Question

Examine the three dimensional structures of each of the following molecules in the simulation, which can be found in the Real Molecules mode. Then, identify which molecules are polar and which are nonpolar. Assume that every bond in each molecule is polar covalent.

ClF₃, SF₄, VeF₂, XeF₄, SO₂, SF₆, NH₃, PCl₅.

Answer 1

To determine whether molecules are polar or nonpolar, one must analyze their molecular geometry and the distribution of electron density. Molecules such as SF6 and XeF4 are nonpolar due to their symmetric structure, while NH3 and SO2 are polar because of asymmetry and lone pairs creating a net dipole.

Step-by-step explanation:

To determine whether the molecules ClF₃, SF₄, VF₂, XeF₄, SO₂, SF₆, NH₃, and PCl₅ are polar or nonpolar, one must take into account the molecule's symmetry and the presence of different electronegativities among bonded atoms. Molecular polarity depends on both of these factors. For instance, molecules like SF₆ and XeF₄ are nonpolar because they have a symmetrical geometrical structure where the bond dipole moments cancel out. However, molecules like NH₃ and SO₂ are polar due to their asymmetric shapes and the presence of lone pairs that contribute to a net dipole moment.

To identify which molecules are polar, one should:

1. Draw the Lewis structure.
2. Determine the geometry using VSEPR theory.
3. Assess if there is an asymmetric distribution of electron density.

Drawing upon the principles of VSEPR theory and the concept of electronegativity, molecules like NH₃ (ammonia) with a trigonal pyramidal shape and a lone pair on the nitrogen atom are definitely polar. In contrast, SF₆ (sulfur hexafluoride) with its octahedral shape and equal bonding to sulfur, is nonpolar. While VF₂ is not a commonly recognized molecule and may represent a typo, if it were to be VF₄ (vanadium tetrafluoride) with a square planar geometry, it would be nonpolar.