Revision: Ionic Equilibria Chemistry HSC Science (General) 12th Standard Board Exam Maharashtra State Board

It is the fraction of total number of moles of the electrolyte that dissociates into its ions when the equilibrium is attained.

It is denoted by symbol α.

\[\alpha=\frac{\text{Number of moles dissociated}}{\text{Total number of moles}}\]

The degree of dissociation (α) of an electrolyte is defined as a fraction of the total number of moles of the electrolyte that dissociates into its ions when the equilibrium is attained.

Degree of dissociation is defined as the fraction of the total number of moles of solute which undergoes dissociation in the solution.

The substance that conducts electric current when dissociated into positively and negatively charged ions is called an electrolyte.

The materials which indicate the presence of an acid or a base in a solution. These are called Acid-Base Indicators or sometimes simple indicators.

A substance that donates a proton \[\ce{(H+)}\] to another substance is known as an acid.

A pair of an acid and a base differing by a proton is called conjugate acid-base pair.

pH scale is a scale for measuring the hydrogen ion concentration in a solution.

The pOH of a solution can be defined as the negative logarithm to the base 10, of the molar concentration of OH⁻ ions in solution.

pOH = -log₁₀[OH^-]

The pH of a solution is defined as the negative logarithm to the base 10, of the concentration of H⁺ ions in solution in mol dm^–3.

pH is expressed mathematically as

pH = -log₁₀ [H⁺] or pH = -log₁₀ [H₃O⁺]

Hydrolysis of salt is defined as the reaction in which cations or anions or both ions of a salt react with ions of water to produce acidity or alkalinity (or sometimes even neutrality).

Hydrolysis of salt is defined as the reaction in which cations or anions or both ions of a salt react with ions of water to produce acidity or alkalinity (or sometimes even neutrality).

The reaction of an anion or cation of a salt with water that produces an acidic or basic solution is called hydrolysis.

A buffer solution having a pH more than 7 is called a basic buffer. Weak base with its salt of strong acid gives basic buffer.

e.g. NH₄OH + NH₄Cl, C₆H₅NH₂ + C₆H₅NH₃Cl

A solution containing a weak acid and its salts with strong base is called an acidic buffer solution.

A buffer solution is defined as a solution which resists drastic changes in pH when a small amount of strong acid, strong base, or water is added to it.

The solution maintains its pH constant or retains an acidic or basic nature even upon the addition of small amounts of acid or base.

The ability of a buffer solution to resist changes in pH on the addition of acid or base is called buffer action.

A buffer solution of pH less than 7 is called an acidic buffer. Weak acid with its salt of strong base gives acidic buffer.

e.g. CH₃COOH + CH₃COONa; HCN + NaCN

It is defined as the product of molar concentration of its ions in a saturated solution each concentration terms raised to the power equal to the number of ions produced on dissociation of one molecule of an electrolyte.

\[A_{x}B_{y}\rightleftharpoons xA^{y+}+yB^{x-}\]

\[K_{\mathrm{sp}}=[A^{y^{+}}]^{x-}[B^{x^{-}}]^{y}\]

The number of moles of a compound that dissolves to give one litre of saturated solution is called its molar solubility.

\[\text{Molar solubility (mol/L)}=\frac{\text{Solubility in g/L}}{\text{Molar mass in g/mol}}\]

The electrolytes which dissociate to a smaller extent in aqueous solution are called weak electrolytes.

A base is a substance which contains OH group and produces hydroxide ions (OH⁻) in aqueous solution is called Arrhenius base.

A substance that donates a proton (H⁺) to another substance is called Bronsted–Lowry acid.

A substance that accepts a proton (H⁺) from another substance is called Bronsted–Lowry base.

Any species that accepts a share in an electron pair is called Lewis acid.

Any species that donates a share in an electron pair is called Lewis base.

The reaction in which cations or anions or both ions of a salt react with ions of water to produce acidity or alkalinity (or sometimes even neutrality) is called hydrolysis of salt.

A solution containing a weak base and its salt with strong acid is called basic buffer solution.

The equilibrium between ions and unionized molecules in solution is called ionic equilibrium.

The substances which give rise to ions when dissolved in water are called electrolytes.

An acid is a substance which contains hydrogen and gives rise to H⁺ ions in aqueous solution is called Arrhenius acid.

