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Before the twentieth century, chemists recognised acids by their sour taste, their ability to turn litmus red, and their reactions with metals. Bases were substances that neutralised acids. These observations were useful, but they did not explain why these substances behaved the way they did.
In 1923, Johannes Brønsted and Thomas Lowry independently proposed a definition that has dominated chemistry ever since. Their framework is elegant and powerful:
This definition does not require water. It applies to reactions in the gas phase, in organic solvents, and in aqueous solution alike. It is the definition you need for Edexcel A-Level Chemistry.
Consider the reaction between hydrochloric acid and water:
HCl(g) + H₂O(l) → H₃O⁺(aq) + Cl⁻(aq)
HCl donates a proton to water, so HCl is acting as an acid. Water accepts the proton, so water is acting as a base. The product H₃O⁺ is the hydroxonium ion (sometimes called the hydronium ion), and it is the species responsible for acidic properties in aqueous solution.
Now consider ammonia dissolving in water:
NH₃(g) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq)
Here, ammonia accepts a proton from water, so ammonia is the base. Water donates a proton, so water is acting as an acid. Notice that water can be either an acid or a base depending on the reaction partner — this makes water amphoteric (or amphiprotic).
Every Brønsted-Lowry acid-base reaction involves the transfer of a proton from one species to another. When an acid donates a proton, it forms its conjugate base. When a base accepts a proton, it forms its conjugate acid.
For the HCl example:
| Species | Role | Conjugate |
|---|---|---|
| HCl | Acid | Cl⁻ (conjugate base) |
| H₂O | Base | H₃O⁺ (conjugate acid) |
For the ammonia example:
| Species | Role | Conjugate |
|---|---|---|
| H₂O | Acid | OH⁻ (conjugate base) |
| NH₃ | Base | NH₄⁺ (conjugate acid) |
A conjugate acid-base pair differs by exactly one proton. This concept is central to understanding buffer solutions later in this course.
A strong acid is one that fully dissociates in aqueous solution. Every molecule donates its proton. Examples include:
The equilibrium lies so far to the right that we write a single forward arrow (→) rather than the equilibrium symbol (⇌).
A weak acid is one that only partially dissociates. At equilibrium, most of the acid molecules remain undissociated. Examples include:
For ethanoic acid:
CH₃COOH(aq) + H₂O(l) ⇌ CH₃COO⁻(aq) + H₃O⁺(aq)
The equilibrium lies well to the left, meaning most molecules remain as CH₃COOH. The extent of dissociation is quantified by the acid dissociation constant Ka, which you will study in detail in Lesson 3.
Similarly, a strong base fully dissociates:
NaOH(s) → Na⁺(aq) + OH⁻(aq)
A weak base only partially accepts protons. Ammonia is the classic example — the equilibrium with water lies to the left, and only a small proportion of NH₃ molecules become NH₄⁺ at any given time.
Important: strong and weak refer to the degree of dissociation, not to the concentration. You can have a dilute solution of a strong acid or a concentrated solution of a weak acid.
Water is the most important amphoteric substance, but it is not the only one. The hydrogen carbonate ion, HCO₃⁻, can act as an acid:
HCO₃⁻(aq) + H₂O(l) ⇌ CO₃²⁻(aq) + H₃O⁺(aq)
or as a base:
HCO₃⁻(aq) + H₂O(l) ⇌ H₂CO₃(aq) + OH⁻(aq)
Amino acids are also amphoteric, which is critical to their behaviour at different pH values.
| Substance | Acid or Base | Strong or Weak |
|---|---|---|
| HCl | Acid | Strong |
| HNO₃ | Acid | Strong |
| H₂SO₄ | Acid | Strong (1st dissociation) |
| CH₃COOH | Acid | Weak |
| H₂CO₃ | Acid | Weak |
| NaOH | Base | Strong |
| KOH | Base | Strong |
| NH₃ | Base | Weak |
| CH₃NH₂ | Base | Weak |
The Brønsted-Lowry framework underpins everything that follows. pH calculations depend on knowing whether an acid is strong or weak. Buffer solutions rely on conjugate acid-base pairs. Titration curves look different depending on the strengths of the acid and base involved. Indicators are themselves weak acids with coloured conjugate bases.
Understanding proton transfer is the single most important foundation for this entire topic.