Write Lewis structures, including typical contributions to the resonance structure (where appropriate, allow for the possibility of octet expansion, including double bonds in different positions), for (a) sulfite ion; (b) hydrogen sulfite ion; (c) perchlorate ion; (d) nitrite ion.

Understanding resonance structures is crucial for predicting the behavior of molecules and ions in chemistry. Resonance occurs when a molecule has multiple valid Lewis structures that can depict its electron arrangement. These various structures, known as resonance structures or resonance contributors, are not distinct entities but rather hypothetical extremes that blend to give an actual, intermediate structure.

Let's imagine it as a molecule having alternate 'faces', much like a person having different profile pictures. Just as a person is not captured completely by a single photograph, a molecule's true electronic structure isn't perfectly described by just one Lewis structure.

When dealing with resonance, one must follow specific rules: First, only the positions of the electrons can differ between the structures; the positions of the nuclei must remain constant. Also, the structures must have the same number of unpaired electrons and the same net charge. Lastly, the actual distribution of electrons in the molecule is a hybrid of those depicted by the resonance structures, and this hybrid averages the properties of the resonance forms.

In drawing resonance structures for ions like sulfite, hydrogen sulfite, perchlorate, and nitrite, as discussed in the exercise, it's not just about drawing each possible structure but also understanding that the actual ion will manifest characteristics of all the possible drawings.

Valence electrons are the outermost electrons of an atom and play a primary role in chemical reactions and bonding. Think of them as the 'social butterflies' of an atom, interacting and connecting with other atoms to form compounds.

To determine the number of valence electrons, simply look at the group number of the elements in the periodic table. For example, oxygen is in group 16 and has 6 valence electrons. Calculating the total number of valence electrons in the molecule or ion, as seen in step 2 of the solution, is the foundation for drawing accurate Lewis structures.

Adding electrons for negative charges or subtracting for positive charges when dealing with ions reflects the changing social circle of our atom: gaining or losing friends. This balance of electrons is essential to achieving the correct configuration that allows for the most stable interactions between atoms, leading to the formation of structures like those of the ions mentioned in the exercise.

The concept of formal charge is a bookkeeping tool that helps chemists understand the distribution of electrons in molecules and ions. It's like assigning responsibility to each atom for its electrons. The formal charge is calculated using the formula:

\f\( FC = V - (N + \frac{B}{2}) \f\)

where \f\( FC \f\) is the formal charge, \f\( V \f\) is the number of valence electrons in the isolated atom, \f\( N \f\) is the number of nonbonding electrons, and \f\( B \f\) is the number of electrons in covalent bonds to the atom. Observing changes in formal charge can be like a detective investigating movements in a bank account: it can lead to insights about where the electrons prefer to reside in the molecule or ion.

During the exercise's Step 5, calculating the formal charge for each atom in each possible resonance structure allows us to judge which structure (or which blend of structures, in reality) is most stable. The structures with the lowest formal charges are typically favored because they reflect a more balanced distribution of electron responsibility among the atoms, akin to an even spread of tasks in a well-functioning team.