Chemical Bonding and Molecular Structure — NEET Notes

Kossel-Lewis approach — the octet idea

In 1916, working independently, the German chemist Walther Kossel and the American chemist Gilbert N. Lewis arrived at the same insight: atoms combine in order to attain the electronic configuration of the nearest noble gas. Lewis pictured the atom as a positively charged kernel (nucleus plus inner electrons) surrounded by an outer shell that could hold up to eight valence electrons. He drew these as dots at the corners of an imaginary cube — the Lewis dot symbol. Kossel pointed out that the highly electropositive alkali metals and the highly electronegative halogens sit on opposite sides of the noble gases in the periodic table; sodium loses an electron to become Na⁺ with the Ne configuration, chlorine gains one to become Cl⁻ with the Ar configuration, and the resulting ions stabilise each other by electrostatic attraction. From this came the octet rule: atoms gain, lose, or share electrons to attain eight in their outer shell. Hydrogen and helium are the exception — they aim for a duplet (two electrons, the He configuration).

Langmuir refined the picture in 1919 by introducing the term covalent bond — a shared pair of electrons. A single bond is one shared pair; a double bond two; a triple bond three. Lewis dot structures became the universal shorthand for showing bonding pairs and lone pairs.

"Octet rule with exceptions" — useful as a sketch, false as a law.

The compromise that runs through this whole chapter

The octet rule is a sketch, not a law. It has three classes of well-known exceptions:

Formal charge is a bookkeeping device for Lewis structures. For an atom in a molecule, formal charge = (valence electrons of free atom) − (non-bonding electrons) − ½ (bonding electrons). It does not represent real charge separation; it helps you pick the most plausible Lewis structure among alternatives. The structure with the smallest formal charges on each atom is usually the lowest in energy.

Ionic bond — the lattice that holds it together

An ionic bond forms when one atom transfers electrons to another. The classic example is sodium chloride: Na (1s²2s²2p⁶3s¹) loses one electron to become Na⁺; Cl (1s²2s²2p⁶3s²3p⁵) gains it to become Cl⁻; both attain noble-gas configurations and the resulting ions are bound by Coulombic attraction. Crucially, the bond does not exist in isolation — in the solid, Na⁺ and Cl⁻ pack into a three-dimensional lattice in which every Na⁺ has six Cl⁻ neighbours and every Cl⁻ has six Na⁺ neighbours. This is the rock-salt structure.

The arithmetic of ionic-bond formation looks alarming at first. The ionisation enthalpy of Na is +495.8 kJ/mol and the electron-gain enthalpy of Cl is −348.7 kJ/mol; the sum is positive at +147.1 kJ/mol. So why does NaCl form spontaneously? The answer is that we have not yet added the term that dominates the whole process — the lattice enthalpy, which is −788 kJ/mol for NaCl. Once the gaseous ions snap into the crystal lattice, far more energy is released than was needed to make them. The ionic compound is stable not because of the octet but because of the lattice.

Lattice enthalpy — the master variable

Lattice enthalpy is defined as the energy required to completely separate one mole of a solid ionic compound into its gaseous constituent ions. It depends on two simple factors:

Born-Haber cycle — measuring what we cannot weigh

Lattice enthalpy cannot be measured directly — you cannot scoop a mole of gaseous Na⁺ and a mole of gaseous Cl⁻ onto a balance. The Born-Haber cycle gets around this by applying Hess's law: the enthalpy change for the overall reaction Na(s) + ½ Cl₂(g) → NaCl(s) equals the sum of the enthalpy changes of any pathway connecting reactants and products. Break the reaction into measurable steps, leave lattice enthalpy as the only unknown, and solve.

Bond parameters — length, angle, enthalpy, order

Four numbers characterise any covalent bond. Bond length is the equilibrium internuclear distance, measured in picometres — about 74 pm for H–H, 154 pm for C–C, 134 pm for C=C, 120 pm for C≡C. Each atom contributes its covalent radius to this length. Bond angle is the angle between two bonds at a central atom — 104.5° in H₂O, 107° in NH₃, 109.5° in CH₄, 120° in BF₃, 180° in BeCl₂. Bond enthalpy is the energy required to break one mole of a particular bond in the gas phase: 436 kJ/mol for H–H, 498 kJ/mol for O=O, 946 kJ/mol for N≡N — that last figure being one of the strongest known and the reason atmospheric nitrogen is so inert. Bond order is simply the number of shared electron pairs: 1 for a single, 2 for a double, 3 for a triple.

