Covalent bonding

Key Definition A covalent bond is the strong electrostatic attraction between two nuclei and a shared pair of electrons between them.

The model is supported by physical evidence from giant atomic (macromolecular) structures, which show high melting points and great hardness, and by electron density maps that reveal high electron density between adjacent nuclei.

Dot-and-cross diagrams represent the shared pairs in a covalent bond as combinations of dots and crosses to show which atom each electron originates from.

Single, double, and triple bonds correspond to one, two, and three shared pairs of electrons respectively.

Key Definition A dative (coordinate) covalent bond forms when both electrons in the shared pair come from the same atom, as seen in NH₄⁺ and Al₂Cl₆.

Carbon forms three giant covalent structures with very different properties.

Diamond

• Each carbon atom is covalently bonded to four others in a tetrahedral arrangement, forming a rigid three-dimensional lattice
• Extremely hard with a very high melting point, so it is used in cutting tools
• Does not conduct electricity, because all four outer electrons of each carbon are used in bonding (no delocalised electrons)

Graphite

• Has hexagonal layers in which each carbon atom is covalently bonded to three others
• One delocalised electron per carbon atom moves between the layers, allowing graphite to conduct electricity
• Weak forces between the layers let them slide over one another, so graphite acts as a lubricant

Graphene

• A single, one-atom-thick layer of graphite (one hexagonal sheet of carbon atoms)
• Very strong and an excellent electrical conductor, so it is used in advanced electronics and composites

Electronegativity

Key Definition Electronegativity is the ability of an atom to attract the bonding pair of electrons in a covalent bond.

It increases across a period and decreases down a group.

Differences in electronegativity create bond polarity, giving partial charges δ⁺ and δ⁻ on the bonded atoms.

Pure covalent and pure ionic bonding are the two extremes of a bonding continuum, and the size of the electronegativity difference determines where a particular bond falls on this continuum.

Polarising effects also explain the partial covalent character found in some “ionic” compounds.

A polar molecule must contain polar bonds and have an asymmetric arrangement so that the bond dipoles do not cancel; predicting polarity therefore requires both an electronegativity comparison and a shape analysis.

CO₂ (linear) and CCl₄ (tetrahedral) are non-polar despite containing polar bonds, because their symmetric shapes cause the bond dipoles to cancel.

H₂O and NH₃ are polar because their bent and pyramidal shapes leave a net dipole.

Exam Tip The exam consistently rewards specific language: “shared pair”, “delocalised electrons”, “electronegativity difference”, and “dipoles cancel”. Imprecise wording loses marks even when the underlying chemistry is correct.