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Chapter 12: Problem 97

Consider the following exothermic reaction at equilibrium: $$\mathrm{N}_{2}(g)+3 \mathrm{H}_{2}(g) \rightleftharpoons 2 \mathrm{NH}_{3}(g)$$. Predict how the following changes affect the number of moles of each component of the system after equilibrium is reestablished by completing the table below. Complete the table with the terms increase, decrease, or no change.

Short Answer

Expert verified
When the pressure is increased, the moles of N2 and H2 decrease, and the moles of NH3 increase. When the temperature is increased, the moles of N2 and H2 increase, and the moles of NH3 decrease. When more N2 is added, the moles of N2 and H2 decrease, and the moles of NH3 increase. When more NH3 is added, the moles of N2 and H2 increase, and the moles of NH3 decrease.

Step by step solution

01

Change in the reaction condition

Let's consider the following changes in the reaction conditions: 1. Increasing the pressure. 2. Increasing the temperature. 3. Adding more N2. 4. Adding more NH3. For each change, we will analyze the effect on the moles of N2, H2, and NH3 according to Le Ch芒telier's principle.
02

Increasing Pressure

When we increase the pressure, the reaction will shift in the direction with fewer moles of gas to counteract this change. In this case, the reaction will shift to the right (towards the formation of NH3) because there are 4 moles of gas on the left side (1 N2 + 3 H2) and 2 moles on the right side (2 NH3). The effect of increasing the pressure: N2 and H2 will decrease, and NH3 will increase.
03

Increasing Temperature

As the reaction is exothermic, increasing the temperature will favor the endothermic reaction, which is the reverse reaction in this case. The reaction will shift to the left (towards the formation of N2 and H2) to counteract this change. The effect of increasing the temperature: N2 and H2 will increase, and NH3 will decrease.
04

Adding more N2

When we add more N2, the reaction will shift in the direction that consumes N2 to reestablish equilibrium. In this case, the reaction will shift to the right (towards the formation of NH3). The effect of adding more N2: N2 will decrease, H2 will decrease (as it also gets consumed), and NH3 will increase.
05

Adding more NH3

When we add more NH3, the reaction will shift in the direction that consumes NH3 to reestablish equilibrium. In this case, the reaction will shift to the left (towards the formation of N2 and H2). The effect of adding more NH3: N2 and H2 will increase, and NH3 will decrease. Now let's complete the table based on our predictions. | Change | N2 | H2 | NH3 | |----------|--------------|--------------|-------------| | Pressure | Decrease | Decrease ?>" "?" ?>"> | Temperature | Increase | Increase | Decrease |(NH3) | Add N2 | Decrease | Decrease | Increase | | Add NH3 | Increase | Increase | Decrease |

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Chemical Equilibrium
When a chemical reaction occurs, the reactants are transformed into products. However, some reactions reach a state known as chemical equilibrium, where the forward and reverse reactions occur at the same rate. At this point, the concentrations of reactants and products remain constant.

For the reaction \(\mathrm{N}_{2}(g) + 3 \mathrm{H}_{2}(g) \rightleftharpoons 2 \mathrm{NH}_{3}(g)\), the equilibrium is established between nitrogen (N2), hydrogen (H2), and ammonia (NH3). The amount of these substances doesn't change over time unless external conditions are altered, such as pressure or temperature changes.
Exothermic Reaction
An exothermic reaction is one that releases heat as it progresses. The reaction of nitrogen and hydrogen to form ammonia is a classic example of an exothermic process where energy is given off into the surroundings.

During such a reaction, if the system's temperature is raised, Le Ch芒telier's principle suggests that the equilibrium will shift to absorb this excess heat, favoring the reverse reaction, which is endothermic. Consequently, forming the reactants from the products becomes more favorable, resulting in an increased amount of N2 and H2, and a decrease in NH3.
Equilibrium Shifts
When a system at equilibrium is subjected to a change in conditions, it undergoes what is known as an equilibrium shift. This is the system's way of counteracting the change to re-establish equilibrium. Le Ch芒telier's principle helps predict the direction of this shift. For instance:

