Brønsted–Lowry acids and bases

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⁺).

Example: In the reaction between HCl(aq) and NaOH(aq)

HCl is the acid (proton donor)

NaOH is the base (proton acceptor).

Acids are also classified based on the number of protons they can donate:

• Monobasic acid: releases 1 H⁺ per molecule (e.g. HCl, CH₃COOH)
• Dibasic acid: releases 2 H⁺ per molecule (e.g. H₂SO₄)
• Tribasic acid: releases 3 H⁺ per molecule (e.g. H₃PO₄)

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.

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!

Acid Reactions and Ionic Equations

Acids are proton donors, and this is reflected in their reactions with other substances. These are often written as ionic equations, which simplify the reaction to show only the ions involved. The key feature is that H⁺ is the reactive acid component, not necessarily the full acid molecule (e.g. HCl or H₂SO₄). For example:

• With metals:
  Acids donate protons to metals to form salts and hydrogen gas.

  Example:Mg + 2H⁺ → Mg²⁺ + H₂

• With metal oxides:
  Form salts and water. The oxide acts as a base.

  Example:2H⁺ + O²⁻ → H₂O

• With carbonates:
  Form salts, water, and carbon dioxide.

  Example:2H⁺ + CO₃²⁻ → CO₂ + H₂O

• With alkalis (e.g. NaOH):
  A neutralisation reaction forms water.

  Example:H⁺ + OH⁻ → H₂O

Acid Dissociation Constant (K_a)

A weak acid (HA) only partially dissociates in water:

Unlike strong acids (which fully dissociate), weak acids establish an equilibrium. The position of this equilibrium is based on the strength of the acid - the stronger the acid, the more the position lies to the right (more dissociation), meaning a higher concentration of H⁺ and A⁻ ions in the mixture.

Like with any equilibrium system, the position of this equilibrium can be described using an equilibrium constant, Ka. K_a is the equilibrium constant for acid dissociation.

• Larger K_a = Stronger weak acid (more dissociation).
• Smaller K_a = Weaker acid (less dissociation).

Example: Ethanoic acid (CH₃COOH)

CH₃COOH ⇌ CH₃COO⁻ + H⁺
K_a = 1.75 × 10⁻⁵ mol dm⁻³ (shows weak dissociation).

Relationship Between K_a and pK_a

pK_a is a logarithmic measure of K_a:

• Lower pK_a = Stronger weak acid (more dissociation).
• Higher pK_a = Weaker acid (less dissociation).
• To convert pK_a to K_a, use K_a = 10^−pK_a

Example: Comparing weak acids

• Ethanoic acid: K_a = 1.75 × 10⁻⁵, so pK_a = 4.76.
• Carbonic acid: K_a = 4.3 × 10⁻⁷, so pK_a = 6.37.
• Since pK_a of ethanoic acid is lower, it is stronger than carbonic acid.

Matt’s exam tip

Don’t get confused by K_a and pK_a. Remember, the bigger the value of K_a, the stronger the acid. The bigger the value of pK_a, the the weaker the acid.

Summary

• Brønsted–Lowry acids donate H⁺ and bases accept H⁺.
• Conjugate pairs differ by one proton.
• Write acid reactions as ionic equations focusing on H⁺.
• Ka measures weak acid dissociation and pK_a = −log(K_a).
• Stronger acids have larger K_a and smaller pK_a.