IB DP Chemistry Reactivity 3.1: Proton transfer: acids, bases and pH

This study guide teaches the standard level part of IB DP Chemistry Reactivity 3.1, “Proton transfer reactions”, from first principles. Its source is the IB Diploma Programme Chemistry guide (first assessment 2025; February 2023 edition), and every understanding from Reactivity 3.1.1 to 3.1.8 gets a worked example. All of it is for both SL and HL; the additional higher level material that follows (Reactivity 3.1.9 onwards) has its own HL study guide. The page applies to sessions from 2025 onwards.

Afterwards, use the revision notes for a fast recap and the practice questions to test yourself. The Reactivity 3 study guide gives the whole theme. Track progress with the course hub and printable checklist, and find gaps with a free diagnostic.

What this unit covers

Understanding	What you must be able to do
Reactivity 3.1.1	Pick out which species is the Brønsted–Lowry acid and which the base; use H⁺(aq) and H₃O⁺(aq); tell a base from an alkali
Reactivity 3.1.2	Write the conjugate acid or base of any species
Reactivity 3.1.3	Write equations showing one species acting as an acid and as a base
Reactivity 3.1.4	Convert between pH and [H⁺]; estimate and measure pH
Reactivity 3.1.5	Use Kw to call a solution acidic, neutral or basic
Reactivity 3.1.6	Explain strong versus weak and concentrated versus dilute; predict where an equilibrium lies
Reactivity 3.1.7	Write neutralization equations; name the parent acid and base of a salt
Reactivity 3.1.8	Sketch and interpret a strong acid–strong base pH curve

Every row is for SL and HL.

Reactivity 3.1.1 Acids donate protons, bases accept them

In the Brønsted–Lowry model, acids donate protons (H⁺ ions) and bases accept them. In water a proton bonds to a water molecule’s lone pair, giving H₃O⁺. You may write it as H⁺(aq) or H₃O⁺(aq).

Base or alkali? An alkali is a soluble base: in water it releases OH⁻ ions, as sodium hydroxide does. Copper(II) oxide is a base, since it neutralizes acids, yet it does not dissolve, so it is not an alkali.

Worked example 1: spotting the acid and the base

(a) HNO₃(aq) + H₂O(l) → H₃O⁺(aq) + NO₃⁻(aq)

(b) HCOOH(aq) + CH₃NH₂(aq) → HCOO⁻(aq) + CH₃NH₃⁺(aq)

Follow the proton: the species that loses H⁺ is the acid; the one that gains it is the base.

(a) HNO₃ → NO₃⁻ loses a proton: HNO₃ is the acid. H₂O → H₃O⁺: water is the base.

(b) Methanoic acid loses a proton: HCOOH is the acid. Methylamine gains it: CH₃NH₂ is the base. No water or OH⁻ is involved, so the model reaches beyond aqueous hydroxides.

Why the definition evolved. Acidity was once linked to oxygen, until oxygen-free acids such as hydrogen chloride were studied. Arrhenius linked acids to H⁺ and bases to OH⁻ in water; Brønsted–Lowry dropped the need for water; Lewis (Reactivity 3.4) widened it to electron pairs.

Reactivity 3.1.2 Conjugate acid–base pairs

Two species that differ by exactly one proton form a conjugate acid–base pair. Removing H⁺ gives the conjugate base and lowers the charge by one; adding H⁺ gives the conjugate acid and raises it by one.

Worked example 2: the polyatomic anions

Give the conjugate acid of CO₃²⁻, PO₄³⁻, OH⁻ and CH₃COO⁻.

Add one H⁺ to each: HCO₃⁻, HPO₄²⁻, H₂O, CH₃COOH. Likewise SO₄²⁻ gives HSO₄⁻ and NO₃⁻ gives HNO₃.

Every acid–base equation contains two pairs. In example 1(b) they are HCOOH / HCOO⁻ and CH₃NH₃⁺ / CH₃NH₂.

Reactivity 3.1.3 Species that can be both

A species that can donate or accept a proton is amphiprotic. It needs a proton to lose and a lone pair to take one. Water, HCO₃⁻, HSO₄⁻ and H₂PO₄⁻ are examples.

