Why atoms share electrons
Atoms form bonds to achieve a lower energy state. When two non-metal atoms approach each other, their valence electrons can interact. Instead of transferring electrons (ionic bonding), non-metals with similar electronegativities share electrons — each atom effectively "sees" a full outer shell.
The shared pair is attracted to both nuclei simultaneously, pulling them together. The resulting bond length and bond energy reflect the balance between nuclear repulsion (too close) and electron sharing (too far apart).
For the full context of how covalent bonding compares to ionic and metallic bonding, see our chemical bonding guide. For the underlying electron shells that determine bonding capacity, see our atomic structure guide.
Single, double and triple bonds
- Single bond (σ bond) — one shared electron pair. Example: H–H (bond energy 436 kJ/mol, bond length 74 pm). Free rotation around the bond axis is possible.
- Double bond (σ + π bond) — two shared electron pairs. Example: C=C in ethene (bond energy 614 kJ/mol, bond length 134 pm). Restricts rotation, causing cis/trans isomerism.
- Triple bond (σ + 2π bonds) — three shared electron pairs. Example: N≡N (bond energy 945 kJ/mol, bond length 110 pm). Makes nitrogen gas extremely stable and unreactive — which is why nitrogen fixation (see our quantum chemistry article) is so energetically demanding.
Common covalent molecules
Covalent bonds in biology
Life is built almost entirely from covalent bonds. The C–C and C–H bonds in organic molecules form the backbone of lipids, sugars, proteins and DNA. The C=O bond in glucose is broken during cellular respiration to release energy — see our cellular respiration guide. Peptide bonds (C–N, with partial double-bond character) link amino acids in proteins.
The phosphodiester bonds linking nucleotides in DNA are covalent — they provide the stable backbone of the double helix. The bases are held together by the weaker hydrogen bonds, which can be disrupted during DNA replication and transcription.