paraphrase the lab
May 2015
Thermodynamics of the Solubility of KNO3 LAB
Abstract:
This experiment determined the Ksp, and some thermodynamic properties of potassium nitrate based off of its solubility. The results for temperature, Ksp, ΔH, ΔS, and ΔG were 341K-310K, 21.87M (avg), 24.92 kJ/mol (28.59% error), 0.1018 kJ/mol (12.24% error), and -8.253 kJ/mol respectively.
Introduction:
The purposed of the experiment was to determine the Ksp of the reaction and then use it to determine several thermodynamic properties. The reaction is likely to be endothermic making the reaction only spontaneous at a certain temperature.
The Ksp of potassium nitrate is given by the following reaction:
Ksp = [K+][NO3-] (0)
The thermodynamics of a reaction relate in these equations:
ΔG = ΔH - TΔS (1)
ΔG = -RTlnKsp (2)
(3)
The experiment allowed the observing of the moment of first precipitation as the reaction proceeded in reverse. As ions first became solid again, the temperature was recorded and defined the temperature of the given Ksp value of each trial. After measuring the Ksp of the reaction (1), the given temperature and Ksp values were plugged into equation (2) to obtain the ΔG value. Plotting a linear graph of equation (3) allowed determining the linear equation from which the ΔH and ΔS terms to be determined algebraically.
Data:
KNO3 mass: 2.051g KNO3 mol: 0.0203
Table 1. Trial Data
|
|
Trial 1 |
Trial 2 |
Trial 3 |
Trial 4 |
Trial 5 |
|
Volume (mL) |
3.50 |
4.00 |
4.60 |
5.00 |
5.40 |
|
[K+], [NO3-] |
5.79 |
5.07 |
4.41 |
4.06 |
3.75 |
|
Temp (K) |
341.07 |
333.85 |
325.15 |
319.25 |
310.15 |
Graph 1. ln(Ksp) vs Inverse Temperature
Results:
Table 2. Thermodynamic results lit ΔH value: 34.9 lit ΔS value: 0.20
|
|
Trial 1 |
Trial 2 |
Trial 3 |
Trial 4 |
Trial 5 |
|
Error |
|
[KNO3] |
5.79 |
5.07 |
4.41 |
4.06 |
3.75 |
|
NA |
|
Ksp |
33.59 |
25.72 |
19.45 |
16.46 |
14.11 |
|
NA |
|
ΔG (kJ) |
-9.97 |
-9.01 |
-8.02 |
-7.43 |
-6.83 |
|
NA |
|
ΔH (kJ/K) |
NA |
NA |
NA |
NA |
NA |
24.92 |
28.59% |
|
ΔS (kJ/K) |
NA |
NA |
NA |
NA |
NA |
0.108 |
12.24% |
The ΔH and ΔS were 24.92 kJ/K (28.59% error) and 0.108 kJ/K (12.24% error). The Ksp, ΔG, and temperatures are dependent on the trial and are given in Table 2.
Discussion:
The experiment was successful in determining the thermodynamic values as intended, with percent errors being below 30%. As expected, the reaction was endothermic with a positive enthalpy value. Also, the relationship between inverse temperature and the ln(Ksp) produced a linear graph that coincided with equation (3) and allowed to calculate the rest of the thermodynamic values.
One legitimate source of error was determining the precise moment of precipitation of solid. Due to their opacity, the flakes were sometimes difficult to spot and were confused with the occasional bubble that was on the edge of the tube. Allowing the temperature to drop lower than the point of equilibrium varied the T value in equation (1) which in turn altered the point of spontaneity of the reaction making it at a lower temperature than actual.
Because thermodynamic values of reactions are simply the negative values of those in the reverse direction, the ΔHo and ΔSo values for the dissociation of KNO3 are 494.6 kJ/mol and -0.133 kJ/molK.
Questions:
The data confirmed that the reaction is endothermic and spontaneous only at a certain temperature. As the temperature dropped more salt precipitated. The dissociation of KNO3 results in greater disorder as there are more individual ions. This is confirmed by a positive ΔS value and negative ΔG value.
Sample Calculations:
Trial 1.
Ksp = (5.79M)(5.79M) = 33.59
ΔG = -(0.008314 kJ/molK)(341.07K)ln(33.59) = -9.97 kJ/mol
Ln(Ksp) = -2997.5x + 12.244
-ΔH/R = -2997.5
ΔH = 24.92 kJ
ΔS = 0.108 kJ