Question

Consider the nitryl cation, \(\mathrm{NO}_{2}^{+}\). (a) Write one or more appropriate Lewis structures for this ion. (b) Are resonance structures needed to describe the structure? (c) With what familiar molecule is it isoelectronic?

Short answer

The Lewis structure for NO₂⁺ is O--N--O with 3 lone pairs on each oxygen atom and 1 lone pair on nitrogen. Resonance structures are needed to describe NO₂⁺ as it can be represented by more than one equally valid Lewis structure. The familiar molecule isoelectronic with the nitryl cation (NO₂⁺) is carbon dioxide (CO₂).

Step-by-step solution

Finding the Lewis Structure for NO₂⁺

First, we need to count the total number of valence electrons. Nitrogen has 5 valence electrons and oxygen has 6 valence electrons. Since the nitryl cation has a +1 charge, we need to subtract one electron from the total count. Therefore, we have a total of \(5 + (2 \times 6) - 1 = 16\) valence electrons. The central atom will be Nitrogen, as it is the least electronegative element in the molecule. Start by placing a single bond between Nitrogen and the two oxygen atoms. It leaves 12 electrons to be distributed. Next, we can form non-bonding pairs around the oxygen atoms. Each oxygen atom will need three lone pair electrons to satisfy the octet rule. Now, we will have 6 valence electrons left. Place one lone pair of electrons on Nitrogen, leaving Nitrogen with 4 unshared electrons which fulfill its octet. This will be the Lewis structure for NO₂⁺: O--N--O (3 lone pairs on each oxygen atom and 1 lone pair on nitrogen).

Checking for Resonance Structures

To determine if resonance structures are needed, we will check if the species can be represented by more than one equally valid Lewis structure. In the current Lewis structure, nitrogen has a formal charge of +1, and each oxygen has a formal charge of 0. We can try forming a double bond between Nitrogen and one of the oxygen atoms, making the nitrogen's formal charge 0 and the other oxygen -1. Now we can draw an alternative structure with a double bond between Nitrogen and the other oxygen atom. Since we can draw multiple valid Lewis structures with different distributions of electrons, there are resonance structures needed to describe NO₂⁺.

Finding an Isoelectronic Molecule

Isoelectronic molecules contain the same number of electrons and have the same structure, but usually not the same atoms. As we know, the nitryl cation (NO₂⁺) has 16 valence electrons. To find an isoelectronic molecule, we need to find out another molecule with the same electron count and structure. Carbon dioxide (CO₂) is a molecule with the same structure as NO₂⁺. Let's check the valence electrons. Carbon has 4 valence electrons, and each oxygen atom has 6 valence electrons. Therefore, CO₂ has a total of 12 valence electrons. However, the nitrite ion (NO₂⁻) is isoelectronic with the nitryl cation (NO₂⁺). NO₂⁻ has an extra electron compared to NO₂⁺ due to its -1 charge, giving it a total of 17 valence electrons, while CO₂ only has 16 valence electrons. Thus, the familiar molecule isoelectronic with the nitryl cation (NO₂⁺) is carbon dioxide (CO₂).

Key concepts

Valence Electrons

Valence electrons are the outermost electrons of an atom and are crucial for understanding chemical bonding. They are the primary electrons involved in forming chemical bonds with other atoms.
For the nitryl cation (NO₂⁺), understanding valence electrons is essential to draw its Lewis structure.

• Nitrogen has 5 valence electrons.
• Each oxygen atom contributes 6 valence electrons.
• Because NO₂⁺ has a +1 charge, we subtract 1 electron.

In total, the ion has 5 + (2×6) - 1 = 16 valence electrons. Knowing how many valence electrons are available helps in constructing accurate Lewis structures by showing how the electrons are distributed around the atoms. This process also helps ensure that each atom achieves a stable electron configuration, typically an octet for molecules, except in cases where more or fewer electrons are common, such as in hydrogen or transition metals.

Resonance Structures

Resonance structures are multiple ways to draw a Lewis structure for a molecule, which describe the delocalization of electrons. In molecules where electrons can be distributed in different ways while the arrangement of atoms remains constant, resonance structures are used to depict all the possible configurations.
For NO₂⁺, the existence of resonance structures illustrates this phenomenon. When drawing NO₂⁺, it is possible to depict the molecule with one oxygen atom double bonded to the nitrogen, while the second oxygen holds a single bond. Consequently, the double bond can shift between the two oxygen atoms resulting in different equivalent Lewis structures.

• Structure 1: N has a double bond with one oxygen, and a single bond with the other.
• Structure 2: N's double bond shifts to the second oxygen.

Resonance maintains that the actual structure is an average of these possibilities, indicating that the bonding electrons are spread out or "delocalized" across the entirety of the molecule. This influences the chemical properties and stability of NO₂⁺, showing how distribution of electrons can change formal charges and affect reactivity.

Isoelectronic Molecules

Isoelectronic molecules are those with the same number of electrons or electron configuration but different elements or compositions. These molecules often have similar structural features and sometimes, similar properties due to their electron configurations.
For the nitryl cation (NO₂⁺), identifying an isoelectronic molecule helps us understand its properties better. Although the solution mentioned carbon dioxide (CO₂) initially, a closer look indicates that CO₂ has 16 electrons in total, matching NO₂⁺ perfectly in electron count and configuration: both have a linear structure.
To clarify, carbon dioxide is constructed as:

• One carbon atom contributes 4 valence electrons.
• Each oxygen atom contributes 6 valence electrons.

Adding these gives us 4 + (2×6) = 16 electrons, the same as NO₂⁺. This shows that CO₂ and NO₂⁺ are isoelectronic. Recognizing such similarities aids in drawing parallels across chemical behaviors and reactivity patterns of different substances.