Tweet |
Custom Search
|
|
As stated in the previous chapter, the number of electrons in the outer, or valence, shell determines the relative activity of the element. The arrangement of electrons in the outer shell explains why some elements are chemically very active, some are not very active, and others are inert. In general, the fewer electrons an element must lose, gain, or share to reach a stable shell structure, the more chemically active the element is. The likelihood of elements forming compounds is strongly influenced by the completion of the valence shell and by the stability of the resulting molecule. The more stable the resulting molecules are, the more likely these molecules are to form. For example, an atom that "needs" two electrons to completely fill the valence shell would rather react with another atom which must give up two electrons to satisfy its valence. In the case of H+
+ Br, this is likely to take place because the exchange would satisfy
the needs of both atoms. Although there is far more to consider than
just the number of valence electrons, this is a good rule of thumb. If the atom needed
two electrons and only picked up one, it would still actively seek out
an additional electron. The reaction of H+ + Te-2
is far less likely to take place because the resulting molecule would
still have an incomplete valence shell. Of course, the combining of two
atoms, when both want to release or gain electrons, may take place (for
example; H2 or O2)
but is less probable when other atoms are available. Atoms are joined or
bonded together through this interaction of their electrons. There are
several types of chemical bonds that hold atoms together; three will be
discussed, ionic, covalent, and metallic. <%CUT%> Ionic Bonds An ionic bond
is formed when one or more
electrons is wholly transferred from one element to another, and the
elements are held together by the force of attraction due to the
opposing charges. An example of ionic bonding is shown in Figure 6(A)
for sodium chloride (table salt).
Figure 6 Ionic Bond, Sodium Chloride The sodium atom
loses the one electron in its outer shell to the chlorine atom, which
uses the electron to fill its outer shell. When this occurs, the sodium
atom is left with a +1 charge and the chlorine atom a -1 charge. The
ionic bond is formed as a result of the attraction of the two
oppositely-charged particles. No single negatively-charged ion has a
greater tendency to bond to a particular positively-charged ion than to
any other ion. Because of this,
the positive and negative ions arrange themselves in three dimensions,
as shown in Figure 6(B), to balance the charges among several ions. In
sodium chloride, for example, each chloride ion is surrounded by as many
sodium ions as can easily crowd around it, namely six. Similarly, each
sodium ion is surrounded by six chloride ions. Therefore, each chloride
ion is bonded to the six nearest sodium ions and bonded to a lesser
extent to the more distant sodium ions. Accordingly, the ionic bond is a
force holding many atoms or ions together rather than a bond between two
individual atoms or ions. A covalent bond is formed when one or more electrons from an atom pair off with one or
more electrons from another atom and form overlapping electron shells in
which both atoms share the paired electrons. Unlike an ionic bond, a
covalent bond holds together specific atoms. Covalent bonding can be
single covalent, double covalent, or triple covalent depending on the
number of pairs of electrons shared. Figure 7 shows the bonding that
occurs in the methane molecule, which consists of four single covalent
bonds between one carbon atom and four hydrogen atoms.
Figure 7 Covalent Bond, Methane CH4 Two double
covalent bonds result when carbon dioxide, which consists of one carbon
atom and two oxygen atoms, is formed. Four pairs of electrons are shared
by the carbon atom, two from each of the two oxygen atoms as shown in
Figure 8. A combination of two electrons form a combination of lower
energy than their energy when separated. This energy difference
represents the force that binds specific atoms together.
Figure 8 Formation of the Carbon Dioxide Molecule When both shared
electrons in a covalent bond come from the same atom, the bond is called
a coordinate covalent bond. Although both shared electrons come from the
same atom, a coordinate covalent bond is a single bond similar in
properties to a covalent bond. Figure 9 illustrates the bonds of the
negatively-charged chlorate ion. The ion consists of one chlorine atom
and three oxygen atoms with a net charge of -1, and is formed with two
coordinate covalent bonds and one covalent bond. The chlorine atom has
effectively gained an electron through the covalent bond, which causes
the overall negative charge.
Figure 9 Coordinate Covalent Bond, Chlorate Ion CIO4 Covalent bonds can
be either polar or nonpolar. When the shared pair of electrons is not
shared equally, one end of the bond is positive, and the other end is
negative. This produces a bond with two poles called a polar covalent
bond. Molecules having
polar covalent bonds are called dipolar or
polar molecules. Water is an
example of a polar molecule. When two atoms of the same element share
one or more pairs of electrons (such as H or N), each atom exerts the
same attraction for the shared electron pair or pairs. When the electron
pairs are distributed or shared equally between the two like atoms, the
bond is called a nonpolar covalent bond. If all the bonds in a molecule
are of this kind, the molecule is called a nonpolar covalent molecule.
|
||