Question

Write three reasonable resonance structures for the azide ion \(\mathrm{N}_{3}^{-}\) in which the atoms are arranged as NNN. Show formal charges.

Short answer

The three resonance structures of azide ion differ in bond orders and formal charges, with variations in the distribution of double bonds and formal charges across the nitrogen atoms.

Step-by-step solution

Count Total Valence Electrons

First, we need to determine the total number of valence electrons available for the azide ion. Each nitrogen atom has 5 valence electrons, so three nitrogen atoms have a total of 15 electrons. Since the azide ion carries a -1 charge, we add one more electron, giving a total of 16 valence electrons.

Arrange Atoms and Initial Bonds

Arrange the nitrogen atoms linearly as N-N-N. Initially, assume single bonds between the consecutive nitrogen atoms: N-N-N. Account for these initial bonds by using 2 electrons per bond, so you've used 4 electrons out of the 16.

Distribute Remaining Electrons as Lone Pairs

Distribute the remaining 12 valence electrons (16 total - 4 used in bonds) as lone pairs. Start by placing lone pairs to complete the octet around each nitrogen atom while keeping track of the total electrons being used.

Create Double Bonds to Minimize Formal Charges

To minimize formal charges, adjust lone pairs into double bonds. Create possible structures by moving lone pairs into bonding positions, resulting in double bonds between nitrogen atoms to reduce high formal charges on any nitrogen atom. Adjust as follows: 1. Structure 1: N≡N-N⁻ 2. Structure 2: ⁻N-N≡N 3. Structure 3: N=N=N (where one end nitrogen gets a negative formal charge and middle nitrogen tends toward a positive formal charge).

Assign and Verify Formal Charges

Assign formal charges using the formula: Formal Charge = (Valence Electrons) - (Non-bonding Electrons + 0.5 * Bonding Electrons). Calculate for each nitrogen in all three structures, ensuring the total charge equals -1. Example for Structure 1, middle nitrogen: 5 - (0 + 0.5*8) = +1.

Validate Resonance Structures

Ensure all resonance structures obey the octet rule where possible, and all have a net charge of -1. Check that each structure can be connected via electron shifts (resonance arrows). All structures are resonance forms contributing to the resonance hybrid of N₃⁻.

Key concepts

Formal Charge Calculation

When dealing with resonance structures, understanding formal charge is key. Formal charge helps us determine the most stable configuration of a molecule or ion.

To calculate the formal charge, use the formula: \( ext{Formal Charge} = ( ext{Valence Electrons}) - ( ext{Non-bonding Electrons} + 0.5 \times \text{Bonding Electrons}) \).

• Valence Electrons: Total electrons in the outer shell.
• Non-bonding Electrons: Electrons not involved in bonding, typically found as lone pairs.
• Bonding Electrons: Electrons involved in covalent bonds with other atoms.

Understanding the formal charge for each atom in resonance structures helps us identify the most likely arrangement of electrons that stabilizes the molecule or ion, such as in the azide ion, \( ext{N}_3^- \). This can reduce extreme charges, which are generally less favorable.

Valence Electrons

Valence electrons play a vital role in chemical bonding. They are the electrons found in an atom's outermost shell and are key in forming bonds.

To determine valence electrons, simply look at an atom's group number in the periodic table. For nitrogen, which belongs to Group 15, it has 5 valence electrons.

In the azide ion, \( ext{N}_3^- \), each nitrogen atom contributes its 5 valence electrons, totaling 15 for the three nitrogen atoms. Additionally, due to the -1 charge of the ion, we add one extra electron, making it 16 valence electrons in total.

This count is crucial for distributing and forming bonds within the different resonance structures.

Octet Rule

The octet rule is a simple guideline in chemistry where atoms tend to bond in such a way that they achieve eight electrons in their valence shell, resembling the electron configuration of nobel gases.

• Nitrogen needs three additional electrons to fulfill its octet.
• It often forms three covalent bonds, either in single, double, or sometimes even triple bonds.

In the context of the azide ion, each nitrogen atom should attempt to satisfy the octet rule. Lone pairs and bonds are adjusted to maintain this rule while calculating formal charges for stability across the resonance structures.

Nitrogen Atoms

In molecules and ions, nitrogen atoms are common and versatile elements. Each nitrogen atom has 5 valence electrons and can form various types of bonds to achieve stability.

This flexibility makes them crucial in compounds like the azide ion, \( ext{N}_3^- \).

• In nitrogen bonding, it's not uncommon to see triple bonds, given nitrogen's tendency to form strong bonds.
• Nitrogen strives to complete its octet while maintaining minimal formal charge, leading to multiple possible configurations in resonance structures.

In the azide ion, nitrogen atoms are strategically bonded to form structures such as N≡N-N⁻, emphasizing their ability to form multiple bonds and resonate among different bonding arrangements.

Electron Distribution

Electron distribution refers to how valence electrons are spread across a molecule or ion, contributing to its bonding and resonance structures.

For the azide ion (\( ext{N}_3^- \)), it involves allocating 16 total valence electrons across the molecule.
By first assigning electrons to form basic bonds and then distributing remaining electrons as lone pairs or additional bonds, we arrive at stable resonance forms.

• Lone pairs are typically placed on atoms to ensure the octet rule is fulfilled.
• Electron shifts can result in different resonant configurations, allowing the molecule to stabilize through shared electron positioning.

It is crucial in achieving an understanding of how molecules can resonate between structures to reflect the most stable, electron-balanced state.