OCR GCSE Chemistry A: Predicting chemical reactions (J248)

Sub-topic C4.1 Predicting chemical reactions opens Topic C4 of the OCR GCSE (9–1) Chemistry A (Gateway Science) J248 specification (Version 4.0, August 2026). This guide teaches all of it: learning outcomes C4.1a to C4.1f, the mathematical learning outcome CM4.1i, and the practical suggestions linked to PAG C1, C5, C7 and C8. None of the C4.1 statements is Higher Tier only, so everything here is assessed at both tiers. The specification marks C4.1c (transition metals) as content found only in the separate chemistry qualification, which is the one you are taking. In the assessment overview, Topic C4 is examined on Paper 2 (Foundation Tier) and Paper 4 (Higher Tier).

Afterwards, condense it with the revision notes then try the practice questions. The OCR GCSE Chemistry hub lists every J248 page; tick off statements on the printable J248 checklist. A free 10-minute diagnostic shows where your chemistry is weakest.

What this sub-topic covers

Reference	Skill	Tier
C4.1a	Recall the simple physical and chemical properties of Groups 1, 7 and 0	Both tiers
C4.1b	Explain those properties using the outer shell of electrons (ease of losing or gaining electrons), and predict properties from trends down a group	Both tiers
C4.1c	Recall the general properties of transition metals and their compounds, using a few examples	Both tiers (chemistry only)
C4.1d	Predict possible reactions and probable reactivity from an element’s position in the Periodic Table	Both tiers
C4.1e	Link a metal’s reactivity with water or dilute acid to its tendency to form a positive ion	Both tiers
C4.1f	Deduce an order of reactivity of metals from experimental results	Both tiers
CM4.1i	Use arithmetic and ratio to find empirical formulae and balance equations	Both tiers

Background: atomic structure and bonding.

Why position predicts behaviour

Members of one group share an outer-electron count, so they react alike. Down a group, each element has one more occupied shell; that explains most trends here.

• Metals (left and centre of the table) react by losing electrons to form positive ions. Their oxides are basic; soluble ones make alkaline solutions.
• Non-metals (right of the table) react with metals by gaining electrons to form negative ions. Many non-metal oxides, such as carbon dioxide and sulfur dioxide, are acidic.
• Group 0 atoms start with a complete outer shell and have nothing to gain by reacting.

A single sodium atom is not soft or silvery; those are bulk properties of the metal. Explain with particles, not bulk properties.

Group 1: the alkali metals

Rubidium and caesium are too reactive for school use, so you predict their behaviour.

Physical properties. Group 1 metals are soft, shiny when freshly cut and tarnish quickly in air. They have low densities (lithium, sodium and potassium float on water) and low melting points, which decrease down the group.

Chemical properties. Each one is stored under oil because it reacts with oxygen and water vapour in the air.

Reaction	General equation	Example
With water	metal + water → metal hydroxide + hydrogen	2Li + 2H₂O → 2LiOH + H₂
With oxygen	metal + oxygen → metal oxide	4K + O₂ → 2K₂O
With chlorine	metal + chlorine → metal chloride (a white solid)	2Na + Cl₂ → 2NaCl

With water, lithium fizzes steadily, sodium melts into a ball and darts about, and potassium’s hydrogen ignites with a lilac flame. The hydroxide dissolves, leaving an alkaline solution.

Explaining the trend. All Group 1 atoms have one outer electron, which they lose to form a 1+ ion. Each step down adds a shell, so the outer electron sits further out with more inner shells screening it. The nucleus holds it less tightly, so it is lost more easily and the metal is more reactive.

Worked example 1: predicting for caesium

In a filmed demonstration, equal-sized pieces of Group 1 metals are dropped into water. Times to disappear: Li 75 s, Na 21 s, K 6 s, Rb 2 s. Predict the time for caesium and describe its reaction with water.

