4.4Unit 4. Rates, Equilibria and Further Organic Chemistry

Acid-base Equilibria — IAL Edexcel Modular (Chemistry) Revision Notes & Topical Questions

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4.4 Acid-base Equilibria — IAL EDEXCEL MODULAR (CHEMISTRY) notes: Identify Brønsted–Lowry acids, bases and conjugate pairs, calculate pH using Ka, pK_a and Kw, interpret titration curves and choose indicators, and explain and calculate buffer action.Brønsted acid = a proton donor…

Conjugate acid–base pair
Amphiprotic
What the symbols mean
Strong acid
Acid dissociation constant, Ka
Strong monobasic acid

Section 1 of 13Brønsted–Lowry theory

Brønsted acid

Brønsted acid a proton donor; Brønsted base = a proton acceptor

HCl is an acid because it donates H⁺.
OH⁻ is a base because it accepts H⁺ to form water.

Conjugate acid–base pair

Conjugate acid–base pair two species related by the transfer of one proton
In CH₃COOH + H₂O ⇌ CH₃COO⁻ + H₃O⁺, the pairs are CH₃COOH/CH₃COO⁻ and H₃O⁺/H₂O.

Amphiprotic

Amphiprotic water can act as an acid or a base

Section 2 of 13pH and acid strength

pH and acid strength

−log[H⁺],

so [H⁺] = 10^−pH
What the symbols mean: pH has no unit · [H⁺] = concentration of hydrogen ions (mol dm⁻³); each 1 pH unit is a ten-fold change in [H⁺]
The scale is logarithmic, so each pH unit is a factor of 10 in [H⁺].
pH is usually given to 2 decimal places.

Strong acid

Strong acid dissociates almost completely, such as HCl, HNO₃ and H₂SO₄
The equilibrium lies so far right that it is treated as irreversible.

Weak acid

Weak acid partially dissociates, so an equilibrium is established
Examples are ethanoic acid, HCN, H₂S and H₂CO₃.

Acid dissociation constant, Ka

[H⁺][A⁻] ÷ [HA], in mol dm^-3

What the symbols mean: K_a = acid dissociation constant (mol dm⁻³) · [H⁺], [A⁻], [HA] = equilibrium concentrations (mol dm⁻³); larger K_a = stronger acid

A larger K_a means a more dissociated and stronger acid.
Ethanoic acid has K_a = 1.74 × 10^-5 mol dm^-3.

pKa

−log K_a

A smaller pK_a means a stronger acid.
Most weak acids have pK_a between 3 and 7.

Section 3 of 13pH of strong and weak acids

Strong monobasic acid

[H⁺] equals the acid concentration
Water's own H⁺ is negligible.

Dibasic acid

Dibasic acid sulfuric acid ionises in two steps, H₂SO₄ → HSO₄⁻ + H⁺ and HSO₄⁻ ⇌ SO₄²⁻ + H⁺
The second step is incomplete, so [H⁺] is less than twice the acid concentration.

Weak acid

[H⁺] = √(K_a × [HA]) and pH = −log[H⁺]
Two assumptions are made[A⁻] = [H⁺], and [HA] at equilibrium equals the initial [HA] because dissociation is very small.

Finding Ka

Finding Ka from a measured pH, K_a = [H⁺]² ÷ [HA]

Section 4 of 13Ionic product of water

[H⁺][OH⁻] = 1 × 10^-14 mol² dm^-6 at 298 K

What the symbols mean: K_w = ionic product of water (1.00 × 10⁻¹⁴ mol² dm⁻⁶ at 298 K) · [H⁺], [OH⁻] = ion concentrations (mol dm⁻³)

pK_w = 14.00 at 298 K.

Pure water

[H⁺] = [OH⁻], so it is neutral at all temperatures
K_w increases with temperature because dissociation is endothermic, so the neutral pH falls below 7 when hot.

Strong base

[OH⁻] equals the base concentration
[H⁺] = K_w ÷ [OH⁻], then pH = −log[H⁺].

Finding [OH−] from pH

[H⁺] = 10^−pH, then [OH⁻] = K_w ÷ [H⁺]

Section 5 of 13Analysing pH data

Relative strength

Relative strength compare the pH of equimolar solutions of acids at the same temperature
The higher the pH, the weaker the acid; the higher the pH, the stronger the base.

Salts

NaCl and KNO₃ are neutral (pH 7) from a strong acid and strong base.
CH₃COONa is alkaline, from a weak acid and strong base.
NH₄Cl is acidic, from a strong acid and weak base.

Dilution

Dilution a strong acid diluted ten times rises by 1 pH unit, and a weak acid rises by about 0.5
A very dilute strong acid (10^-8 mol dm^-3) has a pH close to 7, because the H⁺ from water is no longer negligible.

Calibration

Calibration pH meters are calibrated with buffer solutions

Section 6 of 13Titration curves

pH curve
- a graph of pH against volume of titrant, with an S-shape
- The equivalence point is at the centre of the steep section.
- The curve gives the starting pH, the equivalence volume, the pH at equivalence and the pH range of the vertical section.

