Determination of the Solubility Product Constant of Calcium Hydroxide
DETERMINATION OF THE SOLUBILITY PRODUCT CONSTANT OF CALCIUM HYDROXIDE
ABSTRACT This experiment aimed to determine the solubility product constant (Ksp) of Ca(OH)2 as well as to evaluate the effects of common and non-common ions on its solubility. Ca(OH)2 solids were dissolved in eight various media: distilled water, 1.0 M KCl, 0.5 M KCl, 0.1 M KCl, 0.05 M KCl, 0.005 M KCl, 0.001 M KCl, and 0.1 M Ca(NO3)2. The concentration of dissociated OH- concentrations was determined by means of titrimetric analysis with 0.1 M HCl. The [OH-] values were then used to calculate the solubility and Ksp of Ca(NO3)2 in each medium. Furthermore, a plot of the ionic strength and the solubility values for the KCl media was made. In conclusion, although there were several errors in the execution of the procedures, the experiment can be considered as successful due to the minimal percent error in the experimental Ksp and solubility of Ca(OH)2 . It is recommended to properly perform the procedures of the experiment, particularly in the preparation of Ca(OH)2 suspensions and the titration of solutions for better results. INTRODUCTION Equilibrium in a system is achieved through many ways. One way is dissolving a slightly soluble solid in water. Since this is an equilibrium process, an equilibrium constant can be assigned. In this case, the Keq is the solubility product constant, Ksp. In general, the solubility product constant is the equilibrium constant used for solid substances dissolved in an aqueous solution.[1] It is the product of the equilibrium concentrations of the dissolved ions in a saturated solution of an ionic salt. The more the solid substance dissolves, the higher the Ksp value of the system will be. Moreover, each concentration is raised to the power of the respective stoichiometric coefficient of the ion in the balanced equation. There is no denominator in the Ksp expression because the concentration of pure solids effectively remains constant. [2] 1 For example, when a
ABSTRACT This experiment aimed to determine the solubility product constant (Ksp) of Ca(OH)2 as well as to evaluate the effects of common and non-common ions on its solubility. Ca(OH)2 solids were dissolved in eight various media: distilled water, 1.0 M KCl, 0.5 M KCl, 0.1 M KCl, 0.05 M KCl, 0.005 M KCl, 0.001 M KCl, and 0.1 M Ca(NO3)2. The concentration of dissociated OH- concentrations was determined by means of titrimetric analysis with 0.1 M HCl. The [OH-] values were then used to calculate the solubility and Ksp of Ca(NO3)2 in each medium. Furthermore, a plot of the ionic strength and the solubility values for the KCl media was made. In conclusion, although there were several errors in the execution of the procedures, the experiment can be considered as successful due to the minimal percent error in the experimental Ksp and solubility of Ca(OH)2 . It is recommended to properly perform the procedures of the experiment, particularly in the preparation of Ca(OH)2 suspensions and the titration of solutions for better results. INTRODUCTION Equilibrium in a system is achieved through many ways. One way is dissolving a slightly soluble solid in water. Since this is an equilibrium process, an equilibrium constant can be assigned. In this case, the Keq is the solubility product constant, Ksp. In general, the solubility product constant is the equilibrium constant used for solid substances dissolved in an aqueous solution.[1] It is the product of the equilibrium concentrations of the dissolved ions in a saturated solution of an ionic salt. The more the solid substance dissolves, the higher the Ksp value of the system will be. Moreover, each concentration is raised to the power of the respective stoichiometric coefficient of the ion in the balanced equation. There is no denominator in the Ksp expression because the concentration of pure solids effectively remains constant. [2] 1 For example, when a
References: [1] Petrucci, R.H.; Herring, F.G.; Madura, J.D.; Bissonnette, C.; General Chemistry: Principles and Modern Applications, 10th ed.; Pearson Canada Inc.: Canada, 2010; [2] Oracle ThinkQuest Education Foundation. Equilibrium: Solubility Equilibrium. 2000. Web. 9 May 2013. . [3] Lower, Stephen. Solubility Equilibria of Salts. 2011. Web. 9 May 2013. . [4] Institute of Chemistry. General Chemistry II Laboratory Manual. Quezon City: University of the Philippines Diliman, 2011. Print. [5] Brown, Theodore L., et al. Chemistry: The Central Science, 12th ed. Glenview, IL: Pearson Prentice Hall, 2012. Electronic. pp 722-726. [6] Whitten, K.W., Davis, R.E., Peck, M.L., Stanley. G.G. Chemistry (8th ed.). Thomson Brooks/Cole. California. 2007. 6