1. Each time is roughly a third of the one before (75 → 21 → 6 → 2 s).
2. Continuing the pattern predicts less than 1 s for caesium.
3. Caesium is lower in the group, so its outer electron is lost even more easily than rubidium’s. It reacts explosively with water.
4. Equation: 2Cs + 2H₂O → 2CsOH + H₂. The solution formed is alkaline.

A prediction longer than 2 s would break the trend.

Group 7: the halogens

The halogens exist as diatomic molecules (F₂, Cl₂, Br₂, I₂).

Halogen	State and colour at room temperature
Chlorine	pale green gas
Bromine	red-brown liquid (orange in water)
Iodine	grey-black solid (purple vapour; brown in water)

Melting and boiling points increase down the group, as bigger molecules attract each other more strongly. Colours darken down the group.

Chemical properties. Halogens react with metals to form ionic salts called halides, for example 2Fe + 3Cl₂ → 2FeCl₃. They react with hydrogen to form hydrogen halides such as HCl, which dissolve in water to give acidic solutions.

Explaining the trend. A halogen atom has seven outer electrons and gains one to form a 1− ion. Lower in the group the atom is bigger and its nucleus is screened by more inner shells, so it pulls an extra electron in less strongly. Electrons are gained less easily, so reactivity decreases down Group 7. This is the opposite direction to Group 1, because halogens gain electrons while alkali metals lose them.

Displacement reactions (PAG C1)

A halogen higher in the group pushes a lower halogen out of its salt solution, and you watch for a colour change.

Worked example 2: bromine water and sodium iodide

Orange bromine water is added to colourless sodium iodide solution. Predict what you would see, and write the equation.

1. Bromine is above iodine in Group 7, so bromine is more reactive.
2. Bromine displaces iodine: the solution turns brown as iodine forms.
3. Word equation: bromine + sodium iodide → sodium bromide + iodine.
4. Symbol equation: Br₂ + 2NaI → 2NaBr + I₂.

Bromine water added to sodium chloride solution stays orange: bromine is less reactive than chlorine.

Group 0: the noble gases

Helium, neon, argon, krypton and xenon are colourless gases at room temperature. They exist as single atoms (monatomic). Boiling points and densities increase down the group as the atoms get larger.

They are very unreactive because each atom has a stable full outer shell (two electrons for helium, eight for the others), so it does not lose, gain or share electrons, even with another atom of its own kind.

Transition metals (C4.1c)

The transition metals sit in the central block, between Groups 2 and 3. Iron, copper, nickel and manganese are good examples. Compared with Group 1 metals, they:

Property	Transition metals (iron, copper)	Group 1 (sodium)
Melting point	high (iron 1538 °C, copper 1085 °C)	low (98 °C)
Density	high (iron 7.87 g/cm³, copper 8.96 g/cm³)	low (0.97 g/cm³)
Reactivity	much less reactive; iron rusts slowly, copper does not react with water	reacts vigorously with cold water
Ions	ions with different charges: Fe²⁺ and Fe³⁺, Cu⁺ and Cu²⁺	one ion only: Na⁺
Colour of compounds	often coloured	white solids, colourless solutions
Catalysts	often used as catalysts (iron, nickel)	not typical catalysts

Coloured ions. Copper(II) sulfate solution is blue; iron(II) compounds are pale green and iron(III) compounds orange-brown. The Roman numeral gives the ion’s charge.

Catalysts. Iron is the catalyst in the Haber process for ammonia. Nickel catalyses the reaction of alkenes with hydrogen. Manganese(IV) oxide makes hydrogen peroxide break down faster. In PAG C5 and C8 work you can test whether a transition metal compound speeds up a reaction and identify transition metal ions.

