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Ch. 2 - Acids and Bases; Functional Groups
Wade - Organic Chemistry 9th Edition
Wade9th EditionOrganic ChemistryISBN: 9780135213728Not the one you use?Change textbook
All textbooksWade 9th EditionCh. 2 - Acids and Bases; Functional GroupsProblem 1a
Chapter 2, Problem 1a

The C=O double bond has a dipole moment of about 2.4 D and a bond length of about 1.23 Å.
a. Calculate the amount of charge separation in this bond.

Verified step by step guidance
1
Understand the relationship between dipole moment (μ), charge (q), and distance (d) using the formula: μ = q × d. Here, μ is given as 2.4 D (Debye) and d is given as 1.23 Å (angstroms).
Convert the dipole moment from Debye to coulomb-meters. Recall that 1 D = 3.33564 × 10^-30 C·m.
Convert the bond length from angstroms to meters. Recall that 1 Å = 1 × 10^-10 m.
Rearrange the formula to solve for the charge (q): q = μ / d.
Substitute the converted values of μ and d into the equation to calculate the charge separation in coulombs.

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Key Concepts

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

Dipole Moment

The dipole moment is a measure of the separation of positive and negative charges in a molecule. It is a vector quantity, often expressed in Debye units (D), and is calculated as the product of the charge magnitude and the distance between the charges. In the context of a C=O bond, the dipole moment indicates the polarity of the bond due to the difference in electronegativity between carbon and oxygen.
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Bond Length

Bond length is the average distance between the nuclei of two bonded atoms. It is typically measured in angstroms (Å) and is influenced by the size of the atoms and the bond order. In a C=O double bond, the bond length is shorter than a single bond due to the increased electron sharing, which pulls the atoms closer together, affecting the dipole moment calculation.
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Charge Separation

Charge separation in a bond refers to the distribution of electron density between two atoms, leading to partial positive and negative charges. This concept is crucial for calculating the dipole moment, as it involves determining the effective charge difference across the bond. In a C=O bond, the charge separation arises from the higher electronegativity of oxygen compared to carbon, resulting in a partial negative charge on oxygen and a partial positive charge on carbon.
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Related Practice
Textbook Question

The C=O double bond has a dipole moment of about 2.4 D and a bond length of about 1.23 Å.

b. Use this information to evaluate the relative importance of the following two resonance contributors:

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Textbook Question

The N—F bond is more polar than the N—H bond, but NF3 has a smaller dipole moment than NH3. Explain this curious result.

NF3

μ= 0.2 D

NH3

μ = 1.5 D

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Textbook Question

1. Draw the Lewis structure.

2. Show how the bond dipole moments (and those of any nonbonding pairs of electrons) contribute to the molecular dipole moment.

3. Estimate whether the compound will have a large, small, or zero dipole moment.

d. CH3F

e. CF4

f. CH3OH

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