Introduction to Acids and Bases
Properties of Acids and Bases
| Acids | Bases |
|---|---|
| Have sour taste | Have bitter taste |
| Change litmus from blue to red | Change litmus from red to blue |
| Change phenolphthalein from pink to colorless | Change phenolphthalein from colorless to pink |
| React with bases to form salts | React with acids to form salts |
The reaction between an acid and a base is called neutralization. During neutralization, acids and bases react with each other to produce ionic substances called salts.
Example: HCl (acid) and NaOH (base) react to produce NaCl (salt) and water.
Brønsted-Lowry Concept
Definition
In 1923, J.N. Brønsted (Danish) and T.M. Lowry (English) independently expanded the Arrhenius theory.
According to Brønsted-Lowry concept:
An acid is the species donating a proton in a proton-transfer reaction.
A base is the species accepting the proton in a proton-transfer reaction.
Example: When HCl dissolves in water:
HCl (acid) donates H⁺, H₂O (base) accepts H⁺
Conjugate Acid-Base Pairs
Conjugate acid: Species formed after a Brønsted base accepts a proton.
Example: H₃O⁺ is conjugate acid of H₂O
Conjugate base: Species formed when an acid donates a proton.
Example: Cl⁻ is conjugate base of HCl
Conjugate acid-base pairs: Species on opposite sides of an equation that differ by a proton.
Lewis Concept of Acids and Bases
Definition
In 1923, G.N. Lewis proposed a generalized definition:
Lewis acid: Any species (molecule or ion) that can accept a pair of electrons.
Lewis base: Any species (molecule or ion) that can donate a pair of electrons.
Example 1: BF₃ (Lewis acid) + NH₃ (Lewis base) → BF₃NH₃
Example 2: H⁺ (Lewis acid) + OH⁻ (Lewis base) → H₂O
pH and pOH Calculations
Definitions
pH = -log[H⁺] (negative log of hydrogen ion concentration)
pOH = -log[OH⁻] (negative log of hydroxyl ion concentration)
Key Relationship: pH + pOH = 14 (at 25°C)
For neutral water: [H⁺] = [OH⁻] = 10⁻⁷ M, so pH = pOH = 7
Ionic Product of Water (Kw)
Kw = [H⁺][OH⁻] = 10⁻¹⁴ at 25°C
Kw increases with temperature (75 times from 0°C to 100°C)
| Temperature (°C) | Kw |
|---|---|
| 0 | 0.11 × 10⁻¹⁴ |
| 10 | 0.30 × 10⁻¹⁴ |
| 25 | 1.0 × 10⁻¹⁴ |
| 40 | 3.00 × 10⁻¹⁴ |
| 100 | 7.5 × 10⁻¹⁴ |
Ionization Constants of Acids (Ka)
Definition and Significance
Ka = [H₃O⁺][A⁻]/[HA]
Quantitative measure of acid strength
| Ka Value | Acid Strength | Examples |
|---|---|---|
| > 1 | Strong | HCl, HNO₃, H₂SO₄ |
| 1 to 10⁻³ | Moderately strong | HSO₄⁻ |
| < 10⁻³ | Weak | CH₃COOH, HF, H₂CO₃ |
| ~10⁻¹⁶ | Very weak | H₂O |
Calculating [H₃O⁺] from Ka
For weak acid HA: HA ⇌ H⁺ + A⁻
Ka = x²/(C – x) ≈ x²/C (for x << C)
[H₃O⁺] = x ≈ √(Ka × C)
Buffer Solutions and Common Ion Effect
Buffer Solutions
Solutions that resist pH change when small amounts of acid/base are added.
Types:
1. Acidic buffer: Weak acid + its salt with strong base (pH < 7)
Example: CH₃COOH + CH₃COONa
2. Basic buffer: Weak base + its salt with strong acid (pH > 7)
Example: NH₄OH + NH₄Cl
Common Ion Effect
Suppression of ionization of weak electrolyte by adding common ion.
Example: Adding HCl to NaCl solution causes NaCl precipitation.
Henderson Equation
For acidic buffer: pH = pKa + log([salt]/[acid])
For basic buffer: pOH = pKb + log([salt]/[base])