Question

Predict whether each of the following molecules is polar or nonpolar: (a) \(\mathrm{CCl}_{4}\), (b) \(\mathrm{NH}_{3}\), (c) \(\mathrm{SF}_{4}\), (d) \(\mathrm{XeF}_{4}\), (e) \(\mathrm{CH}_{3} \mathrm{Br}\), (f) \(\mathrm{GaH}_{3}\)

Short answer

Based on the molecular geometry and electronegativity of the atoms involved, the molecules can be classified as: (a) \(CCl_4\): Nonpolar (b) \(NH_3\): Polar (c) \(SF_4\): Polar (d) \(XeF_4\): Nonpolar (e) \(CH_3Br\): Polar (f) \(GaH_3\): Nonpolar

Step-by-step solution

Determine the molecular geometry of each molecule

Using VSEPR theory, determine the molecular geometry of each molecule based on the arrangements of their electron pairs: (a) CCl4: Tetrahedral (b) NH3: Trigonal pyramidal (c) SF4: See-saw (d) XeF4: Square planar (e) CH3Br: Tetrahedral (f) GaH3: Trigonal planar

Check the electronegativity difference between the atoms

Check the electronegativity difference between the atoms in each molecule to determine if they have any dipole moments: (a) CCl4: C-Cl bond is polar due to the electronegativity difference. (b) NH3: N-H bond is polar due to the electronegativity difference. (c) SF4: S-F bond is polar due to the electronegativity difference. (d) XeF4: Xe-F bond is polar due to the electronegativity difference. (e) CH3Br: C-H bond is nonpolar, but C-Br bond is polar due to the electronegativity difference. (f) GaH3: Ga-H bond is polar due to the electronegativity difference.

Evaluate if the dipole moments cancel out or not

Evaluate whether the dipole moments from the polar bonds in the molecules cancel each other out or not: (a) CCl4: Tetrahedral geometry and equal electronegativity around carbon cause all the dipole moments to cancel out, creating a nonpolar molecule. (b) NH3: Trigonal pyramidal geometry and unequal electronegativity around nitrogen result in a net dipole moment, creating a polar molecule. (c) SF4: See-saw geometry and unequal electronegativity around sulfur result in a net dipole moment, creating a polar molecule. (d) XeF4: Square planar geometry and equal electronegativity around xenon cause all the dipole moments to cancel out, creating a nonpolar molecule. (e) CH3Br: Tetrahedral geometry with a polar bond and three nonpolar bonds create an unequal distribution of charge, resulting in a net dipole moment, creating a polar molecule. (f) GaH3: Trigonal planar geometry and equal electronegativity around gallium cause all the dipole moments to cancel out, creating a nonpolar molecule.

Final results:

Based on the molecular geometry and electronegativity of the atoms involved, the molecules can be classified as: (a) CCl4: Nonpolar (b) NH3: Polar (c) SF4: Polar (d) XeF4: Nonpolar (e) CH3Br: Polar (f) GaH3: Nonpolar

Key concepts

VSEPR Theory

Understanding the shape of molecules is crucial in determining many of their physical and chemical properties, including polarity. The Valence Shell Electron Pair Repulsion (VSEPR) theory is a model that helps us predict the three-dimensional arrangement of atoms in a molecule.

VSEPR theory is based on the principle that electron pairs around a central atom will arrange themselves as far apart as possible to minimize repulsion between their negative charges. This includes both bonding pairs (which are shared with other atoms) and lone pairs (which are not shared). For example, a molecule with four bonding pairs, like methane (CH₄), adopts a tetrahedral shape, where the angle between the bonds is approximately 109.5 degrees.

The resulting geometry is essential in determining whether a molecule is polar or nonpolar. A nonpolar molecule is one where the charge distribution is symmetric, while a polar molecule has an asymmetric charge distribution, leading to a net dipole moment. With VSEPR theory, students can predict a molecule's shape and thus get a better sense of its polarity.

Electronegativity Difference

Electronegativity refers to the ability of an atom to attract bonding electrons towards itself. When two atoms form a bond, the difference in their electronegativities determines the bond's character. To determine if a bond is polar or nonpolar, one must look at this electronegativity difference.

A large electronegativity difference usually implies a polar bond, because the electrons are more attracted to one atom than the other, creating a dipole. This is the case in molecules like hydrochloric acid (HCl), where chlorine is much more electronegative than hydrogen.

In the exercise provided, for instance, the C-Cl bond in CCl₄ is polar because of the electronegativity difference between carbon and chlorine. However, when looking at the molecule as a whole, its symmetry leads to the dipoles canceling each other out, resulting in a nonpolar molecule. Understanding the concept of electronegativity difference is therefore essential for identifying the polarity of individual bonds and, in combination with VSEPR theory, the polarity of the entire molecule.

Dipole Moments

Dipole moments are quantitative measures of the polarity of a molecule. They arise from differences in electronegativity and result in molecules having a partial positive end and a partial negative end, just like a magnet with a north and south pole.

The presence of a dipole moment is often a clear indicator that a molecule is polar. For instance, water (H₂O), due to its bent shape and the electronegativity difference between hydrogen and oxygen, has a significant dipole moment and is polar. It is important to note that while individual bonds may be polar, symmetrical molecules may have no net dipole moment because the individual bond dipoles cancel each other out, as demonstrated in the molecules CCl₄ and XeF₄ from the exercise.

When evaluating molecules, we examine whether the vector sum of all dipole moments results in a net dipole. If the geometry allows for the dipoles to cancel out, the molecule is nonpolar. Hence, understanding dipole moments is key for assessing the overall polarity of molecules after analyzing their shapes and electronegativity differences.