Example 2 $sp$ 0 Linear $180^\circ$ $BeCl_2$, $CO_2$ 3 $sp^2$ 0 Trigonal Planar $120^\circ$ $BF_3$, $SO_3$ 3 $sp^2$ 1 Bent $<120^\circ$ $SO_2$, $O_3$ 4 $sp^3$ 0 Tetrahedral $109.5^\circ$ $CH_4$, $CCl_4$ 4 $sp^3$ 1 Trigonal Pyramidal $107^\circ$ $NH_3$, $PCl_3$ 4 $sp^3$ 2 Bent $104.5^\circ$ $H_2O$, $H_2S$ 5 $sp^3d$ 0 Trigonal Bipyramidal $90^\circ, 120^\circ$ $PCl_5$ 5 $sp^3d$ 1 Seesaw $<90^\circ, <120^\circ$ $SF_4$ 5 $sp^3d$ 2 T-shaped $<90^\circ$ $ClF_3$ 5 $sp^3d$ 3 Linear $180^\circ$ $XeF_2$, $I_3^-$ 6 $sp^3d^2$ 0 Octahedral $90^\circ$ $SF_6$ 6 $sp^3d^2$ 1 Square Pyramidal $<90^\circ$ $BrF_5$ 6 $sp^3d^2$ 2 Square Planar $90^\circ$ $XeF_4$

3.3 Valence Bond Theory & Hybridization

Hybridization Formula: $H = \frac{1}{2}(V + M - C + A)$
Where $V$ = valence electrons of central atom, $M$ = number of monovalent atoms, $C$ = charge of cation, $A$ = charge of anion.

If $H=2 \rightarrow sp$, $H=3 \rightarrow sp^2$, $H=4 \rightarrow sp^3$, $H=5 \rightarrow sp^3d$, $H=6 \rightarrow sp^3d^2$, $H=7 \rightarrow sp^3d^3$.
Sigma ($\sigma$) Bonds: Head-on overlap (stronger).
Pi ($\pi$) Bonds: Sideways overlap (weaker).
- Single bond: $1 \sigma$.
- Double bond: $1 \sigma, 1 \pi$.
- Triple bond: $1 \sigma, 2 \pi$.

3.4 Molecular Orbital Theory (MOT)

Bond Order (BO):
$BO = \frac{1}{2}(N_b - N_a)$

Where $N_b$ = number of electrons in bonding MOs, $N_a$ = number of electrons in antibonding MOs.

If $BO