Ionic Equilibrium lonization and Dissociation

Ionic equilibrium is the balance between ions and molecules in a solution. It happens when the rate of ion formation equals the rate at which ions recombine to form molecules. This state is dynamic, meaning ions keep forming and recombining, but the overall concentration of ions and molecules remains constant.

What is Ionization?

Ionization is the process where neutral molecules turn into charged ions. For example, acids, bases, and salts ionize in water.

1. Strong Electrolytes: They ionize completely in water. Examples:
   • Hydrochloric acid (HCl):
     $HCl → H⁺ + Cl⁻$
2. Weak Electrolytes: They only ionize partially. Examples:
   • Acetic acid (CH₃COOH):
     $CH₃COOH ⇌ CH₃COO⁻ + H⁺$

What is Dissociation?

Dissociation is when an ionic compound splits into its ions in water. For example:

• Sodium chloride (NaCl):
  $NaCl → Na⁺ + Cl⁻$

This is slightly different from ionization because the ions already exist in the compound; they just separate.

Factors Affecting Ionization and Dissociation

1. Type of electrolyte: Strong electrolytes ionize completely, weak ones do not.
2. Solution concentration: Higher concentrations can reduce ionization (common ion effect).
3. Temperature: Higher temperatures usually increase ionization.
4. Solvent nature: Water is good at ionizing substances because it stabilizes ions.
5. Common ion presence: If an ion is already present, it suppresses further ionization.

Important Constants

1. Acid Dissociation Constant (Ka):
   Measures the strength of an acid. For example:
   $HA ⇌ H⁺ + A⁻$
   $K_a = \frac{[H⁺][A⁻]}{[HA]}$
   Higher Ka means stronger acid.
2. Base Dissociation Constant (Kb):
   Measures the strength of a base:
   $BOH ⇌ B⁺ + OH⁻$
   $K_b = \frac{[B⁺][OH⁻]}{[BOH]}$
3. Water Ionization Constant (Kw):
   Water ionizes slightly:
   $H₂O ⇌ H⁺ + OH⁻$
   $K_w = [H⁺][OH⁻] = 1 × 10⁻¹⁴$

pH and pOH

pH is a measure of hydrogen ion concentration:
$pH = -\log[H⁺]$

• Acidic: pH < 7
• Neutral: pH = 7
• Basic: pH > 7

Similarly, pOH is:
$pOH = -\log[OH⁻]$
$pH + pOH = 14$

Buffer Solutions

Buffers resist changes in pH when small amounts of acid or base are added. They are made by combining:

1. A weak acid with its salt (e.g., acetic acid + sodium acetate).
2. A weak base with its salt (e.g., ammonia + ammonium chloride).

Applications

1. Industrial Processes: Used in manufacturing chemicals, detergents, and drugs.
2. Biological Systems: Helps maintain pH balance in the body.
3. Water Treatment: Removes impurities.
4. Titration and Analysis: Helps measure acidity and alkalinity.