Covalent Bonding

Covalent bonding forms when two adjacent atoms have the same affinity for their outer electrons - neither one can simply take an electron from the other, so the only way to reach the stable, filled-shell configuration of a noble gas is to share electrons between them. ^[1] The sharing works because atoms must be close enough for an orbital of one to physically overlap with an orbital of the other. Once that overlap exists, the electrons in those orbitals no longer belong to a single atom - they spend their time in the region between both nuclei, binding the pair together. The great hardness of diamond demonstrates just how strong this arrangement can be. ^[1]

The source of attraction in covalent bonding is electrostatic. The shared electrons are concentrated in the space between the bonded atoms, giving that region a net negative charge. This negative charge draws the positively charged atomic cores of both atoms toward each other, holding them close enough for the sharing to continue. ^[1] Unlike ionic bonding, where a transfer permanently separates charge into two distinct ions, covalent bonding keeps the electrons in dynamic shared residence between the bonded atoms. The distinction is one of degree rather than kind: bonding in most minerals is actually intermediate between the two extremes, with characteristics of both. ^[1]

Diamond: sp³ Hybridisation and Sigma Bonds

Diamond is the clearest geological example of pure covalent bonding, and understanding it requires following the orbital geometry step by step. Carbon’s ground-state electron configuration is 1s²2s²2p², with two electrons in the 2s orbital and two more split between two of the three 2p orbitals. To acquire a noble-gas configuration, carbon would need to either gain four electrons (reaching a neon-like configuration) or lose four (reaching a helium-like one). ^[1] Neither ionic path is available because all carbon atoms have identical electron configurations and equal affinity for electrons - one carbon atom cannot rob an electron from another. ^[1] Sharing four electrons is the only option, but sharing requires four unpaired orbitals, and the ground state does not provide them.

The solution is hybridisation. First, one of the two electrons in the 2s orbital is promoted to the vacant 2p orbital, leaving four singly occupied orbitals available for sharing. ^[1] Then, because all four bonds to adjacent carbon atoms must be identical, the four orbitals - one from the 2s subshell and three from the 2p subshell - hybridise into four equivalent sp³ orbitals, each consisting of a large lobe pointing in one direction and a small lobe pointing the opposite way. ^[1] These four large lobes arrange themselves so that they point to the four corners of a regular tetrahedron - 109.5° apart in every direction - making tetrahedral coordination the natural geometry of sp³-bonded carbon.

Each sp³ orbital on one carbon overlaps end-to-end with an equivalent sp³ orbital on each of four neighbouring carbon atoms. The resulting bonds are sigma bonds (σ) - bonds with high rotational symmetry about the axis connecting the two nuclei. ^[1] The continuous three-dimensional network of tetrahedral sigma bonds gives diamond its exceptional hardness: every carbon atom is locked in place by four strong bonds pointing to four neighbours, with no weak planes of easy separation.

Graphite: sp² Hybridisation, Sigma Bonds, and Pi Bonds

Graphite is also made entirely of carbon, but a different hybridisation scheme produces a radically different structure and set of properties. In graphite, only the 2px and 2py orbitals hybridise with the 2s orbital, forming three sp² hybrids, while the 2pz orbital remains unhybridised. ^[1] Each carbon atom therefore bonds to exactly three neighbours, not four, using its three sp² orbitals to form sigma bonds within continuous flat sheets. ^[1]

The unhybridised 2pz orbitals project at right angles to the plane of the carbon sheets. Rather than overlapping end-to-end, these orbitals overlap side-to-side with 2pz orbitals on all three neighbouring atoms, forming pi bonds (π). ^[1] The critical consequence is that the electrons in these pi bonds are not localised between any one pair of atoms - they form a continuous delocalised band above and below each carbon sheet, able to move from atom to atom with little resistance. This makes graphite an electrical conductor, because an applied voltage can drive those mobile electrons along the sheets. ^[1]

The combined sigma and pi bonding within each sheet is actually stronger than the bonding in diamond, as shown by the fact that C-C bonds in graphite are shorter than in diamond - shorter bond length means a stronger bond. ^[1] Graphite’s famously low hardness therefore has nothing to do with weak bonding within its sheets. It comes entirely from the bonding that holds adjacent sheets together - van der Waals bonds - which are much weaker than either the sigma or pi bonds within the sheets. ^[1] Graphite cleaves perfectly parallel to the sheets because breaking van der Waals bonds requires very little force, while the covalent network within each sheet remains intact.

Directionality and Its Structural Consequences

Because covalent bonds require specific orbital overlap, they are highly directional: atoms must be positioned at precise angles relative to each other for the orbitals to overlap. ^[1] This is the key geometric difference between covalent and ionic structures. In ionic bonding, the attractions are non-directional - any cation can point toward any anion regardless of angle - and the resulting structures can usually be described as close-packed arrangements of spheres. Covalent structures follow orbital geometry instead. The tetrahedral sp³ arrangement of diamond, for instance, cannot be described as any standard sphere-packing scheme: it is an open structure with large voids dictated by the 109.5° bond angles, not by the drive to minimise empty space. This directionality ultimately governs crystal morphology, density, and cleavage behaviour in covalently bonded minerals.

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