《材料科学导论 Introduction to Materials Science》参考书籍：Materials Science and Engineering - An Introduction（7th）

496c0215-3712/20/057:19Page1e 2nd REVISE PaGES EQA Learning objectives After careful study of this chapter you should be able to do the following: 1. Name the two atomic models cited, and note (b)Note on this plot the equilibrium separation the differences between them and the bonding energy. 2. Describe the important quantum-mechanical 4.(a) Briefly describe ionic, covalent, metallic, principle that relates to electron energies hydrogen, and van der Waals bonds. 3.(a)Schematically plot attractive, repulsive, and (b) Note which materials exhibit each of these net energies versus interatomic separation ding type for two atoms or ions 2.1 INTRODUCTION Some of the important properties of solid materials depend on geometrical atomic arrangements, and also the interactions that exist among constituent atoms or mol- ecules. This chapter, by way of preparation for subsequent discussions, considers several fundamental and important concepts-namely, atomic structure, electron configurations in atoms and the periodic table, and the various types of primary and secondary interatomic bonds that hold together the atoms comprising a solid These topics are reviewed briefly, under the assumption that some of the material is familiar to the reader Atomic Structure 2.2 FUNDAMENTAL CONCEPTS Each atom consists of a very small nucleus composed of protons and neutrons, which is encircled by moving electrons. Both electrons and protons are electrically charged the charge magnitude being 1.60 X 10C, which is negative in sign for electrons nd positive for protons; neutrons are electrically neutral. Masses for these sub atomic particles are infinitesimally small; protons and neutrons have approximately the same mass, 1.67X 10-kg, which is significantly larger than that of an elec tron,9.11×10-31kg Each chemical element is characterized by the number of protons in the nu- cleus, or the atomic number(z). For an electrically neutral or complete atom, the atomic number also equals the number of electrons. This atomic number ranges in integral units from 1 for hydrogen to 92 for uranium, the highest of the naturally occurring elements. The atomic mass(A)of a specific atom may be expressed as the sum of the asses of protons and neutrons within the nucleus. Although the number of protons is the same for all atoms of a given element, the number of neutrons(N) may be isotope variable. Thus atoms of some elements have two or more different atomic masses. which are called isotopes. The atomic weight of an element corresponds to the weighted average of the atomic masses of the atoms naturally occurring isotopes. The atomic mass unit(amu) may be used for computations of atomic weight. A Terms appearing in boldface type are defined in the Glossary, which follows Appendix E The termatomic mass"is really more accurate than"atomic weight"inasmuch as, in this context, we are dealing with masses and not weights. However, atomic weight is, by conven- tion, the preferred terminology and will be used throughout this book. The reader shor note that it is not necessary to divide molecular weight by the gravitational constant

2.1 INTRODUCTION Some of the important properties of solid materials depend on geometrical atomic arrangements, and also the interactions that exist among constituent atoms or mol￾ecules. This chapter, by way of preparation for subsequent discussions, considers several fundamental and important concepts—namely, atomic structure, electron configurations in atoms and the periodic table, and the various types of primary and secondary interatomic bonds that hold together the atoms comprising a solid. These topics are reviewed briefly, under the assumption that some of the material is familiar to the reader. Atomic Structure 2.2 FUNDAMENTAL CONCEPTS Each atom consists of a very small nucleus composed of protons and neutrons, which is encircled by moving electrons. Both electrons and protons are electrically charged, the charge magnitude being which is negative in sign for electrons and positive for protons; neutrons are electrically neutral. Masses for these sub￾atomic particles are infinitesimally small; protons and neutrons have approximately the same mass, which is significantly larger than that of an elec￾tron, Each chemical element is characterized by the number of protons in the nu￾cleus, or the atomic number (Z).1 For an electrically neutral or complete atom, the atomic number also equals the number of electrons. This atomic number ranges in integral units from 1 for hydrogen to 92 for uranium, the highest of the naturally occurring elements. The atomic mass (A) of a specific atom may be expressed as the sum of the masses of protons and neutrons within the nucleus.Although the number of protons is the same for all atoms of a given element, the number of neutrons (N) may be variable. Thus atoms of some elements have two or more different atomic masses, which are called isotopes. The atomic weight of an element corresponds to the weighted average of the atomic masses of the atom’s naturally occurring isotopes.2 The atomic mass unit (amu) may be used for computations of atomic weight. A 9.11 1031 kg. 1.67 1027 kg, 1.60 1019 C, Learning Objectives After careful study of this chapter you should be able to do the following: 1. Name the two atomic models cited, and note the differences between them. 2. Describe the important quantum-mechanical principle that relates to electron energies. 3. (a) Schematically plot attractive, repulsive, and net energies versus interatomic separation for two atoms or ions. (b) Note on this plot the equilibrium separation and the bonding energy. 4. (a) Briefly describe ionic, covalent, metallic, hydrogen, and van der Waals bonds. (b) Note which materials exhibit each of these bonding types. 1 Terms appearing in boldface type are defined in the Glossary, which follows Appendix E. 2 The term “atomic mass” is really more accurate than “atomic weight” inasmuch as, in this context, we are dealing with masses and not weights. However, atomic weight is, by conven￾tion, the preferred terminology and will be used throughout this book. The reader should note that it is not necessary to divide molecular weight by the gravitational constant. atomic number isotope atomic weight atomic mass unit 1496T_c02_15-37 12/20/05 7:19 Page 16 2nd REVISE PAGES

