OxfordAQA International GCSE Chemistry: Metals (9202)

This study guide teaches section 3.3.1 Metals of the OxfordAQA International GCSE Chemistry (9202) specification, Version 6.2, for exams May/June 2018 onwards. It covers the opening statements 3.3.1 a to c, the reactivity series (3.3.1.1 a to i) and metal carbonates (3.3.1.2 a and b). The qualification is not tiered, so everything here applies to you whatever grade you are aiming for. The specification says content from any part of it may be assessed on either paper.

Section 3.3.1 has no required practical; the copper sulfate electrolysis practical belongs to section 3.3.2.

The course hub lists every topic and the printable checklist lets you tick off each statement. Not sure where your gaps are? Try a free diagnostic first.

What this unit covers

Specification	What you must be able to do
3.3.1 a	Explain why metals conduct and can be bent or hammered into shape
3.3.1 b	Define an alloy; evaluate uses of alloys from composition data; describe steels
3.3.1 c	Give the properties that make copper good for wiring and plumbing
3.3.1.1 a	Describe reactions of eight metals with water or dilute acids; order them
3.3.1.1 b	Use displacement to place metals; describe it as oxidation and reduction; write ionic equations
3.3.1.1 c	Explain why gold is found as the metal but most metals are found as compounds
3.3.1.1 d	Explain reduction of oxides with carbon, including iron
3.3.1.1 e	Explain why metals above carbon are extracted by electrolysis and why this is expensive
3.3.1.1 f	Describe phytomining and bioleaching for low-grade copper ores
3.3.1.1 g	Obtain copper from copper salt solutions by electrolysis; write both half equations
3.3.1.1 h	Obtain copper by displacement with scrap iron; write the ionic equation
3.3.1.1 i	Give the benefits of recycling metals
3.3.1.2 a	Describe thermal decomposition of the carbonates of Mg, Cu, Zn, Ca and Li
3.3.1.2 b	Describe reactions of metal carbonates with acids

Why metals are useful (3.3.1 a to c)

A metal is a giant structure of positive ions with delocalised outer electrons between them. Two properties follow from this.

• Conduction. The delocalised electrons can move through the whole structure, so they carry charge and transfer heat energy.
• Malleability. The ions are arranged in layers. The layers can slide over each other without the structure breaking apart, so a metal can be bent or hammered into shape.

Alloys

An alloy is a mixture of at least two elements, at least one of which is a metal. Alloys often have different properties from the metals in them, which can make them more useful than the pure metals. Atoms of a different size make it harder for the layers to slide, so many alloys are harder than the pure metal.

Steels are mixtures of iron with carbon and sometimes other metals. Changing the amount of carbon, or adding metals such as chromium or nickel, changes the hardness, strength and resistance to corrosion.

You may be given the composition of an unfamiliar alloy and asked to judge a use for it. Link each property in the data to the job.

Worked example 1: alloy composition

A decorative spoon is made from an alloy that is 92.5% silver and 7.5% copper by mass. The spoon has a mass of 24.0 g. Calculate the mass of copper in it, and suggest why the alloy is used instead of pure silver.

1. Mass of copper = 7.5/100 × 24.0 = 1.80 g.
2. Pure silver bends and scratches easily. The added copper makes the alloy harder, so the spoon keeps its shape in use.

Copper for wiring and plumbing

Copper suits household wiring, pipes and tanks because:

• it is a good conductor of heat and electricity;
• it can be bent, but it is hard enough to make pipes or tanks;
• it does not react with water.

The reactivity series (3.3.1.1 a)

You arrange metals in order of reactivity by watching how vigorously they react with water and with dilute acids.

