Acid–Base Theory and Core Definitions

Specification Reference Topic 12, points 1–3, 7, 14i–ii

Quick Notes

Brønsted–Lowry acid: Proton (H⁺ ion) donor
Brønsted–Lowry base: Proton (H⁺ ion) acceptor
Acid–base reaction: Involves transfer of protons (H⁺)
Conjugate acid–base pair: Two species that differ by a single H⁺
Strong acid: Fully dissociates in water (e.g., HCl)
Weak acid: Partially dissociates in water (e.g., CH₃COOH)
Strong base: Fully dissociates to release OH⁻ ions (e.g., NaOH)
Weak base: Partially dissociates (e.g., NH₃)
pH of strong acids increases more upon dilution than weak acids
Strong acids have a lower pH than weak acids of the same concentration

Full Notes

There are several different ways to describe acids and bases in chemistry. At this level, we use the Brønsted–Lowry theory.

Brønsted–Lowry acid: A substance that donates a proton (H⁺).
Brønsted–Lowry base: A substance that accepts a proton (H⁺).

Matt’s exam tip

Remember a H⁺ ion is a proton - meaning both terms can be used when talking about acids and bases.

For Example: Reaction between HCl(aq) and NaOH(aq)

Edexcel A-Level Chemistry example showing acid–base reaction between HCl and NaOH.

HCl is the acid (proton donor)
NaOH is the base (proton acceptor).

Conjugate Acids and Bases

A conjugate acid–base pair consists of two species that differ by a single proton (H⁺).

When an acid donates a proton, the conjugate base is what remains after the acid has lost a proton.

Edexcel A-Level Chemistry diagram illustrating conjugate acid–base pairs and proton transfer.

Example: HCl → Cl⁻ + H⁺
Here, HCl is the acid and Cl⁻ is its conjugate base.

Tracking conjugate pairs helps us follow proton transfer in acid–base reactions.

Identifying Conjugate Pairs

In any acid–base reaction:

The acid and its conjugate base differ by one H⁺.
The base and its conjugate acid differ by one H⁺.

Example:
NH₄⁺ ⇌ NH₃ + H⁺
NH₄⁺ acts as the acid (donates H⁺).
NH₃ is the conjugate base (can accept H⁺).

Matt’s exam tip

When trying to determine conjugate pairs, always look for whether a proton has been lost or gained. Don’t worry about the rest of the formula or how complicated something might look, you are only interested in whether it has gained or lost a H+ ion!

Strong vs Weak Acids and Bases

Acids and bases can be classified as either strong or weak, depending on how they behave when dissolved in water.

Strong acids and bases completely dissociate in water

For Example:
HCl(aq) → H⁺(aq) + Cl⁻(aq)
NaOH(aq) → Na⁺(aq) + OH⁻(aq)

Weak acids and bases partially dissociate, forming an equilibrium system

The equilibrium contains the original weak acid, its conjugate base and H⁺ ions (for a weak acid) or the orginal weak base, its conjugate acid and H⁺ ions.

For Example:
CH₃COOH(aq) ⇌ H⁺(aq) + CH₃COO⁻(aq)

Edexcel A-Level Chemistry ethanoic acid dissociation showing equilibrium between CH3COOH and CH3COO− plus H+.

This means that strong acids release more H⁺ ions in solution than weak acids of the same concentration.

Comparing pH of Strong and Weak Acids

pH is covered in more detail here.

Even at the same concentration, strong acids have a lower pH than weak acids because they release more H⁺ ions.

Example: Comparing pH at the same concentration

1 mol dm⁻³ HCl has a pH of 0
1 mol dm⁻³ CH₃COOH has a pH around 2–3

Summary

Brønsted–Lowry acids donate protons and Brønsted–Lowry bases accept protons.
Conjugate pairs differ by a single proton and help track proton transfer.
Strong acids and bases dissociate completely while weak ones partially dissociate.
At the same concentration strong acids have a lower pH than weak acids.