Chemical changes

1. Metals and oxygen: two useful reaction words

When magnesium burns, it gives a bright white light and forms white magnesium oxide. Metal + oxygen → metal oxide. For magnesium: 2Mg + O₂ → 2MgO.

Oxidation means gaining oxygen. Magnesium is oxidised when it forms MgO. Reduction means losing oxygen. Copper oxide is reduced when oxygen is removed to form copper.

A metal oxide is a compound, not a mixture of loose metal and oxygen. A chemical reaction is needed to remove its oxygen.

2. Which metals react most readily?

A reactive metal forms positive ions readily. Metal atoms form positive ions by losing electrons. The more readily this happens, the more reactive the metal.

Reactive metals with water form a metal hydroxide and hydrogen. For example: 2Na + 2H₂O → 2NaOH + H₂. These examples do not require reactions with steam. Alkali metals and calcium require teacher-controlled demonstrations and appropriate risk assessment.

3. A more reactive metal can take another metal’s place

Put zinc in copper sulfate solution and copper is deposited. Zinc is more reactive than copper, so it displaces copper from its compound: Zn + CuSO₄ → ZnSO₄ + Cu.

Copper cannot displace zinc from zinc sulfate because copper is less reactive. A displacement result therefore gives evidence about the order of two metals.

Compare clean surfaces, equal solution volumes and concentrations, similar metal amounts, temperature and observation time. Different surface areas could change the speed and make an observation difficult to interpret.

4. Choose an extraction method

Very unreactive metals, such as gold, may occur naturally as the element. Most metals occur in compounds in rocks. To extract a metal, separate it from the other elements in its compound.

Many metals below carbon in the reactivity series can be extracted by reducing their oxides with carbon. Carbon takes oxygen from the oxide, leaving the metal. Metals too reactive for this method, such as aluminium, are extracted using electrolysis of molten compounds.

In an unfamiliar extraction question, use the information given rather than memorising every industrial process. Consider whether the reducing agent can remove oxygen, the energy required, cost, waste and environmental effects.

You do not need the detailed stages of a blast furnace in this Trilogy topic. You do need to identify which substance gains or loses oxygen in a supplied equation.

Electrolysis can also be used when the metal would react with carbon during extraction, contaminating the product.

5. Higher only: follow the electrons

Higher only. A second definition of oxidation and reduction applies even when oxygen is not present: oxidation is loss of electrons; reduction is gain of electrons. Remember OIL RIG.

For zinc displacing copper:
Zn → Zn²⁺ + 2e⁻ (oxidation)
Cu²⁺ + 2e⁻ → Cu (reduction).

The overall ionic equation is Zn + Cu²⁺ → Zn²⁺ + Cu. Sulfate ions stay unchanged and are left out; they are spectator ions.

Check both atoms and charge. Electrons must cancel when half equations are added. Never describe a copper ion gaining electrons as “oxidation” simply because it changes into a metal.

6. Acid + suitable metal → salt + hydrogen

Magnesium, zinc and iron react with dilute hydrochloric or sulfuric acid. Bubbles show hydrogen being made. Copper does not react with these dilute acids under normal classroom conditions.

Metal + acid → salt + hydrogen.
Mg + 2HCl → MgCl₂ + H₂
Zn + H₂SO₄ → ZnSO₄ + H₂

With dilute hydrochloric acid, iron forms iron(II) chloride: Fe + 2HCl → FeCl₂ + H₂. With dilute sulfuric acid it forms iron(II) sulfate.

Hydrogen gives a squeaky pop with a lighted splint at the mouth of a small collected gas sample. Follow teacher instructions; keep flames away from larger hydrogen mixtures and the reaction vessel.

Higher only: these are redox reactions. For magnesium: Mg → Mg²⁺ + 2e⁻, so magnesium is oxidised. Hydrogen ions gain electrons: 2H⁺ + 2e⁻ → H₂, so H⁺ is reduced.

7. Name the salt and build its formula

A salt contains a positive ion, usually from a metal, and a negative ion from the acid. The acid decides the ending of the salt’s name.

A formula must have zero overall charge. For Mg²⁺ and Cl⁻, two chloride ions balance one magnesium ion, so the formula is MgCl₂. For Ca²⁺ and NO₃⁻, the formula is Ca(NO₃)₂; brackets show that the whole nitrate ion is repeated.

8. Bases, alkalis and carbonates

A base neutralises an acid. Metal oxides and metal hydroxides are bases. An alkali is a base that dissolves in water, such as sodium hydroxide. Copper oxide is a base but is insoluble, so it is not an alkali.

