⚖️ Equilibrium

1. Chemical Equilibrium

A state in a reversible reaction where the rate of the forward reaction equals the rate of the backward reaction. At equilibrium, the concentrations of reactants and products become constant.

Law of Chemical Equilibrium: For a general reaction aA + bB ⇌ cC + dD

Kc = ([C]^c [D]^d) / ([A]^a [B]^b) Where Kc is the equilibrium constant in terms of concentration.

For gaseous reactions, we use Kp (partial pressures). Relation: Kp = Kc(RT)^Δn.

2. Le Chatelier's Principle

⚖️ Le Chatelier's Principle

If a system in equilibrium is subjected to a change of concentration, temperature, or pressure, the equilibrium shifts in a direction that tends to undo the effect of the change.

Concentration: Adding reactants shifts equilibrium forward.
Pressure: Increasing pressure shifts equilibrium towards fewer gaseous moles.
Temperature: Increasing temperature shifts equilibrium in the direction of the endothermic reaction.
Catalyst: Has no effect on the equilibrium state, only helps achieve it faster.

3. Ionic Equilibrium in Solution

Acids, bases, and salts undergo ionization in water. Strong electrolytes ionize completely. Weak electrolytes ionize partially.

Arrhenius, Brønsted-Lowry, and Lewis Concepts

Arrhenius: Acid gives H⁺ in water. Base gives OH⁻ in water.

Brønsted-Lowry: Acid is a proton (H⁺) donor. Base is a proton acceptor. (Conjugate acid-base pairs).

Lewis: Acid is an electron pair acceptor. Base is an electron pair donor.

4. pH Scale and Buffer Solutions

pH = -log[H⁺]. For water at 298K, Kw = [H⁺][OH⁻] = 10⁻¹⁴. Therefore, pH + pOH = 14.

Buffer Solutions

Solutions which resist change in pH on dilution or with the addition of small amounts of acid or alkali. (E.g., Acetic acid + Sodium acetate).

Henderson-Hasselbalch equation: pH = pKa + log([Salt]/[Acid])

5. Solubility Product (Ksp)

For a sparingly soluble salt AB ⇌ A⁺ + B⁻, the solubility product Ksp = [A⁺][B⁻]. Precipitation occurs if the Ionic Product > Ksp.