Reversible and irreversible reactions

1. Chemical Equilibrium

1. Chemical equilibrium is the state in which the rate of the forward reaction equals the rate of the reverse reaction.
2. It is a dynamic state, meaning that reactions continue to occur, but there is no net change in concentrations of reactants and products.
3. Equilibrium can occur in closed systems where no matter enters or leaves.
4. At equilibrium, the concentrations of reactants and products remain constant but are not necessarily equal.
5. Equilibrium can be achieved in both physical and chemical processes.

2. Dynamic Nature of Equilibrium

1. The term dynamic refers to the fact that reactions are still occurring at the molecular level, even though macroscopic properties remain constant.
2. At equilibrium, the rate of the forward reaction equals the rate of the reverse reaction.
3. The system appears static externally but is continuously active internally.
4. This dynamic nature ensures the concentrations of reactants and products are maintained.
5. Examples of dynamic equilibrium include:
   • Evaporation and condensation in a closed container.
   • Dissolution of a solute and precipitation in a saturated solution.

3. Reversible Reactions

1. Reversible reactions are those that can proceed in both the forward and reverse directions.
2. They are represented by a double arrow (⇌) in the chemical equation.
3. Examples of reversible reactions:
   • H₂ + I₂ ⇌ 2HI
   • N₂ + 3H₂ ⇌ 2NH₃ (Haber process).
4. Reversible reactions tend to achieve equilibrium under suitable conditions.
5. In such reactions, neither reactants nor products are completely consumed.

4. Irreversible Reactions

1. Irreversible reactions proceed in one direction only, from reactants to products.
2. They are represented by a single arrow (→) in the chemical equation.
3. Examples of irreversible reactions:
   • Combustion reactions, e.g., CH₄ + 2O₂ → CO₂ + 2H₂O.
   • Neutralization reactions, e.g., HCl + NaOH → NaCl + H₂O.
4. Irreversible reactions typically proceed to completion, leaving no reactants.
5. They are more common in open systems, where products are removed as they form.

5. Factors Affecting Equilibrium

1. Equilibrium is affected by changes in:
   • Concentration of reactants or products.
   • Temperature.
   • Pressure (for gaseous reactions).
   • Presence of a catalyst.
2. Le Chatelier’s Principle explains how equilibrium shifts to counteract changes in these conditions.

6. Key Points

1. Equilibrium is a dynamic state with no net change in concentrations.
2. Reversible reactions can achieve equilibrium, while irreversible reactions proceed to completion.
3. The rate of the forward reaction equals the rate of the reverse reaction at equilibrium.
4. Reversible reactions are represented by a double arrow (⇌), and irreversible reactions by a single arrow (→).
5. Equilibrium is affected by concentration, temperature, and pressure.
6. Le Chatelier’s Principle predicts the direction of equilibrium shifts.
7. Examples of equilibrium include gas-phase reactions and phase changes.
8. The dynamic nature of equilibrium is crucial for understanding chemical processes.
9. Applications of equilibrium principles include the Haber process and industrial synthesis.
10. Recognizing equilibrium states is essential for solving physical chemistry problems.