(I once held up a piece of pure aluminium and asked a class, "Do you know how much energy it took to make this?" They shrugged. Then I told them: extracting aluminium from its ore requires vast amounts of electricity — so much that aluminium smelters are often built next to hydroelectric dams. And yet, once you have the aluminium, it can be recycled using only 5% of the energy. That is the story of metallurgy — the chemistry of extracting metals from their ores, and the reason why recycling matters so much. In this chapter, we go deeper than ever before into the world of metals. We will learn about the reactivity series, how metals are extracted, how alloys are designed, and how corrosion can be prevented. This is the chemistry that built the modern world. Let us begin.) In your earlier studies, you learned about the physical properties of metals, the reactivity series, and the basic principles of extraction. That is a solid foundation — but at the advanced level, we go deeper. We ask: How exactly do we extract a metal from its ore? What are the thermodynamic principles behind the extraction? How do we choose the right method for each metal? How do we design alloys with specific properties? How does corrosion work at the atomic level, and how can we prevent it? These are the questions that this chapter answers. We will explore the thermodynamics of metal extraction, the chemistry of the blast furnace and electrolysis, the design of alloys, and the electrochemistry of corrosion. By the end, you will have a deep understanding of the metals that build our world. Let us begin.
Metals have a set of characteristic physical properties that distinguish them from non-metals. These properties arise from metallic bonding — a lattice of positive ions in a sea of delocalised electrons.
| Property | Description | Explanation |
|---|---|---|
| Lustrous (shiny) | Metals reflect light and appear shiny when polished. | Delocalised electrons reflect light. |
| Malleable | Can be hammered into thin sheets without breaking. | Layers of ions slide over each other without breaking the metallic bond. |
| Ductile | Can be drawn into wires. | Same reason as malleability. |
| Good conductors of heat | Metals conduct heat quickly. | Delocalised electrons transfer kinetic energy rapidly. |
| Good conductors of electricity | Metals conduct electricity. | Delocalised electrons are free to move and carry charge. |
| High melting and boiling points | Most metals have high melting points. | Strong metallic bonds require a lot of energy to break. |
| High density | Most metals are dense. | Atoms are closely packed in a regular lattice. |
| Sonorous | Metals make a ringing sound when struck. | Vibration of the lattice. |
Exceptions: Mercury is a liquid at room temperature. Sodium, potassium, and lithium are soft, have low melting points, and low densities.
Metals react by losing electrons to form positive ions. The ease with which they lose electrons determines their reactivity.
Most metals react with oxygen to form metal oxides. The reaction is called oxidation.
2Mg + O₂ → 2MgO
4Na + O₂ → 2Na₂O
2Cu + O₂ → 2CuO (slowly, on heating)
| Metal | Reaction with Cold Water | Reaction with Steam |
|---|---|---|
| Potassium, sodium, calcium | React vigorously, producing hydrogen and a metal hydroxide. | — |
| Magnesium | Very slow reaction with cold water. | Reacts vigorously with steam to form magnesium oxide and hydrogen. |
| Zinc, iron | No reaction with cold water. | React with steam to form the metal oxide and hydrogen. |
| Copper, silver, gold | No reaction. | No reaction. |
Metals above hydrogen in the reactivity series react with dilute acids to produce a salt and hydrogen gas.
Mg + 2HCl → MgCl₂ + H₂
Zn + H₂SO₄ → ZnSO₄ + H₂
Metals below hydrogen (copper, silver, gold) do not react with dilute acids.
A more reactive metal can displace a less reactive metal from its compound.
Zn + CuSO₄ → ZnSO₄ + Cu
The reactivity series arranges metals in order of their reactivity, from most reactive to least reactive.
| Metal | Reactivity | Extraction Method |
|---|---|---|
| Potassium (K) | Most reactive | Electrolysis of molten chloride |
| Sodium (Na) | Very reactive | Electrolysis of molten chloride |
| Calcium (Ca) | Very reactive | Electrolysis of molten chloride |
| Magnesium (Mg) | Reactive | Electrolysis of molten chloride |
| Aluminium (Al) | Reactive | Electrolysis of molten oxide |
| Zinc (Zn) | Moderately reactive | Reduction with carbon |
| Iron (Fe) | Moderately reactive | Reduction with carbon (blast furnace) |
| Lead (Pb) | Less reactive | Reduction with carbon |
| Copper (Cu) | Least reactive | Heating in air or displacement |
| Silver (Ag) | Very unreactive | Found native |
| Gold (Au) | Most unreactive | Found native |
(The reactivity series is one of the most useful tools in chemistry. Once you know where a metal sits in the series, you can predict how it will react, how it will be extracted, and how it will behave in the environment. It is a shortcut to understanding a huge amount of chemistry.)
