Oxalic acid (S); Benzoic acid (S); 500 ml of approximately 0.5 N NaOH which has been standardized using Benzoic acid; (P); phenolphthalein indicator (S); ice (L). By permission of instructor, boric acid or benzoic acid may be used in place of oxalic acid. If boric acid (S) is used, 250 ml glycerin (S) is required in addition to above chemicals; and if benzoic acid (S) is used, 250 ml of 95% ethanol (S) is required.

Note: For disposal purposes, the chemicals used in this experiment may be flushed down the drain with copious amounts of water.

CAUTION: NaOH(s) is caustic! Use a spatula to transfer the pellets when preparing the solution! NO HANDS! Weigh the pellets into a glass container! No weighing paper is to be used!

First, the sodium hydroxide solution that has been propared (» 0.5N) must be standardized using benzoic acid. The solubility of oxalic acid is determined at 25, 20, 15,10, 5, and 0° by analyzing for the number of moles of oxalic acid per 100 g of water in the saturated solution. A thermostat is generally available for 25°, and an ice bath is used for 0°. The intermediate temperatures are obtained with an improvised water bath made of a large beaker of water provided with a small motor stirrer. For temperatures below that of the room, ice or cold water is added as needed.

The saturated solution of oxalic acid can be placed in a medium-sized test tube, which in turn is surrounded by a still larger test tube to provide an insulating air jacket to reduce the fluctuations in temperature. The 0.1° thermometer is immersed in the solution, which is stirred vigorously by the hand operation of a vertical ring stirrer which fits closely in the inner test tube so that the projecting ring cannot break the thermometer at the center of the tube.

The distilled water in large test tubes is saturated with oxalic acid by shaking with an excess of crystals at a higher temperature and cooling the solution down to the thermostat temperature so that some of the dissolved material is crystallized out. The solubility is less at the lower temperatures. When the equilibrium is approached in this way from a supersaturated solution, it is achieved rapidly, whereas it may be achieved slowly if heated from a lower temperature so that crystals must be dissolved to give an equilibrium concentration.

Two 10-ml samples are removed at each temperature with a pipette, drained into weighing bottles, and weighed to 0.01 g. To prevent drawing small crystals into the pipette along with the saturated solution a filter should be provided. The filter is removed before the pipette is drained. Alternatively the filtering operation can be carried out with a small piece of filter paper slipped over the bottom of the pipette and fastened with a rubber band. The solution is then titrated with the standardized sodium hydroxide using phenolphthalein as an indicator. Duplicate measurements are to be made at each temperature.

Boric acid can be used as an alternate to oxalic acid. The 10 ml aliquots are titrated using phenolphthalein indicator. For each titration, adding 20 ml glycerin produces a sharper endpoint.

Question: Why is molality or mole fraction used as the concentration unit in the calculations rather than molarity or weight percent?

The desired concentration of NaOH can not be prepared by accurately weighing out NaOH(s) on the analytical balance. Prepare the 0.5 N solution by adding approximately 10 grams of NaOH(s) pellets to 500 ml H2O. Standardize this solution by titrating against weighed samples of benzoic acid using Phenolphthalein as the indicator.

The determination of solubility^1-5 and the calculation of the differential heat of solution at saturation are illustrated in this experiment.

THEORY. One of the simplest cases of equilibrium is that of a saturated solution in contact with excess solute; molecules leave the solid and pass into solution at the same rate at which molecules from the solution are deposited on the solid. The term Id solubility refers to a measure, on some arbitrarily selected scale, of the concentration of the solute in the saturated solution. Here the molal concentration scale will be used, and the solubility then becomes equal to the molality m_s of the solute in the saturated solution.

