Brønsted–Lowry theory of acids and bases

Specification Reference Physical Chemistry: Equilibria 7.2

Quick Notes

Brønsted-Lowry acids are proton (H⁺) donors.
Brønsted-Lowry bases are proton (H⁺) acceptors.
Acid-base equilibrium reactions involve conjugate acid-base pairs.
Strong acids fully dissociate, while weak acids partially dissociate in solution.
pH scale (at 298 K):
- Acidic: pH < 7
- Neutral: pH = 7
- Alkaline: pH > 7
Neutralisation reactions occur when H⁺ ions combine with OH⁻ ions: H⁺ + OH⁻ → H₂O
Salts form during neutralisation
Titration curves depend on acid/base strength
Indicators in titrations are chosen based on the equivalence point of the titration (not all indicators work for all titrations)

Full Notes

Brønsted–Lowry theory of acids and bases

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⁺).

Common Acids (names and formulas)

You need to know the names and chemical formulas of a few key acids:

Hydrochloric acid – HCl
Sulfuric acid – H₂SO₄
Nitric acid – HNO₃
Ethanoic acid – CH₃COOH

These acids are commonly used in titrations and reactions.

Common Alkalis (names and formulas)

You need to know the names and chemical formulas of a few key bases:

Sodium hydroxide – NaOH
Potassium hydroxide – KOH
Ammonia – NH₃

These substances dissolve in water and cause the OH⁻ ion concentration in the solution to increase.

Strong vs Weak Acids and Bases

Strong acids and bases completely dissociate in water (e.g. HCl, NaOH)

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

Weak acids/bases partially dissociate, forming an equilibrium (e.g. CH₃COOH, NH₃)

CH₃COOH(aq) ⇌ H⁺(aq) + CH₃COO⁻(aq)
NH₃(aq) + H⁺(aq) ⇌ NH₄⁺(aq)

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

pH Scale Basics

The pH scale shows how acidic or basic a solution is, based on H⁺ ion concentration.

pH < 7 = acidic [H⁺] is greater than [OH^-]
pH = 7 = neutral (pure water) [H⁺] = [OH^-]
pH > 7 = alkaline [OH^-] is greater than [H⁺]

Note this is for a temperature of 298 K (pH changes with temperature - see here for more detail).

Comparing Strong and Weak Acids

Strong acids:

React faster with metals
Give lower pH readings
Show higher electrical conductivity (more ions) at a given concentration of acid

Weak acids:

React more slowly
Show higher pH values
Have lower electrical conductivity(fewer ions) at a given concentration of acid

Neutralisation Reactions

A neutralisation is when H⁺ ions from an acid reacts with OH⁻ from a base to form H₂O.

The net-ionic equation is always H⁺(aq) + OH⁻(aq) → H₂O(l)

This reaction removes acidity and creates a neutral solution.

Salt Formation

In neutralisation reactions, a salt gets formed.

The H⁺ from the acid is replaced by a metal ion (or NH₄⁺ ion) from the base, forming a salt.

For example:

HCl + NaOH → NaCl + H₂O

Here, the Na⁺ ion from the base (NaOH) replaces the H⁺ ion from the acid (HCl).

The salt NaCl is formed from the ions left over in the mixture after H⁺ and OH⁻ ions have reacted.

Salts are ionic compounds formed from acid–base reactions.

pH Titration Curves

The background theory behind titration curves has been covered in more detail here.

Titration curves show how pH changes as one solution is added to another.

Strong acid + strong base → sharp jump at pH 7

CIE A-Level Chemistry titration curve for strong acid plus strong base showing a sharp jump at pH 7.

Weak acid + strong base → jump at higher pH (~9)

CIE A-Level Chemistry titration curve for weak acid plus strong base with the equivalence jump around pH 9.

Strong acid + weak base → jump at lower pH (~5)

CIE A-Level Chemistry titration curve for strong acid plus weak base with the equivalence jump around pH 5.

Weak acid + weak base → no sharp jump, titration doesn’t work

CIE A-Level Chemistry titration curve for weak acid plus weak base showing no sharp equivalence jump.

These curves help you choose the right indicator.

Choosing Indicators

During a titration, an indicator is used that changes colour at a certain pH.

When the solution reaches this pH, the indicator changes colour – this is how the person carrying out the titration knows it is ‘complete’.

The end point of a titration is when enough ‘titrant’ has been added to make the indicator change colour.

Indicator Colour Changes:

CIE A-Level Chemistry diagram for phenolphthalein indicator colour change (colourless in acid to pink in base, pH ~8.3–10).

CIE A-Level Chemistry diagram for methyl orange indicator colour change (red in acid to yellow in base, pH ~3.5–4.5).

Methyl Orange: Red (acid) → Yellow (base) (pH ~3.5 - 4.5).
Phenolphthalein: Colourless (acid) → Pink (base) (pH ~8.3 - 10).

Different indicators can change colour at different pH values, this is why the same indicators aren’t always used for different titrations.

Indicators should be chosen that change colour at a pH that falls within the sharp peak area of a titration curve.

Acid–Base Combination	Typical Equivalence pH	Suitable Indicator(s)
Strong acid + Strong base	~7	Methyl orange or Phenolphthalein
Weak acid + Strong base	>7 (around 8–10)	Phenolphthalein
Strong acid + Weak base	<7 (around 3–6)	Methyl orange
Weak acid + Weak base	No sharp jump	None suitable

Summary

Brønsted-Lowry acids are proton donors and bases are proton acceptors; reactions involve conjugate acid-base pairs.
Strong acids/bases fully dissociate; weak acids/bases partially dissociate and set up equilibria.
pH at 298 K: <7 acidic, 7 neutral, >7 alkaline; pH varies with temperature.
Neutralisation: H⁺ + OH⁻ → H₂O; salts are formed.
Titration curves depend on acid/base strength; choose indicators whose colour change range lies within the sharp jump region.