Definitions [21]
The equilibrium which is established between the unionised molecules and the ions present in a solution of weak electrolytes is called ionic equilibrium.
Define the degree of dissociation.
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.
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.
\[\alpha=\frac{\text{Number of moles dissociated}}{\text{Total number of moles}}\]
Define conjugate acid-base pair.
A pair of an acid and a base differing by a proton is called conjugate acid-base pair.
Define acids according to Bronsted-Lowry theory.
A substance that donates a proton \[\ce{(H+)}\] to another substance is known as an acid.
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.
Define pOH.
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-]
pH scale is a scale for measuring the hydrogen ion concentration in a solution.
Define pH.
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+]
Define Hydrolysis of salt.
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.
Define hydrolysis.
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).
Define Acidic buffer solution.
A solution containing a weak acid and its salts with strong base is called an acidic buffer solution.
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
Define 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
The number of moles of a compound that dissolves to give one litre of saturated solution is called its molar solubility.
\[\text{Molar solubility (mol/L)}=\frac{\text{Solubility in g/L}}{\text{Molar mass in g/mol}}\]
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.
\[A_{x}B_{y}\rightleftharpoons xA^{y+}+yB^{x-}\]
\[K_{\mathrm{sp}}=[A^{y^{+}}]^{x-}[B^{x^{-}}]^{y}\]
Theorems and Laws [1]
The degree of dissociation of a weak electrolyte is inversely proportional to the square root of its concentration.
| Feature | Weak Acid | Weak Base |
|---|---|---|
| Dissociation | HA ⇌ H⁺ + A⁻ | BOH ⇌ B⁺ + OH⁻ |
| Constant | Acid dissociation constant (Ka) | Base dissociation constant (Kb) |
| Expression | \[K_{a}=\frac{[H^{+}][A^{-}]}{[HA]}\] | \[K_{b}=\frac{[B^{+}][OH^{-}]}{[BOH]}\] |
| Ostwald’s Dilution Law | \[\alpha=\sqrt{\frac{K_{a}}{C}}\] or \[\alpha=\sqrt{K_{a}\cdot V}\] | \[\alpha=\sqrt{\frac{K_{b}}{C}}\] or \[\alpha=\sqrt{K_{b}\cdot V}\] |
Key Points
| Type | Nature | Extent of Ionisation | Examples |
|---|---|---|---|
| Strong electrolytes | Ionise completely or almost completely in solution | Near 100% | HCl, H₂SO₄, NaOH, KOH, NaCl, KCl |
| Weak electrolytes | Ionise partially; poor conductors | Small extent | CH₃COOH, H₃BO₃, NH₄OH, HCN |
Three Theories Compared:
| Theory | Acid | Base |
|---|---|---|
| Arrhenius | Contains H; produces H⁺ ions in aqueous solution | Contains OH group; produces OH⁻ ions in aqueous solution |
| Bronsted–Lowry | Proton donor (H⁺) | Proton acceptor |
| Lewis | Accepts a share in an electron pair | Donates a share in an electron pair |
All Bronsted bases are Lewis bases, but not all Bronsted acids are Lewis acids.
- Acids ionise in water to give H⁺/H₃O⁺ ions; bases give OH⁻ ions.
- The extent of ionisation depends on the strength + concentration.
- Strong → almost complete ionisation; weak → partial (equilibrium exists).
- Represented as: HA ⇌ H⁺ + A⁻
- The degree of ionisation, α, increases with dilution.
- Basis of pH and equilibrium calculations.
Classification Based on Extent of Dissociation
| Type | Examples |
|---|---|
| Strong acids | HCl, H₂SO₄, HNO₃ |
| Weak acids | HF, CH₃COOH, H₂S |
| Strong bases | NaOH, KOH |
| Weak bases | Fe(OH)₃, Cu(OH)₂ |
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.
| Salt | Hydrolysis | Solution Nature | Example |
|---|---|---|---|
| Strong acid + strong base | Does not hydrolyse | Neutral (pH = 7) |
NaCl, KNO3, NaSO4 |
| Strong acid + weak base | Cation hydrolyses | Acidic |
NH4Cl, NH4NO3, (NH4)2SO4 |
| Weak acid + strong base | Anion hydrolyses | Basic |
CH3COONa, KCN, Na2SO3 |
| Weak acid + weak base | Both hydrolyse | Depends: acidic if Ka > Kb; basic if Ka < Kb; neutral if Ka = Kb |
CH3COONH4, NH4CN, (NH4)2CO3 |
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
