LABORATORY EXPERIMENT pubs.acs.org/jchemeduc
Solubility and Solubility Product Determination of a Sparingly Soluble Salt: A First-Level Laboratory Experiment Raffaele P. Bonomo,*,† Giovanni Tabbì,‡ and Laura I. Vagliasindi†,§ †
Dipartimento di Scienze Chimiche, Universita degli Studi di Catania, Viale A, Doria 6, 95125 Catania, Italy Istituto di Biostrutture e Bioimmagini-CNR-Catania, Viale A. Doria 6, 95125 Catania, Italy § Consorzio Interuniversitario C.I.R.C.S.M.B., Via Celso Ulpiani, 27, 70125 Bari, Italy ‡
bS Supporting Information ABSTRACT: A simple experiment was devised to let students determine the solubility and solubility product, Ksp, of calcium sulfate dihydrate in a first-level laboratory. The students experimentally work on an intriguing equilibrium law: the constancy of the product of the ion concentrations of a sparingly soluble salt. The determination of solubility is proposed either in an equimolar precipitation of CaSO4 3 2H2O, from which the Ksp is obtained, or working with an excess of one of the two reagents. KEYWORDS: First-Year Undergraduate/General, Inorganic Chemistry, Laboratory Instruction, Physical Chemistry, Hands-On Learning/Manipulatives, Aqueous Solution Chemistry, Precipitation/Solubility
T
he solubility product law is one of the most difficult concepts to deal with in a general chemistry syllabus. It is difficult for the students to understand what is happening in solution and it is not easy to foresee all the consequences of varying the reagent concentrations. It is hard to understand that there are dissolved ions in solution in equilibrium with a solid precipitate and the ion concentrations are linked together and to the solid precipitate amount through the constant called solubility product, indicated by Ksp. Many general chemistry textbooks treat this problem in an excellent fashion. For this simple experiment, calcium sulfate dihydrate, CaSO4 3 2H2O (gypsum), is precipitated. Other simple experiments involving the same salt have been proposed in this Journal such as an ion exchange complexometric titration experiment,1 a quantitative determination of calcium sulfate (anhydrite) by gravimetric analysis,2 and a titration of hydrogen ions displaced by the calcium sulfate saturated aqueous solution when passing through a ion-exchange column.3
could be affected by an excess of saline solution and the precipitation could be influenced by the presence of ion-pair formation. Given the small, but significant solubility of CaSO4 3 2H2O, the precipitation is very slow. If the solution is allowed to sit overnight, large perfect crystals are obtained that can be easily filtered. However, the gypsum crystals are not of interest, only the determination of calcium ion content in the supernatant solution. To speed up the precipitation process, the beaker is warmed to 50 °C and allowed to cool to room temperature. While cooling, the beaker is covered with paraffin film (or a watch glass) to avoid evaporation as the solution and the precipitate must be accurately assessed. (It may be convenient to prepare the solution at the end of a laboratory session and start the filtration procedure in the next laboratory session.) The crystals of CaSO4 3 2H2O are filtered off by means of glass or polypropylene short-stem funnel and moistened filter paper. The supernatant solution is filtered into another 50 mL beaker; this filtration is the most critical operation and not a single drop of solution should be lost. The crystals are washed inside the beaker with two 2 mL of cold water portions at ∼4 °C to avoid dissolving the calcium sulfate. It is important to know the exact volume of the washing solvent because it will be added to the known 20 mL initial volume. The solution is then heated in a hot plate at a temperature slightly higher than 80 °C (the temperature can be monitored by putting a thermometer in another 50 mL beaker containing water) and treated with 10 mL of sodium oxalate, Na2C2O4, 0.1 M solution to precipitate the calcium oxalate monohydrate, CaC2O4 3 H2O, the solubility product of which is Ksp = 2.6 10 9. Precipitation from a hot solution favors the formation of easily filterable crystals and it is best to leave the beaker with the white precipitate on the hot plate for a while.4 Calcium oxalate is filtered by means of a preweighed filter paper sheet (on an analytical balance having 0.1 mg precision) in a funnel or by means of a preweighed glass
’ EXPERIMENT OVERVIEW A simple laboratory experiment is proposed that can be performed in an approximately two- or three-day laboratory session (8 12 h) and allows students to understand that the supernatant solution still contains ions to induce a subsequent precipitation. It is possible to approximately determine the Ksp associated with the sparingly soluble salt by means of a gravimetric analysis of the further precipitate. ’ EXPERIMENT DETAILS Equimolar Precipitation
To obtain the precipitation of calcium sulfate dihydrate, CaSO4 3 2H2O (gypsum), 10 mL of a CaCl2 3 2H2O 0.1 M solution is put in a 50 mL beaker and charged with 10 mL of sodium sulfate, Na2SO4 0.1 M solution (total initial volume 20 mL). More concentrated CaCl2 3 2H2O or Na2SO4 solutions are not recommended because the subsequent washing procedure Copyright r 2011 American Chemical Society and Division of Chemical Education, Inc.
