Chemistry (0620) | Uses of metals

Cambridge IGCSE Chemistry (0620) Study Notes: Uses of Metals

Hello Future Chemists!

Welcome to the fascinating world of metals! In earlier sections, we learned how to extract metals and what makes them chemically reactive. Now, we are going to look at the practical side: Why do we choose specific metals (and their mixtures, called alloys) for specific jobs? This chapter connects the fundamental properties you learned—like conductivity, density, and hardness—directly to real-world applications, which is essential for your exams!

1. Uses of Pure Metals (Linking Property to Application)

The core principle of using a metal is simple: we match the job requirement to the metal's best physical property. We will focus on two key examples: Aluminium and Copper.

1.1 Aluminium (Al)

Aluminium is known for being lightweight and highly resistant to rust, despite being quite reactive chemically.

Use 1: Manufacture of Aircraft

• Key Property: Low density.
• Explanation: Aircraft need to be as light as possible to save fuel and fly efficiently. Aluminium's low density means it has a high strength-to-weight ratio.

Use 2: Overhead Electrical Cables

• Key Properties: Low density and good electrical conductivity.
• Explanation: While copper is a better conductor, aluminium is much lighter. Using light cables means the support pylons don't need to be as massive, saving construction costs. Aluminium conducts electricity well enough for large-scale power transmission.

Use 3: Food Containers (Foil/Cans)

• Key Property: Resistance to corrosion.
• Explanation: Aluminium reacts very quickly with oxygen in the air to form a thin, tough layer of aluminium oxide (\(\text{Al}_2\text{O}_3\)) on its surface. This oxide layer is impervious (cannot be penetrated) and prevents the metal underneath from reacting with food, air, or moisture. This is why aluminium seems unreactive, despite its high position in the reactivity series.
• (Self-Correction/Deep Dive: This explains the apparent unreactivity of aluminium mentioned later in the syllabus.)

1.2 Copper (Cu)

Copper is found lower in the reactivity series and is famous for its excellent conductivity.

• Use: Electrical wiring (inside houses and devices).
• Key Properties: Good electrical conductivity and ductility.
• Explanation: Copper is one of the best conductors after silver. Its high conductivity ensures minimal energy loss. Ductility means it can easily be drawn out (pulled) into very thin, flexible wires without breaking.

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2. Alloys and Their Enhanced Properties

2.1 What is an Alloy?

Pure metals are often too soft for many applications because their atoms are arranged in perfect, regular layers that can slide past each other easily.

An alloy is defined as a mixture of a metal with one or more other elements (usually other metals, or sometimes non-metals like carbon).

2.2 Why are Alloys Stronger? (The Structure Explanation)

Don't worry if this seems tricky at first—the explanation relies on understanding the metal structure you covered earlier (the giant metallic lattice).

1. In a pure metal, all the atoms are roughly the same size and are arranged neatly in layers.
2. When a force is applied (like hitting it with a hammer), these uniform layers can slide over each other easily. This makes the pure metal malleable (bendy) or ductile (can be drawn into wire), but often too soft.
3. In an alloy, the atoms of the added element are usually different sizes (either slightly larger or smaller) than the main metal atoms.
4. These different-sized atoms disrupt the neat, regular layering structure.
5. Because the layers are no longer perfectly regular, they cannot slide over each other easily.

Key Takeaway: This disruption makes the alloy harder and stronger than the pure metal. This is why alloys are generally more useful for structural purposes.

2.3 Examples of Common Alloys

Brass

• Composition: Mixture of Copper (Cu) and Zinc (Zn).
• Uses: Musical instruments, ornaments, and plumbing fittings (because it is corrosion resistant and easily shaped).

Stainless Steel

• Composition: Mixture of Iron (Fe) with elements like Chromium (Cr), Nickel (Ni), and Carbon (C).
• Uses: Cutlery, kitchen sinks, and surgical instruments.
• Key Properties (for use): Extreme hardness (due to carbon) and high resistance to rusting/corrosion (due to chromium, which forms a protective oxide layer).

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3. Corrosion of Metals (Rusting and Protection)

Corrosion is the destructive chemical reaction of a metal with substances in its environment, often leading to the formation of undesirable compounds (like oxides). The most common example is the rusting of iron.

3.1 Rusting: The Process

Rusting is the corrosion of iron (or steel, which is an iron alloy) to form hydrated iron(III) oxide.

For rusting to occur, two specific conditions are required simultaneously:

1. Oxygen (usually from the air)
2. Water

If either oxygen or water is completely excluded from the iron surface, rusting cannot take place.

The chemical formula for rust is typically written as \(\text{Fe}_2\text{O}_3\cdot x\text{H}_2\text{O}\), where \(x\) indicates a variable amount of water molecules (hence the term 'hydrated').

3.2 Prevention of Rusting: Barrier Methods

The simplest way to prevent rusting is to use a barrier method. These methods physically stop oxygen and water from reaching the iron surface.

• Painting: Used for large structures like car bodies or bridges. Effective as long as the paint layer remains unbroken.
• Greasing or Oiling: Used for moving parts in machinery or tools, as the grease not only protects the metal but also lubricates it.
• Coating with Plastic: Often used on wire fences or refrigerator shelves where a durable, long-lasting barrier is needed.

3.3 Prevention of Rusting: Sacrificial Protection (Extended/Deep Dive)

Barrier methods work well, but if the barrier is scratched, corrosion starts immediately. Sacrificial protection offers a 'self-healing' mechanism.

How Galvanising Works

Galvanising is a common example where iron or steel is coated with a layer of zinc. Zinc performs two roles:

1. It acts as a physical barrier.
2. If the zinc layer is scratched, it provides sacrificial protection.

The Chemistry of Sacrifice

Sacrificial protection relies entirely on the Reactivity Series.

• Zinc is more reactive than Iron (Zinc is higher in the reactivity series: ...Magnesium, Aluminium, Zinc, Iron, Hydrogen...).
• When the iron is exposed, both metals are in contact with water and oxygen.
• Because Zinc is more reactive, it has a stronger tendency to form positive ions (i.e., it loses electrons more easily) than iron.
• The Zinc metal (\(\text{Zn}\)) will react and corrode *instead* of the Iron (\(\text{Fe}\)). The zinc sacrifices itself to protect the iron.

Even when the coating is damaged, the zinc continues to react, and the iron remains protected until all the zinc layer is consumed.

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