Topic 2: Structure 2: Models of bonding and structure

Chemists use simplified models to explain how atoms join and why substances behave the way they do. This theme builds three core bonding models (ionic, covalent and metallic), then uses them to predict shape, polarity and the properties of solids and liquids.

The ionic model and lattices

Ionic bonding forms when a metal transfers one or more electrons to a non-metal, producing positive cations and negative anions. These oppositely charged ions are held together by strong electrostatic attraction acting in all directions, so they pack into a regular three-dimensional giant lattice rather than discrete molecules. The strength of the attraction grows with larger ionic charges and smaller ionic radii, which is why magnesium oxide (Mg2+ and O2-) has a far higher melting point than sodium chloride (Na+ and Cl-). Ionic compounds typically have high melting and boiling points, are brittle (a layer shift brings like charges together and the lattice splits), conduct electricity only when molten or dissolved (ions become mobile), and many are soluble in water because polar water molecules surround and stabilise the ions.

The covalent model and Lewis formulas

Covalent bonding is the sharing of one or more pairs of electrons between non-metal atoms, allowing each atom to reach a more stable electron arrangement. A single bond shares one pair, a double bond two pairs, and a triple bond three pairs; more shared pairs mean shorter, stronger bonds. Lewis (electron dot) formulas show every bonding pair and every lone pair, helping you account for all valence electrons. A useful tool is the dative (coordinate) covalent bond, where both shared electrons come from the same atom, as in the ammonium ion NH4+. Some species, such as boron in BF3 or sulfur in SF6, do not obey the simple octet rule, so the model has clear limits.

VSEPR and molecular shapes

Valence Shell Electron Pair Repulsion theory predicts geometry by assuming electron domains (bonding pairs and lone pairs) around a central atom repel and spread out as far apart as possible. Two domains give a linear shape (180 degrees), three give trigonal planar (120 degrees), and four give tetrahedral (109.5 degrees). Lone pairs repel more strongly than bonding pairs, so they compress bond angles: water has four domains but a bent shape at about 104.5 degrees, and ammonia is trigonal pyramidal at about 107 degrees. Always count domains first, then decide the shape based on how many of them are bonding versus lone pairs.

Bond and molecular polarity

A covalent bond is polar when the two atoms differ in electronegativity, pulling the shared electrons toward the more electronegative atom and creating partial charges (a dipole). Whether the whole molecule is polar depends on both bond polarity and shape: if individual bond dipoles cancel by symmetry the molecule is non-polar overall. Carbon dioxide has two polar C=O bonds but is linear and symmetric, so the dipoles cancel and the molecule is non-polar; water also has polar bonds but its bent shape means the dipoles add up, making it strongly polar. Tetrachloromethane CCl4 is symmetrical and non-polar, while chloromethane CH3Cl is polar.

Intermolecular forces and physical properties

The forces between molecules, not the covalent bonds within them, control melting points, boiling points and volatility of molecular substances. The weakest are London (dispersion) forces, present in all molecules and stronger for larger, more electron-rich molecules. Permanent dipole-dipole forces act between polar molecules. The strongest is hydrogen bonding, which occurs when hydrogen is bonded to nitrogen, oxygen or fluorine and is attracted to a lone pair on N, O or F in a neighbouring molecule. This explains why water boils far higher than expected, why alcohols are less volatile than alkanes of similar mass, and why boiling point generally rises down a group of similar molecules.

Giant covalent structures

Some covalent substances form continuous networks of atoms joined by strong covalent bonds throughout, called giant covalent or covalent network structures. Diamond has each carbon bonded tetrahedrally to four others, giving extreme hardness and a very high melting point but no electrical conductivity. Graphite has layered sheets of carbon each bonded to three others, with delocalised electrons between layers, so it conducts electricity and the layers can slide as a lubricant. Silicon dioxide (quartz) is a hard, high-melting network of Si-O bonds. Because so many strong bonds must break to melt them, all giant covalent solids have very high melting points and are usually insoluble.

The metallic model and properties

Metallic bonding is described as a lattice of positive metal ions surrounded by a sea of delocalised valence electrons that are free to move throughout the structure. The electrostatic attraction between the cations and this shared electron sea holds the metal together. The strength of metallic bonding increases with greater ionic charge (more delocalised electrons) and smaller ionic radius, so transition metals are generally stronger and higher melting than group 1 metals. Mobile electrons make metals excellent electrical and thermal conductors, the non-directional bonding allows layers of ions to slide so metals are malleable and ductile, and the dense packing gives most metals high melting points and lustre.

HL: The bonding continuum and triangle

In reality bonding is not strictly ionic, covalent or metallic but lies on a continuum between these extremes. A bonding triangle (van Arkel-Ketelaar triangle) places compounds according to the average electronegativity of the elements (along the base) and the difference in electronegativity (up the side). Large electronegativity differences indicate predominantly ionic character, small differences between non-metals indicate covalent bonding, and low average electronegativity points to metallic bonding. This model explains polar covalent and partially ionic intermediate cases, and reminds you that labels like ionic or covalent describe the dominant character rather than an absolute category.

Alloys and polymers

Alloys are mixtures of a metal with one or more other elements, where differently sized atoms disrupt the regular layers and make it harder for them to slide, so alloys such as bronze, brass and steel are usually stronger and harder than the pure metal while keeping metallic conductivity. Polymers are very large molecules built by joining many small repeating units called monomers; addition polymers form when monomers with double bonds (such as ethene) link together with no atoms lost, and condensation polymers form when monomers join with the loss of a small molecule like water. The properties of a polymer depend on chain length, branching and the intermolecular forces between chains.