OCR GCSE Chemistry A: Types of chemical reactions (J248)

Sub-topic C3.3 Types of chemical reactions is part of Topic C3 (Chemical reactions) in the OCR GCSE (9–1) Chemistry A (Gateway Science) J248 specification (Version 4.0, August 2026). Here it is taught in full: outcomes C3.3a to C3.3k, the mathematical learning outcome CM3.3i, and the practical suggestions linked to PAG C6 (titration) and PAG C7 (production of salts). Outcomes C3.3b, C3.3g, C3.3i and C3.3j are Higher Tier only; everything else is assessed at both tiers. The assessment overview places Topic C3 on Paper 1 (Foundation Tier) and Paper 3 (Higher Tier).

Ionic and half equations were introduced in Introducing chemical reactions. Afterwards, condense this guide with the Types of chemical reactions revision notes and test yourself with the Types of chemical reactions practice questions. See the OCR GCSE Chemistry hub, tick statements off on the printable J248 checklist, and try a free 10-minute diagnostic.

Coverage at a glance

Statement	Skill	Tier
C3.3a	Define oxidation and reduction by oxygen gained or lost; name the oxidising and reducing agents	Both tiers
C3.3b	Explain oxidation and reduction as loss or gain of electrons	Higher Tier only
C3.3c	Recall that acids form H⁺ ions in water and alkalis contain OH⁻ ions	Both tiers
C3.3d	Describe neutralisation: acid + alkali or base → salt + water	Both tiers
C3.3e	Write aqueous neutralisation as H⁺ + OH⁻ → H₂O	Both tiers
C3.3f	Recall acid reactions with carbonates and some metals; write balanced equations	Both tiers
C3.3g	Use dilute/concentrated and weak/strong correctly	Higher Tier only
C3.3h	Recall that pH measures relative acidity and alkalinity	Both tiers
C3.3i	Link pH to hydrogen ion concentration; read pH titration curves	Higher Tier only
C3.3j	Ten times more H⁺ means the pH is one lower	Higher Tier only
C3.3k	Describe how to measure pH with universal indicator, pH meters and probes	Both tiers
CM3.3i	Arithmetic, ratio, percentage and multistep calculations (M1a, M1c, M1d)	Both tiers

Oxidation and reduction by oxygen (C3.3a)

• Oxidation is the gain of oxygen.
• Reduction is the loss of oxygen.

The two always happen together, in a redox reaction.

Two more terms:

• An oxidising agent gives oxygen to another substance. It oxidises the other substance and is itself reduced.
• A reducing agent takes oxygen from another substance. It reduces the other substance and is itself oxidised.

Worked example 1: extracting tin

Heating tin(IV) oxide with carbon gives tin:

SnO₂ + 2C → Sn + 2CO

1. Tin(IV) oxide, SnO₂, ends up as tin, Sn. It has lost oxygen, so tin(IV) oxide is reduced.
2. Carbon, C, ends up as carbon monoxide, CO. It has gained oxygen, so carbon is oxidised.
3. Carbon took the oxygen, so it is the reducing agent; tin(IV) oxide supplied it, so it is the oxidising agent.

Name the whole substance: “tin(IV) oxide is reduced”, not “tin is reduced”.

Oxidation and reduction by electrons (C3.3b, Higher Tier only)

Everything here is Higher Tier only.

Many redox reactions involve no oxygen, so a wider definition uses electrons:

• Oxidation: electrons are lost.
• Reduction: electrons are gained.

OIL RIG (Oxidation Is Loss, Reduction Is Gain) helps you remember which is which. In electron terms, an oxidising agent accepts electrons (and is reduced), and a reducing agent gives electrons (and is oxidised).

Worked example 2: zinc in silver nitrate solution

Zn + 2Ag⁺ → Zn²⁺ + 2Ag

Split the ionic equation into two half equations:

• Zn → Zn²⁺ + 2e⁻. Each zinc atom loses two electrons, so zinc is oxidised. Zinc is the reducing agent.
• 2Ag⁺ + 2e⁻ → 2Ag. Each silver ion gains one electron, so silver ions are reduced. Silver ions are the oxidising agent.

Check: electrons lost (2) = electrons gained (2).

