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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° 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’s 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

Which molecule in each pair has the greater dipole moment? Give the reason for your choice. (a) \(\mathrm{SO}_{2}\) or \(\mathrm{SO}_{3}\) (b) ICl or IF (c) \(\mathrm{SiF}_{4}\) or \(\mathrm{SF}_{4}\) (d) \(\mathrm{H}_{2} \mathrm{O}\) or \(\mathrm{H}_{2} \mathrm{~S}\)

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

Ethylene, \(\mathrm{C}_{2} \mathrm{H}_{4},\) and tetrafluoroethylene, \(\mathrm{C}_{2} \mathrm{~F}_{4},\) are used to make the polymers polyethylene and polytetrafluoroethylene (Teflon), respectively. (a) Draw the Lewis structures for \(\mathrm{C}_{2} \mathrm{H}_{4}\) and \(\mathrm{C}_{2} \mathrm{~F}_{4},\) and give the ideal \(\mathrm{H}-\mathrm{C}-\mathrm{H}\) and \(\mathrm{F}-\mathrm{C}-\mathrm{F}\) bond angles. (b) The actual \(\mathrm{H}-\mathrm{C}-\mathrm{H}\) and \(\mathrm{F}-\mathrm{C}-\mathrm{F}\) bond angles are \(117.4^{\circ}\) and \(112.4^{\circ},\) respectively. Explain these deviations.

Draw the Lewis structure with lowest formal charges, and determine the charge of each atom in (a) \(\mathrm{CN}^{-} ;\) (b) \(\mathrm{ClO}^{-}\).

Phosphorus pentachloride, a key industrial compound with annual world production of about \(2 \times 10^{7} \mathrm{~kg},\) is used to make other compounds. It reacts with sulfur dioxide to produce phosphorus oxychloride \(\left(\mathrm{POCl}_{3}\right)\) and thionyl chloride \(\left(\mathrm{SOCl}_{2}\right)\). Draw a Lewis structure and name the molecular shape of each of these products.
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