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Ch. 1 - Remembering General Chemistry: Electronic Structure and Bonding (Part 1)
Bruice - Organic Chemistry 8th Edition
Bruice8th EditionOrganic ChemistryISBN: 9780135213711Not the one you use?Change textbook
All textbooksBruice 8th EditionCh. 1 - Remembering General Chemistry: Electronic Structure and Bonding (Part 1)Problem 16c,d
Chapter 1, Problem 16c,d

Give each atom the appropriate formal charge:
c.
d.

Verified step by step guidance
1
Step 1: Recall the formula for calculating formal charge: Formal charge = (Valence electrons) - (Non-bonding electrons) - (1/2 × Bonding electrons). This formula will be applied to each atom in the given structures.
Step 2: For the first structure (tetramethylammonium ion), focus on the nitrogen atom. Nitrogen typically has 5 valence electrons. In this structure, nitrogen is bonded to four methyl groups, forming four single bonds. Since nitrogen has no lone pairs and is sharing 8 bonding electrons, calculate its formal charge using the formula.
Step 3: For the first structure, examine the carbon atoms in the methyl groups. Carbon typically has 4 valence electrons. Each carbon is bonded to three hydrogens and one nitrogen, forming four single bonds. Since there are no lone pairs on these carbons, calculate their formal charges.
Step 4: For the second structure (ammonia-borane complex), focus on the nitrogen atom. Nitrogen has 5 valence electrons. In this structure, nitrogen is bonded to three hydrogens and one boron atom. Nitrogen also has a lone pair of electrons. Use the formal charge formula to determine its charge.
Step 5: For the second structure, examine the boron atom. Boron typically has 3 valence electrons. In this structure, boron is bonded to three hydrogens and one nitrogen atom. Boron has no lone pairs, and it is sharing 8 bonding electrons. Use the formal charge formula to determine its charge.

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Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Formal Charge

Formal charge is a theoretical charge assigned to an atom in a molecule, calculated based on the number of valence electrons, the number of bonds, and the number of non-bonding electrons. It helps in determining the most stable structure of a molecule by indicating how electrons are distributed among atoms. The formula for calculating formal charge is: Formal Charge = Valence Electrons - (Non-bonding Electrons + 1/2 Bonding Electrons).
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Lewis Structures

Lewis structures are diagrams that represent the bonding between atoms in a molecule and the lone pairs of electrons that may exist. They provide a visual representation of the molecule's structure, allowing for the identification of formal charges and the overall electron distribution. Understanding how to draw and interpret Lewis structures is essential for predicting molecular geometry and reactivity.
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Hybridization

Hybridization is the concept of mixing atomic orbitals to form new hybrid orbitals that can accommodate the bonding requirements of a molecule. In the case of nitrogen in the given structure, it typically undergoes sp³ hybridization, resulting in four equivalent hybrid orbitals that form sigma bonds with surrounding atoms. Recognizing hybridization is crucial for understanding molecular shape and bond angles.
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Related Practice
Textbook Question

An atom with a formal charge does not necessarily have more or less electron density than the atoms in the molecule without formal charges. We can see this by examining the potential maps for H2O, H3O+, and HO−.

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a. Which atom bears the formal negative charge in the hydroxide ion?

b. Which atom has the greater electron density in the hydroxide ion?

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Textbook Question

An atom with a formal charge does not necessarily have more or less electron density than the atoms in the molecule without formal charges. We can see this by examining the potential maps for H2O, H3O+, and HO−.

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c. Which atom bears the formal positive charge in the hydronium ion?

d. Which atom has the least electron density in the hydronium ion?

Textbook Question

Which of the atoms in the molecular models in [Problem 20] have

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a. three lone pairs?

b. two lone pairs?

c. one lone pair?

d. no lone pairs?

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Textbook Question

Draw the Lewis structure for each of the following:

e. CH3N+H3

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Textbook Question

Give each atom the appropriate formal charge:

a.

b.

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Textbook Question

Draw the Lewis structure for each of the following:

d. +C2H5

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