Equilibrium
Definition and meaning of Equilibrium in chemistry.
Equilibrium is the state of a reversible chemical reaction in which the forward and reverse reactions proceed at equal rates, so the concentrations of reactants and products no longer change with time.
In more detail
Equilibrium is dynamic, not static: molecules continue converting between reactants and products, but because both reactions occur at the same rate, no net change in composition is observed. The position of equilibrium is described quantitatively by the equilibrium constant, K, calculated from the concentrations (or partial pressures) of products divided by reactants, each raised to its stoichiometric coefficient. Le Chatelier's principle predicts how an equilibrium system responds to disturbances such as changes in concentration, pressure, volume, or temperature, shifting to partially counteract the change.
Key facts
| Field | Physical Chemistry |
|---|---|
| Notation | ⇌ (double half-arrow) |
| Governing quantity | Equilibrium constant, K |
| Related principle | Le Chatelier's principle |
In the Haber process, N2(g) + 3H2(g) ⇌ 2NH3(g), placed in a sealed vessel at fixed temperature reaches equilibrium when ammonia is formed at exactly the rate it decomposes back into nitrogen and hydrogen, leaving the concentrations of all three gases constant.
Frequently asked questions
Does equilibrium mean the reaction has stopped?
No. Equilibrium is dynamic, both the forward and reverse reactions continue, but since they occur at equal rates, the measurable concentrations remain constant.
What is the difference between Kc and Kp?
Kc is the equilibrium constant written in terms of molar concentrations, while Kp uses partial pressures of gaseous species; for a reaction they are related by Kp = Kc(RT)^Δn, where Δn is the change in moles of gas.