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

These notes condense IB DP Chemistry Reactivity 3.1.1 to 3.1.8, “Proton transfer reactions”, for the final weeks before your exams. They are aligned to the IB Diploma Programme Chemistry guide (first assessment 2025; February 2023 edition). Everything here is for SL and HL; pOH, Ka, Kb, indicators and buffers come later in the AHL part of Reactivity 3.1 and are not included. The notes apply to sessions from 2025 onwards.

Every idea here is explained at length, with fuller worked examples, on the study guide page. After that, try the practice questions. The course hub, the syllabus checklist and our free diagnostics show you where to focus next.

Key definitions (Reactivity 3.1.1 to 3.1.3)

Term	Meaning
Brønsted–Lowry acid	Gives away H⁺ (a proton donor)
Brønsted–Lowry base	Takes H⁺ (a proton acceptor)
Alkali	A base that dissolves in water, releasing OH⁻(aq)
Conjugate acid–base pair	Two species whose formulas differ by a single H⁺, e.g. H₂PO₄⁻ / HPO₄²⁻
Amphiprotic	Able to act as a Brønsted–Lowry acid or base, e.g. H₂O, HSO₄⁻
Amphoteric oxide	Reacts with acids and with bases, e.g. Al₂O₃

H⁺(aq) and H₃O⁺(aq) both stand for a proton in water. H₃O⁺ is the oxonium ion: a water molecule that has used a lone pair to bond to H⁺.

Method in steps: naming acid, base and pairs

1. Compare each reactant with the products and find which one loses H⁺. That is the acid.
2. The reactant that gains H⁺ is the base.
3. Pair each reactant with the product that differs from it by one H⁺.
4. Check the charges: losing H⁺ drops the charge by one.

Reminder: in CH₃COOH(aq) + HPO₄²⁻(aq) ⇌ CH₃COO⁻(aq) + H₂PO₄⁻(aq), ethanoic acid is the acid and HPO₄²⁻ the base. Pairs: CH₃COOH / CH₃COO⁻ and H₂PO₄⁻ / HPO₄²⁻.

Amphiprotic species in action

HSO₄⁻ as an acid: HSO₄⁻(aq) + OH⁻(aq) → SO₄²⁻(aq) + H₂O(l)

HSO₄⁻ as a base: HSO₄⁻(aq) + H₃O⁺(aq) → H₂SO₄(aq) + H₂O(l)

Period 3 oxides: Na₂O and MgO are basic, Al₂O₃ is amphoteric, and the non-metal oxides P₄O₁₀, SO₂ and SO₃ are acidic. Nitrogen and sulfur oxides in the air dissolve in rainwater to give acid rain, e.g. SO₂(g) + H₂O(l) → H₂SO₃(aq).

Formulas (Reactivity 3.1.4 and 3.1.5)

Quantity	Formula	Note
pH	pH = −log₁₀[H⁺]	[H⁺] in mol dm⁻³
[H⁺] from pH	[H⁺] = 10⁻ᵖᴴ
Ion product of water	Kw = [H⁺][OH⁻]	1.0 × 10⁻¹⁴ at 298 K
[H⁺] from [OH⁻]	[H⁺] = Kw / [OH⁻]	for alkali solutions

The guide notes that the pH equations, the Kw equation and its 298 K value are in the data booklet.

Method in steps: from concentration to pH

1. Strong acid (monoprotic): [H⁺] = concentration of the acid.
2. Group 1 hydroxide: [OH⁻] = concentration of the hydroxide; then [H⁺] = Kw / [OH⁻].
3. pH = −log₁₀[H⁺].

Reminders:

• 0.0025 mol dm⁻³ HBr: pH = −log₁₀(0.0025) = 2.60
• 0.0040 mol dm⁻³ NaOH: [H⁺] = 1.0 × 10⁻¹⁴ / 0.0040 = 2.5 × 10⁻¹² mol dm⁻³, pH = 11.60
• pH 4.75: [H⁺] = 10⁻⁴·⁷⁵ = 1.8 × 10⁻⁵ mol dm⁻³

Reading the scale

• Each 1 unit of pH is a factor of 10 in [H⁺]. pH 2 has 1000 times the [H⁺] of pH 5.
• Diluting a strong acid by a factor of 10 raises pH by 1.
• Acidic: [H⁺] > [OH⁻]. Neutral: [H⁺] = [OH⁻]. Basic: [H⁺] < [OH⁻].
• Kw rises with temperature because water’s ionization is endothermic. Pure water above 298 K has pH below 7 but is still neutral.

Measuring pH

Method	What you get
Universal indicator paper or solution	An estimate from a colour chart; colours can be hard to judge
pH meter or probe	A precise digital value after calibration with buffers; can record continuously

A sketch of pH against [H⁺] is a falling curve, steep at first and then levelling off, because pH is logarithmic.

Strong and weak (Reactivity 3.1.6)

	Strong	Weak
Ionization in water	Complete	Partial; equilibrium
Acids	HCl, HBr, HI, HNO₃, H₂SO₄	Carboxylic acids, H₂CO₃, HF
Bases	Group 1 hydroxides	NH₃, amines
Equation arrow	→	⇌
Conjugate	Very weak conjugate base	Relatively stronger conjugate base

Must-know distinctions

• Strong vs concentrated. Strong = fully ionized. Concentrated = lots of solute per dm³. Any combination is possible.
• Base vs alkali. CuO and MgO are bases; only soluble bases giving OH⁻ are alkalis.
• Amphiprotic vs amphoteric. Amphiprotic is about donating or accepting a proton; amphoteric describes a substance, such as Al₂O₃, that reacts with both acids and bases.
• Equivalence point vs neutral. For a strong acid with a strong base, they coincide at pH 7 (298 K).

