I
dent Ex~erimentslnvolvina
Elementary experiments concerning solubility product constants normally involve the determination of values already reported in standard textbooks. Campbell (I), in discussing experiments in use at Harvey Mudd College, feels that well designed student experiments should: involve an unknown and ask questions, the answers to which are not readily available to the students; be designed so the students learn to synthesize, separate, purify, and identify chemical species; ask the student to interpret experimental measurements in terms of molecular forces and structures; allow the student, wherever possible, t o learn to operate standard pieces of equipment and to understand their capacities and limitations. These criteria are seldom found in solubility experiments normally involving salts commonly found in any laboratory. Most solubility experiments employ either titrations or gravimetric methods rather than modern standard instrumentation. While some experiments allow for student interpretation of results, only a few solubility experiments allow for studies of ionic effects (I?, S), temperature and concentration changes (b), or solvent effects such as those described by Parker (4). I n designing introductory experiments in solubility equilibrium, it is diicult to allow a student to work on a problem whose answers are not found in their textbooks or in readily available published volumes. If the students are asked to calculate reported values from experimental data, the only unknown is the reason why these constants may not be the same from solution to solution or agree with the reported values. Although the students can be asked t o explore and discuss the reasons why these values are not in precise agreement with the reported values, the time available limits the collection of experimental evidence needed t o support student explorations or discussions. 1 Present Address: University of California, Department of Consumer Sciences, Davis, California 95616.
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Journal of Chemical Education
Two methods which might be used to allow the
K,, values to remain "unknown" to the students are: (1) to require them to work with uncommon salts whose characterizations have not been recorded in the readily available literature or (2) to use solvents or solvent mixtures other than pure water. The nature of the solvent is certainly an important factor in solubility and is discussed in many elementary text books as well as in the article on "super solvents" by Parker (4). The use of uncommon salts could allow students to synthesize and isolate their own "unkno~m." In this study both of the above methods are used. This experiment is designed in three parts. First the student learns to synthesize, separate, purify, and identify some insoluble metallic salts of sorbic acid and cinnamic acid. Although many of these salts are not reported in the readily available literature, the acids and their derivatives are used in industry. Sorbic acid and derivatives are used as fungistatic agents for foods and as additives to drying oils, resin coatings, and rubber. Cinnarnic acid or its derivatives are used in the perfume industry and in the manufacture of materials used in medicine (5). I n the second part of this experiment the purified insoluble salts of these acids are characterized through melting points (or decomposition points), equivalent weight determinations, and chemical spectra. The third part of this study involves investigating the solubility of these "unknowns" in various solvents and under many conditions. These investigations can be easily performed due to the absorption characteristics in the ultraviolet spectra of these two acid anions. Synthesis, Separation, and Purification
The students are provided with sorbic acid or cinnamic acid and a solution containing either silver, c o p per, lead, or zinc cations. They are asked to synthesize, separate, and purify one of the possible salts,
for example, lead sorbate. The students should become aware that the "insoluble" acids dissolve in basic solutions and form the sodium, potassium, or ammonium salts. Once the acid is neutralized into solution, many students will combine the desired cation and anion and expect to see the correct product. However, when ammonia is used to dissolve the acid, the students should he aware of the silver, copper, or zinc ammoniated complex ions which, when formed, inhibit the formation of the desired product. When sodium or potassium hydroxide solutions are used to dissolve the acid, the possibility of the insoluble hydroxides (or oxides) of the mctals should be discussed. The students also need to become aware of the amphoteric nature of both zinc and lead in the strongly basic solutions