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

Electronegativity represents how strongly an atom attracts electrons in a chemical bond. The electronegativity difference between bonded atoms is a critical factor in determining the polarity of the molecule. When two atoms have a substantial difference in their electronegativity, the more electronegative atom will attract the shared electrons more towards itself, resulting in a polar bond.

Consider the molecule IF where iodine (I) and fluorine (F) have significantly different electronegativities; fluorine being much more electronegative than iodine. This difference leads to an unequal sharing of electrons, with the electrons being closer to fluorine, making the IF bond polar. In summary, a large electronegativity difference often points to a polar bond, influencing the overall molecule's polarity.

The shape of a molecule, determined by the positions of its atoms in three-dimensional space, is referred to as its molecular geometry. This geometric arrangement has a direct influence on the molecule's polarity. Even if a molecule contains polar bonds, it might still be nonpolar if its shape is symmetrical.

Take for example carbon disulfide (CS₂), which has a linear molecular geometry. Due to this symmetry, any difference in electronegativity between the sulfur atoms and the central carbon atom does not result in an overall dipole moment. In molecular geometry, the shape of a molecule plays just as crucial a role as the differences in electronegativity when determining the overall molecular polarity.

When evaluating the polarity of a molecule, symmetry can override individual bond polarities. A symmetrical arrangement of polar bonds can result in a nonpolar molecule because the polarities cancel each other out.

The case of sulfur hexafluoride (SF₆) illustrates this concept. Each sulfur-fluorine bond is indeed polar; however, the octahedral geometry of SF₆ is perfectly symmetrical, which means the individual dipoles of the bonds exactly cancel out, leaving no net dipole moment across the molecule. This perfect cancellation is what makes SF₆ nonpolar, despite its polar bonds. Symmetry is thus a determining factor for molecular polarity.

Chemical bond polarity is the concept that describes how the electrons in a bond are distributed between the two bonded atoms. Bonds can be nonpolar, where electrons are shared equally, or polar, where electrons are shared unequally.

For instance, in phosphorus trichloride (PCl₃), the bonds between the phosphorus (P) and chlorine (Cl) atoms are polar because chlorine is more electronegative than phosphorus, thus attracting the shared electrons more. However, in this molecule, the polar bonds do not cancel out due to the trigonal pyramidal geometry, resulting in an overall polar molecule. Chemical bond polarity, combined with molecular geometry, is instrumental in determining the polarity of a molecule.