Metallic Bonding

Setup. In a metal:

1. Each atom donates its outer-shell electrons to a 'sea' shared by all the metal ions.
2. The metal atoms become positive ions (cations), arranged in a regular lattice.
3. The delocalised electrons can move freely throughout the metal.
4. The bonding is the strong electrostatic attraction between the positive ions and the negative sea.

Two parts to remember:

• Lattice of positive ions (where atoms once were).
• Sea of delocalised electrons (the donated outer electrons, free to roam).

This bonding is in 3D. Layers of ions stacked together, with electrons moving in the spaces between.

Cambridge tip. When asked to "describe the bonding in a metal", say BOTH:

• Lattice of positive ions.
• Surrounded by a sea of delocalised electrons.
• Strong electrostatic attraction between them.

1. Electrical conductivity. Delocalised electrons can move through the lattice and carry charge. Metals conduct as solids AND in molten form (no need to melt them, unlike ionic compounds).

2. Thermal conductivity. Delocalised electrons also transfer kinetic energy quickly through the metal. That's why metals feel cold when you touch them — they conduct heat away from your skin.

3. Malleability and ductility.

• Malleable: can be hammered into thin sheets.
• Ductile: can be drawn into wires.

When force is applied, the layers of metal ions slide over each other. The delocalised electrons follow along — the metallic bonding is preserved. The metal deforms WITHOUT shattering. Compare with ionic compounds, which crack along cleavage planes when ions are displaced.

4. High melting and boiling points. Strong electrostatic attraction across the entire lattice → lots of energy needed to break it. Most metals melt at $> 600\degree\text{C}$.

• Notable exception: mercury (m.p. $-39\degree\text{C}$) — only liquid metal at room temp.

5. Shiny appearance. Delocalised electrons absorb and re-emit photons of all visible wavelengths → silvery reflective lustre.

6. High density. Atoms tightly packed in the lattice.

Worked qualitative. Why is copper a good electrical wire material? It has many delocalised electrons (high conductivity), is ductile (can be drawn into wires), is malleable (easily worked), and is reasonably cheap. It also forms an oxide layer (verdigris) that protects underlying copper from further corrosion.

Worked qualitative. Why does an iron rod feel cold when you touch it but a wooden stick at the same temperature feels warm? Iron's delocalised electrons quickly conduct heat AWAY from your fingers. Wood is a poor conductor; little heat flows.

Alloy. A mixture of a metal with one or more other elements (often another metal).

Why alloy? Pure metals are often too soft for engineering uses. Alloying makes them harder by DISRUPTING the regular layer arrangement: when atoms of different sizes are mixed in, the layers can't slide as easily.

Common alloys.

• Brass = copper + zinc. Used for instruments, fittings.
• Bronze = copper + tin. Used for statues, bearings.
• Steel = iron + carbon (and trace others). Stronger than iron.
• Stainless steel = iron + chromium + nickel. Resistant to corrosion.

Worked qualitative. Why is steel harder than pure iron? Carbon atoms are smaller than iron atoms; they fit in the gaps in the iron lattice and prevent layers from sliding past each other.

Cambridge tip. When asked "explain why an alloy is harder than the pure metal", mention DIFFERENT-SIZED atoms disrupting layer sliding.