Problem 80 - Chemistry: An Atoms-Focused Approach

Question

Oxygen and sulfur combine to form a variety of different sulfur oxides. Some are stable molecules and some, including \(\mathrm{S}_{2} \mathrm{O}_{2}\) and \(\mathrm{S}_{2} \mathrm{O}_{3},\) decompose when they are heated. Draw Lewis structures for these two compounds and show all resonance forms.

Step-by-Step Solution

Answer

Question: Draw the Lewis structures and identify all resonance forms for the compounds S2O2 and S2O3. Answer: S2O2: Resonance form 1: O S=O=S O .. .. Resonance form 2: O S=O=S O .. .. S2O3: Resonance form 1: O S S=O .. .. O .. .. Resonance form 2: O S=O S .. .. .. O .. ..

Step 1: Determine the total number of valence electrons in each molecule.

To do this, find the number of valence electrons for each atom in the molecule and add them up. Sulfur has 6 valence electrons and oxygen has 6 valence electrons. For S2O2: 2 * 6 (sulfur) + 2 * 6 (oxygen) = 24 valence electrons For S2O3: 2 * 6 (sulfur) + 3 * 6 (oxygen) = 30 valence electrons

Step 2: Draw the Lewis structures of both molecules.

First, arrange the atoms in each molecule. For S2O2, start with a sulfur atom, connect it to one oxygen atom, and then connect the oxygen to another sulfur atom. For S2O3, start with two sulfur atoms and connect each to one oxygen atom, and then connect these two oxygen atoms to another oxygen atom. Now, fill the remaining valence electrons as lone pairs on each atom. S2O2: O S S O .. .. .. .. .. S2O3: O S S O .. .. .. .. .. O ..

Step 3: Check the octet rule and adjust for formal charges.

Make sure that each atom has 8 valence electrons (except for hydrogen, which should have 2). If any atoms have too few electrons, form a double or triple bond to satisfy the octet rule. For S2O2, it's necessary to form a double bond between the central oxygen atom and one of the sulfur atoms: O S=O=S O .. .. For S2O3, form a double bond between the central oxygen atom and one of the sulfur atoms: O S S=O .. .. O .. ..

Step 4: Identify and draw resonance forms.

If a Lewis structure with multiple equivalent arrangements of the double or triple bond exists, these structures are called resonance forms. For S2O2, there's one additional resonance form where the double bond is formed between the central oxygen and the other sulfur atom: O S=O=S O .. .. For S2O3, there's another resonance form where the double bond is formed between the central oxygen atom and other sulfur atom: O S=O S .. .. .. O .. .. Now, the Lewis structures including all resonance forms are complete: S2O2: Resonance form 1: O S=O=S O .. .. Resonance form 2: O S=O=S O .. .. S2O3: Resonance form 1: O S S=O .. .. O .. .. Resonance form 2: O S=O S .. .. .. O .. ..

Key Concepts

Resonance Forms

In chemistry, resonance forms are different ways to draw the same molecule. They show how electrons can be arranged in various structures. Resonance is useful because it gives us a better idea of the molecule's stability. It spreads out electrons over multiple structures, making the molecule more stable than if it were just one structure.
For sulfur oxides like \( \text{S}_2\text{O}_2 \) and \( \text{S}_2\text{O}_3 \), resonance forms illustrate possible arrangements of electrons. Imagine a set of drawings, each showing the molecule's bonds differently. No single drawing shows the true molecule. Instead, the real structure is a blend of these forms.

• In \( \text{S}_2\text{O}_2 \), you can switch which sulfur and oxygen are connected by a double bond. These switches show different resonance forms.
• For \( \text{S}_2\text{O}_3 \), two different sulfur-oxygen pairings can create double bonds. This variation results in multiple resonance structures.

Understanding resonance is crucial. It explains molecule behavior beyond just one simple structural drawing.

Sulfur Oxides

Sulfur oxides are compounds composed of sulfur and oxygen. They are part of the chemical family known as oxides. Sulfur oxides can vary widely in their chemical structure and properties. Some are gases and some are solids. Common examples include sulfur dioxide (\( \text{SO}_2 \)) and sulfur trioxide (\( \text{SO}_3 \)).
These oxides are important in both natural and industrial processes. They can form from volcanic eruptions or burning fossil fuels. In the context of the exercise, we focus on specific oxides \( \text{S}_2\text{O}_2 \) and \( \text{S}_2\text{O}_3 \), which may not be as stable as their more common counterparts.

• \( \text{S}_2\text{O}_2 \) and \( \text{S}_2\text{O}_3 \) can decompose when heated. This instability is a factor in understanding their chemistry.
• Their formation and decomposition are of interest in various chemical reactions and applications.

Sulfur oxides play significant roles in environmental and atmospheric chemistry, affecting weather patterns and air quality.

Valence Electrons

Valence electrons are the outermost electrons of an atom. They are responsible for forming bonds with other atoms. Drawing Lewis structures depends heavily on counting these electrons. In the case of sulfur oxides, each sulfur and oxygen atom contributes valence electrons to total pool.
Here’s how to calculate valence electrons for \( \text{S}_2\text{O}_2 \) and \( \text{S}_2\text{O}_3 \):

• Each sulfur atom has 6 valence electrons.
• Each oxygen atom also has 6 valence electrons.

To find the total number:

• In \( \text{S}_2\text{O}_2 \): Multiply the valence electrons of sulfur by 2 and oxygen by 2, then add them: \( 2 \times 6 + 2 \times 6 = 24 \) electrons.
• In \( \text{S}_2\text{O}_3 \): Do the same, but with 3 oxygen atoms: \( 2 \times 6 + 3 \times 6 = 30 \) electrons.

Using these electrons, we can draw proper Lewis structures showing how atoms are connected by bonds and shared electrons. Valence electrons are a fundamental concept in understanding and predicting chemical bond formation.