Acid + base → salt + water.
CuO + H₂SO₄ → CuSO₄ + H₂O
HCl + NaOH → NaCl + H₂O

Acid + carbonate → salt + water + carbon dioxide.
CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂

Carbon dioxide turns limewater milky/cloudy. Bubbles alone do not identify a gas: compare the reactants and carry out the appropriate test.

9. Required practical 8: make copper sulfate crystals

Aim: prepare a pure, dry sample of a soluble salt from an insoluble oxide or carbonate. Here, black copper oxide reacts with dilute sulfuric acid to make blue copper sulfate solution. Work under teacher supervision.

1. Measure dilute sulfuric acid into a beaker. Warm it gently using a Bunsen burner, tripod and gauze; do not boil. Turn off the flame before adding powder.
2. Add copper oxide in small portions, stirring. Continue until some black solid remains and no more dissolves. The excess solid shows the acid has been used up.
3. Filter the mixture. Excess copper oxide stays as the residue; copper sulfate solution passes through as the filtrate.
4. Put the filtrate in an evaporating basin. Heat gently using a water bath or electric heater to remove some water and concentrate the solution. Do not heat to dryness.
5. Leave the concentrated solution to cool so crystals form. Collect them by filtration, rinse with a small amount of cold distilled water and dry between filter papers or in a suitable warm place.

Why it works: adding excess insoluble oxide removes leftover acid, and filtration removes the excess oxide. Cooling allows salt to crystallise without boiling all the water away. Too much rinsing water dissolves the product again.

Safety: wear eye protection; avoid inhaling powder; handle hot glass carefully; use small quantities and appropriate acid concentration. Collect copper-containing waste as instructed rather than pouring it away.

The same logic works for other suitable soluble salts when the acid and insoluble oxide/carbonate are specified. It is not suitable for making a salt by simply filtering a mixture of two soluble reactants.

10. pH and neutralisation at particle level

In water, acids produce hydrogen ions, H⁺. Alkalis contain hydroxide ions, OH⁻. The pH scale describes how acidic or alkaline a solution is: below 7 acidic, 7 neutral, above 7 alkaline.

Add a few drops of universal indicator to a small solution sample, or place a drop on indicator paper. Compare the colour with its chart. A pH probe should be calibrated and rinsed between samples to reduce contamination.

During acid–alkali neutralisation, H⁺ and OH⁻ react to form water: H⁺(aq) + OH⁻(aq) → H₂O(l). Adding acid to an alkali lowers the pH; adding enough extra acid eventually makes the mixture acidic.

11. Higher only: strong is not the same as concentrated

Higher only. A strong acid is completely ionised in aqueous solution. Hydrochloric, nitric and sulfuric acids are strong acids. A weak acid is only partly ionised; examples include ethanoic, citric and carbonic acids.

Concentrated/dilute describes how much acid is present per volume of solution. Strong/weak describes the proportion ionised. A strong acid can be dilute; a weak acid can be concentrated.

For acids that can each release one H⁺ ion per molecule, at the same concentration the stronger acid produces more H⁺ ions and has a lower pH. Do not assume strength alone determines pH if concentrations differ.

12. Use electricity to break down an ionic compound

Electrolysis uses an electric current to break down an ionic compound. The liquid or solution carrying the current is the electrolyte. It must contain ions that can move.

In a solid ionic compound, ions are held in fixed positions, so it does not conduct. When molten or dissolved in water, its ions can move and carry charge. Electrons carry charge through the external wires; ions carry charge through the electrolyte.

Positive ions move towards the negative cathode. Negative ions move towards the positive anode. At electrodes, ions are discharged and products form. An inert electrode, such as graphite in these classroom investigations, does not take part in the intended reaction.

13. Molten compounds: only the compound’s ions

A molten binary ionic compound contains ions from just two elements. The metal forms at the cathode and the non-metal at the anode.

These are prediction examples, not instructions for unsupervised practical work. Molten compounds require high temperatures, and some products are toxic or very reactive.

Higher only: for molten PbBr₂, Pb²⁺ + 2e⁻ → Pb at the cathode and 2Br⁻ → Br₂ + 2e⁻ at the anode. Each half equation balances atoms and electrical charge.

14. Extracting aluminium

Aluminium is too reactive to be extracted from aluminium oxide by reduction with carbon. Electrolysis is used instead.

