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 actually performed at a higher temperature of approximately 450°C in industry to enhance the reaction rate. At elevated temperatures, the reaction proceeds faster, allowing for a higher throughput of sulfur trioxide production. Although the conversion percentage is lower at higher temperatures, the rate of reaction is significantly increased, making the process more feasible for industrial use.

Increasing the pressure will shift the equilibrium towards the side with fewer moles of gas, according to Le Chatelier's principle. In this reaction, there are 2 moles of gas on the left (reactants) and 2 moles on the right (products). While the number of moles remains the same, the reaction can still favor the production of sulfur trioxide at higher pressures, due to increased collisions between molecules.

Atmospheric pressure is often used in industry due to the economic and safety considerations associated with high-pressure systems. Operating under high pressures can be costly and requires robust equipment to ensure safety and integrity. Thus, the balance between yield and operational efficiency is often achieved by working at atmospheric pressure.

The balanced chemical equation is:

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

The balanced chemical equation is:

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

The balanced chemical equation is:

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