JEE Chemistry Chemical Bonding And Molecular Structure Notes Set 05

Why Atoms Bond

Atoms combine to lower their energy and, usually, to reach a stable noble-gas arrangement of eight valence electrons (octet) or two (duplet for hydrogen). Bond formation always releases energy. Kossel explained the ionic bond by transfer of electrons and Lewis the covalent bond by sharing of electrons.
Exceptions to the octet rule: electron-deficient molecules (BeCl₂, BF₃, AlCl₃, B₂H₆), expanded octets (PCl₅, SF₆, IF₇) and odd-electron molecules (NO, NO₂).

Ionic (Electrovalent) Bond

An ionic bond forms between an element of low ionisation energy (metal) and one of high electron affinity and high electronegativity (non-metal), and is favoured by a large lattice energy.
Lattice energy rises with ionic charge and falls with ionic size, so MgO > NaF > NaCl > KCl. Born-Haber cycles use Hess's law to calculate it.
Properties: hard, brittle crystals with high melting and boiling points; conduct when molten or dissolved but not as solids; soluble in polar solvents like water, where hydration energy overcomes lattice energy; non-directional bonds, so no space isomerism. Ionic compounds with the same charge show higher melting points for smaller ions (NaF has a higher melting point than NaI).

Covalent Bond and Fajan's Rules

A covalent bond is a shared pair of electrons. Valency is the number of electrons an atom shares. Bond length, bond energy and bond order describe a bond: more shared pairs give a shorter and stronger bond (C≡C < C=C < C–C in length).
Fajan's rules say that covalent character in an ionic bond increases with a small, highly charged cation (high polarising power), a large, highly charged anion (high polarisability), and a cation with an 18-electron (pseudo-noble-gas) core. Hence LiCl is more covalent than NaCl, BeCl₂ is more covalent than MgCl₂, AlCl₃ more covalent than NaCl, and AgCl more covalent than NaCl.
Coordinate (dative) bond: both bonding electrons come from one atom, as in NH₄⁺, H₃O⁺, NH₃→BF₃ and O₃. Compounds such as CuSO₄.5H₂O and K₄[Fe(CN)₆] contain ionic, covalent and coordinate bonds together.

Polarity and Dipole Moment

A bond between atoms of different electronegativity is polar. Dipole moment μ = charge × distance (unit debye, 1 D = 3.34 × 10⁻³⁰ C m). Percentage ionic character = (observed μ / μ for 100% ionic) × 100.
A molecule's dipole moment is the vector sum of its bond dipoles. Zero: CO₂, CS₂, BeF₂ (linear), BF₃, SO₃ (trigonal planar), CH₄, CCl₄ (tetrahedral), SF₆, XeF₄, benzene, p-dichlorobenzene, trans-1,2-dichloroethene. Non-zero: H₂O, NH₃, SO₂, CHCl₃, CH₃Cl, o- and m-dichlorobenzene, cis-dichloroethene. NF₃ has a much smaller dipole moment than NH₃ because the N–F bond dipoles oppose the lone pair.

Sigma and Pi Bonds, Hybridisation

A sigma bond is formed by head-on overlap (s–s, s–p, p–p, hybrid–hybrid) and is strong; a pi bond is formed by sidewise overlap of parallel p orbitals and is weaker. A single bond is one sigma; a double bond is one sigma and one pi; a triple bond is one sigma and two pi.
Hybridisation is the mixing of orbitals of similar energy of one atom to give equivalent hybrid orbitals. sp: linear, 180°, 50% s; sp²: trigonal planar, 120°, 33% s; sp³: tetrahedral, 109.5°, 25% s; sp³d: trigonal bipyramidal; sp³d²: octahedral; sp³d³: pentagonal bipyramidal. Bond length decreases and bond angle increases as s-character increases.
To find the hybridisation, count sigma bonds plus lone pairs on the central atom: 2 gives sp, 3 gives sp², 4 gives sp³, 5 gives sp³d, 6 gives sp³d², 7 gives sp³d³.

