G.N. Lewis proposed that atoms combine to achieve a stable electronic configuration, usually resembling that of the nearest noble gas — with 8 electrons in the outermost shell (the octet rule). Hydrogen and helium are exceptions, being stable with just 2 electrons (the duplet rule).
A Lewis symbol represents an element's chemical symbol surrounded by dots equal to the number of valence electrons. Bonds form when atoms share or transfer these valence electrons to complete their octets.
Some stable molecules don't follow it: $\text{BF}_3$ has only 6 electrons around boron (incomplete octet); $\text{SF}_6$ has 12 electrons around sulphur (expanded octet); and odd-electron species like $\text{NO}$ can never satisfy it exactly.
An ionic bond forms when one atom completely transfers one or more electrons to another, creating oppositely charged ions that are held together by strong electrostatic attraction. Typically forms between a metal (low ionisation enthalpy) and a non-metal (high electron gain enthalpy).
Higher ionic charges and smaller ionic sizes give stronger, more stable ionic lattices.
A covalent bond forms when two atoms share one or more pairs of electrons, typically between two non-metals of similar electronegativity. A single shared pair gives a single bond, two pairs a double bond, and three pairs a triple bond.
Used to identify the most stable Lewis structure among several possibilities — the structure with formal charges closest to zero, and negative charges on the more electronegative atoms, is usually preferred.
| Parameter | Meaning | Trend |
|---|---|---|
| Bond length | Equilibrium distance between the nuclei of two bonded atoms | Decreases as bond order increases (triple < double < single) |
| Bond angle | Angle between two adjacent bonds at a common atom | Depends on hybridisation & lone pair repulsion |
| Bond enthalpy | Energy needed to break one mole of a particular bond in the gaseous state | Increases as bond order increases |
| Bond order | Number of bonds (single/double/triple) between two atoms | Higher bond order = shorter, stronger bond |
Resonance occurs when a single Lewis structure cannot accurately represent a molecule's actual structure, so it's described as a hybrid of two or more valid structures (resonance structures/canonical forms) that differ only in the position of electrons, not atoms. The real molecule is a blend — a resonance hybrid — more stable than any single contributing structure.
Classic example: ozone (O₃), where the two O–O bonds are experimentally found to be identical in length (intermediate between a single and double bond) — explained by two resonance structures with the double bond on alternating sides.
Resonance structures are not different molecules interconverting rapidly — they don't actually exist separately; only the resonance hybrid is real.
When two atoms of different electronegativity form a covalent bond, the shared electron pair is pulled closer to the more electronegative atom, creating partial charges ($\delta^+$ and $\delta^-$) — this is a polar covalent bond.
where $q$ = magnitude of partial charge, $d$ = distance between the charges. Measured in Debye (D); $1\ \text{D} = 3.33 \times 10^{-30}\ \text{C m}$. Dipole moment is a vector quantity — pointing from the positive to the negative end.
The overall dipole moment is the vector sum of all individual bond dipoles. Even if individual bonds are polar, a molecule can be non-polar overall if the bond dipoles cancel due to symmetry — e.g., $\text{CO}_2$ (linear, dipoles cancel) is non-polar despite having polar C=O bonds, while $\text{H}_2\text{O}$ (bent) is polar.
Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shape based on the idea that electron pairs (bonding and lone) around a central atom arrange themselves as far apart as possible to minimise repulsion. Repulsion strength follows: lone pair–lone pair > lone pair–bond pair > bond pair–bond pair.
| Bond pairs + Lone pairs | Shape | Bond angle | Example |
|---|---|---|---|
| 2 + 0 | Linear | 180° | BeCl₂ |
| 3 + 0 | Trigonal planar | 120° | BF₃ |
| 2 + 1 | Bent | <120° | SO₂ |
| 4 + 0 | Tetrahedral | 109.5° | CH₄ |
| 3 + 1 | Trigonal pyramidal | <109.5° | NH₃ |
| 2 + 2 | Bent (V-shaped) | <109.5° | H₂O |
| 5 + 0 | Trigonal bipyramidal | 90° & 120° | PCl₅ |
| 6 + 0 | Octahedral | 90° | SF₆ |
Valence Bond Theory (VBT) explains covalent bond formation as the overlap of atomic orbitals from two atoms, each contributing one unpaired electron. Greater orbital overlap gives a stronger bond.
A sigma bond forms by head-on (axial) overlap of orbitals along the internuclear axis — allows free rotation and is generally stronger. A pi bond forms by sideways (lateral) overlap of parallel p-orbitals above and below the internuclear axis — restricts rotation and is generally weaker than a sigma bond. A single bond is always one σ bond; a double bond is one σ + one π; a triple bond is one σ + two π.
Hybridisation is the mixing of atomic orbitals of similar energy on the same atom to form new, equivalent hybrid orbitals, better suited for bond formation and matching the observed molecular geometry.
| Type | Orbitals Mixed | Geometry | Example |
|---|---|---|---|
| sp | 1 s + 1 p | Linear (180°) | BeCl₂, C₂H₂ |
| sp² | 1 s + 2 p | Trigonal planar (120°) | BF₃, C₂H₄ |
| sp³ | 1 s + 3 p | Tetrahedral (109.5°) | CH₄, NH₃, H₂O |
| sp³d | 1 s + 3 p + 1 d | Trigonal bipyramidal | PCl₅ |
| sp³d² | 1 s + 3 p + 2 d | Octahedral | SF₆ |
Determine the hybridisation of the central atom in NH₃.
Nitrogen has 3 bond pairs (to 3 H atoms) + 1 lone pair = 4 electron domains → sp³ hybridisation, with a tetrahedral electron arrangement (though the molecular shape itself is trigonal pyramidal because one position is a lone pair).
Unlike VBT, Molecular Orbital Theory (MOT) treats electrons as belonging to the molecule as a whole, not to individual atoms. Atomic orbitals combine (via Linear Combination of Atomic Orbitals, LCAO) to form new molecular orbitals spread over the entire molecule.
Combining two atomic orbitals gives two molecular orbitals: a lower-energy bonding molecular orbital (BMO), formed by constructive overlap (in-phase addition), and a higher-energy antibonding molecular orbital (ABMO), formed by destructive overlap (out-of-phase subtraction, with a node between nuclei).
where $N_b$ = electrons in bonding MOs, $N_a$ = electrons in antibonding MOs. Higher bond order means a shorter, stronger, more stable bond; a bond order of zero means the molecule cannot exist.
If all electrons in the molecular orbitals are paired, the species is diamagnetic (weakly repelled by a magnetic field). If one or more molecular orbitals have unpaired electrons, the species is paramagnetic (attracted by a magnetic field) — this is how MOT correctly predicts that O₂ is paramagnetic, something VBT cannot explain.
Calculate the bond order of the nitrogen molecule, N₂ (14 electrons total).
Molecular orbital filling for N₂: $\sigma1s^2\, \sigma^*1s^2\, \sigma2s^2\, \sigma^*2s^2\, \pi2p_x^2\, \pi2p_y^2\, \sigma2p_z^2$ → bonding electrons = 10, antibonding electrons = 4.
Bond order $= \tfrac{1}{2}(10-4) = 3$ — consistent with the known N≡N triple bond.
A hydrogen bond is a special, relatively weak attraction between a hydrogen atom (already covalently bonded to a highly electronegative atom — F, O, or N) and a lone pair on another highly electronegative atom, either on a nearby molecule or within the same molecule. Because F, O, and N are small and highly electronegative, they pull electron density strongly away from hydrogen, leaving it with a significant partial positive charge that then attracts a nearby lone pair.