Review the key concepts, formulae, and examples before starting your quiz.
🔑Concepts
Metallic bonding is defined as the strong electrostatic attraction between a lattice of positive metal ions (cations) and a 'sea' of delocalised electrons.
Metals have a giant metallic lattice structure. The valence electrons are not bound to specific atoms but are free to move throughout the entire structure, which is why they are termed 'delocalised'.
Metals are excellent conductors of electricity because the delocalised electrons are free to move and carry an electrical charge when a potential difference is applied.
Thermal conductivity in metals is high because delocalised electrons and closely packed ions can efficiently transfer kinetic energy through the lattice.
Metals are malleable (can be hammered into sheets) and ductile (can be drawn into wires). This is because the layers of positive ions can slide over each other while the 'sea' of electrons maintains the binding force, preventing the structure from shattering.
The melting and boiling points of metals are generally high because of the strong electrostatic forces between the positive ions and the delocalised electrons, which require significant energy to overcome.
An alloy is a mixture of a metal with one or more other elements (usually other metals or carbon). Alloys like Steel or Brass are harder than pure metals because atoms of different sizes disrupt the regular arrangement of the lattice, making it harder for layers to slide over each other.
📐Formulae
💡Examples
Problem 1:
Explain why Magnesium () has a higher melting point than Sodium ().
Solution:
Magnesium ions have a higher charge () compared to Sodium ions () and contribute more delocalised electrons per atom to the 'sea'. Additionally, the ion is smaller than the ion.
Explanation:
Because the ions have a higher charge density and there are more delocalised electrons, the electrostatic attraction between the cations and the delocalised electrons is much stronger in than in . More energy is therefore required to break these bonds.
Problem 2:
Describe why the addition of Carbon () to Iron () to make Steel results in a material that is less malleable than pure Iron.
Solution:
The Carbon atoms have a different atomic radius than the Iron atoms. When they are introduced into the lattice, they occupy spaces between Iron atoms or replace them, disrupting the regular, layered arrangement of the atoms.
Explanation:
In pure Iron, the layers of atoms are of uniform size and can easily slide over one another when force is applied. In Steel, the different-sized Carbon atoms 'lock' the layers in place, preventing them from sliding easily, which increases hardness and reduces malleability.
Problem 3:
Why do metals conduct electricity in both solid and liquid states?
Solution:
In both the solid and molten (liquid) states, the delocalised electrons are free to move throughout the structure.
Explanation:
Unlike ionic compounds, which only conduct when molten or aqueous because the ions themselves must move, metals rely on delocalised electrons which remain mobile regardless of whether the metal is a solid or a liquid.