Lab Report: Equilibrium Reactions and Le Châtelier's Principle

Purpose

The purpose of this report is to investigate how reactions in equilibrium respond to changes and to apply Le Châtelier's principle to industrial processes aiming to enhance product yields. This study examines reactions in equilibrium and their responses to various types of changes.

Introduction

Traditionally, chemical reactions are assumed to proceed to completion, meaning the reaction continues in the forward direction until one of the reactants is completely consumed. However, many reactions are not so straightforward and can occur in both the forward and reverse directions simultaneously.

Such reactions are referred to as reversible reactions.

In reversible reactions, a point is reached where the forward reaction rate equals the reverse reaction rate. This point is known as equilibrium. At equilibrium, the concentrations of reactants and products remain constant over time. It's essential to understand that even though the concentrations of reactants and products don't change, the reaction is ongoing. Equilibrium is a dynamic state that persists as long as the reaction conditions remain constant.

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Because the rates of the forward and reverse reactions are equal at equilibrium, there is no net change in the concentrations of reactants and products over time. To an observer, it may seem as though the reaction has stopped, but in reality, the reaction is still proceeding; the rates of both forward and reverse reactions have simply become equal.

Procedure

A. Initial Concentrations Determination

1. Measure 31.5 mL (0.5 mol) of glacial acetic acid and 29.1 mL (0.5 mol) of ethyl alcohol in separate clean, dry 50-mL graduated cylinders.
2. Pour both reactants simultaneously into a round-base flask and mix thoroughly.
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3. Immediately withdraw 1 mL of the reaction mixture using a 1-mL pipette.
4. Place the 1 mL sample in a 125-mL Erlenmeyer beaker containing 30 mL of deionized water.
5. Add 3 drops of phenolphthalein indicator and titrate the sample with the standard 0.25 M NaOH solution.
6. Record the volume of NaOH required.
7. Calculate the initial concentration of acetic acid using the formula:
   [CH₃COOH] = (Volume of NaOH (mL) × Molarity of NaOH) / Volume of acetic acid (mL)
8. Place 2 to 3 boiling chips in the reaction mixture to ensure smooth boiling.
9. Carefully add 20 drops of concentrated sulfuric acid.
10. Reconnect the condenser to the flask and start heating the mixture. Ensure a snug fit of the condenser and that water is flowing through it to prevent vapor escape. Maintain gentle boiling for 1 to 1.5 hours to allow the reaction to reach equilibrium.

B. Blank Determination

11. Prepare a blank solution by adding 20 drops of H₂SO₄ (the same as added to the reaction mixture) to 60.6 mL of deionized water in a 125-mL Erlenmeyer beaker.
12. Mix thoroughly.
13. Pipette 1 mL of this blank solution into a second 125-mL Erlenmeyer beaker.
14. Add 30 mL of deionized water and 3 drops of phenolphthalein indicator.
15. Titrate as before with the 0.25 M NaOH solution.
16. Record the volume of NaOH needed to reach the endpoint.

C. Final Concentrations Determination

When the reaction has reached equilibrium, turn off the heat and allow the reaction mixture to cool to room temperature. Then disconnect the condenser.

17. Pipette 1 mL of the reaction mixture into a 125-mL Erlenmeyer beaker.
18. Add 30 mL of deionized water and 3 drops of phenolphthalein indicator.
19. Titrated as before with the 0.25 M NaOH solution.
20. Record the volume of NaOH needed to reach the endpoint. This volume represents the amount of NaOH solution required to neutralize the acetic acid and sulfuric acid in the reaction mixture. Subtract the volume needed to neutralize the sulfuric acid, determined from the blank, to obtain the volume of NaOH solution required to neutralize the acetic acid.
21. Use the corrected volume and the formula:
    [CH₃COOH] = (Volume of corrected NaOH (mL) × Molarity of NaOH) / Volume of acetic acid (mL)

D. Equilibrium Constant Determination

Since the initial concentrations of acetic acid and ethyl alcohol are equal, and the two reactants have a 1:1 ratio, the equilibrium concentration of ethyl alcohol is also equal to the equilibrium concentration of acetic acid. Given that one mole of ethyl acetate is formed for each mole of acetic acid reacted, the equilibrium concentration of ethyl acetate is equal to the change in acetic acid concentration, which is the initial acetic acid concentration minus the acetic acid equilibrium concentration. Additionally, one mole of water is formed for each mole of ethyl acetate produced, making the water equilibrium concentration equal to the ethyl acetate equilibrium concentration. The equilibrium constant, Keq, for this reaction can be calculated from these equilibrium concentrations.

