JEE Chemistry Chemical Bonding and Molecular Structure Notes Set 01

Why Atoms Bond

Atoms combine to lower their energy. Kossel and Lewis explained this through the octet rule: atoms tend to reach eight electrons (two for hydrogen) in the valence shell by losing, gaining or sharing electrons. Exceptions are electron-deficient molecules (\( \ce{BeCl2} \), \( \ce{BF3} \)), odd-electron molecules (\( \ce{NO} \), \( \ce{NO2} \)) and expanded-octet molecules (\( \ce{PCl5} \), \( \ce{SF6} \)) that use d-orbitals of period 3 and later elements. Period 2 elements never expand the octet, which is why \( \ce{NF5} \) and \( \ce{OF4} \) do not exist.

Ionic Bond

An ionic bond forms by complete transfer of electrons. It is favoured by a low ionisation energy of the metal, a high electron affinity of the non-metal and, above all, a high lattice energy. Lattice energy cannot be measured directly and is found with the Born-Haber cycle (Hess's law). Ionic solids are hard, brittle, have high melting and boiling points, are poorly soluble in non-polar solvents and conduct only when molten or dissolved.

A salt dissolves when hydration energy exceeds lattice energy. Small, highly charged ions have large hydration energies. Hence \( \ce{MgSO4} \) and \( \ce{BeSO4} \) are soluble, while sulphate solubility falls steadily from \( \ce{CaSO4} \) to \( \ce{BaSO4} \). Hydrated ions are bigger for smaller bare ions, so the mobility of alkali ions rises from \( \ce{Li+} \) to \( \ce{Cs+} \).

Fajans' Rules and Partial Covalent Character

Covalent character in an ionic compound grows when (1) the cation is small, (2) the cation carries a high charge, (3) the anion is large and (4) the cation has an 18-electron (pseudo-noble) shell instead of a noble-gas shell. Examples: \( \ce{LiCl} \) is more covalent than \( \ce{NaCl} \); \( \ce{AlI3} \) is more covalent than \( \ce{AlF3} \); \( \ce{SnCl4} \) is molecular while \( \ce{SnCl2} \) is more ionic; \( \ce{BeCl2} \) is covalent while \( \ce{BaCl2} \) is ionic. Carbonate stability follows the same logic: a strongly polarising cation pulls oxide character out of the carbonate and decomposes it to \( \ce{CO2} \) more easily, so stability rises down group 2 and group 1 carbonates (except lithium) do not decompose.

Covalent Bond, Sigma and Pi Overlap

A covalent bond is the sharing of electron pairs. Axial (head-on) overlap gives a sigma bond; sideways overlap of p or d orbitals gives a pi bond. Overlap strength (and so bond strength) is s-s lower than p-p lower than s-p for the directional case, and sigma is stronger than pi. A double bond is one sigma plus one pi; a triple bond is one sigma plus two pi. Pi bonds are never made from hybrid orbitals. Bond order, bond length and bond energy are linked: the higher the order, the shorter and stronger the bond. A coordinate (dative) bond is a covalent bond in which one atom supplies both electrons, as in \( \ce{NH4+} \), \( \ce{H3O+} \), \( \ce{BF3.NH3} \), \( \ce{O3} \) and \( \ce{SO3} \).

Hybridisation and VSEPR Shapes

Count the sigma bonds plus lone pairs on the central atom to find hybridisation. The shape follows from the arrangement of electron pairs with lone pairs removed. The order of repulsion is lone pair-lone pair larger than lone pair-bond pair larger than bond pair-bond pair, so lone pairs compress bond angles. Greater s-character in a hybrid orbital widens its angle (sp 180, sp2 120, sp3 109.5 degrees).

Bond angles in hydrides fall as the central atom becomes less electronegative: \( \ce{NH3} \) (107) larger than \( \ce{PH3} \) (93). In fluorides and oxides the opposite pattern holds: \( \ce{NF3} \) is smaller than \( \ce{NH3} \) because the electronegative ligands draw the bond pairs away from the centre. Angles of the nitrogen dioxide family rise \( \ce{NO2-} \) (115) \( \lt \) \( \ce{NO2} \) (134) \( \lt \) \( \ce{NO2+} \) (180).

Polarity and Dipole Moment

Dipole moment is charge times distance, \( \mu = q \times d \), measured in debye (1 D = 3.336 x 10^-30 C m). It is a vector sum of bond dipoles, so symmetrical molecules (\( \ce{CO2} \), \( \ce{BF3} \), \( \ce{CCl4} \), \( \ce{PCl5} \), \( \ce{SF6} \), \( \ce{XeF4} \), benzene, trans-dichloroethene, p-dichlorobenzene) are non-polar, while \( \ce{H2O} \), \( \ce{SO2} \), \( \ce{NH3} \), \( \ce{SF4} \) and the cis isomer are polar. In \( \ce{NH3} \) the lone-pair moment adds to the N-H bond moments (1.47 D) whereas in \( \ce{NF3} \) it opposes the N-F moments (0.2 D). Percentage ionic character equals (observed moment / calculated moment for full charges) x 100. Chloromethanes: \( \ce{CH3Cl} \) 1.9, \( \ce{CH2Cl2} \) 1.6, \( \ce{CHCl3} \) 1.0, \( \ce{CCl4} \) 0 D.

Molecular Orbital Theory

Atomic orbitals combine linearly into bonding (lower energy) and antibonding (higher energy, marked with a star) molecular orbitals. Electrons fill in order of energy, obeying the Aufbau principle, Pauli exclusion and Hund's rule. Bond order is half of the difference between bonding and antibonding electrons. A molecule exists only if bond order is above zero (\( \ce{Be2} \) and \( \ce{He2} \) do not).

For molecules up to \( \ce{N2} \) (14 electrons) the order is sigma 1s, sigma* 1s, sigma 2s, sigma* 2s, pi 2px = pi 2py, sigma 2pz, pi* 2px = pi* 2py, sigma* 2pz. From \( \ce{O2} \) onwards sigma 2pz drops below the pi orbitals. Species with unpaired electrons are paramagnetic; those with all electrons paired are diamagnetic.

Isoelectronic species have identical bond orders. Removing an electron from an antibonding orbital (\( \ce{O2 -> O2+} \), NO to \( \ce{NO+} \)) raises the order and shortens the bond; removing a bonding electron (\( \ce{N2 -> N2+} \)) lowers it.

Hydrogen Bonding and Other Forces

A hydrogen bond is an attraction between H bonded to F, O or N and a lone pair on another such atom. Strength: F-H...F is greater than O-H...O is greater than N-H...N. Intermolecular hydrogen bonding raises boiling point, viscosity and solubility in water (HF, \( \ce{H2O} \), \( \ce{NH3} \), alcohols, glycerol, carboxylic acid dimers, p-nitrophenol). Intramolecular hydrogen bonding (o-nitrophenol, salicylaldehyde) makes a chelate ring and increases volatility. Ice has an open cage in which each water molecule forms four hydrogen bonds, so ice is less dense than water; the anion \( \ce{HF2-} \) in \( \ce{KHF2} \) is held by a very strong F-H...F bond. Other forces: London dispersion (all molecules, grows with size), dipole-dipole and dipole-induced dipole. Dry ice, solid methane, solid neon and solid \( \ce{SO2} \) are molecular solids, whereas \( \ce{SiO2} \), diamond and SiC are covalent network solids.