Chemistry (university) · Part 02
Chemical bonding and molecular shape
Ionic and covalent bonds, Lewis structures, molecular geometry, hybridization and molecular orbitals.
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- ElectronegativityElectronegativity is how strongly an atom pulls shared electrons toward itself in a bond. It rises across a period and falls down a group. Fluorine is the highest on the Pauling scale, at 4.0.
- Ionic bonding and lattice energyIn an ionic bond, one atom hands electrons to another, making positive and negative ions that attract. The ions pack into a crystal lattice. Lattice energy is the energy released when gaseous ions form one mole of the solid.
- Lewis structureA Lewis structure shows every valence electron as bonds (shared pairs) or lone pairs. Count total valence electrons, join atoms with single bonds, then add lone pairs so each atom nears an octet. Hydrogen needs only two.
- Formal chargeFormal charge is a bookkeeping number for an atom in a Lewis structure: valence electrons minus lone-pair electrons minus number of bonds. The best structure has formal charges as close to zero as possible.
- ResonanceSometimes one Lewis structure cannot show a molecule correctly. Several equal structures, called resonance forms, differ only in where electrons are placed. The real molecule is a blend (hybrid) of them.
- Bond polarity and dipole momentA bond is polar when electrons are shared unequally, giving one end a partial negative charge (δ−) and the other a partial positive charge (δ+). A molecule is polar if its bond dipoles do not cancel.
- VSEPR theoryVSEPR stands for valence shell electron pair repulsion. Electron groups around a central atom (bonds and lone pairs) repel and spread as far apart as possible. The count of groups sets the geometry.
- HybridizationHybridization mixes atomic orbitals on one atom to make new, equal orbitals that match the observed shape. One s plus three p gives four sp³; one s plus two p gives three sp²; one s plus one p gives two sp.
- Molecular orbital theoryIn molecular orbital (MO) theory, atomic orbitals combine to make orbitals spread over the whole molecule. Combining in phase gives a lower-energy bonding orbital. Out of phase gives a higher-energy antibonding orbital. Bond order = (bonding − antibonding electrons) / 2.