Experiment 5

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Determination of the Solubility Product Constant of Calcium Hydroxide Austin Raniel Tan
Institute of Chemistry, College of Science, University of the Philippines, Diliman, Quezon City 1101 Philippines -------------------------------------------------
Christian Tica, Ryan Tabernilla, Michael Siao, Ron Mabunga, Jaime Olivares, and rest of Team Pogi

Abstract
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By measuring the concentration of the hydroxide ion from a solution saturated with Ca(OH)2 titration analysis against HCl, the experiment’s objective is to determine the solubility constant of calcium hydroxide. But, there are few limitations in solubility constant concept, like the Diverse Ion effect and the Common ion effect. Sources of error probably were the assumptions made throughout the experiment and can be explained further by the limitations of the solubility constant. -------------------------------------------------

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Introduction
In a reaction where a slightly soluble ionic solid is dissolved in water, it ionizes into its respective ions. It signifies that equilibrium between the undissolved and dissolved ions in a saturated solution has been established when a precipitate forms.

The solubility of the solid ionic compound Calcium Hydroxide was observed in this experiment. In the equilibrium reaction of Ca(OH)2 Ca(OH)2 (s)⇌Ca(aq)2++ OH(aq)- [1] ,
the solubility product constant, Ksp, is represented by
Ksp=Ca2+OH- [2]
Ksp is related to Gibb’s free energy change of the dissolution process. Given ∆G° = -RT lnKsp [3]
We can literally say that the solubility constant is affected by temperature.
The solubility constant of the system can also be compared to the reaction quotient Qsp like other systems, given by Qsp=[Ca2+]i[OH-]i [4]
When compared with Qsp, the solubility constant can be used as a basis for the formation of a precipitate. If Qsp = Ksp, the solution is saturated, while Qsp < Ksp signifies that the solution is unsaturated and no precipitate forms. The solution is supersaturated and precipitate will form when Qsp > Ksp. [1] The common-ion effect in solubility equilibria was also observed. Based on the Le Chatelier’s principle, an equilibrium mixture responds to an increase in the concentration of one of its reactants by shifting in the direction in which that reactant is consumed. For reactions of ionic compounds, the addition of a common ion will cause the system to shift to where that common ion is consumed.

Another factor that affects the Ksp besides the temperature is the ionic strength of the solution. For a solution with n number of moles, the ionic strength, μ, is represented by μ =12i=1ncizi2 [5]

where ci is the molar concentration of each ion and zi is the charge of each ion.[1]
The objective was to determine the solubility constant of calcium hydroxide by measuring the [OH-] from a solution saturated with solid Ca(OH)2. The solubility of calcium hydroxide in a solution with common ions was also evaluated. In the first part of the experiment, the preparation fo Ca(OH)2(s) was done by mixing 10 mL of 1.0M Ca(NO3)2 and 20 mL of 1.0M NaOH. The precipitate, Ca(OH)2 obtained was then placed into 7 different prepared KCl media to make calcium hydroxide suspensions.

Using 0.1M HCl as titrant, 25 mL aliquot from each media was then titrated in order to determine the concentration of OH-. By stoichiometry, it was assumed that [Ca2+] was half of [OH-] obtained in the titration.

Molar solubility values were then calculated using the formulas [OH-] = MH+VH+ [6]
VOH-

[OH-]2=Ca2+=S [7]

which is equal to the equilibrium concentration of Ca2+ and one half of [OH-] based on the balance equation.

Using molar solubility, Ksp values were calculated using the formula Ksp=4s3.

The solubility values were plotted versus the ionic strengths in...
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