Key Points

A chemical bond is an attractive force that holds atoms or ions together in a chemical species. The octet rule states that atoms tend to combine in such a way that they each have eight electrons in their valence shells, achieving a stable noble gas configuration.

An ionic bond is formed by the complete transfer of one or more electrons from a metallic atom to a non-metallic atom, resulting in the electrostatic attraction between the resulting positive and negative ions, for example, in NaCl.

A covalent bond is formed by the mutual sharing of one or more electron pairs between atoms to complete their octets. Lewis structures use dots to represent valence electrons and lines to represent shared pairs (bonds).

Bond length is the equilibrium distance between the nuclei of two bonded atoms. Bond angle is the angle between orbitals containing bonding electron pairs. Bond enthalpy is the energy required to break one mole of a specific type of bond in a gaseous state.

In Lewis theory, bond order is the number of bonds between two atoms (e.g., 1 for single, 2 for double, 3 for triple). As bond order increases, bond enthalpy increases and bond length decreases.

When a single Lewis structure cannot accurately represent a molecule, its actual structure is described as a resonance hybrid of multiple contributing structures, called canonical forms. For example, the ozone molecule, $O_3$, has two resonance structures.

A covalent bond between atoms of different electronegativity is polar, creating a dipole moment ($\mu$). The net dipole moment of a polyatomic molecule depends on both the individual bond dipoles and the molecular geometry.

The Valence Shell Electron Pair Repulsion (VSEPR) theory states that the shape of a molecule is determined by the repulsions between electron pairs in the valence shell of the central atom, which arrange themselves to be as far apart as possible.

The magnitude of repulsion between electron pairs decreases in the order: Lone pair-Lone pair (lp-lp) > Lone pair-Bond pair (lp-bp) > Bond pair-Bond pair (bp-bp). This explains why molecules like $H_2O$ have a bent shape.

Valence Bond (VB) theory describes the formation of a covalent bond as the overlap of atomic orbitals. The strength of the bond depends on the extent of this overlap.

A sigma ($\sigma$) bond is formed by the direct, head-on overlap of orbitals along the internuclear axis. A pi ($\pi$) bond is formed by the sideways overlap of p-orbitals above and below the internuclear axis. A double bond consists of one $\sigma$ and one $\pi$ bond.

Hybridisation is the process of mixing atomic orbitals of slightly different energies to form a new set of equivalent orbitals, called hybrid orbitals. The type of hybridisation determines the geometry of the molecule.

Key types include sp (linear, $180^\circ$, e.g., $BeCl_2$), $sp^2$ (trigonal planar, $120^\circ$, e.g., $BCl_3$), and $sp^3$ (tetrahedral, $109.5^\circ$, e.g., $CH_4$).

MO theory describes electrons in a molecule residing in molecular orbitals that extend over the entire molecule. Combining atomic orbitals creates an equal number of bonding (lower energy) and antibonding (higher energy) molecular orbitals.

Bond order is defined as one-half the difference between the number of electrons in bonding orbitals ($N_b$) and antibonding orbitals ($N_a$). The formula is B.O. = $\frac{1}{2}(N_b - N_a)$. A positive bond order indicates a stable molecule.

MO theory explains the magnetic behavior of molecules. If all molecular orbitals are doubly occupied, the substance is diamagnetic. If there are one or more singly occupied orbitals, it is paramagnetic, like the $O_2$ molecule.

A hydrogen bond is an attractive force between a hydrogen atom covalently bonded to a highly electronegative atom (F, O, or N) and another electronegative atom. It is weaker than a covalent bond but significantly affects properties like boiling point.