Question

Draw the Lewis resonance structures for the following molecules. $$ \mathrm{O}_{3} $$

Short answer

The two Lewis resonance structures for ozone (O3) are as follows: Structure 1: \( O = O - O^(+) \) with three lone pairs on the left oxygen atom and two lone pairs on the right oxygen atom. Structure 2: \( O^(+) - O - O \) with two lone pairs on the left oxygen atom and three lone pairs on the right oxygen atom. Both structures have formal charges of 0 for the central and left oxygen atoms, and +1 and -1 for the right oxygen atoms (with charges switching places in each resonance structure).

Step-by-step solution

Determine the total number of valence electrons

To draw the resonance structures for O3, we first need to count the total number of valence electrons for all three oxygen atoms. Each oxygen atom has 6 valence electrons, so we have a total of 18 valence electrons for O3.

Draw the central atom and connect the other oxygen atoms

Put one of the oxygen atoms (O) at the center, and connect the other two oxygen atoms to it using single covalent bonds. This will use up 2x2 = 4 valence electrons, and we will have 14 valence electrons left.

Distribute the remaining electrons as lone pairs

Now, we have to distribute the remaining 14 electrons as lone pairs to complete the octet for each oxygen atom. First, we need to complete the octet rule for the central oxygen atom. We have used 2 of its 6 valence electrons, so we can add two lone pairs (4 electrons) to it. This completes the octet for the central oxygen atom. Now, we can distribute the remaining 10 electrons (5 lone pairs) to the other two oxygen atoms. Place three lone pairs (6 electrons) on one oxygen atom and two lone pairs (4 electrons) on the other oxygen atom. Now that we have placed all 18 valence electrons on the oxygen atoms, we have drawn a Lewis structure for O3.

Check formal charges and draw resonance structures

To check whether we've drawn a correct resonance structure, we will calculate the formal charges for each oxygen atom. Formal Charge = Valence Electrons - Non-bonding Electrons - (1/2 x Bonding Electrons) For the central atom (O): Formal Charge = 6 - 4 - (1/2 x 4) = 0 For the left oxygen atom (O) with 3 lone pairs: Formal Charge = 6 - 6 - (1/2 x 2) = 0 For the right oxygen atom (O) with 2 lone pairs: Formal Charge = 6 - 4 - (1/2 x 4) = 0 Now that we've confirmed the formal charges are correct, we can draw the resonance structure by swapping the positions of the oxygen atoms with 2 and 3 lone pairs. This creates the second resonance structure of O3, where the only difference is the location of the oxygen atoms with 2 and 3 lone pairs. Thus, we have drawn the two resonance structures for ozone (O3).

Key concepts

valence electrons

Valence electrons are crucial for determining how atoms interact and bond with each other. They are the electrons located in the outermost shell of an atom, and are responsible for the chemical properties of the element. In the context of drawing Lewis structures, knowing the number of valence electrons helps in figuring out how many electrons are available for forming bonds and completing octets.
For ozone (O\(_3\)), each oxygen atom possesses 6 valence electrons. Since there are three oxygen atoms in O\(_3\), the molecule collectively has a total of 18 valence electrons. Identifying these electrons allows us to distribute them correctly and ensure that each atom achieves a stable configuration through bonding or lone pairs.

octet rule

The octet rule is a chemical rule of thumb that states atoms are most stable with eight electrons in their valence shell. This concept is essential when drawing Lewis structures, as it ensures each atom achieves a complete and stable electron configuration.
In the case of O\(_3\) (ozone), each oxygen atom tries to complete its octet either through bonding or non-bonding electrons. When distributing 18 valence electrons among the three oxygen atoms:

• Single bonds are used initially to connect the atoms, consuming some valence electrons.
• Additional electrons are then allocated as lone pairs to fulfill the octet for each oxygen atom, ensuring stability.

The central oxygen atom gets extra attention to balance its electron count with eight, allowing for a proper octet while maintaining the bonds.

formal charge

Formal charge assists in verifying the most likely distribution of electrons in a molecule's Lewis structure. It is calculated using the formula:\(\text{Formal Charge} = \text{Valence Electrons} - \text{Non-bonding Electrons} - \frac{1}{2} \text{Bonding Electrons}\)Assessing formal charge ensures the structure makes sense and is electrically neutral or carries the least charge possible.
In the ozone molecule's structure calculation:

• The central oxygen atom has a formal charge of zero, indicating a balanced electron distribution.
• The same applies to the other oxygen atoms once single bonds and lone pairs are accounted for correctly.

By ensuring the formal charges are optimal (close to zero), you can determine that the structure is plausible and a likely depiction of the real molecule.

resonance structures

Resonance structures are multiple ways to represent a molecule's bonding using Lewis structures, which differ only in the positioning of electrons (not atoms). They help illustrate the delocalization of electrons within a molecule.
For ozone (O\(_3\)), the existence of resonance is expressed by alternating the position of the double bond among the different oxygen atoms. In one structure, the double bond may be between one set of atoms, and in another, it's between another set.
Resonance structures reflect the real scenario where electrons are not static but are shared over multiple paths. This phenomenon aids in achieving stability by allowing for the spread of electron density over several atoms, explaining properties that a single Lewis structure cannot portray. The resulting structure for O\(_3\) is then understood as a hybrid of these resonance forms, rather than a depiction based on just one form.