Question

The cation \(\mathrm{N}_{2} \mathrm{F}^{+}\) is isoelectronic with \(\mathrm{N}_{2} \mathrm{O}\). a. What does it mean to be isoelectronic? b. Draw the Lewis structure of \(\mathrm{N}_{2} \mathrm{F}^{+}\). (Hint: The molecule contains a nitrogen-nitrogen bond.) c. Which atom has the +1 formal charge in the structure you drew in part (b)? d. Does \(\mathrm{N}_{2} \mathrm{F}^{+}\) have resonance forms? e. Could the middle atom in the \(\mathrm{N}_{2} \mathrm{F}^{+}\) ion be a fluorine atom? Explain your answer.

Short answer

Answer: The Lewis structure of N2F+ is: F || N = N - N ^ The atom with the +1 formal charge is the nitrogen atom (N1) that is bonded to the fluorine atom.

Step-by-step solution

a. Isoelectronic definition

Isoelectronic refers to two or more chemical species that have the same number of electrons and the same electronic structure, and hence exhibit similar chemical properties.

b. Drawing the Lewis structure of N2F+

Step 1. Calculate the total number of valence electrons: Nitrogen (N) has 5 valence electrons, and Fluorine (F) has 7 valence electrons. Since N2F+ is a cation with a +1 charge, it has one less electron. Total valence electrons = (2 * 5) + 7 - 1 = 16 valence electrons Step 2. Arrange the atoms: Place a Nitrogen-Nitrogen (N-N) bond as hinted in the center and connect the Fluorine (F) atom to one of the Nitrogen atoms. - N-N-F Step 3. Distribute the valence electrons as pairs: - First, complete the octet rule for the outer atoms (F) - Then place remaining electron pairs around the central atoms (N-N) The Lewis structure for N2F+ is: F || N = N - N ^

c. Identifying the atom with +1 formal charge

To determine the formal charge of each atom, we'll use the formula: Formal Charge = Valence Electrons - (Bonds + Lone Electrons). For the central nitrogen (N2) atom: Formal Charge = 5 - (3 + 4) = -2 For the other nitrogen (N1) atom: Formal Charge = 5 - (2 + 2) = +1 For the fluorine (F) atom: Formal Charge = 7 - (1 + 6) = 0 The atom with the +1 formal charge is the nitrogen atom (N1) that is bonded to the fluorine atom.

d. Resonance forms of N2F+

Since all octets are fulfilled and there are no other possible structures that can be drawn without violating the octet rule, N2F+ does not have any resonance forms.

e. Possibility of the middle atom being a fluorine atom

No, the middle atom in N2F+ cannot be a fluorine atom. Fluorine forms only single bonds due to its high electronegativity and its completion of the octet rule with a single bond. In addition, nitrogen can form multiple bonds and satisfy its octet with different configurations, which is essential in forming the N2F+ species. Thus, it is more reasonable to have a nitrogen-nitrogen bond (N-N) in the center instead of a nitrogen-fluorine bond (N-F).

Key concepts

Isoelectronic Species

Isoelectronic species are fascinating in chemistry because they share the same number of electrons, leading to similar electronic structures. This means they often exhibit comparable chemical behaviors.
For example, the cation \(\mathrm{N}_2 \mathrm{F}^{+}\) is isoelectronic with \(\mathrm{N}_2 \mathrm{O}\). Both species have an identical electron count, making them chemically similar in some respects.
When determining isoelectronic species, consider both neutral atoms and ions, as long as the total electron count matches. This can provide insights into their potential reactivity and properties.
By understanding this concept, you can predict and compare the behaviors of different molecules or ions, enhancing your ability to solve related chemical problems.

Formal Charge

Formal charge is a crucial concept in understanding the stability and reactivity of molecules.
It allows chemists to assess how electrons are distributed in a molecule.
To calculate the formal charge of an atom, use the formula:

• Formal Charge = Valence Electrons - (Bonds + Lone Electrons)

This calculation helps identify the most plausible Lewis structure by minimizing charges.
For example, in the \(\mathrm{N}_2 \mathrm{F}^{+}\) ion, the nitrogen atom (N1) bonded to fluorine has a +1 formal charge.
This indicates that it provides one more electron than it receives, contributing to the molecule’s overall positive charge.
Understanding which atoms carry formal charges can pinpoint reactive sites within a molecule, important for predicting chemical reactions.

Resonance Structures

Resonance structures are representations of a molecule where the electron position varies across different structures while the arrangement of atoms remains constant.
These structures illustrate the delocalization of electrons within a molecule, which can enhance stability.
However, not all molecules have resonance forms. In the case of \(\mathrm{N}_2 \mathrm{F}^{+}\), there are no resonance forms available.
All atoms fulfill the octet rule, and rearranging electrons would violate this rule or leave atoms without complete shells.
Resonance is crucial when multiple valid Lewis structures exist, helping accurately depict a molecule’s electron distribution.
It’s a powerful concept to explain phenomena like bond energy and lengths that are not evident from a single structure.

Valence Electrons

Valence electrons are the outermost electrons of an atom, vital in bonding and chemical reactivity.
They are the electrons involved in forming bonds between atoms. Knowing the number of valence electrons helps predict how atoms bond.
For the \(\mathrm{N}_2 \mathrm{F}^{+}\) ion, knowing that nitrogen has 5 and fluorine has 7 valence electrons guides the drawing of its Lewis structure.
You start by adding up the valence electrons, accounting for any charges to find the total electron count available for bonding. This guides where you can place electron pairs to satisfy bond requirements and the octet rule.
Understanding valence electrons allows chemists to propose structures and predict how molecules interact with each other.