Acid-Base Theories

Acids and bases are among the most fundamental concepts in chemistry, governing everything from industrial manufacturing to biological processes in living cells. Over the past two centuries, chemists have developed increasingly sophisticated theoretical frameworks to explain what makes a substance an acid or a base. Three major theories — Arrhenius, Brønsted-Lowry, and Lewis — each capture different aspects of acid-base behaviour, and each successive theory broadens the definition to encompass a wider range of chemical phenomena. Understanding acid-base theories is essential for mastering ionic equilibrium, because acids and bases exist in solution as ions that establish dynamic equilibria. The strength of an acid or base, the pH of a solution, and the behaviour of buffer systems all depend on which theory you apply and how completely a substance donates or accepts protons or electrons. Each theory is not wrong — rather, they are nested frameworks, with later theories generalising the earlier ones. In this article, you will explore the historical development of acid-base theories, learn to identify conjugate acid-base pairs in Brønsted-Lowry reactions, understand the electron-pair donor-acceptor model of Lewis, and apply these concepts to real chemical equations. These ideas underpin titrations, buffer calculations, solubility equilibria, and much of organic and inorganic reaction chemistry you will encounter throughout your studies.

Key takeaways
  • Arrhenius Theory (1884)
  • Brønsted-Lowry Theory (1923)
  • Conjugate Acid-Base Pairs
  • Lewis Theory (1923)
  • Comparing the Three Theories

Arrhenius Theory (1884)

Svante Arrhenius proposed the first modern acid-base theory. An Arrhenius acid is any substance that produces hydrogen ions (H⁺) when dissolved in water. An Arrhenius base is any substance that produces hydroxide ions (OH⁻) in aqueous solution. For example, hydrochloric acid dissociates: HCl → H⁺ + Cl⁻, and sodium hydroxide dissociates: NaOH → Na⁺ + OH⁻. Neutralisation is simply: H⁺ + OH⁻ → H₂O. The main limitation of this theory is that it is restricted to aqueous solutions and cannot explain bases like ammonia (NH₃) that contain no hydroxide group yet still behave as bases.

Brønsted-Lowry Theory (1923)

Johannes Brønsted and Thomas Lowry independently proposed a broader definition: an acid is a proton (H⁺) donor, and a base is a proton (H⁺) acceptor. This applies to any solvent, not just water. The concept of conjugate pairs is central: when an acid donates a proton, it becomes its conjugate base; when a base accepts a proton, it becomes its conjugate acid. In the reaction NH₃ + H₂O ⇌ NH₄⁺ + OH⁻, water acts as the acid (donating H⁺) and NH₃ acts as the base. The conjugate acid of NH₃ is NH₄⁺, and the conjugate base of H₂O is OH⁻. A strong acid has a weak conjugate base, and vice versa. Amphoteric species like water or HCO₃⁻ can act as either acid or base depending on the reaction partner.

Conjugate Acid-Base Pairs

A conjugate acid-base pair differs by exactly one proton (H⁺). In the general reaction: HA + B ⇌ A⁻ + BH⁺, the pair HA/A⁻ is one conjugate pair (HA is the acid, A⁻ is its conjugate base), and B/BH⁺ is the other conjugate pair (B is the base, BH⁺ is its conjugate acid). For acetic acid in water: CH₃COOH + H₂O ⇌ CH₃COO⁻ + H₃O⁺. Here, CH₃COOH/CH₃COO⁻ is one conjugate pair and H₂O/H₃O⁺ is the other. The stronger the acid, the weaker its conjugate base, because a stronger acid releases its proton more readily, leaving behind an anion with little tendency to recapture it.

Lewis Theory (1923)

Gilbert N. Lewis proposed the most general theory: a Lewis acid is an electron-pair acceptor, and a Lewis base is an electron-pair donor. This definition encompasses all Brønsted-Lowry acids and bases, but also includes reactions with no proton transfer. For example, BF₃ (boron trifluoride) accepts an electron pair from NH₃ (which donates its lone pair on nitrogen) to form the adduct F₃B←NH₃. BF₃ is a Lewis acid (electron-pair acceptor) even though it cannot donate a proton. Metal ions acting as central atoms in coordination compounds are Lewis acids, accepting lone pairs from ligands (Lewis bases). Lewis theory explains the formation of coordination complexes and many organic reactions such as Friedel-Crafts alkylation.

