Metallic Bonding
This lesson explains metallic bonding and how the structure of metals accounts for their characteristic properties. You also need to understand why alloys are harder than pure metals. This topic is covered in the Edexcel GCSE Chemistry (1CH0) specification under Structure, Bonding and Properties.
What Is Metallic Bonding?
In a metal:
- The atoms are arranged in a regular pattern (a lattice).
- The outer-shell electrons of each metal atom are delocalised — they are no longer associated with a single atom.
- This creates a structure of positive metal ions (the atoms that have lost their outer electrons) surrounded by a "sea" of delocalised electrons.
- The metallic bond is the strong electrostatic attraction between the positive metal ions and the delocalised electrons.
Exam Tip: The definition of metallic bonding is: "the strong electrostatic attraction between positive metal ions and a sea of delocalised electrons." Learn this word for word.
The Structure of Metals
The structure can be described as:
- Positive metal ions arranged in a regular lattice (rows and layers).
- Delocalised electrons that are free to move throughout the entire structure.
- The delocalised electrons came from the outer shells of the metal atoms.
The delocalised electrons act as a "glue" — the attraction between the positive ions and the electrons holds the structure together.
Properties of Metals Explained
High Melting and Boiling Points
- The strong electrostatic attraction between the positive ions and the sea of delocalised electrons requires a large amount of energy to overcome.
- Therefore, metals generally have high melting and boiling points.
- Metals with more delocalised electrons or higher charges on the ions tend to have even higher melting points (e.g. iron has a higher melting point than sodium).
| Metal | Melting Point (°C) | Outer Electrons (delocalised) |
|---|
| Sodium (Na) | 98 | 1 |
| Magnesium (Mg) | 650 | 2 |
| Aluminium (Al) | 660 | 3 |
| Iron (Fe) | 1538 | 2 (transition metal) |
Exam Tip: When comparing melting points of metals, consider: (1) the number of delocalised electrons per atom, and (2) the charge on the metal ion. More delocalised electrons and higher charge = stronger metallic bonding = higher melting point.
Electrical Conductivity
Metals are good conductors of electricity.
Why?
- The delocalised electrons are free to move throughout the metal structure.
- When a voltage is applied, these electrons can flow in one direction, carrying electrical charge.
- This is an electric current.
Exam Tip: State that delocalised electrons are "free to move through the structure" and can "carry charge." Don't just say "free electrons" — explain that they can move to carry charge.
Thermal Conductivity
Metals are good conductors of heat (thermal energy).
Why?
- When one end of a metal is heated, the delocalised electrons gain kinetic energy.
- These energetic electrons move quickly through the metal and transfer energy to other parts of the structure by colliding with metal ions.
- Energy is rapidly transferred through the metal.
Malleability and Ductility
Metals are malleable (can be hammered into shape) and ductile (can be drawn into wires).
Why?
- The layers of metal ions can slide over each other when a force is applied.
- Crucially, the delocalised electrons can move with the ions and continue to hold the structure together after the layers have shifted.
- The metallic bond is not broken when the layers slide — the sea of electrons re-adjusts.
This is very different from ionic compounds, where displacement of layers causes like-charged ions to be adjacent, resulting in repulsion and shattering (brittleness).
Alloys
An alloy is a mixture of two or more elements, at least one of which is a metal.
Why Are Alloys Harder Than Pure Metals?
In a pure metal, all the atoms (ions) are the same size, so the layers are regular and can slide over each other easily.
In an alloy, atoms of different sizes are introduced into the lattice. These differently-sized atoms disrupt the regular arrangement of layers, making it more difficult for the layers to slide.
This is why alloys are generally harder and stronger than pure metals.
Common Alloys
| Alloy | Made From | Property / Use |
|---|
| Steel | Iron + carbon (+ other metals) | Stronger than pure iron — used in construction, tools |
| Stainless steel | Iron + chromium + nickel | Resistant to corrosion — cutlery, surgical instruments |
| Bronze | Copper + tin | Hard and resistant to corrosion — statues, bells |
| Brass | Copper + zinc | Hard, gold-coloured — musical instruments, fittings |
| Gold alloys (e.g. 9 carat) | Gold + copper/silver/zinc | Harder than pure gold — jewellery |
Exam Tip: A common question asks: "Explain why alloys are harder than pure metals." The key answer is: atoms of different sizes disrupt the regular layers, preventing them from sliding over each other.
Comparing Metallic Bonding with Ionic Bonding
| Feature | Metallic Bonding | Ionic Bonding |
|---|
| Between | Metal atoms (metal + metal) | Metal and non-metal |
| Particles involved | Positive ions + delocalised electrons | Positive and negative ions |
| Type of attraction | Electrostatic (ions ↔ electrons) | Electrostatic (cation ↔ anion) |
| Conduct electricity (solid) | Yes (delocalised electrons) | No (ions in fixed positions) |
| Malleable | Yes (layers slide, electrons re-adjust) | No (brittle — layers shift, same-charge repulsion) |
Worked Example: Comparing Sodium and Iron
Question: Iron has a much higher melting point (1538 °C) than sodium (98 °C). Both are metals. Explain this difference.
Answer:
- Both iron and sodium have metallic bonding — a strong electrostatic attraction between positive metal ions and a sea of delocalised electrons.
- Sodium atoms have 1 outer electron that is delocalised, forming Na⁺ ions with a +1 charge.
- Iron atoms contribute 2 delocalised electrons (as a transition metal), forming Fe²⁺ ions with a +2 charge.
- Iron has more delocalised electrons per atom and a higher charge on its ions.
- This means the electrostatic attraction between the ions and the sea of electrons is stronger in iron.
- Therefore, more energy is needed to overcome the metallic bonding in iron, giving it a much higher melting point.
Key Points
- Metallic bonding is the strong electrostatic attraction between positive metal ions and a sea of delocalised electrons.
- Metals have high melting points because of the strong metallic bonds.
- Metals conduct electricity because delocalised electrons are free to move and carry charge.
- Metals conduct heat because delocalised electrons transfer kinetic energy through the structure.
- Metals are malleable and ductile because layers can slide without breaking the metallic bond.
- Alloys are harder than pure metals because differently-sized atoms disrupt the layers, preventing sliding.
Practice Questions
- Describe metallic bonding using the terms "positive ions" and "delocalised electrons."
- Explain why magnesium has a higher melting point than sodium.
- Explain why copper can conduct electricity.
- Explain why metals can be bent into shape without breaking.
- Explain why adding tin to copper (to make bronze) makes the metal harder.
- Aluminium is used to make aircraft bodies. Suggest why aluminium is suitable for this purpose in terms of its properties.
Worked Example: Why Are Metals Malleable but Ionic Compounds Brittle?
Question (6 marks): Copper is malleable (can be hammered into shape) while sodium chloride is brittle (shatters under force). Explain this difference in terms of bonding and structure.
Answer:
Copper (metallic bonding):
- Copper has a giant metallic structure — layers of positive Cu²⁺ ions in a sea of delocalised electrons.
- When a force is applied, the layers of ions can slide over each other.
- Crucially, the delocalised electrons move with the ions and continue to hold the structure together.
- The metallic bond is not broken when the layers slide, so copper can be bent, hammered and shaped without breaking.