Revise Metallic bonding for Cambridge 9701 (3.3)

Metallic bonding is found in metallic elements and alloys. It consists of a giant lattice structure. Unlike ionic lattices, which have alternating positive and negative ions, a metallic lattice consists only of positive ions (cations). These cations are formed when metal atoms lose their outer-shell electrons. These electrons are no longer bound to a specific atom; instead, they become delocalised and form a mobile 'sea' of electrons that surrounds the fixed positive ions. The bonding itself is the strong, non-directional electrostatic force of attraction between the positive cations and the negative delocalised electrons.

The 'sea of electrons' model is incredibly successful at explaining the characteristic properties of metals.

Electrical and Thermal Conductivity: For a substance to conduct electricity, it must contain mobile charged particles. In metals, the delocalised electrons are these particles. When a voltage is applied, the electrons are attracted to the positive terminal and flow through the lattice, carrying charge and creating a current. Similarly, metals are good thermal conductors because the mobile electrons can quickly transfer kinetic energy throughout the lattice. The close packing of ions also allows for efficient energy transfer via vibrations.

Malleability and Ductility: Malleability is the ability to be hammered into shape, while ductility is the ability to be drawn into a wire. When a force is applied to a metal, layers of positive ions can slide over one another. The delocalised electrons are able to move with the ions, maintaining the electrostatic attraction and preventing the structure from shattering. This is in stark contrast to ionic compounds, which are brittle and fracture when layers are displaced, causing ions of like charge to repel each other.

The strength of the metallic bond determines a metal's melting point, boiling point, and hardness. The strength depends on the magnitude of the electrostatic attraction between the cations and the delocalised electrons. Two key factors are the charge on the ion and its size.

$Strength of attraction \propto \frac{\text{charge on ion}}{\text{ionic radius}}$

A higher charge on the cation (e.g., Mg$^{2+}$ vs Na$^{+}$) leads to a stronger attraction to the electrons. Additionally, more electrons are contributed to the delocalised sea per atom (two from Mg vs one from Na). A smaller ionic radius means the positive charge is more concentrated (higher charge density) and the electrons are closer to the nucleus, resulting in a stronger attraction. Therefore, metals with smaller, more highly charged ions will have stronger metallic bonds and higher melting points.

An alloy is a mixture containing a metal and at least one other element. Common examples include brass (copper and zinc), bronze (copper and tin), and steel (iron and carbon). Pure metals have a regular, ordered lattice structure. In an alloy, the atoms of the different elements have different sizes. This disrupts the regular arrangement of the layers of ions. Consequently, it becomes much more difficult for the layers to slide over one another. This makes most alloys harder, stronger, and less malleable than their constituent pure metals.