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19 facts

01Octet exceptions: incomplete (BeH₂, BCl₃, LiCl), odd-electron (NO, NO₂), expanded (PF₅, SF₆, H₂SO₄).
02Formal charge = V − L − ½S.
03Lattice enthalpy of NaCl = 788 kJ mol⁻¹.
041 D = 3.33564 × 10⁻³⁰ C m.
05Zero dipole moment: BeF₂, BF₃, CO₂, CH₄, CCl₄; non-zero: H₂O, NH₃, NF₃.
06Dipole moment of NH₃ > NF₃ because the lone-pair moment adds in NH₃ and opposes in NF₃.
07Repulsion order: lp-lp > lp-bp > bp-bp.
08Bond angles: CH₄ 109.5° > NH₃ ~107° > H₂O ~104.5°.
09Shapes: SF₄ see-saw, ClF₃ T-shaped, XeF₂ linear, BrF₅ square pyramidal, XeF₄ square planar, SF₆ octahedral.
10PCl₅: sp³d, axial bonds longer than equatorial bonds.
11sp³d uses dz²; sp³d² uses dz² and dx²−y².
12s character: sp 50%, sp² 33%, sp³ 25%.
13σ bond is stronger than π bond; double bond = 1σ + 1π, triple bond = 1σ + 2π.
14Bond order = ½(Nb − Na).
15σ2pz lies above π2p up to N₂, and below π2p for O₂ and F₂.
16O₂ is paramagnetic with bond order 2; N₂ is diamagnetic with bond order 3.
17Bond orders: O₂⁺ 2.5, O₂ 2, O₂⁻ 1.5, O₂²⁻ 1.
18B₂ is paramagnetic; C₂ is diamagnetic.
19Hydrogen bonding needs H attached to F, O or N; o-nitrophenol shows intramolecular H-bonding.

Common traps

4 to know

Formal charge

FC = V − L − ½S

V = valence electrons of free atom, L = non-bonding (lone) electrons, S = shared (bonding) electrons.

Dipole moment

μ = Q × r

Unit C m or Debye; 1 D = 3.33564 × 10⁻³⁰ C m. Net μ is the vector sum of bond moments.

Bond order (MO theory)

Bond order = ½(Nb − Na)

Nb = electrons in bonding MOs, Na = electrons in antibonding MOs. Zero means the molecule is not stable.

Hybridisation count

Steric number = σ bonds + lone pairs on central atom

2 sp, 3 sp², 4 sp³, 5 sp³d, 6 sp³d².

13 terms

Octet rule: Tendency of atoms to reach eight valence electrons by losing, gaining or sharing electrons.
Formal charge: Charge assigned to an atom in a Lewis structure by assuming shared electrons are split equally.
Lattice enthalpy: Energy needed to break one mole of an ionic solid completely into gaseous ions.
Bond enthalpy: Energy needed to break one mole of a particular bond in gaseous molecules.
Resonance hybrid: The single real structure that several Lewis structures together approximate.
Dipole moment: Product of the separated charge and the distance between charge centres; a vector.
Sigma bond: Bond formed by head-on overlap of orbitals along the internuclear axis.
Pi bond: Bond formed by sideways overlap of parallel p orbitals.
Hybridisation: Mixing of atomic orbitals of similar energy on one atom into equivalent new orbitals.
LCAO: Building molecular orbitals by adding and subtracting atomic orbital wave functions.
Antibonding orbital: A molecular orbital higher in energy than the parent atomic orbitals, which weakens bonding when occupied.
Paramagnetic: Attracted into a magnetic field because of unpaired electrons.
Hydrogen bond: Attraction between an H atom bonded to F, O or N and a lone pair on another such atom.