In the Laboratory
The Solubility of Ionic Solids and Molecular Liquids
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Carl Baer* and Sheila M. Adamus Department of Chemistry, Providence College, Providence, Rhode Island 02918; *
[email protected] An understanding of solubility is usually taken for granted by a practicing chemist. Solubility is usually explained in introductory texts in terms of intermolecular forces and solubility behavior is exploited to the chemist’s advantage in almost all areas of chemistry. The most common solubility problems that confront the chemist are choosing appropriate solvents for dissolution of reactants, precipitation of products, and purification of products via recrystallization. However, the most common experiments dealing with solubility at the introductory level involve determination of the solubility product of a sparingly soluble salt (for examples, see refs 1–5). In the two-part experiment described here, students observe solubility phenomena more applicable to the preceding problems both by measuring the solubility of three salts in water at several temperatures and by mixing pairs of molecular liquids and observing their miscibilities. Experiments have been reported in this Journal (6–8) for the measurement of solubilities of highly soluble ionic compounds, but the experiment described below utilizes a different, simpler technique for measuring solubility, is broader in scope (with the inclusion of molecular liquids), and makes use of computer spreadsheets to pool data and give the students an opportunity for much broader data analysis. Because of the last point the experiment is best utilized in courses where a full lab report is required. Both parts of this experiment are very easy to carry out and require little in the way of equipment. The solid solubility determination does require magnetic stirrers and small stir bars for each student. For the liquid miscibility determination, the experiment should be conducted under an efficient fume hood, and no flames should be present. Description of the Experiment
Part I. Solubility of Ionic Salts The salts used in this experiment are sodium chloride, lead(II) chloride, and potassium aluminum sulfate dodecahydrate. Their solubilities in water are measured at approximately 5 °C, room temperature, 50 °C, and 80 °C (temperature control need not be rigorous). Solubilities are determined by addition of small portions of each preweighed salt to a test tube containing 5 mL of water and a small stirring bar until the saturation point is reached. Thus the solubility point is found to be between the amount of salt that totally dissolves and that which does not. This method has advantages over a previously published experiment (6 ) in which a hot solution is cooled until crystallization is observed, although both methods give acceptable accuracy. The method described here allows the student to perform other tasks (see below) while their latest salt addition is dissolving and avoids the problem of supersaturation. Additionally, waste is minimized, since at least three of the temperature points for a given salt can be obtained from a single test tube. The set of salts in this experiment has been chosen so that very different solubility behaviors are observed. Sodium chloride’s solubility changes
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little with increased temperature, whereas the solubility of the other two salts changes greatly (more than an order of magnitude); however, lead(II) chloride’s solubility is about two orders of magnitude lower than potassium aluminum sulfate dodecahydrate’s. Data treatment is an important aspect of this experiment. Since obtaining the data at the four temperatures takes some time, each student is assigned to work with only one of the three salts in the lab. Student data from each laboratory section are placed on a server accessible from the college’s computer labs. Each student is then required to (i) determine if any data are faulty,W (ii) decide on an appropriate format and graph the solubility vs. temperature profile for each salt, and (iii) compare and contrast the results for each salt. A final feature of this part of the experiment is the introduction of the concept of recrystallization. After the high-temperature solubility is measured, the student cools the solution and observes the result. In the lab report each student is required to give a specific procedure for recrystallizing 100 g of potassium aluminum sulfate dodecahydrate, based on the solubility-versus-temperature curves obtained.
Part II. Molecular Liquids and Solubility The liquid solubility portion of this experiment is easily carried out concurrently with the solid solubility measurements, since the solid must be allowed time to dissolve after each portion is added. Seven liquids are tested for miscibility: hexane, toluene, diethyl ether, acetone, isopropanol, methanol, and water. Students simply make all possible combinations of pairs of these and observe whether each pair is miscible or forms two layers in a test tube. Using the dielectric constant of each substance as a measure of polarity, students are asked to identify any trends in the solubilities and discuss them in terms of the polarities of the molecules. An interesting result is that the generally useful rule “like dissolves like” does not seem to describe the solubility trends very well: methanol and toluene, with dielectric constants of 33 and 2.4 D, respectively, are miscible, as are water (79 D) and isopropanol (18 D). The problem lies in exactly what is meant by a “like” solvent and can be resolved if the definition of “like” is broadened. In realizing this dilemma and then thinking about it, students come to a greater understanding of the role of intermolecular forces and entropy in determining solubility.W Further, they are exposed, perhaps for the first time, to the limitations of a general rule in chemistry, which can be a valuable learning experience in and of itself. Summary This experiment has been conducted at Providence College in the second semester of our first-year chemistry courses since Spring 1993. Results for the salt solubilities are generally good, although it is typical for several students in a section of 16–20 to obtain poor solubility results for their
Journal of Chemical Education • Vol. 76 No. 11 November 1999 • JChemEd.chem.wisc.edu
In the Laboratory
assigned salts.W This is, in fact, desirable, as it forces each student to make decisions about which data are poor among the fairly large number of solubilities reported by the class. The utilization of a relatively large amount of data is one of the major strengths of the experiment, as students in introductory chemistry usually only get to see their individual results. In the liquid solubility portion, results are also good. Both parts of the experiment give students ample opportunity to reach their own conclusions about the solubility behavior of a variety of chemical substances. The lab report is a crucial part of the experiment, as the students have to select the best data and decide on the best way to present it.W The experiment is especially well suited for group learning where the group is responsible for submitting one report. It is also well suited to illustrating the power of computer spreadsheet analysis of data.
Note W Supplementary materials for this article are available on JCE Online at http://jchemed.chem.wisc.edu/Journal/issues/1999/Nov/ abs1540.html.
Literature Cited Thomsen, M. W. J. Chem. Educ. 1992, 69, 328. Gotlib, L. J. J. Chem. Educ. 1990, 67, 937. Wruck, B.; Reinstein, J. J. Chem. Educ. 1989, 66, 515. Sawyer, A. K. J. Chem. Educ. 1983, 60, 416. Gasparro, F. P. J. Chem. Educ. 1976, 53, 98. Wolthius, E.; Pruiksma, A. B.; Heerema, R. P. J. Chem. Educ. 1960, 37, 137. 7. Pacer, R. A. J. Chem. Educ. 1971, 48, 225. 8. Pacer, R. A. J. Chem. Educ. 1984, 61, 467. 1. 2. 3. 4. 5. 6.
JChemEd.chem.wisc.edu • Vol. 76 No. 11 November 1999 • Journal of Chemical Education
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