Chemistry Notes

Study Guide: Redox Reactions & Equilibrium (Unit 3 and 4 QCE)

Redox Reactions

  • Definition: Redox (reduction-oxidation) reactions involve the transfer of electrons between chemical species. In these reactions, one substance is oxidized (loses electrons) and another is reduced (gains electrons).

  • Key Concepts:

    • Oxidation: Process of losing electrons, results in an increase in oxidation state.

    • Reduction: Process of gaining electrons, results in a decrease in oxidation state.

    • Oxidizing Agent: The substance that gets reduced and causes oxidation in another substance.

    • Reducing Agent: The substance that gets oxidized and causes reduction in another substance.

  • Identifying Redox Reactions:

    • Changes in oxidation states of elements in the reactants and products are indicators of redox processes.

    • Balancing redox reactions often involves half-reactions.

  • Examples:

    • Combustion of Hydrocarbons: e.g., CH4 + 2O2 → CO2 + 2H2O (C is oxidized, O is reduced)

    • Corrosion: e.g., Rusting of iron (Fe + O2 + H2O → Fe2O3•nH2O)

Equilibrium

  • Definition: A state in a reversible reaction where the rates of the forward and reverse reactions are equal, leading to constant concentrations of reactants and products.

  • Key Principles:

    • Dynamic Equilibrium: Although the overall concentration remains constant, individual molecules continuously change through the reaction.

    • Le Châtelier’s Principle: If an external change is made to a system at equilibrium, the system will adjust to counteract the change and restore equilibrium.

  • Factors Affecting Equilibrium:

    • Concentration: Increasing concentration of reactants produces more products, thereby shifting the equilibrium towards products.

    • Temperature: For exothermic reactions, increasing temperature shifts equilibrium towards reactants; for endothermic, it shifts towards products.

    • Pressure: In gaseous reactions, increasing pressure shifts equilibrium towards the side with fewer moles of gas.

  • Equilibrium Constant (K):

    • Expression describing the ratio of the concentrations of products to reactants at equilibrium.

    • K = [Products]/[Reactants] (each raised to the power of their coefficients in the balanced equation).

  • Example Reactions:

    • Ammonium Chloride Equilibrium: NH4Cl(s) ⇌ NH3(g) + HCl(g)

    • Synthesis of Nitrogen Dioxide: N2(g) + 2O2(g) ⇌ 2NO2(g)

Study Tips

  • Practice balancing redox reactions using half-reaction methods.

  • Work through examples of equilibrium shifts by changing concentration, pressure, and temperature.

  • Familiarize yourself with common redox reactions and their applications in real life (battery operation, biochemical processes).