R. J. Kokes, M. K. Dorfman,
and 1. Mathia The Johns Hopkins University Baltimore, Maryland
I I
Experiments for general chemistry V
Equilibria in Ionic Solutions
In conjunction with lectures on equilibria in electrolytic solutions, students perform a set of laboratory experiments in which they determine both the ionization constant of chloroacetic acid (a weak electrolyte) and the solubility of its silver salt. The experiments are quantitative and involve corrections for volume changes on mixing. Thus, they serve The new freshmm laboratory course at Johns Hopkins has been described in THIS JOURNA&38, 16 (1962). Many of the experiments are innovations in an introductory course. This series of articles describes the experimental procedures in some detail.
both to verify the mass action law and to give the students experience in the mathematical expression of equilibrium conditions in solution. The experiments require several laboratory periods; the students work in pairs. First, the solubility product of silver chloroacetate is determined in a nearly neutral solution containing varying concentrations of silver ion. Secondly, the solubility of silver chloroacetate is determined as a function of pH; the ionization constant of chloroacetic acid is calculated from the variation of the solubility with pH and the previously determined value of the solubility product. Thus, the accuracy of the student's calculation of the ionization
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constant is contingent upon the accuracy of his determination of the solubility product. For the determination of the solubility product of silver chloroacetate, the students prepare three solutions containing varying proportions of silver nitrate and sodium chloroacetate a t such concentrations that the solubility product of silver chloroacetate is exceeded. Stock solutions of standard silver nitrate and sodium chloroacetate are supplied. Precipitates of silver chloroacetate form immediately, and after 24 hours equilibrium is attained. Each solution is then filtered and a 10-ml sample of the filtrate is pipetted into a beaker containing 10 ml of distilled water and 5 ml of 6 N HN03. After the solution is heated to boilmg, 5 ml of 6 N HC1 is added to precipitate silver chloride from the solution. The precipitate is digested briefly in the boiling solution and is transferred to a weighed fritted glass filter where it is washed with methanol and dried to constant weight by suction. From the known added amounts of silver nitrate and sodium chloroacetate, (Ag+)o and (C~AC-)~, and the weight of silver chloride precipitated from the sample, the solubility product of silver chloroacetate can be calculated. Sample calculations from a student's notebook are shown below. (Ag+), = 0.0976 M ( C ~ A C -= ) ~0.0965 M We~ghtof AgCl precipitated from a 10 ml sample = 0.120 WsJna
0.120 Moles of AgCl precipitated = -= 8.37 X lo-' 143.3 (Age). - (AgC), = Concentration of silver precipitated a s the chloroacetate = 0.0976 - 0.0837 = 0.0139 M (ClAc-), = (CIAc-)* - 0.0139 = 0.0965 - 0.0139 = 0.0826 M K, = ( A g + ) , ( c l A ~ -= ) ~(0.0826) (0.0837) = 6.91 X lo-'
For the determination of the ionization constant of chloroacetic acid a more complex procedure is involved. First, each pair of students carefully prepares solutions of 2.0, 1.0, and 0.5 N nitric acid; the concentrations of these solutions are accurately known since they are made up by quantitative dilution of a standardized stock solution. Enough solid silver chloroacetatel is added to each solution to assure that some solid remains after equilibrium has been established. After the solutions have stood for a t least 24 hours a 5-ml sample of the supernatant liquid is pipetted into a cork-stoppered Erlenmeyer flask of known weight. The filled flask is then reweighed to determine the density of the solution. The silver ion concentration of each solution is then determined by the procedure described above.
' Silver chloroacetate for use in this experiment was prepared by the method described by MACDOUQALL, F. H., A N D REHNER, J., J R J . Ant. Chem. Soe., 56,368-72 (1934).
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In the treatment of the data it is essential that a correction be made for the change in volume of the solution when silver chloroacetate is dissolved. A sample calculation is shown below. Initial concentration of HNO. = (Ht). = 2.1088 M Measured density of solution = d, = 1.280 g/ml Weight of AgClprecipihted from a 5 ml sample = 1.192 g Moles of AgClprecipitated =
c 2 0.83 X 10-Pmoles 143.3 =
Conc. of silver = 2 X 100 X 0.83 X 10-2 = 1.66 formal K., = 5.9 X lo-' (ClAc-) = - 3.55 X 10-'F (Apt) 1.66 (HClAc) = (Ag+) - (ClAc3 = 1.660 - 0.004 = 1.6568 Correction of ( H + )fmvolume change: Density of 2.1088 MHNOs = 1.069 grams/ml4 Weight of 10ml2.1088M HN03 = 10.69 gram Weight of AgClAc dissolved in 10 ml = (.0166) (201.4) = 3.34 P .~.--
+
Total weight = 10.69 3.34 = 14.03 g Measured density = d. = 1.280 glml 14.03 Find volume = - = 10.98 ml 1.280
Final(H+) = (Ht).., - (HClAc) = 1.924 - 1.656 = 0.268M (H+) (CIAc-) - (268) (3.55 X 10-8) = 5,74 Ki,, = (HClAc) 1.656 Table 1. Comparison of Student and Literature Values (Student values are averages of six determinations)
Quantity
pK (student)
pK (literature)
K., of AgClAc Ki of HClAc
2.195 0.036 3.093 *0.086
2.265" 2.921'
*
We have found the experiment a particularly good illustration of the principles of solution equilibrium. The first part is simple in procedure and in calculation. The second part provides the student with a classic example of complex equilibrium, one that is very important in qualitative analysis. The procedure and computation force him to become acquainted with the meaning of the defined units of solution chemistry. Finally, these experiments require quantitative analytical techniques. Small errors lead to negative concentrations in the second experiment. Thus, the careless student often obtains au indisputable illustration of the conseauences of s l o ~ work. ~v '"Handbook of Chemistry and Physics," 3.rd ed., Chemical Rubber Publishing Company, Cleveland, Ohio, 1951-52, p. 1686. 3 Calculsted from density data in HILL, A. E., A N D SIMMONS, J . P., J . Am. Chem. Sac., 31,821 (1909). 'I % t e n a r i m l C ~ X c aTables l 6 : 245 (19331.
Acknowledgment Grateful acknowledgment is made of a grant from the National Science Foundation for new teaching aids in chemistry, which made the development of these experiments possible.
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Journal of Chemical Education
(f:,)
Corrected acid cone. = (H+),,, = 2.1088 - = 1.924M
