Question

Some of the important pollutants in the atmosphere are ozone \(\left(\mathrm{O}_{3}\right)\), sulfur dioxide, and sulfur trioxide. Write Lewis structures for these three molecules. Show all resonance structures where applicable.

Short answer

The Lewis structures for the three atmospheric pollutants are: 1. Ozone (O3): O = O - O ↔ O - O = O (two resonance structures) 2. Sulfur dioxide (SO2): O = S - O (no resonance structures) 3. Sulfur trioxide (SO3): S = O - O - O → O = S - O - O → O - O = S - O (three resonance structures)

Step-by-step solution

Determine the total number of valence electrons for each molecule

We can find the number of valence electrons for each non-metal element using the periodic table. Oxygen (O) has 6 valence electrons, and sulfur (S) has 6 valence electrons as well. For ozone (O3), we have three oxygen atoms, so our total number of valence electrons is: 3 atoms × 6 electrons = 18 electrons For sulfur dioxide (SO2), we have one sulfur atom and two oxygen atoms. The total number of valence electrons is: 1 (S atom) × 6 electrons + 2 (O atoms) × 6 electrons = 18 electrons For sulfur trioxide (SO3), we have one sulfur atom and three oxygen atoms. The total number of valence electrons is: 1 (S atom) × 6 electrons + 3 (O atoms) × 6 electrons = 24 electrons

Draw the Lewis structures for ozone (O3)

For ozone, we have 18 valence electrons to distribute between the three oxygen atoms to satisfy the octet rule. We can arrange the atoms in a straight line, with the single atom Oxygen as the central atom. In this structure, we will form a single bond between the central O and each of its neighboring O atoms, and then distribute the remaining electrons evenly as lone pairs. Each O atom should have two lone pairs and we will have two more electrons left to form a double bond between the central O and one of the neighboring O atoms. Finally, we can draw the resonance structure by shifting the double bond to the other neighboring O. This gives us two resonance structures for ozone as shown: O = O - O ↔ O - O = O

Draw the Lewis structures for sulfur dioxide (SO2)

For sulfur dioxide, we have 18 valence electrons to distribute between one sulfur atom and two oxygen atoms. We can arrange the atoms with sulfur (S) in the center, bonded to each of the oxygen (O) atoms. We distribute all six valence electrons for each of the oxygen atoms, including a double bond (four shared electrons) with the sulfur atom. Each of the oxygen atoms will have two lone pairs, and the sulfur atom will have one lone pair. The resulting Lewis structure for SO2 is as follows: O = S - O There are no resonance structures for sulfur dioxide since there are no additional possible electron configurations that satisfy the octet rule for all atoms.

Draw the Lewis structures for sulfur trioxide (SO3)

For sulfur trioxide, we have 24 valence electrons to distribute among the sulfur atom and the three oxygen atoms. With sulfur (S) in the center, we arrange the three oxygen (O) atoms around it. We will form a double bond between sulfur and each of the oxygen atoms. This will use up 6 electrons for each of the oxygen atoms (18 for all of the oxygen atoms). Now, each oxygen atom has two lone pairs, and we're left with 6 more electrons. We can distribute them as a lone pair for the sulfur atom, making the sulfur atom have 12 electrons. With this, we can draw the resonance structures, which involve shifting the double bond between the sulfur and one of the neighboring oxygens. The resonance structures for SO3 are as follows: S = O - O - O → O = S - O - O → O - O = S - O In conclusion, we have drawn the Lewis structures for ozone, sulfur dioxide, and sulfur trioxide, as well as their resonance structures where applicable.

Key concepts

Ozone (O3) Lewis Structure

Understanding the Lewis structure for ozone ((O_3)) involves visualizing how its 18 valence electrons are shared among the three oxygen atoms. In this arrangement, one oxygen atom is sandwiched between two others, forming a bent molecular shape. Each outer oxygen atom is connected to the central oxygen by a single bond, accompanied by three pairs of 'lone' electrons that are not involved in bonding.

However, to satisfy the octet rule, which dictates that atoms seek to have eight electrons in their valence shell, one of the single bonds is converted to a double bond, resulting in a structure with a single and double bond emanating from the central oxygen atom. The true distribution of electrons is a hybrid of these possibilities, leading to what we call resonance structures, which we will delve into in a later section.

Ozone's intriguing properties and relevance in the atmosphere as a UV shield make it an essential compound to understand, and its Lewis structure gives us insight into its reactive nature.

Sulfur Dioxide (SO2) Lewis Structure

The Lewis structure of sulfur dioxide ((SO_2)) is determined by arranging its 18 valence electrons in a way that fulfills the octet rule. At the center, the sulfur atom (S) forms double bonds with each of the two oxygen atoms (O). These double bonds account for 16 of the electrons, with the remaining 2 forming a lone pair on the sulfur atom.

The arrangement leaves each oxygen atom with two lone pairs, creating a bent molecular geometry similar to that of water ((H_2O)). The SO2 molecule does not exhibit resonance because all atoms already comply with the octet rule, making further electron shifts unnecessary. Sulfur dioxide's structure, featuring a central sulfur atom with two oxygens attached via double bonds, plays a fundamental role in its properties, including its role as a pollutant and its acidic behavior when it combines with water in the atmosphere.

Sulfur Trioxide (SO3) Lewis Structure

Sulfur trioxide ((SO_3)), with its 24 valence electrons, presents a more complex Lewis structure. The sulfur atom is at the center of the molecule, surrounded symmetrically by three oxygen atoms. Each oxygen is doubly bonded to the sulfur, with two lone pairs of electrons resting on each oxygen atom.

The double bonds utilize 18 of the valence electrons, leaving the sulfur atom with an extra lone pair. This arrangement satisfies all octet rules; however, sulfur ends up with an expanded octet, a phenomenon possible due to sulfur's position in the third period of the periodic table, allowing it to utilize d-orbitals to accommodate more than eight electrons. While the written Lewis structure shows a single configuration, SO3 actually has resonance structures that portray the different possible arrangements of the double bonds, indicating a delocalized electron system across the entire molecule.

Resonance Structures

Resonance structures are alternate ways of drawing the Lewis structure of a molecule, showing different possible arrangements of the electrons. For molecules like ozone ((O_3)) and sulfur trioxide ((SO_3)), resonance is crucial as it represents the molecule's true, delocalized electron distribution rather than a fixed arrangement with a single-double bonds.

Resonance structures are drawn using double-headed arrows indicating equivalence, not a reaction. They must all have the same number of electrons, the same molecular geometry, and comply with the octet rule where possible. These structures imply that the bonds are a hybrid of the depicted limits, ultimately leading to more stability within the molecule and influencing its chemical behavior.

Valence Electrons

Valence electrons are the electrons located in the outermost shell of an atom and play a pivotal role in chemical bonding and reactions. They are the agents of chemical change, and their count determines how an atom interacts with others to form molecules.

In the context of drawing Lewis structures, knowing the number of valence electrons for each element is the first step to determine how atoms will bond and whether there will be lone pairs. For example, both oxygen and sulfur typically have six valence electrons, allowing them to form multiple bonds or have lone pairs, depending on the requirement of the molecule's geometry to achieve a stable electronic configuration. The careful distribution of these electrons according to the octet rule (and occasionally the expanded octet rule) is essential for creating accurate Lewis structures.