Based on formal charges, draw the most preferred Lewis structure for the chlorate ion, Clio, Chemistry can anyone please draw correctly Based on formal charges, draw the most preferred Lewis structure for the chlorate ion, ClO3. To add lone pairs, click the button before clicking on the molecule. To add bonds connect atoms with a line 7 To add formal charges, click the or button before clicking on the molecule or brackets. Draw the molecule by placing atoms on the grid and connecting them with bonds. Include lone pairs of electrons. Overall charge may be specified with square brackets or formal charges. â–º View Available Hint(s) ANNIE A :O: + 😀 0: : 0:
The Correct Answer and Explanation is :
The most preferred Lewis structure for the chlorate ion (ClO₃⁻) can be determined by considering the formal charges and the octet rule. Here is how we approach drawing the structure:
Step-by-Step Process:
- Count the total valence electrons:
- Chlorine (Cl) is in Group 7A (halogens), so it contributes 7 valence electrons.
- Each Oxygen (O) atom is in Group 6A, so each oxygen contributes 6 valence electrons.
- The chlorate ion has a negative charge (ClO₃⁻), which means we add 1 extra electron to the total valence count. Total valence electrons = 7 (from Cl) + 3 × 6 (from O) + 1 (extra electron due to the negative charge) = 26 valence electrons.
- Arrange the atoms:
- The chlorine atom is the central atom because it is less electronegative than oxygen. Place the chlorine atom in the center and connect it to the three oxygen atoms with single bonds.
- Distribute electrons to complete the octet:
- Each single bond between chlorine and oxygen consists of 2 electrons, so 3 bonds use up 6 electrons. After placing these bonds, we have 26 – 6 = 20 electrons left to place.
- Distribute the remaining electrons as lone pairs on the oxygen atoms to complete their octets. Each oxygen will get 6 more electrons in the form of lone pairs (since each oxygen atom needs 2 more electrons to complete its octet).
- Double check formal charges:
- To minimize the formal charges, we want to make sure the atoms hold the least possible formal charge. Formal charge (FC) is calculated using the formula:
[
FC = \text{Valence electrons} – \left( \text{Lone pair electrons} + \frac{1}{2} \times \text{Bonding electrons} \right)
] - If all oxygen atoms are surrounded by lone pairs and single bonds, and chlorine is surrounded by single bonds, each oxygen will have a formal charge of -1, and chlorine will have a formal charge of +1. This satisfies the overall charge of -1 for the ion.
- Final Lewis structure:
- Chlorine is the central atom with three single bonds to three oxygen atoms.
- Each oxygen atom has three lone pairs of electrons (except for one oxygen that may form a double bond to chlorine, if needed).
- The overall charge of -1 is placed in square brackets around the structure.
Thus, the most stable Lewis structure for ClO₃⁻ looks like this:
- Chlorine (Cl) in the center with three bonds to oxygen atoms.
- One oxygen with a double bond to chlorine and the other two oxygen atoms with single bonds.
- Formal charges: Chlorine (+1), two oxygens with single bonds (-1), one oxygen with a double bond (neutral).
- Overall charge: -1.
Explanation:
- The structure follows the octet rule for oxygen atoms and minimizes formal charges.
- Chlorine can exceed the octet rule since it is in period 3 and has available d orbitals for bonding. This makes it possible for chlorine to bond with three oxygens.
- The negative charge is best placed on the oxygens because oxygen is more electronegative than chlorine.
This structure minimizes the formal charges and satisfies the octet rule for all atoms, making it the most stable configuration for the chlorate ion (ClO₃⁻).