Draw three resonance structures for each of the following: (a) \({ }^{-} \mathrm{CH}_{2} \mathrm{NO}_{2}\) (b) \(\mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{CH}_{3}\) (c) \({ }^{-} \mathrm{CH}_{2} \mathrm{CO}_{2}{ }^{-}\) (d) \(\mathrm{HOSO}_{2} \mathrm{O}^{-}\)

Lewis structures are a valuable tool in understanding the bonding in molecules. They represent atoms using their chemical symbols and show bonding and non-bonding electrons as lines and dots, respectively. When drawing a Lewis structure, it's crucial to follow the octet rule, which states that atoms tend to form bonds until they are surrounded by eight electrons.
Begin by counting the total number of valence electrons in the molecule. For instance, in \(\text{CH}_2\text{NO}_2^-\), combine the valence electrons from carbon (4), hydrogen (1 each), nitrogen (5), and oxygen (6 each), plus the additional electron due to the negative charge.
Place the atoms with the least electronegativity (usually central atoms like carbon or nitrogen) in the center, and arrange other atoms around them. Distribute the electrons to satisfy the octet rule for each atom. This includes sharing electrons to form bonds or placing lone pairs where necessary. Using this approach lays a clear foundation for identifying resonance structures.

Electron delocalization occurs when electrons are not confined to a single bond or atom but are instead spread out across multiple atoms or bonds. This is a key feature of resonance structures. In resonance structures, the arrangement of atoms remains the same, but the placement of electrons can differ. For example, in the molecule \(\text{CH}_2\text{NO}_2^-\), the negative charge can be shifted between the oxygen atoms while nitrogen remains bonded to both.
This delocalization contributes to the stability of the molecule, as electrons are spread out more evenly, reducing electron-electron repulsion. Visualizing delocalization requires drawing multiple resonance structures and using double-headed arrows to indicate that the actual structure is a hybrid of these possibilities. This hybrid structure represents the more accurate depiction of the electron distribution in the molecule.
Understanding delocalization helps explain why some molecules react the way they do and their properties, like stability and reactivity.

Formal charges are used to determine the most stable resonance structure of a molecule. They represent the difference between the valence electrons in an isolated atom and the electrons assigned to it in a molecule. To calculate the formal charge of an atom in a Lewis structure, use the formula:
\text{Formal charge} = (\text{Valence electrons}) - (\text{Non-bonding electrons}) - (\text{Bonding electrons}/2) \
Minimizing the formal charge across atoms leads to the most stable structure. Typically, resonance structures where atoms have a formal charge as close to zero as possible are more stable. For example, in \(\text{CH}_2\text{NO}_2^-\), you should aim to place the negative charge on the more electronegative atoms (oxygen, in this case) whenever possible.
Understanding formal charges also helps in predicting the behavior of molecules during chemical reactions, as areas of high formal charge can be sites of reactivity.