Question

Consider the following reaction which is at equilibrium at 100°C: A(g) + B(g) ↔ AB (g) ΔH=420 kJ/mol Predict the direction where the reaction would proceed, and explain your reasoning: a. The reaction vessel were pressurized to 5atm with helium b. The temperature were increased to 200°C c. The reaction were coupled with another reaction: AB + D2 → AD + BD d. The volume of the reaction vessel were increased from 2.0L to 5.0L without removing any gas e. A catalyst were added to the reaction vessel

Answer 1

A(g) + B(g) ↔ AB (g) ΔH = +420 kJ/mol = exothermic

Δn = moles of products - moles of reactants = 1- (1+1) = -1

a) Δn = -1 , then forward reaction favors with high pressure.

If the reaction vessel were pressurized to 5 atm with helium, equilibrium shifts to right.

b) Exothermic reactions are favoured by decrease in temperature.

If  temperature were increased to 200°C , equilibrium shifts to left.

c) The reaction were coupled with another reaction: AB + D2 → AD + BD

i.e addition of product AB.

Addition of product shifts equilibrium to left .

d)  Δn = -1 , then forward reaction favors with low volume.

So if volume of the reaction vessel were increased from 2.0L to 5.0L without removing any gas, equilibrium shifts to left.

e) If a catalyst were added to the reaction vessel, no change in the position of equilibrium

Catalyst do not effect the position of equilibrium. It helps to attain the equilibrium quickly.