Solubility and K,, of Calcium Sulfate: A General Chemistry Laboratory Experiment For several years we have used this laboratory experiment to accompany the lecture topics of solubility and of chemical equilibrium, which are in sequence in our lecture format. However, the experiment could he used t o accompany either topic, independently. The experiment is a simple, direct determination of the solubility of CaS04 by the evaporation t o dryness in a 50-mL beaker of a 10.0-mL sample of a saturated solution of CaSOa using a hot plate. From the weight of the empty beaker and the weight of the beaker with the residue, the weight uf anhydrous CaSOn can be found. After evaporation t o dryness the solid residue is heated for about 5 min a t a hot plate setting of about 250DCto ensure that the CaSOd is anhydrous, and care is also taken to protect the anhydrous salt from rehydrating prior to weighing by covering the beaker with a piece of aluminum foil. Two determinations are carried out and the solubility of CaSOh in moles per liter of solution is calculated for each. From the average value the K,, of CaSOa is calculated. Whereas we use analytical balances with a precision of 0.1 mg, a less precise balance could he used if more sample were taken. Aprelab exercise is used to require that thestudent reads over and understands what is to he done heforecoming tolab and to focus attention on fundamental ideas such as the nature of chemical equilibrium, LeChatelier's principle, solubility, saturation, and molarity. In the introduction to the experiment there is a section summarizing the nature of chemical equilihrium, the general expression for the equilibrium constant, and a specific example of the expression for the solubility product constant far silver chloride. We have the students complete the experiment and the writeup in the same 3-hr laboratory period. If it were desired, the students could he asked to make their own saturated solutions; however, this takes more time and results are not as consistent. In addition t o weighing practice and protection of a sample from taking up moisture, pipeting technique is also introduced using the simple inexpensive pipeting device described by the author in J. CHEM.EDUC.,45,733 (1968). The class averages obtained for KSp,ranging from about 2.2 X 10W4to 3.1 X lo-', agree well with that of 3.1 X 10-4calculated from solubility data in the "Handbook of Chemistry and Physics." K,, values listed in tables in qualitative analysis texts vary considerably and are low by a power of ten. However, one text by C. J. Brockman [Ginn and Co., 19301,lists a value of 2.3 X 10-4close to that which we have obtained experimentally over a period of about four years. We do not a t this time know the source of the considerable variation found in some of these tables. Of particular interest is the low cost of instituting and carrying out the experiment. All residues of anhydrous CsS04 are collected in one flask and used for preparation of saturated solutions in subsequent years. Three hot plates can take care of a 24-person lab, and we estimate the one-time cost for hot plates, beakers, and chemicals t o he under $240. In our case, the same laboratory will he operated approximately 10 times in one year, so that the cost per lab meeting time becomes rather small and is made proportionally less in subsequent years as the continuing cost is only breakage. Albert K. Sawyer University of New Hampshire Durham. NH 03824
416
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of Chemical Education
