Hybridisation

Energy Minimization: Atoms undergo hybridisation because the resulting geometry allows valence electrons to be as far apart as possible, minimizing electrostatic repulsion. This state is energetically more favorable than using unhybridised orbitals for bonding.

Orbital Character: Each type of hybrid orbital possesses a specific percentage of 's' and 'p' character. For instance, an $sp$ orbital is $50\%$ $s$ and $50\%$ $p$, while an $sp^3$ orbital is $25\%$ $s$ and $75\%$ $p$; higher $s$-character typically results in shorter, stronger bonds.

Promotion of Electrons: In many cases, such as carbon, an electron is 'promoted' from a filled $s$ orbital to an empty $p$ orbital before hybridisation occurs. This provides more unpaired electrons available for bonding, increasing the atom's valency.

To determine the hybridisation of a central atom, first count the number of electron domains (regions of high electron density) surrounding it. An electron domain is defined as either a single/double/triple bond or a lone pair of electrons.

Match the number of electron domains to the required number of hybrid orbitals: 2 domains require $sp$, 3 domains require $sp^2$, and 4 domains require $sp^3$. For expanded octets in the third period or below, 5 domains require $sp^3d$ and 6 domains require $sp^3d^2$.

Identify the role of unhybridised orbitals: Any $p$ orbitals not used in hybridisation remain available to form pi ($\pi$) bonds. For example, an $sp^2$ hybridised atom has one unhybridised $p$ orbital, which it uses to form one $\pi$ bond in a double bond system.

Sigma vs. Pi Bonds: Sigma bonds are formed by the head-on overlap of hybrid orbitals (or $s$ orbitals), whereas pi bonds are formed by the lateral (side-on) overlap of unhybridised $p$ orbitals. Sigma bonds allow for free rotation, while pi bonds restrict it.

Electron vs. Molecular Geometry: Hybridisation determines the electron domain geometry. The actual molecular geometry (the shape of the atoms) may differ if lone pairs are present, as lone pairs occupy hybrid orbitals but are not 'seen' as part of the molecular shape.

The 'Domain Count' Rule: Always count a double or triple bond as a single electron domain when determining hybridisation. A common mistake is counting each individual bond in a multiple bond as a separate domain.

Lone Pair Inclusion: Never forget to include lone pairs in your domain count. For example, in ammonia ($NH_3$), the nitrogen has three bonding pairs and one lone pair, totaling four domains, which necessitates $sp^3$ hybridisation.

Predicting Bond Angles: Use the hybridisation type to predict the base bond angle, then adjust slightly for lone pairs. Lone pairs exert more repulsion than bonding pairs, typically compressing the angles between bonding atoms (e.g., $107^{\circ}$ in $NH_3$ instead of $109.5^{\circ}$).

Multiple Bonds and Hybridisation: Students often think that a triple bond requires $sp^3$ hybridisation because it involves many electrons. In reality, a triple bond requires $sp$ hybridisation because it only occupies one directional domain, leaving two $p$ orbitals free for $\pi$ bonding.

Hybridisation of Terminal Atoms: While we usually focus on the central atom, terminal atoms (like Oxygen in $C=O$) can also be hybridised. If an oxygen atom has one double bond and two lone pairs, it is $sp^2$ hybridised.

Energy Levels: It is a misconception that hybrid orbitals exist in isolated atoms. Hybridisation is a model that describes the state of an atom only when it is bonding with other atoms.