Summary

Kössel-Lewis approach: Atoms achieve stable noble-gas configurations by sharing (covalent bond) or transferring (ionic bond) electrons. The octet rule (8 valence electrons) is the guiding principle, with exceptions like $H$ (duplet), $Li$, $Be$, $B$ (incomplete octet), and expanded octets ($PCl_5$, $SF_6$).

Lewis dot structures: Represent valence electrons as dots around the element symbol. Formal charge = $V - (L + \frac{1}{2}S)$, where $V$ = valence electrons, $L$ = lone-pair electrons, $S$ = shared electrons. The most stable structure has the smallest formal charges (preferably zero) and negative charge on the more electronegative atom.

Ionic bond: Electrostatic attraction between oppositely charged ions. Lattice energy ($U$) depends on ion charges ($Q_1, Q_2$) and interionic distance ($r$): $U \propto \frac{Q_1 Q_2}{r}$. Higher charge and smaller size → stronger lattice → higher melting point.

Covalent bond: Formed by sharing of electron pairs. Bond parameters: bond length (equilibrium distance between nuclei), bond angle, bond enthalpy (energy required to break one mole of bonds in gaseous state), and bond order ($\frac{1}{2}[\text{bonding electrons} - \text{antibonding electrons}]$). Higher bond order → shorter, stronger bond.

Valence Bond Theory (VBT): A covalent bond forms when atomic orbitals overlap. Greater overlap → stronger bond. Sigma ($\sigma$) bonds result from end-to-end overlap (s-s, s-p, p-p along the internuclear axis); pi ($\pi$) bonds from sidewise overlap of parallel p-orbitals. A single bond is one $\sigma$; a double bond is one $\sigma$ + one $\pi$; a triple bond is one $\sigma$ + two $\pi$.

Hybridisation: Mixing of atomic orbitals of the same atom to form equivalent hybrid orbitals. Key types: $sp$ (linear, 180°), $sp^2$ (trigonal planar, 120°), $sp^3$ (tetrahedral, 109.5°), $sp^3d$ (trigonal bipyramidal, 90°/120°), $sp^3d^2$ (octahedral, 90°). The number of hybrid orbitals equals the number of $\sigma$-bonds + lone pairs on the central atom.

Molecular Orbital Theory (MOT): Atomic orbitals combine to form molecular orbitals (MOs) — bonding (lower energy) and antibonding (higher energy). For homonuclear diatomic molecules:

• For $Li_2$ to $N_2$: $\sigma_{1s} < \sigma^*_{1s} < \sigma_{2s} < \sigma^*_{2s} < \pi_{2p_x} = \pi_{2p_y} < \sigma_{2p_z} < \pi^*_{2p_x} = \pi^*_{2p_y} < \sigma^*_{2p_z}$
• For $O_2$, $F_2$, $Ne_2$: $\sigma_{1s} < \sigma^*_{1s} < \sigma_{2s} < \sigma^*_{2s} < \sigma_{2p_z} < \pi_{2p_x} = \pi_{2p_y} < \pi^*_{2p_x} = \pi^*_{2p_y} < \sigma^*_{2p_z}$
• Bond order = $\frac{N_b - N_a}{2}$. $O_2$ has bond order 2 and is paramagnetic (two unpaired electrons in $\pi^*$ orbitals).

VSEPR Theory: Predicts molecular geometry based on repulsion between electron pairs (bonded and lone) around the central atom. Lone pair–lone pair repulsion > lone pair–bond pair > bond pair–bond pair. Examples: $CH_4$ (tetrahedral), $NH_3$ (trigonal pyramidal), $H_2O$ (bent), $SF_4$ (see-saw), $ClF_3$ (T-shaped), $XeF_4$ (square planar). …