Determination of the Solubility Product Constant of Calcium Hydroxide

The equilibrium constant for the solubility equilibrium between an ionic solid and its ions is called solubility constant [1] , Ksp of the solute. For example, the solubility product is defined by

MxAy(s) ⇋xM(aq)y++ yA(aq)x- (1)
Where M is the metal cation, A is the anion, x and y are the corresponding charges of the ions. The equilibrium expression is
Ksp=[MY+]x[AX-]Y (2)

In the example, MxAy(s) does not appear in the equilibrium constant expression since its activity is 1 and does not appear in the equation.
The expression for the reaction quotient of
the solid MxAy(s) is
Qsp=[MY+]x[AX-]Y(3)

If Qsp<Ksp, there is no precipitate formed and ions are dissociated in the solution. Likewise, if Qsp>Ksp, ions tend to form precipitate and eventually appear in the solution.

The ionic strength is another factor that affects the Ksp value. It depends on the concentration of the ions present in the solution and their charges. The ionic strength is expressed as  μ=12cizi2 (4)

Where μ is the ionic strength of the solute in the solution is, ci is the molar concentration of each ion and zi is the charge of the constituent ions. In the experiment, the solute Ca(OH)2 is taken into consideration. The Ksp of Ca(OH)2 is expressed as Ksp=[Ca2+][OH-]2 (5)

And was alsp determined by calculating the hydroxide ion concentrationfrom solutions saturated with Ca(OH)2. The effect of diverse and  common ions on the solubility was also examined in the experiment.

Methodology

A precipitate was prepared by mixing 10 mL of 1.0 M Ca(NO3)2 and 20 mL of 1.0 M NaOH. The class was divided into five groups with each group assigned to prepare a calcium hydroxide suspension, each with a different media: distilled water, 1.0 M KCl, 0.5 M KCl, 0.1 M KCl, 0.01 M KCl, 0.005 M KCl, and 0.001 M KCl. The 1.0 M KCl was prepared using solid KCL, while the lower concentrations were prepared by diluting the 1.0 M KCl solutions. The assigned medium was then isolated to 100 mL for each group and then Ca(OH)2 was gradually added to the solution until it is saturated

C. Common-Ion Effect:

100 mL of 0.1 M Ca(NO3)2 in a 250 mL beaker was added with Ca(OH)2 solid to the assigned media media until it was saturated. 50 mL of the solution was then filtered and an aliquot of 25 mL of the solution with 3 drops of phenolphthalein, giving the solution a dark pink color, was titrated with 0.1 M HCl until the dark pink color of the phenolphthalein vanished.

Results and Discussion

The hydroxide ion concentrations can be solved using the relationship MH+VH+= MOH-Valiquot (6) where MH+ and MOH- are the concentrations of H+ and OH- ions, respectively.

Table 1: 0.1 M HCl Volumes Used to Neutralize Saturated Ca(OH)2 Solutions in Different Media Medium| HCl Volume (mL)|
Distilled Water| 9.7|
0.001 M KCl| 8.3|
0.005 M KCl| 8.8|
0.010 M KCl| 9.0|
0.050 M KCl| 8.3|
0.100 M KCl| 6.3|
0.500 M KCl| 12.3|
0.100 M Ca2+| 9.7|

Table 1 shows the volume by which hydroxide ions saturated in the different solutions by hydronium ions present in 0.10 M HCl.

Table 2: Ca2+ and OH- Concentrations
Medium| [OH-]| [Ca2+]|
Distilled Water| 0.0388| 0.0194|
0.001 M KCl| 0.0332| 0.0166|
0.005 M KCl| 0.0352| 0.0176|
0.010 M KCl| 0.036| 0.018|
0.050 M KCl| 0.0332| 0.0166|
0.100 M KCl| 0.0252| 0.0126|
0.500 M KCl| 0.0492| 0.0246|
0.100 M Ca2+| 0.0388| 0.0194|

Table 2 shows the calcium and hydroxide ion concentrations obtained at different media. Based from the table, as the concentration of KCl increases, the Ca2+ and OH- concentrations also increase. This is due to the diverse non-common ion effect. The ions of KCl tend to surround the participating ions, thus increasing the solubility and the Ksp of the product.

Decreased solubility of calcium hydroxide due to the common ion Ca2+ present in the calcium nitrate medium increases the calcium ion concentration, thus by Le Chatelier’s principle, shifting the reaction to form more precipitate and decreases the solubility compared with the solubility n the distilled water medium.

Ionic Strength

Figure 1: Scatter plot of Ionic Strength against Solubility

Conclusion and Recommendations
The experimental value of Ksp obtained is 2.92055 x 10-5 at distilled water medium, having a percentage error of 431% from the literature value 5.5 x 10-6 at 25 oC. It was observed that as the ionic strength increases, the solubility increases. The presence of the common ion showed a decrease in solubility product from the experiment.

Different factors may have affected the experimental value of the solubility product. The possible contamination of reagents in the apparatus used in the experiment may have affected the outcome of the value. Since the literature value is set at 25 oC, the experiment should have been done at the same temperature to minimize the possible error. The titration procedure may have affected also the outcome of the experiment since the equivalence point should have a color of mixed bright pink pigment and colorless throughout the titration process.

The researcher recommends the following for the improvement of the experiment or for alternative research related to the experiment: *
Appendix

1. Ebbing, D. D., & Gammon, S. D. (2009). General Chemistry (9th ed.). Boston: Houghton Mifflin Company. 2. Monster, F. (2000–2006 Pearson Education, publishing as Fact Monster.). Chemistry Equilibrium Constants. Retrieved January 20, 2013, from Fact Monster: http://www.factmonster.com/cig/chemistry/equilibrium-constants.html 3.
Brian, L. (2007, February 2). Hows Does a Spectrophotometer Work. Retrieved January 2013, 2012, from NSF Center for Biophotonics Science & Technology: http://cbst.ucdavis.edu/education/courses/spring-2007-ist9/lewfinaldraft.doc

4. “Spectroscopy.” Encyclopedia Britannica. 2009. Encyclopedia Britannica Online. 21 January 2013. <http://www.britannica.com/EBchecked/topic/537853/William-Shakespeare>.

5. “Beer-Lambert Law.” Everything Bio. nd. Everything Bio. 21 January 2013. <http://www.everythingbio.com/glos/definition.php?word=Beer-Lambert+Law>.