Chemistry · General chemistry I · Concept
Lewis structures and formal charge
A Lewis structure shows how a molecule’s valence electrons are arranged as bonds and lone pairs. Count the valence electrons, join the atoms with single bonds, complete the outer atoms’ octets and put leftovers on the central atom; if it is still short, form double or triple bonds. Formal charge picks the best structure, and resonance describes molecules no single structure can.
Count the valence electrons
Add the valence electrons of every atom. For main-group elements other than helium, this is the last digit of the group number: 1 for hydrogen, 4 for carbon, 5 for nitrogen, 6 for oxygen, 7 for fluorine and chlorine. For an ion, add one electron for each negative charge and remove one for each positive charge.
Build the skeleton
The least electronegative atom, never hydrogen, usually goes in the center with the others around it. Hydrogen forms only one bond. Join each outer atom to the center with a single bond, which uses two electrons.
Complete the octets
Place the remaining electrons as lone pairs on the outer atoms until each has an octet; hydrogen needs only two electrons. Put any electrons left over on the central atom.
Use multiple bonds when the center is short
If the central atom still has fewer than eight electrons, turn a lone pair on a neighboring atom into a shared pair, making a double bond; repeat for a triple bond. The total number of electrons does not change.
Formal charge
Formal charge compares the electrons an atom owns in the structure with its valence electrons: V valence electrons, minus N nonbonding electrons, minus half of the B electrons in its bonds. The formal charges add up to the overall charge. The best structure keeps formal charges closest to zero and puts any negative formal charge on the more electronegative atom.
Resonance
When two or more structures differ only in where the electrons sit, with the atoms in the same places, no single one describes the molecule. The real structure is a blend of them, the resonance hybrid. In the nitrate ion all three nitrogen–oxygen bonds are identical, between a single and a double bond; the ion does not flip between forms. Resonance forms describe one compound: the same atoms joined the same way. If a hydrogen has moved to another atom, the structures are tautomers, which are different compounds.
Exceptions to the octet rule
Hydrogen holds two electrons. Boron is often drawn with six, as in BF₃, and beryllium with four, as in BeH₂. Central atoms from period 3 down can be drawn with more than eight, as in SF₆, but second-period atoms such as carbon, nitrogen and oxygen never are. Sulfate can be drawn with octets and formal charges or with an expanded octet on sulfur; both are accepted, and many texts expand the octet to lower the formal charges. A molecule with an odd number of electrons, such as NO, cannot give every atom an octet.
Common mistakes
- Counting lone pairs instead of electrons: each lone pair is two electrons.
- Forgetting an ion’s charge when counting electrons.
- Giving hydrogen more than one bond.
- Giving a second-period atom such as nitrogen or oxygen more than eight electrons.
- Adding electrons to complete octets: a double bond moves a lone pair into a bond, and the total stays fixed.
- Treating resonance structures as a molecule switching between forms.
Key terms
- Lewis structure
- A drawing of a molecule or ion that shows bonds as lines and lone pairs as dots. Its electrons add up to the atoms’ valence electrons, plus one for each negative charge or minus one for each positive charge; an ion is drawn in brackets with its charge. It shows how the valence electrons are shared, but not the 3D shape, which VSEPR predicts.
- Valence electron
- An electron in an atom’s outermost shell, the electrons that take part in bonding. A main-group element has as many as its A-group number, such as 6 for oxygen (group 6A).
- Lone pair
- A pair of valence electrons that belongs to one atom instead of a bond, drawn as two dots. Lone pairs take up space and push bonds closer together, changing a molecule’s shape.
- Octet rule
- Main-group atoms tend to bond until they have eight valence electrons around them; hydrogen needs only two. The exceptions: molecules with an odd number of electrons, such as NO; boron and beryllium with fewer than eight, as in BF₃; and atoms from period 3 down with 10 or 12, as in SF₆. Second-period atoms never go past eight.
- Formal charge
- The charge an atom would have if bonding electrons were shared equally: valence electrons − lone-pair electrons − half the bonding electrons. The formal charges add up to the overall charge. The best Lewis structure keeps them near zero and puts any negative charge on the most electronegative atom.
- Resonance
- Two or more valid Lewis structures for one molecule that differ only in where the electrons are, not the atoms. The real molecule is a blend (a hybrid) of them; it does not flip between them.
Work through an example
Draw the Lewis structure of formaldehyde, CH₂O, and check the formal charges.
Draw the Lewis structure of CH₂O →Sources and scope
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