Three patterns recur across NEET problems and are worth memorising. First, for the same pair of atoms, bond order ↑ → bond length ↓ → bond enthalpy ↑. Hence C–C (154 pm, 348 kJ/mol) < C=C (134 pm, 614 kJ/mol) < C≡C (120 pm, 839 kJ/mol). Second, isoelectronic species have identical bond orders: N₂, CO, CN⁻, NO⁺ all carry triple bonds with bond order 3 — NEET 2017 Q.6 tested exactly this. Third, resonance averages bond lengths: all three N–O bonds in NO₃⁻ are equal (between single and double), and all three C–O bonds in CO₃²⁻ are equal too.

Bond polarity and dipole moment

No bond is purely covalent or purely ionic. When the two atoms are different, the more electronegative one pulls the shared pair toward itself, creating a polar covalent bond. The result is a dipole moment µ = Q × r, measured in Debye units (1 D = 3.34 × 10⁻³⁰ C m). For diatomic molecules the bond dipole is the molecular dipole. For polyatomic molecules, the molecular dipole is the vector sum of the bond dipoles, and symmetry can cancel everything out: BeF₂, BF₃, CCl₄, CO₂, and 1,4-dichlorobenzene all have µ = 0 despite having polar bonds. NEET 2020 Q.150 and NEET 2021 Q.91 both turn on this distinction.

A subtle case worth noticing: NH₃ has µ = 1.47 D but NF₃ has µ = 0.23 D, even though fluorine is more electronegative than hydrogen. The reason is the lone pair on nitrogen. In NH₃ the N–H bond dipoles point toward N and add to the lone-pair dipole. In NF₃ the N–F bond dipoles point away from N and cancel the lone-pair dipole. Geometry decides the answer, not electronegativity alone.

The covalent bond — sharing as architecture

Covalent bonds are formed by sharing one or more pairs of electrons between atoms. Each atom contributes at least one electron to the shared pair, and each ends up with a stable noble-gas configuration. The hallmark of a covalent compound is the molecule — a discrete unit, often with definite shape, soluble in non-polar solvents, with low melting points and poor electrical conductivity. Contrast this with ionic compounds, which form three-dimensional lattices, dissolve in water, conduct in the molten or aqueous state, and melt only at very high temperatures.

The Kossel-Lewis picture says nothing about the actual mechanism of covalent bonding — why does H₂ form at all? What does "sharing" mean physically? Three theories of increasing sophistication answer these questions: VSEPR (which predicts geometry), valence bond theory (which explains bonding via orbital overlap), and molecular orbital theory (which treats electrons as belonging to the whole molecule). We will take them in turn.

VSEPR theory — geometry from electron-pair repulsion

The Valence Shell Electron Pair Repulsion theory was proposed by Sidgwick and Powell in 1940 and refined by Nyholm and Gillespie in 1957. It is the most-tested idea in this chapter. The core postulate is brutally simple: electron pairs in the valence shell of the central atom repel each other and arrange themselves to maximise the distance between them. Whether the pair is bonded or unbonded does not matter for placement; both kinds count.

A second postulate sharpens the prediction: lone pairs occupy more space than bonded pairs because a lone pair is held by only one nucleus while a bonded pair is shared between two. The order of repulsion strength is therefore:

Lone pair – Lone pair > Lone pair – Bond pair > Bond pair – Bond pair

The Nyholm-Gillespie repulsion order — NEET 2016 Q.2 tested this directly

This is why H₂O (two lone pairs on O) has an H–O–H angle of 104.5°, NH₃ (one lone pair on N) has 107°, and CH₄ (no lone pairs on C) keeps the ideal tetrahedral 109.5°. The lone pairs squeeze the bond pairs together. NEET 2016 Q.19 — "the H–O–H bond angle in H₂O is larger than the H–C–H bond angle in CH₄" — turns on this single chain.