A solution which resists drastic changes in pH when a small amount of strong acid or strong base or water is added to it is called buffer solution.

A solution containing a weak acid and its salt with strong base is called acidic buffer solution.

The substances which do not ionize and exist as molecules in aqueous solutions are called nonelectrolytes.

The electrolytes ionizing completely or almost completely are called strong electrolytes.

\[K_a=\frac{[H^+][A^-]}{[HA]}\]

For weak acid:

\[K_a=\frac{\alpha^2c}{1-\alpha}\]

If α is very small:

\[K_a=\alpha^2c\]

\[\alpha=\sqrt{\frac{K_a}{c}}\]

\[\alpha=\frac{\text{number of moles dissociated}}{\text{total number of moles}}\]

Percent dissociation:

% dissociation = α × 100

\[K_b=\frac{[B^+][OH^-]}{[BOH]}\]

For weak base:

\[K_b=\frac{\alpha^2c}{1-\alpha}\]

If α is small:

\[K_b=\alpha^2c\]

\[\alpha=\sqrt{\frac{K_b}{c}}\]

K_w​ = [H_3​O⁺][OH⁻]

At 298 K:

K_w = 1.0×10⁻¹⁴

pH = −log[H⁺]

pOH = −log⁡[OH⁻]

Relationship:

pH + pOH = 14

According to Arrhenius theory, acids and bases are defined on the basis of ion formation in aqueous solution.

• An acid is a substance which produces H⁺ ions in aqueous solution.
• A base is a substance which produces OH⁻ ions in aqueous solution.
• This theory explains the acidic and basic nature only in aqueous medium.

Example of acid:

\[\mathrm{H}\mathrm{C}\mathrm{l}(\mathrm{a}\mathrm{q})\xrightarrow{water}\mathrm{H}^+(\mathrm{a}\mathrm{q})+\mathrm{C}\mathrm{l}^-(\mathrm{a}\mathrm{q})\]

Example of base:

NaOH(aq) → Na⁺(aq) + OH⁻(aq)

Bronsted and Lowry proposed a more general theory based on proton transfer.

• An acid is a substance that donates a proton (H⁺).
• A base is a substance that accepts a proton (H⁺).
• Acid-base reactions involve transfer of proton from acid to base.

Example reaction:

HCl + NH₃ ⇌ NH₄⁺ + Cl⁻

A pair of acid and base differing by a proton is called a conjugate acid–base pair.

Ostwald expressed the quantitative relationship between concentration and degree of dissociation of weak electrolytes.

It applies only to weak electrolytes.

The degree of dissociation increases on dilution.

For weak acid HA:

HA ⇌ H⁺ + A⁻

The dissociation constant is:

\[K_a=\frac{[H^+][A^-]}{[HA]}\]

For small α:

\[K_a=\alpha^2c\quad\mathrm{and}\quad\alpha=\sqrt{\frac{K_a}{c}}\]

Three Theories Compared:

Classification Based on Extent of Dissociation

Pure water is a weak electrolyte that self-ionises:

Ionic product of water:

At 298 K (pure water): [H₃O⁺] = [OH⁻] = 1.0 × 10^-7 mol/L

pH = negative logarithm of H₃O⁺ ion concentration (mol/L).

• The pH scale (0–14) measures the concentration of H⁺ ions in a solution; values < 7 indicate acids, > 7 indicate bases, and 7 is neutral.
• A universal indicator shows different colours at different pH levels, helping to determine the strength of an acid or base.
• Strong acids/bases produce more H⁺ or OH⁻ ions in solution, while weak acids/bases produce fewer ions at the same concentration.

Ionisation of a weak electrolyte is suppressed when a strong electrolyte with a common ion is added. According to Le Chatelier’s Principle, equilibrium shifts left due to increased concentration of a common ion.

Example 1 (Weak Acid):

• Reaction:
  CH₃COOH ⇌ H⁺ + CH₃COO⁻
• Add: CH₃COONa (gives CH₃COO⁻)
• Effect: Ionisation of CH₃COOH decreases

Example 2 (Weak Base):

• Reaction:
  NH₄OH ⇌ NH₄⁺ + OH⁻
• Add: NH₄Cl (gives NH₄⁺)
• Effect: Ionisation of NH₄OH decreases