Equilibrium Constant of an Esterification Reaction1
Author: J. M. McCormick
Last Update: August 12, 2009
Introduction
A carboxylic acid will react with an alcohol to form an ester and water in a reversible reaction (i. e., the ester will react with water to form a carboxylic acid and an alcohol).2 The forward and reverse processes will proceed until there is no net change in the amount of each chemical species, which is when we say that chemical equilibrium has been attained. Note that even though nothing apparently changes on the macroscopic scale, the reaction is still proceeding on the microscopic scale.
The equilibrium constant, K, for the reaction of ethanol and acetic acid to form ethyl acetate (shown as Eqn. 1) will be studied. This reaction is sufficiently slow that the amount of acid present at equilibrium may be determined by direct titration of the reaction mixture without upsetting the equilibrium to an appreciable extent. Unfortunately, this also means that reaching equilibrium takes a long time. Therefore, aqueous HCl will be added as a catalyst, but even then it will take about a week at room temperature for the reaction to come to equilibrium.
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Procedure
Prepare an approximately 3 M aqueous HCl solution in distilled water. Prepare at least three replicates of the solutions given in Table 1. Note that although you may use pipets to prepare these solutions, you will need to know the mass of each reactant as precisely as possible. Place the appropriate amounts of each reagent into a clean bottle or flask that can be tightly sealed. Store the sealed bottles at a constant 25 °C for at least a week. Prepare additional sets of samples and store these at other temperatures (it is advisable to use temperatures above 25 °C, why?).
| No. | Volume of HCl (mL) | Volume of Ethyl Acetate (mL) | Volume of Water (mL) | Volume of Absolute Ethanol (mL) | Volume of Acetic Acid (mL) |
|---|---|---|---|---|---|
| 1 | 5.00 | 0.00 | 5.00 | 0.00 | 0.00 |
| 2 | 5.00 | 5.00 | 0.00 | 0.00 | 0.00 |
| 3 | 5.00 | 4.00 | 1.00 | 0.00 | 0.00 |
| 4 | 5.00 | 2.00 | 3.00 | 0.00 | 0.00 |
| 5 | 5.00 | 4.00 | 0.00 | 1.00 | 0.00 |
| 6 | 5.00 | 4.00 | 0.00 | 0.00 | 1.00 |
Table 1. Solutions, and their compositions, used in this exercise.
Prepare and standardize a 0.6 M NaOH (or KOH) solution in distilled water. This solution may be prepared ahead of time and stored in a sealed bottle under nitrogen.
Just before titrating bottles equilibrated above room temperature, rapidly cool the bottle to room temperature. This makes it easier for you to handle the bottle during the titration, but should not effect the equilibrium constant measured at the higher temperature. The reaction is slow enough that it cannot react to the change in temperature and is thus kinetically trapped with a set of concentrations that reflect the equilibrium constant at the higher temperature. Titrate each bottle with the standard strong base solution to a phenolphthalein endpoint.
Results and Analysis
From the total amount of acid present at equilibrium and how much was present to begin with, one can determine the amount of each substance present at equilibrium and thus the equilibrium constant (remember to use molalities in the calculation of K). Report the value of K at 95% confidence and perform a propagation of error analysis. If you make any assumptions, be sure to note them in your report. Comment on what important consideration was ignored in this analysis, how we might include it and how its inclusion might change the results.
Use the van’t Hoff equation and the values of K and different temperatures to determine the enthalpy of reaction, ΔrH, at 95% confidence. These data should be presented graphically, if you have a sufficient number of points. Perform a propagation of error analysis and compare your results to the literature value.
- 1. James, A. M. and Prichard, F. E. Practical Physical Chemistry, 3rd Ed.; Longman: Burnt Mill, England, 1974, p. 220-221.
- 2. Solomons, T. W. G. Organic Chemistry, 4th Ed.; Wiley: New York, 1988, p. 837-843.