Worked example 3: hydrogencarbonate both ways

(a) As an acid: HCO₃⁻(aq) + OH⁻(aq) → CO₃²⁻(aq) + H₂O(l)

(b) As a base: HCO₃⁻(aq) + H₃O⁺(aq) → H₂CO₃(aq) + H₂O(l). The H₂CO₃ breaks down to CO₂ and water, so hydrogencarbonates fizz with acids.

Water does this to itself: H₂O(l) + H₂O(l) ⇌ H₃O⁺(aq) + OH⁻(aq).

Oxides. Across period 3, oxides change from basic to acidic. Na₂O(s) + H₂O(l) → 2NaOH(aq) gives an alkali; SO₃(g) + H₂O(l) → H₂SO₄(aq) gives an acid. Aluminium oxide reacts with both acids and bases (amphoteric). Oxides of nitrogen and sulfur from burning fuels form nitric and sulfuric acids in clouds, causing acid rain.

Reactivity 3.1.4 Measuring acidity: the pH scale

• pH = −log₁₀[H⁺]
• [H⁺] = 10⁻ᵖᴴ

The guide says these equations are in the data booklet. One pH unit is a tenfold change in [H⁺], and lower pH means more H⁺. A sketch of pH against [H⁺] is a curve, not a line: steep at low [H⁺], flattening as [H⁺] rises.

Worked example 4: both directions

(a) [H⁺] = 0.0045 mol dm⁻³: pH = −log₁₀(0.0045) = 2.35

(b) pH 3.60: [H⁺] = 10⁻³·⁶⁰ = 2.5 × 10⁻⁴ mol dm⁻³

(c) Hydroiodic acid, 0.040 mol dm⁻³, is fully ionized: pH = 1.40. Dilute it tenfold and pH = 2.40. Each tenfold dilution of a strong acid raises pH by one.

Measuring pH. Universal indicator gives an estimate from a colour chart (red strongly acidic, green near neutral, purple strongly alkaline). A pH meter or probe, calibrated with buffer solutions, gives a precise digital value and can log readings throughout a titration.

Reactivity 3.1.5 The ion product of water, Kw

Water ionizes slightly, so every aqueous solution holds both H⁺ and OH⁻:

Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 298 K

The guide says this equation and its 298 K value are in the data booklet. With the product fixed, more H⁺ means less OH⁻.

• Acidic: [H⁺] > [OH⁻] (pH below 7 at 298 K)
• Neutral: [H⁺] = [OH⁻] (pH 7 at 298 K)
• Basic: [H⁺] < [OH⁻] (pH above 7 at 298 K)

Worked example 5: pH of an alkali

0.020 mol dm⁻³ KOH at 298 K is fully ionized, so [OH⁻] = 0.020 mol dm⁻³.

[H⁺] = (1.0 × 10⁻¹⁴) / 0.020 = 5.0 × 10⁻¹³ mol dm⁻³, so pH = 12.30.

Temperature. Water’s ionization is endothermic, so heating shifts it right and Kw rises. Hot pure water has pH below 7 yet is still neutral, because [H⁺] still equals [OH⁻].

Reactivity 3.1.6 Strength of acids and bases

A strong acid or base ionizes completely in water; a weak one only partly, setting up an equilibrium.

• Strong acids listed in the guide: HCl, HBr, HI, HNO₃ and H₂SO₄. Strong bases: group 1 hydroxides.
• Weak acids include carboxylic acids and carbonic acid; weak bases include ammonia and amines.

Strength is not concentration. Concentration is how much is dissolved per dm³; strength is the fraction that ionizes. 5 mol dm⁻³ ethanoic acid is concentrated but weak; 0.001 mol dm⁻³ nitric acid is dilute but strong.

Worked example 6: how much is ionized?

0.20 mol dm⁻³ of a weak monoprotic acid HA has pH 2.70.

[H⁺] = 10⁻²·⁷⁰ = 2.0 × 10⁻³ mol dm⁻³

Percentage ionized = (2.0 × 10⁻³ / 0.20) × 100 = 1.0%. A strong acid of equal concentration would have pH 0.70.

In the lab, a strong acid at the same concentration shows a lower pH, higher conductivity and faster reaction with magnesium or a carbonate, but not more gas in total.