as these complex anions will not allow precipitation of the desired product. These factors emphasize the importan~eof pH control which should he used during the elementary synthesis. Many of the above phemonena can he demonstrated during the formation of the white silver sorhate salt. The soluble silver ammoniated complex ion, formed when excess ammonia is used to dissolve the acid, prevents the formation of the desired salt; and the dark brown solid, silver oxide, is obtained instead of the desired white preripitat,e if the anion solution is too basic. The desired compounds are then separated by employing a Biichner funnel and vacuum filtration and are purified by washing with small amounts of water followed by methanol or ethanol. The organic solvent helps remove most of the water as well as any free acid which may be present. The melting points (or decomposition points), and the visible and ultraviolet spectra can he determined to help characterize the product. The identity of the products may be checked by reacting a portion of the compound with some dilute acid. This frees the organic acid which can be removed by filtration. A positive qualitative test for the cation on the remaining solution identifies the starting salt. The approximate equivalent weight of the new compound can be determined by titrating the cation using standard methods or by employing a spectrometric method. Both sorbic and cinnamic acids absorb in the ultraviolet regions (257 and 273 mp respectively), while the concentration of copper salts can he determined by the intensity of the blue color of the copper-ammonia complex in the visible regions. In both cases the concentrations of cations or anions depend upon the absorbance of light at a particular wavelength. Beer's law can he introduced without the students having had calculus (6, 7). More explicit directions should be given to the students in the laboratory. Spectroscopic Determination of the K.,
Each student is asked to determine the K,, for a sorbate or cinnamate salt which he has either been
assigned or synthesized as described above. A solvent or a dilution of two solvents is then assigned in which the K,, can he determined. After the solute and solvent are allowed to come to equilibrium (these mixtures should be kept in the dark overnight), the undissolved molecules can be filtered from the solution. Some of the solution can then be transferred into the sample side of the cuvet of the spectrophotometer so that the absorbance can be measured. Proper dilution should he made so that the absorbance is between 0.1 and 1.0. If the solvent itself absorbs, dilutions can he made with methanol or water. In these cases identical dilutions with the pure solvents are made to use as the Solubility Product Constants, K . , for Lead Sorbate at Room Temperatures (22°C) in Various Solvents Solvents Acetic acid Dimethylsulfaxide Methanol Water Ethenrd Diethyl ether Iaooctane
2 2 3 2 2 1
K.. Large x 10-3 x 10-7 X lo-" X 10-9 X x 10-17
reference samples. A calibration graph used to calculate the concentration of ions in solution can either be made available or made by all the students, each using a different concentration of dissolved sorbic or cinnamic acid in a slightly basic solution. After determining the concentration of the anions, the K,, in the appropriate solvent can be calculated. Using these anions and the spectrophotometer, the K,;s can be determined over a large range. The table gives these values ranging from 10-"10-l8 for lead sorbate at room temperature in the various solvents listed. Acetic acid was wed as a solvent to stress to the students that the solvent is a chemical and that it can react directly with the solute molecules. Because of the ease in determining the K,, using spectrometric methods, the students, collectively or individually, can investigate several equilibrium mixtures and discover the effectof ionic environment on solubility (3). Investigations similar to those st,udyinginteractions of salts in methanol-water systems ( 2 ) and those studying the nature of dissolved silver acetate (8)can also be easily performed following these procedures. Literature Cited (1) CAMPBELL, J. A., J. CHEM.EDUC.,42,488 (1965). R. T., J. CEEEM. EDUC.,42, (2) NEIDID,H. A,, AND YINGLIND, 476 lQR.5~ -.- l\A*"v,. (3) GOODMAN, ROBERT C., AND PETRUCCI, RALPHH., J. CHEM. Enuc., 42,104 (1965). (4) P~RKER, A. J., Intern. Sci. and Technol., 28, (1965). (5) STECHER, PAULG., "THEMERCK INDEX,"Merck and CO., Ine.. Rahwrw. N. J.. 7th ed.. 1960. D. 264 and D. 968. R;CH.& C., J. CHEM.'EDUC., 41, i66 (1964). (6) PINK~ERTON, C. C., J. CHEM.EDUC.,42,226 (1965). (7) SOLOMONB, (8) RAMETTE, RICHARD W., J. CHEK.EDUC.,43,299 (1966).
Volume 44, Number 2, February 1967
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