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 is performed at a higher temperature of approximately 450°C due to the following reasons:

1. Higher temperatures increase the rate of reaction, making the process more efficient.
2. Although the equilibrium yield of product (sulfur trioxide) decreases with temperature according to Le Chatelier's Principle, the increased reaction rate allows for a larger quantity of product to be produced in a reasonable timeframe.
3. At lower temperatures, the reaction could take an excessively long time, making it impractical for industrial processes.

i. High pressures would increase the equilibrium yield of sulfur trioxide in this reaction because the reaction results in a decrease in the number of gas molecules. According to Le Chatelier's Principle, the system will shift towards the side with fewer gas molecules when pressure is increased, favoring the formation of $SO_3$.

ii. Atmospheric pressure is usually used in industry, even though high pressures increase the equilibrium yield of sulfur trioxide, because such high pressures can lead to increased safety risks, equipment costs, and more complex handling requirements. Operating at atmospheric pressure is safer and more cost-effective.

i. Dilute sulfuric acid is added to sodium carbonate solution:

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

ii. Sulfur trioxide gas is bubbled through concentrated sulfuric acid:

$SO_3(g) + H_2SO_4(l) \rightarrow H_2S_2O_7(l)$

iii. A piece of zinc metal is added to 6 M sulfuric acid:

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