496c0215-3711/10/0510:42page1 REVISED PAGES EQA 2.3 Electrons in Atoms 17 scale has been established whereby 1 amu is defined as f of the atomic mass of the most common isotope of carbon, carbon 12(C)(A=12.00000). Within this cheme, the masses of protons and neutrons are slightly greater than unity, and The atomic weight of an element or the molecular weight of a compound may be specified on the basis of amu per atom(molecule) or mass per mole of material In one mole of a substance there are 6.023 X 10(Avogadro's number)atoms or molecules. These two atomic weight schemes are related through the following equation 1 amu/atom(or molecule)=1 g/mol For example, the atomic weight of iron is 55.85 amu/atom, or 55.85 g/mol. Sometimes se of amu per atom or molecule is convenient; on other occasions g(or kg)/mol is preferred. The latter is used in this book Concept Check 2.1 Why are the atomic weights of the elements generally not integers? Cite two [thEanswermaybefoundatwww.wileycom/college/callister(studentCompanionSite).I 2.3 ELECTRONS IN ATOMS Atomic Models During the latter part of the nineteenth century it was realized that many phe nomena 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 and subatomic entities that came to be known as quantum mechanics. An understanding of the behavior of electrons in atoms and crystalline olds necessarily involves the discussion of quantum-mechanical concepts. How ever, a detailed exploration of these principles is beyond the scope of this book, and only a very superficial and simplified treatment is given 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 Fi Another important quantum-mechanical principle stipulates that the energies of electrons are quantized; that is, electrons are permitted to have only specific val- ues of energy. An electron may change energy, but in doing so it must make a quan- tum 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 al lowed electron energies as being associated with energy levels or states. These states do not vary continuously with energy; 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 will fill only one of these states

scale has been established whereby 1 amu is defined as of the atomic mass of the most common isotope of carbon, carbon Within this scheme, the masses of protons and neutrons are slightly greater than unity, and (2.1) The atomic weight of an element or the molecular weight of a compound may be specified on the basis of amu per atom (molecule) or mass per mole of material. In one mole of a substance there are (Avogadro’s number) atoms or molecules. These two atomic weight schemes are related through the following equation: For example, the atomic weight of iron is 55.85 amu/atom, or 55.85 g/mol. Sometimes use of amu per atom or molecule is convenient; on other occasions g (or kg)/mol is preferred. The latter is used in this book. 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 phe￾nomena 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 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. How￾ever, a detailed exploration of these principles is beyond the scope of this book, and only a very superficial and simplified treatment is given. 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 val￾ues of energy. An electron may change energy, but in doing so it must make a quan￾tum 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 al￾lowed electron energies as being associated with energy levels or states. These states do not vary continuously with energy; 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 will fill only one of these states. 1 amu/atom 1or molecule2  1 g/mol 6.023 1023 A Z  N 12 1 12C2 1A  12.000002. 1 12 2.3 Electrons in Atoms • 17 quantum mechanics Bohr atomic model mole 1496T_c02_15-37 11/10/05 10:42 Page 17 REVISED PAGES

