OCR GCSE Chemistry A: Predicting chemical reactions (J248) – Revision Notes

For full explanations and longer worked examples, use the Predicting chemical reactions study guide.

These notes cover sub-topic C4.1 Predicting chemical reactions of the OCR GCSE (9–1) Chemistry A (Gateway Science) J248 specification (Version 4.0, August 2026): learning outcomes C4.1a–C4.1f, mathematical learning outcome CM4.1i and the linked practical suggestions (PAG C1, C5, C7, C8). Every C4.1 statement is assessed at both tiers; none is Higher Tier only. C4.1c is marked in the specification as separate-chemistry content. Topic C4 is examined on Paper 2 (Foundation Tier) and Paper 4 (Higher Tier).

When you are ready, try the practice questions. The OCR GCSE Chemistry hub and the J248 checklist show where C4.1 fits, and a free 10-minute diagnostic helps you decide what to revise next.

The one idea behind C4.1

Same group → same number of outer electrons → similar reactions. Down a group → one extra shell each time → outer electrons sit further out, behind more inner shells.

• Losing electrons (metals) gets easier down a group → reactivity rises.
• Gaining electrons (non-metals) gets harder down a group → reactivity falls.
• Full outer shell (Group 0) → no tendency to react.

Group 1 summary (C4.1a, C4.1b)

Feature	Detail
Outer electrons	1; forms a 1+ ion
Physical	soft, low density, low melting point; melting point falls down the group
Storage	under oil (reacts with air and water vapour)
+ water	metal hydroxide + hydrogen; alkaline solution
+ oxygen	metal oxide, e.g. 4Na + O₂ → 2Na₂O
+ chlorine	white metal chloride, e.g. 2K + Cl₂ → 2KCl
Reactivity	increases down the group

Water observations: lithium fizzes steadily; sodium melts into a ball and whizzes around; potassium gives a lilac flame.

Why reactivity rises: outer electron further from the nucleus + more shielding → weaker attraction → electron lost more easily.

Group 7 summary (C4.1a, C4.1b)

Halogen	Room temperature	In water
Chlorine, Cl₂	pale green gas	very pale green
Bromine, Br₂	red-brown liquid	orange
Iodine, I₂	grey-black solid	brown

• Diatomic molecules. Melting and boiling points rise down the group (bigger molecules, stronger forces between them).
• React with metals to give halide salts; with hydrogen to give hydrogen halides (acidic in water).
• Reactivity decreases down the group: an extra electron is attracted less strongly by a nucleus that is further away and more shielded.

Displacement (PAG C1): a halogen pushes a lower halogen out of its halide salt solution.

Example: Cl₂ + 2NaBr → 2NaCl + Br₂ (colourless solution turns orange).

Added → to solution ↓	chlorine	bromine	iodine
chloride	–	no change	no change
bromide	orange	–	no change
iodide	brown	brown	–

Group 0 summary (C4.1a, C4.1b)

• Colourless, monatomic gases; very unreactive.
• Full outer shell: 2 electrons (helium) or 8 (the rest).
• Boiling point and density increase down the group.

Transition metals (C4.1c)

Typical examples: iron, copper, nickel, manganese, chromium.

Property	Example
High melting point	iron 1538 °C
High density	copper 8.96 g/cm³
Low reactivity compared with Group 1	copper does not react with water or dilute acid
Ions with different charges	Fe²⁺ / Fe³⁺; Cu⁺ / Cu²⁺
Coloured compounds	Cu²⁺ blue; Fe²⁺ pale green; Fe³⁺ orange-brown
Catalysts	iron (Haber process); nickel (adding hydrogen to alkenes); manganese(IV) oxide (decomposing hydrogen peroxide)

Predicting from position (C4.1d)

Position	Likely behaviour
Group 1 or 2, left side	metal; forms positive ions; basic oxide; may react with water
Central block	transition metal; coloured compounds; variable charge
Groups 6 and 7, right side	non-metal; forms negative ions with metals; acidic oxide (Group 6)
Group 0	no reaction expected

A reaction is likely when a metal meets a non-metal, or when a more reactive element meets a compound of a less reactive one in the same group or series.

Reactivity of metals (C4.1e, C4.1f)

A useful order, most reactive first: potassium, sodium, lithium, calcium, magnesium, zinc, iron, copper, silver.

• metal + water → metal hydroxide + hydrogen
• metal + dilute acid → salt + hydrogen, e.g. Zn + 2HCl → ZnCl₂ + H₂
• The more easily the metal forms its positive ion (Zn → Zn²⁺ + 2e⁻), the faster the reaction.

