Consider the chemical equation in equilibrium: [tex]CH_4(g)+H_2 O(g) \Leftrightarrow CO(g)+3 H_2(g)[/tex]
What will happen to the equilibrium of this reaction if the pressure is increased?
A. The equilibrium will shift to the left to favor the reverse reaction.
B. The equilibrium will not be affected by changing the pressure.
C. The equilibrium will not be reestablished after this kind of stress.
Answer (2)
The reactant side ( C H 4 ( g ) + H 2 O ( g ) ) has 2 moles of gas.
The product side ( CO ( g ) + 3 H 2 ( g ) ) has 4 moles of gas.
Increasing pressure favors the side with fewer moles of gas (Le Chatelier's principle).
The equilibrium shifts to the left to favor the reverse reaction. T h ee q u i l ib r i u m w i ll s hi f tt o t h e l e f tt o f a v or t h ere v ersere a c t i o n .
Explanation
Analyzing the Chemical Equation Let's analyze the given chemical equation and determine how the equilibrium will shift when the pressure is increased. The equation is: C H 4 ( g ) + H 2 O ( g ) ⇔ CO ( g ) + 3 H 2 ( g )
Determining Moles of Gas First, we need to determine the number of moles of gas on each side of the equation. On the reactant side (left side), we have 1 mole of C H 4 ( g ) and 1 mole of H 2 O ( g ) , for a total of 2 moles of gas. On the product side (right side), we have 1 mole of CO ( g ) and 3 moles of H 2 ( g ) , for a total of 4 moles of gas.
Applying Le Chatelier's Principle Now, we apply Le Chatelier's principle, which states that if a change of condition is applied to a system in equilibrium, the system will shift in a direction that relieves the stress. In this case, the stress is an increase in pressure. According to Le Chatelier's principle, increasing the pressure will favor the side with fewer moles of gas.
Determining the Shift in Equilibrium Since the reactant side has 2 moles of gas and the product side has 4 moles of gas, increasing the pressure will shift the equilibrium to the left, favoring the reverse reaction. This is because the reverse reaction reduces the number of gas molecules, thus relieving the pressure.
Final Answer Therefore, the equilibrium will shift to the left to favor the reverse reaction.
Examples
Consider a scenario where you're optimizing a chemical reaction in a closed container. Understanding how pressure affects equilibrium is crucial. For instance, in the Haber-Bosch process ( N 2 ( g ) + 3 H 2 ( g ) ⇔ 2 N H 3 ( g ) ), increasing the pressure favors the production of ammonia ( N H 3 ) because the product side has fewer moles of gas. This knowledge allows you to manipulate reaction conditions to maximize yield and efficiency, saving resources and energy in industrial processes. By applying Le Chatelier's principle, you can predict and control the equilibrium shift, leading to better outcomes in chemical synthesis.
Increasing the pressure of the reaction C H 4 ( g ) + H 2 O ( g ) ⇔ CO ( g ) + 3 H 2 ( g ) will cause the equilibrium to shift to the left, favoring the reverse reaction. This is because the left side has fewer moles of gas. Therefore, the correct answer is A: The equilibrium will shift to the left to favor the reverse reaction.
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