Why do atoms bond?
Atoms combine because the bonded arrangement has lower energy than the separate atoms. The Kössel–Lewis approach says that atoms tend to attain a noble-gas configuration (usually eight valence electrons, the octet) by losing, gaining or sharing electrons.
Ionic bonds
An ionic (electrovalent) bond forms when electrons are transferred from an atom with low ionisation enthalpy (typically a metal) to one with high electron-gain enthalpy (typically a non-metal). The oppositely charged ions attract each other.
Na (2, 8, 1) → Na⁺ (2, 8) + e⁻, and Cl (2, 8, 7) + e⁻ → Cl⁻ (2, 8, 8).
The stability of an ionic solid comes mainly from its lattice enthalpy: the energy released when the gaseous ions come together to form the crystal lattice.
Covalent bonds and Lewis structures
A covalent bond forms when atoms share a pair of electrons. Sharing one, two or three pairs gives single, double or triple bonds (as in H₂, O₂ and N₂).
Formal charge helps choose the best Lewis structure:
FC = (valence electrons) − (non-bonding electrons) − ½(bonding electrons)
The best structure has the smallest formal charges on the atoms.
Limitations of the octet rule
Incomplete octet: BeCl₂ (4 electrons around Be) and BF₃ (6 around B).
Expanded octet: elements of period 3 and beyond can exceed 8, as in PCl₅ (10) and SF₆ (12).
Odd-electron molecules: NO and NO₂.
VSEPR theory: predicting shapes
According to Valence Shell Electron Pair Repulsion theory, the electron pairs around a central atom arrange themselves as far apart as possible. Lone pairs repel more strongly than bond pairs:
lone pair–lone pair > lone pair–bond pair > bond pair–bond pair
Molecule | Bond pairs | Lone pairs | Shape | Bond angle |
|---|---|---|---|---|
BeCl₂ | 2 | 0 | linear | 180° |
BF₃ | 3 | 0 | trigonal planar | 120° |
CH₄ | 4 | 0 | tetrahedral | 109.5° |
NH₃ | 3 | 1 | trigonal pyramidal | 107° |
H₂O | 2 | 2 | bent (V-shaped) | 104.5° |
PCl₅ | 5 | 0 | trigonal bipyramidal | 120° and 90° |
SF₆ | 6 | 0 | octahedral | 90° |
The extra lone pairs on NH₃ and H₂O squeeze the bond angles below the ideal 109.5°.
Hybridisation
Atomic orbitals of similar energy on the central atom mix to form an equal number of identical hybrid orbitals.
Hybridisation | Geometry | Example |
|---|---|---|
sp | linear | BeCl₂, C₂H₂ |
sp² | trigonal planar | BF₃, C₂H₄ |
sp³ | tetrahedral | CH₄, NH₃, H₂O |
sp³d | trigonal bipyramidal | PCl₅ |
sp³d² | octahedral | SF₆ |
Sigma and pi bonds. A single bond is one σ bond. A double bond is one σ + one π bond, and a triple bond is one σ + two π bonds. σ bonds come from head-on overlap and are stronger than π bonds, which come from sideways overlap.
Bond parameters
Bond length: the distance between the nuclei of bonded atoms. Triple < double < single.
Bond enthalpy: the energy needed to break one mole of bonds. Triple > double > single.
Bond order: the number of bonds between two atoms. Higher bond order means a shorter, stronger bond.
Polarity and dipole moment
When atoms of different electronegativity share electrons, the bond is polar: one end is slightly negative (δ−) and the other slightly positive (δ+). The dipole moment μ = charge × distance measures this, in debye (D).
A molecule can have polar bonds but zero overall dipole moment if its shape makes the bond dipoles cancel. CO₂ (linear), BF₃ and CCl₄ are non-polar. H₂O (bent) and NH₃ (pyramidal) are polar.
Hydrogen bonding
A hydrogen bond is an attraction between a hydrogen atom bonded to a highly electronegative atom (F, O or N) and a lone pair on another such atom.
Intermolecular H-bonding (between molecules) raises boiling points. This is why water is a liquid at room temperature while H₂S is a gas.
Intramolecular H-bonding (within one molecule, as in o-nitrophenol) lowers the boiling point compared with isomers that bond between molecules.