In the Laboratory
Preparation and Analysis of Multiple Hydrates of Simple Salts Richard W. Schaeffer,* Benny C. Chan, Shireen R. Marshall, Brian Blasiole, Neetha Khan, Kendra L. Yoder, Melissa E. Trainer, and Claude H. Yoder Department of Chemistry, Franklin and Marshall College, Lancaster, PA 17604; *
[email protected] Many salts form more than one hydrate: MnC12ⴢ4H2O crystallizes from water at room temperature, whereas MnCl2ⴢ2H2O can be obtained by crystallization at 65 °C; CoCl2ⴢ6H2O crystallizes at room temperature, but above 52 °C, CoCl2ⴢ2H2O forms and above 90 °C, CoCl2ⴢH2O is stable; CaCl2ⴢ2H2O is available commercially as a desiccant, but CaCl2ⴢ6H2O is readily formed at room temperature (1). We have developed a laboratory project in which the student prepares a series of hydrates of simple salts and then determines the extent of hydration of the product(s). In conjunction with previously published articles in this Journal, we believe this project provides a good introduction to the concepts of solubility, saturation, recrystallization, relative compound stability (e.g., a dihydrate vs tetrahydrate at elevated temperature), and simple gravimetric analysis (2, 3). Moreover, the project lends itself to many variations. For example, a student could be given a “starting” hydrated salt and asked to prepare another hydrate within the specified temperature range. Or students could be given the formulas of several hydrates stable over different temperature ranges and be asked to “discover” a method of preparation. If it is desired to extend the project, the cation could be determined quantitatively by gravimetric precipitation of the oxalate, for example, and the anion, chloride or sulfate, could be determined by precipitation with silver or barium ion, respectively, or by a variety of common volumetric and colorimetric methods (4 ).
Table 1. Summary of Student Data for Hydrated Salts Target Hydrate
Stability Starting Species Synth Range/°C (CAS No.) Method
Hydration (%) Exptl
Theor
ZnSO4ⴢ6H2O
39-60
ZnSO4ⴢ7H2O (7446-20-0)
EM
39.0
40.1
ZnSO4ⴢH2O
6–l00
ZnSO4ⴢ7H2O (7446-20-0)
EM
9.68
10.0
CaCl2ⴢ6H2O
55–29
CaCl2ⴢ2H2O (10035-04-8)
TDM
47.6
49.3
CoCl2ⴢ2H2O
49–137
CoCl2ⴢ6H2O (7791-13-1)
DH
21.9
21.7
CoCl2ⴢ2H2O
49–137
CoCl2ⴢ6H2O (7791-13-1)
EM
24.3
21.7
MnCl2ⴢ2H2O
5–135
MnCl2ⴢ4H2O (13446-34-9)
EM
22.8
22.3
MnCl2ⴢ2H2O
5–135
MnCl2ⴢ4H2O (13446-34-9)
DH
23.2
22.3
FeCl2ⴢ2H2O
11–155
FeCl2ⴢ4H2O (13478-10-9)
DH
21.3
22.1
FeCl2ⴢ2H2O
11–155
FeCl2ⴢ4H2O (13478-10-9)
SAM
22.4
22.1
The preparation of ionic hydrates can be accomplished by four methods: 1. slow evaporation of the solvent from a near saturated solution of the starting hydrate at a temperature within the stability range of the desired hydrate 2. crystallization within the temperature range of the target hydrate from a saturated solution prepared at higher temperature (assuming an endothermic heat of solution) 3. crystallization within the temperature stability range from mixed solvents (“knocking the compound out of solution” by adding a solvent such as 2-propanol to a nearly saturated aqueous solution) 4. heating a higher hydrate to the temperature range of the desired lower hydrate
We have found that the hydrates given in Table 1 are easily prepared by the method indicated. Determination of the water of hydration by heating gently in the Bunsen flame produced values generally within 1–5% of the theoretical, clearly sufficient to determine an integral number of waters of hydration. Probably the greatest experimental problem resides in filtration and removal of excess water and simultaneously protecting the lower hydrates from adsorption of atmospheric water. Hazards C AUTION : Some of the compounds used in this experiment are considered “moderately toxic” (e.g., cobalt chloride LD50, rat, oral = 766 mg/kg) and should be handled with care to avoid ingestion or inhalation of vapors or dust. 2-Propanol (CAS 67-63-0) is volatile (0.05 atm at 20 °C) and flammable, and should not be ingested (LD50, rat, oral = 5840 mg/kg). We recommend that all reagents be handled in an appropriate hood. We recommend the use of gloves and approved safety goggles. Synthesis of Hydrates
Evaporation Method (EM) A nearly saturated solution of approximately 2 g of the commercially available ionic hydrates from Table 1 was prepared in an 8-in. test tube or 50-mL beaker by adding small amounts of water (usually about 3–5 mL) until the solid dissolved. A 0.5-mL portion of 2-propanol was added to the solution and the test tube was placed in a water bath or sand bath maintained at a fairly constant temperature (an ordinary hot-plate provided sufficient temperature control) within the stability range of the desired hydrate (see Table 1). The solution was allowed to evaporate until crystals formed. In
JChemEd.chem.wisc.edu • Vol. 77 No. 4 April 2000 • Journal of Chemical Education
509
In the Laboratory
no case was the solution allowed to evaporate to dryness. The rate of evaporation was slowed in some cases by covering the test tube or beaker with Parafilm tape into which several holes were punched. The crystals were isolated by suction filtration, washed with 2-propanol, and placed in a desiccator overnight.