Published: October 31, 2011 545
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Journal of Chemical Education Gooch-type crucible with a capacity of about 30 mL and a porosity of 4 6 μm and connecting it through a filtering flask to a water aspirator pump. The precipitate must be washed many times with hot water (given the minimal solubility of this salt, it is better to be excessive in the washing procedure, at least 10 washes with 20 30 mL portions) inside the beaker. The precipitate is stored in a drying oven for 1 h without exceeding the temperature of 100 °C, to avoid partial dehydration of calcium oxalate monohydrate. Then, the filter sheet or the filtering crucible containing the precipitate is taken to room temperature and weighed and reweighed until the value is practically constant. From the mass of CaC2O4 3 H2O, it is possible to determine the number of moles, which are equal to those of Ca2+ present in the solution after the precipitation of CaSO4 3 2H2O. Given the volume (24 mL), the concentration of the Ca2+ ions can be calculated. Once the calcium ion concentration is known, the solubility, s, is obtained, and the solubility product of CaSO4 3 2H2O can be determined by means of the relationship, Ksp = [Ca2+][SO42‑] = s2. Values of solubility ranging from 0.013 to 0.015 M are determined by most of the students and consequently values of the solubility product from 1.7 10 4 to 2.2 10 4 can be calculated. Each student writes his or her own report, introducing the solubility product concept, the relationship between solubility and solubility product, and commenting on his experimental results. Excess Reagent Precipitation
Once the solubility product is calculated, it might be interesting to obtain the solubility of calcium sulfate dihydrate when working with an excess of one of the reagents, to allow students to investigate the intrinsic relationship between the relative concentrations of the two ions, Ca2+ and SO42 . Therefore, 10 mL of the CaCl2 3 2H2O 0.1 M solution and 20 mL of Na2SO4 0.1 M solution are put together (total volume 30 mL) in a 50 mL beaker, thus, precipitating calcium sulfate dihydrate in a solution containing an excess of SO42 ions (in order to optimize time in this laboratory session, this second precipitation could be carried out while washing calcium oxalate of the first precipitation). After the same filtering and washing procedure, the precipitation of CaC2O4 3 H2O is induced by the addition of 10 mL of 0.1 M sodium oxalate solution. From the calcium oxalate monohydrate mass, the initial Ca2+ ionic concentrations (total volume 34 mL) can be easily calculated and the new value of the solubility determined. In the report, each student is required to compare and draw conclusions about this new solubility value and to comment on the pertinent results. Also in this case, the washing procedure of the CaC2O4 3 H2O precipitate is the critical step, and the number of washes must be at least doubled: 20 washes with 20 30 mL portions with hot water. In these new conditions, values of solubility for CaSO4 3 2H2O ranging from 0.005 to 0.007 M are obtained. These values, which are about 1/3 or 1/2 of the previous solubility values, are a little larger than the theoretical value of 0.003 M, which can be calculated by using an average CaSO4 3 2H2O solubility product constant of 2 10 4.
’ HAZARDS Calcium and barium chlorides and oxalate compounds may cause irritation to the skin, eyes, and respiratory tract and may be harmful if swallowed or inhaled. Preparation of the reaction mixtures in each of the procedures can be safely performed with
LABORATORY EXPERIMENT
standard laboratory protective equipment, such as a laboratory coat, safety glasses, and gloves.
