University of tennessee knoxville chem 138 lab report common ion effect equilibrium quotient PDF

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University of tennessee knoxville chem 138 lab report common ion effect equilibrium quotient...

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University of Tennessee Knoxville – CHEM 138 Lab Report – Common Ion Effect, Equilibrium Quotients Abstract When dissolving solids in solution, it is important to consider the common ion effect. The purpose of this experiment was to observe the common ion effect by dissolving KHT, which dissociates into K+ and HT-, in a solution containing KCl, which dissociates into K+ and Cl-. We accomplished this by making nine mixtures, each with a different concentration of KCl, and dissolved KHT in each solution. We then performed three titrations of each solution using a strong base to react with the weak acid HT-, allowing us to determine the solubility of KHT in each solution. As concentrations of KCl increased, the solubility of KHT decreased, but the calculated Ksp values had little correlation with KCl concentration. The average calculated value of Ksp of KHT for all the data was 9.97*10-4, which is 9.4% away from the theoretical Ksp of KHT, which is 1.1*10-3, suggesting that we underestimated the amount of KHT dissolved in the mixtures.

Introduction The purpose of this experiment was to examine how the solubility of a solid is affected by the presence of other dissolved compounds in the solution which share an ion in common with the solid. This is known as the common ion effect. For the solid KHT, the equilibrium quotient for dissolution Ksp = [K+][HT-]. To observe the common ion effect, we made nine solutions, each with the same volume, with different concentrations of KCl and NaCl. We then dissolved 1.0 gram of solid KHT in each solution and performed titrations of each mixture using NaOH solution to react with the weak acid HT- produced by the dissolution of KHT. Knowing

that the NaOH reacts with the HT- in a 1:1 ratio, we were able to calculate the amount of KHT dissolved in the solution using the amount of NaOH it took to reach the equivalence point, which was indicated by phenolphthalein.

Calculations For Mixture 5, Trial 1:

For Mixture 5, Trial 2:

Average solubility for Mixture 5:

Mixture 6, Trial 1:

Mixture 6, Trial 2:

Average solubility for Mixture 6:

Ksp for Mixture 5 KHT:

Ksp for Mixture 6 KHT:

Mixtur Concentration of e KCl 1 0

Solubility 0.0311

2

0.02

0.0229

3 4 5 6

0.03 0.04 0.05 0.06

0.0209 0.0176 0.0154 0.0135

7

0.07

0.0121

8

0.08

0.0108

9

0.1

0.00891

Ksp 0.00096 6 0.00098 3 0.00106 0.00101 0.00100 0.00099 1 0.00099 3 0.00098 0 0.00097 1

The average Ksp across all mixtures is 9.97*10-4.

Graphs

Solubility vs. Concentration of KCl 0.04 0.03

Solubility

0.03 0.02 0.02 0.01 0.01 0

0

0.02

0.04

0.06

Concentration of KCl

0.08

0.1

0.12

Ksp vs. Concentration of KCl 0 0 0

Ksp

0 0 0 0 0 0 0 0

0.02

0.04

0.06

0.08

0.1

0.12

Concentration of KCl

Questions 1. Compared to the theoretical Ksp of KHT of 1.1*10-3, the results from Mixture 5 had an error of 7.3% and the results from Mixture 6 had an error of 15%. Finally, the average Ksp from the pooled data is 9.97*10-4 which is 9.4% off from the theoretical value. The most likely source of error would be failing to stop the drip of the NaOH solution into the mixture exactly when the phenolphthalein changes color, but that would lead to an overestimation of solubility and therefore Ksp, while our results underestimated it. It is also possible that the amounts of KCl and NaCl solutions may have been measured inaccurately, or that the KHT might not have been mixed thoroughly enough to reach equilibrium, especially if some of it got stuck to the bottom of the beaker. That would lead to underestimation of the solubility and Ksp like what we observed in the data. 2. As the concentration of KCl increases, the solubility of KHT decreases, which makes sense because the dissolved KCl also contains K+ ions, and since and it doesn’t

matter where the K+ ions come from, the K+ ions from the KCl would reduce the room for K+ ions from the KHT. On the other hand, there doesn’t seem to be a correlation between Ksp and KCl concentration, and Ksp remains relatively similar for each mixture, even though there is some variation. This makes sense given that Ksp for a substance does not change.

Conclusion Our results showed decreasing solubility as the concentration of KCl increased. Trials 1 and 2 of Mixture 5 led to an average solubility of 0.0156 and an average Ksp of 1.02*10-3 (7.3% error compared to the theoretical value) while Trials 1 and 2 of Mixture 6 led to an average solubility of 0.0134 and an average Ksp of 9.34*10-4 (15% error). The average Ksp for all the data was 9.97*10-4 (9.4% error). We obtained these results by making nine mixtures, each with a different concentration of KCl and NaCl, and adding KHT and stirring it until it reached equilibrium. We then performed three titrations of each mixture using NaOH as the base and phenolphthalein as the indicator and used this to determine how much of the HT-, a weak acid, was present in each solution. Potential sources of error include mistakes in mixing the solutions or in stirring the KHT to ensure that it had a chance to reach equilibrium, and mistakes in stopping the titrations quickly enough after the phenolphthalein indicator changed color....