Determine whether each molecule is polar or nonpolar. Lewis structures and determine molecular geometry to answer this question correctly a. PBr5 b. SO2 c. SF6
The Correct Answer and Explanation is:
To determine whether a molecule is polar or nonpolar, we need to consider its Lewis structure, molecular geometry, and the electronegativity differences between atoms.
a. PBr₅ (Phosphorus pentabromide)
- Lewis Structure: Phosphorus (P) is the central atom with five bromine (Br) atoms bonded to it. It has 10 electrons in its valence shell, which is allowed for elements in period 3 or beyond.
- Molecular Geometry: Trigonal bipyramidal. Three Br atoms are in the equatorial plane and two are in the axial positions.
- Polarity: Nonpolar. Although P–Br bonds are polar due to the electronegativity difference, the symmetry of the trigonal bipyramidal geometry causes the dipoles to cancel out. The overall molecule has no net dipole moment.
b. SO₂ (Sulfur dioxide)
- Lewis Structure: Sulfur (S) is bonded to two oxygen (O) atoms. There is one lone pair on the sulfur, and both S–O bonds are double bonds or have resonance.
- Molecular Geometry: Bent (or V-shaped) with an angle of about 120 degrees.
- Polarity: Polar. The molecule has an asymmetric shape due to the lone pair on sulfur, causing a net dipole moment. The polar bonds do not cancel out, making the whole molecule polar.
c. SF₆ (Sulfur hexafluoride)
- Lewis Structure: Sulfur (S) is the central atom bonded to six fluorine (F) atoms. Sulfur uses an expanded octet with 12 valence electrons.
- Molecular Geometry: Octahedral. The six F atoms are placed symmetrically around the sulfur atom.
- Polarity: Nonpolar. Despite the polar S–F bonds, the octahedral shape distributes the dipoles evenly, canceling them out. The molecule has no net dipole moment.
Summary:
- PBr₅ – Nonpolar
- SO₂ – Polar
- SF₆ – Nonpolar
Molecular geometry is crucial for determining polarity. Even if individual bonds are polar, the overall shape can lead to a nonpolar molecule when the bond dipoles cancel out.
