Question

In a chemical reaction 4mol of Nitrogen monoxide NO and 2.5mol of Oxygen O2 are mixed together in a container with a volume of 1000ml to produce Nitrogen Dioxide NO2. If equilibrium constant Kc for this reaction is 0.001.

1. Write a balanced chemical equation for the above reaction?

2. Determine the equilibrium concentrations for each substances?

3. If the overall enthalpy of the above reaction is 49kJ/mol. Predict the outcome of the reaction?

4. If the volume of the product NO2 is being increased, where will equilibrium shift to and why will it have to shift that way?

Answer 1

[AI-generated answer — Bot GPT 3.5]

1. The balanced chemical equation for the reaction is:

2NO(g) + O2(g) -> 2NO2(g)

2. Let x be the amount of NO that reacts to form NO2.
Initial concentrations:
[NO] = 4 mol / 1 L = 4 M
[O2] = 2.5 mol / 1 L = 2.5 M
[NO2] = 0 M

At equilibrium:
[NO] = 4 - 2x M
[O2] = 2.5 - x M
[NO2] = 2x M

Substitute into the equilibrium expression:
Kc = [NO2]^2 / [NO]^2[O2] = (2x)^2 / (4 - 2x)^2(2.5 - x) = 0.001

Solving the equation will give x = 0.537 M. Therefore, at equilibrium:
[NO] = 4 - 2(0.537) = 2.926 M
[O2] = 2.5 - 0.537 = 1.963 M
[NO2] = 2(0.537) = 1.074 M

3. Since the overall enthalpy change of the reaction is positive (endothermic), increasing the temperature will favor the forward reaction in order to absorb the excess heat, resulting in an increase in the concentration of NO2.

4. If the volume of NO2 is increased, the equilibrium will shift to the left to reduce the amount of NO2 produced. This is because according to Le Chatelier's principle, the system will respond to the increase in NO2 concentration by favoring the reverse reaction to consume some of the excess NO2, ultimately reaching a new equilibrium.