Study of Solubility Equilibrium
Experiment 1: Study of Solubility Equilibrium
Data Treatment and Analysis
Section 1: Solubility Product Constant Temperature (˚C) | Volume of NaOH used (mL) | | | | Titration 1 | Titration 2 | Average | 28 | 12.7 | 12.8 | 12.75 | 9 | 10.5 | 10.5 | 10.5 | 19 | 11.3 | 11.2 | 11.25 | 40 | 16.2 | 16.2 | 16.2 | 50 | 22.8 | 22.9 | 22.85 |
Table 1: The volume of NaOH used in the titration at various temperatures.
No. of moles of KHC4H4O6 = 1.45 g ÷ 188.177g/mol = 7.71 x 10-3mol
Molarity of KHC4H4O6 before filtration = 7.71 x 10-3mol ÷ 0.1L = 7.7 x 10-2M
No. of moles of NaOH: (0.07415M x 12.75) /1000L = 9.454 x 10-4 mol
∴ No. of moles of HC4H4O6- = No. of moles K+ = 9.454 x 10-4 mol
[HC4H4O6-] = [K+] = (9.454 x 10-4 mol) / 0.025L = 0.03785 mol/L
Ksp = [K+][HC4H4O6-] = (0.03785 ÷ 1mol/L) x (0.03785 ÷ 1 mol/L) = 1.433 x 10-3
1/T = 1/(28+273.15) = 3.6 x 10-3 ln Ksp = ln (1.433 x 10-3 )= -6.548 1/Temperature(1/K) (10-3 ) | ln Ksp | 3.3 | -6.548 | 3.5 | -6.92 | 3.4 | -6.801 | 3.2 | -6.07 | 3.1 | -5.384 | Table 2: The inverse of different temperatures and the ln for their respective Ksp. Graph 1: Graph of ln Ksp against 1/T
Graph 1 is a graph of ln Ksp against 1/T. It is a linear graph with a R2- value of 0.916 which is close to 1 and a standard error of regression of 0.2103 which is relatively close to zero, indicating that the model of the graph chosen is suitable and the fit of the trendline is good.. The negative gradient shows that the ln Ksp is inversely proportional to 1/T. The gradient and intercept from the graph is used to calculate the enthalpy and entropy change of the reaction. As demonstrated in the calculations below the change in enthalpy is -31.62 kJ (negative) and the change in entropy is 51.5kJ (positive). Therefore the reaction is spontaneous and the Gibb’s free energy is less than zero, pushing the reaction to the right until equilibrium is achieved (Appendix)
RT ln Ksp = = Δ H - T Δ S y
Data Treatment and Analysis
Section 1: Solubility Product Constant Temperature (˚C) | Volume of NaOH used (mL) | | | | Titration 1 | Titration 2 | Average | 28 | 12.7 | 12.8 | 12.75 | 9 | 10.5 | 10.5 | 10.5 | 19 | 11.3 | 11.2 | 11.25 | 40 | 16.2 | 16.2 | 16.2 | 50 | 22.8 | 22.9 | 22.85 |
Table 1: The volume of NaOH used in the titration at various temperatures.
No. of moles of KHC4H4O6 = 1.45 g ÷ 188.177g/mol = 7.71 x 10-3mol
Molarity of KHC4H4O6 before filtration = 7.71 x 10-3mol ÷ 0.1L = 7.7 x 10-2M
No. of moles of NaOH: (0.07415M x 12.75) /1000L = 9.454 x 10-4 mol
∴ No. of moles of HC4H4O6- = No. of moles K+ = 9.454 x 10-4 mol
[HC4H4O6-] = [K+] = (9.454 x 10-4 mol) / 0.025L = 0.03785 mol/L
Ksp = [K+][HC4H4O6-] = (0.03785 ÷ 1mol/L) x (0.03785 ÷ 1 mol/L) = 1.433 x 10-3
1/T = 1/(28+273.15) = 3.6 x 10-3 ln Ksp = ln (1.433 x 10-3 )= -6.548 1/Temperature(1/K) (10-3 ) | ln Ksp | 3.3 | -6.548 | 3.5 | -6.92 | 3.4 | -6.801 | 3.2 | -6.07 | 3.1 | -5.384 | Table 2: The inverse of different temperatures and the ln for their respective Ksp. Graph 1: Graph of ln Ksp against 1/T
Graph 1 is a graph of ln Ksp against 1/T. It is a linear graph with a R2- value of 0.916 which is close to 1 and a standard error of regression of 0.2103 which is relatively close to zero, indicating that the model of the graph chosen is suitable and the fit of the trendline is good.. The negative gradient shows that the ln Ksp is inversely proportional to 1/T. The gradient and intercept from the graph is used to calculate the enthalpy and entropy change of the reaction. As demonstrated in the calculations below the change in enthalpy is -31.62 kJ (negative) and the change in entropy is 51.5kJ (positive). Therefore the reaction is spontaneous and the Gibb’s free energy is less than zero, pushing the reaction to the right until equilibrium is achieved (Appendix)
RT ln Ksp = = Δ H - T Δ S y
References: Chemical dymanics (n.d). Retrieved 3rd February,2012 from http://www.shodor.org/unchem/advanced/thermo/#enthalphy Laird, Brian. B (2009), University Chemistry. Acid-Base Equilibria and Solubility, pp. 611-663. New York. McGraw-Hill. Ranken, M.D., Kill, R.C. and Baker, C.J.C. (1997), Food Industries Manual (Rev. ed.). Alcoholic Beverages, pp. 236-272. Great Britain. Chapman & Hall.