Direction of equilibrium. Acid–base equilibria lie towards the weaker acid and weaker base. For HCN(aq) + F⁻(aq) ⇌ CN⁻(aq) + HF(aq), given that HF is the stronger acid, CN⁻ is the stronger base, so the equilibrium lies to the left. A strong acid has a large ionization equilibrium constant; a weak acid’s is far below 1.

Hydrogen halides. Acid strength rises from HF to HI because the H–X bond gets longer and weaker down group 17. HF is weak; the other three are strong.

Reactivity 3.1.7 Neutralization

Base	Example equation
Metal oxide	2HNO₃(aq) + MgO(s) → Mg(NO₃)₂(aq) + H₂O(l)
Metal hydroxide	CH₃COOH(aq) + KOH(aq) → CH₃COOK(aq) + H₂O(l)
Carbonate	H₂SO₄(aq) + Na₂CO₃(aq) → Na₂SO₄(aq) + H₂O(l) + CO₂(g)
Hydrogencarbonate	HCOOH(aq) + NaHCO₃(aq) → HCOONa(aq) + H₂O(l) + CO₂(g)
Ammonia	H₂SO₄(aq) + 2NH₃(aq) → (NH₄)₂SO₄(aq)
Amine	HBr(aq) + CH₃CH₂NH₂(aq) → CH₃CH₂NH₃Br(aq)

Ammonia and amines just gain a proton, so no water forms. The ionic equation for a strong acid with a strong base is H⁺(aq) + OH⁻(aq) → H₂O(l). The reaction is exothermic because forming the O–H bond releases energy.

Worked example 7: parent acid and base

The anion of a salt comes from the acid; the cation comes from the base.

(a) Calcium nitrate: nitric acid and calcium hydroxide (or oxide or carbonate).

(b) Ammonium sulfate: sulfuric acid and ammonia.

(c) Sodium propanoate: propanoic acid and sodium hydroxide (or carbonate or hydrogencarbonate).

Soluble salts are recovered by evaporating some water and crystallizing; insoluble ones by filtration. A metal releasing hydrogen from an acid is a redox reaction (Reactivity 3.2), not neutralization.

Reactivity 3.1.8 pH curves for strong acid–strong base titrations

A pH curve plots pH against volume of titrant added. The guide limits these to monoprotic neutralizations.

Worked example 8: building the curve

20.00 cm³ of 0.0800 mol dm⁻³ HNO₃ is titrated with 0.100 mol dm⁻³ KOH.

Amount of acid = 0.0800 × 0.02000 = 1.60 × 10⁻³ mol, reacting 1 : 1, so the equivalence point is at 1.60 × 10⁻³ / 0.100 = 16.00 cm³.

KOH / cm³	0.00	8.00	15.90	16.00	16.10	30.00
pH	1.10	1.54	3.56	7.00	10.44	12.45

Check: at 15.90 cm³, excess H⁺ = 1.0 × 10⁻⁵ mol in 35.90 cm³, so [H⁺] = 2.8 × 10⁻⁴ mol dm⁻³ and pH = 3.56.

Features to sketch and interpret:

• pH-axis intercept: the starting pH (1.10). For a strong acid it gives the concentration: 10⁻¹·¹⁰ ≈ 0.080 mol dm⁻³.
• A slow rise while acid is in excess.
• A near-vertical section: pH jumps from about 3.6 to 10.4 for 0.20 cm³ of KOH.
• Equivalence point at pH 7 (298 K), mid-way up the vertical section. It is also called the stoichiometric point because acid and base are present in the ratio of the equation.
• The curve then flattens towards the pH of the KOH (13.0) without reaching it.

Adding acid to base gives the mirror image: high start, falling curve.

Common errors

• Calling OH⁻ the conjugate acid of water. It is water’s conjugate base; H₃O⁺ is its conjugate acid.
• Not changing the charge when adding or removing a proton.
• Treating “weak” and “dilute” as the same.
• Defining neutral as pH 7 at any temperature.
• Putting water in the products of acid + ammonia.
• Reading the equivalence point at the top of the vertical section, not its midpoint.

Official syllabus

This unit is drawn from Reactivity 3.1 (understandings 3.1.1 to 3.1.8) of the Diploma Programme Chemistry guide, first examined in 2025 and issued by the International Baccalaureate Organization in February 2023.