Question

Which of the following molecules is(are) not polar? For each polar molecule, indicate the direction of polarity-that is, which is the negative end and which is the positive end. (a) \(\mathrm{BCl}_{3}\) (b) \(\mathrm{CF}_{4}\) (c) \(\mathrm{PCl}_{3}\) (d) \(\operatorname{GeH}_{4}\)

Short answer

Nonpolar molecules: (a) \(\mathrm{BCl}_{3}\), (b) \(\mathrm{CF}_{4}\), (d) \(\mathrm{GeH}_{4}\). Polar molecule: (c) \(\mathrm{PCl}_{3}\), negative end at chlorine, positive end at phosphorus.

Step-by-step solution

Understand Molecular Geometry

Determine the molecular geometry of each molecule. - (a) \(\mathrm{BCl}_{3}\): Trigonal planar.- (b) \(\mathrm{CF}_{4}\): Tetrahedral. - (c) \(\mathrm{PCl}_{3}\): Trigonal pyramidal.- (d) \(\mathrm{GeH}_{4}\): Tetrahedral.

Assess Electronegativity Differences

Evaluate the difference in electronegativity between the central atom and surrounding atoms:- (a) Boron (\(\mathrm{B}\)) and Chlorine(\(\mathrm{Cl}\)). Chlorine is more electronegative than boron.- (b) Carbon (\(\mathrm{C}\)) and Fluorine(\(\mathrm{F}\)). Fluorine is more electronegative than carbon.- (c) Phosphorus (\(\mathrm{P}\)) and Chlorine (\(\mathrm{Cl}\)). Chlorine is more electronegative than phosphorus.- (d) Germanium (\(\mathrm{Ge}\)) and Hydrogen (\(\mathrm{H}\)). Germanium is slightly more electronegative, but the difference is small.

Determine Symmetry and Polarity

Determine if the molecule is symmetrical or if there is a net dipole moment:- (a) \(\mathrm{BCl}_{3}\): Symmetrical, dipoles cancel out.- (b) \(\mathrm{CF}_{4}\): Symmetrical, dipoles cancel out.- (c) \(\mathrm{PCl}_{3}\): Not symmetrical, net dipole moment present.- (d) \(\mathrm{GeH}_{4}\): Symmetrical, dipoles cancel out.

Identify Polar and Nonpolar Molecules

Identify which molecules are polar and which are nonpolar:- (a) \(\mathrm{BCl}_{3}\): Nonpolar.- (b) \(\mathrm{CF}_{4}\): Nonpolar.- (c) \(\mathrm{PCl}_{3}\): Polar; Negative end towards chlorine, positive towards phosphorus.- (d) \(\mathrm{GeH}_{4}\): Nonpolar.

Key concepts

Molecular Geometry

Molecular geometry is all about the three-dimensional arrangement of atoms in a molecule. This concept is key to predicting how molecules will interact and what their properties might be. For example:

• Localized around a central atom, common geometries include linear, bent, trigonal planar, tetrahedral, and trigonal pyramidal.
• In the exercise, for instance,
  • Boron trichloride ( \( \text{BCl}_3 \) ) presents a trigonal planar geometry.
  • Carbon tetrafluoride ( \( \text{CF}_4 \) ) and germanium tetrahydride ( \( \text{GeH}_4 \) ) both exhibit tetrahedral geometries.
  • Phosphorus trichloride ( \( \text{PCl}_3 \) ) is shaped as a trigonal pyramid.

Understanding the geometry helps in identifying the potential polarity of the molecule since the spatial arrangement influences how dipole moments might cancel out or add up.

Electronegativity Differences

Electronegativity refers to the tendency of an atom in a molecule to attract electrons to itself. Differences in electronegativity between atoms determine the distribution of charge:

• For molecules, if one atom is significantly more electronegative than another, it pulls electrons towards itself, forming a dipole.
• Consider these pairs:
  • In \( \text{BCl}_3 \) and \( \text{PCl}_3 \) , chlorine with high electronegativity attracts electrons more than boron or phosphorus.
  • In \( \text{CF}_4 \) , fluorine is more electronegative than carbon.
  • For \( \text{GeH}_4 \) , although germanium is slightly more electronegative than hydrogen, the difference is minimal.

These differences help predict the polarity of individual bonds within the molecule.

Symmetry in Molecules

Symmetry in a molecule is crucial for determining its overall polarity. A symmetrical molecule often has nonpolar characteristics because any dipoles are balanced out:

• In symmetrical molecules, individual dipoles from polar bonds cancel each other out.
• For instance:
  • The molecules \( \text{BCl}_3 \) , \( \text{CF}_4 \) , and \( \text{GeH}_4 \) have symmetrical structures, causing any individual bond dipoles to cancel out and result in nonpolar molecules.
  • In contrast, \( \text{PCl}_3 \) lacks symmetry due to the trigonal pyramidal shape and thus has a net dipole moment, making it polar. Here, the bonding pair does not equally cancel out since the lone pair disturbs the symmetry.

Thus, recognizing and analyzing the symmetry in a molecule helps in understanding if a molecule will behave as polar or nonpolar.