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5) Calculate the formal charges in both of the Lewis structures depicted below and explain why the left structure is how sulfate is typically represented, even though it violates the octet rule. :0: :ö: I| : :O:

5) Calculate the formal charges in both of the Lewis structures depicted below and explain why the left structure is how sulfate is typically represented, even though it violates the octet rule. :0: :ö: I| : :O:

Chemistry
Chemistry
10th Edition
ISBN:9781305957404
Author:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste
Publisher:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste
Chapter1: Chemical Foundations
Section: Chapter Questions
Problem 1RQ: Define and explain the differences between the following terms. a. law and theory b. theory and...
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**Task:** 5) Calculate the formal charges in both of the Lewis structures depicted below and explain why the left structure is how sulfate is typically represented, even though it violates the octet rule. **Explanation of Diagrams:** The image shows two Lewis structures for the sulfate ion (\( \text{SO}_4^{2-} \)): 1. **Left Structure:** - The sulfur atom (\( \text{S} \)) is at the center and is double-bonded to two oxygen atoms and single-bonded to two other oxygen atoms. - Each oxygen atom has two lone pairs of electrons, except for the oxygen atoms in double bonds, which have two lone pairs. - This structure does not obey the octet rule for sulfur as it exceeds eight electrons; however, it is a common representation due to resonance stability. 2. **Right Structure:** - The sulfur atom is at the center and is single-bonded to all four oxygen atoms. - Each oxygen atom carries three lone pairs of electrons. - This structure obeys the octet rule but presents higher formal charges. **Calculating Formal Charges:** - **Left Structure:** - **Sulfur:** Formal charge = \( 6 - (4 + 2) = 0 \) - **Double-bonded Oxygen:** Formal charge = \( 6 - (4 + 2) = 0 \) - **Single-bonded Oxygen:** Formal charge = \( 6 - (6 + 1) = -1 \) (each of the two is -1) - **Right Structure:** - **Sulfur:** Formal charge = \( 6 - (0 + 8) = -2 \) - **Oxygen (all single-bonded):** Formal charge = \( 6 - (6 + 1) = -1 \) (each of four) **Conclusion:** The left structure is typically used to represent \(\text{SO}_4^{2-}\) as it resonates among several structures, minimizing formal charge despite violating the octet rule for sulfur. This system of delocalized electrons provides a more stable configuration overall.
Transcribed Image Text:**Task:** 5) Calculate the formal charges in both of the Lewis structures depicted below and explain why the left structure is how sulfate is typically represented, even though it violates the octet rule. **Explanation of Diagrams:** The image shows two Lewis structures for the sulfate ion (\( \text{SO}_4^{2-} \)): 1. **Left Structure:** - The sulfur atom (\( \text{S} \)) is at the center and is double-bonded to two oxygen atoms and single-bonded to two other oxygen atoms. - Each oxygen atom has two lone pairs of electrons, except for the oxygen atoms in double bonds, which have two lone pairs. - This structure does not obey the octet rule for sulfur as it exceeds eight electrons; however, it is a common representation due to resonance stability. 2. **Right Structure:** - The sulfur atom is at the center and is single-bonded to all four oxygen atoms. - Each oxygen atom carries three lone pairs of electrons. - This structure obeys the octet rule but presents higher formal charges. **Calculating Formal Charges:** - **Left Structure:** - **Sulfur:** Formal charge = \( 6 - (4 + 2) = 0 \) - **Double-bonded Oxygen:** Formal charge = \( 6 - (4 + 2) = 0 \) - **Single-bonded Oxygen:** Formal charge = \( 6 - (6 + 1) = -1 \) (each of the two is -1) - **Right Structure:** - **Sulfur:** Formal charge = \( 6 - (0 + 8) = -2 \) - **Oxygen (all single-bonded):** Formal charge = \( 6 - (6 + 1) = -1 \) (each of four) **Conclusion:** The left structure is typically used to represent \(\text{SO}_4^{2-}\) as it resonates among several structures, minimizing formal charge despite violating the octet rule for sulfur. This system of delocalized electrons provides a more stable configuration overall.
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