Question

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

Short answer

(a) Increasing the concentration of NH3 will increase the yield of NO. (b) Increasing the concentration of H2O will decrease the yield of NO. (c) Decreasing the concentration of O2 will decrease the yield of NO. (d) Decreasing the volume of the container will decrease the yield of NO. (e) Adding a catalyst will have no effect on the yield of NO. (f) Increasing temperature will decrease the yield of NO.

Step-by-step solution

(a) Increasing the concentration of NH3

Increasing the concentration of NH3 will result in a shift in the equilibrium to the right, favoring the formation of products, including NO. Therefore, the yield of NO will increase.

(b) Increasing the concentration of H2O

Increasing the concentration of H2O will result in a shift in the equilibrium to the left, favoring the formation of reactants and consuming some of the products, including NO. Therefore, the yield of NO will decrease.

(c) Decreasing the concentration of O2

Decreasing the concentration of O2 will result in a shift in the equilibrium to the left, favoring the formation of reactants and consuming some of the products, including NO. Therefore, the yield of NO will decrease.

(d) Decreasing the volume of the container

Decreasing the volume of the container will increase pressure, causing the equilibrium to shift towards the side with fewer moles of gas. In this case, there are 9 moles of gas on the left side of the equation and 10 moles of gas on the right side. Hence, the reaction will shift to the left, favoring the formation of reactants, and the yield of NO will decrease.

(e) Adding a catalyst

Adding a catalyst will speed up both the forward and reverse reactions, making the system reach equilibrium more quickly. However, it will have no effect on the position of the equilibrium or the yield of NO.

(f) Increasing temperature

Since the reaction is exothermic (ΔH < 0), increasing the temperature will shift the equilibrium to the left, favoring the formation of reactants, as this will counteract the temperature increase and cool the system down. As a result, the yield of NO will decrease.

Key concepts

Understanding Le Chatelier's Principle

Le Chatelier's Principle is a fundamental concept in chemistry that predicts how a change in conditions affects a chemical equilibrium. It states that if a dynamic equilibrium is disturbed by changing conditions such as concentration, temperature, or pressure, the system will respond to counteract the disturbance and re-establish equilibrium. Let's dive into some of these concepts:
\[ \bullet \] Increasing the concentration of a reactant like NH extsubscript{3} causes the reaction to shift right. This will favor the production of more products, including NO. The principle here is to 'relieve' the stress of added NH extsubscript{3}.
\[ \bullet \] Conversely, adding more water (H extsubscript{2}O) to the products will cause the equilibrium to shift to the left. This will create more reactants to 'consume' the excess product.
\[ \bullet \] Decreasing the concentration of a reactant such as O extsubscript{2} means the system will again move to the left to regenerate O extsubscript{2} at the cost of the product yield.
Understanding this principle helps us predict how a reaction will respond to changes in its conditions, maintaining chemical equilibrium.

Reaction Yield Explained

Reaction Yield refers to the amount of product formed in a chemical reaction. It varies based on equilibrium conditions and how they are manipulated. In the context of the NO-producing reaction given:
\[ \bullet \] Increasing the concentration of NH extsubscript{3} boosts the reaction yield of NO. This is because the reaction shifts towards the products.
\[ \bullet \] Factors like pressure can change the yield too. Decreasing the container volume increases the pressure, making the reaction shift where there are fewer moles of gas. Since our reaction has more moles on the product side, this means the yield will decrease in NO production.
\[ \bullet \] A change in temperature also influences yield. With an exothermic reaction (release of heat), increasing temperature causes more reactant formation, reducing NO yield.
Maximizing yield requires understanding and manipulating these variables carefully.

The Impact of Equilibrium Shift

Equilibrium Shift is about the direction in which an equilibrium moves when changes occur in a reaction system. Every shift has implications for the concentrations of reactants and products.
\[ \bullet \] Pressure changes influence shifts significantly. If pressure is increased by reducing volume, the reaction shifts to the side with fewer gas molecules. Here, reducing volume shifts to the left side, decreasing NO production.
\[ \bullet \] Adding a catalyst affects the speed, not the direction, of equilibrium shifts. A catalyst speeds up the attainment of equilibrium without favoring any side more than the other.
\[ \bullet \] Temperature shifts also have a profound effect. For exothermic reactions, greater heat (higher temperature) prompts the value to shift to the left, diminishing the NO yield.
The key to manipulating reactions effectively is managing these shifts to obtain desired amounts of products.