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Univenity of Northern Colorado
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Crystallization of Supersaturated Sodium Acetate and the Temperature Dependence of the Autoionization Constant of Water Joseph A. Pergler, Ronald 0. Ragsdale, and Thomas G. ~ichmond' University of Utah, Salt Lake City, UT 84112 Crystallization of a supersaturated van't Hoff Plot for Acetate Ion solution of sodium acetate (NaOAc) is a popular lecture demonstration a n d even in a large auditorium students are fascinated by the rapid solidification of the solution upon seeding with a crystal of the solid ( I ) . The crystallization is exothermic and heat is evolved (AH= -19.7 kJ1mole) as the system achieves its thermodynamically stable state. We sought a simple means of qualitatively indicating the temperature rise in this experiment that would be visible to a large audience. Because the colorlessto-pink transition of phenolphthalein occurs in the same pH range as the hydrolysis constant of acetate ion, we chose to examine the behavior of this indicator. Our studies of this system ultimately resulted i n the discovery of a n effective means of qualitatively demonstrating the variation of the autoionization constant of water with temoerature and quantitative measurements Kb (a),and KdO). confirm our interpretation of this phe- Figure 1. Temperature dependence of Ka (0). nomenon. Several reports have described the use of acid-base indicators to demonVisible Absorbance Data for NaOAc System s t r a t e variation of equilibrium cons t a n t s with respect to temperature. These include using bromothymol blue in very pure water and phenolphthalein in dilute aqueous ammonia (2, 3). Both of these systems proved difficult to reproduce in our hands. In the former case, very pure water is required (for example, atmospheric COz reducing the pH of the solution becomes a 51 problem) and the color change is not a easily visible. In the latter case, the d -0.6equilibrium i s quite sensitive to the 2 concentration of the ammonia solution and again only a small color change is 0.8observed.
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Procedure Addition of several milliliters of phenolphthalein indicator solution to a su-
-1.0 1 2.8
2.9
3.0
3.1
3.2
3.3
3.4
looorr (K)
'Authorto whom correspondenceshould be Figure 2. Visible Absorbance Data for 5.0 M NaOAc with phenolphthalein (A = 554 nm) as a addressed. function of 1000/T(K) Volume 72 Number 11 November 1995
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persaturated solution (3.5 g NaOAc.3H20/mL Hz0,-8 M) of sodium acetate in a large test tube affords a solution that is almost imperceptibly faint pink. A distinct pink color that is easily visible against the white background of the solid salt appears as crystallization takes place and the temperature of the system rises to about 50 "C. This pink color fades over the course of a 45-min lecture period as the solid returns to room temnerature. An even more dramatic brilliant pink color is observed if the solution i s warmed in a water bath to 80-90 "C: cooline in ice causes the color to fade. The process is reversible and the viscous nature of the solution makes i t possible to produce a nearly colorless baud below the hot pink layer by cooling only the lower portion of the tube. Similar results are obtained in concentrated solutions i n the range of 2.0-5.0 M. Literature Background Eouilibrium constants are. of course. temperature dependent, but i t is not obvious which of the relevant equilibrium constants is most responsible for the observed color change. Interestingly, literature data indicate that & for HOAc is essentially independent of temperature over the range of 0-50 "C, with a maximum K, measured a t 25 "C and slight decreases a t higher and lower temperatures (4). Other &nple carboxylic acids show similar behavior that may be a consequence of hydrogen-bonded dimeric structurks for these acids. I n contra&, the autoionization constant of water, K, changes by nearly two orders of magnitude and follows van't Hoff behavior over this temperature range as shown i n the Figure l ( 5 ) . Thus, the temperature dependence of the derived equilibrium constant Kb for hydrolysis of acetate ion is controlled by the value of K,. Spectroscopic Measurements Quantitative absorbance measurements were performed on sodium acetate solutions that were 1.0 and 5.0 M and supersaturated as a function of temperature i n the presence of phenolphthalein indicator (2. max = 554 nm). The measured nH values for these solutions were i n eood agreement with those calculated from Kb. Spectra were recorded on a n H P 8452AUV-VIS diode array spectrometer using 1.00-em quartz cells and a thermostatted cell holder that was calibrated with a thermocouple. Typical results for a 5.0 M solution are illustrated i n Figure 2. The plot of absorbance versus 11T follows a logarithmic relationship (R2= 0.997). Similar absorbance versus temperature plot's were obtained for other concentrations and the overall magnitude of the absorbance change increased with increasing acetate concentration.
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Summary I t is perhaps surprising that, given the high "non-ideal" concentrations of the solutions studied, the experimental data obey a simple van't Hoff relationship. One advantage of high concentration of NaOAc in this experiment i s the pH of the more concentrated solutions i s closer to the pK, of phenolphthalein than i n dilute solutions and thus more dramatic color changes are easily observed. I n addition, the system is robust and not sensitive to atmospheric CO2 so a n open tube may be stored and easily reused. Finally
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Journal of Chemical Education
this experiment points to the importance of considering all the eouilibria associated with a svstem in attemotine . to explain the observed behavior. " Literature Cited 1. Shakhashiti. 8.2.Chemicd DemonLmtionr, Val. 1: University ofWiaeonsin: Madison, WI, 1981; p. 27. 2. Campbell, J. A. J. Chrm. Educ. 1910.47, 273. 3. Summeilin, L. R.; Ealy Jr, J. L. Chemical Demondrulhns; Ametiran Chemical soeietv: Washineion.. DC.. 1985:..o 62. 4. Harned, H. S; Embree. D. D. J. Am. Chem. Sac 1934.56, 1050. For a classic discursion ofionization equilibria seeHarned. H. S.:Owen. 8. B. ThePhysied Chsmialry ofElrctroiyficSolutions.3rd ed.: Reinhold: New York. 1958. Chapter 15. 5. Marshall. W L:Franck, E. U. J Phya. Chem. Ref. Data 1981. 10,295.
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