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1800-102-2727Acids and bases are among the most fundamental categories of chemical substances. Their behaviour influences ecosystems, industry, everyday life, and environmental processes. From the sour taste of lemonade to the alkaline nature of cleaning products, acid–base chemistry is everywhere.
Understanding acid–base concepts is essential in equilibrium chemistry because most acid–base reactions are reversible and governed by equilibrium principles. Concepts such as ionisation, pH, buffer action, solubility, and reaction strength are all closely connected to chemical equilibrium.
Acid–base chemistry connects molecular structure, ionisation, equilibrium constants, and proton transfer. Over time, several scientists proposed definitions to explain acid–base behaviour more accurately.
Proposed by Svante Arrhenius in 1884, this theory focuses on aqueous solutions.
A substance that produces hydrogen ions (H⁺) in water.
HCl(aq) + H₂O(l) → H₃O⁺(aq) + Cl⁻(aq)
A substance that produces hydroxide ions (OH⁻) in water.
NaOH(aq) → Na⁺(aq) + OH⁻(aq)
Proposed in 1923 by Brønsted and Lowry, this theory is based on proton transfer.
NH₃ + H₂O ⇌ NH₄⁺ + OH⁻
| Acid | Conjugate Base |
|---|---|
| HCl | Cl⁻ |
| H₂O | OH⁻ |
Proposed by Gilbert Lewis in 1923, this theory focuses on electron pair transfer.
BF₃ + :NH₃ → F₃B←NH₃
| Concept | Acid | Base | Limitation |
|---|---|---|---|
| Arrhenius | Produces H⁺ | Produces OH⁻ | Aqueous only |
| Brønsted–Lowry | Proton donor | Proton acceptor | Proton required |
| Lewis | Electron pair acceptor | Electron pair donor | Very broad |
Acid and base strength depends on the degree of ionisation and is measured using dissociation constants Ka and Kb.
pH = −log[H⁺]
| pH | Nature |
|---|---|
| < 7 | Acidic |
| 7 | Neutral |
| > 7 | Basic |
Amphoteric substances can act as both acids and bases.
Example: H₂O
| Salt Type | Example | Nature of Solution |
|---|---|---|
| Strong Acid + Strong Base | NaCl | Neutral |
| Strong Acid + Weak Base | NH₄Cl | Acidic |
| Weak Acid + Strong Base | CH₃COONa | Basic |
| Weak Acid + Weak Base | NH₄CH₃COO | Depends on Ka and Kb |
A buffer solution resists changes in pH when small amounts of acid or base are added.
Example: CH₃COOH and CH₃COONa
pH = pKa + log([Salt]/[Acid])
pOH = pKb + log([Salt]/[Base])
Because it can both donate and accept a proton.
HCl ionises completely, while acetic acid ionises partially.
Conjugate-base stability and bond dissociation energy.
Yes. BF₃ is a Lewis acid because it accepts an electron pair even though it does not produce H⁺ ions in water.
Acid–base chemistry has evolved from the Arrhenius definition to the Brønsted–Lowry proton-transfer theory and finally the broader Lewis electron-pair concept. Understanding ionisation, pH, equilibrium, conjugate acid–base pairs, and buffer systems is essential for predicting chemical behaviour in laboratory and real-world applications.