The same idea covers a non-metal displacement such as Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂. Chlorine gains electrons (reduced, the oxidising agent). Bromide ions lose electrons (oxidised).

Acids and alkalis in water (C3.3c)

• Acids form hydrogen ions, H⁺, when they dissolve in water.

• Solutions of alkalis contain hydroxide ions, OH⁻.

• HNO₃ → H⁺ + NO₃⁻

• H₂SO₄ → 2H⁺ + SO₄²⁻

• KOH → K⁺ + OH⁻

The specification lists two misconceptions here. First, the H⁺ ions do not stay part of the acid molecule: in solution they are separate and move freely. Second, alkalis are not less corrosive than acids: concentrated sodium hydroxide solution damages skin and eyes badly. Wear eye protection with both.

Neutralisation (C3.3d and C3.3e)

Neutralisation is an acid reacting with an alkali or a base to form a salt plus water. A base is a substance that neutralises an acid, such as a metal oxide or metal hydroxide. An alkali is a soluble base.

• acid + alkali: KOH + HNO₃ → KNO₃ + H₂O
• acid + base: MgO + H₂SO₄ → MgSO₄ + H₂O

The metal gives the first part of the salt’s name; the acid gives the second:

Acid	Salt ending	Example salt
Hydrochloric acid, HCl	chloride	potassium chloride, KCl
Nitric acid, HNO₃	nitrate	potassium nitrate, KNO₃
Sulfuric acid, H₂SO₄	sulfate	magnesium sulfate, MgSO₄

The ionic equation

In any aqueous neutralisation, the reaction that actually happens is:

H⁺(aq) + OH⁻(aq) → H₂O(l)

In KOH + HNO₃, the K⁺ and NO₃⁻ ions are unchanged, so they are left out. Every acid–alkali neutralisation in water reduces to this one equation.

Acids with carbonates and metals (C3.3f)

Two general patterns:

• acid + metal carbonate → salt + water + carbon dioxide
• acid + metal → salt + hydrogen (only some metals; copper, silver and gold do not react with dilute hydrochloric or sulfuric acid)

Worked example 3: writing balanced equations

(a) Sodium carbonate with nitric acid.

1. Products: sodium nitrate, water, carbon dioxide.
2. Formulae: Na₂CO₃ + HNO₃ → NaNO₃ + H₂O + CO₂
3. Two Na on the left, so put 2 in front of NaNO₃. That needs two NO₃, so put 2 in front of HNO₃.
4. Na₂CO₃ + 2HNO₃ → 2NaNO₃ + H₂O + CO₂. Check H: 2 = 2. Check O: 3 + 6 = 6 + 1 + 2 = 9.

(b) Iron with dilute sulfuric acid, forming iron(II) sulfate.

Fe + H₂SO₄ → FeSO₄ + H₂, balanced as written.

You will see fizzing in both.

Making a pure, dry salt (PAG C7)

The specification’s practical suggestion for C3.3d is the production of a pure dry sample of a salt (PAG C7). Here is the method for magnesium sulfate from magnesium oxide and dilute sulfuric acid:

1. Gently warm some dilute sulfuric acid in a beaker.
2. Add magnesium oxide powder a spatula at a time, stirring, until some solid stays undissolved. The excess base makes sure all the acid has reacted.
3. Filter to remove the leftover magnesium oxide. The filtrate is magnesium sulfate solution.
4. Boil away part of the water from the filtrate in an evaporating basin, stopping when crystals begin to appear at the edge.
5. Leave it to cool so crystals grow. Filter them off and dry them between sheets of filter paper or in a warm oven.

Excess solid can be filtered off; excess acid would stay in the solution and contaminate the salt.

Dilute, concentrated, weak and strong (C3.3g, Higher Tier only)

Everything here is Higher Tier only. The specification flags a misconception: thinking strength and concentration mean the same thing.

• Concentrated and dilute describe the amount of substance of acid in a given volume of solution (the ratio of amount of acid to volume). Concentrated means a large amount per dm³; dilute means a small amount.
• Strong and weak describe the degree of ionisation. Hydrochloric, nitric and sulfuric acids are strong: in water every molecule ionises (complete ionisation). Ethanoic, citric and carbonic acids are weak: only some of their molecules ionise (partial ionisation).