• Metal + water → metal hydroxide + hydrogen
• Metal + dilute hydrochloric acid → metal chloride + hydrogen
• Metal + dilute sulfuric acid → metal sulfate + hydrogen

Metal	With cold water	With dilute acid
Potassium	Floats, melts into a ball, whizzes about, hydrogen ignites with a lilac flame	Explosively violent: never done in school
Sodium	Floats, melts into a ball, fizzes and moves quickly	Violent: never done in school
Lithium	Floats and fizzes steadily; does not melt	Very vigorous
Calcium	Sinks and fizzes steadily; solution turns cloudy	Vigorous fizzing
Magnesium	A very slow trickle of bubbles over a long time	Rapid fizzing; tube gets warm
Zinc	No visible reaction	Steady bubbles
Iron	No visible reaction	Slow bubbles
Copper	No reaction	No reaction

From the table, the order (most reactive first) is: potassium, sodium, lithium, calcium, magnesium, zinc, iron, copper. Carbon is placed between magnesium and zinc when you decide how to extract a metal (see below).

Sample equations: 2Li + 2H₂O → 2LiOH + H₂ and Fe + 2HCl → FeCl₂ + H₂.

Displacement reactions (3.3.1.1 b)

Put a metal into a solution containing ions of a less reactive metal and it takes that metal’s place in the compound. If nothing happens, the metal added is less reactive than the metal in the solution.

In displacement, the atoms of the more reactive metal lose electrons, so they are oxidised. The ions of the less reactive metal gain electrons, so they are reduced. Both happen at once, so this is a redox reaction.

Worked example 2: an ionic equation

A strip of magnesium ribbon is placed in zinc chloride solution. A grey coating forms on the ribbon. Write the ionic equation and say what is oxidised.

1. Symbol equation: Mg + ZnCl₂ → MgCl₂ + Zn.
2. The chloride ions are unchanged (spectator ions), so leave them out.
3. Ionic equation: Mg + Zn²⁺ → Mg²⁺ + Zn.
4. Magnesium atoms lose two electrons each: magnesium is oxidised. Zinc ions gain two electrons each: zinc ions are reduced.

Where metals come from (3.3.1.1 c to e)

Gold and other very unreactive metals occur in the ground uncombined, because they do not react with oxygen, water or other substances in rocks. Most metals are found as compounds, so a chemical reaction is needed to extract the metal.

The method depends on where the metal sits relative to carbon:

Position	Method	Example
Less reactive than carbon	Heat the oxide with carbon	Iron, zinc
More reactive than carbon	Electrolysis of a molten compound	Aluminium

Reduction with carbon. For this topic, reduction means the removal of oxygen. Carbon removes oxygen from the metal oxide, and the carbon itself gains oxygen.

Iron in the blast furnace

You do not need details of the furnace, but you must know the raw materials and the simple chemistry.

• Raw materials: iron ore (mainly iron(III) oxide, Fe₂O₃), coke (carbon), limestone and hot air.
• Coke burns in the hot air and releases heat: C + O₂ → CO₂
• Carbon dioxide reacts with more coke to make carbon monoxide: CO₂ + C → 2CO
• Carbon monoxide removes oxygen from the ore: Fe₂O₃ + 3CO → 2Fe + 3CO₂
• Limestone removes impurities such as sand as slag: CaCO₃ → CaO + CO₂, then CaO + SiO₂ → CaSiO₃

The overall reduction by carbon can be summed up as 2Fe₂O₃ + 3C → 4Fe + 3CO₂. The iron(III) oxide is reduced (it loses oxygen); the carbon is oxidised (it gains oxygen).

Worked example 3: carbon needed

Using 2Fe₂O₃ + 3C → 4Fe + 3CO₂, calculate the mass of carbon needed to make 1120 kg of iron. (Aᵣ: Fe = 56, C = 12)

1. Amount of iron = 1120 ÷ 56 = 20.0 kmol.
2. Ratio C : Fe = 3 : 4, so amount of carbon = 20.0 × 3/4 = 15.0 kmol.
3. Mass of carbon = 15.0 × 12 = 180 kg.

Electrolysis. Metals more reactive than carbon, such as aluminium, cannot have the oxygen removed by carbon. They are extracted by electrolysis of molten compounds. Melting the compound and passing the current take large amounts of energy, so these metals are expensive. The aluminium details are in section 3.3.2.

Questions may describe the extraction of another metal; read the information and apply the reactivity ideas.

Copper from low-grade ores (3.3.1.1 f to h)

Traditional mining of low-grade copper ores moves huge amounts of rock for little copper. New methods are being researched to limit the environmental impact.

