08 Lewis Structures and Resonance

Learn how to build Lewis structures, check electron counts and formal charges, recognize octet-rule exceptions, and interpret resonance as electron delocalization.

Building Lewis Structures

A tracks in a molecule or polyatomic ion. A line represents a shared electron pair, or covalent bond; dots represent lone pairs or unpaired electrons. Treat the diagram as a useful electron-counting model rather than a literal map of electron positions.

A reliable drawing sequence

Count the available . Add the atoms’ , then adjust for charge: add one electron per negative charge or subtract one per positive charge.
Choose a skeleton. Hydrogen is terminal; usually, the least electronegative atom other than hydrogen is central.
Connect the atoms with single bonds. Each bond accounts for two electrons.
Complete terminal-atom shells: hydrogen has a duet, while most other terminal atoms are given an octet using lone pairs.
Put remaining electrons on the central atom. If it lacks an octet, consider converting lone pairs on adjacent atoms into multiple bonds.
Check the electron total, shells, and formal charges. Put an ion in brackets and show its overall charge.

Example: carbon dioxide

For $\mathrm{CO_2}$ , count $16$ : $4$ from carbon and $6$ from each oxygen. The single-bond skeleton $\mathrm{O-C-O}$ uses $4$ electrons. Giving both oxygens octets uses another $12$ , leaving carbon short of an octet. Convert one lone pair from each oxygen into a bonding pair to obtain $\mathrm{O=C=O}$ . Each oxygen then has two lone pairs, every atom has an octet, and each is zero.

Takeaway: Count first, distribute electrons systematically, and check the completed structure rather than assuming the first skeleton is sufficient.

Checking Formal Charges

A helps compare plausible Lewis structures by assigning each atom its share of the electrons. In the bookkeeping model, bonding electrons are divided equally between the two atoms in a bond.

\text{Formal charge} = \text{valence electrons} - \text{nonbonding electrons} - \frac{1}{2}(\text{bonding electrons})

A convenient equivalent form is:

\text{Formal charge} = \text{valence electrons} - \text{lone-pair electrons} - \text{number of bonds}

In $\mathrm{CO_2}$ , each oxygen has $6$ , $4$ nonbonding electrons, and $2$ bonds, giving $6-4-2=0$ . Carbon has $4$ , no nonbonding electrons, and $4$ bonds, giving $4-0-4=0$ . The sum of all formal charges must equal the molecule’s or ion’s overall charge.

When choosing between plausible structures, favor those that satisfy octets for second-period atoms, keep formal-charge magnitudes small, and place negative on more electronegative atoms when other factors are similar. is a model value, not a measured atomic charge.

Takeaway: Use formal charges to check electron bookkeeping and compare structures; they do not describe measured charges on individual atoms.

Octet-Rule Exceptions

The is a useful guide, especially for second-period elements such as carbon, nitrogen, oxygen, and fluorine, but it does not apply to every .

Duet: Hydrogen is stable with two electrons in its valence shell.
Incomplete octet: Some central atoms are commonly represented with fewer than eight electrons. Boron in $\mathrm{BF_3}$ has six shared electrons, and beryllium in $\mathrm{BeH_2}$ has four.
Odd-electron species: An odd total number of means not every electron can be paired. Nitric oxide, $\mathrm{NO}$ , has $11$ and one unpaired electron.
More than eight electrons: Familiar Lewis representations of $\mathrm{PCl_5}$ and $\mathrm{SF_6}$ place more than eight electrons around the central atom.

These are reasons to follow the electron count and chemically plausible bonding rather than adding or removing electrons just to force every atom into an octet.

Takeaway: Apply the as a guideline, then let the electron count and plausible bonding determine whether an exception is needed.

Resonance and Delocalized Electrons

When one cannot represent an electron distribution adequately, show alternative placements of electrons for the same arrangement of atoms. Connect contributors with the resonance arrow $\leftrightarrow$ . Keep atom positions fixed; move only electrons, typically lone pairs and $\pi$ bonds. Every contributor must have the same total number of electrons and the same overall charge.

Example: nitrite ion

Nitrite, $\mathrm{NO_2^-}$ , has $18$ : $5$ from nitrogen, $12$ from oxygen, and $1$ for the negative charge. With nitrogen central, its two equivalent can be represented as:

\left[\mathrm{O=N-O}\right]^- \leftrightarrow \left[\mathrm{^-O-N=O}\right]

In each form, nitrogen has one lone pair. The double-bonded oxygen has two lone pairs, and the single-bonded oxygen has three. The single-bonded oxygen has $-1$ ; nitrogen and the double-bonded oxygen have $0$ . The charges add to the required overall charge of $-1$ .

The actual species is a , not a molecule rapidly switching between the two drawings. Its electrons are delocalized. Because the two contributors are equivalent, the two nitrogen–oxygen bonds are equivalent in the hybrid, with bonding intermediate between a single and a double bond.

Takeaway: Resonance changes electron placement, not atom positions; the hybrid describes the delocalized species represented by the contributors.