Chemistry Reference
In-Depth Information
We can deduce the shapes of many molecules from
their electron dot diagrams and the fact that the electron
groups on the central atom tend to get as far apart as pos-
sible while remaining attached to that atom. We then at-
tach the outer atoms and describe the molecular shape by
the positions of the atoms. The electron groups and their
attached atoms are not necessarily limited to a planar
arrangement (Section 13.4).
A molecule with more than one polar bond in which all
the effects of the polar bonds do not cancel out has a dipole
moment—an electrical dissymmetry that causes an inter-
molecular attraction between this molecule and other simi-
lar ones. This attraction is called a dipolar attraction, and it
lowers the ability of the substance to exist in the gas phase.
However, if the polar bonds in a molecule are oriented so
that their effects are canceled out, as in carbon dioxide, then
a molecule with no dipole results (Section 13.5).
Intermolecular attractions include dipolar attrac-
tions, as well as van der Waals forces and hydrogen bond-
ing. van der Waals forces are similar to dipolar attractions
but result from instantaneous dissymmetry of charge,
which may disappear the next instant. The more electrons
in the molecule, the greater is the van der Waals force.
However, van der Waals forces tend to be lower in mag-
nitude than dipolar attractions.
Hydrogen bonding is an intermolecular force
between a hydrogen atom bonded to a fluorine, oxygen,
or nitrogen atom and an unshared pair of electrons on an-
other such atom in an adjacent molecule (or sometimes
even in the same molecule). The hydrogen atom may be
bonded to one small electronegative atom and attracted to
another at one instant and then bonded to the second and
attracted to the first at the next instant (Section 13.6).
Items for Special Attention
Be sure to distinguish between atomic size and ionic size.
Molecular geometry is actually determined experimen-
tally. We can deduce the geometry of many molecules from
their electron dot diagrams, with correct results in a great
majority of cases.
■
■
Ionization energy and electronegativity vary in the same
ways in the periodic table, but opposite to the way atomic
size varies.
■
Remember the difference between polar bonds and polar
molecules. A molecule can have polar bonds without being
a polar molecule.
■
If a bond is either polar or nonpolar, it is covalent.
■
Adding a new shell of electrons considerably increases the
size of the atom; breaking into a complete octet to ionize
an electron requires a huge quantity of energy.
■
Hydrogen bonding is the intermolecular force between
molecules that have hydrogen atoms attached to very
small, highly electronegative atoms with unshared pairs of
electrons (nitrogen, oxygen, or fluorine atoms). It has noth-
ing to do with the chemical bond between the two hydro-
gen atoms in
■
The term
ionization energy
means “first ionization energy”
unless otherwise specified.
■
Any molecule with one or more lone pairs on the central
atom will have a different shape of the electron groups
from the shape of the molecule (defined by the locations of
the atoms).
■
H
2
molecules.
Answers to Snapshot Reviews
13.1 A. Opposite charges attract, and the more protons there
are, the greater will be the attraction for the electrons.
13.2 A. (a)
13.4 A. (a) Tetrahedral
(b) Trigonal pyramidal
(c) Angular
13.5 A. (a) (It has a symmetrical structure.)
13.6 A. (a) Chemical bond
Sr
6
Ca
6
Mg
(b)
As
6
S
6
F
CCl
4
(c)
13.3 A. (a) F
Sb
6
Se
6
S
(b) Intermolecular force
(b) N
(c) Br
Self-Tutorial Problems
13.1
(a) In which periodic group does each atom have a new
shell of electrons?
(b) In which periodic group does the size of each ele-
ment increase markedly from the preceding atom in
the periodic table?
13.2
Make a table of the number of protons, the number of elec-
trons, and the relative sizes for each of the following parts:
(a) F and
F
O
2
(b) Ne and
Na
(d) Ar and Cl
(c) Na and
