We said previously that the Lewis model is a localized electron model, which means valence electrons are assumed to be either shared between two atoms or as a lone pair on a single atom. Not all molecules fit this model well. For some molecules, electrons are delocalized, which means they may be found throughout multiple areas of the molecule. For these cases, we can use resonance to represent the delocalization of electrons. An example is ozone, which is O3: ozone has two resonance structures. Resonance structures are Lewis structures that differ only by the distribution of electrons. In these two resonance structures, we have one single bond and one double bond. They differ by the location of the single and double bonds as well as the location of one lone pair of electrons. However, the actual ozone molecule does not have one single bond and one double bond. Rather, it has two identical bonds that are somewhere between a single and a double bond. The actual ozone molecule exists as a weighted average of the resonance structures. To represent the actual molecule more clearly, we can use a resonance hybrid structure. A resonance hybrid includes only bonds or lone pairs that show up in all resonance structures. Then, for any bonds that do not show up in all residence structures, we added a dashed line to show the delocalization of those electrons. Although it's not shown here, we can also represent any non-zero formal charges that show up in one resonance structure but not all using partial charges. We'll discuss formal charges next. The formal charge is the charge an atom would have within a Lewis structure if lone pairs of electrons are counted as being only on the atom in question and the electrons within bonds are split evenly between the two atoms. To calculate the formal charge on a particular atom, we take the number of valence electrons in the neutral atom, subtract the number of lone pairs of electrons on that atom, and the number of bonds with that atom. Formal charge can help us evaluate different resonance structures to determine if one is better than another. When assigning formal charges and comparing resonance structures, we can use the following four guidelines: 1. First, the sum of all formal charges in a neutral molecule is zero. 2. Second, the sum of all formal charges in a polyatomic ion is equal to the charge on the ion. 3. Third, smaller numbers for formal charges are better, with a formal charge of zero being the best, if possible. 4. And fourth, if you must have one or more non-zero formal charges in a Lewis structure, it's better to put a more negative formal charge on the more electronegative atom.