2.2 The covalent model

Build a Lewis formula by counting all valence electrons, choosing a skeleton, completing outer octets and placing any remainder on the central atom. Then recount electrons and formal charge; a shared pair contributes to the attraction between both nuclei rather than belonging exclusively to either atom.

Charged-species check: NH₄⁺ has 5 + 4(1) − 1 = 8 valence electrons, so draw four N–H shared pairs, no lone pair on N, enclose the ion in brackets and write the overall + charge. A final audit must match the available electron total, complete the required outer-shell arrangements and reproduce the species charge; lone pairs, shared pairs and formal charges answer different parts of that audit.

For the same pair of atoms, increasing bond order generally shortens and strengthens the bond because more electron density lies between the nuclei. Use that comparison locally: bond enthalpy also depends on the atoms and molecular environment, so any triple bond is not automatically stronger than every unrelated single bond.

In NH₃→BF₃ the nitrogen lone pair supplies both bonding electrons, so the arrow starts at N and ends at B. The arrow records how the pair originated; after formation it is a shared covalent pair, and in a complex the same donor logic identifies each ligand–metal bond.

NH₃ has four electron domains around N, so its electron-domain geometry is tetrahedral but its molecular shape is trigonal pyramidal; the lone pair compresses the H–N–H angle below 109.5°. Count a double bond as one domain and distinguish electron geometry from the shape named using atoms only.

For H–Cl, chlorine is more electronegative, so label Hδ⁺–Clδ⁻ and point the dipole arrow toward Cl. Electronegativity difference predicts unequal sharing within that bond; it does not by itself decide the polarity of the whole molecule.

CO₂ contains polar C=O bonds, but their equal opposite dipoles cancel in a linear molecule. In bent H₂O they do not cancel, so the molecule has a net dipole. Always establish the three-dimensional geometry before adding dipoles as vectors.

Diamond is hard because each carbon is held in a three-dimensional covalent network, while graphite conducts along layers through delocalized electrons and its layers can slide. Silicon dioxide is also an extended network: describe network atoms, not discrete SiO₂ molecules, when explaining its high melting point.

All molecules have London dispersion forces. Add permanent dipole–dipole attraction when a net molecular dipole exists, and add hydrogen bonding only when the required H–N, H–O or H–F environment and an acceptor lone pair are present. Name every relevant force before deciding which dominates.

Compare like evidence: pentane has stronger dispersion forces and a higher boiling point than butane because its electron cloud is larger. The simple London < dipole–dipole < hydrogen-bond ordering is a guide for comparable molecules, not a rule that ignores molecular size and the number of interaction sites.

If a spot moves 3.2 cm while the solvent front moves 8.0 cm, Rf = 0.40. A larger Rf means greater relative affinity for the mobile phase under those conditions, but values from different solvents, stationary phases or temperatures are not directly interchangeable.