Bonding & structure

Ionic, covalent, metallic · shapes of molecules · intermolecular forces

Shapes of molecules (electron-pair repulsion)

• Pairs of electrons repel to maximum separation; lone pairs repel more than bonding pairs, closing bond angles by ≈2.5° each.
• 2 pairs — linear, 180° (CO₂). 3 — trigonal planar, 120° (BF₃). 4 — tetrahedral, 109.5° (CH₄). 3+1 lone — pyramidal, 107° (NH₃). 2+2 lone — bent, 104.5° (H₂O). 6 — octahedral, 90° (SF₆).

Intermolecular forces, weakest → strongest

• Induced dipole–dipole (London/dispersion) — in everything; grows with electron count and surface contact (why boiling points rise down group 7 and fall with branching).
• Permanent dipole–dipole — between polar molecules (e.g. HCl).
• Hydrogen bonding — H bonded to N, O or F, attracted to a lone pair on N/O/F of another molecule. Explains water's high boiling point and why ice is less dense than water (open lattice held by H-bonds).

Structure determines properties

• Giant ionic: high melting point, conducts only molten/aqueous.
• Giant covalent (diamond, graphite, SiO₂): very high melting point; graphite conducts via delocalised electrons between layers.
• Simple molecular: low melting point — only weak intermolecular forces break, not covalent bonds.

3.1.3.5Shapes summary — learn this table cold

3.1.3.2Co-ordinate (dative covalent) bonding

• Both shared electrons come from one atom: NH₄⁺ (N donates to H⁺), H₃O⁺, Al₂Cl₆, and every metal–ligand bond in complex ions.
• Once formed it is identical to any other covalent bond — all four N–H bonds in NH₄⁺ are equivalent.

3.1.3.1The four crystal types

Extended🎓 Beyond the standard course

• Dipoles add as vectors. CO₂'s two large bond dipoles cancel; CHCl₃'s don't. cis/trans isomers can differ in polarity — trans-1,2-dichloroethene is non-polar, the cis isomer polar.
• Bond order sets length and strength: C–C ≈ 154 pm, C=C ≈ 134 pm, C≡C ≈ 120 pm. Benzene's six equal 140 pm bonds — between single and double — are the structural evidence for delocalisation.
• Polarisability: big, soft anions (I⁻) are easily distorted by small, highly-charged cations — the origin of covalent character in "ionic" solids like AgI, and the deeper reason London forces grow with electron count.
• Hydrogen bonds are ~20 kJ mol⁻¹ — an order of magnitude weaker than covalent bonds (~350) but decisive in bulk: water's boiling point sits ~200 °C above the group trend. Water out-boils HF because each H₂O averages more hydrogen bonds per molecule (two donors, two acceptors).
• Ice's density anomaly is a lattice constraint: the tetrahedral H-bond network enforces open hexagonal channels that collapse on melting.

Deep dive📚 The rest of the chapter, in full

The complete shapes table

Each lone pair squeezes bonding angles by about 2.5° — lone-pair/lone-pair repulsion > lone-pair/bond > bond/bond. For unfamiliar ions, count electrons: for NH₄⁺, N has 5 outer electrons + 4 from H − 1 for the positive charge = 8 = 4 bonding pairs → tetrahedral.

Polarity: bond vs molecule

A bond is polar when the two atoms differ in electronegativity (F 4.0 > O 3.4 > N ≈ Cl 3.0). A molecule is only polar if the bond dipoles do not cancel. CO₂ (linear) and CCl₄ (tetrahedral) contain polar bonds but are non-polar molecules — the dipoles cancel by symmetry. H₂O and CHCl₃ are polar because theirs do not. This distinction is worth a mark almost every year.

Ranking boiling points — the examiner's favourite

• Hydrogen bonding needs H bonded directly to N, O or F, plus a lone pair on the acceptor. It explains why H₂O boils at 100 °C while H₂S boils at −60 °C, why ice is less dense than water (an open hydrogen-bonded lattice), and why NH₃, H₂O and HF break their groups' trends.
• Permanent dipole–dipole forces act between polar molecules (e.g. CH₃Cl).
• Induced dipole (London) forces act between everything and grow with electron count and contact area: boiling points climb along the alkanes, and branched isomers boil lower than straight chains (2,2-dimethylpropane −10 °C vs pentane 36 °C) because branching reduces surface contact.

The four crystal types

Graphite conducts because each carbon bonds to only three others, leaving one delocalised electron per atom between layers; the layers slide (lubricant). Diamond's four bonds per carbon make it the hardness benchmark. Graphene is a single graphite layer — the strongest known material and an excellent conductor.

Mastery vault🏛 Every remaining spec point, banked

Ionic bonding, quantitatively

Lattice attraction grows with charge and shrinks with ionic radius: MgO (2+/2−, small ions) melts at 2852 °C vs NaCl (1+/1−) at 801 °C — which is why MgO lines furnaces. Isoelectronic series N³⁻ > O²⁻ > F⁻ > Na⁺ > Mg²⁺ > Al³⁺ shrink steadily: same electrons, rising nuclear charge. Evidence for ions: molten and aqueous conductivity, and electrolysis migrating coloured ions (Cu²⁺ blue to the cathode, CrO₄²⁻ yellow to the anode).

Electronegativity values worth memorising

ΔEN ≳ 1.8 → predominantly ionic; 0.4–1.8 → polar covalent; < 0.4 → essentially non-polar. C–H (Δ0.4) counts as non-polar — the reason alkanes have only London forces. The ionic/covalent boundary is a continuum: large ΔEN with a polarising cation (AlCl₃) lands in between, subliming at 180 °C like a covalent solid.

Force magnitudes — why rankings work

Melting a simple molecular solid breaks only the intermolecular forces — the covalent bonds survive (I₂ vapour is still I₂). Only macromolecular melting breaks covalent bonds. This single distinction settles most "explain the melting point" questions.

Solubility logic

• Ionic solids dissolve in water when hydration energy (ion–dipole attraction) repays lattice energy; the δ− oxygens surround cations, δ+ hydrogens surround anions.
• Alcohols mix with water via hydrogen bonding — but solubility falls as the non-polar chain grows (hexan-1-ol barely dissolves).
• Non-polar solutes dissolve in non-polar solvents (iodine in cyclohexane, purple) — "like dissolves like" is the pattern, argued through which forces form vs break.

Harder shape calls

• ClF₃: 7 + 3 = 10 electrons around Cl → 3 bond pairs + 2 lone pairs → T-shaped, ~87.5°.
• SF₆ vs SF₄: 6 pairs octahedral 90° vs 5 pairs with one lone pair → seesaw.
• Ions: subtract electrons for +, add for −. NO₃⁻: N has 5 + 3 shared considerations → trigonal planar 120° with delocalised π — quote 120° and equal bond lengths.
• Two lone-pair rule of thumb: each lone pair −2.5° from the parent angle (109.5 → 107 → 104.5).