Data

Results

The equilibrium concentrations of the products (water and ethyl acetate) are calculated as follows:

[CH₃COOH] equilibrium concentration = 2.27 M

[C₂H₅OH] equilibrium concentration = 2.27 M

[CH₃COOC₂H₅] equilibrium concentration = 5.9225 M

[H₂O] equilibrium concentration = 5.9225 M

The equilibrium constant, Keq, is calculated using the equation:

Keq = ([CH₃COOC₂H₅] × [H₂O]) / ([CH₃COOH] × [C₂H₅OH])

Keq = (5.9225²) / (2.27²) = 6.81

Discussion

As mentioned earlier, the value of Kc depends on temperature. The Kc's temperature dependence relies on the change in enthalpy (∆ °) of the reaction. For an endothermic reaction (∆ ° > 0), the value of Kc increases with increasing temperature. In contrast, for an exothermic reaction (∆ ° < 0), the value of Kc decreases with increasing temperature. Summarizing the Kc temperature dependence:

1. If ∆ ° > 0: Kc increases with increasing temperature.
2. If ∆ ° < 0: Kc decreases with increasing temperature.

The temperature dependence of the equilibrium constant, Kc, plays a crucial role in understanding how reactions respond to changes in temperature. In this experiment, we have determined the equilibrium constant (Keq) at a specific temperature. The calculated Keq value of 6.81 indicates that the reaction strongly favors the formation of products, ethyl acetate and water, at the given temperature and conditions.

It's important to note that Keq is a dimensionless quantity, and its value provides insight into the position of equilibrium. When Keq is greater than 1, as in this case, it suggests that the equilibrium lies to the right, meaning that the formation of products is favored. Conversely, if Keq were less than 1, it would indicate that the equilibrium favors the reactants' formation.

The temperature dependence of Keq, as mentioned earlier, is tied to the change in enthalpy (∆ °) of the reaction. Endothermic reactions, where heat is absorbed (∆ ° > 0), exhibit an increase in Keq with rising temperature. This means that raising the temperature in an endothermic reaction favors the formation of products, leading to a higher equilibrium constant. On the other hand, exothermic reactions, which release heat (∆ ° < 0), experience a decrease in Keq with increasing temperature, indicating that higher temperatures shift the equilibrium toward the reactants' side.

In the context of industrial processes, understanding the temperature dependence of equilibrium reactions is crucial for optimizing reaction conditions to maximize product yields. By manipulating temperature, pressure, and other factors, industries can fine-tune reactions to achieve the desired product concentrations efficiently. Le Châtelier's principle provides a valuable framework for predicting and controlling how changes in conditions affect equilibrium positions.

Conclusion

In this experiment, we investigated reactions in equilibrium and applied Le Châtelier's principle to understand how changes in conditions can affect the position of equilibrium. The determination of the equilibrium constant (Keq) allowed us to assess the extent to which the reaction favored the formation of products. The calculated Keq value of 6.81 indicated a strong preference for the products at the specific temperature of the experiment.

We observed that the temperature dependence of Keq is closely related to the change in enthalpy (∆ °) of the reaction. Endothermic reactions, with positive ∆ ° values, exhibit an increase in Keq with rising temperature, while exothermic reactions, with negative ∆ ° values, experience a decrease in Keq with increasing temperature.

Our findings underscore the importance of understanding and manipulating equilibrium reactions in industrial processes. By optimizing reaction conditions, such as temperature and pressure, industries can enhance product yields and improve the efficiency of chemical processes. Le Châtelier's principle serves as a valuable tool in predicting and controlling how changes in conditions impact equilibrium positions, ultimately leading to better process design and resource utilization.

This experiment not only deepened our understanding of chemical equilibria but also highlighted the practical implications of these principles in real-world applications.