Polar C–X bonds
Carbon is only slightly more electronegative than
hydrogen. Therefore, we expect C–F, C–O, and C–N
bonds to be polar, just like the corresponding H–F, H–O,
and H–N bonds. Unfortunately, polar bonds involving carbon
are almost impossible to detect using potential maps.
Methyl propionate, CH3CH2C(=O)OCH3,
furnishes us with a good example of this phenomenon. Two carbons
in this molecule are bonded to oxygen. The methyl carbon
makes one CO bond, O-CH3, and the carbonyl carbon
makes three CO bonds, C(=O)-O. Since all of the CO bonds are polar,
all of them transfer electron density from carbon to oxygen and
we expect to find large positive potentials around these carbons.
The following potential map of methyl propionate standard color
scale) shows that this does not happen. The potential near each
carbon is very small.
of methyl propionate, CH3CH2C(=O)OCH3
(standard color scale)
Two factors account for these observations. First,
the “local” atom principle breaks down near carbon whenever
it is sandwiched between several other atoms. Notice that the carbonyl
carbon in methyl propionate lies in between two oxygens, both of
which create large local negative potentials. These negative potentials
offset the positive potentials created by carbon.
A second important factor is atom size. Larger atoms
have larger size density surfaces. This means we calculate and display
potential at regions farther away from the atom. Since potential
falls off as 1/r, a potential map will show smaller potentials near
a large atom than it will near a small atom, other factors being
The effect of distance (or atom radius) on potential
is dramatically illustrated by the potential maps of the halide
anions (following figure, color scale = –200 to –120
Potential maps of
(red = –200, dark blue = –120 kcal/mol): F–,
Cl–, Br–, I–.
All of these ions carry identical charges of –1,
but the potential is most negative around the smallest anion, fluoride.
As ion and map size increase, the map’s potential becomes
steadily less negative.
These observations, when taken together, suggest several
principles that should be used when interpreting potential maps:
The “local” atom interpretation
must take account of atom size. We should only compare
“local” potentials for atoms of similar size.
“Local” potentials are also
affected by neighboring atoms (a common problem with carbon)
and by non-spherical patterns in the atom’s electron
density cloud (a common problem for nitrogen and oxygen).
We should never compare carbon and hydrogen
potentials. Even when a carbon and hydrogen carry similar
charges, the “local” potential near carbon
will always be much smaller.