1. Count valence electrons. 2. Draw the basic structure with single bonds. 3. Adjust for octets by converting single bonds to double bonds where needed. 4. Draw all possible resonance structures.

Bond Order = (Total number of bonds) / (Number of bond locations)

1. Draw all possible Lewis structures. 2. Calculate the formal charge for each atom in each structure. 3. Choose the structure with the lowest formal charges, with negative charges on the most electronegative atoms.

1. Draw N connected to three O atoms with single bonds. 2. Add lone pairs to complete octets. 3. Convert one single bond to a double bond, and draw all three possible arrangements with a double-headed arrow between them.

Calculate formal charges, and then reduce formal charges on atoms by forming double or triple bonds if possible. The most stable structure usually has the lowest formal charges, with negative charges on the most electronegative atoms.

Resonance leads to bond lengths that are intermediate between single and double bonds, and bond strength that is greater than a single bond but less than a double bond.

A large formal charge generally decreases the stability of a Lewis structure.

You may lose points, as it indicates an incomplete understanding of resonance.

Minimizing formal charges generally increases the stability of a Lewis structure.

The Lewis structure becomes less stable.

Resonance Structures: Represent different ways of drawing the same molecule, differ only in electron arrangement. | Isomers: Are different molecules with the same molecular formula, but different arrangements of atoms.

Formal Charge: Assumes equal sharing of electrons in a bond. | Oxidation Number: Assumes the more electronegative atom gets all the shared electrons.