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Acid–Base Concepts: Arrhenius, Brønsted–Lowry & Lewis Theory Explained

Acid–Base Concepts: Arrhenius, Brønsted–Lowry & Lewis Theory Explained

Acid–Base Concepts: From Simple Neutralisation to Chemical Equilibrium

Acids 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.

Introduction

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.

  • Arrhenius Concept
  • Brønsted–Lowry Concept
  • Lewis Concept

The Arrhenius Acid–Base Concept

Proposed by Svante Arrhenius in 1884, this theory focuses on aqueous solutions.

Arrhenius Acid

A substance that produces hydrogen ions (H⁺) in water.

HCl(aq) + H₂O(l) → H₃O⁺(aq) + Cl⁻(aq)

Arrhenius Base

A substance that produces hydroxide ions (OH⁻) in water.

NaOH(aq) → Na⁺(aq) + OH⁻(aq)

Limitations

  • Applicable only to aqueous solutions.
  • Cannot fully explain ammonia's basic nature.
  • Does not include non-aqueous systems.

Brønsted–Lowry Concept

Proposed in 1923 by Brønsted and Lowry, this theory is based on proton transfer.

  • Acid: Proton (H⁺) donor
  • Base: Proton (H⁺) acceptor

NH₃ + H₂O ⇌ NH₄⁺ + OH⁻

Conjugate Acid–Base Pairs

Acid Conjugate Base
HCl Cl⁻
H₂O OH⁻

Lewis Concept

Proposed by Gilbert Lewis in 1923, this theory focuses on electron pair transfer.

  • Lewis Acid: Electron pair acceptor
  • Lewis Base: Electron pair donor

BF₃ + :NH₃ → F₃B←NH₃

Comparison of Acid–Base Theories

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

Strength of Acids and Bases

Acid and base strength depends on the degree of ionisation and is measured using dissociation constants Ka and Kb.

  • pKa = −log Ka
  • pKb = −log Kb
  • Higher Ka = Stronger Acid = Lower pKa
  • Ka × Kb = Kw

pH Scale

pH = −log[H⁺]

pH Nature
< 7 Acidic
7 Neutral
> 7 Basic

Amphoteric Substances

Amphoteric substances can act as both acids and bases.

Example: H₂O

Salt Hydrolysis

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

Buffer Solutions

A buffer solution resists changes in pH when small amounts of acid or base are added.

Example: CH₃COOH and CH₃COONa

Henderson–Hasselbalch Equation

pH = pKa + log([Salt]/[Acid])

pOH = pKb + log([Salt]/[Base])

Acid–Base Titration Curves

  • Strong Acid vs Strong Base: Equivalence point pH = 7
  • Weak Acid vs Strong Base: Equivalence point pH > 7
  • Weak Base vs Strong Acid: Equivalence point pH < 7

Applications of Acid–Base Concepts

  • Industrial processes
  • Biological systems (blood pH)
  • Pharmaceutical chemistry
  • Environmental monitoring
  • Titration analysis

Frequently Asked Questions (FAQs)

Why is water amphoteric?

Because it can both donate and accept a proton.

Why is HCl a strong acid but CH₃COOH weak?

HCl ionises completely, while acetic acid ionises partially.

What determines acid strength?

Conjugate-base stability and bond dissociation energy.

Can a substance be a Lewis acid without being an Arrhenius acid?

Yes. BF₃ is a Lewis acid because it accepts an electron pair even though it does not produce H⁺ ions in water.

Conclusion

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.

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