A useful model: picture 400 molecules of a weak acid in water. If only 4 of them have split into ions, the degree of ionisation is 4 ÷ 400 × 100 = 1%. For a strong acid, every one of the 400 is ionised. So you can have a dilute strong acid or a concentrated weak acid.

Worked example 4: comparing concentrations (CM3.3i)

Solution A has 0.40 mol of hydrochloric acid in 500 cm³. Solution B has 0.15 mol in 100 cm³. Which is more concentrated?

1. Convert volumes to dm³: 500 cm³ = 0.500 dm³; 100 cm³ = 0.100 dm³.
2. A: 0.40 ÷ 0.500 = 0.80 mol/dm³.
3. B: 0.15 ÷ 0.100 = 1.5 mol/dm³.
4. B is more concentrated, even though it contains less acid in total.

Estimate first (M1d): B has over a third as much acid in a fifth of the volume, so it should be more concentrated.

pH and hydrogen ion concentration (C3.3h, C3.3i, C3.3j)

C3.3h (both tiers): relative acidity and alkalinity are measured by pH. The scale usually runs from 0 to 14. pH 7 is neutral. Below 7 is acidic: the lower the number, the more acidic. Above 7 is alkaline: the higher the number, the more alkaline.

C3.3i and C3.3j are Higher Tier only. A neutral solution has equal amounts of H⁺ and OH⁻. An acidic solution has more H⁺ than OH⁻. The higher the hydrogen ion concentration, the lower the pH. The specification works in whole-number pH values, and the rule is:

Each time the H⁺ concentration increases by a factor of 10, the pH decreases by 1.

Worked example 5: factor of ten (Higher Tier only)

(a) A solution changes from pH 4 to pH 1. The pH has dropped by 3, so the H⁺ concentration has increased by 10 × 10 × 10 = 1000 times.

(b) 20 cm³ of a strong acid at pH 2 is diluted with water to 2000 cm³. The volume is 100 times larger, so the H⁺ concentration is 100 times smaller. 100 = 10 × 10, so the pH rises by 2, to pH 4.

pH titration curves (Higher Tier only)

When an alkali is added from a burette to an acid while a pH probe records the pH, the readings make a titration curve. These values are for 25.0 cm³ of a strong acid with a strong alkali of equal concentration:

Alkali added (cm³)	0	10	20	24	24.9	25.0	25.1	26	30	50
pH	1	1	2	3	4	7	10	11	12	13

The pH rises slowly at first, because plenty of H⁺ remains. Near exact neutralisation it jumps steeply through pH 7 over a tiny volume. Then excess OH⁻ makes the solution alkaline and the pH levels off high. The middle of the steep part marks the volume that neutralises the acid, 25.0 cm³ here.

Measuring pH (C3.3k, PAG C6)

The specification lists determining the pH of unknown solutions and use of pH probes as PAG C6 practical suggestions.

• Universal indicator (solution or paper): add a few drops, or dip the paper, and match the colour against the chart. Roughly: red for strong acid, orange and yellow for weaker acid, green at 7, blue for weak alkali, purple for strong alkali. Colours vary a little between brands, so always use the chart that comes with it. It gives only an approximate whole-number pH, and the colour judgement is subjective.
• pH meter or pH probe: calibrate it in buffer solutions of known pH, rinse the probe with distilled water between samples, dip it in and read the display. It gives a more precise value, works in coloured solutions, and with a data logger can record pH continuously during a titration.

Common errors

• Calling the oxidising agent “oxidised”. The oxidising agent gets reduced.
• Naming “tin” instead of “tin(IV) oxide” as the substance reduced.
• (Higher Tier only) Writing Zn + 2e⁻ → Zn²⁺. Electrons lost appear on the right.
• Leaving water or carbon dioxide out of acid + carbonate products.
• Using “weak” to mean “dilute” (Higher Tier only).
• Saying a fall of 2 in pH means twice as much H⁺. It is 100 times (Higher Tier only).

Official syllabus

OCR GCSE (9–1) Chemistry A (Gateway Science) J248 specification, Version 4.0 (August 2026), for first assessment in 2018, published by OCR. Sub-topic C3.3 Types of chemical reactions, learning outcomes C3.3a–C3.3k, mathematical learning outcome CM3.3i and practical suggestions (PAG C6, PAG C7).