Covalent bonding
The idea
Nonmetals solve the octet problem differently: instead of transferring electrons, two nonmetal atoms share pairs of them, and each shared pair is a covalent bond. Both atoms count the shared electrons toward their own octet, so both can reach a noble-gas arrangement without anyone fully giving anything up. Sharing produces true molecules — discrete units like H₂O or CO₂ — which is why covalent substances are often gases, liquids, or soft solids with low melting points.
A Lewis structure is the bookkeeping tool: count all valence electrons in the molecule, connect atoms with single bonds in a sensible skeleton (the less electronegative atom usually goes in the center), spread the remaining electrons as lone pairs starting with the outer atoms, and if any atom still falls short of an octet, convert lone pairs into extra shared pairs to make double or triple bonds.
Two corrections worth making early: a double bond shares four electrons (two pairs), not two; and sharing is rarely equal. When the bonded atoms differ in electronegativity, the shared pair sits closer to the stronger puller, making the bond polar — covalent and ionic bonding are two ends of one sliding scale, not separate worlds.
Worked example
Draw the Lewis structure of carbon dioxide, CO₂: find the total valence electrons, place the bonds and lone pairs, and confirm every atom has an octet.
- Count the supply: carbon brings 4 valence electrons and each oxygen brings 6, so the total is 4 + 2 × 6 = 16 electrons to place.
- Build the skeleton with carbon in the center (it is less electronegative than oxygen) and a single bond to each oxygen. Two single bonds use 4 electrons, leaving 12.
- Distribute the remaining 12 as lone pairs on the outer atoms: 6 on each oxygen. Now each oxygen has an octet (2 bonding + 6 nonbonding), but carbon has only 4 electrons around it — 4 short.
- Fix carbon's deficit by sharing more: move one lone pair from each oxygen into the bonding region, turning both bonds into double bonds, written O=C=O.
- Audit the result: carbon now shares four pairs (8 electrons), each oxygen has two bonding pairs plus two lone pairs (8 electrons), and the total is still 16. Every atom has an octet, so the structure is complete.
Answer. CO₂ is O=C=O — two double bonds to the central carbon, with two lone pairs left on each oxygen.
Check your understanding
- What signals during the Lewis procedure tell you a molecule needs a double or triple bond?
- Why do molecular (covalent) substances usually melt at far lower temperatures than ionic ones?
- How does a difference in electronegativity between two bonded atoms change where the shared pair spends its time?
- Where does the octet guideline break down, and why should you treat it as a strong default rather than a law?
Build the foundations first
Covalent bonding builds on these concepts. If any feel shaky, start there.