Question

Write a Lewis structure for each of the following simple molecules. Show all bonding valence electron pairs as lines and all nonbonding valence electron pairs as dots. For those molecules that exhibit resonance, draw the various possible resonance forms. a. \(\mathrm{NO}_{2}\) b. \(\mathrm{H}_{2} \mathrm{SO}_{4}\) c. \(\mathrm{N}_{2} \mathrm{O}_{4}\)

Short answer

The Lewis structures for the given molecules are as follows: a. \(\mathrm{NO}_{2}\) has resonance structures N=O--O and N--O=O. b. \(\mathrm{H}_{2}\mathrm{SO}_{4}\) has the Lewis structure: O= S =O | | O-H H-O c. \(\mathrm{N}_{2}\mathrm{O}_{4}\) has the Lewis structure: O= N--N =O | | O O

Step-by-step solution

Determine total number of valence electrons

N has 5 valence electrons, and each oxygen atom (O) has 6 valence electrons, so the total number of valence electrons for this molecule is: \[5 + 6 + 6 = 17\]

Create the Lewis structure

The nitrogen atom (N) is the central atom, and it forms two bonds with the oxygen atoms. One of the oxygen atoms have a double bond with N, and the other oxygen atom has a single bond with N: N=O--O There are a total of 2 lone pairs (nonbonding valence electron pairs) on the oxygen atom with a single bond, and 1 lone pair on the oxygen atom with the double bond.

Identify and draw resonance forms

Because there is a difference in bond length (double and single bond) in the molecule, the resonance will be possible. The other resonance structure is obtained by swapping the double and single bond between the two O atoms: N--O=O Thus, the resonance forms of \(\mathrm{NO}_{2}\) are N=O--O and N--O=O. For molecule \(\mathrm{H}_{2} \mathrm{SO}_{4}\):

Determine total number of valence electrons

Sulfur (S) has 6 valence electrons, oxygen (O) has 6, and hydrogen (H) has 1 valence electron. Therefore, the total number of valence electrons for this molecule is: \[6 + 4(6) + 2(1) = 32\]

Create the Lewis structure

Sulfur is the central atom, and it has four oxygen atoms surrounding it, two with double bonds and two with single bonds. The two single-bonded oxygens will each be bonded with a hydrogen atom: O= S =O | | O-H H-O There are no lone pairs on the sulfur atom (S), two lone pairs on each oxygen atom with a double bond, and three lone pairs on each oxygen atom with a single bond. No resonance forms exist for this molecule since all atoms have a full octet. For molecule \(\mathrm{N}_{2} \mathrm{O}_{4}\):

Determine the total number of valence electrons

Each nitrogen atom (N) has 5 valence electrons, and each oxygen atom (O) has 6 valence electrons. The total number of valence electrons for this molecule is: \[2(5) + 4(6) = 34\]

Create the Lewis structure

In this molecule, the two nitrogen atoms are bonded to each other, and each nitrogen atom is bonded to two oxygen atoms, one with a single bond and one with a double bond: O= N--N =O | | O O There are two lone pairs (nonbonding valence electron pairs) on each nitrogen atom, and one lone pair on each oxygen atom with a double bond, and three lone pairs on each oxygen atom with a single bond. No resonance forms exist for this molecule since all atoms have a full octet.

Key concepts

Molecular Resonance

Molecular resonance is a fascinating concept in chemistry related to Lewis structures. When we draw a molecule's Lewis structure, we arrive at the most stable arrangement of electrons and bonds. However, some molecules can be represented by more than one valid Lewis structure. These variations are known as resonance structures.
In the case of the molecule \( \mathrm{NO}_{2} \), there is a difference in bond lengths between nitrogen and the oxygen atoms. As a result, the molecule exhibits molecular resonance. This means that instead of being fixed in one single form, the electrons can "resonate" or shift between two different structures.

• For \( \mathrm{NO}_{2} \), the resonance forms are shown as \( \mathrm{N=O-O} \) and \( \mathrm{N-O=O} \).
• Resonance forms help in understanding the real structure, which is a hybrid of all possible forms.

Understanding that resonance can stabilize molecules by distributing electrons more evenly is essential. It's important to note that the actual molecule isn't alternating between forms; rather, it exists as a stable "blend" of the possible structures.

Valence Electrons

Valence electrons are the outermost electrons of an atom and play a pivotal role in chemical bonding. These electrons are usually involved in forming bonds between atoms, either by sharing or transferring.

• Each element has a different number of valence electrons based on its group position in the periodic table.
• For example, nitrogen has 5 valence electrons, while oxygen has 6.

When crafting Lewis structures, counting the total number of valence electrons for all participating atoms is the first step. This count determines how many electrons are available to create bonds and the overall configuration of the molecule. Understanding valence electrons is crucial in predicting how atoms will bond and the molecule structure.

Bonding Pairs

In a Lewis structure, bonding pairs are pairs of valence electrons shared between two atoms, resulting in a chemical bond. These pairs can be visualized as lines between the atomic symbols.
Bonding pairs form two main types of bonds:

• Single Bonds: Represented by a single line, it's a pair of electrons shared between two atoms. In \( \mathrm{N_{2}O_{4}} \), the single line indicates a single bond between atoms such as N and its adjacent O with single bonds.
• Double Bonds: Depicted by two lines, indicating the sharing of two pairs (four electrons) between atoms. In the \( \mathrm{H_{2}SO_{4}} \) structure, sulfur forms double bonds with two oxygen atoms.

The concept of bonding pairs is central to understanding molecular geometry and predicts the molecule's shape. Knowing this is valuable for determining the molecule's properties and reactivity.

Nonbonding Pairs

Nonbonding pairs, also known as lone pairs, are valence electrons not involved in chemical bonding. In Lewis structures, they appear as dots surrounding an atom.
These electrons are crucial as they can influence the shape and polarity of molecules:

• Lone pairs can create regions of electron density that cause repulsions, affecting the molecule's geometrical shape.
• For instance, in \( \mathrm{NO_2} \), the oxygen atom with a single bond to nitrogen has two lone pairs, which can impact its reactivity and interaction with other molecules.

Although they don't partake in bonding, nonbonding pairs significantly affect molecular structure and the physical properties of the compounds. Recognizing the role of these electrons helps in predicting molecular behavior and interactions.