Speca › Topics › GCSE Combined Science

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• Assessment mark schemeWord, editable
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Foundation and Higher tier

• Explain that metals react with oxygen to form metal oxides, and that this is oxidation because the metal gains oxygen, with the reverse (loss of oxygen) being reduction.
• Recall and describe the reactions (if any) of potassium, sodium, lithium, calcium, magnesium, zinc, iron, and copper with water and with dilute acids (limited to room temperature, not steam), and place these metals, plus hydrogen and carbon, often included for comparison, in a reactivity series, explaining that a metal's reactivity relates to how readily it forms positive ions, and deducing an order of reactivity from experimental results; know that a more reactive metal can displace a less reactive one from a compound.
• Explain that unreactive metals (e.g. gold) occur naturally as the metal itself, while most metals occur as compounds requiring chemical extraction, and that metals less reactive than carbon can be extracted from their oxides by reduction with carbon (reduction = loss of oxygen): interpreting or evaluating a given metal-extraction process and identifying what's oxidised/reduced in terms of oxygen gain or loss. The detailed extraction processes themselves are not required.
• Describe how acids react with some metals to produce a salt and hydrogen (reactions limited to magnesium, zinc, and iron with hydrochloric and sulfuric acid).
• Explain that acids are neutralised by alkalis (soluble metal hydroxides) and bases (insoluble metal hydroxides/oxides) to form a salt and water, and by metal carbonates to form a salt, water, and carbon dioxide, and that the salt produced depends on the acid used (hydrochloric → chlorides, nitric → nitrates, sulfuric → sulfates) and the positive ion in the base/alkali/carbonate; predict the products of a given acid reaction and use common ion formulae to deduce a salt's formula.
• Describe how to prepare a pure, dry sample of a named soluble salt: adding an insoluble solid (metal, metal oxide, hydroxide, or carbonate) to an acid until no more reacts, filtering off the excess solid, then crystallising the resulting salt solution.
• Explain that acids produce hydrogen ions (H⁺) and alkalis produce hydroxide ions (OH⁻) in aqueous solution, describe the pH scale (0–14, measured with universal indicator or a pH probe) with pH 7 neutral, acidic solutions below 7, and alkaline solutions above 7, and describe neutralisation as hydrogen ions reacting with hydroxide ions to form water.
• Describe electrolysis: in a molten or dissolved ionic compound, ions become free to move, making the liquid or solution an electrolyte able to conduct electricity; passing a current makes positive ions move to the negative electrode (cathode) and negative ions move to the positive electrode (anode), where they're discharged to form elements.
• Predict the products of electrolysing a molten binary ionic compound (e.g. lead bromide) with inert electrodes: metal at the cathode, non-metal at the anode.
• Explain why electrolysis, rather than reduction with carbon, is used to extract metals too reactive for carbon reduction or that react with carbon, using aluminium's extraction (electrolysis of molten aluminium oxide and cryolite, with a carbon anode that must be continually replaced) as the worked example, and explain why a mixture (rather than pure aluminium oxide) is used as the electrolyte.
• Predict the products of electrolysing an aqueous solution containing a single ionic compound with inert electrodes: hydrogen at the cathode if the metal is more reactive than hydrogen, and oxygen at the anode unless halide ions are present (in which case the halogen forms instead), explained by water molecules themselves breaking down to supply hydrogen and hydroxide ions available for discharge.

Higher tier only

• Define oxidation as loss of electrons and reduction as gain of electrons; write ionic equations for displacement reactions, and identify which species are oxidised and which are reduced in a given reaction, symbol equation, or half equation.
• Explain acid-metal reactions as redox reactions in terms of electron loss/gain, identifying the oxidised and reduced species in given equations.
• Distinguish strong acids (completely ionised in solution, e.g. hydrochloric, nitric, sulfuric) from weak acids (only partially ionised, e.g. ethanoic, citric, carbonic), and dilute from concentrated (in terms of amount of substance present); explain that for a given concentration, a stronger acid has a lower pH, and that each one-unit drop in pH corresponds to a tenfold increase in hydrogen ion concentration; describe neutrality and relative acidity in terms of hydrogen ion concentration and whole-number pH values.
• Write half equations for the reactions occurring at each electrode during electrolysis (throughout this topic's electrolysis content), and complete/balance supplied half equations: recognising that at the cathode, positive ions gain electrons (reduction), and at the anode, negative ions lose electrons (oxidation).

Required practicals

• Required practical 8: prepare a pure, dry sample of a soluble salt from an insoluble oxide or carbonate, heating dilute acid with a Bunsen burner and evaporating the resulting solution using a water bath or electric heater.
• Required practical 9: investigate, by developing and testing a hypothesis, what happens when aqueous solutions are electrolysed using inert electrodes.

Where students go wrong

• Confusing oxidation and reduction, particularly forgetting which electrode (or which direction of electron/oxygen transfer) each one corresponds to.
• Assuming "strong acid" means "concentrated acid" (Higher tier): these describe two independent properties: strength is about the degree of ionisation, concentration is about how much acid is dissolved in a given volume.
• Believing pH is a linear scale in terms of hydrogen ion concentration, rather than understanding (Higher tier) each pH unit represents a tenfold change in concentration.
• Mixing up which electrode is the cathode and which is the anode, and consequently which ion type (positive or negative) moves to which.
• Assuming the products of electrolysing an aqueous solution are the same as electrolysing the molten compound, rather than understanding water itself contributes ions that compete with the dissolved compound's own ions.

How it gets asked in the exam

"Explain why [metal] reacts/doesn't react with...", "Predict the products of the electrolysis of...", "Describe how you would prepare a pure, dry sample of...", "Use the pH scale to identify...", "Write a half equation for the reaction at the [electrode]" (Higher), "Explain, in terms of electrons, why this is a redox reaction" (Higher).

Key vocabulary

Oxidation, reduction, reactivity series, displacement, extraction, salt, neutralisation, pH scale, alkali, base, electrolysis, electrolyte, electrode, cathode, anode, ion, half equation (HT), strong acid (HT), weak acid (HT).

Assumed prior knowledge

• Ionic bonding and ion charge from Bonding, Structure and the Properties of Matter (this taxonomy's topic 09): needed throughout electrolysis and salt-formula content.
• Balanced symbol equations and relative formula mass from Quantitative Chemistry (this taxonomy's topic 10).
• KS3-level familiarity with the idea that acids and alkalis exist and can be identified with indicators.

How Speca scaffolds this topic

• The reactivity series benefits from a single, consistently ordered reference strip (potassium down to copper, hydrogen and carbon marked in their comparative positions) used throughout every reactivity, extraction, and displacement example in this topic, rather than re-presented differently each time.
• A shared "oxidation and reduction, two ways of describing it" reference card, oxygen-based definition alongside the electron-based one (Higher), makes explicit that these are two lenses on the same idea, not two unrelated topics to learn separately.
• pH work benefits from a fixed, always-visible 0–14 number line with colour bands matching universal indicator, used consistently across every acid/alkali example, so pH values are always anchored to the same visual reference.
• For electrolysis, a consistent labelled diagram template (electrolyte, two labelled electrodes, direction-of-ion-movement arrows) applied to every worked example (molten binary compound, aluminium extraction, aqueous solution) helps students transfer the same underlying process across genuinely different specific cases.

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