’ DISCUSSION Unfortunately, the literature value4 11 of the CaSO4 3 2H2O solubility product is not tabulated with a high degree of certainty, probably because of the salt precipitation in media with different ionic strength or the presence of different levels of ion association, and it varies from 1.42 10 4 to 2.58 10 5. The thermodynamic value of the CaSO4 3 2H2O solubility product is 2.4 10 5, because, as explained by Meites et al.,12 when dealing with 2:2 electrolytes, extensive pair formation occurs and therefore the activity of Ca2+ and SO42 ions cannot be considered equal to their concentration and the values of activity coefficients are needed as well as of the soluble CaSO4 species concentration. The same discrepancy was noted by those authors who previously proposed the determination of CaSO4 3 2H2O with different procedures.1,2 Most of students were surprised that their results did not coincide with their textbook value and this serves as a good opportunity to make them aware of the simplification adopted into these calculations as described by Meites et al.12 However, the task of this simple experiment is not the exact determination of the solubility product of gypsum, but the understanding of the chemical consequences of the solubility product law by the first-level students. In principle, the SO42 concentration can be evaluated in the same way too, determining it by BaSO4 precipitation (Ksp = 1.1 10 10) induced by the addition of a barium chloride solution. After dividing the initial 24 mL solution in two equal portions, the addition of a 5 mL of BaCl2 3 2H2O 0.1 M to one of the solutions produces an abundant precipitate of barium sulfate, but the washing and weighing of BaSO4 is a tiring procedure and not recommended. Analogous CaSO4 3 2H2O solubility results were obtained after a tedious washing procedure (often more than 1 L of hot water). Moreover, it is difficult to get rid of BaSO4, being one of the most insoluble salts of sulfate ions, so the subsequent cleaning procedure of the filtering crucibles could be arduous for first-level chemistry students. ’ SUMMARY This procedure is not far superior to previously published laboratory experiments for Ksp determination, but it is another way of approaching the problem. The CaSO4 3 2H2O solubility product is determined focusing on the equilibrium between the solid and its ion concentrations, which could be used for a further precipitation. It is appropriate for a first-year chemistry laboratory and uses simple glassware and low-cost equipment. This experiment has been used in the general and inorganic chemistry laboratory for many years. Students have appreciated the laboratory and have gained an understanding of the equilibrium between an insoluble salt and its ions in the supernatant solution. ’ ASSOCIATED CONTENT
bS
Supporting Information Student handout; notes for the instructor. This material is available via the Internet at http://pubs.acs.org.
’ AUTHOR INFORMATION Corresponding Author
*E-mail:
[email protected]. 546
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’ REFERENCES (1) Koubek, E. J. Chem. Educ. 1976, 53, 254. (2) Sawyer, A. K. J. Chem. Educ. 1983, 60, 416. (3) Masterman, D. J. Chem. Educ. 1987, 64, 408. (4) Kolthoff, I. M.; Sandell, E. B.; Meehan, E. J.; Bruckenstein, S. Quantitative Chemical Analysis; Macmillan: New York, 1969. (5) Hogness, T. R.; Johnson, W. C.; Armstrong, A. R., Qualitative Analysis and Chemical Equilibrium, Holt, Rinehart & Winston, New York, 1966. (6) Clifford, A. F. Inorganic Chemistry of Qualitative Analysis; Prentice Hall, Inc.: Englewood Cliffs, NJ, 1961. (7) Atkins, P.; Jones, L. Chemistry, Molecules, Matter, and Change; W. H. Freeman and Company: New York, 1997. (8) Handbook of Chemistry and Physics, 43rd ed.; The Chemical Rubber Publishing Company: Cleveland, OH, 1961 62. (9) Block, E. Can. J. Chem. 1961, 39, 1746. (10) Bennet, A. C.; Adams, F. Soil Sci. Soc. Am. J. 1972, 36, 288. (11) CRC Handbook of Chemistry and Physics, 77th ed.; CRC Press: New York, 1996 1997. (12) Meites, L.; Pode, J. S. F.; Thomas, H. C. J. Chem. Educ. 1966, 43, 667.
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