Lecture 24 © slg CHM 151 TOPICS: QUIZ 6 1. Bond and Molecular Polarity

Key, Q 6 Tetrahedral ~109.5 o “bent” Same as BrO 2 - Lewis Structure Shape

Key, Q 6 Trigonal bipyramidal 90, 180 o “T shaped” Same as IF 3 Lewis Structure Shape

Oxidation Numbers (from formula) Formal charge (from Lewis structure)

Oxidation Numbers (from formula) Formal charge (from Lewis structure)

UNIT FIVE Bond and Molecular Polarity (Test 5) Gas Laws (Test 5) Intermolecular Attractions (ACS test only)

We have classified bonds “ionic” and “covalent”, depending on whether electron pairs are shared or electrons are completely transferred from one atom to another. In actuality, there is no sharp dividing line between the two types but rather a continuum: Evenly shared electrons Unevenly shared electrons Transferred electrons To determine where a bond lies in this “continuum”, it is useful to consider the difference in electronegativity ( X) between the two atoms making up the bond:

When the difference ( X) is less than 0.5, sharing is fairly even and electrons are not much closer to one atom than the other. The bonds are considered “non-polar.” When the difference is between 0.5 and about 1.7, the electrons are closer to the more electronegative atom and partial charge buildup, polarization, develops. When ( X) is greater than 1.7 or so, ionic bonding becomes the more likely bond type and valence electrons are transferred to the more electronegative atom.

So, we need to consider a third more specialized type of bond, “the polar covalent bond:” This type of bond will be the important factor to be considered when we look at molecular polarity, which arises from molecular shape and bond polarity. The polar molecular in turn will exhibit different solubilities and boiling points than non polar molecules.

Let us consider the bond between H and Cl in a molecule of hydrogen chloride (only hydrochloric acid when in water!): E pair closer to Cl, more electronegative Orbital between H and Cl

The electron cloud from the pair of shared electrons is more dense closer to the chlorine, and much less dense closer to the hydrogen. The bond has become “polarized”: it has developed a region (or “pole”) of partial positive charge buildup and a region (or “pole”) of partial negative buildup.

Major portion of electron density

Arrow to indicate polar bond, pointing to more (-) atom “partially positive” “partially negative”

The molecule has only one bond, and it is polar. This makes the entire molecule a “dipole”, one which has a positive and negative pole and will align in an electrical or magnetic field: All diatomic molecules with polarized bonding between the two atoms are DIPOLES.

Other examples of diatomic dipoles:

MOLECULAR POLARITY, LARGER MOLECULES All diatomic molecules with polar bond(s) are dipoles, but the situation is not so simple for larger molecules. There are two factors to consider: Are the bonds polar? Are they arranged so that the center of positive charge and the center of negative charge do not “coincide”?

BOND POLARITY MOLECULAR POLARITY

Center of +,- charges coincide, center of molecule

POLAR BONDS, NON-POLAR MOLECULE

Centers Coincide, no dipole

In conclusion, to be a dipole, a polar molecule (or polyatomic ion), the presence of polar bonds is required. However, in addition, the polar bonds must be arranged so that they are not canceling. Molecular shape must be such that the center of the negative charge buildup does not coincide with the center of positive charge buildup.

Group Work Determine the polarity of each of the below: Draw to shape Check  en Draw arrow, if dipole

Hybrid structure

Importance of Polarity As it turns out, it is the difference in polarity which determines, for the same sized species, whether it is soluble in water and whether it is a gas at room temperature or a volatile liquid which evaporates quickly or a high boiling liquid which does not evaporate at all. The attractions between molecules which causes these differences all arise from increasing polarity or its complete lack...