Aluminium oxide is dissolved in molten cryolite. The mixture melts at a lower temperature than pure aluminium oxide, reducing the energy needed to keep the electrolyte molten. Electrical energy is still needed to drive electrolysis, making extraction expensive.

Aluminium forms at the negative electrode. Oxygen forms at the positive carbon electrodes. The oxygen reacts with carbon to make carbon dioxide, so the carbon anodes wear away and must be replaced.

Higher only: Al³⁺ + 3e⁻ → Al is reduction. 2O²⁻ → O₂ + 4e⁻ is oxidation. Oxygen then reacts with the carbon anode; the electrode reaction and the reaction consuming carbon are linked but distinct steps.

15. Aqueous solutions: water adds another choice

Aqueous means dissolved in water. Water provides H⁺ and OH⁻ ions as well as the ions from the dissolved compound. This changes which products can form.

1. At the cathode: if the metal is more reactive than hydrogen, hydrogen is produced. If the metal is less reactive than hydrogen, the metal is deposited.
2. At the anode: a halogen is produced if halide ions (chloride, bromide or iodide) are present in the solution used. Otherwise oxygen is produced.

These are the product predictions required for this course. Real product proportions can depend on conditions; use the information given in the question. In particular, do not confuse aqueous sodium chloride, which forms hydrogen at the cathode, with molten sodium chloride, which forms sodium.

16. Required practical 9: investigate aqueous electrolysis

Aim: investigate products formed when aqueous solutions are electrolysed using inert electrodes. Start with a testable hypothesis, for example: “With inert electrodes, copper sulfate solution will deposit copper at the cathode, but sodium sulfate solution will produce hydrogen there.” This prediction uses the metals’ positions relative to hydrogen.

1. Wear eye protection. Add a measured volume of the first teacher-approved solution to a beaker or electrolysis cell.
2. Place clean graphite electrodes into the liquid, without allowing them to touch. Connect them to the positive and negative terminals of a low-voltage d.c. supply.
3. Switch on and observe each electrode separately for a fixed time. Record gas bubbles, colour changes and any metal deposit. Do not assume every bubble is the same gas.
4. Where appropriate, collect small gas samples using suitable teacher-provided collection apparatus. Test hydrogen with a lighted splint, oxygen with a glowing splint and chlorine with damp litmus paper. Never smell chlorine; it is harmful, so use the teacher’s specified ventilation and small-scale procedure.
5. Switch off before changing the apparatus. Rinse/clean electrodes and use fresh solution for each repeat and each different electrolyte.
6. Compare observations with the hypothesis. Repeat to check consistency. Explain products using ion movement, water-derived ions and reactivity.

Variables: change the identity of the electrolyte. Observe the products at each electrode. Keep the voltage, testing time, solution volume and concentration, electrode material, immersed area and spacing consistent where possible. Product identity is different from measuring the rate of gas production.

Gas-test results: hydrogen gives a squeaky pop; oxygen relights a glowing splint; chlorine bleaches damp litmus paper white (blue paper may first turn red). Switch off the current and test only a small collected sample using the teacher’s procedure.

Possible limitations include contaminated electrodes, products dissolving in water, leaks during gas collection and tests carried out on mixed gases. A missing visible deposit or bubble is not enough on its own to show that no reaction occurred.

17. Higher only: electrode half equations

Higher only. At the cathode, positive ions gain electrons: reduction. At the anode, negative ions lose electrons: oxidation.

1. Balance the atoms, including the two atoms in H₂, O₂ and Cl₂.
2. Add electrons to balance charge.
3. Check that reduction has electrons on the left and oxidation has electrons on the right.

18. Make an observation, then explain it

A question asking what you would observe needs something visible: fizzing, a coloured deposit or a change in solution colour. “Zinc is more reactive” is an explanation, not an observation.

For a practical plan, write the order, quantities or measurements, controls, safety, repeats and how evidence will test the prediction. For a calculation, show the equation or subtraction and give a unit.

Six-mark explanations need linked ideas. For example: “Excess copper oxide remains, showing the acid has reacted. Filtering removes this insoluble solid, leaving copper sulfate solution.” Higher-only labels identify specification content; optional challenges can also assess Foundation ideas.

Explain the two required practicals

Task 1: Explain the purpose of excess solid, filtration, gentle concentration and cooling in salt preparation.

Task 2: Predict what forms at A and B for aqueous copper sulfate. Explain how you would test the gas, keep the comparison fair and work safely. Use these two practicals to build a clear written plan; your answer saves with this topic.

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One final thought

Use a scientific idea and a specific example in your reflection.