Nitrous oxide (N2 O) has three possible Lewis structures:...

Nitrous oxide $\left(\mathrm{N}_{2} \mathrm{O}\right)$ has three possible Lewis structures: Given the following bond lengths, $$ \begin{array}{llll} \mathrm{N}-\mathrm{N} & 167 \mathrm{pm} & \mathrm{N}=\mathrm{O} & 115 \mathrm{pm} \\ \mathrm{N}=\mathrm{N} & 120 \mathrm{pm} & \mathrm{N}-\mathrm{O} & 147 \mathrm{pm} \\ \mathrm{N} \equiv \mathrm{N} & 110 \mathrm{pm} & & \end{array} $$ rationalize the observations that the $\mathrm{N}-\mathrm{N}$ bond length in $\mathrm{N}_{2} \mathrm{O}$ is $112 \mathrm{pm}$ and that the $\mathrm{N}-\mathrm{O}$ bond length is $119 \mathrm{pm}$. Assign formal charges to the resonance structures for $\mathrm{N}_{2} \mathrm{O}$. Can you eliminate any of the resonance structures on the basis of formal charges? Is this consistent with observation?

Nitrous oxide $\left(\mathrm{N}_{2} \mathrm{O}\right)$ has three possible Lewis structures:
Given the following bond lengths,
$$
\begin{array}{llll}
\mathrm{N}-\mathrm{N} & 167 \mathrm{pm} & \mathrm{N}=\mathrm{O} & 115 \mathrm{pm} \\
\mathrm{N}=\mathrm{N} & 120 \mathrm{pm} & \mathrm{N}-\mathrm{O} & 147 \mathrm{pm} \\
\mathrm{N} \equiv \mathrm{N} & 110 \mathrm{pm} & &
\end{array}
$$
rationalize the observations that the $\mathrm{N}-\mathrm{N}$ bond length in $\mathrm{N}_{2} \mathrm{O}$ is $112 \mathrm{pm}$ and that the $\mathrm{N}-\mathrm{O}$ bond length is $119 \mathrm{pm}$. Assign formal charges to the resonance structures for $\mathrm{N}_{2} \mathrm{O}$. Can you eliminate any of the resonance structures on the basis of formal charges? Is this consistent with observation?

Key Concepts

Resonance Structures

Resonance structures are alternative Lewis representations that illustrate the delocalization of electrons in a molecule. They show different ways to arrange electrons without altering the overall connectivity of the atoms. The actual molecule is best described as a hybrid of these contributing forms, which helps to explain intermediate bond lengths when bonds have characteristics of more than one bond order.

Formal Charge Assignment

Formal charge is a bookkeeping tool used in Lewis structures to indicate the electron distribution for each atom. It is calculated by comparing the number of valence electrons in an isolated atom with the number assigned in the molecule. Favorable resonance structures are usually those that minimize formal charges, avoid placing negative charges on less electronegative atoms, and reduce overall charge separation, thereby reflecting a more stable electronic arrangement.

Bond Order and Bond Length

Bond order refers to the number of chemical bonds between a pair of atoms, with higher bond orders generally correlating with shorter bond lengths and stronger bonds. In molecules where resonance contributes to the bonding, observed bond lengths are often intermediate between those expected for single, double, or triple bonds, which reflects partial bonding character due to electron delocalization.

Stability of Resonance Forms

The stability of individual resonance structures is assessed by how well they distribute formal charges and by avoiding unfavorable charge placements. Resonance forms that involve high or mismatched formal charges are less favorable and contribute less to the overall resonance hybrid. This principle helps rationalize why some drawings of a molecule are considered less likely and is consistent with experimental observations such as bond lengths.

Transcript

00:01 To answer this question, first let's draw the possible lewis structures for the resonance forms of n2o.

00:05 So we have two nitrogen, and each of them is going to have five valence electrons, plus oxygen, and we have six valence electrons, so that's 16 electrons available.

00:15 And we want each of these three atoms to have eight valence electrons to satisfy the octet rule.

00:20 So we want to see 24 electrons, which means that we need to share eight, the difference between the two.

00:27 And so that's going to involve four bonds.

00:44 So one way that we could have four bonds is like that.

00:47 Another way is two and two.

00:52 And another way is one and three like that.

00:57 Then let's add the loan pairs.

00:58 If there's three bonds, there should be one lone pair.

01:01 If there's one bond, there should be three loan pairs.

01:02 That way you get to eight.

01:03 In the middle, one already has four bonds, so we don't do anything to it.

01:06 All right, two bonds means two loan pairs.