The concept of acids and bases was historically defined by their physical properties, but modern chemistry uses structural definitions. The most important model for AQA A-Level Chemistry is the Brønsted-Lowry theory, which defines acids and bases entirely in terms of the transfer of hydrogen ions (protons).
A species (molecule or ion) that acts as a proton donor, releasing hydrogen ions (H+) in solution.
A species (molecule or ion) that acts as a proton acceptor, gaining hydrogen ions (H+) from a donor.
Key Principle
An acid-base reaction always involves a proton transfer. An isolated proton (H+) cannot exist in aqueous solution; it is immediately accepted by a base. In water, H+ binds to a water molecule to form the hydronium ion (H3O+).
Conjugate Acid-Base Pairs
When a Brønsted-Lowry acid loses a proton, it forms a species that is capable of gaining that proton back. This new species is called its conjugate base. Similarly, when a base gains a proton, it forms its conjugate acid.
Every acid-base reaction features two conjugate pairs. For example, when hydrogen chloride dissolves in water:
HCl(aq) + H₂O(l) ⇌ H₃O⁺(aq) + Cl⁻(aq)
- HCl (acid 1) donates a proton to form Cl⁻ (base 1). They form one conjugate pair.
- H₂O (base 2) accepts a proton to form H₃O⁺ (acid 2). They form the second conjugate pair.
Two chemical species that differ only by the presence or absence of one transferrable proton, H+.
When asked to identify conjugate pairs, remember that the acid of the pair always contains exactly one more H+ ion than its conjugate base. For example, in a pair containing H₂SO₄ and HSO₄⁻, H₂SO₄ is the acid and HSO₄⁻ is the base. Note that H₂SO₄ and SO₄²⁻ do not form a conjugate pair because they differ by two protons, not one!
Strong vs. Weak Acids
Acids differ in how easily they donate protons. We classify them as either strong or weak based on their level of dissociation (splitting into ions) in water:
An acid that dissociates completely in aqueous solution. The position of equilibrium lies entirely to the right: HA(aq) → H⁺(aq) + A⁻(aq).
An acid that only partially dissociates in aqueous solution. The position of equilibrium lies far to the left, and a reversible equilibrium is set up: HA(aq) ⇌ H⁺(aq) + A⁻(aq).
Common strong acids include hydrochloric acid (HCl), nitric acid (HNO₃), and sulfuric acid (H₂SO₄). Common weak acids include carboxylic acids such as ethanoic acid (CH₃COOH) and carbonic acid (H₂CO₃).
Step 1: Track the proton transfer:
We observe that water (H₂O) loses a proton (H⁺) to become hydroxide (OH⁻). Thus, water acts as a proton donor (acid) and hydroxide is its conjugate base.
Ammonia (NH₃) gains a proton to become ammonium (NH₄⁺). Thus, ammonia acts as a proton acceptor (base) and ammonium is its conjugate acid.
Step 2: Group the conjugate pairs:
- Conjugate pair 1: NH₄⁺ (acid) and NH₃ (base).
- Conjugate pair 2: H₂O (acid) and OH⁻ (base).
Step 1: Classify the acids:
Hydrogen chloride (HCl) is a strong acid. Ethanoic acid (CH₃COOH) is a weak acid.
Step 2: Describe dissociation:
HCl dissociates completely in water. The reaction goes to completion: HCl(g) + H₂O(l) → H₃O⁺(aq) + Cl⁻(aq). The resulting solution contains 1.0 mol dm^-3 of H₃O⁺ ions.
CH₃COOH dissociates only partially in water. A reversible equilibrium is established that lies far to the left: CH₃COOH(aq) + H₂O(l) ⇌ CH₃COO⁻(aq) + H₃O⁺(aq). Typically, only around 1% of the weak acid molecules dissociate, so the concentration of H₃O⁺ ions is much lower than 1.0 mol dm^-3 (around 0.004 mol dm^-3).
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