# Metallic Bonding and Giant Structures
Metals are everywhere — from the steel in buildings to the copper in electrical wires. The properties that make metals so useful (strength, conductivity, malleability) are all explained by metallic bonding. In this guide, we'll explore how metallic bonds work, why metals have their characteristic properties, and how alloys improve on pure metals.
1. What Is Metallic Bonding?
In a metal, the atoms are packed closely together in a regular arrangement. The outer-shell electrons become delocalised — they detach from their atoms and are free to move throughout the entire structure.
This creates:
- A lattice of positive metal ions (the atoms that have lost their outer electrons)
- A "sea of delocalised electrons" that moves freely between the ions
The metallic bond is the strong electrostatic attraction between the positive metal ions and the sea of negatively charged delocalised electrons.
Definition: Metallic bonding is the strong electrostatic force of attraction between positively charged metal ions and a sea of delocalised electrons.
2. Structure of Metals
Metals have a giant metallic structure (also called a giant metallic lattice):
- Regular arrangement of positive ions in rows and layers
- Delocalised electrons fill the spaces between ions
- The structure extends in all directions
- The bonding is strong throughout the entire structure
3. Properties of Metals
3.1 High Melting and Boiling Points
Why: The electrostatic attraction between the positive ions and the sea of delocalised electrons is strong and acts throughout the giant lattice. A large amount of energy is needed to overcome these forces.
The stronger the metallic bond (more delocalised electrons, smaller ions), the higher the melting point.
3.2 Good Electrical Conductivity
Why: The delocalised electrons are free to move through the metal lattice. When a potential difference (voltage) is applied, electrons flow through the metal, carrying charge. This is how electric current works in metals.
3.3 Good Thermal Conductivity
Why: When one part of a metal is heated, the delocalised electrons gain kinetic energy and move faster. They transfer this energy quickly to other parts of the metal by colliding with other electrons and ions.
3.4 Malleability and Ductility
Malleable = can be hammered into shape Ductile = can be drawn into wires
Why: When a force is applied, the layers of metal ions can slide over each other without breaking the metallic bond. The sea of delocalised electrons shifts with the ions, maintaining the attraction. The bond doesn't break — it just reforms in the new position.
This is very different from ionic compounds, where displacing layers causes like charges to be adjacent, leading to cracking.
Summary Table
| Property | Explanation |
|---|---|
| High melting point | Strong attraction between ions and delocalised electrons throughout the giant lattice |
| Good electrical conductor | Delocalised electrons free to move and carry charge |
| Good thermal conductor | Delocalised electrons transfer kinetic energy quickly |
| Malleable | Layers slide without bonds breaking; electrons maintain attraction |
| Ductile | Same as malleable — layers slide |
| Shiny | Delocalised electrons reflect light |
4. Factors Affecting Metallic Bond Strength
Not all metals are the same. The strength of metallic bonding depends on:
- Number of delocalised electrons: More delocalised electrons → stronger metallic bonding
- Size of the metal ion: Smaller ions → electrons closer to nucleus → stronger attraction
- Charge on the ion: Higher charge → stronger attraction
| Metal | Group | Ion Charge | Delocalised Electrons | Relative MP |
|---|---|---|---|---|
| Sodium | 1 | +1 | 1 per atom | Low (98°C) |
| Magnesium | 2 | +2 | 2 per atom | Higher (650°C) |
| Aluminium | 3 | +3 | 3 per atom | Higher (660°C) |
Magnesium has a higher melting point than sodium because:
- Mg²⁺ has a higher charge than Na⁺
- Mg has more delocalised electrons per atom (2 vs 1)
- Mg²⁺ ions are smaller than Na⁺ ions
- All of these mean stronger electrostatic attraction
5. Alloys
An alloy is a mixture of two or more elements, at least one of which is a metal. Most alloys are mixtures of metals.
Why Are Alloys Harder Than Pure Metals?
In a pure metal, all atoms are the same size, arranged in regular layers. These layers can slide over each other relatively easily.
In an alloy, atoms of different sizes are mixed in. The different-sized atoms disrupt the regular arrangement of layers. This makes it harder for layers to slide, so the alloy is harder and stronger than the pure metal.
