Chapter 4: Chemical Bonding and Molecular Structure | Summary

Lewis dot structures provide a picture of bonding in molecules by showing shared pairs and lone pairs. Steps to draw Lewis structures include: counting total valence electrons, placing the least electronegative atom at the centre, distributing single bonds first, and then assigning lone pairs or multiple bonds to satisfy the octet. For ions, electrons are added (for anions) or removed (for cations) before drawing the structure.

Formal charge is the difference between valence electrons in a free atom and the electrons assigned to it in a Lewis structure. It is calculated as: Formal Charge = (valence electrons) - (lone pair electrons) - (1/2)(bonding electrons). Formal charges help in selecting the most stable (lowest energy) Lewis structure from multiple possibilities. The structure with the smallest formal charges on atoms is preferred.

The octet rule is not universally applicable. There are three main exceptions: (1) Incomplete octet - some molecules like LiCl, BeH2, BCl3, and AlCl3 have fewer than eight electrons around the central atom. (2) Odd-electron molecules - species like NO and NO2 have an odd number of electrons and cannot satisfy the octet rule for all atoms. (3) Expanded octet - elements in the third period and beyond (e.g., P, S) can accommodate more than eight electrons around them due to the availability of 3d orbitals, as seen in PF5, SF6, and H2SO4.

Additional drawbacks include: the octet rule is based on the chemical inertness of noble gases, yet some noble gases like xenon and krypton form compounds (XeF2, KrF2). The theory also fails to explain molecular shapes and says nothing about the relative stability or energy of molecules.

Ionic or electrovalent bonds form due to electrostatic attraction between oppositely charged ions. The formation depends on the ease of forming positive and negative ions (governed by ionization enthalpy and electron gain enthalpy) and the stability of the crystal lattice. Ionic bonds form more easily between elements with low ionization enthalpy and high electron gain enthalpy.

Lattice enthalpy is defined as the energy required to completely separate one mole of a solid ionic compound into its gaseous constituent ions. For NaCl, the lattice enthalpy is 788 kJ mol-1. Even if the sum of ionization enthalpy and electron gain enthalpy is positive, the crystal lattice is stabilized by the energy released in lattice formation. Hence, lattice enthalpy (not just achieving the octet) is the true qualitative measure of stability for ionic compounds.

Bond length is the equilibrium distance between the nuclei of two bonded atoms. Covalent radius is half the distance between two similar atoms joined by a covalent bond. Van der Waals radius represents the overall size of the atom in a non-bonded situation. Bond angle is the angle between bonding orbital pairs around the central atom, measured in degrees. It provides information about the distribution of orbitals and helps determine molecular shape.

Bond enthalpy is the energy required to break one mole of a specific bond in the gaseous phase (unit: kJ mol-1). Greater bond dissociation enthalpy indicates a stronger bond. In polyatomic molecules, the mean bond enthalpy is used. Bond order is the number of bonds between two atoms; higher bond order means greater bond enthalpy and shorter bond length. Resonance structures arise when a single Lewis structure cannot adequately represent the true bonding; the resonance hybrid (e.g., for O3, CO3 2-) has bond characteristics intermediate between the canonical forms.

When a covalent bond forms between two identical atoms, the electron pair is shared equally and the bond is nonpolar. When atoms differ in electronegativity, the shared pair shifts toward the more electronegative atom, creating a polar covalent bond. Dipole moment (mu) is the product of the magnitude of charge and the distance of separation: mu = Q x r. It is expressed in Debye units (D), where 1 D = 3.33564 x 10-30 C m.

In polyatomic molecules, the net dipole moment is the vector sum of all individual bond dipoles and also depends on molecular geometry. For example, CO2 and BeF2 have zero dipole moments due to symmetric cancellation of bond dipoles, while H2O has a net dipole moment of 1.85 D due to its bent shape. Fajans' rules describe the partial covalent character of ionic bonds: smaller cations and larger anions give greater covalent character; higher cation charge also increases covalent character.

The Valence Shell Electron Pair Repulsion (VSEPR) theory, proposed by Sidgwick and Powell (1940) and refined by Nyholm and Gillespie (1957), predicts molecular geometry based on the repulsion between electron pairs in the valence shell. The main postulates are: molecular shape depends on the number of electron pairs around the central atom; electron pairs repel each other and occupy positions minimizing repulsion; lone pairs occupy more space than bond pairs.

