To determine which change would be favorable to the forward reaction, we need to analyze the effects of the different options on the system at equilibrium, based on Le Chatelier's principle.
The balanced chemical equation is:
\[ \text{CO(g)} + \text{H}_2\text{O(g)} ⇄ \text{CO}_2\text{(g)} + \text{H}_2\text{(g)} + \text{energy} \]
This reaction shows that one mole of gas is on the reactant side (CO) and one mole of gas (H2) and one mole of gas (CO2) on the product side. Therefore, the total number of moles of gas on both sides is:
- Reactants: 1 mole (CO) + 1 mole (H2O) = 2 moles of gas
- Products: 1 mole (CO2) + 1 mole (H2) = 2 moles of gas
Given the options:
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An increase in pressure: This would favor the side with fewer moles of gas, but both sides have the same number of moles (2), so this change does not favor either direction.
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A decrease in pressure: This would favor the side with more moles of gas, which again does not help the forward reaction since the mole count is equal.
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An increase in the concentration of CO: This would shift the equilibrium to the right, favoring the forward reaction.
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An increase in the concentration of CO2: This would shift the equilibrium to the left, favoring the reverse reaction.
Therefore, the change that would be favorable to the forward reaction is:
An increase in the concentration of CO.