Dynamic equilibrium is the state reached when the rate of the forward reaction exactly equals the rate of the reverse reaction ().
It is termed 'dynamic' because the reactions have not stopped; rather, the molecules continue to react in both directions at the same speed, resulting in no net change.
At this point, the concentrations or partial pressures of all reactants and products remain constant over time, though they are rarely equal to each other.
Macroscopic properties, such as color intensity, pressure, or temperature, become stable once equilibrium is established.
| Feature | Reversible Reaction | Irreversible Reaction |
|---|---|---|
| Direction | Bidirectional () | Unidirectional () |
| Completion | Never reaches 100% completion | Proceeds until reactants are gone |
| Final State | Dynamic Equilibrium | Static end-point |
| System Type | Requires a closed system for gases | Can occur in open or closed systems |
Dynamic vs. Static: In dynamic equilibrium, individual particles are constantly changing form, whereas in a static state, all motion or change has ceased entirely.
Rate vs. Concentration: Equilibrium is defined by equal rates, not equal concentrations. The ratio of products to reactants at equilibrium is determined by the equilibrium constant ().
Identify Equilibrium on Graphs: Look for the 'plateau' on concentration-time graphs. Equilibrium is established the moment all lines become horizontal, indicating constant concentrations.
Check the System: If a question involves a gas-phase reaction and mentions an 'open beaker,' the reaction is likely not at equilibrium because gaseous products can escape, preventing the reverse reaction.
Rate Equality: Always remember that at equilibrium, the forward and reverse rates are equal. If a graph shows two rates that are constant but different, the system is NOT at equilibrium.
Verify the Arrows: Use the double half-arrow symbol specifically for equilibrium systems; using a single arrow in an equilibrium context is a common technical error that loses marks.
The 'Stopped' Fallacy: A common mistake is assuming the reaction has stopped at equilibrium. In reality, the forward and reverse processes are occurring rapidly; they just cancel each other out.
Equal Concentration Myth: Students often incorrectly believe that at equilibrium. Equilibrium only requires that concentrations are constant, not equal.
Catalyst Misconception: Adding a catalyst increases the rate of both the forward and reverse reactions equally. It helps reach equilibrium faster but does not change the final equilibrium position.
