Chemical Bonding & Molecular Structure

Atoms bond to achieve stable, noble-gas-like electron arrangements — the essence of the octet rule (eight electrons in the valence shell). The Kossel-Lewis approach explains two main ways this happens: by transferring electrons (ionic bonding) or by sharing them (covalent bonding). Which route an element takes depends largely on its ionisation enthalpy and electron gain enthalpy, linking this chapter directly to periodicity.

An ionic (electrovalent) bond forms when a metal transfers electrons to a non-metal, producing oppositely charged ions held together by electrostatic attraction, as in $\text{Na}^+\text{Cl}^-$. The strength of an ionic crystal is measured by its lattice enthalpy, which increases with higher ionic charges and smaller ionic sizes. Fajans' rules add nuance: a small, highly charged cation with a large anion polarises the anion's cloud, giving ionic compounds partial covalent character — a frequent NEET reasoning point.

A covalent bond forms when two atoms share one or more electron pairs, as in $\text{H}_2$, $\text{O}_2$ and $\text{CH}_4$. These are represented by Lewis (electron-dot) structures, drawn by counting total valence electrons, forming bonds, and completing octets (with duplets for hydrogen). Multiple bonds (double, triple) form when atoms share two or three pairs, as in $\text{O}=\text{O}$ and $\text{N}\equiv\text{N}$.

To choose the best Lewis structure, chemists use formal charge, calculated as (valence electrons) − (non-bonding electrons) − ½(bonding electrons). The most stable structure has formal charges closest to zero. The octet rule, while powerful, has clear exceptions NEET likes to test: incomplete octets ($\text{BeCl}_2$, $\text{BF}_3$), odd-electron species ($\text{NO}$), and expanded octets ($\text{PCl}_5$, $\text{SF}_6$) where the central atom holds more than eight electrons.

Lewis structures show which atoms are bonded but not the three-dimensional shape of a molecule. The VSEPR (Valence Shell Electron Pair Repulsion) theory fills this gap with one simple idea: electron pairs around a central atom repel one another and arrange themselves as far apart as possible. This single principle predicts the geometry of most NEET molecules quickly.

The key step is to count the electron pairs (bond pairs + lone pairs) on the central atom. Two pairs give a linear shape ($180^{\circ}$); three give trigonal planar ($120^{\circ}$); four give tetrahedral ($109.5^{\circ}$); five give trigonal bipyramidal; and six give octahedral ($90^{\circ}$). These idealised angles apply when all the pairs are bonding.

The crucial refinement is the effect of lone pairs, which occupy more space than bonding pairs and so repel more strongly. The repulsion order is lone pair-lone pair $>$ lone pair-bond pair $>$ bond pair-bond pair. This compresses bond angles when lone pairs are present, giving the classic NEET series: methane $\text{CH}_4$ ($109.5^{\circ}$, no lone pairs) $>$ ammonia $\text{NH}_3$ ($107^{\circ}$, one lone pair) $>$ water $\text{H}_2\text{O}$ ($104.5^{\circ}$, two lone pairs).

Shapes are often described by their actual atom positions, ignoring lone pairs: $\text{NH}_3$ is pyramidal and $\text{H}_2\text{O}$ is bent, even though both are based on a tetrahedral arrangement of electron pairs. Practising the count-pairs-then-adjust-for-lone-pairs routine on common molecules ($\text{CO}_2$, $\text{BF}_3$, $\text{CH}_4$, $\text{NH}_3$, $\text{H}_2\text{O}$, $\text{PCl}_5$, $\text{SF}_6$) covers nearly every shape question NEET asks.

Valence bond theory (VBT) pictures a covalent bond as the overlap of half-filled atomic orbitals, with greater overlap giving a stronger bond. A head-on overlap forms a strong sigma ($\sigma$) bond, while the sideways overlap of $p$ orbitals forms a weaker pi ($\pi$) bond. A single bond is one $\sigma$; a double bond is one $\sigma$ + one $\pi$; a triple bond is one $\sigma$ + two $\pi$ — a counting rule NEET uses constantly.

Pure atomic orbitals often cannot explain observed shapes (for instance, carbon's four identical bonds in $\text{CH}_4$). The fix is hybridisation — the mixing of atomic orbitals of similar energy to form new, equivalent hybrid orbitals that point toward the bonded atoms. The type of hybridisation fixes the geometry, so identifying it is the master key to molecular shape.

The common hybridisations and their geometries are: $sp$ (2 hybrids, linear, $180^{\circ}$, as in $\text{BeCl}_2$ and $\text{C}_2\text{H}_2$); $sp^2$ (3 hybrids, trigonal planar, $120^{\circ}$, as in $\text{BF}_3$ and $\text{C}_2\text{H}_4$); $sp^3$ (4 hybrids, tetrahedral, $109.5^{\circ}$, as in $\text{CH}_4$, $\text{NH}_3$, $\text{H}_2\text{O}$); $sp^3d$ (5 hybrids, trigonal bipyramidal, as in $\text{PCl}_5$); and $sp^3d^2$ (6 hybrids, octahedral, as in $\text{SF}_6$).

A fast NEET shortcut finds hybridisation by counting the sigma bonds plus lone pairs on the central atom: a total of 2 means $sp$, 3 means $sp^2$, 4 means $sp^3$, 5 means $sp^3d$, and 6 means $sp^3d^2$. Lone pairs occupy hybrid orbitals too, which is why $\text{NH}_3$ and $\text{H}_2\text{O}$ are both $sp^3$ despite being pyramidal and bent rather than tetrahedral. This count-and-classify method handles almost every hybridisation question.

VBT explains shapes well but fails for some molecules, so molecular orbital theory (MOT) offers a deeper picture. Atomic orbitals combine to form molecular orbitals spread over the whole molecule — a lower-energy bonding orbital and a higher-energy antibonding orbital. Electrons fill these by the same rules as atoms (Aufbau, Pauli, Hund). The key quantity is the bond order $= \tfrac{1}{2}(N_b - N_a)$, where $N_b$ and $N_a$ are the electrons in bonding and antibonding orbitals.

Bond order predicts stability and existence: a bond order of zero means the molecule does not exist, which is why $\text{He}_2$ is not formed. A higher bond order means a stronger, shorter bond. MOT's most celebrated success is explaining the paramagnetism of $\text{O}_2$ — it has two unpaired electrons in its $\pi^*$ antibonding orbitals, a fact the Lewis and valence-bond pictures cannot show. This single example is a NEET favourite.

Whether a molecule is polar depends on its shape and bond dipoles. A bond between atoms of different electronegativity is polar, but the molecule is non-polar if the bond dipoles cancel by symmetry. So $\text{CO}_2$ (linear) is non-polar despite polar bonds, while $\text{H}_2\text{O}$ (bent) is strongly polar — the asymmetry leaves a net dipole moment. Predicting polarity from geometry is a routine NEET skill.

The most biologically important interaction is the hydrogen bond — a strong dipole attraction when hydrogen is bonded to the small, highly electronegative atoms N, O or F. Hydrogen bonding gives water its unusually high boiling point and surface tension, and, crucially for medical aspirants, holds the two strands of the DNA double helix together through base pairing and stabilises the folded shapes of proteins. Intermolecular hydrogen bonds (as in water) raise boiling points, while intramolecular ones (as in ortho-nitrophenol) can lower them — a subtle distinction NEET sometimes probes.