Metallic Bonding

Aluminium has a giant metallic structure: a regular 3D lattice of Al³⁺ ions surrounded by a 'sea' of delocalised electrons. The strong electrostatic attraction between these ions and the delocalised electrons is the metallic bond. Every metallic property below comes back to this picture.

1. Good electrical conductor. Electricity is carried by delocalised electrons that are NOT bound to any specific atom. When connected to a potential difference, the electrons drift through the lattice from one end to the other, carrying charge → aluminium conducts electricity. This makes it suitable for the overhead power cables that transport electricity from power stations to homes. (Aluminium is preferred over copper for overhead use because it is LIGHTER for the same conductivity.)

2. Malleable and ductile. When a force is applied to aluminium, the layers of Al³⁺ ions can slide past each other without breaking the metallic bond — the delocalised electrons flow with the ions and maintain the electrostatic attraction. So aluminium can be:

• Hammered into thin sheets (malleable) — used to make drinks-can bodies and food foil.
• Drawn into thin wires (ductile) — used for power cables.

3. Low density (for a metal). Aluminium has a relatively low density (~2.7 g/cm³) compared to other structural metals like iron (~7.9 g/cm³). This is partly because of the small Al³⁺ ions and partly because of their packing. Low density + reasonable strength → aircraft bodies are made of aluminium alloys (e.g. duralumin = Al + Cu).

4. High melting point. Aluminium melts at 660 °C — much higher than room temperature. The strong electrostatic attraction between Al³⁺ ions (large 3+ charge) and the delocalised electrons requires a great deal of energy to overcome. So aluminium remains solid at all normal working temperatures.

5. (Bonus) Resistant to corrosion. Aluminium quickly forms a thin layer of Al₂O₃ on its surface, which is impermeable and protects the metal underneath. This is why aluminium is used for drinks cans and outdoor structures — though this is a chemical property, not a metallic-bond property.

(Any FOUR of the above, each with a metallic-bond-based explanation, gives 8 marks.)

Conclusion. Substance X has metallic bonding (and a giant metallic structure). [Most likely: copper, mp = 1085 °C.]

Justification — each observation in turn.

(1) Shiny when freshly cut. Metals are lustrous because their surface delocalised electrons reflect light. Ionic and covalent substances tend to be dull (white salts) or coloured powders. The shine immediately suggests a metal.

(2) Malleable. When force is applied, layers of positive ions can slide past each other without breaking the metallic bond — the delocalised electrons flow with them and maintain the electrostatic attraction. Ionic compounds are HARD but BRITTLE (they shatter when layers slide because like-charge ions end up adjacent). Simple molecular and giant covalent substances do not deform like metals at all. So malleability strongly indicates metallic bonding.

(3) Conducts when SOLID and when molten. This is the decisive observation:

• An IONIC compound conducts ONLY when molten or aqueous (ions free to move) — NOT as a solid.
• A SIMPLE MOLECULAR substance does NOT conduct in any state (no charge carriers).
• A GIANT COVALENT substance generally does NOT conduct, with graphite being the only exception.
• A METAL conducts in BOTH solid and molten states because the delocalised electrons are free to move at all times (no need to break the lattice).

The fact that X conducts as a solid rules out ionic and most non-metals → metallic.

(4) Melting point 1085 °C. High mp is consistent with strong bonding throughout the structure. Both metals and ionic compounds (and giant covalent) can have high mp — so this on its own doesn't distinguish — but a value around 1000 °C is typical for a metal (combined with the other clues, metallic bonding is confirmed).

Identity guess. Copper's mp is 1085 °C, so X is very likely copper.

Definition of alloy. An alloy is a mixture of a metal with one or more other elements (usually other metals, but sometimes non-metals like carbon). The components are mixed in the molten state, then allowed to solidify. Alloys are NOT compounds — there is no chemical reaction; the metals retain their individual identities, but the atoms of different sizes disrupt the regular lattice, making the alloy harder and often stronger than the pure metal.

Three examples.

1. Steel = iron + carbon (≤ 2 % C).

• Pure iron is relatively soft and easily corroded.
• Carbon atoms are smaller than Fe atoms and slot into the Fe lattice, disrupting the regular arrangement. Layers of Fe atoms can no longer slide easily → steel is much harder and stronger than pure iron.
• Use: structural use in buildings, bridges, car bodies. The combination of strength and relatively low cost makes steel the most-used alloy in the world.

2. Brass = copper + zinc.

• Pure copper is too soft for many uses.
• Zn atoms are slightly larger than Cu atoms. In brass, the differently-sized atoms disrupt the regular lattice → brass is harder and stronger than pure copper.
• Use: door handles, plumbing fittings, musical instruments (trumpets, trombones), decorative items — combines hardness with corrosion resistance and an attractive golden colour.

3. Stainless steel = iron + chromium (≥ 10.5 %) + nickel.

• Pure iron rusts when exposed to water and oxygen.
• Adding chromium causes a thin, impermeable layer of Cr₂O₃ to form on the surface — this protects the iron underneath. Nickel adds further strength and corrosion resistance.
• Use: cutlery, sinks, medical instruments, food-preparation surfaces — any application needing corrosion resistance in a wet environment.

(Other valid alloys: bronze = Cu + Sn (statues, bearings); duralumin = Al + Cu (aircraft bodies); solder = Sn + Pb (joining electronic components).)

Structure. A metal has a regular lattice of positive ions surrounded by a sea of delocalised electrons.

Mechanism of heat conduction.

1. When one end of a metal is heated, the positive ions at that end gain kinetic energy and vibrate more vigorously in their lattice positions.
2. The vibrating ions transfer some of this kinetic energy to the delocalised electrons in their vicinity. The electrons gain kinetic energy and move FASTER.
3. Because the delocalised electrons are free to move throughout the lattice, they carry their increased kinetic energy rapidly to other parts of the metal.
4. When these fast-moving electrons collide with cooler ions elsewhere in the metal, they transfer kinetic energy to them — heating up that part of the metal.
5. The energy transfer is fast and efficient because:
   • Delocalised electrons are very mobile (much faster than vibrating ions).
   • The lattice is held together by strong electrostatic forces, so vibrations of one ion are quickly passed to neighbouring ions (lattice-vibration conduction is also present, but less important).

Result. Heat travels rapidly along the metal — making metals excellent thermal conductors. Non-metals (e.g. wood, plastic, glass) lack delocalised electrons, so they rely on slow lattice vibrations alone → poor thermal conductors (insulators).

Saucepan use. A copper or aluminium base conducts heat from the stove evenly across the bottom of the pan, ensuring food is cooked uniformly. Plastic handles are added because plastic is a thermal INSULATOR (poor heat conduction) — the user can hold the pan without burning their hand.