Getting Started
Many chemical reactions appear to simply "stop" before all the reactants are used up. This is not a failure of the reaction, but rather the establishment of a balanced, two-way process. This chapter explores the concept of chemical equilibrium, a dynamic state where reactions occur in both the forward and reverse directions simultaneously, leading to a stable, unchanging system on the macroscopic scale.
What You Should Be Able to Do
After completing this section, you will be able to:
Explain why the observable properties of a reversible reaction (like color or pressure) become constant when the system reaches equilibrium.
Describe the activity at the molecular level for a system at equilibrium, emphasizing that reactions have not stopped.
Interpret graphs of concentration or partial pressure versus time to determine when a system has reached equilibrium.
Identify common physical and chemical processes that are reversible and can establish an equilibrium state.
Key Concepts & Analysis
We can understand the establishment of equilibrium by examining the system's state as it changes over time.
Dynamics of Reaching Equilibrium
Baseline Condition: The Initial State
A reversible reaction is a chemical reaction that can proceed in both the forward (reactants to products) and reverse (products to reactants) directions. It is typically represented with a double arrow (⇌). Consider the reversible decomposition of dinitrogen tetroxide (a colorless gas) into nitrogen dioxide (a brown gas):
N₂O₄(g) ⇌ 2NO₂(g)
Initially, if we place only N₂O₄ in a sealed container, the system's baseline condition is a concentration of N₂O₄ and zero concentration of NO₂. At this instant (time = 0), only the forward reaction can occur. The rate of this forward reaction is at its maximum because the concentration of the reactant is at its highest. The rate of the reverse reaction is zero because no product exists yet.
The Process: Approaching Equilibrium
As the forward reaction N₂O₄ → 2NO₂ proceeds, the concentration of the reactant (N₂O₄) decreases, while the concentration of the product (NO₂) increases. According to principles of chemical kinetics, a decrease in reactant concentration causes the rate of the forward reaction to slow down.
Simultaneously, as NO₂ molecules begin to populate the container, the reverse reaction 2NO₂ → N₂O₄ can begin. Initially slow, the rate of this reverse reaction speeds up as the concentration of NO₂ builds. The system is now in a state of flux: the forward reaction is slowing down while the reverse reaction is speeding up.
The Resulting Change: The Equilibrium State
Eventually, the system reaches a point where the rate of the forward reaction has decreased enough, and the rate of the reverse reaction has increased enough, that the two rates become equal.
Rateforward = Ratereverse
This is the definition of chemical equilibrium. At this point, N₂O₄ is being consumed by the forward reaction at the exact same rate it is being produced by the reverse reaction. The same is true for NO₂. Because the molecules of each species are being produced and consumed at the same rate, there is no net change in their concentrations.
This results in constant macroscopic properties. For our example, the brown color of the gaseous mixture stops getting darker, and the total pressure inside the container becomes constant. This state is not static but is a dynamic equilibrium—a balance between two opposing processes that continue to occur at equal rates.
Key Models & Representations
Graphs are powerful tools for visualizing the journey to equilibrium. The two most common types plot concentration vs. time and rate vs. time.
Visualizing Equilibrium: Concentration and Rate
| Stage of Reaction | Macroscopic View (Observations) | Microscopic View (Molecular Level) | Graphical Representation (Concentration vs. Time) |
|---|---|---|---|
| Initial (t=0) | Colorless gas (pure N₂O₄). Pressure is at its initial value. | Only N₂O₄ molecules exist. Forward reaction begins. Reverse reaction rate is zero. | [N₂O₄] is at its maximum. [NO₂] is zero. |
| Approaching Equilibrium | Gas mixture turns progressively darker brown. Pressure increases. | N₂O₄ molecules convert to NO₂. NO₂ molecules begin to convert back to N₂O₄. Rateforward > Ratereverse. | [N₂O₄] curve decreases. [NO₂] curve increases. The slopes are changing. |
| At Equilibrium | The brown color of the gas becomes constant. The pressure is constant. | Forward and reverse reactions occur at the same rate. Molecules are constantly interconverting. | The concentration curves for both [N₂O₄] and [NO₂] become horizontal lines (zero slope). |