Page 84 - Theory and Problems of BEGINNING CHEMISTRY

CHAP. 5] CHEMICAL BONDING 73

There are exceptions to the so-called octet rule, but only a few will be encountered within the scope of this
book. Such cases will be discussed in detail as they arise.
The constituent atoms in polyatomic ions are also linked by covalent bonds. In these cases, the net charge
on the ion is determined by the total number of electrons and the total number of protons. For example, the
ammonium ion, NH 4 , formed from five atoms, contains one fewer electron than the number of protons. A
+
nitrogen atom plus 4 hydrogen atoms contains a total of 11 protons and 11 electrons, but the ion has only
10 electrons, 8 of which are valence electrons:

+
H
H N H
H

Similarly, the charge on a hydroxide can be determined by counting the total numbers of protons and electrons.
The hydroxide ion contains one valence electron more than the total in the two individual atoms—oxygen and
hydrogen:

_
O H

Covalent Bonding of More Than Two Atoms
Writing electron dot diagrams for molecules or ions containing only two atoms is relatively straightforward.
When several atoms are to be represented as being linked by means of covalent bonds, the following procedure
may be used to determine precisely the total number of electrons which must be shared among the atoms. The
procedure, useful only for compounds in which the atoms obey the octet rule, will be illustrated using sulfur
dioxide as an example.

Steps Example
1. Determine the number of valence electrons available. S 6
2O 12
Total 18
2. Determine the number of electrons necessary to satisfy the octet (or duet) S 8
rule with no electron sharing. 2O 16
Total 24
3. The difference between the numbers obtained in steps 2 and 1 is the number Required 24
of bonding electrons. Available −18
To be shared 6
4. Place the atoms as symmetrically as possible. (Note that a hydrogen atom OSO
cannot be bonded to more than one atom, since it is capable of sharing only
two electrons.)
5. Place the number of electrons to be shared between the atoms, a pair at a O S O
time, at first one pair between each pair of atoms. Use as many pairs as
remain to make double or triple bonds.
6. Add the remainder of the available electrons to complete the octets (or duets) O S O
of all the atoms. There should be just enough if the molecule or ion follows
the octet rule.

EXAMPLE 5.7. Draw electron dot diagrams showing the bonding in the following compounds: (a) CO, (b)CO 2 ,(c)CH 4 ,
and (d) MgCl 2 .