Page 60 - Materials Science and Engineering An Introduction

32 • Chapter 2 / Atomic Structure and Interatomic Bonding

a number of material properties depend on E 0 , the curve shape, and bonding type.
For example, materials having large bonding energies typically also have high melt-
ing temperatures; at room temperature, solid substances are formed for large bonding
energies, whereas for small energies, the gaseous state is favored; liquids prevail when
the energies are of intermediate magnitude. In addition, as discussed in Section 6.3, the
mechanical stiffness (or modulus of elasticity) of a material is dependent on the shape
of its force–versus–interatomic separation curve (Figure 6.7). The slope for a relatively
stiff material at the r r 0 position on the curve will be quite steep; slopes are shallower
for more flexible materials. Furthermore, how much a material expands upon heating
or contracts upon cooling (i.e., its linear coefficient of thermal expansion) is related to
the shape of its E–versus–r curve (see Section 19.3). A deep and narrow “trough,” which
typically occurs for materials having large bonding energies, normally correlates with
a low coefficient of thermal expansion and relatively small dimensional alterations for
changes in temperature.
primary bond Three different types of primary or chemical bond are found in solids—ionic, cova-
lent, and metallic. For each type, the bonding necessarily involves the valence electrons;
furthermore, the nature of the bond depends on the electron structures of the constitu-
ent atoms. In general, each of these three types of bonding arises from the tendency of
the atoms to assume stable electron structures, like those of the inert gases, by com-
pletely filling the outermost electron shell.
Secondary or physical forces and energies are also found in many solid materials;
they are weaker than the primary ones but nonetheless influence the physical proper-
ties of some materials. The sections that follow explain the several kinds of primary and
secondary interatomic bonds.

2.6 PRIMARY INTERATOMIC BONDS
Ionic Bonding
ionic bonding Ionic bonding is perhaps the easiest to describe and visualize. It is always found in
compounds composed of both metallic and nonmetallic elements, elements situated
at the horizontal extremities of the periodic table. Atoms of a metallic element easily
give up their valence electrons to the nonmetallic atoms. In the process, all the atoms
acquire stable or inert gas configurations (i.e., completely filled orbital shells) and, in
addition, an electrical charge—that is, they become ions. Sodium chloride (NaCl) is
the classic ionic material. A sodium atom can assume the electron structure of neon
(and a net single positive charge with a reduction in size) by a transfer of its one va-
lence 3s electron to a chlorine atom (Figure 2.11a). After such a transfer, the chlorine
ion acquires a net negative charge, an electron configuration identical to that of argon;
it is also larger than the chlorine atom. Ionic bonding is illustrated schematically in
Figure 2.11b.
coulombic force The attractive bonding forces are coulombic—that is, positive and negative ions, by
virtue of their net electrical charge, attract one another. For two isolated ions, the attrac-
tive energy E A is a function of the interatomic distance according to
Attractive energy—
interatomic E A = - A (2.9)
separation r
relationship
Theoretically, the constant A is equal to

1
A = ( Z 1 e)( Z 2 e) (2.10)
4pP 0