Metallic Bonding Exam Questions

• Metal atoms lose their outer electrons to become positive ions / cations [1]
• The electrons become delocalised / form a “sea” of free electrons [1]
• There is a strong electrostatic attraction between the positive metal ions and the delocalised electrons — this is the metallic bond [1]

Examiner note: The phrase “sea of electrons” is acceptable but must be linked to electrostatic attraction. Writing “atoms share electrons” is wrong — that describes covalent bonding. The electrons are delocalised across the whole structure, not shared between two atoms.

• Metals have delocalised / free electrons [1]
• These electrons can move through the metal structure [1]
• When a potential difference is applied, the electrons flow in one direction, carrying charge / creating a current [1]

Examiner note: “Metals have free electrons” alone is worth 1 mark. You must explain that these electrons can move when a voltage is applied. Do not confuse with ionic conduction — metals conduct through electron movement, not ion movement.

• In metals, the positive ions are arranged in layers [1]
• When a force is applied, the layers can slide over each other [1]
• The metallic bond is maintained because the delocalised electrons can move with the ions / the sea of electrons continues to hold the ions together in the new positions [1]

Examiner note: Contrast this with ionic compounds: in a metal, the bonding is non-directional, so shifting layers does not create repulsion. In an ionic lattice, shifting layers brings like charges together, causing the material to shatter.

• Metallic: positive metal ions and delocalised electrons; Ionic: positive and negative ions [1]
• Metallic: electrostatic attraction between positive ions and the sea of delocalised electrons; Ionic: electrostatic attraction between oppositely charged ions [1]
• Metals conduct electricity when solid because delocalised electrons are free to move [1]
• Ionic compounds only conduct when molten or dissolved because ions must be free to move [1]

Examiner note: Both bonding types involve electrostatic attraction, but the particles involved differ. The key property difference tested is conductivity in the solid state: metals conduct (free electrons), ionic solids do not (ions fixed in lattice).

• Metals have a giant metallic structure [1]
• The strong electrostatic attraction between the positive ions and the delocalised electrons requires a large amount of energy to overcome [1]

Examiner note: “Strong bonds” without specifying the nature of the attraction scores 0. State “electrostatic attraction between positive ions and delocalised electrons.”

• The delocalised electrons gain kinetic energy when heated [1]
• They move faster and transfer kinetic energy to other parts of the metal by collisions [1]
• The closely packed positive ions can also transfer vibrations to neighbouring ions [1]

Examiner note: Thermal conductivity in metals has two components: electron transfer (the main mechanism) and lattice vibrations. Most marks come from the electron explanation. Saying “heat travels through the metal” is not a mechanism.

(a) Aluminium has more delocalised electrons per atom (3 vs 1) [1], so the electrostatic attraction between the positive ions (Al3+) and the larger sea of electrons is stronger / more energy needed to break [1]

(b) Sodium has fewer delocalised electrons and a larger ionic radius / lower charge on its ions (Na+ vs Al3+) [1], so the metallic bonds are weaker and the layers are more easily displaced / the metal is more easily deformed [1]

Examiner note: More delocalised electrons = stronger metallic bond = higher melting point and greater hardness. Higher charge on the ion also increases the attraction. Na+ (1+) vs Al3+ (3+) is the key comparison.