Definitions [14]
Define acids according to Bronsted-Lowry theory.
A substance that donates a proton \[\ce{(H+)}\] to another substance is known as an acid.
Define conjugate acid-base pair.
A pair of an acid and a base differing by a proton is called conjugate acid-base pair.
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.
pH scale is a scale for measuring the hydrogen ion concentration in a solution.
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-]
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 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.
Define Acidic buffer solution.
A solution containing a weak acid and its salts with strong base is called an acidic buffer solution.
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
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
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.
\[\text{Molar solubility (mol/L)}=\frac{\text{Solubility in g/L}}{\text{Molar mass in g/mol}}\]
Key Points
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.
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
