Question

Place the species below in order of the shortest to the longest nitrogen–oxygen bond. $$\mathrm{H}_{2} \mathrm{NOH}, \quad \mathrm{N}_{2} \mathrm{O}, \quad \mathrm{NO}^{+}, \quad \mathrm{NO}_{2}^{-}, \quad \mathrm{NO}_{3}^{-}$$ \(\left(\mathrm{H}_{2} \mathrm{NOH} \text { exists as } \mathrm{H}_{2} \mathrm{N}-\mathrm{OH} .\right)\)

Short answer

The order of the species with the shortest to the longest nitrogen-oxygen bond is: $$NO^{+}, \quad NO_{2}^{-}, \quad NO_{3}^{-}, \quad H_{2}NOH, \quad N_{2}O$$

Step-by-step solution

Determine Molecular Structures

Draw the Lewis structures for each of the given compounds, taking into account the formal charge on each atom. - For H2NOH (exists as H2N-OH): \(H - N - O - H\), where nitrogen has a lone pair of electrons. - For \(N_2O\): A nitrogen atom is triple-bonded to another nitrogen atom, which is single-bonded to an oxygen atom. Nitrogen atoms have one lone pair of electrons each. - For \(NO^+\): Nitrogen is double-bonded to oxygen. Nitrogen has no lone pairs, and oxygen has three lone pairs. - For \(NO_2^-\): One of the oxygen atoms is double-bonded to nitrogen, while the other oxygen is single-bonded to nitrogen with a formal negative charge. Both oxygen atoms have two lone pairs, while nitrogen has one lone pair. - For \(NO_3^-\): This ion has a resonance structure where all nitrogen-oxygen bonds are identical, with one single and two-thirds bonds. Each oxygen atom has two lone pairs, and nitrogen has no lone pairs.

Determine Bond Orders

Calculate the bond order for each nitrogen-oxygen bond in the given compounds. - H2N-OH: Single bond, bond order = 1. - \(N_2O\): Single bond, bond order = 1. - \(NO^+\): Double bond, bond order = 2. - \(NO_2^-\): One double bond and one single bond, average bond order = 1.5. - \(NO_3^-\): Resonance structure, bond order = 1.67 (two-thirds bonds).

Arrange Compounds by Nitrogen-Oxygen Bond Length

Based on the bond orders calculated in Step 2, arrange the compounds in the order of shortest to longest nitrogen-oxygen bond lengths. Shortest bond (highest bond order) to longest bond (lowest bond order): 1. \(NO^+\) (bond order = 2) 2. \(NO_2^-\) (bond order = 1.5) 3. \(NO_3^-\) (bond order = 1.67) 4. H2N-OH and \(N_2O\) (bond order = 1) The order of the species with the shortest to the longest nitrogen-oxygen bond is: $$NO^{+}, \quad NO_{2}^{-}, \quad NO_{3}^{-}, \quad H_{2}NOH, \quad N_{2}O$$

Key concepts

Nitrogen-Oxygen Bonds

Nitrogen-oxygen bonds are an essential aspect of chemical bonding that play a significant role in the structure and stability of various compounds. These types of bonds involve a nitrogen atom and an oxygen atom connecting through shared electrons.
The nature of the nitrogen-oxygen bond can vary from single to triple bonds, influencing the bond length and strength. The higher the bond order, the stronger and shorter the bond.

• Single bonds have the longest and weakest interactions, typical in molecules where nitrogen shares only one pair of valence electrons with oxygen.
• Double bonds exhibit moderate strength and length, involving two pairs of shared electrons.
• Triple bonds, being the strongest and shortest, are rare but indicate a high level of electron sharing, providing great stability to the compound.

Understanding these variations will help analyze the behavior and reactivity of different nitrogen-oxygen compounds.

Lewis Structures

Lewis structures are a straightforward method for visualizing the arrangement of electrons in molecules. They illustrate the distribution of bonding and non-bonding (lone) electron pairs among the atoms in a compound.
In constructing a Lewis structure, recognize the following steps:

• Count the total number of valence electrons present in the molecule.
• Determine the central atom, usually the least electronegative, excluding hydrogen.
• Connect the atoms using single lines to represent bonds, and distribute the rest of the electrons to satisfy the octet rule for all atoms involved.

Lewis structures are essential for predicting molecular geometry, the type of bonds present, and the likely resonance structures in cases where multiple bonding arrangements are possible, such as in nitrate ion \(NO_3^-\).

Molecular Bond Order

Molecular bond order is a concept used to describe the stability of chemical bonds within a molecule. It refers to the average number of bonds between a pair of atoms.
Calculating bond order involves counting the number of bonding pairs between two atoms and dividing by the number of equivalent positions.
In simple terms:

• A bond order of 1 signifies a single bond, indicating one count of bonding electrons.
• A bond order of 2 represents a double bond, where two pairs of electrons are shared.
• A bond order of 3 indicates a triple bond, with three pairs of electrons shared between the atoms.

Higher bond orders typically translate to shorter and stronger bonds, meaning that as the bond order increases, the bond length decreases. Thus, the concept aids in predicting the comparative lengths and strengths of nitrogen-oxygen bonds in different molecular species.

Resonance Structures

Resonance structures are a way to represent molecules where the electron distribution can be depicted in more than one valid Lewis structure. In such cases, electrons are delocalized, meaning they are shared among more than two atoms.
These structures demonstrate the "true" hybrid spread of electron density, often resulting in more stable molecular formations.
For example, the nitrate ion \(NO_3^-\) is a classic example, as it can be seen as having several resonance structures. This results in bonds that are not purely single or double but rather a hybrid, shown by the bond order of 1.67.
Resonance doesn't imply the molecule shifts between structures; rather, it explains a more accurate depiction of its electron distribution, which stabilizes the molecule more than any single structure could suggest.