Titration	Start pH	Equivalence pH	Vertical range
Strong acid + strong base	Low (about 1)	7	Large, about 3 to 11
Weak acid + strong base	About 3	Above 7	About 7 to 11
Strong acid + weak base	Low (about 1)	Below 7	About 3 to 7
Weak acid + weak base	About 3	About 7	No sharp vertical section

1Weak acid with strong base

Weak acid with strong base a buffer region forms before equivalence as ethanoate ions form
At half-equivalence [HA] = [A⁻], so pH = pK_a.

2Titrating in reverse

Titrating in reverse adding acid to base gives the mirror image of the curve

Section 7 of 13Choosing an indicator

Indicators

Indicators weak acids whose conjugate forms have different colours: HIn ⇌ H⁺ + In⁻
The constant is K_In, used as pK_In.
The working range is about pK_In ± 1.

Rule

Rule the colour change range must lie within the vertical section of the titration curve
Methyl orange (pK_In 3.7) is unsuitable for weak acid with strong base, because its change is far below the vertical section.
Phenolphthalein, with a higher pH range, is suitable.

Section 8 of 13Buffer solutions

Buffer

Buffer a solution that resists changes in pH when small amounts of acid or alkali are added

An acidic buffer is a weak acid and its conjugate base; a basic buffer is a weak base and its conjugate acid.
Ethanoic acid with sodium ethanoate is a common example.

Action of an acidic buffer HA / A⁻

The acid is partly ionised and the salt is fully ionised, so there are large reserves of both HA and A⁻.
Added H⁺ reacts with A⁻ to form HA, so [H⁺] hardly changes.
Added OH⁻ reacts with H⁺ to form water, and HA ionises to replace it.
The ratio [HA] ÷ [A⁻] changes only slightly, so the pH stays almost constant.

Making a buffer

Making a buffer partially neutralise a weak acid with NaOH, so that weak acid and its conjugate base remain
A buffer below pH 7 uses a weak acid and its salt.

Applications

Applications blood is buffered at pH 7.35–7.45 by CO₂ + H₂O ⇌ H⁺ + HCO₃⁻
Too much acid in the blood causes acidosis.
Food with more protein has a higher buffer capacity, so it spoils more slowly.

Section 9 of 13Buffer calculations

Buffer calculations

[H+]

K_a × [acid] ÷ [salt]

The concentrations of the weak acid and its conjugate base at equilibrium are used.

Henderson–Hasselbalch equation

Henderson–Hasselbalch equation pH = pK_a + log([A⁻] ÷ [HA])
Equal concentrations give pH = pK_a.

Designing a buffer of required pH

Find [H⁺] = 10^−pH.
Rearrange to [acid] ÷ [base] = [H⁺] ÷ K_a.
Choose concentrations of acid and salt in that ratio.

Section 10 of 13More from the specification

1Brønsted-Lowry theory

An acid donates a proton and a base accepts one; each acid has a conjugate base.

2Ionisation of water

Kw = [H+][OH-] = 1.0 × 10^-14 at 298 K; the expression applies to aqueous solutions.

3Weak acid constant

Ka = [H+][A-]/[HA]; calculations assume equilibrium concentrations and negligible ionisation.

4Strong versus weak

Strong acids ionise fully; weak acids partly, so the position of equilibrium lies to the left.

5Buffers

A weak acid with its salt resists pH change, as excess added hydroxide is neutralised by the acid.

6Aqueous

Dissolved in water.

7Hydroxide

A compound containing the OH⁻ ion, formed by Group I metals with water.

8Ionisation

The removal of an electron from an atom or molecule.

9Excess

More than enough of a reactant.

10Equilibrium position

Point of minimum free energy.

Section 11 of 13More on this topic

Buffers and titration curves
- buffers resist pH change, and indicators are chosen to match the vertical section of the curve

Section 12 of 13Key terms

Term	Meaning
Brønsted acid	A proton donor.
Brønsted base	A proton acceptor.
Conjugate pair	Two species that differ by one proton.
Acid dissociation constant (K_a)	The equilibrium constant for the dissociation of a weak acid.
Ionic product of water (K_w)	K_w = [H⁺][OH⁻], equal to 1 × 10⁻¹⁴ mol² dm⁻⁶ at 298 K.
Buffer	A solution that resists changes in pH when small amounts of acid or alkali are added.
Equivalence point	The point at which the acid and base have reacted in the ratio of the equation.
pK_a	−log K_a.

Section 13 of 13Exam technique

Confusing strength with concentration — strength is the extent of dissociation

Using [H⁺] = [HA] for a weak acid — use √(K_a × [HA])

Using an unsuitable indicator — its range must lie within the vertical section

Saying pure water is only neutral at 25 °C — it is neutral whenever [H⁺] = [OH⁻]

Thinking a buffer stops pH changing — it only resists change

Last updated: 18 October 2026