Chemical Bonding and Structure - Further aspects of covalent bonding and structure (HL only)

Sigma ($\sigma$) bonds are formed by the head-on (axial) overlap of atomic orbitals, resulting in electron density concentrated between the nuclei of the bonding atoms.

Pi ($\pi$) bonds are formed by the sideways overlap of parallel $p$ orbitals, resulting in electron density above and below the internuclear axis. A double bond consists of one $\sigma$ and one $\pi$ bond, while a triple bond consists of one $\sigma$ and two $\pi$ bonds.

Formal Charge ($FC$) is a tool used to determine the most stable Lewis structure. The preferred structure is the one where the formal charges on atoms are closest to zero.

Hybridization is the mixing of atomic orbitals (such as $s$ and $p$) to form new hybrid orbitals ($sp$, $sp^2$, $sp^3$) that are degenerate (equal in energy).

$sp^3$ hybridization results in tetrahedral geometry ($109.5^{\circ}$), $sp^2$ results in trigonal planar geometry ($120^{\circ}$), and $sp$ results in linear geometry ($180^{\circ}$).

Delocalization occurs when $\pi$ electrons are shared by more than two nuclei, often represented by resonance structures. This leads to intermediate bond lengths and increased stability (e.g., in $C_6H_6$ or $O_3$).

The catalytic depletion of ozone ($O_3$) by $NO_x$ and CFCs involves the breaking of weaker bonds in ozone by UV radiation. Oxygen ($O_2$) requires higher energy UV-C (shorter wavelength) to break its double bond compared to the resonance-stabilized $1.5$ bond order in ozone, which is broken by UV-B.