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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 Bronsted 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.
The older Arrhenius definition (acid = produces H⁺ in water, base = produces OH⁻ in water) is more limited because it only works in aqueous solutions. The Bronsted-Lowry definition is broader and explains reactions that the Arrhenius model cannot, such as the reaction of HCl gas with NH₃ gas to produce NH₄Cl smoke.
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).
Why the equilibrium sign matters: The double arrow (⇌) for ammonia tells you the reaction does not go to completion. Most NH₃ molecules remain unreacted at equilibrium. The single arrow (→) for HCl tells you the reaction goes essentially to completion. This distinction between complete and incomplete proton transfer is the foundation of the strong/weak acid classification.
Every Bronsted-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.
Identifying conjugate pairs in an equation: In any proton-transfer reaction, there are always two conjugate pairs. For the general reaction HA + B ⇌ A⁻ + BH⁺, the pairs are HA/A⁻ and BH⁺/B.
Question: In the reaction HSO₄⁻(aq) + H₂O(l) ⇌ SO₄²⁻(aq) + H₃O⁺(aq), identify the acid, base, and both conjugate pairs.
Solution:
Notice that HSO₄⁻ is acting as an acid here, but in a different reaction it could act as a base (accepting a proton to form H₂SO₄). This makes HSO₄⁻ amphoteric.
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.
Why does strength matter? The distinction determines every calculation you will do. For a strong acid, [H⁺] equals the initial acid concentration. For a weak acid, [H⁺] is much less than the initial concentration and must be calculated using Ka.
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.
| Base | Type | Dissociation |
|---|---|---|
| NaOH | Strong | NaOH → Na⁺ + OH⁻ (100%) |
| KOH | Strong | KOH → K⁺ + OH⁻ (100%) |
| Ba(OH)₂ | Strong | Ba(OH)₂ → Ba²⁺ + 2OH⁻ (100%) |
| NH₃ | Weak | NH₃ + H₂O ⇌ NH₄⁺ + OH⁻ (~1%) |
| CH₃NH₂ | Weak | CH₃NH₂ + H₂O ⇌ CH₃NH₃⁺ + OH⁻ (~1%) |
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. The zwitterionic form of an amino acid has both a protonated amine group (NH₃⁺, acting as a potential acid) and a deprotonated carboxyl group (COO⁻, acting as a potential base).
Other amphoteric species you should recognise: HSO₄⁻, H₂PO₄⁻, HPO₄²⁻, and Al(OH)₃.
| Mistake | Correction |
|---|---|
| Calling a weak acid "dilute" | Weak refers to degree of dissociation, not concentration |
| Saying NaOH is a Bronsted-Lowry base because it releases OH⁻ | The Bronsted-Lowry definition says the OH⁻ ion is the base because it accepts protons |
| Confusing conjugate pairs with reactant/product pairs | A conjugate pair differs by exactly one proton |
| Forgetting water can be an acid | Water donates H⁺ to NH₃, acting as a Bronsted-Lowry acid |
| Stating H₂SO₄/SO₄²⁻ is a conjugate pair | They differ by two protons; the correct pair is H₂SO₄/HSO₄⁻ |
| Substance | Acid or Base | Strong or Weak |
|---|---|---|
| HCl | Acid | Strong |
| HNO₃ | Acid | Strong |
| H₂SO₄ | Acid | Strong (1st dissociation) |
| CH₃COOH | Acid | Weak |
| H₂CO₃ | Acid | Weak |
| HF | Acid | Weak |
| NaOH | Base | Strong |
| KOH | Base | Strong |
| Ba(OH)₂ | Base | Strong |
| NH₃ | Base | Weak |
| CH₃NH₂ | Base | Weak |
The Bronsted-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.
flowchart TD
A[Is the substance donating or accepting H⁺?] --> B{Donating H⁺}
A --> C{Accepting H⁺}
B --> D[It is an ACID]
C --> E[It is a BASE]
D --> F{Does it fully dissociate?}
E --> G{Does it fully accept protons?}
F --> H[Yes = STRONG ACID]
F --> I[No = WEAK ACID]
G --> J[Yes = STRONG BASE]
G --> K[No = WEAK BASE]