Question

A linear nitrogen anion, \(\mathrm{N}_{5}\), was isolated for the first time in 1999. a. Draw the Lewis structures for four resonance forms of linear \(\mathrm{N}_{5}^{-}.\) b. Assign formal charges to the atoms in the structures in part (a), and identify the structures that contribute the most to the bonding in \(\mathrm{N}_{5}^{-}\). c. Compare the Lewis structures for \(\mathrm{N}_{5}^{-}\) and \(\mathrm{N}_{3}^{-}\). In which ion do the nitrogen-nitrogen bonds have the higher average bond order?

Short answer

Question: Draw four resonance forms of linear N5-, assign formal charges to the atoms in the structures, identify the structures with the most contribution to the bonding, and compare the Lewis structures of N5- and N3- to determine which ion has a higher average bond order for nitrogen-nitrogen bonds. Answer: The four resonance forms of linear N5- have different distributions of coordinate bonds and formal charges, with some structures being more stable and having charges closer to zero. The total formal charge of all structures should be -1, as N5- is an anion. Comparing the Lewis structures and calculating the average bond orders for N5- and N3-, we can determine that N3- has a higher average nitrogen-nitrogen bond order than N5-.

Step-by-step solution

Draw the four resonance forms of linear N5-

To draw the resonance forms, we need to understand the molecular structure of the ion. N5- is an anion, and its linear structure means that the nitrogen atoms are connected in a line. Each nitrogen atom has 5 valence electrons. In the N5- ion, there's an extra electron, so we have a total of 26 electrons to distribute among the nitrogen atoms. Coordinate bond pairs can be moved throughout the ion to create different resonance forms. Make sure all structures are aligned linearly and that the octet rule is satisfied for each nitrogen atom.

Assign formal charges and identify the most contributing structures

To assign formal charges, use the formula: Formal charge = Valence electrons - Non-bonding electrons - 0.5 * Bonding electrons. Calculate the formal charge for each nitrogen atom in each resonance form. Then, evaluate the structures for stability. Structures with charges closer to zero are generally more stable. If possible, multiple structures may contribute to the bonding of N5-. It's crucial to consider the sum of the formal charges in a resonance structure when identifying the structures with the most contribution.

Compare Lewis structures of N5- and N3- and find average bond order

N3-, also known as the azide ion, has 3 nitrogen atoms and -1 charge. Draw the Lewis structures for both N5- and N3-. To calculate the average bond order for nitrogen-to-nitrogen bonds, understand the bond order as the number of chemical bonds between a pair of atoms. For the Lewis structure of N5- and N3-, calculate the bond orders for N-to-N bonds. Add the bond orders and divide by the total number of N-to-N bonds in each ion. Finally, compare the average bond orders between N5- and N3- and conclude which ion has the higher nitrogen-nitrogen bond order.

Key concepts

Resonance Forms

Resonance forms are different possible arrangements of electrons in a molecule. They explain how a molecule can have multiple, legitimate structures. For linear \(\text{N}_{5}^{-}\), these forms show how electrons are distributed across nitrogen atoms.
Creating resonance structures involves moving electron pairs without altering the actual arrangement of atoms. In the case of \(\text{N}_{5}^{-}\), the molecule maintains its linear shape, and each nitrogen atom adheres to the octet rule, meaning they hold 8 electrons in their outer shell.

• It's important to keep all resonance forms aligned linearly.
• Each structure reflects the same overall electron count (26 electrons for \(\text{N}_{5}^{-}\)).
• By drawing these different forms, we appreciate how the molecule can vary while still representing the same \(\text{N}_{5}^{-}\) ion.

Understanding resonance forms helps in visualizing the stability of molecules and offers insights into their potential energy states.

Formal Charge

The formal charge concept is pivotal in assessing the most stable Lewis structure among resonance forms. It is calculated for each atom using the formula:
\[\text{Formal charge} = \text{Valence electrons} - \text{Non-bonding electrons} - 0.5 \times \text{Bonding electrons}\]
Assigning formal charges helps determine which resonance structures are more plausible by evaluating stability. Generally, structures with formal charges closest to zero are considered more stable.

• In \(\text{N}_{5}^{-}\), calculate this for each nitrogen and aim for charges closest to zero.
• Structures where formal charges add up to the ion's total charge without extreme separation of charge will contribute more significantly to the bonding.

This principle aids in predicting which resonance form of \(\text{N}_{5}^{-}\) describes the molecule most accurately in a chemical context.

Average Bond Order

Average bond order is a crucial factor in understanding the strength and length of bonds within molecules. It is calculated by dividing the total number of bonds by the number of bond locations between a pair of atoms.
For \(\text{N}_{5}^{-}\) and \(\text{N}_{3}^{-}\), drawing their Lewis structures is essential to evaluate bond orders.

• For \(\text{N}_{5}^{-}\), you look at the number of nitrogen-nitrogen bonds across resonance forms.
• Calculate each bond's order within these forms, clarify the average, then do the same for \(\text{N}_{3}^{-}\), which has 3 nitrogen atoms.
• Compare these averages to see which ion possesses stronger nitrogen bonds.

By doing so, you identify which structure features higher bond strength, offering insights into molecular stability and reactivity.

Molecular Structure

The molecular structure of a compound like \(\text{N}_{5}^{-}\) reveals its arrangement of atoms in space and influences its dynamic characteristics.
For \(\text{N}_{5}^{-}\), the nitrogen atoms align linearly, giving a straightforward insight into how the electrons might influence the shape. Despite resonance changes, the molecule's linear geometry remains.

• The arrangement allows understanding of spatial atomic positions and electron sharing.
• It's critical in predicting molecular interactions with other compounds.
• A molecule's structure also impacts its physical and chemical properties, like solubility and reactivity.

Comparing with \(\text{N}_{3}^{-}\), we deepen the understanding of such structural influences, acknowledging identical chemistry with varied structural forms.