Bonding and properties in metals and alloys

Metallic bonding is the force that holds metal atoms together in a solid structure. Unlike ionic or covalent bonding, metallic bonding involves electrons that are free to move throughout the entire metal structure.

The Structure of Metals

Metals consist of a giant structure where atoms are arranged in a regular, repeating pattern (known as a lattice). The atoms are said to have regular arrangements because they have lots of atoms - exactly how many depends on how big the piece of metal is.

Key Features of Metal Structure:

• Atoms are arranged in a regular way (often in rows or layers)

• The structure is described as a "giant metallic structure" or "giant metallic lattice"

• Metals contain huge numbers of atoms all bonded together

Bonding in Metals

In metals, the electrons in the outer shells of the atoms are delocalised. This means that they aren't associated with a particular atom or bond, so they're free to move through the whole structure.

What Happens to the Atoms?

• Metal atoms lose their outer shell electrons

• They become positively charged metal ions

• These ions remain fixed in position in the lattice

What Happens to the Electrons?

• The outer electrons become delocalised (or "free")

• They form a "sea of electrons" that can move throughout the structure

• These mobile electrons are sometimes called "free electrons"

• Think of a ping pong balls in a box arranged in fixed positions that represent the positive ions, surrounded by jelly which represents the sea of delocalised electrons.

The Metallic Bond is defined as the electrostatic force of attraction between:

• The positive metal ions, and

• The delocalised (negative) electrons

These forces hold the metal structure together.

1. High Melting and Boiling Points

Metals have very high melting and boiling points because:

• The electrostatic forces between the metal ions and delocalised electrons are very strong

• A lot of energy is needed to break these forces

• Most compounds with metallic bonds have very high melting and boiling points, so they're generally solid at room temperature

Therefore, metals are solid at room temperature (except mercury).

2. Electrical and Thermal Conductivity

Electrical Conductivity: Metals conduct electricity well because:

• They have delocalised electrons that are free to move through the whole structure

• These electrons can carry electric charge through the metal

• The electrons act as charge carriers

Thermal Conductivity: Metals are good conductors of heat because:

• The delocalised electrons are free to move

• They can transfer thermal energy quickly through the structure

• The electrons carry the thermal energy through the metal

Science in Action: The high thermal conductivity of metals makes them ideal for cooking pots and pans. Steel and aluminium cookware needs to be durable, have a very high melting point and be unreactive so that the pan does not react with the food being cooked. The metal chosen for pans also needs to conduct heat well so that the heat transfers efficiently from the hob to the food being cooked.

3. Malleability

Metals are malleable - they can be bent, shaped, hammered or rolled into different forms without breaking.

Why are metals malleable?

• Metals consist of atoms held together in a regular structure

• The atoms form layers that are able to slide over each other

• When force is applied, the layers can move while the metallic bonding remains intact

• The delocalised electrons continue to hold the structure together even after the layers have moved

This means metals can be bent and shaped without shattering, and they can also be hammered or rolled into different shapes.

Pure metals often aren't quite suitable for certain jobs - they're often too soft when they're pure. To overcome this, metals are mixed with other metals (or sometimes other elements) to make them harder. These mixtures are called alloys.

Definition: An alloy is a mixture of two or more metals, or a metal and another element. Alloys are harder than pure metals.

Examples of Alloys:

Why Are Alloys Harder Than Pure Metals?

Different elements have different sized atoms. When another element is mixed with a pure metal, the new element atoms will distort (disrupt) the layers of metal atoms.

This distortion makes it more difficult for the layers to slide over each other, which is why alloys are harder than pure metals.

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