Chapter 4 — Chemical Bonding and Molecular Structure
Why atoms stick together at all — and why the result is sometimes a hard, high-melting crystal of ions and sometimes a gas of small, shaped molecules. Lewis structures, ionic bonds and bond parameters, the VSEPR shapes, hybridisation, molecular orbitals and hydrogen bonding.
Pour liquid nitrogen between the poles of a strong magnet and it splashes straight through. Pour liquid oxygen and something strange happens: the pale-blue liquid clings to the poles and hangs there, bridging the gap, until it boils away. Oxygen is attracted by a magnet, which means its molecules contain unpaired electrons. Yet the simple picture of the $\mathrm{O_{2}}$ molecule you may have drawn in Class 10 — two oxygen atoms sharing two pairs of electrons, every electron neatly paired — predicts no such thing. To explain one bucket of pale-blue liquid we need a better theory of the chemical bond.
This chapter builds that theory step by step. It starts with the idea of Kössel and Lewis that atoms bond to reach a stable octet, either by transferring electrons (ionic bonds) or by sharing them (covalent bonds), and with the numbers that describe a bond: its length, angle, enthalpy, order and polarity. Then it asks what shape a molecule has — why $\mathrm{CO_{2}}$ is straight, $\mathrm{H_{2}O}$ bent and $\mathrm{NH_{3}}$ a pyramid — and answers with electron-pair repulsion (VSEPR). Valence bond theory explains bonds as overlapping atomic orbitals and introduces hybridisation, the language of all of organic chemistry. Finally, molecular orbital theory explains the magnetism of oxygen and predicts which diatomic molecules can exist at all, and hydrogen bonding explains why water, with a molar mass of only $18\ \mathrm{g\,mol^{-1}}$, is a liquid at room temperature. Shapes, hybridisation and bond order appear in almost every Board, JEE and NEET paper.
- explain the Kössel–Lewis approach and the octet rule, draw Lewis structures of molecules and ions, assign formal charges, and recognise the three kinds of exceptions to the octet rule;
- describe the formation of an ionic bond, the role of lattice enthalpy, and the conditions that favour ionic bonding;
- define bond length, bond angle, bond enthalpy and bond order, relate them to each other, and use resonance structures to find average bond orders;
- explain bond polarity and dipole moment, and use Fajans' rules to compare the covalent character of ionic compounds;
- predict the shapes and bond angles of molecules and ions with 2 to 6 electron pairs using VSEPR theory, and decide whether a molecule is polar;
- explain bond formation by orbital overlap, distinguish σ and π bonds, and assign sp, sp², sp³, sp³d and sp³d² hybridisation;
- draw molecular orbital energy-level diagrams for $\mathrm{H_{2}}$ to $\mathrm{Ne_{2}}$ and their ions, calculate bond orders and predict magnetic behaviour;
- explain inter- and intramolecular hydrogen bonding and its effect on boiling points and other properties.
Sections in this chapter
Lewis Structures, Ionic Bonds and Bond Parameters
The octet rule and Lewis structures, formal charge and the limits of the octet rule; ionic bonds and lattice enthalpy; bond length, angle, enthalpy and order; resonance; bond polarity, dipole moment and Fajans' rules.
VSEPR Theory, Molecular Shapes and Polarity
Electron pairs repel and spread out: linear to octahedral arrangements, the shapes produced by lone pairs (bent, pyramidal, see-saw, T-shaped, square planar…), and dipole moment as a vector sum.
Valence Bond Theory and Hybridisation
Bonding as orbital overlap; the $\mathrm{H_{2}}$ energy curve; σ and π bonds; sp, sp², sp³, sp³d and sp³d² hybridisation in $\mathrm{BeCl_{2}}$, $\mathrm{BCl_{3}}$, $\mathrm{CH_{4}}$, $\mathrm{C_{2}H_{4}}$, $\mathrm{C_{2}H_{2}}$, $\mathrm{PCl_{5}}$ and $\mathrm{SF_{6}}$.
Molecular Orbital Theory and Hydrogen Bonding
LCAO; bonding and antibonding, σ and π molecular orbitals; energy-level orders with and without s–p mixing; bond order and magnetism of $\mathrm{H_{2}}$ to $\mathrm{Ne_{2}}$ and their ions; inter- and intramolecular hydrogen bonds.