The effect of temperature on the solubility product constant, Ksp, of potassium hydrogen tartrate in water was investigated in the temperature range of 285K to 318K at normal atmospheric pressure. It was found that the solubility of potassium hydrogen tartrate decreases with a decrease in temperature and consequently a smaller volume of sodium hydroxide is needed to neutralize it. The molar solubility of potassium hydrogen tartrate was calculated from the volume of sodium hydroxide used. The experimental result for the solubility at 298K was compared to literature data, to verify the reliability of this method.
The experimental value is in agreement with the literature data. The values for the solubility product constant at each temperature were obtained from the respective molar solubilities, and fitted to the Van’t Hoff equation. The corresponding enthalpy change for the dissolution is positive and this is consistent with solubility decreasing with decreasing temperature. Introduction Solubility is the amount of a solute that can dissolve in a given volume of solvent at a given temperature, producing a saturated solution1.
A saturated solution refers to a solution that is so concentrated that no more solute that is added will dissolve in it2. Hence an equilibrium exists between the dissolved ions and the undissolved salt3 and the equilibrium constant is the solubility product constant, Ksp. Potassium hydrogen tartrate (KHT) is sparingly soluble in water, and the equilibrium for its dissolution favours the undissolved salt. The dissolution of potassium hydrogen tartrate in water can be represented by the equation below: KHC4H4O6 (s) – K+ (aq) + HC4H4O6- (aq)
It follows that the concentration of HC4H4O6- when the solution is saturated is the molar solubility of KHC4H4O6. The solubility product constant, Ksp, is then determined from the concentrations of the two dissolved ions when the solution is saturated. In general, for a sparingly soluble salt, AxBy, forming a saturated solution at equilibrium, AxBy (s) – xAy+ (aq) + yBx- (aq) Ksp = [A]x [B]y Thus the exponents x and y in the Ksp calculation are the stoichiometric coefficients of the dissolved ions in the equation representing the reaction.
Solids are excluded from the Ksp calculation because their concentration is invariant. Hence the Ksp of potassium hydrogen tartrate can be expressed as follows: Ksp = [K+][HC4H4O6-] The product [A]x [B]y at any time is called the ionic product. The ionic product is equal to the solubility product constant only when the solution is saturated. If the ionic product exceeds the solubility product constant at any time, it means that the solution is saturated and the excess ions will precipitate out of the solution. The solubility of a solute is affected by temperature changes.
The solute may become more or less soluble, depending on the sign of the enthalpy change of dissolution. Le Chatelier’s principle states that when a system in equilibrium is subjected to a change such as in its temperature or concentration of products or reactants, the equilibrium will shift in the direction which counteracts this change. If the reaction is endothermic, heat is lost by the surroundings to the solution during the dissolution. Hence, if the temperature of the surroundings is raised, the equilibrium will shift in favour of the dissolved ions, in accordance with Le Chatelier’s principle, to remove the excess heat.
Titration is used to determine the solubility of potassium hydrogen tartrate. This is because the hydrogen tartrate ions can be neutralized by a base and hence if the concentration of the base is known, the concentration of hydrogen tartrate ions can be computed using the following equation: C1V1 = C2V2 The volume of base used is necessary for the calculation and can be found with the aid of an indicator. Phenolphthalein is chosen because it has a distinct colour change as more base is added within its pH range. The Ksp of KHC4H4O6 can then be calculated using the value of [HC4H4O6-].
The temperature dependence of the solubility product constant is described by the Van’t Hoff equation: InK= -?HRT+?SR ?H and ?S represent the standard enthalpy and entropy change of the reaction, and can be found from the gradient and y-intercept respectively, if the experimental values for temperature, T, and each corresponding Ksp value are used to plot a graph of InKsp versus 1/T. Experimental Procedure The procedure involved titrating saturated solutions of KHC4H4O6 at different temperatures against a known concentration of NaOH. Dried potassium hydrogen phthalate, KHC8H4O4, was required for the initial standardization of the NaOH. . 5g to the nearest 0. 0001g of dried potassium hydrogen phthalate, KHC8H4O4, was weighed using a top pan balance, and then transferred into a 250mL conical flask. 25 mL of deionized water was added, followed by 2 drops of phenolphthalein indicator. The solution was titrated against NaOH with swirling until the endpoint was reached. The endpoint was when the solution had a permanent light pink colour. The titration was repeated and the volume of NaOH required for each trial was tabulated, and the average volume of NaOH used to calculate the concentration of the NaOH solution.
This standardized NaOH solution was used for the whole experiment. 100mL of deionized water was transferred to a 250mL conical flask and approximately 1. 0-1. 5g of KHC4H4O6 was added, followed by swirling for 5 minutes. The solution was then filtered to remove undissolved KHC4H4O6. A 25. 0mL pipette was used to transfer the filtrate into another conical flask and 2 drops of phenolphthalein were added. The solution was titrated with NaOH. The volume of NaOH required for titration was recorded. For temperatures below room temperature, an ice water bath was prepared and the temperature of the KHC4H4O6 solution was maintained.
The solution was filtered in small portions and the temperature of the solution in the filter funnel was measured throughout. 25. 0mL portions of the filtrate were titrated with NaOH and the volume of NaOH used was recorded. This was done for the temperatures 12°C and 19°C. For temperatures above room temperature, a hot water bath was set up. The solution was filtered and the temperature of the solution in the filter funnel was measured throughout. 25. 0mL portions of the filtrate were titrated with NaOH and the volume of NaOH used was recorded. This was done for the temperatures 36°C and 46°C.