Covalent bonding

Covalent Bonding and Supporting Evidence

A covalent bond is the strong electrostatic attraction between two nuclei and the shared pair of electrons between them. This attraction forms because both positively charged nuclei are pulled towards the same region of concentrated negative charge, which is the shared electron pair.

Two main lines of evidence support the covalent bonding model: the physical properties of giant atomic structures and electron density maps of simple molecules.

Giant atomic structures such as diamond have extremely high melting points, showing that large amounts of energy are required to break the strong covalent bonds throughout the lattice. Their hardness and rigidity demonstrate that the directional shared-electron attraction holds atoms in fixed positions, providing structural evidence for shared-pair bonding.

Electron density maps reveal regions of high electron density located directly between adjacent nuclei in molecules such as Cl₂ or HCl. This concentration of electrons between the nuclei provides direct experimental proof that bonding electrons are shared, not fully transferred from one atom to another.

Dot-and-Cross Diagrams for Covalent Species

Dot-and-cross diagrams show how outer electrons are shared, using dots for one atom’s electrons and crosses for the other’s. Each shared pair represents one covalent bond, so a single bond shows one shared pair, a double bond shows two shared pairs, and a triple bond shows three shared pairs.

In O₂, two pairs of electrons are shared between the oxygen atoms, producing a double bond, while in N₂ three pairs are shared to form a triple bond.

A dative covalent bond, also called a coordinate bond, is a covalent bond in which both shared electrons originate from the same atom. Once formed, a dative bond is identical in strength and behaviour to any ordinary covalent bond, and is often represented by an arrow pointing from donor to acceptor.

In the ammonium ion (NH₄⁺), the lone pair on nitrogen donates into an empty orbital on H⁺, giving four equivalent N–H bonds in the resulting cation.

In Al₂Cl₆, two AlCl₃ units join because two chlorine atoms each donate a lone pair into an empty orbital on aluminium, forming two dative bonds across the dimer.

Examiner InsightFor dative bonds, the two electrons in the shared pair must both come from the same atom in the diagram. Splitting the pair as one dot and one cross is a marking error.

Exam TipAlways use an arrow from donor to acceptor when showing a dative bond.

Giant Carbon Lattices and Their Uses

Carbon forms three examinable giant covalent structures — diamond, graphite, and graphene — each with distinct properties arising from different bonding arrangements.

In diamond, every carbon atom is covalently bonded to four others in a tetrahedral arrangement, producing a rigid three-dimensional network. This network gives diamond a very high melting point and extreme hardness, so diamond is used in cutting tools, drill tips, and abrasive surfaces. Diamond does not conduct electricity because all four outer electrons on each carbon are localised in covalent bonds, leaving no free electrons.

In graphite, each carbon forms three covalent bonds within flat hexagonal layers, leaving one electron per atom delocalised between the layers. These delocalised electrons allow graphite to conduct electricity along its layers, so graphite is used in electrodes for electrolysis and in batteries. The layers in graphite are held together only by weak intermolecular forces, so layers slide over each other easily, allowing graphite to act as a lubricant and as the marking material in pencils.

Graphene is a single isolated layer of graphite, just one atom thick, retaining the same hexagonal arrangement and delocalised electrons. Graphene is exceptionally strong, transparent, and highly conductive, so it is used in flexible electronics, composite reinforcement materials, and high-speed transistors.

Electronegativity in a Covalent Bond

Electronegativity is the ability of an atom to attract the bonding pair of electrons in a covalent bond. It is measured on the Pauling scale, where fluorine has the highest value of 4.0 and is therefore the most electronegative element.

Electronegativity increases across a period because nuclear charge increases while the inner shielding stays roughly the same, so the nucleus attracts the bonding pair more strongly.

Electronegativity decreases down a group because atoms become larger and inner electron shielding increases, so the bonding pair is held less strongly by the nucleus.

When two atoms in a covalent bond differ in electronegativity, the bonding pair is pulled towards the more electronegative atom, producing a polar bond.

Bonding Continuum and the Origin of Polarity

Pure covalent and pure ionic bonding lie at the extremes of a bonding continuum, with most real bonds falling somewhere between the two.

When two bonded atoms have identical electronegativity, the bonding pair is shared equally and the bond is purely covalent, as in H₂ or Cl₂.

A small electronegativity difference produces a polar covalent bond, where the more electronegative atom carries a partial negative charge δ⁻ and the other carries δ⁺.

A larger electronegativity difference shifts more electron density towards the more electronegative atom, increasing the ionic character of the bond. If the electronegativity difference is large enough, the electrons are effectively transferred and the bond becomes ionic, as in NaCl where the difference is approximately 2.1.

The continuum view explains why some “ionic” compounds, such as AlCl₃, show significant covalent character because the small cation polarises the larger anion.

MisconceptionBonds are not strictly “ionic” or “covalent”. The same compound can show partial character of both, depending on the electronegativity gap between the bonded atoms.

Exam TipRefer to “ionic character” or “covalent character” rather than placing bonds in absolute categories.

Polar Bonds Versus Polar Molecules

A polar bond has a permanent dipole because the bonded atoms differ in electronegativity, but a polar molecule requires that these bond dipoles do not cancel overall. A molecule is polar only if it contains polar bonds AND its shape leaves the individual bond dipoles uncancelled.

In CO₂, both C=O bonds are polar, but the molecule is linear, so the two equal dipoles point in opposite directions and cancel. CO₂ is therefore a non-polar molecule despite containing polar bonds.

In H₂O, both O–H bonds are polar, and the molecule has a bent shape, so the dipoles do not cancel and a net dipole points towards the oxygen. Water is therefore a polar molecule.

To predict polarity, first identify polar bonds using electronegativity differences, then consider molecular shape and symmetry to decide whether the bond dipoles cancel.

A symmetrical arrangement of identical polar bonds, such as in CCl₄ (tetrahedral) or BF₃ (trigonal planar), gives a non-polar molecule because the dipoles cancel by symmetry. An asymmetric arrangement, or the presence of lone pairs that disrupt symmetry, leaves a net dipole and produces a polar molecule.