Worked example 3: an empirical formula showing variable charge (CM4.1i)

A student heats 1.12 g of iron wool in chlorine. The product has a mass of 3.25 g. Find its empirical formula. (Ar: Fe = 56, Cl = 35.5)

1. Mass of chlorine combined = 3.25 − 1.12 = 2.13 g.
2. Moles of Fe = 1.12 ÷ 56 = 0.020 mol.
3. Moles of Cl = 2.13 ÷ 35.5 = 0.060 mol.
4. Divide by the smaller value: Fe 0.020 ÷ 0.020 = 1; Cl 0.060 ÷ 0.020 = 3.
5. Empirical formula FeCl₃, so the iron is present as Fe³⁺.
6. Balanced equation: 2Fe + 3Cl₂ → 2FeCl₃.

Iron also forms FeCl₂, so iron has ions with different charges. The product is heavier than the iron because chlorine has combined with it; likewise a metal oxide weighs more than the metal burned to make it.

Metals with water and dilute acid (C4.1e)

When a metal reacts with water or acid, its atoms lose electrons to form positive ions, for example Mg → Mg²⁺ + 2e⁻. The more easily it forms its positive ion, the more vigorous the reaction.

• metal + water → metal hydroxide + hydrogen (the most reactive metals)
• metal + dilute hydrochloric acid → metal chloride + hydrogen

A reactivity series covering the common laboratory metals, most reactive first: potassium, sodium, lithium, calcium, magnesium, zinc, iron, copper, silver. Copper and silver have so little tendency to form ions that they do not react with dilute hydrochloric acid at all.

Worked example 4: comparing rates of hydrogen production (PAG C8)

Calcium is added to cold water and 72 cm³ of hydrogen is collected in 40 s. Magnesium ribbon in cold water gives only about 3 cm³ in 10 minutes. Calculate the mean rate for each and explain the difference.

1. Calcium: 72 ÷ 40 = 1.8 cm³/s.
2. Magnesium: 10 minutes = 600 s; 3 ÷ 600 = 0.005 cm³/s.
3. Equation for calcium: Ca + 2H₂O → Ca(OH)₂ + H₂.
4. Calcium forms Ca²⁺ more easily than magnesium forms Mg²⁺, so reacts faster.

For a fair comparison, keep the volume and temperature of the liquid (and the acid’s concentration) the same and use similar-sized pieces of metal. Collect gas in a gas syringe or upturned measuring cylinder.

Deducing reactivity from experiments (C4.1f)

A metal higher in the series displaces a lower one from its dissolved salt. Zinc in copper(II) sulfate gives zinc sulfate and a brown coating of copper: Zn + CuSO₄ → ZnSO₄ + Cu. The blue colour fades as Cu²⁺ ions are removed.

Two kinds of result build an order:

1. Displacement grids (PAG C1, C7). Count how many salts each metal displaces. The metal that displaces the most is the most reactive.
2. Temperature rises (PAG C8). Displacement reactions are exothermic. Using the same mass of each metal powder and the same volume of a salt solution, a larger temperature rise means a more reactive metal.

Worked example 5: three unknown metals

Equal masses of metal powders X, Y and Z are added to 25 cm³ portions of copper(II) sulfate solution at 21.0 °C. Final temperatures: X 33.5 °C, Y 21.0 °C, Z 26.0 °C. Put X, Y, Z and copper in order of reactivity.

1. Rises: X 12.5 °C, Y 0.0 °C, Z 5.0 °C.
2. X and Z displaced copper, so both are more reactive than copper; X gave the larger rise.
3. Y gave no reaction, so Y is less reactive than copper.
4. Order, most reactive first: X > Z > copper > Y.

Repeat each test, and use a lidded polystyrene cup to reduce energy loss.

Common errors

• Saying Group 1 metals “gain a full shell” when they react; they lose their single outer electron.
• Explaining the halogen trend with “more shells, so more reactive”. For halogens it is the opposite.
• Writing Cl instead of Cl₂ in equations, then failing to balance.
• Giving a transition metal a single fixed charge.

Next steps

Use the revision notes for a quick summary and self-test, then the practice questions.

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).