QCE Chemistry - Unit 3 - Chemical equilibrium systems

Bronsted-Lowry Model | QCE Chemistry

Understand acids, bases and conjugate pairs using the Bronsted-Lowry proton transfer model.

Updated 2026-05-17 - 4 min read

The Bronsted-Lowry model defines acids and bases by proton transfer.

• An acid donates a proton, $\mathrm{H^+}$.
• A base accepts a proton, $\mathrm{H^+}$.

This model is more flexible than "acids contain hydrogen and bases contain hydroxide" because it explains reactions involving species such as ammonia, carbonate ions and water.

Conjugate pairs

When an acid donates a proton, it becomes its conjugate base. When a base accepts a proton, it becomes its conjugate acid.

Example:

$ \mathrm{NH_3(aq)} + \mathrm{H_2O(l)} \rightleftharpoons \mathrm{NH_4^+(aq)} + \mathrm{OH^-(aq)} $

$\mathrm{NH_3}$ accepts a proton, so it is the base. $\mathrm{NH_4^+}$ is its conjugate acid.

$\mathrm{H_2O}$ donates a proton, so it is the acid. $\mathrm{OH^-}$ is its conjugate base.

How to identify acid and base roles

Use this process:

1. Compare formulas on the left and right.
2. Find which species lost $\mathrm{H^+}$. That species was the acid.
3. Find which species gained $\mathrm{H^+}$. That species was the base.
4. Pair each species with its conjugate partner.

Charges must still balance. Losing $\mathrm{H^+}$ makes the species one charge more negative. Gaining $\mathrm{H^+}$ makes it one charge more positive.

For example, when $\mathrm{NH_3}$ gains $\mathrm{H^+}$, it becomes $\mathrm{NH_4^+}$. When $\mathrm{CH_3COOH}$ loses $\mathrm{H^+}$, it becomes $\mathrm{CH_3COO^-}$. The atoms have barely changed, but the charge has changed because a proton has been transferred.

Amphiprotic species

An amphiprotic species can donate or accept a proton depending on what it reacts with.

Water is the classic example:

$ \mathrm{H_2O} + \mathrm{H^+} \rightarrow \mathrm{H_3O^+} $

Water acts as a base because it accepts a proton.

$ \mathrm{H_2O} \rightarrow \mathrm{H^+} + \mathrm{OH^-} $

Water acts as an acid because it donates a proton.

Hydrogen carbonate, $\mathrm{HCO_3^-}$, is also amphiprotic:

• as an acid: $\mathrm{HCO_3^-} \rightarrow \mathrm{CO_3^{2-}} + \mathrm{H^+}$
• as a base: $\mathrm{HCO_3^-} + \mathrm{H^+} \rightarrow \mathrm{H_2CO_3}$

Buffers

A buffer is a solution that resists pH change when small amounts of acid or base are added. In QCE Chemistry, buffers are important because they are conjugate in nature: they contain a weak acid and its conjugate base, or a weak base and its conjugate acid.

If extra hydrogen ions are added to an ethanoic acid/ethanoate buffer, ethanoate ions can react with them:

$ \mathrm{CH_3COO^-(aq)} + \mathrm{H^+(aq)} \rightarrow \mathrm{CH_3COOH(aq)} $

If extra hydroxide ions are added, ethanoic acid can neutralise them:

$ \mathrm{CH_3COOH(aq)} + \mathrm{OH^-(aq)} \rightarrow \mathrm{CH_3COO^-(aq)} + \mathrm{H_2O(l)} $

You do not need buffer calculations here. The main skill is explaining the response using conjugate pairs and Le Chatelier's principle.

Worked example

Exam traps

Other traps:

• pairing species that differ by more than one proton
• forgetting that water can act as either acid or base
• changing atoms other than $\mathrm{H}$ when finding conjugate pairs
• ignoring charge changes after proton transfer
• treating buffer solutions as if they stop pH change completely; they only resist small changes

Quick check

Sources

• QCAA Chemistry subject page
• QCAA Chemistry 2025 syllabus
• OpenStax Chemistry 2e: Bronsted-Lowry Acids and Bases
• OpenStax Chemistry 2e: Relative Strengths of Acids and Bases