Buvi Equilibrium

Equilibrium

Reversible and Irreversible Reactions

  • Reversible Reactions:

    • Defined as reactions in which the complete amount of reactants does not convert wholly into products.

    • Example: Neutralisation reaction between a weak acid and a weak base.

      • Reaction: ⇌ (Weak acid and strong base such as NaOH).

  • Irreversible Reactions:

    • Involve complete conversion of reactants into products.

    • Example: Neutralisation reactions between strong acids and strong bases.

      • Reaction: CH3COOH + NaOH → CH3COONa + H2O

      • Further example: HCl + NaOH → NaCl + H2O

Equilibrium and Its Dynamic Nature

  • Equilibrium:

    • Defined as the state where the concentrations of reactants and products remain constant over time.

    • Law of Mass Action:

      • The rate of a chemical reaction is proportional to the product of the concentrations of the reactants at constant temperature.

    • For a simple reversible reaction:

      • A + B ⇌ C + D

      • At equilibrium: Rate of forward reaction = Rate of backward reaction

      • Equilibrium Constant (K) can be expressed as:

        • K = kf [A]^a [B]^b / kr [C]^c [D]^d

Relation between Kp, Kc, and Kx

  • Δn: Number of moles of gaseous products - Number of moles of gaseous reactants.

  • Relationship between Gibbs Free Energy (ΔG) and K:

    • ΔG = ΔG° + 2.303 RT log Q

    • Also:

      • ΔG° = -RT ln K

      • At a certain temperature, Kp = Kc (RT)^Δn

Le-Chatelier's Principle

  • Defined as:

    • Any change in the factors determining the equilibrium will shift the equilibrium to reduce or counteract the effect of that change.

Applications of Le-Chatelier's Principle

  1. Synthesis of Ammonia (Haber Process):

    • Conditions favoring forward reaction: High pressure, low temperature, excess N2 and H2, removal of NH3.

  2. Formation of Sulfur Trioxide:

    • Exothermic reaction; forward reaction favored under high pressure.

  3. Synthesis of Nitric Oxide:

    • Endothermic; favorable conditions include high temperature and excess reactants.

  4. Dissociation of Phosphorus Pentachloride:

    • High temperature and low pressure favors dissociation.

Applications to Physical Equilibrium

  1. Melting of Ice (Ice-Water System):

    • Example of physical equilibrium with changes in volume and temperature affecting the state.

    • Absorption of heat produces more water.

  2. Melting of Sulfur:

    • Equilibrium shifts with temperature and pressure affecting melting points.

  3. Boiling of Water:

    • More vapor produced at high temperatures; pressure increases boiling point.

  4. Solubility of Salts:

    • Heat absorption affects solubility; different behaviors based on heat development.

Arrhenius Theory of Electrolytic Dissociation

  • Postulates:

    1. Electrolytes dissociate in aqueous solution to form ions.

    2. There exists equilibrium between ions and undissociated molecules.

  • Degree of Ionization (α):

    • Higher dielectric constants lead to better ionization.

    • Degree of ionization increases with temperature.

Common Ion Effect

  • Defined as the suppression of dissociation of weak electrolytes by the presence of a common ion from a strong electrolyte.

Solubility Product

  • In a saturated solution of sparingly soluble electrolyte, two equilibria exist based on the solubility and dissociation.

Relative Strength of Acids and Bases

Weak Acids

  • Relative strength derived from ionization constants at the same concentration.

    • Kb(HA1)/Kb(HA2) = a1/a2

Weak Bases

  • Similar relationship for bases based on their dissociation constants.

pH Scale

  • pH is the negative logarithm of the concentration of H+ ions in a solution.

  • Relationship between pH and pOH:

    • pH + pOH = 14

Buffer Solutions

  • Defined as solutions that resist changes in pH when small amounts of acid or base are added.

  • Acidic buffer calculations can be made using the Henderson-Hasselbalch equation.