Write the Lewis structure, 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 the role of valence electrons is fundamental when analyzing chemical bonding and the creation of molecules. Valence electrons are the outermost electrons of an atom and are responsible for the chemical properties of the element. In the context of Lewis structures, these electrons are the key players in forming bonds.

Each atom has a specific number of valence electrons based on its position in the periodic table. For instance, atoms in Group 1 have 1 valence electron, while atoms in Group 17 have 7. When writing Lewis structures, the total count of valence electrons sets the stage for how atoms will bond.

Counting Valence Electrons

In Step 1 of the textbook solution, the total number of valence electrons for each ion is determined. This is a critical first step as it provides a count that will be used to adhere to the octet rule and create the correct structure, including the resonance structures if appropriate.

The octet rule is a principle that atoms tend to combine in such a way that they each have eight electrons in their valence shell, achieving a noble gas configuration. This rule is a guide for the formation of stable molecular structures.

Following the octet rule underscores Step 4 and Step 6 in the solution process. Firstly, electrons are distributed around outer atoms to complete their octet. Hydrogen is an exception, content with two electrons. For central atoms, achieving an octet can require shuffling of lone pairs from the outer atoms to form double or triple bonds, ensuring maximum stability.

Exceptions to the Octet Rule

While the octet rule applies to many situations, there are exceptions. Elements in the third period and beyond can accommodate more than eight electrons—this is known as the octet expansion, as seen in the case of the central atom in sulfite and perchlorate ions.

Resonance structures are alternate Lewis structures for the same molecule or ion, showing different possible arrangements of electrons that still maintain the overall distribution of charges. These structures highlight the delocalization of electrons within a molecule.

Step 7 of the exercise deals with resonance structures by examining if the double or triple bonds can be relocated among the same atoms while preserving the octet rule. This is important because while each individual resonance structure may not accurately represent the molecule's true form, the actual structure is a hybrid of all possible resonance forms.

Visualizing Resonance

To visualize this, think of resonance structures like different dance steps to the same tune—they offer a different look, but the overall dance remains the same. Drawing all valid resonance structures and connecting them with double-headed arrows demonstrate this electron fluidity, giving a fuller picture of the molecule's bonding characteristics.