Comparing the Three Theories

Every Arrhenius acid is also a Brønsted-Lowry acid, and every Brønsted-Lowry acid is also a Lewis acid — the theories are nested. Arrhenius theory is limited to aqueous solutions. Brønsted-Lowry theory works in any solvent and introduces the powerful conjugate pair concept. Lewis theory is the broadest and covers all reactions involving electron-pair donation, including those with no proton transfer at all. For most equilibrium calculations at high school level, Brønsted-Lowry is the most useful, but Lewis theory becomes essential in transition metal chemistry, organic mechanisms, and understanding why certain molecules act as catalysts.

The Key Equation

\ceCH3COOH+H2O<=>CH3COO−+H3O+\ce{CH3COOH + H2O <=> CH3COO^{-} + H3O^{+}}
  • \ce{CH3COOH} — Acetic acid (mol/L): The Brønsted-Lowry acid; donates H⁺ to water. A weak acid that only partially dissociates.
  • \ce{H2O} — Water (—): Acts as the Brønsted-Lowry base in this reaction, accepting the proton from acetic acid.
  • \ce{CH3COO^{-}} — Acetate ion (mol/L): The conjugate base of acetic acid; formed after acetic acid donates its proton.
  • \ce{H3O^{+}} — Hydronium ion (mol/L): The conjugate acid of water; formed when water accepts a proton from acetic acid.

Worked Example

Problem. Identify the Brønsted-Lowry acid, base, conjugate acid, and conjugate base in the following reaction, and state whether HSO₄⁻ is acting as an acid or a base: HSO₄⁻ + NH₃ ⇌ SO₄²⁻ + NH₄⁺

\ceHSO4−+NH3<=>SO42−+NH4+\ce{HSO4^{-} + NH3 <=> SO4^{2-} + NH4^{+}}

Answer. HSO₄⁻ is the Brønsted-Lowry acid; NH₃ is the Brønsted-Lowry base. Conjugate pairs: HSO₄⁻/SO₄²⁻ and NH₃/NH₄⁺.

Figure
Diagram
Conjugate Acid-Base Pairs in Brønsted-Lowry TheoryCH₃COOHAcid (H⁺ donor)CH₃COO⁻Conjugate Base− H⁺Conjugate Pair 1H₂OBase (H⁺ acceptor)H₃O⁺Conjugate Acid+ H⁺Conjugate Pair 2H⁺ transferLewis Acid-Base: BF₃ + NH₃ → F₃B←NH₃BFFFLewis Acid (empty orbital)NHHHLewis Base (lone pair)lone pair donated
Origins · 1 minSvante Arrhenius and how this was discovered
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Svante Arrhenius
1859–1927 · Swedish

Arrhenius proposed in his 1884 doctoral thesis that electrolytes dissociate into ions in solution, and defined acids as substances producing H⁺ ions and bases as substances producing OH⁻ ions in water. His thesis was initially awarded a low grade by his examiners, but he was later vindicated when he received the Nobel Prize in Chemistry in 1903 for his theory of electrolytic dissociation. Arrhenius was one of the first scientists to connect carbon dioxide emissions from burning fossil fuels to global warming, predicting in 1896 that doubling atmospheric CO₂ would raise Earth's temperature by 5–6°C — a remarkably prescient calculation for over a century ago.

  • pH and pOH Calculations — pH is defined using H₃O⁺ concentration, which arises directly from Brønsted-Lowry acid dissociation equilibria. Understanding acid-base theories is a prerequisite for calculating pH.
  • Weak Acids and the Acid Dissociation Constant Ka — Ka quantifies the extent of proton donation in Brønsted-Lowry equilibria. The relationship between acid strength and conjugate base weakness is central to Ka calculations.
  • Buffer Solutions — Buffers consist of a weak acid and its conjugate base — a concept that comes directly from Brønsted-Lowry conjugate pair theory.
  • Coordination Compounds and Ligands — Ligands are Lewis bases that donate electron pairs to metal ions (Lewis acids). Lewis acid-base theory is essential for understanding transition metal complex formation.