To apply VSEPR to any molecule, follow a four-step process:

In trigonal bipyramidal and octahedral cases, lone pairs preferentially occupy the positions that minimise lp–bp repulsions. In a trigonal bipyramid, lone pairs go to equatorial positions (only 2 lp–bp repulsions at 90° instead of 3). This explains the see-saw shape of SF₄ (sp³d, 1 lp, equatorial), the T-shape of ClF₃ (sp³d, 2 lp, both equatorial), and the linear shape of XeF₂ (sp³d, 3 lp, all equatorial). NEET 2022 Q.62 asked which of IF₅, SF₄, XeF₂, ClF₃ has maximum lp–lp repulsion — the answer is XeF₂ with three lone pairs in the equatorial plane.

Valence bond theory — the orbital overlap picture

VSEPR predicts where the atoms sit but says nothing about why a bond exists or how strong it is. Valence Bond Theory, introduced by Heitler and London in 1927 and developed by Linus Pauling, fills that gap. Its starting point is the formation of H₂. When two H atoms approach, two new attractive forces appear (each nucleus attracts the other's electron) and two new repulsive forces (electron–electron and nucleus–nucleus). The attractive forces dominate until the atoms reach an equilibrium separation of 74 pm, where the potential energy hits a minimum. At that minimum, the two 1s orbitals have partially merged — overlapped — and the electrons, with opposite spins, are shared between the nuclei. The energy released, 436 kJ/mol, is the bond enthalpy.

From this single picture, three properties of covalent bonds emerge naturally. Directionality: p-orbitals point along axes, so p–p overlap is directional. Strength scales with overlap: bigger overlap, stronger bond. Bond order matters: when two atoms share more than one pair, additional sideways overlaps form on top of the axial overlap.

Sigma and pi bonds

VBT distinguishes two kinds of covalent bond by the geometry of the overlap. A sigma (σ) bond is formed by end-on, head-on overlap along the internuclear axis — s–s, s–p, or p–p axial overlap. A pi (π) bond is formed by sideways overlap of two parallel p-orbitals perpendicular to the internuclear axis, producing two saucer-shaped clouds above and below the bond. Because head-on overlap concentrates electron density between the nuclei, sigma bonds are stronger than pi bonds. The structural recipe for multiple bonds is now clear: a single bond = 1 σ, a double bond = 1 σ + 1 π, a triple bond = 1 σ + 2 π. NEET 2023 Q.53 asked you to count σ-bonds, π-bonds, and lone pairs in pyridine — the answer (11 σ, 3 π, 1 lp) flows from this recipe.

Hybridisation — Pauling's geometry-from-orbitals fix

Plain VBT runs into trouble with polyatomic molecules. Carbon in its ground state is 2s² 2p², with only two unpaired electrons — yet it forms four equivalent C–H bonds in methane, all pointing to the corners of a tetrahedron at 109.5°. Plain p–p overlap would predict 90° between any two C–H bonds and would leave one H attached anywhere. Pauling's resolution was hybridisation: atomic orbitals of slightly different energy on the same atom can mix to produce a new set of equivalent orbitals — hybrid orbitals — that point in specific directions.

The number of hybrid orbitals always equals the number of atomic orbitals mixed. Each hybrid orbital has a fixed s-character and p-character. Hybridisation lets you read off geometry directly: count the steric number of the central atom (bond pairs + lone pairs), assign the matching hybridisation, and the geometry follows.

A small caution: the choice of hybridisation is read off the steric number, not the molecular shape. SF₄ and CH₄ are both sp³d... no, SF₄ is sp³d (steric number 5, one lone pair, see-saw shape). CH₄ is sp³ (steric number 4, no lone pairs, tetrahedral). The steric number tells you the hybrid set; the lone-pair count tells you the molecular shape.

