Study of Solubility Equilibrium of Khc4H4O6

Topics: Titration, Solubility, Acid Pages: 10 (2634 words) Published: March 8, 2013
Name: Beatrice Yeo Zi HuiLab Group: B1
Fume Hood Number: B4Matriculation Number: A0102491R
Email Address: 1st February 2013 CM1191
Experiment 1: Study of Solubility Equilibrium

1. Abstract
The aim of this experiment is to determine the relationship between the solubility of potassium hydrogen tartrate (KHC4H4O6) and temperature. Titrate a known concentration of NaOH against a saturated solution of KHC4H4O6 at different temperatures to obtain the concentrations of KHC4H4O6, and hence the solubility product constant of KHC4H4O6 at various temperatures. It was found that the solubility product constant of KHC4H4O6 is higher at higher temperatures, thus it can be concluded that the salt is more soluble when temperature increases. 2. Introduction

The aim of this experiment is to determine how solubility of a salt is affected by temperature differences. In this experiment, KHC4H4O6, a sparingly soluble salt is used. The dissociation reaction is

KHC4H4O6 (s) ⇌ K+ (aq) + HC4H4O6- (aq)

and the solubility product constant, Ksp, expression can be written as

Ksp = [K+][HC4H4O6-]

As KHC4H4O6 dissociates, it gives the same amount of HC4H4O6- and K+ ions, so the Ksp expression may be rewritten as the square of [HC4H4O6-], i.e.

Ksp = [HC4H4O6-]2

The hydrogen tartrate ion, HC4H4O6-, acts as a weak monoprotic acid. The concentration of HC4H4O6- can be determined by titrating it against a standardised strong base.

HC4H4O6-(aq) + OH-(aq) → C4H4O6-2(aq) + H2O(l)

In this experiment, NaOH is used. As the HC4H4O6- and OH- ions will react in a 1:1 molar ratio, the concentration of HC4H4O6- can be obtained from calculating the moles of OH- used. From here, the Ksp values can be obtained by repeating the experiment at various temperatures.

The Ksp and temperature relationship may be shown through the Jacobus Henricus van’t Hoff equation. The equation is derived from the Gibbs free energy definition (1) and the Gibbs-Helmholtz equation (2).

ΔG°reaction = -RT ln Ksp (1)
ΔG°reaction = ΔH°reaction - TΔS°reaction (2)

Equation both equations and divide throughout by -RT to obtain the van’t Hoff equation.

-RT ln Ksp = ΔH°reaction - TΔS°reaction
ln Ksp = -ΔH°reaction RT + ΔS°reactionR

Although H and S are dependent on temperature, in this experiment there is an assumption that the difference between the heat capacities of the reactants and products are negligible. Therefore, ΔH°reaction and ΔS°reaction may be assumed to be independent of temperature and can be considered as constants. From here, a graph of ln Ksp against 1/T with a gradient of -ΔH°reaction RT and an intercept of ΔS°reactionR can be plotted to determine the relationship between Ksp and temperature. The ΔH and ΔS of the reaction may also be deduced.

3. Experimental Procedures
The experiment comes in 2 parts. The first part was the standardization of a NaOH solution and the second part consists of finding the concentration of HC4H4O6- using the NaOH solution at various temperatures through titration. NaOH is a strong base while HC4H4O6- is a weak organic acid. Throughout the experiment, phenolphthalein was used as an indicator. Phenolphthalein turns from colourless to pink at around pH 8.2 to 9.6.

4.1 Standardization of NaOH solution
An analytical balance was used to weigh accurately about 0.5g of dried potassium hydrogen phthalate. The mass was recorded up to 4 decimal places and the potassium hydrogen phthalate was transferred into a 250mL conical flask. A measuring cylinder was used to add 25mL of distilled water into the conical flask. 2 drops of phenolphthalein was added. A burette was rinsed twice with 4-5 mL of NaOH solution before it was filled up and set up on a retort stand. A trial titration was done to estimate the volume of NaOH required for the neutralization reaction. The whole procedure was repeated thrice more and the volume of NaOH used was recorded. The...
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