Bonding
For elements in the s and p blocks, the number of valence electrons can easily be
determined from the group number. In the s block, Group 1 elements have one
valence electron, while Group 2 elements have two valence electrons. In the p
block, the number of valence electrons is equal to the group number minus ten.
Group 13 elements have three valence electrons, Group 14 elements have four, and
so on. The noble gases in Group 18 have eight valence electrons, and the full outer s
and p sublevels are what give these elements their special stability.
• When examining chemical bonding, it is necessary to keep track of the
valence electrons on each atom. An electron dot diagram shows the
valence electrons of an atom as dots distributed around the element’s
symbol. For example, a beryllium atom, which has two valence
electrons, would have the electron dot diagram below.
• Electron dot diagrams for all elements in a given group of
representative elements are identical (except for the elemental symbol)
because each element in one of those groups has the same number of
valence electrons. Most transition metals have two valence electrons in
their ground state, though some of the elements with unusual electron
configurations have only one.
• The Octet Rule
• The noble gases are unreactive because of their electron
configurations. American chemist, Gilbert Lewis (1875- 1946), used
this observation to explain the types of ions and molecules that are
formed by other elements. He called his explanation the octet rule. The
octet rule states that elements tend to form compounds in ways that give
each atom eight valence electrons. An exception to this rule is the
elements in the first period, which are particularly stable when they
have two valence electrons. A broader statement that encompasses both
the octet rule and this exception is that atoms react in order to achieve
the same valence electron configuration as that of a noble gas
• There are two ways in which atoms can satisfy the octet rule. One way
is by sharing their valence electrons with other atoms, which is called
Covalent Bonding. The second way is by transferring valence
electrons from one atom to another, which is called Ionic Bonding.
Atoms of metallic elements tend to lose all of their valence electrons,
which leaves them with an octet from the next lowest principal energy
level. Atoms of nonmetallic elements tend to gain electrons in order to
fill their outermost principal energy level with an octet.
• Cations are the positive ions formed when an atom loses one or more
electrons. The cations most commonly formed by the representative
elements are those that involve the loss of all valence electrons.
Consider the alkali metal, sodium (Na). It has one valence electron in
the third principal energy level. Upon losing that electron, the sodium
ion now has an octet of electrons from the second principal energy
level. The equation below illustrates this process.
Na
1s22s22p63s1
→
Na+ + e−
1s22s22p6 (octet)
• The electron configuration of the sodium ion is now the same as that of
the noble gas, neon. The term isoelectronic is used to describe two
atoms or ions that have the same electron configuration. The sodium
ion is isoelectronic with the neon atom.
• Consider the analogous processes for magnesium and aluminum:
Mg
1s22s22p63s2
Al
1s22s22p63s23p1
→
→
Mg2+ + 2e−
1s22s22p6 (octet)
Al3+ + 3e−
1s22s22p6 (octet)
• Both of these atoms form ions by losing all of their valence electrons,
two in the case of magnesium, and three in the case of aluminum. The
same noble gas configuration is achieved by all of these ions. In other
words, the Mg2+ ion, the Al3+ ion, the Na+ ion, and the Ne atom are
all isoelectronic. Under typical conditions, the representative elements
form cations by losing a maximum of three electrons.
• We can also show the loss of valence electron(s) with an electron dot
diagram.
• Na• → Na+ + e−
• Anions are the negative ions formed by gaining one or more electrons.
When nonmetal atoms gain electrons, they often do so until they reach
an octet of valence electrons in their outermost principal energy level.
This process is illustrated below for the elements fluorine, oxygen,
and nitrogen.
• All of these anions are isoelectronic with each other and with neon.
They are also isoelectronic with the three cations from the previous
section. Under typical conditions, a maximum of three electrons will be
gained during the formation of anions.
• Outer electron configurations are constant within a group, so this
pattern of ion formation repeats itself for Periods 3, 4, and so on
It is important not to misinterpret the concept of being isoelectronic. A
sodium ion is very different from a neon atom ( Figure 8.2) because
their nuclei contain different numbers of protons. One is an essential
ion that is a part of table salt, while the other is an unreactive gas that
makes up a very small part of the atmosphere. Likewise, sodium ions
are very different than magnesium ions, fluoride ions, and all the other
members of this isoelectronic series (N3−, O2−, F−, Ne, Na+, Mg2+,
Al3+).
• Transition metals belong to the d block, meaning that the d sublevel of
electrons is in the process of being filled with up to 10 electrons.
Many transition metals cannot lose enough electrons to attain a noble
gas electron configuration. Additionally, you have learned that the
majority of transition metals are capable of adopting ions with different
charges. Iron, which can form either Fe2+ or Fe3+ ions, loses electrons,
as shown below.
• According to the Aufbau process, the electrons fill the 4s sublevel
before beginning to fill the 3d sublevel. However, the outermost s
electrons are always the first to be removed when forming transition
metal cations. Because most transition metals have two valence
electrons, a charge of 2+ is very common for transition metal ions, as we
have already seen in the case of iron. A half-filled d sublevel (d5) is also
particularly stable. This type of configuration is obtained when an iron
atom loses a third electron.
• Some transition metals that have relatively few d electrons, such as
scandium, may be able to attain a noble gas electron configuration.
Others may attain configurations that include a full d sublevel, such as zinc and copper.
The resulting configuration above, with 18 electrons in the outermost principal energy
level, is referred to as a pseudo noble gas electron configuration. It gives particular
stability to the Zn2+ and Cu+ ions.
SUMMARY
••An electron dot diagram shows the chemical symbol of an element with dots that
represent valence electrons evenly distributed around the symbol.
••The octet rule states that elements form chemical compounds so that each atom will
acquire the electron configuration of a noble gas. Most noble gases have eight valence
electrons, except for helium, which has only two.
••Representative metals generally lose all of their valence electrons when forming ions,
leaving them with a complete octet of electrons from the next-lowest energy level. Most
nonmetals gain electrons when forming ions until their outer energy level has acquired
an octet.
••Atoms and ions that have the same electron configuration are called isoelectronic.
Common ions of representative elements are isoelectronic with a noble gas.
••When forming ions, transition metals lose their valence s-sublevel electrons before they
lose their d-sublevel electrons. Half-filled or completely filled d sublevels give transition
metal ions greater stability.