Molecular orbital theory — electrons belong to the whole molecule

Hybridisation rescues VBT for polyatomic molecules, but even VBT runs into problems with simple diatomics. The most famous failure: VBT predicts O₂ to be diamagnetic (all electrons paired in two double-bond electron pairs), yet liquid O₂ is observably paramagnetic — it sticks to a magnet. The theory that explains this, and much more, is Molecular Orbital Theory, developed by Hund and Mulliken in 1932.

MOT abandons the idea that bonding electrons belong to particular atoms. Instead, the atomic orbitals of all atoms in a molecule combine — by linear combination of atomic orbitals (LCAO) — to produce a new set of molecular orbitals that extend over the whole molecule. Two atomic orbitals always combine to make two molecular orbitals: a bonding MO (lower in energy, electron density between the nuclei) by addition of wavefunctions, and an antibonding MO (higher in energy, with a node between the nuclei) by subtraction. We label them σ, π, σ*, π*; the asterisk marks the antibonding orbital.

For atomic orbitals to mix, three conditions must be met: similar energy (1s with 1s, not 1s with 2s), same symmetry about the bond axis (2p_z with 2p_z if the bond is along z), and appreciable overlap. Once filled, MOs follow the same rules as atomic orbitals — Aufbau, Pauli, Hund.

Bond order, stability, magnetic character

From the MO diagram you read off three quantities:

• Bond order = ½ (N_b − N_a), where N_b is the number of electrons in bonding MOs and N_a in antibonding MOs. Bond order > 0 → stable molecule; bond order = 0 → does not exist.
• Bond length and enthalpy: higher bond order → shorter, stronger bond.
• Magnetic character: any unpaired electron makes the molecule paramagnetic; all paired makes it diamagnetic.

MO diagrams — H₂, He₂, N₂, O₂, F₂

Now to the second-row diatomics that NEET asks every year. There are two energy orderings of MOs to remember, and the difference is critical.

Now the per-molecule rundown — exactly what NEET asks:

A useful family of related calculations to memorise: O₂ → O₂⁺ → O₂⁻ → O₂²⁻. Each electron added goes into the π* antibonding MOs, so bond order drops: O₂⁺ has BO 2.5, O₂ has BO 2, O₂⁻ (superoxide) has BO 1.5, O₂²⁻ (peroxide) has BO 1. Bond length therefore increases in the same order, and stability decreases. NEET 2018 Q.61 — "which has highest bond order among NO, CN⁻, CN⁺, CN?" — turns on counting electrons and reading off bond order from the MO diagram.

Hydrogen bonding — the bond that runs biology

The last bond in this chapter is by far the most consequential one in the world outside the chemistry lab. A hydrogen bond is the attractive force between a hydrogen atom covalently bonded to a small, highly electronegative atom (F, O, or N) and another electronegative atom — in the same molecule (intramolecular) or in a different one (intermolecular). The shared electron pair in the H–X bond is pulled toward X, leaving H with a partial positive charge δ+ and X with a partial negative charge δ−. The δ+ hydrogen of one molecule is then attracted to the δ− atom of the next.

Hydrogen bonds are weaker than covalent bonds (typically 10–40 kJ/mol vs 200–800 kJ/mol) but much stronger than ordinary van der Waals dispersion forces (2–5 kJ/mol). They are responsible for:

• The anomalously high boiling point of water (100 °C; H₂S boils at −60 °C despite being heavier).
• The fact that ice floats on water — H-bonds force water molecules into an open lattice that is less dense than liquid water.
• The dimerisation of carboxylic acids and the high boiling points of alcohols.
• The alpha-helix and beta-sheet structures of proteins, and the double helix of DNA — every base pair held by 2 or 3 H-bonds.
• The mutarotation of sugars and the solubility of NH₃, ethanol, glucose in water.

Intramolecular H-bonds (e.g. in o-nitrophenol, salicylaldehyde) lower boiling points by reducing intermolecular association, while intermolecular H-bonds (e.g. in p-nitrophenol) raise boiling points. NEET 2023 Q.74 grouped hydrogen bonding with dipole–dipole, dipole–induced dipole, and dispersion forces as "intermolecular forces" — explicitly excluding covalent bonds, which act within a molecule.