Enno Wolthuis, Arthur B. ~ruiksma,' and Robert P. ~eererna'
Calvin College Grand Rapids, Michigan
Determination of Solubility A laboratory experiment
During the past several years we have conducted a rcsearch program on the freshman level designed to improve instruction in general chemistry. Specifically, it has been our intention to develop a series of quantitative and/or unknown-type experiments. Such experiments emphasize the quantitative aspects of chemistry and put a premium on good laboratory technique. They interest even the apparently disinterested student and a t the same time challenge the best of them. Our procedure has been to examine critically all published experiments and other literature on a subject, and then, if possible, to improve upon the-best available experiment or devise a new one to accomplish a similar purpose. Whenever possible, it has been our objective to develop a laboratory exercise which is adaptable to an unknowu substance, and to challenge the student to identify it through his work in the laboratory. Two such experiments, the products of our freshman research program, have been d e ~ c r i b e d . ~ . ~ There is a scarcity of good, elementary experiments on the properties of liquids and solutions. One found in a few laboratory manuals deals with the determination of the solubility curve of a salt whose identity is known to the student. I t was our opinion that such an experiment has merit because it effectively teaches the meauing of a saturated solution and it affords the
' This paper reports the work done by Pruiksma and Heereme.
as part of their freshman year's work at Calvin College.
2 WOLTHUIS, E., DE VEIES,D., AND POUTSMA,M., J. CHEM. Eouc., 34,133-4 (1957). a WOLTHUIS, E.,VISSER,M., AND OPPENHUIZEN, I., J. CHEM. Eouc., 35, 412-14 (1958).
student an opportunity to learn the value of a graphical presentation of laboratory data. Eddy has recently described such an experiment on the determination of the solubility curve for borax by tit ratio^.^ We proposed to study the methods for determining the solubility curve of a solute, particularly with a view to finding a procedure which could be used by the student on an uuknown solute. For the purpose of this study only water was cousidered for the solvent and the common salts as solutes. Preliminary investigations in the laboratory showed that, if fairly accurate results are to be expected in the usual 2-3 hour laboratory period, the preferred method for determining the solubility of a solute a t various temperatures is to find the temperature a t which a solution of known composition is saturated, that is, at which the solute begins to crystallize out of solution. In the time available for this work seven salts were tried. The solubility curves of these were determined in triplicate. Four of these salts (KBr, KC108, KN03, and K2Cr2O7)were found to give results very close to the values listed in the ninth edition of Lange's "Handbook of Chemistry." The three other salts which were tried (K2SOl, NHCI, and AIPOI) gave erratic results, probably because of hydrate formation or supersaturation of their solutions. It was our objective to work out the details of procedure so that reliable results could be obtained by the average student with an unknown salt. This was especially important because of the wide variation in solubilities of these salts. The procedure given below 'EDDY,R. D., J. CHEM.EDUC.,35, 364-5 (1958).
Volume 37, Number
3, March 1960
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Now add 1 ml more distilld water, again heat as before to dissolve d l solid, allow to cool, and again determine the saturation temperature. Further additions of water, 1 ml at a time, should be made, and saturation temperature8 determined down to 30°C or less. If it is found that the saturation temperature drops rapidly after the first addition of 1 rnl water, it is advisable to add only 0.5 ml in order to get ss many points as possible for the solubility curve. On the other hand, if it appears that the saturation temperature drops very slowly with additions of 1 ml water, you may increase the water additions to 2 or 3 ml. When all the readings have been made, record them on the report sheet, noting carefully the total amount of water as solvent for each temperature recorded. Calculate the concentration of emh solution in terms of grams of salt per 100 g water, and record these values d m Then plot on a sheet of graphing paper, properly laheled, the concentration as ordinate against the temperature as abscissa.
Typical Student Results Temperature, 'C
A typical set of student data for the solubility of the four salts is found in the table. This tabular form for reporting experimental data and the calculated solubilities is required of each student, as well as a plot of the results.
has been tested in several classes during the past two years with satisfactory results.
Data Sheet and T v ~ i c a lResults
Solute (9) 5.33
The Experiment Fit an Gin. test tuhe with a two-hole stopper holding a 110°C thermometer extending to within ' / r in. of the bottom of the tube. A glass stirring rod with a loop a t its lower end extends out of the test tuhe through the second hole in the stopper, and should he of such diameter as to slide earily throxgh the stopper. Fill a. 400ml beaker a/, full of water, set it on a tripod or iron ring with an asbestos screen, and start heating it to boiling. Meanwhile weigh the test tube, without stopper, on the analytical balance to the nearest centigram. (This is most conveniently done by hanging the tuhe with a loop of copper wire from the balance pan support.) Obtain from your instructor a small envelope containing shout 5 g of an unknown salt, empty it into the weighed test tube, and weigh it again. Add to the tuhe exactly 3 ml distilled water, memured with a buret or graduated pipet. Attach the stopper with thermometer and stirrer, and clamp it in such 8. way that the level of solution in the tuhe is the same as the level of water in the beaker. While stirring well, warm the solution until all the solid is dissolved. If it has not dissolved when the water in the bath is boiling, add exactly 1ml more distilled water, or as much more as is necesssly to dissolve the salt. Be careful not to add more water than is absolutely necessary, and keep a record of the amount of water added. Also, do not keep the solution hot any longer than is necessary in order to prevent excessive evaporation of the water. After the salt has dissolved, raise the tuhe out of the water bath and allow it to cool, stirring the solution continuously. Record the temperature at which the first crystds appear in the solution. Repeat the warming and cooling to ohtain a cheek result.
KBr
Water (g)
6 6.5 7 7.5
Solution concentration (g solute per 100 g water) 88.8 82.0 76.2 71.0
-R.
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Journal of Chemical Education
70 55 41.5 32
The figure shows how the determined values (circled) compare with the data in the literature, represented by the curves.
Ingenious theoretical superstructures live in constant dread of factual termites that continually gnaw at their foundations. They topple at the first inconsistency with observation. Concepts glory only in a relatively short term of ofice. We have witnessed the kaleidoscopic changes from Genesis to Damin, from Ptolemy to Copernicus, from B q l e to uan der Waals,from Newton to Einstein. This does not necessarily mean that no statements can be made about reality. It just means that science herself deals in temporary hypotheses of perfection at any given moment.
138
Saturation temp. ("c)
G . H. S m in "The Tao of Science"