The method used to extract a metal from its ore depends on its position in the reactivity series. The key thermodynamic principle is that a more reactive metal has a more negative standard electrode potential (or a more positive enthalpy of formation of its oxide). To extract a metal, we must reduce it — that is, we must supply electrons to its ions.
Methods of extraction:
Iron is extracted from its ore (haematite, Fe₂O₃) in a blast furnace.
Raw materials:
Key reactions:
Products:
Types of iron produced:
| Type | Carbon Content | Properties | Uses |
|---|---|---|---|
| Pig iron | 3–4% | Hard, brittle | Making steel |
| Cast iron | 2–4% | Hard, brittle, corrosion-resistant | Engine blocks, pipes, cookware |
| Wrought iron | Soft, malleable, ductile | Decorative ironwork, chains | |
| Steel | 0.1–1.5% | Hard, strong | Construction, tools, machinery |
Aluminium is extracted from bauxite (Al₂O₃) by electrolysis. The bauxite is dissolved in molten cryolite (Na₃AlF₆) to lower the melting point from about 2000°C to about 1000°C, saving energy.
Electrolysis of molten aluminium oxide:
Overall: 2Al₂O₃ → 4Al + 3O₂
Key points:
Copper is extracted from its ores (e.g., chalcopyrite, CuFeS₂) by a combination of processes:
Zinc is extracted from its ore (zinc blende, ZnS) by roasting and then reduction with carbon:
An alloy is a mixture of two or more metals, or a metal and a non-metal, designed to have improved properties. Alloys are usually harder and stronger than pure metals.
Why are alloys harder than pure metals? In a pure metal, the atoms are all the same size and can slide over each other easily. In an alloy, atoms of different sizes disrupt the regular lattice, preventing the layers from sliding. This makes the alloy harder and stronger.
| Alloy | Composition | Properties | Uses |
|---|---|---|---|
| Steel | Iron + carbon | Harder than pure iron | Construction, tools, machinery |
| Stainless steel | Iron + chromium + nickel | Resists rusting | Cutlery, surgical instruments, kitchen sinks |
| Brass | Copper + zinc | Hard, shiny, corrosion-resistant | Musical instruments, door handles, coins |
| Bronze | Copper + tin | Hard, corrosion-resistant | Statues, medals, bells |
| Solder | Lead + tin | Low melting point | Joining electrical wires |
| Duralumin | Aluminium + copper + magnesium | Light, strong | Aircraft, cars |
| Nichrome | Nickel + chromium + iron | High electrical resistance, heat-resistant | Heating elements in toasters, kettles |
Corrosion is the gradual destruction of a metal by chemical reactions with substances in its environment.
| Metal | Corrosion Product | Conditions Needed |
|---|---|---|
| Iron | Rust (hydrated iron(III) oxide) | Oxygen and water |
| Copper | Green patina (basic copper carbonate) | Oxygen, water, carbon dioxide |
| Aluminium | Aluminium oxide (protective layer) | Oxygen (forms protective layer that prevents further corrosion) |
| Silver | Tarnish (silver sulphide) | Hydrogen sulphide in air |
Rusting is an electrochemical process. It occurs through the following steps:
Factors that accelerate rusting: salt (acts as an electrolyte), acid rain (provides H⁺ ions), and contact with less reactive metals (creates a galvanic cell).
| Method | How It Works | Examples |
|---|---|---|
| Painting | Creates a barrier preventing oxygen and water from reaching the iron. | Gates, bridges, car bodies |
| Oiling / greasing | Same principle — a barrier. | Tools, bicycle chains, machine parts |
| Galvanising (coating with zinc) | Zinc is more reactive than iron, so it corrodes first, protecting the iron underneath (sacrificial protection). | Dustbins, roofing sheets, buckets |
| Electroplating (coating with tin, chromium, nickel) | Barrier coating of a less reactive metal. | Food cans (tin-plated), taps (chrome-plated) |
| Alloying (making stainless steel) | Mixing iron with chromium and nickel changes its structure and prevents rusting. | Cutlery, surgical instruments, kitchen sinks |
Recycling metals is important for both economic and environmental reasons:
(The next time you throw away a drinks can, remember: that can is made of aluminium, and recycling it uses only 5% of the energy it took to make it. That is a staggering saving. Recycling is not just a good habit — it is a chemical imperative.)