• Phytomining: plants grow on land containing copper compounds and absorb them. The plants are harvested and burned, and the ash contains the copper compounds.
• Bioleaching: bacteria produce a leachate, a solution that contains copper compounds.

Copper metal is then obtained from the copper salt solution in one of two ways.

1. Electrolysis. Using inert carbon (graphite) electrodes, copper ions move to the negative electrode (cathode) and gain electrons:

• Cathode: Cu²⁺ + 2e⁻ → Cu (reduction; pink-brown copper coats the electrode)
• Anode, if the solution is copper sulfate: 4OH⁻ → O₂ + 2H₂O + 4e⁻ (oxidation; bubbles of oxygen)
• Anode, if the solution is copper chloride: 2Cl⁻ → Cl₂ + 2e⁻ (oxidation; chlorine)

If the anode is made of copper instead of carbon, the anode dissolves: Cu → Cu²⁺ + 2e⁻. Read the question carefully for the electrode material it gives.

2. Displacement with scrap iron. Iron is more reactive than copper, so cheap scrap iron displaces copper from the solution:

• Ionic equation: Fe + Cu²⁺ → Fe²⁺ + Cu
• Iron atoms lose electrons (oxidised); copper ions gain electrons (reduced).

Worked example 4: scrap iron

How much iron reacts to produce 3.20 g of copper? (Aᵣ: Fe = 56, Cu = 64)

1. Amount of copper = 3.20 ÷ 64 = 0.0500 mol.
2. Ratio Fe : Cu = 1 : 1, so 0.0500 mol of iron reacts.
3. Mass of iron = 0.0500 × 56 = 2.80 g.

Recycling metals (3.3.1.1 i)

We should recycle metals because extracting them:

• uses up limited resources (ores are finite);
• is expensive in energy (heating furnaces, electricity for electrolysis);
• harms the environment (mining scars land, transport and furnaces release carbon dioxide, waste rock must be dumped).

Metal carbonates (3.3.1.2)

Thermal decomposition

The carbonates of magnesium, copper, zinc, calcium and lithium all break down on heating in the same way:

metal carbonate → metal oxide + carbon dioxide

• MgCO₃ → MgO + CO₂
• CuCO₃ → CuO + CO₂ (green solid turns black)
• ZnCO₃ → ZnO + CO₂ (white solid is yellow while hot and white again when cold)
• CaCO₃ → CaO + CO₂
• Li₂CO₃ → Li₂O + CO₂

Not all Group 1 carbonates decompose at Bunsen burner temperatures. Sodium carbonate and potassium carbonate do not, even though lithium carbonate does.

Worked example 5: mass of residue

2.10 g of magnesium carbonate is heated until its mass stops changing. Calculate the mass of magnesium oxide left. (Aᵣ: Mg = 24, C = 12, O = 16)

1. Mᵣ of MgCO₃ = 24 + 12 + 3 × 16 = 84; Mᵣ of MgO = 24 + 16 = 40.
2. Ratio MgCO₃ : MgO = 1 : 1, so 84 g gives 40 g.
3. Mass of MgO = 2.10 × 40/84 = 1.00 g. The other 1.10 g leaves as carbon dioxide.

Carbonates with acids

Metal carbonates react with acids to give a salt, water and carbon dioxide:

• CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂
• ZnCO₃ + H₂SO₄ → ZnSO₄ + H₂O + CO₂

You see fizzing; the gas turns limewater cloudy white.

Common errors

• Calling an alloy a compound. It is a mixture.
• Leaving the spectator ions in an ionic equation, or forgetting charges on the ions.
• Saying the copper ions are oxidised in displacement. They gain electrons, so they are reduced.
• Claiming all carbonates decompose in a Bunsen flame. Sodium and potassium carbonates do not.

Where to go next

Condense this page with the Metals revision notes, then test yourself with the Metals practice questions. For the mole steps, see Quantitative chemistry; for exam technique, see exam preparation.

Official syllabus

OxfordAQA International GCSE Chemistry (9202) specification, Version 6.2, for exams May/June 2018 onwards, published by OxfordAQA. Section 3.3.1 Metals, including 3.3.1.1 The reactivity series and 3.3.1.2 Metal carbonates.