Question

Several steps are involved in the industrial production of sulfuric acid. One step involves the oxidation of sulfur dioxide gas to form sulfur trioxide gas. A catalyst is used to increase the rate of production of sulfur trioxide gas. In a rigid cylinder with a movable piston, this reaction reaches equilibrium, as represented by the equation below.

2SO2(g) + O2(g) <==> 2SO3(g) + 392 kJ
Explain, in terms of collision theory, why increasing the pressure of the gases in the cylinder increases the rate of the forward reaction.
Determine the amount of heat released by the production of 1.0 mole of SO3(g).
State, in terms of the concentration of SO3(g) , what occurs when more O2(g) is added to the reaction at equilibrium.

Answer 1

Increasing the pressure of the gas increases the intermolecuar forces of attraction between the molecules because the distance between the molecules are being reduced by the action of the piston. 

Answer 2

Increasing the pressure will mean that these particles will collide more frequently to yield the products hence driving the forward reaction.

What is the collision theory?

According to the collsion theory, reaction occurs as a result of collision between reactants in a system. Now we have the reaction; 2SO2(g) + O2(g) <==> 2SO3(g) + 392 kJ.

We can see that there are more particles on the left hand side than on the right hand side. Increasing the pressure will mean that these particles will collide more frequently to yield the products hence driving the forward reaction.

Now;

2 moles of SO3(g) produces 392 kJ of heat

1 mole of SO3(g) produces 1 mole * 392 kJ /2 moles

= 196 kJ

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