Method in steps: deducing an order from results

1. For a displacement grid, count the reactions for each metal. Most reactions = most reactive.
2. For temperature data, subtract the start temperature from each final value. Biggest rise = most reactive (same mass of metal, same volume and concentration of solution).
3. For gas collection, divide volume by time to get a mean rate in cm³/s. Fastest = most reactive.
4. A metal that does not react with a salt solution is less reactive than the metal in the salt.
5. Check that your order is consistent with every result, not just one.

Worked reminder: sodium in water gives 45 cm³ of hydrogen in 15 s. Mean rate = 45 ÷ 15 = 3.0 cm³/s.

Practical links

PAG	C4.1 activity	Key points
C1	halogen displacement; metals with water and acid; metal displacement	small amounts, careful mixing, record colour changes
C5	investigating transition metal compounds	identify coloured ions; compare with Group 1 compounds
C7	metal and salt reactions	the displacement product is a new salt solution plus a metal
C8	gas volume or temperature change	same mass and surface area of metal; same volume and concentration of solution; repeat and average

Safety: Group 1 metals are handled with tongs, in small pieces, behind a safety screen. Halogens are toxic, so use dilute solutions in a well-ventilated room.

Empirical formulae and balancing (CM4.1i)

Method in steps

1. Mass of the added element = product mass − metal mass.
2. Moles of each element = mass ÷ Ar.
3. Divide both by the smaller number to get the whole-number ratio.

Worked reminder: 1.10 g of manganese gives 1.74 g of an oxide. Oxygen = 0.64 g. Mn: 1.10 ÷ 55 = 0.020 mol; O: 0.64 ÷ 16 = 0.040 mol. Ratio 1 : 2, so MnO₂ (manganese(IV) oxide).

Balancing with ratio: 4Li + O₂ → 2Li₂O. 0.35 g of lithium is 0.050 mol, which makes 0.025 mol of Li₂O (Mr 30), a mass of 0.75 g. The oxide is heavier than the lithium because oxygen has been added.

Must-know distinctions

• Group 1 vs Group 7 trend: reactivity rises down Group 1 but falls down Group 7.
• Chlorine vs chloride: chlorine is the element (Cl₂); chloride is the ion (Cl⁻) in a salt.
• Group 1 vs transition metals: low vs high density and melting point; one ion vs ions with different charges; white vs coloured compounds.
• Displacement vs no reaction: colour change or temperature rise means a reaction; no change means the added element is less reactive.

Quick self-test

1. State the number of outer electrons in a Group 7 atom.
2. Name the gas made when lithium reacts with water.
3. Write a balanced equation for sodium reacting with chlorine.
4. Explain why argon is unreactive.
5. Predict the state of iodine at room temperature and give its colour.
6. Will iodine solution react with potassium bromide solution? Explain.
7. Give two properties that make iron different from potassium.
8. Name the transition metal catalyst used in the Haber process.
9. 1.17 g of potassium reacts with bromine to give 3.57 g of product. Find the empirical formula. (Ar: K = 39, Br = 80)
10. Metal Q displaces copper from copper(II) sulfate but not iron from iron(II) sulfate. Place Q in the series iron, copper.
11. Why does potassium react faster with water than lithium?
12. Copper(II) chloride solution is blue. What property of transition metals does this show?

Answers

1. Seven.
2. Hydrogen.
3. 2Na + Cl₂ → 2NaCl.
4. Its atoms have a full outer shell (stable), so they do not lose, gain or share electrons.
5. Solid; grey-black.
6. No reaction: bromine is the more reactive halogen, so iodine cannot take its place.
7. Any two: higher melting point; higher density; much less reactive; forms ions with different charges (Fe²⁺ and Fe³⁺); coloured compounds.
8. Iron.
9. Br = 3.57 − 1.17 = 2.40 g. K: 1.17 ÷ 39 = 0.030 mol; Br: 2.40 ÷ 80 = 0.030 mol. Ratio 1 : 1, so KBr.
10. Between them: iron > Q > copper.
11. Potassium has two more occupied shells than lithium, so its outer electron sits further out, is shielded more and is lost more readily.
12. Transition metal compounds are often coloured.

Where marks are usually lost

• Saying halogens get more reactive down the group because they “have more shells”.
• Explaining reactivity with “wants a full shell” instead of attraction between the nucleus and the outer electron.
• Writing Na₂ or Cl in equations; sodium is Na, chlorine is Cl₂.
• Describing potassium’s flame as “orange” or “yellow”; it is lilac.
• Calling bromine a gas or a solid at room temperature.
• Forgetting that the added metal must be more reactive for displacement to occur.
• Using a final temperature instead of the temperature change to rank metals.
• Dividing masses rather than moles when finding an empirical formula.
• Saying a metal “loses mass” when it forms an oxide or chloride.

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 C4.1 Predicting chemical reactions, learning outcomes C4.1a–C4.1f, mathematical learning outcome CM4.1i and the linked practical suggestions (PAG C1, C5, C7, C8).