Equilibrium position: an acid–base equilibrium lies towards the weaker acid and weaker base. HI is a stronger acid than HF because the H–I bond is longer and weaker.

Experimental tests at equal concentration: the strong acid gives the lower pH, the higher conductivity and the faster reaction with Mg or CaCO₃. Total gas volume is the same.

Neutralization (Reactivity 3.1.7)

Reactants	Products
acid + metal oxide	salt + water
acid + metal hydroxide	salt + water
acid + carbonate	salt + water + CO₂
acid + hydrogencarbonate	salt + water + CO₂
acid + ammonia	ammonium salt only
acid + amine	alkylammonium salt only

Quick reminders:

• 2HCl(aq) + CaCO₃(s) → CaCl₂(aq) + H₂O(l) + CO₂(g)
• CH₃CH₂COOH(aq) + NaHCO₃(aq) → CH₃CH₂COONa(aq) + H₂O(l) + CO₂(g)
• HNO₃(aq) + NH₃(aq) → NH₄NO₃(aq)
• HCl(aq) + CH₃NH₂(aq) → CH₃NH₃Cl(aq)

Parent acid and base: anion from the acid, cation from the base. Potassium ethanoate comes from ethanoic acid and potassium hydroxide; lithium chloride from hydrochloric acid and lithium hydroxide.

Neutralization is exothermic: forming O–H bonds in water releases energy. Metal + acid giving hydrogen is redox, not neutralization.

pH curves (Reactivity 3.1.8)

Strong acid in the flask, strong base from the burette (monoprotic only):

1. Starts low: the pH-axis intercept is the pH of the acid alone.
2. Rises slowly while acid is in excess.
3. Near-vertical rise around the equivalence point, roughly pH 3 to 11.
4. Equivalence point at pH 7 (298 K), at the middle of the steep section.
5. Levels off just below the pH of the base being added.

Base in the flask gives the mirror image. The equivalence point is also called the stoichiometric point because the reactants are in their equation ratio. Its volume lets you find an unknown concentration: n = c × V for the titrant, then use the 1 : 1 ratio.

Reminder: if 18.75 cm³ of 0.0800 mol dm⁻³ NaOH reaches equivalence with 15.00 cm³ of HCl, then [HCl] = (0.0800 × 18.75) / 15.00 = 0.100 mol dm⁻³, and the curve’s starting pH is 1.00.

Quick self-test

1. State the conjugate base of H₂O and the conjugate acid of H₂O.
2. Give the conjugate acid of C₂H₅NH₂.
3. Write an equation for H₂PO₄⁻ acting as a base with H₃O⁺.
4. Calculate the pH of 0.0085 mol dm⁻³ HNO₃.
5. A solution has pH 8.20. Calculate [H⁺].
6. A solution at 298 K has [OH⁻] = 4.0 × 10⁻⁹ mol dm⁻³. Find [H⁺] and say whether it is acidic, neutral or basic.
7. Calculate the pH of 0.0016 mol dm⁻³ LiOH at 298 K.
8. Why is magnesium oxide a base but not an alkali?
9. Write the equation for propanoic acid reacting with sodium hydrogencarbonate.
10. Name the parent acid and base of potassium ethanoate.
11. On a strong acid–strong base pH curve, what does the pH-axis intercept tell you?

Answers

1. OH⁻ (conjugate base); H₃O⁺ (conjugate acid).
2. C₂H₅NH₃⁺.
3. H₂PO₄⁻(aq) + H₃O⁺(aq) → H₃PO₄(aq) + H₂O(l).
4. pH = −log₁₀(0.0085) = 2.07.
5. [H⁺] = 10⁻⁸·²⁰ = 6.3 × 10⁻⁹ mol dm⁻³.
6. [H⁺] = 1.0 × 10⁻¹⁴ / 4.0 × 10⁻⁹ = 2.5 × 10⁻⁶ mol dm⁻³. [H⁺] > [OH⁻], so acidic (pH 5.60).
7. [OH⁻] = 0.0016 mol dm⁻³, so [H⁺] = 1.0 × 10⁻¹⁴ / 0.0016 = 6.3 × 10⁻¹² mol dm⁻³; pH = 11.20.
8. It neutralizes acids (accepts protons) but is insoluble in water, so it does not give a solution of OH⁻ ions.
9. CH₃CH₂COOH(aq) + NaHCO₃(aq) → CH₃CH₂COONa(aq) + H₂O(l) + CO₂(g).
10. Ethanoic acid and potassium hydroxide.
11. The pH of the acid before any base is added; for a strong monoprotic acid, [acid] = 10⁻ᵖᴴ.

Where marks are usually lost

• Naming the acid and base but writing the conjugate pairs the wrong way round, or pairing species that differ by more than one proton.
• Dropping or mis-stating charges on conjugates, such as writing PO₄²⁻ for the conjugate base of HPO₄²⁻.
• Using pH = −log₁₀[OH⁻] for an alkali instead of first finding [H⁺] from Kw.
• Writing [H⁺] = 10ᵖᴴ, missing the minus sign.
• Saying a weak acid is a dilute acid, or the reverse.
• Claiming a strong acid gives more hydrogen or CO₂ than a weak acid of the same amount; it only gives it faster.
• Including water as a product when an acid reacts with ammonia or an amine.
• Forgetting CO₂ as a product with carbonates and hydrogencarbonates.
• Sketching a pH curve with a gentle slope at equivalence, or placing the equivalence point anywhere other than pH 7 for strong acid with strong base.
• Stating that neutral water always has pH 7, ignoring temperature.

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.