Two Lewis structures are shown below for each species. Determine the formal charge on each atom and then, if appropriate, identify the Lewis structure of lower energy for each species. (a) \(\quad \ddot{O}-\ddot{S}=\ddot{O} \quad \ddot{O}=\ddot{S}=\ddot{O}\) (b) O=S(=O)([O-])S(=O)(=O)[O-]

The concept of formal charge is crucial for understanding the stability of molecules and predicting their behavior. A formal charge is an assumption of charge distribution whereby electrons in covalent bonds are equally shared between the bonded atoms.

To calculate an atom's formal charge, the formula used is:
\( \text{Formal charge} = (\text{Valence electrons}) - (\frac{1}{2} \times \text{Bonding electrons}) - (\text{Lone pair electrons}) \)
Valence electrons are the electrons in the outer shell of an atom that can participate in the formation of chemical bonds. The number of valence electrons typically corresponds to an element's group number on the periodic table.

For example, oxygen (group 16) usually has 6 valence electrons, while sulfur (group 16 as well) also has 6. When calculating the formal charge, bonding electrons (those involved in bonds) are counted once per bond and divided by two since the assumption is they're equidistant from the bonded atoms. Lone pair electrons are those that are localized on one particular atom and not shared.

Applying the Formal Charge Formula

When we applied this formula to species (a), we found out that all atoms had a formal charge of zero, which indicates a stable electronic structure. In other cases, the best Lewis structure will minimize the formal charge on atoms and place any negative charges on the more electronegative atoms.

Bonding electrons are at the heart of covalent bonds, where two atoms share a pair of electrons to achieve a more stable electron configuration. In Lewis structures, bonding electrons are depicted as lines connecting atoms. Each line represents a shared pair of electrons or a 'bond'.

These electrons can be counted in two ways when considering formal charge calculations and Lewis structure representations: as part of the total electron count for each individual atom, or simply as a bond connecting two atoms. When we calculate the formal charge, they are divided by two to represent that they are shared between two atoms.

Understanding Bonding in Lewis Structures

For instance, in species (a)'s Lewis structure, sulfur is bonded to two oxygen atoms. Each bond corresponds to two bonding electrons. Counting all of the bonds associated with sulfur (three in the first structure and two in the second), we divided by two to accurately assign the bonding electrons in calculating formal charge.

Lone pair electrons are valence electrons that are not shared between atoms and are not used in bonding. These electrons belong to a single atom and are often represented as a pair of dots in Lewis structures. They play a significant role in determining the shape of a molecule and can influence the molecule's chemical reactivity.

Lone pairs are critical when calculating formal charges because they are counted in their entirety into the formal charge formula, unlike bonding electrons, which are split between the atoms involved.

Impact of Lone Pairs on Molecule Stability

The presence or absence of lone pairs can also affect the stability of an atom within a molecule. For example, in species (a), the oxygen atoms have lone pairs, which are taken into account when determining the formal charge and thus the overall stability of the molecule. In species (b), the extra lone pairs on the negatively charged oxygens (O-) contribute to the molecule's negative charge and its chemical behavior.