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Chapter 10: Problem 53

How can a molecule with polar covalent bonds not be polar? Give an example.

Short Answer

Expert verified
A molecule can be nonpolar if its shape allows the dipole moments from polar bonds to cancel out, like in CO2.

Step by step solution

01

- Understand Polar Covalent Bonds

A polar covalent bond occurs when two atoms share electrons unequally, causing a dipole moment. One atom becomes partially negative (due to gaining more electron density), and the other becomes partially positive.
02

- Consider Molecular Geometry

The overall polarity of a molecule depends not just on individual polar covalent bonds but also on the molecule's geometry. The shape determines if the dipoles cancel out.
03

- Analyze Dipole Moments

If the molecule's shape allows the dipole moments to cancel each other, the molecule can be nonpolar. For example, if the molecule is symmetrical, the dipoles might point in opposite directions and cancel.
04

- Example of Carbon Dioxide (CO2)

Carbon dioxide (CO2) is a perfect example. It has two polar covalent bonds (C=O) but is linear in shape. The opposite dipoles cancel out due to the 180掳 bond angle, making CO2 a nonpolar molecule.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Polar Covalent Bond
A polar covalent bond forms when two atoms share electrons unevenly.
One atom pulls the shared electrons closer, becoming partially negative.
The other atom becomes partially positive due to less electron density.
For example, in a water molecule (H2O), oxygen holds electrons more tightly than hydrogen.
This creates a polar bond with a negative charge on the oxygen and a positive charge on the hydrogen.
Such bonds are common in molecules where atoms have different electronegativities.
Molecular Geometry
The shape of a molecule also impacts its overall polarity.
Geometry determines how individual dipole moments (created by polar covalent bonds) align within the molecule.
If the shape is symmetrical, these dipole moments may cancel out.
Leading to a nonpolar molecule despite having polar covalent bonds.
For example, carbon dioxide (CO2) has a linear geometry. Despite having polar bonds, the dipoles cancel each other due to symmetry.
Understanding the molecular geometry is key in predicting the polarity of complex molecules.
Dipole Moment
A dipole moment is a measure of the separation of positive and negative charges in a molecule.
It depends on both individual bond polarities and the geometry of the molecule.
Mathematically, it is the product of the charge difference and the distance between the charges.
Molecules with high dipole moments are usually polar.
However, if the dipole moments in a molecule cancel each other out, the molecule is nonpolar.
For instance, despite water's bent shape leading to a significant dipole moment, carbon dioxide鈥檚 linear structure results in no net dipole moment.
Nonpolar Molecules
Nonpolar molecules have no net dipole moment. This happens when the molecule's geometry allows polar bond dipoles to cancel out.
Symmetry is often a key factor. For instance, carbon dioxide (CO2) is nonpolar, thanks to its linear shape.
Even though each C=O bond is polar, they cancel because they point in opposite directions.
Nonpolar molecules are often gases or have low boiling and melting points.
Understanding why molecules are nonpolar helps in fields like chemistry and material science.

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Most popular questions from this chapter

Arrange the following \(\mathrm{ACl}_{n}\) species in order of decreasing \(\mathrm{Cl}-\mathrm{A}-\mathrm{Cl}\) bond angles: \(\mathrm{SCl}_{2}, \mathrm{OCl}_{2}, \mathrm{PCl}_{3}, \mathrm{SiCl}_{4}, \mathrm{SiCl}_{6}^{2-}\)

Determine the electron-group arrangement, molecular shape, and ideal bond angle(s) for each of the following: (a) \(\mathrm{CO}_{3}^{2-}\) (b) \(\mathrm{SO}_{2}\) (c) \(\mathrm{CF}_{4}\)

There are three different dichloroethylenes (molecular formula \(\mathrm{C}_{2} \mathrm{H}_{2} \mathrm{Cl}_{2}\) ), which we can designate \(\mathrm{X}, \mathrm{Y},\) and \(\mathrm{Z}\). Compound X has no dipole moment, but compound \(Z\) does. Compounds \(X\) and \(Z\) each combine with hydrogen to give the same product: $$ \mathrm{C}_{2} \mathrm{H}_{2} \mathrm{Cl}_{2}(\mathrm{X} \text { or } \mathrm{Z})+\mathrm{H}_{2} \longrightarrow \mathrm{ClCH}_{2}-\mathrm{CH}_{2} \mathrm{Cl} $$ What are the structures of \(\mathrm{X}, \mathrm{Y},\) and \(\mathrm{Z} ?\) Would you expect compound \(Y\) to have a dipole moment?

Despite many attempts, the perbromate ion \(\left(\mathrm{BrO}_{4}^{-}\right)\) was not prepared in the laboratory until about \(1970 .\) (In fact, articles were published explaining theoretically why it could never be prepared!) Draw a Lewis structure for \(\mathrm{BrO}_{4}^{-}\) in which all atoms have lowest formal charges.

Draw a Lewis structure for (a) \(\mathrm{PH}_{4}^{+} ;\) (b) \(\mathrm{C}_{2} \mathrm{~F}_{4} ;\) (c) \(\mathrm{SbH}_{3}\).
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