496c0215-3712/20/0513:51Page1 2nd REVISE Pages EQA 18. Chapter 2 / Atomic Structure and Interatomic Bonding Orbital electron Figure 2.1 Schematic representation of the bohi Nucleus Thus, the Bohr model represents an early attempt to describe electrons in oms, in terms of both position(electron orbitals)and energy(quantized energy This Bohr model was eventually found to have some significant limitations because of its inability to explain several phenomena involving electrons. A wave-mechanical considered to exhibit both wave-like and particle-like characteristics. With he resolution was reached with a wave-mechanical model. in which the electron model, an electron is no longer treated as a particle moving in a discrete or- bital; 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 wave mechanical models for the hydrogen atom. Both these models are used through- out the course of this book; the choice depends on which model allows the more simple explanation first three electron 3p energy states for the Bohr hydrogen atom. (b) Electron energy states for the first three shells of the wave-mechanical 5 hydrogen atom. 1x 10-18 s(Adapted from WG Moffatt. G. W. Pearsall, and j. wulff The structure and Properties of Materials, Vol I Structure, p. 10 Copyright o 1964 by John wiley Sons by permission of John Wiley Sons, Inc

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 with a wave-mechanical model, in which the electron is considered to exhibit both wave-like and particle-like characteristics. With this model, an electron is no longer treated as a particle moving in a discrete or￾bital; 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 wave￾mechanical models for the hydrogen atom. Both these models are used through￾out the course of this book; the choice depends on which model allows the more simple explanation. 18 • Chapter 2 / Atomic Structure and Interatomic Bonding Orbital electron Nucleus Figure 2.2 (a) The first three electron energy states for the Bohr hydrogen atom. (b) Electron energy states for the first three shells of the wave-mechanical hydrogen atom. (Adapted from W. G. Moffatt, G. W. Pearsall, and J. Wulff, The Structure and Properties of Materials, Vol. I, Structure, p. 10. Copyright © 1964 by John Wiley & Sons, New York. Reprinted by permission of John Wiley & Sons, Inc.) Figure 2.1 Schematic representation of the Bohr atom. 0 0 –1 × 10–18 –2 × 10–18 (a) (b) –5 –10 –15 n = 1 1s Energy (J) Energy (eV) n = 2 n = 3 2s 3s 2p 3p 3d –1.5 –3.4 –13.6 wave-mechanical model 1496T_c02_15-37 12/20/05 13:51 Page 18 2nd REVISE PAGES

496Tc0215-3712/20/057:19page1 2nd REVISE PaGeS EQA 2.3 Electrons in Atoms 19 Fi he(a) Bohr and(b) in terms of electron distribution.(Adapted from Z.D. Jastrzebski. The Nature and Properties of Engineering Materials, 3r 1987 by John Wiley Sons New York. Reprinted by permission of John Wiley Distance from nucleus -+i- Orbital electron Quantum Numbers sing wave mechanics, every electron in an atom is characterized by four parame- ters called quantum numbers. The size, shape, and spatial orientation of an electrons probability density are specified by three of these quantum numbers. Furthermore, Bohr energy levels separate into electron subshells, and quantum numbers dictate the number of states within each subshell. Shells are specified by a principal quantum number n, which may take on integral values beginning with unity, sometimes these shells are designated by the letters K, L, M, N,O, and so on, which correspond espectively, to n =1, 2, 3, 4, 5,..., as indicated in Table 2. 1. Note also that this quan- tum number, and it only, is also associated with the bohr model. This quantum num- ber is related to the distance of an electron from the nucleus, or its position. The second quantum number, L, signifies the subshell, which is denoted by wercase letter-an s, p, d, or f, it is related to the shape of the electron subshell In addition, the number of these subshells is restricted by the magnitude of n Allowable subshells for the several n values are also presented in Table 2. 1. The number of energy states for each subshell is determined by the third quantum num- ber, mp. For an s subshell, there is a single energy state, whereas for p, d, and f sub- shells, three, five, and seven states exist, respectively(Table 2.1). In the absence of an external magnetic field the states within each subshell are identical. However, when a magnetic field is applied these subshell states split, each state assuming a slightly different energy