An equilibrium-constant relation may be written for the equilibrium considered:

(1)

Here a₂ represents the activity of the solute in the saturated solution and a₂* the activity of the pure solid solute. The conventional choice of standard state for the latter is the pure solute itself at the temperature and pressure involved, making a₂* identically equal to unity. The activity a₂ is related to the molality m of the solute by means of the activity coefficient g , a function of T, P, and composition which approaches unity as m approaches zero. Then

K = [a₂]m=m_s = g _sm_s

where the subscript s indicates that the relation applies to the saturated solution. The Z symbol [a₂]_m=ms denotes the value of the activity a₂ for the saturated solution.

The change in K with temperature at constant pressure reflects a change in m_s,

and also the change in g _s, which is affected by both the variations in temperature and concentration of the solution. The van’ t Hoff equation requires that

(3)

where D H° is the standard enthalpy change for the solution process. This quantity should not be confused with any actual experimentally measurable heat of solution; it can be determined indirectly, however.

Taking into account the effects of temperature and concentration on g _s (Ref. 6), there results for constant pressure

(4)

(5)

In this approximation, then, the differential heat of solution at saturation may be calculated at a given temperature T by multiplying by -R the slope at this temperature of the plot of ln m_s versus 1/T.

(6)

The heat of solution with which we are concerned here is the heat absorbed when 1 mole of the solid is dissolved in a solution that is already practically saturated. It differs from the heat of solution at infinite dilution, which is the heat of solution often given in tables, by an amount equivalent to the heat of dilution from saturation to infinite dilution.

CALCULATIONS. The solubility in moles per 1000 g of solvent is calculated at each of the four temperatures and compared with the accepted values. It is interesting to compare these values with the solubilities calculated in moles per liter.

The logarithm of the solubility in moles per 1000 g of solvent is plotted against the reciprocal of the absolute temperature, and a smooth curve is drawn through the four points. If the value of D H were constant, the line would be straight. Tangents are drawn at 25 and 0° and the heat of solution is determined at those two temperatures with the help of Eq. (5).

Practical applications. The solubility of a substance may be calculated at other temperatures when it has been determined at two different temperatures. The results are more accurate when the heat of solution is not affected by temperature or when the temperature range is small.

Suggestions for further work. The solubility of other materials may be determined in a similar way, e.g., benzoic acid or succinic acid or other solids having low solubility and easy methods of analysis. Nonaqueous solvents may be used also.

Boric acid may be used instead of oxalic acid for this experiment. It is titrated by using phenolphthalein as an indicator and adding 10 to 20 ml of neutral glycerin to give a sharp end point.

The solubility-product rule and the effect of other salts on solubility may be illustrated by determinations of the solubility of a slightly soluble salt such as silver bromate in the presence of common ions, ammonium hydroxide, and other salts.

The influence of salts in reducing the solubility of benzoic acid may be determined.⁶ The salting-out constant thus obtained can be used for calculating activity coefficients.

In Eq. (2) it is assumed that the heat of solution is independent of temperature, but this assumption is not often justified. The equation may be made exact by introducing terms for the heat capacity of the solute and solvent and for the solution.

1. J. H. Hildebrand and R. L. Scott, "Solubilities of Nonelectrolytes," Reinhold Publishing Corporation, New York, 1950.

2. A. Seidell, "Solubilities of Inorganic and Metal Organic Compounds," 3d ed. vol. I, D. Van Nostrand Company, Inc., Princeton, N.J., 1940.

3. A. Seidell, "Solubilities of Organic Compounds," 3d ed., vol. II, D. Van Nostrand Company, Inc., Princeton, N.J., 1941.

4. A. Seidell and W. F. Linke, "Supplement to Solubilities of Inorganic and Organic Compounds," 3d ed., vol. III, D. Van Nostrand Company, Inc., Princeton, N.J., 1952.

5. W. J. Mader, R. D. Mold, and M. J. Mold in A. Weissberger (ed.), "Technique of Organic Chemistry," vol. 1, "Physical Methods of Organic Chemistry," 3d ed., pt. 1, chap. 11, Interscience Publishers, Ins., New York, 1959.

6. A. T. Williamson, Trans. Faraday Soc., 40: 421 (1944).