the metallic model

Metals are found on the left side of the periodic table. They have a small number of electrons in their outer shell, and therefore have relatively low ionisation energies. They typically react by losing their valence electrons and forming positive ions (cations).

Metallic character refers to the tendency of metal atoms to lose control of their outer shell (valence) electrons. In the elemental state, when there is no non-metal available to accept electrons and form an ionic compound, these outer electrons are held only loosely by the metal nucleus and tend to become delocalised.

In a pure metallic element, the valence electrons are not associated closely with any one nucleus. Instead, they become delocalised and are free to move throughout the entire metal structure. The metal atoms that lose these electrons become cations, arranged in a regular lattice.

A metallic bond is the electrostatic attraction between a lattice of cations and delocalised electrons.

in the metallic model, a metal is described as a lattice of positive ions (cations) surrounded by a “sea” of delocalised electrons. The electrons can move freely through the structure, and the electrostatic attraction between these electrons and the cations holds the lattice together.

Metallic bonding is called non-directional because the electrostatic attraction acts in all directions around each cation toward the delocalised electrons. The electrons are not localised between specific pairs of ions, so the bond is not fixed along particular axes.

Metals conduct electricity well because the delocalised electrons are highly mobile. When a potential difference is applied, these electrons can move freely through the metal structure and carry charge.

In the metallic lattice, the delocalised electrons are free to move throughout the structure. The metallic bond does not restrain them to any particular ion, so they can flow under an applied voltage, allowing metals to act as good electrical conductors (e.g. copper in electrical circuits).

Metals conduct heat well because delocalised electrons and closely packed ions enable efficient transfer of thermal energy. Electrons can gain kinetic energy in hotter regions and transfer it rapidly through the lattice, while closely packed cations also help transfer vibrational energy.

Metals are malleable because the movement of delocalised electrons is non-directional and random through the cation lattice. When a force is applied and the atoms change position, the metallic bond remains intact, as cations are still surrounded by delocalised electrons. Layers of ions can slide without breaking the bonding.

Metals are ductile (can be drawn into wires) for the same reason they are malleable: the non-directional metallic bond allows layers of cations to move past each other while remaining bonded by the surrounding delocalised electrons. The structure can change shape without fracturing.

Metals usually have high melting points because a large amount of energy is required to overcome the strong electrostatic forces of attraction between the lattice of cations and the delocalised electrons. Strong metallic bonding → high melting and boiling points

Metals are shiny and lustrous because delocalised electrons in the metal crystal structure can reflect light. Incoming photons are absorbed and re-emitted at the surface, giving metals their characteristic shiny appearance (e.g. jewellery, decorative structures).

Metals typically show: - Good electrical conductivity - Good thermal conductivity - Malleability - Ductility - High melting points - Shiny, lustrous appearance All these can be explained by the metallic model: a lattice of cations and delocalised electrons with non-directional electrostatic attraction.

The strength of a metallic bond is determined by: - The number of delocalised electrons per atom - The charge on the metal cation - The radius of the cation Greater electron density and smaller cation radius → stronger metallic bonding.

• Na: 1s²2s²2p⁶3s¹ → forms Na⁺ with 1 delocalised electron per atom • Mg: 1s²2s²2p⁶3s² → forms Mg²⁺ with 2 delocalised electrons per atom Mg²⁺ is also smaller than Na⁺. So Mg has more delocalised electrons and smaller, more highly charged cations, giving it stronger metallic bonds and a higher melting point.

Melting points: - Na: 98 °C - Mg: 650 °C The much higher melting point of Mg confirms that metallic bonding is stronger in magnesium than in sodium, consistent with Mg²⁺ having a greater charge, smaller radius, and more delocalised electrons.

Metals with stronger metallic bonds (higher melting points) tend to be less reactive, as more energy is required to disrupt their structure or remove electrons. Metals with weaker metallic bonds (lower melting points) tend to be more reactive, especially in reactions where they lose electrons (e.g. Group 1 metals down the group).

Down a group, the melting point generally decreases because the radius of the cation increases, which weakens the attraction between the delocalised electrons and the cations, reducing metallic bond strength.

Down a group, the size of the cation increases. The delocalised electrons are further from the nucleus, so the electrostatic attraction between the cations and electrons is weaker. Therefore, metallic bond strength decreases down a group, and melting points decrease.

Melting points: - Na: 98 °C - K: 63 °C - Rb: 39 °C The decrease in melting point down Group 1 shows that the metallic bond becomes weaker because the cations become larger, reducing the attraction to the delocalised electrons.

Across a period (for the s and p block metals), the melting point generally increases from left to right as the charge on the cation and the number of delocalised electrons increase, leading to stronger metallic bonding.

Transition elements are elements whose atoms have an incomplete d sublevel, or can give rise to cations with an incomplete d sublevel.

Metals with highly charged, small cations and many delocalised electrons form the strongest metallic bonds. That means small ionic radius + large positive charge → high charge density → strong electrostatic attraction.