A classic experiment in equilibrium studies dating from 1862 involved the reaction in solution of ethanol and acetic acid to produce ethyl acetate and water.
The reaction can be followed by analyzing the equilibrium mixture for its acetic acid content.
In one experiment, a mixture of 1.000 acetic acid and 0.5000 ethanol is brought to equilibrium. A sample containing exactly one-hundredth of the equilibrium mixture requires 28.80 0.1040 for its titration.
Calculate the equilibrium constant for the ethanol-acetic acid reaction based on this experiment.
See your post above this one.
I would like to know the answer to this question too
The equilibrium constant is 4.0. Do not know how to get that, just got it from the answer key. Thought giving this would be better than nothing.
To calculate the equilibrium constant (K) for the ethanol-acetic acid reaction based on this experiment, we need to use the information given about the reaction, the concentration of reactants, and the volume of titrant required for titration.
Here are the steps to calculate the equilibrium constant:
Step 1: Write the balanced chemical equation for the reaction:
The reaction can be represented as follows:
Ethanol + Acetic acid ⇌ Ethyl acetate + Water
Step 2: Determine the concentrations of acetic acid and ethanol:
From the information given, the initial concentrations of acetic acid and ethanol were 1.000 M and 0.5000 M, respectively. However, we need the concentration at equilibrium. Let's assume the concentration of acetic acid at equilibrium is x M (as some acetic acid will be converted to ethyl acetate).
Step 3: Calculate the change in concentration for acetic acid:
The change in acetic acid concentration from initial to equilibrium is (1.000 - x) M.
Step 4: Determine the concentration of ethyl acetate and water:
Since the reaction reaches equilibrium, the concentration of ethyl acetate and water will be the same as the change in concentration of acetic acid, which is (1.000 - x) M.
Step 5: Determine the volume of titrant required for the titration:
From the information given, a sample containing exactly one-hundredth of the equilibrium mixture requires 28.80 mL of 0.1040 M titrant for its titration.
Step 6: Calculate the moles of acetic acid in the sample:
The moles of acetic acid in the sample can be calculated using the formula:
moles = concentration × volume
So, the moles of acetic acid = (1.000 - x) M × (0.01 × equilibrium volume)
Step 7: Calculate the concentration of acetic acid in the sample:
The concentration of acetic acid in the sample can be calculated using the formula:
concentration = moles / volume of sample
So, the concentration of acetic acid in the sample = moles / (0.01 × equilibrium volume)
Step 8: Use the concentration of acetic acid in the sample to calculate x:
Since the reaction is in equilibrium, the concentration of acetic acid in the sample is equal to the concentration of acetic acid at equilibrium, which is x M.
Step 9: Write the expression for the equilibrium constant (K):
The equilibrium constant expression for the reaction is:
K = ([Ethyl acetate] × [Water]) / ([Acetic acid] × [Ethanol])
Step 10: Substitute the values into the equilibrium constant expression:
Using the concentrations of ethyl acetate and water as (1.000 - x) M, and the initial concentration of ethanol as 0.5000 M, we get:
K = [(1.000 - x) M × (1.000 - x) M] / (x M × 0.5000 M)
Step 11: Simplify the expression:
K = (1.000 - x)^2 / (0.5x)
Step 12: Calculate the value of x:
To solve for x, we can use the information given that a sample containing exactly one-hundredth of the equilibrium mixture requires 28.80 mL of 0.1040 M titrant for its titration. This means that the moles of acetic acid in the sample are equal to the moles of titrant used:
moles of acetic acid = moles of titrant
(1.000 - x) M × (0.01 × equilibrium volume) = 0.1040 M × 0.02880 L
Step 13: Solve for x:
Solving the equation obtained in step 12 for x will give us the value of x.
Step 14: Substitute the value of x into the equilibrium constant expression:
Using the calculated value of x, substitute it into the equilibrium constant expression from step 11 to find the value of K.
Following these steps will help you calculate the equilibrium constant for the ethanol-acetic acid reaction based on the given experiment.