The formula
- Step rule for drawing the structure
1 count valence electrons, 2 connect atoms with single bonds, 3 complete octets on outer atoms, 4 put leftovers on the central atom, 5 form multiple bonds if the central atom is short- Formal charge on an atom
FC = (valence e⁻) - (lone pair e⁻) - ½(bonding e⁻)- Electron count for an ion
total e⁻ = Σ(valence e⁻) - charge- Octet rule
8 electrons around each main group atom, 2 around hydrogen
What the symbols mean
| Symbol | Meaning | Unit |
|---|---|---|
FC | Formal charge on one atom in a given structure | elementary charge (an integer) |
valence e⁻ | Valence electrons the free atom has, read off the main group number | count |
lone pair e⁻ | Nonbonding electrons drawn on that atom | count |
bonding e⁻ | Electrons in bonds to that atom, split evenly between the two atoms | count |
charge | Charge on the ion, added for an anion and subtracted for a cation | elementary charge |
When it applies
- Molecules and polyatomic ions built from main group elements, where the octet rule is a reliable target.
- The least electronegative atom goes in the center, except hydrogen, which is always terminal because it takes only two electrons.
- Where more than one arrangement satisfies the octet rule, the structure with formal charges closest to zero is the preferred one.
- Odd-electron species, electron-deficient atoms such as boron, and third-row atoms with expanded octets are the documented exceptions.
Worked example
Problem. Draw the Lewis structure of carbon dioxide, CO₂, and check it with formal charges.
- Count valence electrons: carbon contributes 4 and each oxygen 6, so 4 + 2(6) = 16 electrons to place.
- Carbon is the less electronegative atom, so the skeleton is O-C-O. Two single bonds use 4 electrons and leave 12.
- Complete the outer atoms first: three lone pairs on each oxygen uses all 12. Carbon now has only 4 electrons around it, well short of an octet.
- Move one lone pair from each oxygen into a second bond, giving O=C=O. Carbon now has four bonds, which is 8 electrons, and each oxygen keeps two lone pairs.
- Check the total: 8 electrons in the two double bonds plus 8 in the four remaining lone pairs is 16, matching step 1.
- Check formal charges: carbon is 4 - 0 - 8/2 = 0, each oxygen is 6 - 4 - 4/2 = 0. All zero, so this is the preferred structure.
Answer. O=C=O, using all 16 valence electrons, with two lone pairs on each oxygen and a formal charge of zero on every atom.
Common mistakes
- Counting atoms instead of valence electrons at step 1, or forgetting to add electrons for a negative charge and subtract them for a positive one.
- Giving hydrogen an octet. Hydrogen is complete with a single bond and two electrons, and it is never the central atom.
- Leaving the central atom short rather than forming a double or triple bond from a lone pair on an outer atom.
- Drawing a structure whose electron count does not match the total from step 1, which is the quickest check available and the one most often skipped.
Related formulas
- Electron configuration:
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p - Molecular formula:
n = molar mass / empirical formula mass - Covalent bond:
ΔEN = |EN(A) - EN(B)|