Equilibrium Explained

Chemical equilibrium occurs in reversible reactions when the rates of forward and reverse reactions equalize, resulting in constant concentrations. For example, in \( \ce{N2 + 3H2 <=> 2NH3} \) (Haber process), ammonia production balances decomposition. This article explores dynamic equilibrium, the equilibrium constant, Le Chatelier’s principle, and practical uses.

Dynamic Equilibrium

At equilibrium, reactants and products form at the same rate, a dynamic process:

\[ \ce{A + B <=> C + D} \]

Rate forward = \( k_f [\ce{A}][\ce{B}] \); rate reverse = \( k_r [\ce{C}][\ce{D}] \). At equilibrium: \( k_f [\ce{A}][\ce{B}] = k_r [\ce{C}][\ce{D}] \). Concentrations stabilize, but reactions continue.

Equilibrium Constant

The equilibrium constant (\( K \)) quantifies the extent of a reaction:

\[ K_c = \frac{[\ce{C}][\ce{D}]}{[\ce{A}][\ce{B}]} \quad (\text{concentration, mol/L}) \]

For gases, \( K_p = \frac{P_{\ce{C}} P_{\ce{D}}}{P_{\ce{A}} P_{\ce{B}}} \) (partial pressures). Example: \( \ce{H2 + I2 <=> 2HI} \), if \( [\ce{H2}] = 0.1 \, \text{M} \), \( [\ce{I2}] = 0.1 \, \text{M} \), \( [\ce{HI}] = 0.8 \, \text{M} \):

\[ K_c = \frac{(0.8)^2}{(0.1)(0.1)} = 64 \]

\( K > 1 \): products favored; \( K < 1 \): reactants favored.

Le Chatelier’s Principle

This principle predicts equilibrium shifts under stress:

  • Concentration: Adding \( \ce{N2} \) in \( \ce{N2 + 3H2 <=> 2NH3} \) shifts right (more \( \ce{NH3} \)).
  • Pressure: Increasing pressure favors fewer moles (right for \( \ce{N2 + 3H2} \)).
  • Temperature: Exothermic reactions (e.g., \( \Delta H < 0 \)) shift left if heated.

Catalysts speed equilibrium but don’t shift it.

Applications

Equilibrium is key in:

  • Industry: Haber process optimizes \( \ce{NH3} \) yield.
  • Biology: Hemoglobin-\( \ce{O2} \) binding in blood.
  • Environment: \( \ce{CO2} \) dissolution in oceans (\( \ce{CO2 + H2O <=> H2CO3} \)).

It balances efficiency and sustainability.