| Metal Ion | Test | Result |
|---|---|---|
| Cu²⁺ | Add sodium hydroxide solution | Blue precipitate |
| Fe²⁺ | Add sodium hydroxide solution | Green precipitate |
| Fe³⁺ | Add sodium hydroxide solution | Red-brown precipitate |
| Zn²⁺ | Add sodium hydroxide solution | White precipitate, soluble in excess |
| Al³⁺ | Add sodium hydroxide solution | White precipitate, soluble in excess |
| Ca²⁺ | Flame test | Brick-red flame |
| Na⁺ | Flame test | Golden-yellow flame |
| K⁺ | Flame test | Lilac flame |
Question: Which is more reactive — zinc or copper? Explain using the reactivity series.
Solution: Zinc is more reactive than copper. In the reactivity series, zinc is above copper. This means zinc can displace copper from its compounds. For example, if zinc is placed in copper(II) sulphate solution, it will displace the copper: Zn + CuSO₄ → ZnSO₄ + Cu. The blue colour of the solution fades as copper is deposited.
Question: Explain why aluminium is extracted by electrolysis rather than by reduction with carbon.
Solution: Aluminium is high in the reactivity series — it is more reactive than carbon. This means carbon cannot displace aluminium from its oxide. Only electrolysis, which uses electrical energy to break down the compound, is powerful enough to extract aluminium from its ore. Metals below carbon in the reactivity series (like iron and zinc) can be extracted by reduction with carbon.
Question: Explain why galvanising protects iron from rusting, even if the zinc coating is scratched.
Solution: Galvanising coats iron with zinc. Zinc is more reactive than iron, so it corrodes in preference to iron — this is called sacrificial protection. Even if the coating is scratched, the zinc continues to corrode first, protecting the iron underneath. This is why galvanised steel is so durable in outdoor conditions.
Question: Explain why steel is harder than pure iron.
Solution: Steel is an alloy of iron and carbon. In pure iron, the atoms are all the same size and can slide over each other easily, making it relatively soft. In steel, the carbon atoms are different in size and disrupt the regular lattice of iron atoms. This prevents the layers from sliding over each other, making the alloy harder and stronger than pure iron.
1. State three physical properties of metals.
Any three of: lustrous, malleable, ductile, good conductors of heat and electricity, high melting points, high density, sonorous.
2. What is the reactivity series?
The reactivity series is an arrangement of metals in order of their chemical reactivity, from most reactive (potassium) to least reactive (gold).
3. Explain why gold is found in nature as a pure element.
Gold is very unreactive — it is at the bottom of the reactivity series. It does not react with oxygen, water, or most other substances, so it is found in nature as the pure metal.
4. What conditions are needed for iron to rust?
Oxygen and water (moisture).
5. Describe two methods of preventing rusting.
Any two of: painting, oiling/greasing, galvanising (coating with zinc), electroplating, alloying (stainless steel).
6. What is an alloy? Give one example.
An alloy is a mixture of two or more metals, or a metal and a non-metal, designed to have improved properties. Example: steel (iron + carbon).
7. Why is aluminium extracted by electrolysis rather than by reduction with carbon?
Aluminium is more reactive than carbon, so carbon cannot displace it from its oxide. Only electrolysis is powerful enough to extract aluminium from its ore.
Question 1 [10 marks]
(a) Describe the physical properties of metals and explain them using metallic bonding. [4]
(b) Describe the reaction of sodium, magnesium, and copper with water. [3]
(c) Explain why metals conduct electricity. [3]
(a) Metals are lustrous, malleable, ductile, good conductors of heat and electricity, and have high melting points. These properties arise from metallic bonding: a lattice of positive ions in a sea of delocalised electrons. The delocalised electrons are free to move, which explains conductivity. The layers of ions can slide over each other, which explains malleability and ductility. (4 marks: 2 for properties, 2 for explanation.)
(b) Sodium reacts vigorously with cold water to form sodium hydroxide and hydrogen. Magnesium reacts very slowly with cold water but reacts with steam to form magnesium oxide and hydrogen. Copper does not react with water or steam. (3 marks: 1 for each metal.)
(c) Metals conduct electricity because they have delocalised electrons that are free to move throughout the lattice. When a voltage is applied, these electrons carry charge from one end to the other. (3 marks: 1 for delocalised electrons, 1 for free movement, 1 for carrying charge.)
