Question

Predict whether each of the following molecules is polar or nonpolar. (a) \(\mathrm{IF},(\mathrm{b}) \mathrm{CS}_{2},(\mathrm{c}) \mathrm{SO}_{3}\) (d) \(\mathrm{PCl}_{3},(\mathrm{e}) \stackrel{\mathrm{SF}_{6}}\) (f) \(\mathbb{F F}_{5}\)

Short answer

(a) IF is polar, (b) CS2 is nonpolar, (c) SO3 is nonpolar, (d) PCl3 is polar, (e) SF6 is nonpolar, and (f) PF5 is nonpolar.

Step-by-step solution

1. Determine molecule's geometry and electronegativity difference

To begin, we need to identify the molecular geometries for each of the given molecules and look for charge distribution. We must also consider the electronegativity difference between the atoms.

(a) IF

This molecule has a linear geometry with iodine in the center and a single bond to fluorine. The electronegativity difference between I and F is significant, and there is no other bond in this linear molecule to cancel out the charge difference. Thus, IF has a net dipole, making it a polar molecule.

(b) CS2

Carbon disulfide (CS2) has a linear molecular geometry with carbon in the center and two sulfur atoms on either side. The electronegativity difference between the carbon and sulfur atoms is very small, which means there is no significant charge distribution across the molecule. Thus, CS2 has no net dipole and is a nonpolar molecule.

(c) SO3

Sulfur trioxide (SO3) has a trigonal planar molecular geometry with sulfur in the center, surrounded by three oxygen atoms. Sulfur and oxygen have an electronegativity difference, but the trigonal planar shape means that the polarity of each S-O bond cancels each other out, and there is no net dipole. Thus, SO3 is a nonpolar molecule.

(d) PCl3

Phosphorus trichloride (PCl3) has a trigonal pyramidal molecular geometry with phosphorus in the center and three chlorine atoms. There is a significant electronegativity difference between phosphorus and chlorine, and because of the structure and bond angles, the polarity doesn't cancel out. Therefore, PCl3 has a net dipole, and it is a polar molecule.

(e) SF6

Sulfur hexafluoride (SF6) has an octahedral molecular geometry with sulfur in the center and six fluorine atoms on the corners. While there is an electronegativity difference between sulfur and fluorine, the octahedral shape means that any charge distribution is symmetrically cancelled out, and there is no net dipole. Thus, SF6 is a nonpolar molecule.

(f) PF5

Phosphorus pentafluoride (PF5) has a trigonal bipyramidal molecular geometry with phosphorus in the center and five fluorine atoms. The electronegativity difference between phosphorus and fluorine leads to charge distribution, but because of the trigonal bipyramidal geometry, the polarities cancel each other out, leaving no net dipole. Thus, PF5 is a nonpolar molecule.

Key concepts

Electronegativity

Electronegativity is a key concept in understanding molecular polarity. It is a measure of how strongly an atom attracts electrons in a chemical bond. Different elements have different electronegativities, and this affects how electrons are shared between them.
When two atoms form a bond, the atom with higher electronegativity will pull the shared electrons closer, creating a dipole. The difference in electronegativity between atoms can determine bond polarity:

• If the difference is large, the bond is polar.
• If the difference is small or zero, the bond is nonpolar.

This concept is essential in predicting whether a molecule will be polar. For example, in the molecule \( ext{IF}\), iodine and fluorine have a significant electronegativity difference, resulting in a polar bond. By understanding electronegativity, we can anticipate how molecules interact with others, affecting solubility, reactivity, and more.

Molecular Geometry

Molecular geometry is the three-dimensional arrangement of atoms in a molecule. It dictates the spatial distribution of bonds, helping us understand the shape and bond angles within a molecule.
This arrangement is vital in determining molecular polarity, where:

• A symmetric geometry usually leads to nonpolar molecules.
• An asymmetric geometry often results in polar molecules.

Taking \( ext{CS}_2\) as an example, its linear geometry causes any potential dipoles from C-S bonds to cancel out, leaving no net dipole, hence nonpolar. Similarly, \( ext{SO}_3\) is trigonal planar, and although individual S-O bonds are polar, their symmetry cancels out the overall dipole moment, leading to a nonpolar molecule.

Dipole Moment

The dipole moment is a vector quantity that represents the separation of charges within a molecule. It shows both the magnitude and direction of polarity. A molecule with a non-zero dipole moment is polar.

Understanding Dipole Moment:

• It occurs when there is an uneven distribution of electrons.
• Its unit is Debye (D), with a greater value indicating higher polarity.

Considering \( ext{PCl}_3\), the presence of a dipole moment is due to the significant electronegativity difference and the molecular structure that does not allow the dipoles to cancel each other out. By analyzing dipole moments, chemists can predict the behavior of molecules in electric fields, their interaction with other molecules, and their solubility in water.

Trigonometry of Molecules

To fully understand molecular geometry, knowledge of trigonometry is crucial. It helps describe the angles and spatial orientation of bonds within a molecule.
Molecules often have geometries defined by specific angles, dictated by the hybridization of the bonding orbitals:

• Linear geometry features a 180° bond angle.
• Trigonal planar geometry has bond angles of 120°.
• Tetrahedral geometry is characterized by 109.5° angles.

The use of trigonometry becomes important when analyzing molecules like \( ext{SF}_6\), which adopts an octahedral geometry with 90° angles. Despite the electronegativity difference between sulfur and fluorine, the symmetrical octahedral shape ensures that these angles lead to cancellation of any dipole moments, rendering the molecule nonpolar.