One of the steps involved in the industrial preparation of sulfuric acid is the oxidation of sulfur dioxide to sulfur trioxide according to the equation \[ 2SO_2(g) + O_2(g) \rightleftharpoons 2SO_3(g) \] a - VCE - SSCE Chemistry - Question 8 - 2006 - Paper 1

The reaction between sulfur dioxide and oxygen to form sulfur trioxide is an exothermic process. According to Le Chatelier’s principle, increasing the temperature shifts the equilibrium position to favor the reactants. However, higher temperatures are used in industrial processes to increase the reaction rate, overcoming the slower kinetics at lower temperatures. The compromise temperature of around 450°C provides a sufficient rate of reaction while still favoring product formation.

i. High pressures would increase the equilibrium yield of sulfur trioxide in this reaction.

Increasing the pressure favors the side of the equilibrium with fewer moles of gas. In this case, the reaction involves 3 moles of gas on the reactant side (2 moles of ( SO_2 ) and 1 mole of ( O_2 )) and only 2 moles of gas on the product side (2 moles of ( SO_3 )). Thus, increasing the pressure shifts the equilibrium to the right, producing more ( SO_3 ).

ii. Atmospheric pressure is usually used in industry, even though high pressures increase the equilibrium yield of sulfur trioxide in this reaction.

While high pressures can significantly increase yield, there are practical and economic considerations in industrial processes. Maintaining high pressures requires more robust equipment and increases safety risks. Additionally, the cost of operating at high pressures may not be justified by the marginal gains in yield, especially when the reaction can be run efficiently at atmospheric pressure.

[ H_2SO_4(aq) + Na_2CO_3(s) \rightarrow Na_2SO_4(aq) + H_2O(l) + CO_2(g) ]

[ SO_3(g) + H_2SO_4(l) \rightarrow H_2S_2O_7(l) ]

[ Zn(s) + H_2SO_4(aq) \rightarrow ZnSO_4(aq) + H_2(g) ]