Ionic Equilibrium in Solution: Overview, Definitions, Formulas & Class 11 Notes

Ionic Equilibrium is a state where the rate of ion formation and the reverse process (ion coming back together) is equal. This process takes place in reversible reactions, which involve weak acids, bases, and slightly soluble salts. 

 Examples: CH₃COOH ⇌ CH₃COO⁻ + H⁺ 

Important Links:

Electrolytes: A substance that dissolves in water and produces ions (charged particles) to conduct electricity is called an electrolyte. When we dissolve electrolytes in water, they break apart into ions (positive & negative), to allow the electric current to pass through the solution. 

 Types of Electrolytes: 

1. Strong Electrolytes: They completely ionize in water, such as Potassium hydroxide (KOH), Sodium chloride (NaCl), and Hydrochloric acid (HCl).
2. Weak Electrolytes: The substance partially ionizes in water. For example, Ammonium hydroxide (NH₄OH), Acetic acid (CH₃COOH).

 Non-Electrolytes: A substance that completely dissolves in water but does not produce electricity due to no ion production is called a non-electrolyte. These are neutral molecules in the solution. For example, urea, alcohol, and sugar (glucose).

The fraction of the total molecules of an electrolyte that ionized in a solution is the degree of ionization. It is represented by the Greek letter α (alpha).  

 Degree of Ionization formula: 

 α = Number of molecules ionized / Total number of molecules initially present 

 Factors Affecting Degree of Ionization 

1. Nature of Solute: Strong acids and bases
2. Nature of Solvent: A High dielectric solvent increases ionization.
3. Concentration: Lower concentration
4. Temperature: The Higher the temperature, the higher ionization.
5. Presence of Common Ions: Decreases ionization.

This law explains the relationship between the degree of ionization (α) of a weak electrolyte and its concentration in solution. Ostwald's Dilution Law helps to explain how the ionization of a weak acid or base increases as solutions become more dilute.

Important links: NCERT Solutions for Class 11 Chemistry Equilibrium

Expression of the law:

For a weak electrolyte 
Acetic acid: CH₃COOH ⇌ CH₃COO⁻ + H⁺, if 

C = initial concentration 
α = degree of ionization 

Then the ionization constant Ka (or Kb for a base) is expressed as: 

$K_a=\frac{\left[{\mathrm{CH}}_3{\mathrm{COO}}^{-}\right]\left[H^{+}\right]}{\left[{\mathrm{CH}}_3\mathrm{COOH}\right]}=\frac{\left(C\alpha \right)\left(C\alpha \right)}{C(1-\alpha )}=\frac{C{\alpha }^2}{1-\alpha }$

Limitations of Ostwald’s Dilution Law

1. Applicable for weak electrolytes
2. Presume that the solution follows the law of mass action.
3. Not applicable with higher concentration or very strong electrolytes

The common ion effect is an event in which the addition of a strong electrolyte containing a common ion to a weak electrolyte solution suppresses the ionization (or dissociation ), decreasing the concentration of its ions in solution. This effect is based on Le Chatelier's Principle.
Explanation with Example:
Consider acetic acid (CH₃COOH) in water:
CH₃COOH⇌CH₃COO⁻+H⁺

If we add a strong electrolyte, sodium acetate (CH₃COONa), it dissociates completely into CH₃COO⁻ and Na⁺ ions.

CH₃COO⁻ is a common ion; the concentration increases and shifts the equilibrium to the left, reducing the ionization of acetic acid.

Read More: NCERT Solution | Class 11 Chemistry NCERT Solutions 

Key Points:

• Strong electrolytes form a common ion
• Decreases the ionization of the weak electrolyte.
• Common ions lower the concentration of H⁺ or OH⁻, affecting pH.
• Buffer solutions maintain a constant pH.

A buffer solution is a solution that resists the changes in its pH when a small amount of acid or base is added to it.

Buffer Equation (Henderson-Hasselbalch Equation):

For acidic buffer: $\mathrm{pH}={\mathrm{pK}}_a+\log \left(\frac{\left[\mathrm{Salt}\right]}{\left[\mathrm{Acid}\right]}\right)$

For basic buffer: $\mathrm{pOH}={\mathrm{pK}}_b+\log \left(\frac{\left[\mathrm{Salt}\right]}{\left[\mathrm{Base}\right]}\right)$

Application of Buffer Solutions:

• Biological systems
• Chemical industries
• Pharmaceuticals
• Laboratories to control reaction conditions

The pH is a measure of acidity or basicity (alkalinity). It helps to know the concentration of hydrogen ions (H⁺) in the solution.

$\mathrm{pH}=-\log \left[H^{+}\right]$

• The solution is acidic if [H⁺] is high. Here pH is low.
• The solution is basic if [H⁺] is low. The pH is high.

On the other hand, pOH measures the concentration of hydroxide ions (OH⁻) in a solution.

$\mathrm{pOH}=-\log \left[{\mathrm{OH}}^{-}\right]$

• The solution is basic if [OH⁻] is high. Here, pOH is low.

Relation Between pH and pOH

At 25°C (298 K): 
 pH + pOH = 14

To find pOH = 14−pH