Answer:
- <u>Option 2. </u><u><em>Produce more ammonia.</em></u>
Explanation:
The influence of temperature in equilibrium reactions can be predicted from the heat (enthalpy) information.
This is the chemical reaction:
- 3 H₂ (g) + N₂(g) ⇄ 2 NH₃(g) ∆H = −92.2 kJ
The information about the enthalpy of the reaction, ∆H = − 92.2 kJ, indicates that energy (heat) has been released to the surroundings (the products of the forward reaction have less energy than the reactants), which is defined as an exothermic reaction.
Then, you can rewrite the equaition in the form:
- 3 H₂ (g) + N₂(g) ⇄ 2 NH₃(g) + 92.2 kJ
This is, the heat can be seen as a product of the direct reaction (or a reactant of the reverse reaction).
Now, it is quite straight to apply Le Chatelier's principle:
a) Decreasing temperature is equivalent to extract heat or having less heat on the left side.
b) Then, the equilibrium must shift in a way that this lack of heat is compensated. Then, the reaction will shift to the right to produce more heat.
As conclusion, you can tell that in exothermic reactions, a decrase in temperature will cause the equilibrium to shift to the right.
This shift, of course, means the production of more ammonia.
The other choices are discarded following this brief reasoning:
1. increase the velocity of the gas molecules: the average velocity of the particles increases when the average kinetic energy increases, and the average kinetic energy will decrease if the temperature decreases. So, this statement is false.
3. increase the kinetic energy of the gas molecules: no, the average kinetic energy is proportional to the temperature, then reducing the temperature decreasese the average kinetic energy.
4. produce less ammonia: it was shown that reducing the temperature will produce more ammonia.
5. have no effect: no, it does have effect, as shown.
