Why metals matter in IGCSE Chemistry
Metals account for roughly three-quarters of all known elements, and the IGCSE Chemistry syllabus expects you to understand their properties, reactions, and extraction in considerable detail. This topic connects to electrochemistry, redox, and environmental chemistry, so a solid grasp here pays dividends across the entire course.
Let us work through the topic systematically, building from basic properties to the reactivity series and then to extraction methods. At each stage, we will pause to apply the ideas through worked problems - the kind of reasoning Cambridge examiners reward.
Physical properties: metals versus non-metals
The first task is to distinguish metals from non-metals by their physical characteristics. The table below summarises the key contrasts the IGCSE specification requires.
| Property | Metals | Non-metals |
|---|---|---|
| Thermal conductivity | Good conductors | Poor conductors (insulators) |
| Electrical conductivity | Good conductors | Poor conductors (except graphite) |
| Malleability and ductility | Malleable (can be hammered into shape) and ductile (can be drawn into wires) | Brittle when solid |
| Melting and boiling points | Generally high (exception: mercury is liquid at room temperature) | Generally low (exceptions: diamond, silicon) |
| Appearance | Shiny (lustrous) when freshly cut | Dull (except iodine crystals) |
| State at room temperature | Solid (except mercury) | Solid, liquid, or gas |
A common exam pitfall: students forget that graphite conducts electricity even though carbon is a non-metal. If a question asks for an exception, graphite is the classic answer.
Chemical properties of metals
Metals react in predictable patterns with dilute acids, water, and oxygen. The vigour of each reaction depends on where the metal sits in the reactivity series, which we will examine shortly.
Reactions with dilute acids
Metals above hydrogen in the reactivity series react with dilute hydrochloric acid or dilute sulfuric acid to produce a salt and hydrogen gas.
- Magnesium + hydrochloric acid: Mg + 2HCl → MgCl₂ + H₂ (vigorous fizzing, rapid dissolving)
- Zinc + sulfuric acid: Zn + H₂SO₄ → ZnSO₄ + H₂ (steady bubbling)
- Copper + hydrochloric acid: no reaction (copper is below hydrogen)
Reactions with water and steam
- Potassium and sodium react vigorously with cold water, producing a metal hydroxide and hydrogen gas.
- Calcium reacts steadily with cold water.
- Magnesium reacts very slowly with cold water but burns brightly in steam, producing magnesium oxide and hydrogen.
- Iron does not react with cold water but reacts reversibly with steam.
- Copper, silver, gold show no reaction with water or steam.
Reactions with oxygen
Most metals react with oxygen when heated. The more reactive the metal, the more vigorous the combustion. Magnesium burns with a brilliant white flame; iron glows and forms sparks; copper forms a black coating of copper(II) oxide; gold does not react at all.
The reactivity series
The reactivity series ranks metals (and carbon and hydrogen as reference points) from most reactive to least reactive. Memorising this order is non-negotiable for IGCSE Chemistry - it underpins displacement reactions, extraction methods, and electrochemistry.
| Position | Element | Reaction with water | Extraction method |
|---|---|---|---|
| Most reactive | Potassium (K) | Vigorous with cold water | Electrolysis |
| Sodium (Na) | Vigorous with cold water | Electrolysis | |
| Calcium (Ca) | Steady with cold water | Electrolysis | |
| Magnesium (Mg) | Very slow with cold water; vigorous with steam | Electrolysis | |
| Aluminium (Al) | Reacts with steam (protective oxide layer slows reaction) | Electrolysis | |
| Reference | Carbon (C) | - | - |
| Zinc (Zn) | Reacts with steam | Reduction with carbon | |
| Iron (Fe) | Reacts reversibly with steam | Reduction with carbon | |
| Reference | Hydrogen (H) | - | - |
| Copper (Cu) | No reaction | Reduction with carbon | |
| Silver (Ag) | No reaction | Found native or chemical reduction | |
| Least reactive | Gold (Au) | No reaction | Found native |
Displacement reactions: a worked approach
A more reactive metal will displace a less reactive metal from a solution of its salt. This is one of the most tested applications of the reactivity series. Let us work through the logic step by step.
Worked Example 1
Question: Predict what happens when a piece of zinc is placed into copper(II) sulfate solution.
Step 1 - Locate both metals in the reactivity series. Zinc sits above copper.
Step 2 - Apply the rule. The more reactive metal (zinc) displaces the less reactive metal (copper) from its salt solution.
Step 3 - Write the equation.
Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)
Step 4 - Describe the observations. The blue copper(II) sulfate solution gradually becomes colourless as zinc sulfate forms. A reddish-brown deposit of copper metal appears on the zinc strip. The zinc strip decreases in size.
Worked Example 2
Question: Will iron displace aluminium from aluminium chloride solution?
Step 1 - Locate both metals. Aluminium sits above iron in the reactivity series.