  • An increase in pressure will cause the equilibrium to shift towards the side with fewer moles of gas.
  • An increase in temperature for an exothermic reaction drives the equilibrium towards the reactants.
  • Adding more reactants pushes the equilibrium towards the products to use up the added substances.
  • Conversely, adding more products will shift the equilibrium towards the reactants to reduce the concentration of the added products.
Reaction Conditions
The conditions under which a reaction is carried out, such as reaction conditions, can heavily influence the position of the equilibrium. Reaction conditions include temperature, pressure, and the concentration of reactants and products. Each of these factors interact in specific ways:

  • Pressure: By increasing the pressure, equilibrium shifts to the side with fewer gas molecules to minimize pressure.
  • Temperature: An increase for an exothermic reaction favors the endothermic reverse reaction, and vice versa for a decrease.
  • Concentration: Introducing more of a reactant or product will shift the equilibrium to restore balance, typically by consuming the excess substance.

Adjusting these conditions can be a strategic way to maximize desired products or minimize unwanted reactants.

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Most popular questions from this chapter

A sample of solid ammonium chloride was placed in an evacuated container and then heated so that it decomposed to ammonia gas and hydrogen chloride gas. After heating, the total pressure in the container was found to be 4.4 atm. Calculate \(K_{\mathrm{p}}\) at this temperature for the decomposition reaction $$\mathrm{NH}_{4} \mathrm{Cl}(s) \rightleftharpoons \mathrm{NH}_{3}(g)+\mathrm{HCl}(g)$$,

At \(25^{\circ} \mathrm{C}, K_{\mathrm{p}} \approx 1 \times 10^{-31}\) for the reaction $$\mathrm{N}_{2}(g)+\mathrm{O}_{2}(g) \rightleftharpoons 2 \mathrm{NO}(g)$$. a. Calculate the concentration of NO, in molecules/cm \(^{3}\) that can exist in equilibrium in air at \(25^{\circ} \mathrm{C} .\) In air, \(P_{\mathrm{N}_{2}}=0.8 \mathrm{atm}\) and \(P_{\mathrm{O}_{2}}=0.2 \mathrm{atm}\). b. Typical concentrations of NO in relatively pristine environments range from \(10^{8}\) to \(10^{10}\) molecules/cm \(^{3}\). Why is there a discrepancy between these values and your answer to part a?

The synthesis of ammonia gas from nitrogen gas and hydrogen gas represents a classic case in which a knowledge of kinetics and equilibrium was used to make a desired chemical reaction economically feasible. Explain how each of the following conditions helps to maximize the yield of ammonia. a. running the reaction at an elevated temperature b. removing the ammonia from the reaction mixture as it forms c. using a catalyst d. running the reaction at high pressure

A 1.604 -g sample of methane \(\left(\mathrm{CH}_{4}\right)\) gas and 6.400 g oxygen gas are sealed into a 2.50 -L vessel at \(411^{\circ} \mathrm{C}\) and are allowed to reach equilibrium. Methane can react with oxygen to form gaseous carbon dioxide and water vapor, or methane can react with oxygen to form gaseous carbon monoxide and water vapor. At equilibrium, the pressure of oxygen is 0.326 atm, and the pressure of water vapor is 4.45 atm. Calculate the pressures of carbon monoxide and carbon dioxide present at equilibrium.

At a given temperature, \(K=1.3 \times 10^{-2}\) for the reaction \(\mathrm{N}_{2}(g)+3 \mathrm{H}_{2}(g) \rightleftharpoons 2 \mathrm{NH}_{3}(g)\).Calculate values of \(K\) for the following reactions at this temperature. a. \(\frac{1}{2} \mathrm{N}_{2}(g)+\frac{3}{2} \mathrm{H}_{2}(g) \rightleftharpoons \mathrm{NH}_{3}(g)\). b. \(2 \mathrm{NH}_{3}(g) \rightleftharpoons \mathrm{N}_{2}(g)+3 \mathrm{H}_{2}(g)\) c. \(\mathrm{NH}_{3}(g) \rightleftharpoons \frac{1}{2} \mathrm{N}_{2}(g)+\frac{3}{2} \mathrm{H}_{2}(g)\). d. \(2 \mathrm{N}_{2}(g)+6 \mathrm{H}_{2}(g) \rightleftharpoons 4 \mathrm{NH}_{3}(g)\).
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