Effects of Structure on Properties
Physical properties of metals, ceramics, and polymers, such as ductility, thermal expansion, heat
capacity, elastic modulus, electrical conductivity, and dielectric and magnetic properties, are a direct
result of the structure and bonding of the atoms and ions in the material. An understanding of the
origin of the differences in these properties is of great engineering importance.
In single crystals, a physical property such as thermal expansion varies with direction, reflecting
the crystal structure; whereas in polycrystalline and amorphous materials, a property does not vary
with direction, reflecting the average property of the individual crystals or the randomness of the
amorphous structure. Most engineering materials are polycrystalline, composed of many grains, and
thus an understanding of the properties requires not only a knowledge of the structure of the single
grains but also a knowledge of grain size and orientation, grain boundaries, and other phases present;
that is, a knowledge of the microstructure of this material.
Atomic Structure
Atoms consist of electrons, protons, and neutrons. The central nucleus consists of positively charged protons and electrically neutral neutrons. Negatively charged electrons are in orbits about the nucleus
in different energy levels, occupying a much larger volume than the nucleus.
In an atom, the number of electrons equals the number of protons and, hence, an atom is neutral.
The atomic number of an element is given by the number of protons, and the atomic weight is given
by the total number of protons and neutrons. (The weight of the electrons is negligible.) Thus,
hydrogen, H, with one proton and one electron, has an atomic number of 1 and an atomic weight of
1 and is the first element in the periodic chart. Oxygen, O, with atomic number 8, has eight protons
and eight neutrons and, hence, an atomic weight of 16.
Completed electronic shells have a lower energy than partially filled orbitals when bonded to
other atoms. As a result of this energy reduction, atoms share electrons to complete the shells, or
gain or lose electrons to form completed shells. In the latter case, ions are formed in which the
number of electrons is not equal to the number of protons. Thus, O by gaining two electrons, has a
charge of -2 and forms the oxygen ion O2-.
The periodic chart arranges elements in columns of the same electronic configuration. The first
column consists of the alkalies Li, Na, K, Cs, Rb; each has one electron in the outer shell that can
be lost. Similarly, the second column of alkaline-earths can form Mg2+, Ca2+, Sr2+, Ba2+ by losing
two electrons. The seventh column consists of the halogens Fl, Cl, Br, I, which by gaining one
electron become the halides, all with a charge of -1. The eighth column consists of the inert gases
He, Ne, Ar, K, Xe, with completed shells. The bonding of the elements and ions with similar electronic
configurations is similar. Moving down a column increases the number of electrons and, hence,
the atom's size increases even though the outer electronic configuration remains the same.
The outer electrons that are lost, gained, or shared are called valence electrons, and the inner
electrons are called core electrons. For the most part, the valence electrons are important in determining
the nature of the bonding and, hence, the structure and properties of the materials.
Bonding
When two atoms or ions are within atomic distances of each other, distances of 0.5-3.OA, bonding
may occur between the atoms or ions. The resulting reduction in energy due to an attractive force
leads to the formation of polyatomic gas molecules, liquids, and solids. If the energy of the bonds
is large (75-275 kcal/mol), primary bonds are formed—metallic, ionic, or covalent. If the energy of
the bond is smaller (1-10 kcal/mol), secondary bonds are formed—van der Waals and hydrogen. In
addition, combinations of bond types, such as a mixture of ionic and covalent bonds, may occur.
Metallic Bonding
In a metallic crystal, an ordered arrangement of nuclei and their electrons is embedded in a cloud of
valence electrons, which are shared throughout the lattice. The resulting bonding is a nondirectional
primary bond. Since the binding energy of the valence electrons is relatively small, the mobility of
these electrons is high and creates high electrical and thermal conductivity. The atoms are approximately
spherical in shape as a result of the shape of completed inner shell. Examples of metals are
Cu, Au, Ag, and Na.
Ionic Bonding
The strongest type of bonding between two oppositely charged particles is called ionic bonding. The
positively charged ions (cations) attract as many negatively charged ions (anions) as they can and
form ionic bonds. The primary bond formed is nondirectional if the bonding is purely ionic. Li+ and
F~ in LiF form predominately ionic bonds. In general, since the electrons are strongly bonded,
electrical and thermal conductivities are much smaller than in metals and, thus, ionic bonded materials
are classified as insulators or dielectrics.
Covalent Bonding
Covalent bonding results from an overlap or sharing, not from gain or loss of valence electrons. A
net reduction of energy as a result of each atom's completing the other's orbital also results in a
primary bond, but it is directional. The directionality is a result of the shape of the orbitals involved
in the bonding. When C is covalently bonded to four other C's in diamond, the bonding is purely
covalent and the configuration of these four bonds is tetrahedral. When B, however, is bonded to
three other B's, a triangular configuration is formed. Organic polymers and diatomic gases such as
Cl2 are typical examples of covalent bonding. As a result of the strong bonding of the valence
electrons, these materials, for the most part, have low electrical and thermal conductivity.
Van der Waals and Hydrogen Bonding
Van der Waals bonds are secondary bonds, the result of fluctuating dipoles, due to the fact that at an
instant of time the centers of positive and negative charge do not coincide. An example is an inert
gas such as Ar, which below -19O0C forms a solid as a result of these weak attractive forces. Similar
weak forces exist in molecules and solids. Hydrogen bonds are also secondary bonds, but they are
the result of permanent dipoles. For example, the water molecule, H2O, is nonlinear and the bonding
between H and an adjacent O in water results in H2O being a liquid above O0C a 1 ami pressure
rather than a gas, as is the case for other molecules of comparable molecular weight....................