Step 2 - Apply the rule. Iron is less reactive than aluminium, so iron cannot displace aluminium.
Answer: No reaction occurs.
The pattern is straightforward: check the relative positions, and the metal higher up always wins. Exam questions often present an unfamiliar metal and give you experimental data to place it in the series - the reasoning is identical.
Uses of metals linked to properties
Cambridge expects you to connect specific uses to the physical properties that make them possible. This is a frequent source of 2-mark questions.
- Aluminium in aircraft manufacture - low density (the aircraft needs to be light enough to fly)
- Aluminium in overhead electrical cables - low density (lighter cables need fewer support pylons) and good electrical conductivity
- Aluminium in food containers - resistance to corrosion (a thin protective oxide layer forms naturally on the surface)
- Copper in electrical wiring - excellent electrical conductivity and ductility (can be drawn into thin, flexible wires)
When answering, always name the property first, then explain why that property suits the application. Vague answers such as "aluminium is useful" earn zero marks.
Alloys
An alloy is a mixture of a metal with one or more other elements (usually other metals, though carbon is included in steel). Alloys are typically harder and stronger than the pure metals because atoms of different sizes disrupt the regular layered arrangement, making it harder for layers to slide over each other.
Two alloys you must know:
- Brass - a mixture of copper and zinc. Harder than pure copper, resistant to corrosion, used in musical instruments and fittings.
- Stainless steel - a mixture of iron with chromium, nickel, and carbon. Resists rusting, used in cutlery, surgical instruments, and kitchen sinks.
Corrosion of metals
Rusting is the corrosion of iron. It requires both oxygen and water to be present. Remove either one, and rusting stops. This is testable with a simple controlled experiment using three test tubes:
- Iron nail in ordinary water exposed to air - rusts (both water and oxygen present)
- Iron nail in boiled water sealed with oil - does not rust (water present but no dissolved oxygen; oil prevents air re-entering)
- Iron nail in a dry tube with calcium chloride desiccant - does not rust (oxygen present but no water)
Prevention methods fall into two categories:
- Barrier methods - painting, oiling, greasing, plastic coating, electroplating (all prevent water and oxygen from reaching the iron surface)
- Sacrificial protection (galvanising) - coating iron with zinc. Zinc is more reactive than iron, so it corrodes preferentially, protecting the iron underneath even if the coating is scratched.
Extraction of metals
The extraction method depends on the metal's position in the reactivity series. This is a direct, logical connection that examiners test regularly.
- Metals above carbon (K, Na, Ca, Mg, Al) are too reactive to be reduced by carbon. They must be extracted by electrolysis of their molten compounds.
- Metals below carbon but above hydrogen (Zn, Fe, Cu) can be extracted by heating their oxide with carbon (reduction). Carbon removes the oxygen, forming carbon dioxide.
- Very unreactive metals (Ag, Au) are found uncombined in the Earth's crust and need only physical separation.
Worked Example 3
Question: Iron is extracted from iron(III) oxide using carbon in a blast furnace. Write the equation and explain why this method works.
Equation: 2Fe₂O₃ + 3C → 4Fe + 3CO₂
Explanation: Carbon is more reactive than iron (it sits above iron in the reactivity series). Carbon therefore reduces the iron(III) oxide by removing the oxygen, forming carbon dioxide. The iron is left behind as the molten metal.
Aluminium extraction is a special case worth noting. Although aluminium oxide has a very high melting point (over 2000 degrees Celsius), adding cryolite lowers the melting point considerably, reducing energy costs. The molten mixture is then electrolysed, with aluminium deposited at the cathode and oxygen produced at the anode.
Common exam mistakes to avoid
- Confusing rusting conditions. Both water and oxygen are needed. Students who write "water or oxygen" lose the mark.
- Forgetting carbon's position. Carbon is not a metal, but its position in the reactivity series determines which metals can be extracted by carbon reduction.
- Writing unbalanced equations. Always count atoms on both sides. Displacement and extraction equations are favourites for balancing questions.
- Mixing up sacrificial protection and barrier methods. Galvanising works by sacrificial protection (zinc reacts instead of iron), not simply as a physical barrier.
- Vague property-use links. Saying copper is "good for wires" without naming conductivity and ductility earns no marks.
Self-check questions
- Place these metals in order of reactivity from most to least reactive: iron, potassium, copper, magnesium.
- Predict and explain what happens when magnesium ribbon is placed into iron(II) sulfate solution.
- Explain why aluminium is extracted by electrolysis rather than by reduction with carbon.
- An iron gate is galvanised with zinc. Explain how this protects the iron from rusting, even if the zinc coating is scratched.
- State two conditions necessary for iron to rust and describe an experiment to prove one of them is essential.
A methodical guide to the Metals topic in IGCSE Chemistry (0620), covering physical and chemical properties, the reactivity series, alloys, corrosion, and extraction methods. Includes worked examples on displacement reactions and exam-focused comparison tables.
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