.5g
JT 0230-1
O.O867/g
(0.04)
Potassium Dichromate
2.8s
JT 3090-1
O.O534/g
(0.15) 0.02 0.04
Acetic Acid
13.4g
0.0145/g
0.19
Potassium Hydroxide
8-48
EM AXOO73-14
EM PX1480-13
0.0527/g
0.44
Starch
0.02g
JT 4006-4
0.1304/g
Cation Exchange Dowex 50W-X8
5.0g
1927-1
O.O795/g
(0.40) 0.04 0.08
Potassium Iodate
0.16g
PX1500-2
0.2506/g
Sodium Bisulfite
0.88
SXO345-1
O.O329/g
0.04
0.03
Benzoic Acid
2.58
HT
EM
EM
EM
BX0360-1
O.O830/g
(0.21) 0.02 0.04
Methanol
5 . m
JT 9070-11
O.OlSS/g
(0.08) 0.10 0.20
*
ONE CLASS OF THIRTY STUDENTS (15 STUDENT
- 113 -
- PAIR TEAMS)
$3.00$3.45
APPENDIX IV
Suggested Names for Solids and Solutions Used in NOWASTE Manual
(numbers at end refer to experiments)
1.
RECOVERED ALUM1"M
2.
RECOVERED ACIDS TO 30
3.
RECOVERED BASES TO 30
4.
RECOVERED HALIDES TO 25
5.
RECOVERED LEAD IODIDE TO 26
6.
RECOVERED SILVER TO 17
7.
RECOVERED CYCLOHEXANE TO 29
8.
RECOVERED NITRATES TO 27
9.
RECOVERED IODINE IN CYCLOHEXANE TO 26
10.
EXPERIMENT 6:
SALTS TO 31
WASTE PRODUCTS
11.. RECOVERED COPPER (11) CHLORIDE TO 8
12.
RECLAIMED COPPER TO 7
13.
RECLAIMED COPPER WIRE TO 9
14.
RECLAIMED LEAD SHOT TO 9
15.
RECLAIMED LEAD TO 14
16.
RECOVERED SODIUM OXALATE TO 21
17.
RECOVERED POTASSIUM DICHROMATE TO 12
18.
RECOVERED POTASSIUM DICHROMATE TO 11
19.
RECOVERED POTASSIUM DICHROMATE MOTHER LIQUOR TO 11
20.
-
aECOpEBED -LEADIIODIDE-TO-16
21.
22.
SODIUM ACETATE AQUEOUS TO 15
23.
RECLAIMED LEAD NITRATE TO 4
24.
RECLAIMED SILVER NITRATE TO 5
25.
RECOVERED ZINC TO 18
26.
RECOVERED ZINC CARBONATE TO 20
27.
RECLAIMED ZINC TO 17
28.
RECLAIMED OXALIC ACID TO 10
29.
RECLAIMED BENZOIC ACID TO 23
SODIUM THIOSULFATE AQUEOUS TO 15
30 t. RECOVERED IODINE TO 26
- 114 -
APPENDIX IV (cont.)
31.
RECLAIMED AMMONIUM IODIDE TO 4
32.
RECLAIMED CYCLOHEXANE TO 5
33.
RECLAIMED METHANOL TO 23
34.
RECLAIMED ETHANOL TO 12
35.
RECLAIMED AMMONIUM SULFATE TO 33
36.
RECLAIMED SULFATES TO 31
37.
RECLAIMED CHLORIDES TO 32
38.
RECLAIMED ALUMINA TO 33
39.
RECLAIMED BLEACH TO 5
40.
RECLAIMED OXALIC ACID TO 10
- 115 -