Equilibrium (College Board AP® Chemistry): Exam Questions

A student places a sample of solid ammonium chloride (NH₄Cl) into a sealed glass container and heats it. The following reversible reaction occurs:

NH₄Cl (s) ⇌ NH₃ (g) + HCl (g)

After some time, the system reaches equilibrium.

Explain how the reaction shown is an example of a reversible reaction.

Describe the molecular-level processes occurring in this system at dynamic equilibrium.

The student measures the total pressure in the container over time. Sketch a graph showing how pressure changes as the system approaches equilibrium. Label your axes and indicate the trend observed.

Explain why the pressure changes in this way.

A student investigates the reaction between nitrogen dioxide (NO₂​) and dinitrogen tetroxide (N₂O₄​) in a sealed container at a constant temperature:

2NO₂ (g) ⇌ N₂O₄ (g)

Initially, the container contains only NO₂​ gas at a concentration of 0.50 M. The system is allowed to reach equilibrium.

Describe how the concentrations of NO_2​ and N₂O₄​ change as equilibrium is established.

Sketch a graph showing the rate of the forward and reverse reactions over time as equilibrium is established. Explain the trend observed in the graph.

After equilibrium is reached, additional NO₂​ gas is injected into the container. Predict and explain how the rate of the forward reaction changes immediately after this disturbance.

A sealed container initially contains only NO₂(g), a brown gas. Over time, a colorless gas forms, and the intensity of the brown color decreases. Eventually, the mixture appears pale brown in color.

The overall chemical equation is shown below:

2 NO₂ (g) ⇌ N₂O₄ (g)

A student records the concentrations of both gases in the container at regular intervals. A graph of the concentration of NO₂ (g) and N₂O₄ (g) over time is shown.

Identify one observable feature of the system that supports the conclusion that the reaction is reversible.

Define dynamic equilibrium and explain why the system appears unchanged once equilibrium is reached.

Compare the rates of the forward and reverse reactions before equilibrium is reached and at equilibrium.

Using the graph, estimate the time at which equilibrium is first established. Justify your estimate.

Predict how the concentration vs. time graph would change, if at all, and explain whether the final equilibrium concentrations would be affected if a catalyst is added to the system at t = 0.

A sealed, rigid 2.00 L container holds solid ammonium carbamate, NH₄CO₂NH₂ (s), at 298 K. Over time, a gas-phase equilibrium is established for the endothermic reaction of the solid forming the gaseous products. The total pressure in the container eventually levels off at 1.28 atm and remains constant.

Write the balanced equation for the decomposition of ammonium carbamate into ammonia and carbon dioxide.

Explain why, even though the system appears unchanged after a certain point, gas-phase reactions are still occurring.

Sketch a graph of total pressure (y-axis) vs. time (x-axis) for this system. Label the axes and show the shape of the curve.

On your graph in part (c), indicate the point at which equilibrium is established and explain how you identified it.

In terms of the relative rates of the forward and reverse reactions, explain why the system reaches a constant pressure over time.

Explain how a catalyst would change the time taken to reach equilibrium and how the final equilibrium pressure would be affected.

Suppose the temperature is increased while keeping the container volume constant.

Predict and explain how the pressure would change once a new equilibrium is reached.