Question

Write a Lewis structure for each of the following simple molecules. Show all bonding valence electron pairs as lines and all nonbonding valence electron pairs as dots. For those molecules that exhibit resonance, draw the various possible resonance forms. a. \(\mathrm{SO}_{2}\) b. \(\mathrm{N}_{2} \mathrm{O}(\mathrm{N}\) in center) c. \(O_{3}\)

Short answer

Lewis structures for the given molecules are as follows: a. \(\mathrm{SO}_{2}\): Resonance structures: O=S=O <-> O= S-O b. \(\mathrm{N}_{2} \mathrm{O}\) (N in the center): N-N=O (No resonance structures) c. \(O_{3}\): Resonance structures: O=O-O <-> O-O=O

Step-by-step solution

a. Draw the Lewis structure for \(\mathrm{SO}_{2}\)

1. Count the total number of valence electrons: S has 6 valence electrons and O has 6 valence electrons, so the total count is: `6 + 2 * 6 = 18`. 2. Arrange the atoms: O=S=O and form single bonds between them. There are now 12 valence electrons left to distribute. 3. Assign the remaining 12 valence electrons to the oxygen atoms, not forgetting about the octet rule. 4. To complete the octet for sulfur, form a double bond with one of the oxygen atoms. Now each atom has a complete octet. 5. There is a resonance structure between the two possible double bond positions: O=S=O <-> O= S-O

b. Draw the Lewis structure for \(\mathrm{N}_{2} \mathrm{O}\) with N in the center

1. Count the valence electrons: N has 5 valence electrons and O has 6 valence electrons, so the total count is: `2 * 5 + 6 = 16`. 2. Arrange the atoms: N-N=O and form single bonds between them. There are now 10 valence electrons left to distribute. 3. Assign the remaining 10 valence electrons, not forgetting about the octet rule. The first N atom has a complete octet and the last O atom has a complete octet. There is one lone pair (2 valence electrons) left, which is added to the central N atom. 4. To complete the octet for the central N atom, form a double bond with one of the N atoms. Now each atom has a complete octet. 5. There is only one structure, so no resonance structures.

c. Draw the Lewis structure for \(O_{3}\)

1. Count the valence electrons: O has 6 valence electrons, so the total count is: `3 * 6 = 18`. 2. Arrange the atoms: O-O-O and form single bonds between them. There are now 12 valence electrons left to distribute. 3. Assign the remaining 12 valence electrons, not forgetting about the octet rule. 4. To complete the octet for the central O atom, form a double bond with one of the O atoms. Now each atom has a complete octet. 5. There is a resonance structure between the two possible double bond positions: O=O-O <-> O-O=O

Key concepts

Valence Electrons

Imagine valence electrons as the outermost soldiers of an atom's army, ready to bond with neighboring atoms. These electrons are significant because they determine an atom's bonding behavior. When we draw Lewis structures, which are simplistic maps of molecules showing the bonds and the lone pairs of electrons, valence electrons are the central characters.

In the Lewis structure exercise for \(\mathrm{SO}_{2}\), sulfur (S) and oxygen (O) came with 6 valence electrons each. This is because S is in group 16 of the periodic table, and likewise for O. Counting them is the first step to map out who's bringing what to the molecular party. It’s akin to knowing what ingredients you have before baking a cake. This count decides how the atoms bond and helps ensure each atom achieves a stable electron configuration, adhering to the octet rule.

Resonance Structures

A molecule sometimes can't be pinned down to just one layout. It’s like a seating chart for a dinner party that works in many different ways. This is the concept behind resonance structures. They occur when you can draw multiple valid Lewis structures for a molecule, differing only in the position of the electrons, not the position of the atoms.

In the exercise with ozone (\(O_{3}\)), we saw this play out. There were two structures, each having a different double bond location, hence creating a resonance. By envisioning the electrons as social butterflies fluttering between these two structures, we can grasp the concept of molecular resonance. This doesn't mean the molecule is constantly flipping between these forms, but rather exists as a hybrid, benefiting from the stability of both possibilities.

The Octet Rule

The octet rule is like the golden rule of chemistry—it tells us that atoms are generally most stable when they hold eight electrons in their valence shell. It’s the chemistry equivalent of finding balance in life. When drawing Lewis structures, we aim to satisfy this rule for each atom (with a few exceptions like hydrogen).

For instance, in the \(\mathrm{N}_{2} \mathrm{O}\) molecule from the exercise, we had to make sure each nitrogen (N) and oxygen (O) ended up with full outer shells of eight electrons. If an atom doesn't have enough electrons to fulfill the octet rule after forming single bonds, we can pull a switcheroo and form double or triple bonds as necessary. This is exactly what was done in the steps to ensure everybody's happy with their electron arrangement.