Chemical Bonding and Structure
Chemical Bonding and Structure
The exam tests whether you can choose the correct bonding model for a substance, predict molecular shape and polarity, and use structure to explain physical properties such as melting point, solubility and conductivity.
Shapes from electron-pair repulsion
Count electron pairs around the central atom and arrange them to minimise repulsion; lone pairs repel more strongly than bond pairs. Two pairs give linear (180 degrees), three trigonal planar (120), four tetrahedral (109.5), five trigonal bipyramidal, six octahedral (90). Lone pairs compress angles: CH4 109.5, NH3 107, H2O 104.5. Double bonds count as one region of electron density but repel slightly more than single bonds (SO2, 119 degrees).
Bonding models
Ionic lattices have high melting points and conduct only when molten or aqueous. Simple molecular substances have weak intermolecular forces, so low melting points and no conductivity. Giant covalent solids (diamond, SiO2, graphite in-plane) are hard and refractory. Metals conduct because of delocalised electrons and are malleable because layers slide. Intermolecular forces rank: London (induced dipole, stronger with more electrons) < permanent dipole-dipole < hydrogen bonding, which requires H bonded directly to N, O or F.
Worked example
Explain why water boils at 100 degrees C but hydrogen sulfide at about -60 degrees C, although both are bent molecules of similar size. Oxygen is much more electronegative than sulfur, so the O-H bond is strongly polar and the hydrogen atom interacts with a lone pair on oxygen of a neighbouring molecule: water molecules are extensively hydrogen bonded. H2S has only weaker permanent dipole-dipole attractions plus London forces. Boiling breaks intermolecular forces, not covalent bonds, so the difference in those forces explains the 160-degree gap.
Traps
Claiming hydrogen bonding whenever hydrogen is present; confusing electron-pair geometry with molecular shape; forgetting to subtract lone-pair repulsion when quoting angles; comparing bond enthalpies when the question asks about melting; and drawing incorrect total electron counts in dot-and-cross diagrams of ions such as NH4+ and O2-.
Key ideas
- Shape comes from electron-pair repulsion; lone pairs reduce bond angles.
- Hydrogen bonding needs H directly attached to N, O or F.
- Giant covalent and ionic structures melt high; molecular structures melt low.
- Polarity depends on shape as well as bond dipoles: CO2 is non-polar, H2O is polar.