Catalysts in industry
3 marks · EasyThe contact process for making sulfuric acid uses vanadium(V) oxide as a catalyst. In one experiment, 100 cm³ of sulfur dioxide and excess oxygen were reacted at 450°C and 2 atm, both with and without the catalyst. With the catalyst, the reaction reached equilibrium in 2 minutes. Without the catalyst, it took 10 minutes.
How does the catalyst affect the rate of the reaction?
Does the catalyst increase the amount of sulfuric acid produced at equilibrium? Give a reason for your answer.
Explain why using a catalyst is economically important in this reaction.
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(a) It increases the rate of reaction/the reaction happens faster.
(b) No, the catalyst does not change the position of equilibrium or the yield of sulfuric acid, just how quickly equilibrium is reached.
(c) Using a catalyst means lower energy costs and faster production, saving money in industry.
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Marking points
B1 (a) · 1 mark · States that the catalyst increases reaction rate/reaction is faster.
B2 (b) · 1 mark · States the catalyst does not increase the amount of sulfuric acid produced, with correct reason (does not change equilibrium position/yield).
B3 (c) · 1 mark · States that using a catalyst is economically important (faster production and/or lower energy costs).
Worked solution
(a) The catalyst speeds up the reaction/reduces the time taken to reach equilibrium.
(b) No, a catalyst does not increase the equilibrium yield of sulfuric acid; it just makes the reaction reach equilibrium faster.
(c) Using a catalyst saves money by lowering the necessary temperature/pressure and making production quicker in industry.