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Chapter 6: Problem 29

In terms of an operational definition, when is a molecule considered to be polar?

Short Answer

Expert verified
A molecule is considered to be polar if it has polar bonds due to the difference in electronegativity between the atoms involved, an asymmetric molecular geometry, preventing the cancellation of bond dipoles, and a net dipole moment, indicating an uneven distribution of charges within the molecule.

Step by step solution

01

Understanding Polarity

A molecule is considered polar when its electrons are distributed unevenly, resulting in a net dipole moment, which is a measure of the separation of positive and negative charges in the molecule. This occurs due to the difference in electronegativity between the atoms involved and the molecular geometry.
02

Electronegativity Differences

The electronegativity of an atom is a measure of how strongly it attracts the shared electrons in a chemical bond. When there's a significant difference in electronegativity between two atoms in a bond, the electrons will be more attracted to the atom with higher electronegativity, leading to an unequal distribution of electron density and creating a polar bond.
03

Molecular Geometry

The shape of the molecule plays a critical role in determining its polarity. For a molecule to be polar, it must have an asymmetric distribution of polar bonds, which means that it should not have a symmetrical geometry in which the individual bond dipoles cancel each other out.
04

Net Dipole Moment

Once we have established that a molecule has polar bonds and an asymmetric geometry, we need to examine the net dipole moment of the molecule. A molecule with a net dipole moment is considered polar, as it has an uneven distribution of charges. The net dipole moment can be calculated using the bond dipole moments and the geometry of the molecule. In conclusion, a molecule is considered to be polar if it has the following characteristics: 1. Polar bonds due to the difference in electronegativity between the atoms involved. 2. An asymmetric molecular geometry, preventing the cancellation of bond dipoles. 3. A net dipole moment, indicating an uneven distribution of charges within the molecule.

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Most popular questions from this chapter

Consider the molecule \(\mathrm{BrCl}_{3}\). (a) Draw a dot diagram for the molecule. (b) Predict the molecule's electron group geometry. (c) Predict the molecule's shape (name the shape and draw it using wedges, lines, and broken lines, and indicate the ideal value of the bond angles). (d) Predict if the molecule is polar or nonpolar. If polar, draw the dipole moment for the molecule.

Which should have the largest molecular dipole moment: \(\mathrm{H}_{2}, \mathrm{CO}_{2}, \mathrm{CH}_{3} \mathrm{~F}\), or \(\mathrm{CH}_{7} \mathrm{I}\) ?

Draw a combined Lewis dot, molecular-shape diagram for each of the following species. Name each shape, and indicate whether the molecule or ion has an overall dipole moment. If so, draw the dipole moment vector. (a) \(\mathrm{Cl}_{2} \mathrm{O}\) (b) NOF (c) \(\mathrm{PF}_{3}\) (d) \(\mathrm{ICl}_{2}^{+}\) (Hint: See problem statement and hint for Problem 6.69. Hint for \((b): \mathrm{N}=\mathrm{O}\) bonds are common.)

Draw the three-dimensional shape of methanol, \(\mathrm{CH}_{3} \mathrm{OH}\). Indicate the numeric value of all bond angles. Is this molecule polar? If so, draw the molecular dipole moment vector.

What is it about the trigonal bipyramidal shape that distinguishes it from all the other molecular shapes that we have covered? (Hint: Think about lone pairs of electrons.)
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