Question

Consider the reaction $$ 4 \mathrm{NH}_{3}(g)+5 \mathrm{O}_{2}(g) \underset{4 \mathrm{NO}(g)+6 \mathrm{H}_{2} \mathrm{O}(g), \Delta H=-904.4 \mathrm{~kJ}}{\rightleftharpoons} $$ Does each of the following increase, decrease, or leave unchanged the yield of NO at equilibrium? (a) increase \(\left[\mathrm{NH}_{3}\right] ;\) (b) increase \(\left[\mathrm{H}_{2} \mathrm{O}\right] ;(\) c \()\) decrease \(\left[\mathrm{O}_{2}\right]\); (d) decrease the volume of the container in which the reaction occurs: (e) add a catalyst: (f) increase temperature.

Short answer

(a) Increasing [NH3] concentration will increase the yield of NO. (b) Increasing [H2O] concentration will decrease the yield of NO. (c) Decreasing [O2] concentration will decrease the yield of NO. (d) Decreasing the volume of the container will decrease the yield of NO. (e) Adding a catalyst will not change the yield of NO. (f) Increasing temperature will decrease the yield of NO.

Step-by-step solution

(a) Increase [NH3] concentration

Increasing the concentration of NH3 will cause the equilibrium to shift to the right in order to reduce the excess NH3. This will result in an increased yield of NO as the system tries to reestablish equilibrium.

(b) Increase [H2O] concentration

Increasing the concentration of H2O will cause the equilibrium to shift to the left as the system tries to consume the excess H2O. This will result in a decreased yield of NO as the system returns to equilibrium.

(c) Decrease [O2] concentration

Decreasing the concentration of O2 will cause the equilibrium to shift to the left as the system tries to replace the reduced O2. As a result, the yield of NO will decrease as the system reestablishes equilibrium.

(d) Decrease the volume of the container

Decreasing the volume of the container will increase the pressure on the system. The equilibrium will therefore shift to the side with a lower number of moles of gas in order to reduce the pressure. In this case, the left side of the reaction has 9 moles of gas (4 moles NH3 and 5 moles O2), while the right side has 10 moles of gas (4 moles NO and 6 moles H2O). The equilibrium will shift to the left, resulting in a decreased yield of NO.

(e) Add a catalyst

Adding a catalyst will only affect the rate at which the reaction reaches equilibrium. It will not change the position of the equilibrium or the yield of NO.

(f) Increase temperature

Since the given reaction is exothermic (ΔH = -904.4 kJ), increasing the temperature will cause the equilibrium to shift to the endothermic side in order to counteract the added heat. In this case, the endothermic side is the left side of the reaction. As a result, the yield of NO will decrease with the increase in temperature.

Key concepts

Le Chatelier's Principle

Le Chatelier's Principle is a fundamental concept in chemistry that describes how a system at equilibrium responds to changes in concentration, temperature, or pressure. When a change is applied to a system in equilibrium, the system will adjust itself to counteract that change and restore equilibrium balance.

For instance, if the concentration of a reactant is increased, the equilibrium will shift towards the products to reduce this effect. This is why in our example, increasing \(\left[\text{NH}_3\right]\) shifts the equilibrium to the right, increasing the yield of NO. Similarly, if the concentration of a product increases, the system shifts to use up this excess product, as seen with \(\left[\text{H}_2\text{O}\right]\).

Changes in volume and pressure also affect equilibrium. The system shifts towards the side with fewer moles of gas when the volume decreases to reduce pressure. In our example, this means shifting left, resulting in less NO.

• Increase in reactant concentration shifts equilibrium toward products.
• Increase in product concentration shifts equilibrium toward reactants.
• Decreasing volume shifts equilibrium toward fewer moles of gas.

Exothermic Reactions

Exothermic reactions release heat into their surroundings, as indicated by a negative enthalpy change (\(\Delta H\)). When the temperature of a system with an exothermic reaction increases, the equilibrium will shift to absorb the added heat, thus moving towards the endothermic direction.

Based on our example reaction with \(\Delta H = -904.4 \, \text{kJ}\), increasing temperature causes the equilibrium to move towards the reactants, reducing the yield of NO. This shift occurs because the reaction tries to counterbalance the temperature change by favoring the direction that absorbs heat.

Exothermic reactions are quite common in everyday chemistry, contributing to various processes such as combustion and metabolism.

• Exothermic reactions release heat (\(\Delta H < 0\)).
• Increasing temperature shifts equilibrium towards endothermic side (consumes heat).
• In an exothermic reaction, higher temperatures generally decrease the yield of products.

Catalysts in Chemistry

A catalyst is a substance that speeds up the rate of a chemical reaction without being consumed in the process. Importantly, while a catalyst can accelerate the achievement of equilibrium, it does not alter the position of the equilibrium itself. This means the concentrations of reactants and products at equilibrium remain unchanged in the presence of a catalyst.

In our reaction example, adding a catalyst allows the system to reach equilibrium faster but does not affect the yield of NO. This happens because catalysts lower the activation energy required for the reaction to proceed, allowing it to proceed more quickly.

• Catalysts speed up reactions without being consumed.
• They do not change equilibrium positions but lessen the time to reach equilibrium.
• They work by lowering activation energy, making processes more efficient.

Catalysts are widely used in industry to enhance efficiency in chemical manufacturing processes, allowing reactions to occur under milder conditions and reducing energy costs.