Understanding Thermodynamics Of Titrated Borate
Understanding the Thermodynamics of Titrated Borate. Alexis Cervantes and Ak Young. CHEM 114-07 April 30, 2024.
Introduction
The purpose of this experiment was to evaluate thermodynamic values and understand their relationship to equilibrium constants. By measuring the change in enthalpy (H°) and the change in entropy (S°), the equilibrium constant (Ksp) was determined for the dissolution of borax in water. In procedure one of the experiment, the temperature of the borax solution was varied and its saturation level was measured at different temperatures (10°, 30°, 40°, 50°). In procedure two, the concentration of the borate anion was measured through a titration with 1.00 M HCl. Equation 6: Ksp= [Na+]2 [[B4O5(OH)4]2-]= (s)(2s)2= …show more content…
Sample Calculation for S:
y-intercept (b) = SR S = Ry-intercept (b) S = 8.3145 (29.379) S = 244.257 J/molK.
Discussion
Borax Dissolution Reaction: What is the difference?
The dissolution of borax in water is an endothermic reaction. In Table 2. The positive values of enthalpy, H°, across all groups indicates the reaction has gained more energy than what was released.
The value and size of S° obtained from the experiment seems to align well with the expected increase upon the dissolution of borax. In reference to the General Chemistry Textbook, the value 244.257 J/molK, looks in agreement with similar values of entropy displayed.
Spontaneity Prediction: What is the difference between
According to Gibbs free energy equation (G° = H° -TS°), a reaction will become nonspontaneous when G° becomes positive. In order to predict (calculate) the class average temperature, we can rearrange the equation. T = H°S°.
Calculation Plug in the values of the class average from Table 2. T = 65.750 kJ/mol196.735 J/molK T = 0.334 kJ/K Convert kJ to J T 334K or 61°C Above this temperature, the dissolution of borax in water will no longer be spontaneous.
Theoretical H° and …show more content…
difference from theory = Avgwhole class- theorytheory100 difference from theory= (65.750 kJ/mol)- (78.179 kJ/mol)78.179 kJ/mol100 difference from theory= -15.89 %
Sample calculation for theoretical S°.
difference from theory = Avgwhole class- theorytheory100 difference from theory= (196.735 J/molK)-(244.257 J/molK)(244.257 J/molK) difference from theory= -19.46%
This indicates that the experimental H° is approximately 15.89% lower than the theoretical value. As for the experimental S°, it is approximately 19.46% lower than the theoretical value.
Error Analysis
The negative percent difference for both H° and S° suggests a systematic error between the experimental and theoretical values. This could be attributed to various factors such as incomplete dissolution of borax, inaccuracies in temperature measurements, or experimental limitations. One possible source of error is variations in the temperature readings during data collection. Fluctuations in the temperature of the solution could lead to inaccuracies in determining the equilibrium constants, affecting the slope and intercept of the
Introduction
The purpose of this experiment was to evaluate thermodynamic values and understand their relationship to equilibrium constants. By measuring the change in enthalpy (H°) and the change in entropy (S°), the equilibrium constant (Ksp) was determined for the dissolution of borax in water. In procedure one of the experiment, the temperature of the borax solution was varied and its saturation level was measured at different temperatures (10°, 30°, 40°, 50°). In procedure two, the concentration of the borate anion was measured through a titration with 1.00 M HCl. Equation 6: Ksp= [Na+]2 [[B4O5(OH)4]2-]= (s)(2s)2= …show more content…
Sample Calculation for S:
y-intercept (b) = SR S = Ry-intercept (b) S = 8.3145 (29.379) S = 244.257 J/molK.
Discussion
Borax Dissolution Reaction: What is the difference?
The dissolution of borax in water is an endothermic reaction. In Table 2. The positive values of enthalpy, H°, across all groups indicates the reaction has gained more energy than what was released.
The value and size of S° obtained from the experiment seems to align well with the expected increase upon the dissolution of borax. In reference to the General Chemistry Textbook, the value 244.257 J/molK, looks in agreement with similar values of entropy displayed.
Spontaneity Prediction: What is the difference between
According to Gibbs free energy equation (G° = H° -TS°), a reaction will become nonspontaneous when G° becomes positive. In order to predict (calculate) the class average temperature, we can rearrange the equation. T = H°S°.
Calculation Plug in the values of the class average from Table 2. T = 65.750 kJ/mol196.735 J/molK T = 0.334 kJ/K Convert kJ to J T 334K or 61°C Above this temperature, the dissolution of borax in water will no longer be spontaneous.
Theoretical H° and …show more content…
difference from theory = Avgwhole class- theorytheory100 difference from theory= (65.750 kJ/mol)- (78.179 kJ/mol)78.179 kJ/mol100 difference from theory= -15.89 %
Sample calculation for theoretical S°.
difference from theory = Avgwhole class- theorytheory100 difference from theory= (196.735 J/molK)-(244.257 J/molK)(244.257 J/molK) difference from theory= -19.46%
This indicates that the experimental H° is approximately 15.89% lower than the theoretical value. As for the experimental S°, it is approximately 19.46% lower than the theoretical value.
Error Analysis
The negative percent difference for both H° and S° suggests a systematic error between the experimental and theoretical values. This could be attributed to various factors such as incomplete dissolution of borax, inaccuracies in temperature measurements, or experimental limitations. One possible source of error is variations in the temperature readings during data collection. Fluctuations in the temperature of the solution could lead to inaccuracies in determining the equilibrium constants, affecting the slope and intercept of the