VSEPR Theory and Shapes

Electron pairs around the central atom arrange themselves to be as far apart as possible. The repulsion order is lone pair–lone pair > lone pair–bond pair > bond pair–bond pair, so lone pairs compress bond angles (CH₄ 109.5°, NH₃ 107°, H₂O 104.5°). Bond angles of hydrides of one group fall down the group (NH₃ > PH₃ > AsH₃ > SbH₃).
Examples: linear (BeCl₂, CO₂, XeF₂, I₃⁻); trigonal planar (BF₃, SO₃, CO₃²⁻, NO₃⁻); bent (SO₂, H₂O, ClO₂⁻); tetrahedral (CH₄, NH₄⁺, SO₄²⁻, BF₄⁻); pyramidal (NH₃, PCl₃, H₃O⁺, ClO₃⁻); see-saw (SF₄); T-shaped (ClF₃, ICl₃); trigonal bipyramidal (PCl₅); square pyramidal (BrF₅); octahedral (SF₆); square planar (XeF₄); pentagonal bipyramidal (IF₇); distorted octahedral (XeF₆).

Resonance

When one Lewis structure cannot describe a molecule, the real structure is a resonance hybrid of several contributing structures which differ only in the positions of electrons, not of atoms. The hybrid is more stable than any single structure, and the extra stability is the resonance energy. Examples: benzene (all C–C bonds equal), CO₃²⁻ (three structures, each C–O bond order 1.33), NO₃⁻, O₃, CO₂.

Molecular Orbital Theory

Atomic orbitals combine to form bonding molecular orbitals (lower energy) and antibonding molecular orbitals (higher energy). Electrons fill them by the Aufbau, Pauli and Hund rules. Bond order = ½(Nb − Na); a stable molecule needs a positive bond order, and a higher bond order means a shorter, stronger bond.
Energy order for O₂ and heavier: σ1s, σ*1s, σ2s, σ*2s, σ2p_z, π2p_x = π2p_y, π*2p_x = π*2p_y, σ*2p_z. For B₂, C₂, N₂ (and lighter) the π orbitals lie below σ2p_z.
Results: N₂ bond order 3 (diamagnetic); O₂ bond order 2 with two unpaired electrons in π* (paramagnetic); O₂⁺ 2.5, O₂⁻ 1.5, O₂²⁻ 1; NO 2.5 (paramagnetic), NO⁺ 3; He₂ bond order 0, so it does not exist; H₂⁺ and He₂⁺ have bond order ½. Bond length: O₂⁺ < O₂ < O₂⁻ < O₂²⁻.

Hydrogen Bonding

A hydrogen bond is an attraction between a hydrogen atom bonded to F, O or N and a lone pair on another F, O or N atom. Its strength (10–40 kJ/mol) lies between van der Waals forces and covalent bonds.
Intermolecular hydrogen bonding raises melting and boiling points and solubility in water: HF, H₂O and NH₃ boil higher than the other hydrides of their groups; ethanol boils higher than dimethyl ether. Intramolecular hydrogen bonding (o-nitrophenol, salicylaldehyde) lowers the boiling point and makes the compound steam volatile. Ice is less dense than water because of its open hydrogen-bonded cage, with each water molecule bonded to four others.

Forces in Solids

Ionic solids (NaCl) are held by electrostatic attraction; covalent network solids (diamond, silica, SiC) by a network of covalent bonds and are very hard with very high melting points; metals by the attraction between a sea of mobile electrons and positive ions; molecular solids (dry ice, ice, solid argon) by weak van der Waals or hydrogen bonds and have low melting points.

Solved Examples

Example 1. Find the hybridisation and shape of XeF₄.
Answer:
Xenon has 8 valence electrons; four are used for bonds with F, leaving two lone pairs.
Four bond pairs plus two lone pairs make six electron pairs, so xenon is sp³d² hybridised.
The two lone pairs sit opposite each other, so the shape is square planar.

Example 2. Calculate the bond order of O₂⁻ and say whether it is paramagnetic.
Answer:
O₂⁻ has 17 electrons: 10 in bonding and 7 in antibonding orbitals.
Bond order = (10 − 7)/2 = 1.5.
The last electron is unpaired (π*), so the ion is paramagnetic.

Example 3. The dipole moment of HCl is 1.03 D and its bond length is 1.27 Å. Find the percentage ionic character.
Answer:
For 100% ionic character μ = e × d = 4.8 × 10⁻¹⁰ esu × 1.27 × 10⁻⁸ cm = 6.1 × 10⁻¹⁸ esu cm = 6.1 D.
Percentage ionic character = (1.03/6.1) × 100 = 17%.

Example 4. Arrange NH₃, H₂O, CH₄ and CO₂ in increasing order of bond angle.
Answer:
CO₂ is sp (180°), CH₄ is sp³ with no lone pair (109.5°), NH₃ has one lone pair (107°) and H₂O has two lone pairs (104.5°).
Order: H₂O < NH₃ < CH₄ < CO₂.