The order of repulsion is: lp-lp > lp-bp > bp-bp. This explains why NH3 has a bond angle of 107° (one lone pair) and H2O has 104.5° (two lone pairs), both lower than the tetrahedral angle of 109.5°. VSEPR accurately predicts geometries such as linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral, and distorted shapes like bent, pyramidal, see-saw, T-shape, square pyramidal, and square planar.

Valence Bond (VB) theory, introduced by Heitler and London (1927) and developed by Pauling, explains covalent bond formation in terms of overlap of atomic orbitals containing electrons of opposite spins. When two H atoms approach each other, attractive forces outweigh repulsive forces at the equilibrium bond length (74 pm) and the potential energy is minimum. Covalent bonds are classified as sigma (sigma, end-to-end axial overlap) and pi (pi, sidewise parallel overlap). Sigma bonds are stronger than pi bonds.

Hybridisation, introduced by Pauling, is the intermixing of orbitals of slightly different energies to form a new set of equivalent hybrid orbitals. Key types include sp (linear, 180°, e.g., BeCl2), sp2 (trigonal planar, 120°, e.g., BCl3, C2H4), sp3 (tetrahedral, 109.5°, e.g., CH4, NH3, H2O), sp3d (trigonal bipyramidal, e.g., PCl5), sp3d2 (octahedral, e.g., SF6), and dsp2 (square planar, e.g., [Ni(CN)4]2-). Hybridisation accounts for the directional properties and characteristic geometries of polyatomic molecules.

Molecular Orbital (MO) theory, developed by Hund and Mulliken (1932), treats bonding electrons as belonging to the entire molecule rather than individual atoms. Molecular orbitals are formed by the Linear Combination of Atomic Orbitals (LCAO). Addition of wave functions gives a bonding MO (lower energy, greater stability) and subtraction gives an antibonding MO (higher energy, nodal plane between nuclei). The number of MOs equals the number of combining AOs.

The electronic configuration of a molecule is written by filling MOs in order of increasing energy (aufbau principle), following Pauli's exclusion principle and Hund's rule. Bond order = 1/2(Nb - Na). A positive bond order indicates stability; zero or negative means the molecule is unstable. Magnetic behaviour depends on unpaired electrons: diamagnetic if all MOs are doubly occupied (e.g., H2, N2), paramagnetic if unpaired electrons are present (e.g., O2, B2). MO theory successfully explains the paramagnetic nature of O2.

Hydrogen bonding occurs when a hydrogen atom, covalently bonded to a highly electronegative atom (F, O, or N), experiences attraction toward another electronegative atom. The electronegative atom pulls the shared electron pair away from H, giving it a partial positive charge (delta+). This positively charged H then attracts a neighbouring electronegative atom, forming a hydrogen bond, which is weaker than a covalent bond but stronger than van der Waals forces.

There are two types of hydrogen bonds: intermolecular (between two different molecules, e.g., HF, H2O, alcohols) and intramolecular (within the same molecule, e.g., o-nitrophenol). Hydrogen bond strength is maximum in the solid state and minimum in the gaseous state. Hydrogen bonds profoundly influence the boiling point, solubility, viscosity, and biological properties of compounds.

Chapter 4 provides a comprehensive framework for understanding how and why atoms combine to form molecules. Beginning with the Kossel-Lewis approach and the octet rule, it progresses through ionic and covalent bonding, bond parameters (length, angle, enthalpy, order, polarity), resonance, VSEPR theory, valence bond theory with hybridisation, and finally molecular orbital theory. Each theory addresses limitations of the previous, together giving a near-complete picture of chemical bonding from the simplest diatomic molecules to complex polyatomic systems.

For board examinations, this chapter is extremely high-weightage. Students must master Lewis structure drawing, formal charge calculation, VSEPR-based geometry prediction, hybridisation assignment (sp, sp2, sp3, sp3d, sp3d2), MO electronic configurations and bond order calculations for homonuclear diatomics (H2 through Ne2), and hydrogen bonding. Conceptual clarity about why O2 is paramagnetic, how resonance stabilises molecules, and the factors favouring ionic bond formation is essential for scoring high marks.