Question

Explain why the nitrogen-oxygen bond lengths in \(\mathrm{N}_{2} \mathrm{O}_{4}\) (which has a nitrogen-nitrogen bond) and \(\mathrm{N}_{2} \mathrm{O}\) are nearly identical ( 118 and 119 pm, respectively).

Short answer

In summary, the nitrogen-oxygen bond lengths in N₂O₄ and N₂O are nearly identical (118 and 119 pm, respectively) because of the influence of resonance structures in N₂O₄, which causes the N-O bond character to be between a single and a double bond. In contrast, the N₂O molecule's N-N triple bond might pull electrons closer, making its N-O bond slightly electron-deficient. Consequently, both factors contribute to the nearly identical bond lengths despite their different bond orders.

Step-by-step solution

To understand the bond lengths, we first need to know the structure of each molecule. \(\mathrm{N}_{2}\mathrm{O}_{4}\) has a nitrogen-nitrogen bond and two nitrogen-oxygen double bonds, and \(\mathrm{N}_{2}\mathrm{O}\) has a nitrogen-nitrogen triple bond and one nitrogen-oxygen bond. We can draw their Lewis Structures as: \(\mathrm{N}_{2}\mathrm{O}_{4}: \ \ \mathrm{O=N-O-N=O}\) \(\mathrm{N}_{2}\mathrm{O}: \ \ \mathrm{N\equiv N-O}\) #Step 2: Evaluate resonance structures#

Resonance structures involve the delocalization of electrons, which can affect bond lengths. We will analyze if there are any resonance structures for these molecules. For \(\mathrm{N}_{2}\mathrm{O}_{4}\), there are two possible resonance structures: Structure 1: \(\mathrm{O=N-O-N=O}\) Structure 2: \(\mathrm{O-N=O-N=O}\) These structures indicate that electrons are delocalized between nitrogen-oxygen bonds, and the bond character is between a single and a double bond. On the other hand, \(\mathrm{N}_{2}\mathrm{O}\) has no possible resonance structures due to the presence of a nitrogen-nitrogen triple bond. #Step 3: Determine bond order for nitrogen-oxygen bonds#

Now we will calculate the bond order of nitrogen-oxygen bonds in both molecules, which can help us to understand the bond lengths. For \(\mathrm{N}_{2}\mathrm{O}_{4}\), considering the resonance structures, the bond order of nitrogen-oxygen bonds can be calculated as: Bond order = \(\frac{number \ of \ bonding \ electrons - number \ of \ antibonding \ electrons}{2}\) Bond order = \(\frac{4}{2}\) Bond order = \(2\) For \(\mathrm{N}_{2}\mathrm{O}\), the bond order of the nitrogen-oxygen bond is: Bond order = \(\frac{number \ of \ bonding \ electrons - number \ of \ antibonding \ electrons}{2}\) Bond order = \(\frac{2}{2}\) Bond order = \(1\) #Step 4: Analyze bond lengths#

The bond lengths in molecules are affected by bond order, and they are expected to be shorter for higher bond orders. However, in this case, the bond lengths in \(\mathrm{N}_{2}\mathrm{O}_{4}\) and \(\mathrm{N}_{2}\mathrm{O}\) (118 pm and 119 pm, respectively) turn out to be nearly identical. To explain this observation, consider that in the case of \(\mathrm{N}_{2}\mathrm{O}_{4}\), the bond order for nitrogen-oxygen bonds is 2, but due to resonance and electron delocalization, the actual bond character might be between a single and a double bond. Consequently, the bond lengths might be slightly longer than expected for a typical double bond. On the other hand, in the case of \(\mathrm{N}_{2}\mathrm{O}\), the nitrogen-nitrogen triple bond might pull the electrons of nitrogen-oxygen bond closer, making the latter a bit electron-deficient. Consequently, its bond length might be slightly longer than expected for a typical single bond. Therefore, both factors contribute to the nearly identical bond lengths between nitrogen-oxygen bonds in \(\mathrm{N}_{2}\mathrm{O}_{4}\) and \(\mathrm{N}_{2}\mathrm{O}\) even though their bond orders are different.

Key concepts

Resonance Structures

In chemistry, resonance structures are used to represent molecules that cannot be accurately depicted by a single Lewis structure. These structures help explain the delocalization of electrons within certain compounds. For example, in \(\mathrm{N}_2\mathrm{O}_4\), we find two distinct resonance structures consisting of \(\mathrm{O=N-O-N=O}\) and \(\mathrm{O-N=O-N=O}\).

This delocalization indicates that the double bond character of nitrogen-oxygen bonds is spread over the different nitrogen-oxygen pairs rather than being localised to one specific bond. This electron delocalization results in bond lengths that are intermediate between those expected for single and double bonds.

• In \(\mathrm{N}_2\mathrm{O}_4\), electrons are shared across multiple bonds, making the resonance a key player in the molecular stability and bond length similarity.
• For \(\mathrm{N}_2\mathrm{O}\), resonance is absent due to a dominant nitrogen-nitrogen triple bond which restricts electron delocalization and maintains its unique bond length classification.

Bond Order

The concept of bond order is essential as it provides insights into the bond strength and length in molecules. Bond order represents the average number of bonds between two atoms and can be calculated from the number of electrons involved in forming those bonds.

For \(\mathrm{N}_2\mathrm{O}_4\), the bond order for nitrogen-oxygen bonds is around 2 due to resonance. Resonance structures indicate that the bond involves both single and double bond characteristics, which balances the bond order value.

• Bond order of 2 suggests a stronger bond compared to a single bond, supporting partial double bond characteristics.

In contrast, \(\mathrm{N}_2\mathrm{O}\) holds a bond order of 1 for nitrogen-oxygen bonds, indicating features closer to a single bond. This reflects the absence of resonance, reinforcing a simple single bond nature.

• A bond order of 1 typically suggests weaker and longer bonds compared to higher bond orders.

Understanding bond orders aids in grasping why bond lengths may appear deceptively similar despite their differing chemical settings.

Nitrogen-Oxygen Bonds

Nitrogen-oxygen bonds are an intriguing aspect of molecular chemistry due to their involvement in various molecular structures. These bonds can present as single, double, or even resonate between forms, as seen in different compounds like \(\mathrm{N}_2\mathrm{O}_4\) and \(\mathrm{N}_2\mathrm{O}\).

In \(\mathrm{N}_2\mathrm{O}_4\), the resonance between nitrogen-oxygen bonds leads to an average bond length, showcasing characteristics of both single and double bonds.

• This delocalization is central to understanding why the bond length remains uniform at approximately 118 pm.

On the other hand, the nitrogen-oxygen bond in \(\mathrm{N}_2\mathrm{O}\) being influenced by the nitrogen-nitrogen triple bond, retains a length slightly longer than usual for a pure single bond, standing close to 119 pm.

• The proximity in bond lengths between different nitrogen-oxygen environments is fascinating and highlights the importance of considering electron delocalization and resonance effects.

Through these observations, we explore how electron interactions shape molecular structure, emphasizing the diverse nature of nitrogen-oxygen connectivity.