Question 2 [10 marks]
(a) Describe the extraction of iron in a blast furnace, including the raw materials and key reactions. [6]
(b) Explain why aluminium is extracted by electrolysis. [2]
(c) State two uses of aluminium and explain why it is suitable for each. [2]
(a) Raw materials: haematite (Fe₂O₃), coke (C), limestone (CaCO₃), and air. Coke burns to produce CO₂: C + O₂ → CO₂. CO₂ reacts with coke to form CO: CO₂ + C → 2CO. CO reduces iron oxide: Fe₂O₃ + 3CO → 2Fe + 3CO₂. Limestone decomposes: CaCO₃ → CaO + CO₂. CaO removes silica as slag: CaO + SiO₂ → CaSiO₃. Molten iron sinks to the bottom and is tapped off. (6 marks: 1 for raw materials, 1 for coke reaction, 1 for CO formation, 1 for reduction, 1 for limestone, 1 for slag.)
(b) Aluminium is more reactive than carbon, so carbon cannot displace it from its oxide. Electrolysis is used because it can break down the aluminium oxide into aluminium and oxygen using electrical energy. (2 marks: 1 for "more reactive than carbon", 1 for "electrolysis".)
(c) Any two of: aircraft (light and strong), cans (light, corrosion-resistant, non-toxic), window frames (corrosion-resistant), foil (malleable, non-toxic), electrical cables (good conductor, light). (2 marks: 1 for each use + reason.)
Question 3 [10 marks]
(a) Describe the conditions needed for iron to rust and explain how rusting can be prevented. [6]
(b) Explain why galvanising protects iron even when the coating is scratched. [2]
(c) What is an alloy? Explain why alloys are usually harder than pure metals. [2]
(a) Iron rusts when exposed to both oxygen and water. Prevention methods include: painting (barrier), oiling/greasing (barrier), galvanising (coating with zinc), electroplating (coating with tin or chromium), and alloying (stainless steel). (6 marks: 1 for conditions, 1 for each of five prevention methods.)
(b) Galvanising protects iron by sacrificial protection. Zinc is more reactive than iron, so it corrodes in preference to the iron. Even if the coating is scratched, the zinc continues to corrode first, protecting the iron underneath. (2 marks: 1 for "more reactive than iron", 1 for "corrodes first".)
(c) An alloy is a mixture of two or more metals, or a metal and a non-metal. Alloys are harder than pure metals because the atoms of different sizes disrupt the regular lattice, preventing the layers from sliding over each other. (2 marks: 1 for definition, 1 for explanation.)
Challenge 1: Explain why the reactivity series is so important in the extraction of metals.
The reactivity series determines the method used to extract a metal from its ore. Metals above carbon in the series (K, Na, Ca, Mg, Al) are too reactive to be reduced by carbon, so they must be extracted by electrolysis, which uses electrical energy to break down their compounds. Metals below carbon (Zn, Fe, Pb) can be extracted by reduction with carbon, which is cheaper and simpler. Metals at the bottom of the series (Cu, Ag, Au) are so unreactive that they can be found native or obtained by simple heating. Without the reactivity series, we would not know which method to use for each metal. It is the key that unlocks the entire field of metallurgy.
Challenge 2: A student says, "Gold is the best metal because it is the most expensive." Is this statement accurate? Explain, using examples from the chapter.
The statement is not accurate. The "best" metal depends on what you need it for. Gold is expensive because it is rare and unreactive, which makes it useful for jewellery and electronics. But it is not the best metal for every purpose. Iron is much cheaper and stronger, so it is used for construction and machinery. Copper is a better conductor of electricity than gold, so it is used for wiring. Aluminium is lighter than gold, so it is used in aircraft. Each metal has properties that make it suitable for specific uses. The value of a metal is not just in its price — it is in its properties and how well they match the job. This is why we use a wide range of metals, not just gold.
Challenge 3: Explain why the corrosion of iron is a serious economic problem, and describe three methods of preventing it.
The corrosion of iron (rusting) is a serious economic problem because iron and steel are used extensively in construction, transport, and industry. Rust weakens structures, causing bridges to fail, cars to deteriorate, and machinery to break down. Billions are spent every year repairing or replacing rusted iron and steel. Three methods of prevention: (1) Painting — creates a barrier that prevents oxygen and water from reaching the iron. (2) Galvanising — coating with zinc, which corrodes in preference to iron (sacrificial protection). (3) Alloying — making stainless steel by adding chromium and nickel, which changes the structure of the iron and prevents rusting. Each method works by either blocking the reactants or by changing the metal's structure so it cannot rust.
"Metals — the backbone of civilisation. From the ore in the ground to the metal in your hand, you now understand the chemistry of extraction. Go build something."
