Page 50 - Materials Science and Engineering An Introduction

22 • Chapter 2 / Atomic Structure and Interatomic Bonding

Concept Check 2.1 Why are the atomic weights of the elements generally not integers?
Cite two reasons.
[The answer may be found at www.wiley.com/college/callister (Student Companion Site).]

2.3 ELECTRONS IN ATOMS
Atomic Models
During the latter part of the nineteenth century, it was realized that many phenomena
involving electrons in solids could not be explained in terms of classical mechanics. What
followed was the establishment of a set of principles and laws that govern systems of atomic
quantum mechanics and subatomic entities that came to be known as quantum mechanics. An understanding of
the behavior of electrons in atoms and crystalline solids necessarily involves the discussion
of quantum-mechanical concepts. However, a detailed exploration of these principles is
beyond the scope of this text, and only a very superficial and simplified treatment is given.
Bohr atomic model One early outgrowth of quantum mechanics was the simplified Bohr atomic model,
in which electrons are assumed to revolve around the atomic nucleus in discrete orbitals,
and the position of any particular electron is more or less well defined in terms of its
orbital. This model of the atom is represented in Figure 2.1.
Another important quantum-mechanical principle stipulates that the energies of
electrons are quantized—that is, electrons are permitted to have only specific values of
energy. An electron may change energy, but in doing so, it must make a quantum jump
either to an allowed higher energy (with absorption of energy) or to a lower energy (with
emission of energy). Often, it is convenient to think of these allowed electron energies as
being associated with energy levels or states. These states do not vary continuously with en-
ergy—that is, adjacent states are separated by finite energies. For example, allowed states
for the Bohr hydrogen atom are represented in Figure 2.2a. These energies are taken to be
negative, whereas the zero reference is the unbound or free electron. Of course, the single
electron associated with the hydrogen atom fills only one of these states.
Thus, the Bohr model represents an early attempt to describe electrons in atoms, in
terms of both position (electron orbitals) and energy (quantized energy levels).
This Bohr model was eventually found to have some significant limitations because of
its inability to explain several phenomena involving electrons. A resolution was reached
wave-mechanical with a wave-mechanical model, in which the electron is considered to exhibit both wave-
model like and particle-like characteristics. With this model, an electron is no longer treated as
a particle moving in a discrete orbital; rather, position is considered to be the probability
of an electron’s being at various locations around the nucleus. In other words, position is
described by a probability distribution or electron cloud. Figure 2.3 compares Bohr and

Figure 2.1 Schematic representation of the Bohr Orbital electron
atom.

Nucleus