When two nonmetals bond, neither one is willing to fully give up an electron — both want to gain electrons to complete their octet. Instead of transferring electrons, they share them. A shared pair of electrons, attracted to both nuclei at once, is a covalent bond. Think of the simplest case, two hydrogen atoms. Each has one electron and wants two. As the atoms approach, each nucleus starts attracting the other atom's electron as well as its own. At the right distance the attraction of both nuclei for the shared pair balances the repulsion between the two nuclei, and the pair sits between them holding them together. Each hydrogen now "counts" both electrons, so each has a full first shell. Because the sharing happens between specific atoms, covalent bonding produces separate, countable molecules — , , — rather than an endless lattice of ions.
- Electrons are transferred, making ions
- Forms a lattice; the formula is a ratio (formula unit)
- High melting points; solids at room temperature
- Conducts when melted or dissolved
- Electrons are shared, making molecules
- Formula gives the actual atoms in one molecule
- Low melting and boiling points; many are gases or liquids
- Usually does not conduct (no ions), except acids dissolved in water
How many bonds does each atom form?
A nonmetal atom usually forms as many covalent bonds as the number of electrons it needs to reach 8 (or 2 for hydrogen). Carbon, in Group 14, has 4 valence electrons and needs 4 more, so it forms 4 bonds; oxygen has 6 and needs 2, so it forms 2 bonds and keeps 2 lone pairs. Knowing these usual patterns is the fastest way to check a finished structure: if a neutral carbon has only 3 bonds, or an oxygen has 3, something has gone wrong.
| Atom | Valence e⁻ | Usual bonds | Lone pairs | Example |
|---|---|---|---|---|
| H | 1 | 1 | 0 | , HCl |
| F, Cl, Br, I | 7 | 1 | 3 | HF, |
| O, S | 6 | 2 | 2 | , |
| N, P | 5 | 3 | 1 | , |
| C, Si | 4 | 4 | 0 | , |
Drawing Lewis structures: a step-by-step method
- Count total valence electrons: add up the group-number electrons for every atom in the formula. For a polyatomic ion, add 1 electron per negative charge, or subtract 1 per positive charge.
- Choose a central atom: usually the atom that can form the most bonds and is not hydrogen (H is always terminal, never central).
- Connect each outer atom to the central atom with a single bond (a line = 1 shared pair = 2 electrons).
- Subtract the electrons already used for bonds from the total, and place the rest as lone pairs on the outer atoms until each has a full octet (H needs only 2 total).
- If electrons remain after outer atoms have full octets, place them as lone pairs on the central atom.
- Check the central atom: if it doesn't have 8 electrons (a full octet), convert a lone pair on a neighbouring atom into an extra bond (double or triple) until it does.
Structural formulas and multiple bonds
A structural formula shows the same information as a Lewis structure but draws bonds as lines instead of dots between atoms: a single line = single bond, a double line = double bond, a triple line = triple bond. Lone pairs are often left off structural formulas for simplicity but are still understood to be there. Multiple bonds form when there are not enough electrons for every atom to reach an octet with single bonds alone, so neighbouring atoms share two or three pairs instead of one. More shared pairs pull the two nuclei closer together and hold them more tightly, so a triple bond is shorter and stronger than a double bond, which is shorter and stronger than a single bond. The triple bond in nitrogen gas, , is so strong that the nitrogen making up 78% of the air is very unreactive.
| Bond | Shared pairs | Electrons | Carbon–carbon bond length | Example |
|---|---|---|---|---|
| Single | 1 | 2 | 154 pm (longest, weakest) | Ethane, |
| Double | 2 | 4 | 134 pm | Ethene, ; |
| Triple | 3 | 6 | 120 pm (shortest, strongest) | Ethyne, ; |
The octet rule — and its exceptions
| Exception | Example | Why |
|---|---|---|
| Duet rule | H₂, HCl | Hydrogen (and helium) is stable with only 2 electrons — its first shell holds a maximum of 2 |
| Electron-deficient central atom | BF₃, BeCl₂ | Boron and beryllium can be stable with fewer than 8 electrons around them (6 and 4, respectively) |
| Expanded octet | PCl₅, SF₆ | Elements in period 3 or beyond are large enough to accommodate more than 8 electrons around the central atom |
| Odd-electron species | NO | An odd total number of valence electrons makes a full-octet structure impossible for every atom |