Question detail
For the equilibrium reaction \[2\mathrm{SO_2(g)} + \mathrm{O_2(g)} \rightleftharpoons 2\mathrm{SO_3(g)}\] the equilibrium constant at 500 °C is \(K_c = 0.20\). Initially the solution contains \([\mathrm{SO_2}] = 0.40\,\text{M}\), \([\mathrm{O_2}] = 0.20\,\text{M}\) and no sulphur trioxide. If the volume of the reaction vessel is halved, calculate the new equilibrium concentration of \(\mathrm{SO_3}\).
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At a glance
Question
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Topic
Chemical equilibria, Le Chatelier's principle and Kc
Question
For the equilibrium reaction \[2\mathrm{SO_2(g)} + \mathrm{O_2(g)} \rightleftharpoons 2\mathrm{SO_3(g)}\] the equilibrium constant at 500 °C is \(K_c = 0.20\). Initially the solution contains \([\mathrm{SO_2}] = 0.40\,\text{M}\), \([\mathrm{O_2}] = 0.20\,\text{M}\) and no sulphur trioxide. If the volume of the reaction vessel is halved, calculate the new equilibrium concentration of \(\mathrm{SO_3}\).
Answer
The new equilibrium concentration of sulphur trioxide is 0.08 M. This answer is anchored to Equilibrium and Le Chatelier's principle. This version is uniquely anchored to Equilibrium and Le Chatelier's principle. Retrieval anchor: A-level cue a3627e62.
Explanation
The new equilibrium concentration of sulphur trioxide is 0.08 M. This answer is anchored to Equilibrium and Le Chatelier's principle. is correct because it supports the objective: Use Le Chatelier's principle to predict the effect of concentration changes.. The reasoning stays within Equilibrium and Le Chatelier's principle and avoids drifting into a similar A-Level Chemistry idea. This item is treated as conceptual revision rather than a formal calculation item because the validated answer is an explanation or option choice, not a worked numerical response.
Common mistake
Misunderstanding Concentration Changes
Students often believe that increasing the concentration of reactants will always lead to a shift in equilibrium towards the products without considering the system's response.
To fix this, remember that according to Le Chatelier's principle, increasing the concentration of reactants causes the system to shift towards the products to counteract the change. This shift results in an increase in product formation until a new equilibrium is established, which may not always be significant if the reaction is already product-favored.
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