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. \(\mathrm{Cl}_{2} \mathrm{O}\) b. \(\mathrm{CO}_{2}\) c. \(S O_{3}\)

Short answer

The Lewis structures for the given molecules are: a. $\mathrm{Cl}_{2}\mathrm{O}$: O=Cl-O, with each Cl having 3 lone pairs and O having 2 lone pairs. b. $\mathrm{CO}_{2}$: O=C=O c. $S O_{3}$ exhibits resonance with three forms: 1. O=S-O | O (lone pairs on oxygens) 2. -O-S=O | O (lone pairs on oxygens) 3. -O | S-O | O= (lone pairs on oxygens)

Step-by-step solution

a. Lewis structure for \(\mathrm{Cl}_{2} \mathrm{O}\)

Find the total number of valence electrons for \(\mathrm{Cl}_{2} \mathrm{O}\). Chlorine has 7 valence electrons and oxygen has 6, so the total valence electrons are \(7 \times 2 + 6 = 20\). Next, place the oxygen in the center and place the two chlorine atoms on either side. Connect each chlorine to the oxygen with single bonds, using a pair of electrons for each bond. Now distribute the remaining electrons to complete the octets of each atom. Since we used four electrons for the single bonds, we have 16 electrons left to distribute. Assign the remaining electrons to the chlorine atoms first and then to the oxygen atom. The final representation is: O = Cl - O, where each Cl has 3 lone pairs and O has 2 lone pairs.

b. Lewis structure for \(\mathrm{CO}_{2}\)

Find the total number of valence electrons for \(\mathrm{CO}_{2}\). Carbon has 4 valence electrons and oxygen has 6. So the total valence electrons are \(4 + 6 \times 2 = 16\). Place carbon in the center and place the two oxygen atoms on either side. Connect each oxygen to the carbon with double bonds, using two pairs of electrons for each bond. Since we used all the valence electrons (16) to form double bonds, no electrons are left to distribute. The final Lewis structure is: O=C=O

c. Lewis structure for \(S O_{3}\)

Find the total number of valence electrons for \(S O_{3}\). Sulfur has 6 valence electrons, and oxygen has 6. So the total valence electrons are \(6 + 3 \times 6 = 24\). Place sulfur in the center and three oxygen atoms around it. Connect each oxygen to the sulfur with single bonds, using a pair of electrons for each bond. We've now used 6 electrons to form single bonds, so we have 18 electrons left to distribute. Check if the sulfur atom needs to expand its octet due to the presence of more than 8 electrons around it. If so, we can form double bonds. For resonance forms, we keep the central sulfur with single bonds to two oxygens and create a double bond with the third oxygen in one structure. In the next form, we shift the double bond to another oxygen, and in the third form, we move the double bond once more, completing the resonance forms. The final Lewis structures are: Resonance form 1: \(O= S - O\) | \( O \) with lone pairs on oxygens. Resonance form 2: \( - O - S = O\) | \( O \) with lone pairs on oxygens. Resonance form 3: \( - O - \) | \(S - O\) | \( O =\) with lone pairs on oxygens.

Key concepts

Chemical Bonding

Chemical bonding is fundamental to understanding the behavior and properties of molecules. It's the process where atoms or ions are held together in a molecule or compound through the attraction of particles. Two common types of bonds are ionic bonds and covalent bonds.

Ionic bonding occurs when there is a complete transfer of electrons from one atom to another, resulting in positive and negative ions that attract each other. Covalent bonding, on the other hand, involves the sharing of electron pairs between atoms. The shared electrons allow each atom to achieve a more stable electron configuration, often observed in nonmetallic elements.

When drawing Lewis structures, like in the problems of Cl₂O, CO₂, and SO₃, we illustrate the covalent bonds as lines connecting the atoms and the lone pairs of electrons as dots surrounding them. Covalent bonds can be single, double, or triple, depending on the number of shared electron pairs. A deeper understanding of chemical bonding can explain the structure and reactivity of molecules and is crucial for predicting the shape and properties of a substance.

Valence Electrons

Valence electrons play a crucial role in chemical bonding, as they are the outermost electrons of an atom and are involved in forming bonds. Knowing the number of valence electrons is key to drawing Lewis structures which visually represent the arrangement of electrons in a molecule.

In the exercise provided, you saw that for Cl₂O, chlorine has 7 valence electrons and oxygen has 6, giving a total of 20 valence electrons to work with for the Lewis structure. Similarly, carbon in CO₂ has 4 valence electrons and each oxygen has 6. For SO₃, sulfur and oxygen both possess 6 valence electrons.

When creating Lewis structures, you first count the valence electrons available, then you use them to form bonds between atoms. The remaining valence electrons are then distributed as lone pairs to complete the octets, ensuring each nonmetal atom has 8 electrons to achieve a noble gas configuration. This principle is a guide to satisfying the octet rule, with some exceptions such as expanded octets or incomplete octets in certain molecules.

Resonance Structures

Resonance structures are a tool chemists use to depict the delocalization of electrons in certain molecules that cannot be represented by a single Lewis structure. In essence, resonance structures are different representations that show the same arrangement of atoms but with different placements of electrons.

It is important to note that resonance structures do not indicate that the molecule is rapidly changing between forms. Instead, they represent a hybrid of the possible configurations, with the molecule having a structure that is a blend of the resonance forms. The true electronic structure is an average of all the resonance forms and typically more stable than any individual structure could be.

In the provided exercise with SO₃, each Lewis structure showed a different oxygen atom double-bonded to sulfur. These are not discrete structures hopping between forms; rather, the true structure of SO₃ is a blend of these forms. Resonance is a crucial concept as it often stabilizes molecules, impacts molecular shape, and influences properties like bond length and energy.