Metallic Bonding

Electrical Conductivity: Metals conduct electricity because the delocalized electrons are mobile and can carry a charge through the structure when a voltage is applied.

Thermal Conductivity: Heat energy is transferred quickly through metals via the kinetic energy of the mobile electrons and the vibrations of the closely packed ions.

Malleability and Ductility: Metals can be hammered into sheets or drawn into wires because the layers of ions can slide over each other without breaking the bond, as the 'sea' of electrons adjusts to the new positions.

High Melting/Boiling Points: The strong electrostatic forces require significant thermal energy to overcome, resulting in high transition temperatures for most metals.

Number of Delocalized Electrons: Metals with more valence electrons (e.g., $Mg^{2+}$ vs $Na^+$) contribute more electrons to the 'sea', increasing the density of the negative charge and the strength of the attraction.

Charge of the Metal Ion: A higher positive charge on the cation (e.g., $Al^{3+}$) results in a stronger electrostatic pull on the delocalized electrons compared to a lower charge.

Ionic Radius: Smaller metal ions allow the delocalized electrons to get closer to the positive nucleus, resulting in a stronger electrostatic attraction and higher melting points.

Precise Terminology: Always use the phrase 'electrostatic attraction between positive ions and delocalized electrons' to define the bond; avoid saying 'attraction between atoms'.

Explain Properties: When asked why metals are malleable, mention that layers of ions slide and the metallic bond is non-directional, so it does not break during deformation.

Comparison Questions: If comparing melting points (e.g., Sodium vs Magnesium), focus on the charge density: Magnesium has a higher charge ($2+$ vs $1+$) and more delocalized electrons per atom.

Diagram Checks: Ensure your sketches show a regular lattice of ions and that the electrons are clearly labeled as 'delocalized' or 'mobile'.