Why should atoms bond at all? In nature we find that some elements
like He, Ne, and Ar are
never found bonded to other atoms whereas most other elements are only
found bonded to other
elements. What makes the noble gases so special? The answer lies in
their closed shell
electron configurations. Because the valence shell of a noble gas is
completely full, it cannot accept another electron into the shell. The nucleus
is positively
charged and pulls on the electron, so the loss of an
electron from a noble gas is unfavorable. Therefore, like real nobility,
the noble gases do not want
to do anything at all--that is, noble gases are unreactive because they
have filled valence shells.
Any element other than a noble gas has an open shell configuration, which is
unstable relative to the configuration of a noble gas. Non-noble atoms react to
form bonds in an
attempt to achieve a closed
shell electron configuration. For example, when a lithium atom and a
fluorine atom meet, as shown
in , lithium can achieve a noble gas configuration,
1s2, by donating
an electron to fluorine which also achieves the noble gas configuration
1s22s22p6:
Figure %: Transferring an electron from Li to F gives both noble gas
electron configurations.
The above reaction represents the formation of an ionic bond. The
negatively charged
anion, F, and the positively charged cation are held together in
the bond by the attraction
of unlike charges as dictated by Coulomb's law. You may have asked
yourself why two
fluorine atoms don't come together to perform the following reaction:
Even though the reaction may appear to be favorable because of its production of
a
closed shell species, there
is a way to have both F atoms achieve a noble gas configuration. By
sharing their electrons,
each fluorine atoms can have a complete octet in its valence shell.
Such a sharing of
electrons is called a covalent bond and will be discussed in depth in a
separate section.
Properties of a Bond
The way bond properties were chosen to characterize bonds have a historical
basis. Scientists made their first rational attempts to describe bonding by
looking at data they could collect
about bonds. We too will look at the experimental data on bonds to try to
analyze
bonding.
Perhaps the most useful aspect to know of a bond is its strength. Weak
bonds are easily broken and
molecules with such bonds are fairly reactive. Conversely,
strong bonds are difficult to
break and give rise to stable molecules. Therefore, it is sensible to
define bond strength as the
amount of energy needed to break a chemical bond. Trends in bond strength
show that
homoatomic bonds (those formed between atoms of the same element) tend to be
strong. But going across a row in the periodic table, the trend in bond
strength may not be regular. For example, period 2 elements have the following
strength order: Li-Li > Be-Be
< B-B < C-C < N-N > O-O > F-F. This irregular trend is repeated in
period 3 homoatomic
bonds. If we look at bond strength data, we also notice that the Li-F bond
is several times stronger
than the F-F bond or the Li-Li bond. It is not important for you to memorize
such trends. We use them
to show that whatever theory of covalent bonding we propose must account
for these observations.
