Experimental Determination of Calcium Hydroxide's Ksp and Effects of Ionic Strength and Common-Ion Effect on Solubility

The solubility product constant (Ksp) is a fundamental parameter that characterizes the extent of dissolution of a sparingly soluble salt in water. In this experiment, we aimed to determine the Ksp of calcium hydroxide (Ca(OH)2) by conducting a series of titrations and subsequent calculations.

Materials and Methods:

1. Preparation of Calcium Hydroxide Solution:
   • A saturated solution of calcium hydroxide was prepared by adding excess calcium hydroxide to distilled water.
   • The mixture was stirred until no further dissolution occurred, indicating the saturation point.
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2. Titration with Hydrochloric Acid:
   • A standardized hydrochloric acid (HCl) solution was titrated into the calcium hydroxide solution.
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   • The titration was performed until a clear endpoint was observed, indicating the complete neutralization of calcium hydroxide.
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3. Calculation of Solubility Product Constant:
   • The concentration of hydroxide ions (OH-) in the saturated calcium hydroxide solution was determined using the titration results.
   • The Ksp was calculated using the equilibrium expression for the dissociation of calcium hydroxide in water.

Calculations: The equilibrium expression for the dissociation of calcium hydroxide is given by: \text{Ca(OH)_2 (s)} \rightleftharpoons \text{Ca}^{2+} (aq) + 2\text{OH}^- (aq)

Let x be the molar solubility of calcium hydroxide. Then, the concentrations of Ca^{2+} and OH^- ions at equilibrium are x and 2x, respectively.

The expression for the solubility product constant (Ksp) is given by:Ksp=[Ca2+]×[OH−]2

Since [Ca2+]=x and [OH−]=2x, the Ksp expression becomes: Ksp=x×(2x)2

Solving for x and substituting back into the expression, the Ksp can be calculated.

Results: The titration results indicated that the volume of hydrochloric acid required for neutralization was related to the concentration of hydroxide ions in the calcium hydroxide solution. Table 1 presents the raw titration data, including initial and final volumes of the acid, and the volume used in each titration.

Certain systematic studies have demonstrated that, under specific conditions, the solubility product constant (Ksp) of a slightly soluble ionic solid is influenced more by the ionic strength rather than the inherent chemical characteristics of its ions. Ionic strength is quantified as the sum of the product of the concentration (Ci) and the square of the respective charge (Zi) for each ion. Notably, this consideration is applicable only to solutions with an ionic strength of 0.1M or less. Beyond this threshold, the solubility becomes dependent on the ion nature.

This phenomenon arises due to the interplay of attractive and repulsive forces between the ions of the solid and the surrounding medium. The slight charge on the solid's ions prompts the added solution's ions to envelop each solid ion, diminishing the interaction between them. Consequently, the solubility increases when dissolved in a medium with high ionic strength.

The investigated reaction involves the precipitation of calcium hydroxide:

Ca2+(aq)+2OH−⇌Ca(OH)2(s)Ca2+(aq)+2OH−⇌Ca(OH)2​(s)

The corresponding Ksp expression is given by:

Ksp=[Ca2+][OH−]2

Due to the challenge of directly measuring the dissolved ion amounts in the laboratory, a titration method was employed to determine hydroxide ion concentration. The solubility of Ca2+2+ ions is then inferred to be half of the hydroxide ion concentration, allowing the calculation of the solubility product constant.

The experiment revealed that the solubility of calcium hydroxide precipitate decreases when dissolved in a medium containing a common ion, such as calcium ions from calcium nitrate. This "common-ion effect" aligns with Le Chatelier's principle, wherein the system adjusts to counteract the stress imposed until equilibrium is reached.

The investigation also explored the impact of ionic strength on calcium hydroxide solubility by suspending the precipitate in varying concentrations of KNO3 solutions. The results indicated that as the medium's concentration increases, so does the solubility of the precipitate. This can be attributed to the concept of ionic strength, where increasing concentrations lead to more ions surrounding the solid ions, reducing the attraction between them.