Temperature Difference Method (TDM) A nearly saturated solution of the commercially available hydrate was prepared by placing about 2 g of the starting metal hydrate in an 8-in. test tube or 50-mL beaker, which was placed in a water bath or sand bath maintained at a temperature above the maximum limit for the target hydrate. Distilled water was added dropwise to the test tube until all the solid was dissolved. One or two additional drops of distilled water were added and the test tube was again placed into a water or sand bath maintained at a lower temperature within the temperature range for stability of the target hydrate. The target hydrate typically crystallized within 30 min. The crystals were isolated by suction filtration, washed with 2-propanol, and placed in the desiccator overnight. Solvent Addition Method (SAM) A working solution of an ionic hydrate from Table 1 was formed as described in the EM procedure. The test tube or beaker containing the working solution was placed into a water or sand bath maintained at a temperature within the stability range of the target hydrate. After the temperatures equilibrated, 2-propanol was added dropwise until a precipitate formed. (Other weakly complexing solvents that are miscible with water but less effective than water at solubilizing ionic materials could also be used to precipitate the hydrate. 2-propanol is a good choice because it is readily available, reasonably safe, and inexpensive.) The crystals were isolated by suction filtration, washed with 2propanol, and placed in the desiccator overnight. Direct Heating (DH) Approximately 2 g of the starting hydrate was heated in an oven for 24 h at a temperature within the stability range for the hydrate desired. The product hydrate was stored in a desiccator until analyzed. Determination of Water of Hydration A 0.5–1.0 g (± l mg) hydrate sample was weighed into a porcelain crucible and heated gently with a low Bunsen burner flame. After approximately 20 min the heat was increased until there was no further evidence of the release of water. The sample should not be heated so strongly that decomposition occurs (as evidenced, for example, by an extreme change in color—from blue or green to black in the case of iron and copper salts). After cooling in a desiccator the sample was reweighed and the percent water of hydration was calculated from the difference in weight.
510
Information for the Instructor In this project, a series of inorganic ionic hydrates are prepared and analyzed for water content. The concepts of hydration, solubility, saturation, crystallization, compound stability, and simple gravimetric analysis are illustrated and reinforced in a straightforward experiment. We believe the experiment is probably most suited for the general chemistry laboratory, and that the variations possible in the hydrates prepared and analytical methods employed within the framework of the procedure described make this a versatile project. A glance at Table 1 reveals that not every synthetic method was used for each target hydrate. This paper is not meant to be an exhaustive report on the preparation of specific hydrates, but rather a report of an undergraduate-level experiment suited to illustrate the topics mentioned. Each of the synthetic methods will work with a variety of common hydrates with the exception of the thermal difference method (TDM), which may not generate product if the target hydrate has a higher solubility at a lower temperature or has nucleation problems. In one application the instructor could assign a student or small group of students the synthesis and analysis of a hydrate from Table 1 using all four techniques. Alternatively, multiple hydrates could be assigned using one or more synthetic approaches. The student would then “discover” which approaches seem to work best for each hydrate. Each of the synthetic methods described was tested by undergraduate students and the results are summarized in Table 1. Students were very successful in isolating the various hydrates by each of the techniques. Moreover, the analysis of the water of hydration was successful on the basis of a comparison of the expected mass percent of water from the “ideal” formula with the experimental mass percent (see columns 5 and 6 in Table 1.). The average percent error for the comparisons reported was 3.4% (2.3% if one value with an especially large error, cobalt(II) chloride via the evaporation method, is neglected). The instructor will probably want to include a discussion of the nature of hydrated salts (i.e., how is the water incorporated into the structure) and the effect of temperature on the extent of hydration. We recommend the excellent discussion by Kauffman and Baxter (5) as a guide. Literature Cited 1. CRC Handbook of Chemistry and Physics, 75th ed.; Lide, D. R., Ed.; CRC: Boca Raton, FL, 1994. 2. Silber, H. B. J. Chem. Educ. 1972, 49, 586–589. 3. Kielland, J. J. Chem. Educ. 1937, 14, 412–414. 4. Procedures can be readily found in many laboratory manuals. See, for example: Day, R. A.; Underwood, A. L. Quantitative Analysis, 5th ed.; Prentice-Hall: Englewood Cliffs, NJ, 1986. 5. Kauffman, G. B.; Baxter, J. F. J. Chem. Educ. 1981, 58, 349–353.
Journal of Chemical Education • Vol. 77 No. 4 April 2000 • JChemEd.chem.wisc.edu
