Structure 2. Models of bonding and structure - From models to materials

Metallic bonding is the electrostatic attraction between a lattice of positive metal ions (cations) and a 'sea' of delocalized valence electrons. The strength of the bond increases with the number of valence electrons and decreasing ionic radius, e.g., $Mg^{2+}$ has a higher melting point than $Na^{+}$ because of higher charge density.

Ionic bonding results from the electrostatic attraction between oppositely charged ions in a giant lattice. The strength of the ionic bond is determined by Coulomb's Law, where the force of attraction is proportional to the product of charges ($q_1 q_2$) and inversely proportional to the square of the distance between them ($r^2$).

Alloys are mixtures of metals with other elements. They are often harder than pure metals because atoms of different sizes disrupt the regular lattice layers, preventing them from sliding over each other easily. Examples include steel and brass.

Giant covalent structures like diamond, graphite, and $SiO_2$ involve atoms linked by covalent bonds in a continuous network. In diamond, each carbon is $sp^3$ hybridized, forming a tetrahedral structure, while in graphite, carbon is $sp^2$ hybridized, forming layers held by London dispersion forces.

Molecular materials are held together by intermolecular forces (IMFs) such as London dispersion forces, dipole-dipole attractions, and hydrogen bonding. Hydrogen bonding occurs when hydrogen is covalently bonded to highly electronegative elements ($N$, $O$, or $F$).

The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular geometry based on the principle that electron domains (bonding and non-bonding pairs) around a central atom stay as far apart as possible to minimize repulsion.

Formal charge ($FC$) is used to determine the most stable Lewis structure among resonance contributors. The structure where the $FC$ of each atom is closest to zero is generally preferred.