Question

Place the species below in order of the shortest to the longest nitrogen- oxygen bond. \(\begin{array}{lllll}\mathrm{H}_{2} \mathrm{NOH}, & \mathrm{N}_{2} \mathrm{O}, & \mathrm{NO}^{+}, & \mathrm{NO}_{2}^{-}, & \mathrm{NO}_{3}^{-}\end{array}\) \(\left(\mathrm{H}_{2} \mathrm{NOH}\right.\) exists as \(\mathrm{H}_{2} \mathrm{~N}-\mathrm{OH} .\).

Short answer

The order of the species from shortest to longest nitrogen-oxygen bond is: NO+, N2O, NO2-, NO3-, H2NOH.

Step-by-step solution

Identify the structure of each species

Here are the structures of each species: H2NOH (H2N-OH): This molecule has a single bond between N and O. N2O: There are two ways to draw the structure of N2O: N≡N-O (linear) or N=N=O (bent). The N-O bond order in the linear structure is 1, and in the bent structure, it is 2. NO+: This is a diatomic ion with a bond order of 3 due to its electronic configuration. NO2-: The NO2- ion has a resonance structure, and the bond order is 1.5 for each N-O bond. NO3-: The nitrate ion (NO3-) has a resonance structure with three equally contributing forms. The bond order is 4/3 or 1.33 for each N-O bond since there are 2 donated electrons forming bonds with three oxygen atoms.

Compare the bond orders

Now we can compare the bond orders of the different species: H2NOH (H2N-OH): Bond order = 1 N2O (N=N=O or N≡N-O): Bond order = 1 or 2 NO+: Bond order = 3 NO2-: Bond order = 1.5 NO3-: Bond order = 1.33

Arrange the species in order of increasing nitrogen-oxygen bond length

Using the relationship between bond order and bond length (higher bond order corresponds to shorter bond length), we can arrange the species in ascending order of N-O bond length: NO3- (bond order 1.33) < H2NOH (bond order 1) < NO2- (bond order 1.5) < N2O (lowest bond order 1) < NO+ (bond order 3) The final order from shortest to longest N-O bond length is: NO+, N2O, NO2-, NO3-, H2NOH

Key concepts

Chemical Bonding

Chemical bonding is a fundamental aspect of chemistry that explains how atoms connect to form molecules. At its core, chemical bonding involves the sharing or transferring of electrons between atoms to achieve a full outer shell of electrons, which is the most stable electron configuration.

There are three primary types of chemical bonds: ionic, covalent, and metallic. Ionic bonds occur when one atom donates an electron to another, creating oppositely charged ions that attract each other. Covalent bonds arise when atoms share electrons to fill their outer shells. Within covalent bonds, there can be single, double, or triple bonds, indicating the number of shared electron pairs, with triple bonds being the strongest and shortest.

In the case of nitrogen-oxygen compounds, the covalent bonds between N and O can vary in bond order - that is, the number of electron pairs shared between the two atoms. A higher bond order indicates stronger and shorter bonds. For example, \textbf{NO\(^+\)} has a bond order of 3, meaning it has a triple bond that's shorter and stronger than the single bonds found in \textbf{H\(_2\)NOH}. This directly influences the nitrogen-oxygen bond length and stability within different molecules.

Molecular Geometry

Molecular geometry refers to the three-dimensional arrangement of atoms within a molecule. This geometry is determined by factors such as the number of bonds, the types of bonds, and the repulsion between electron pairs in the valence shell of the central atom.

VSEPR (Valence Shell Electron Pair Repulsion) theory is a useful model in predicting molecular geometry. It states that electron pairs around a central atom will arrange themselves as far apart as possible to minimize repulsion. Depending on the number and arrangement of bonding pairs and lone pairs of electrons, molecules can adopt various shapes, such as linear, bent, trigonal planar, tetrahedral, and more.

For instance, \textbf{NO\(_3^-\)} has a trigonal planar geometry with 120-degree bond angles, allowing for equal resonance structures to contribute to its overall shape and bond lengths. On the other hand, \textbf{N\(_2\)O} can have a linear or bent shape, depending on the resonance structure considered, which in turn affects the bond length between the nitrogen and oxygen atoms.

Resonance Structures

Resonance structures are a set of two or more Lewis structures that represent the delocalization of electrons in molecules that cannot be accurately depicted by a single Lewis structure. These structures are not real entities but conceptual tools that reflect the greater stability of the delocalized system.

Electrons in molecules with resonance are not restricted to one location or bond; instead, they are spread over several atoms, often contributing to unusual bond lengths. The actual distribution of electrons in the molecule is a hybrid, or average, of all possible resonance forms. Important factors such as bond order, bond length, and the overall energy of the molecule are derived from the concept of resonance.

For example, \textbf{NO\(_2^-\)} and \textbf{NO\(_3^-\)} both have resonance structures that impact their bond orders and lengths. In \textbf{NO\(_2^-\)}, the two N-O bonds are equivalent, with a bond order of 1.5. In contrast, \textbf{NO\(_3^-\)} has three resonance structures that average out to give a bond order of 1.33 for each N-O bond, slightly elongating them compared to \textbf{NO\(_2^-\)}. Recognizing resonance is crucial in understanding the electronic structure and characteristics of a molecule.