Chapter 4: Chemical Bonding

Ionic bonding is described as the potent electrostatic attraction between positive and negative ions situated within an ionic crystal lattice, exemplified by compounds like sodium chloride or calcium fluoride, typically arising from electron transfer to complete outer shells. Covalent bonding involves the sharing of outer electrons, ranging from single bonds to stronger double and triple bonds. Specialized covalent formations covered include co-ordinate bonding (or dative covalent bonding), where one atom provides both electrons for the shared pair, as seen in the ammonium ion or the dimer aluminium chloride (Al2Cl6). Furthermore, the chapter addresses exceptions to the stable octet configuration, noting electron-deficient atoms (like boron in BF3) and atoms that exhibit expanded octets (such as sulfur in SF6). The energy required to break these bonds is quantified as bond energy, related inversely to the bond length (internuclear distance). Molecular geometry is governed by the Valence Shell Electron Pair Repulsion (VSEPR) theory, which predicts shapes and bond angles by minimizing repulsive forces between electron pairs, with lone pair repulsion being significantly greater than bonding pair repulsion, resulting in shapes like the tetrahedral structure of methane, the pyramidal structure of ammonia, or the non-linear V shape of water. Covalent bond formation is explained through orbital overlap and hybridisation, differentiating between sigma bonds (formed by linear end-on overlap) and pi bonds (formed by sideways overlap of p orbitals); single bonds are purely sigma, while double and triple bonds incorporate one or two pi bonds, respectively. Metallic bonding involves the electrostatic attraction between positive metal ions and a 'sea' of mobile, delocalized electrons, which accounts for characteristic properties like high electrical conductivity. The chapter uses the concept of electronegativity—the ability of an atom in a bond to attract the electron pair—to predict bond character (Pauling values show that a large difference predicts ionic character) and bond polarity. Differences in electronegativity create polar covalent bonds with partial positive and negative charges; the overall molecular polarity depends on both bond polarity and the molecule’s geometry. Finally, the weak intermolecular forces (van der Waals forces) are categorized: instantaneous dipole-induced dipole forces (London dispersion forces, present in all substances and increasing with electron count/surface area), permanent dipole-permanent dipole forces, and the much stronger hydrogen bonding. Hydrogen bonding occurs when hydrogen is attached to highly electronegative atoms (Fluorine, Oxygen, or Nitrogen) and is responsible for the unique physical properties of water, including its anomalous density change upon freezing and its high boiling point.