Problem 98 - Chemistry Principles and Practice

Question

A In the gas phase, the oxide \(\mathrm{N}_{2} \mathrm{O}_{5}\) has a structure with an \(\mathrm{N}-\mathrm{O}-\mathrm{N}\) core, with the other four oxygen atoms in terminal positions. In contrast, in the solid phase, the stable form is \(\left[\mathrm{NO}_{2}\right]^{+}\left[\mathrm{NO}_{3}\right]^{-} .\) Draw one Lewis structure of the molecular form (with \(\mathrm{N}-\mathrm{O}-\mathrm{N}\) single bonds) and all possible resonance structures of both ions observed in the solid. Remember that second-period elements never exceed an octet.

Step-by-Step Solution

Answer

Draw Lewis structures using octet rules and consider resonance for stable formations.

Step 1: Understanding the Structure

The molecular form of dinitrogen pentoxide, \( N_2O_5 \), has an \( N-O-N \) core with the other oxygen atoms at terminal positions. First, note that each nitrogen atom should form a maximum of four bonds, including lone pairs, because it cannot have more than an octet of electrons.

Step 2: Drawing the Lewis Structure for \\( N_2O_5 \\\)

Start by connecting the two nitrogen atoms by an oxygen: \( N-O-N \). Each nitrogen atom will then be bonded to two more oxygen atoms. Assign a single bond between the first nitrogen and an oxygen atom, and from the same nitrogen, add the last oxygen with a double bond to satisfy the octet rule. Repeat the process for the other nitrogen atom. Complete the octets of all oxygen atoms using lone pairs.

Step 3: Lewis Structure of \\[NO_2^+\] Ion

The \( NO_2^+ \) ion can be represented with nitrogen forming a double bond with one oxygen and a coordinate covalent bond with the other. Place one pair of lone pairs on one oxygen and none on nitrogen to satisfy the positive charge, ensuring the nitrogen atom is part of the octet.

Step 4: Lewis Structure of \\[NO_3^-\] Ion

For the \( NO_3^- \) ion, one nitrogen atom is bonded to three oxygen atoms, one with a double bond. Place lone pairs on each oxygen to complete their octets, ensuring one lone pair serves as the charge to make it negative. All three oxygen bonds have resonance, indicating delocalized electrons.

Step 5: Resonance Structures

Draw alternative resonance structures for the \( NO_3^- \) ion, rotating the double bond among the three oxygens while maintaining the negative charge and octets. For \( NO_2^+ \), show resonance by alternating the location of the double bond between two oxygen atoms.

Key Concepts

Lewis Structures

Lewis structures are a vital part of understanding chemical bonding. They provide a visual representation of how atoms are connected in a molecule. In a Lewis structure, bonds are represented by lines between atoms, while electrons that are not involved in bonding, called lone pairs, are shown as dots surrounding the atoms.
In drawing Lewis structures, there are key steps to follow:

• Count the total number of valence electrons available from all the atoms.
• Arrange the atoms to show specific connections, often guided by the molecular formula.
• Use lines to draw bonds, typically a single line for a single bond (sharing of two electrons) between atoms. Longer lines represent double or triple bonds.
• Distribute the remaining electrons to satisfy the octet rule (each atom, except hydrogen, should be surrounded by eight electrons).
• Ensure the number of electrons used equals the count of valence electrons from the first step.

In \(N_2O_5\), for example, each nitrogen atom forms bonds with up to three oxygen atoms: a single bond in the molecular core and double bonds with terminal oxygen atoms. This satisfies nitrogen's need for eight electrons in its valence shell.

Resonance Structures

Resonance structures are alternate ways to represent the distribution of electrons in certain molecules or ions. A set of resonance structures depicts the same molecule but rearranges the location of electrons to better show all possible distributions.
By using resonance, chemists gain insight into how electrons can move within a molecule. Some molecules, like the \([NO_3^-]\) ion, exhibit significant resonance because the electrons are delocalized over the entire molecule. This is particularly useful in the \([NO_3^-]\) ion, where the double bond can appear between any nitrogen and oxygen pair.
Key features of resonance structures include:

• All resonance structures must have the same arrangement of atoms.
• The total number of valence electrons remains constant across structures.
• Only electron positions change — atom positions stay the same.
• Resonance structures collectively describe electron distribution more accurately than any single structure.

This concept is crucial for ions like \([NO_2^+]\) and \([NO_3^-]\), where drawing these structures helps in understanding the true nature of the electron distribution and bonding.

Octet Rule

The octet rule is a chemical rule of thumb that reflects the tendency of atoms in molecules to have eight electrons in their valence shell, giving them the same electron configuration as a noble gas.
For most elements in the second period, like carbon, nitrogen, and oxygen, satisfying the octet rule leads to greater stability. Here’s how it works:

• Each atom seeks to fill its outer shell with eight electrons through bonding (except for hydrogen which seeks two).
• Atoms achieve their octet by sharing, donating, or receiving electrons.
• In molecules like \(N_2O_5\), each nitrogen atom shares electrons with oxygen atoms to meet the octet requirement.

The octet rule drives the formation of specific bond types. For example, in \(N_2O_5\), nitrogen forms multiple bonds with oxygen to satisfy the rule. However, it's crucial to understand that there are exceptions, particularly in elements beyond the second period, but for second-period elements, it's a reliable guide for predicting molecule structure.