Common Alloys
| Alloy | Composition | Properties | Uses |
|---|---|---|---|
| Steel | Iron + carbon (+ other metals) | Harder than pure iron; various types | Construction, vehicles, tools |
| Stainless steel | Iron + chromium + nickel | Resistant to corrosion | Cutlery, medical instruments |
| Bronze | Copper + tin | Hard, resistant to corrosion | Statues, coins, bells |
| Brass | Copper + zinc | Harder than copper, golden colour | Musical instruments, fittings |
| Gold alloys (e.g. 18 carat) | Gold + copper/silver | Harder than pure gold | Jewellery |
Why Not Use Pure Metals?
Pure metals are often:
- Too soft (gold, copper, iron, aluminium)
- Too reactive (iron rusts easily)
- Too expensive (gold)
Alloying solves these problems by making metals harder, stronger, and more resistant to corrosion.
6. Special Metals
Copper
- Excellent electrical conductor
- Used in electrical wiring
- Malleable and ductile
- Does not react with water
Aluminium
- Low density (lightweight)
- Resistant to corrosion (oxide layer protects it)
- Used in aircraft, drink cans, overhead power lines
- Good conductor of heat and electricity
Iron/Steel
- Strong (especially as steel)
- Relatively cheap and abundant
- Rusts unless protected (paint, galvanising, stainless steel)
- Used in construction, vehicles, bridges
Worked Example: Explaining Conductivity
Question: Explain why copper is a good electrical conductor.
Copper has a giant metallic structure. The outer electrons of copper atoms become delocalised — they are free to move throughout the structure. When a voltage is applied, these delocalised electrons flow through the metal, carrying electrical charge. This makes copper an excellent conductor of electricity.
Worked Example: Comparing Metals
Question: Explain why magnesium has a higher melting point than sodium.
Magnesium atoms have 2 outer electrons that become delocalised (vs 1 for sodium), and Mg²⁺ ions have a higher charge than Na⁺ and are smaller. This means the electrostatic attraction between the Mg²⁺ ions and the sea of delocalised electrons is stronger, requiring more energy to overcome. Therefore, magnesium has a higher melting point.
Worked Example: Alloys
Question: Explain why steel is harder than pure iron.
In pure iron, the atoms are all the same size and arranged in regular layers that can slide over each other easily. In steel, carbon atoms (which are smaller than iron atoms) are mixed into the structure. These different-sized atoms disrupt the regular layers, making it more difficult for layers to slide. This makes steel harder and stronger than pure iron.
8. Practice Questions
- Define metallic bonding.
- Explain why metals are good conductors of electricity, referring to their structure.
- Why are metals malleable but ionic compounds are brittle?
- Explain why an alloy is generally harder than the pure metal it is made from.
- Suggest why aluminium is used for aircraft bodies but steel is used for car bodies.
Want to check your answers and get step-by-step solutions?
9. Common Misconceptions
| Misconception | Reality |
|---|---|
| Metallic bonds are weak | Metallic bonds are strong — most metals have high melting points |
| Metals conduct because of ions moving | Electrons (not ions) carry charge in solid metals |
| Alloys are compounds | Alloys are mixtures — the metals are not chemically bonded to each other |
| All metals are hard | Some metals (like sodium and gold) are quite soft |
| Metals are malleable because bonds break | Layers slide while the metallic bond reforms — bonds don't break |
10. Exam Tips
- Always mention delocalised electrons and positive metal ions when describing metallic bonding
- Use the phrase "strong electrostatic force of attraction" in your answer
- For alloy questions, explain that different-sized atoms prevent layers from sliding
- When comparing metals, use the number of delocalised electrons and ion size/charge
- Don't confuse metallic bonding with ionic bonding — both involve electrostatic attraction but the mobile particles are different (electrons vs ions)
Frequently Asked Questions
Why can't metallic bonding explain the hardness of all metals?
While metallic bonding explains general properties, the exact hardness also depends on crystal structure and how the layers are arranged. This is beyond GCSE.
Are delocalised electrons shared between specific atoms?
No — they belong to the entire structure, not to any individual atom. Think of them as a "sea" or "cloud" of electrons.
Why is mercury a liquid at room temperature if metallic bonds are strong?
Mercury has an unusual electron configuration that results in weaker metallic bonding. This is studied at university level.
Summary
- Metallic bonding = strong attraction between positive metal ions and a sea of delocalised electrons
- Properties: high melting point, good conductors (heat and electricity), malleable, ductile
- Stronger bonds with: more delocalised electrons, smaller ions, higher ion charge
- Alloys are harder because different-sized atoms prevent layers from sliding
- Metals conduct electricity via delocalised electrons; ionic compounds conduct via free ions (when molten/dissolved)