Quantum Numbers Using wave mechanics, every electron in an atom is characterized by four parame￾ters called quantum numbers. The size, shape, and spatial orientation of an electron’s probability density are specified by three of these quantum numbers. Furthermore, Bohr energy levels separate into electron subshells, and quantum numbers dictate the number of states within each subshell. Shells are specified by a principal quantum number n, which may take on integral values beginning with unity; sometimes these shells are designated by the letters K, L, M, N, O, and so on, which correspond, respectively, to as indicated in Table 2.1. Note also that this quan￾tum number, and it only, is also associated with the Bohr model. This quantum num￾ber is related to the distance of an electron from the nucleus, or its position. The second quantum number, l, signifies the subshell, which is denoted by a lowercase letter—an s, p, d, or f; it is related to the shape of the electron subshell. In addition, the number of these subshells is restricted by the magnitude of n. Allowable subshells for the several n values are also presented in Table 2.1. The number of energy states for each subshell is determined by the third quantum num￾ber, For an s subshell, there is a single energy state, whereas for p, d, and f sub￾shells, three, five, and seven states exist, respectively (Table 2.1). In the absence of an external magnetic field, the states within each subshell are identical. However, when a magnetic field is applied these subshell states split, each state assuming a slightly different energy. ml . n  1, 2, 3, 4, 5, . . . , 2.3 Electrons in Atoms • 19 1.0 0 (a) (b) Orbital electron Nucleus Probability Distance from nucleus Figure 2.3 Comparison of the (a) Bohr and (b) wave￾mechanical atom models in terms of electron distribution. (Adapted from Z. D. Jastrzebski, The Nature and Properties of Engineering Materials, 3rd edition, p. 4. Copyright © 1987 by John Wiley & Sons, New York. Reprinted by permission of John Wiley & Sons, Inc.) quantum number 1496T_c02_15-37 12/20/05 7:19 Page 19 2nd REVISE PAGES

496c0215-3711/10/0510:42page20 REVISED PAGES EQA 20 .Chapter 2 Atomic Structure and Interatomic Bonding Table 2.1 The Number of available electron states in Some of the electron Shells and subshells Principal Shell Number Number of electrons Number n Subshells of States Per Subshell Per Shell 22626 351357 2 Associated with each electron is a spin moment, which must be oriented either up or down. Related to this spin moment is the fourth quantum number, ms, for which two values are possible(+2 and), one for each of the spin orientations. Thus, the Bohr model was further refined by wave mechanics, in which the in- troduction of three new quantum numbers gives rise to electron subshells within each shell. A comparison of these two models on this basis is illustrated, for the hydrogen atom, in Figures 2.2a and 2.2b A complete energy level diagram for the various shells and subshells using the wave-mechanical model is shown in Figure 2.4. Several features of the diagram are worth noting. First, the smaller the principal quantum number, the lower the energy level; for example, the energy of a ls state is less than that of a 2s state, which in turn is lower than the 3s. Second, within each shell, the energy of a subshell level in- reases with the value of the I quantum number. For example, the energy of a 3d state is greater than a 3p, which is larger than 3s. Finally, there may be overlap in energies of the electrons for the various shells and subshells ( From K. M. Ralls, T.H. Courtney, and J. Wulff, Introduction to Materials d-P- Science and Engineering, p. 22. Copyright o 1976 by John Wiley Sons, New York. Reprinted by permission of John Wiley Sons

Associated with each electron is a spin moment, which must be oriented either up or down. Related to this spin moment is the fourth quantum number, for which two values are possible ( and ), one for each of the spin orientations. Thus, the Bohr model was further refined by wave mechanics, in which the in￾troduction of three new quantum numbers gives rise to electron subshells within each shell. A comparison of these two models on this basis is illustrated, for the hydrogen atom, in Figures 2.2a and 2.2b. A complete energy level diagram for the various shells and subshells using the wave-mechanical model is shown in Figure 2.4. Several features of the diagram are worth noting. First, the smaller the principal quantum number, the lower the energy level; for example, the energy of a 1s state is less than that of a 2s state, which in turn is lower than the 3s. Second, within each shell, the energy of a subshell level in￾creases with the value of the l quantum number. For example, the energy of a 3d state is greater than a 3p, which is larger than 3s. Finally, there may be overlap in 1  2 1 2 ms, 20 • Chapter 2 / Atomic Structure and Interatomic Bonding Table 2.1 The Number of Available Electron States in Some of the Electron Shells and Subshells Principal Quantum Shell Number Number of Electrons Number n Designation Subshells of States Per Subshell Per Shell 1 K s 1 22 s 1 2 2 L p 3 6 8 s 1 2 3 M p 3 618 d 5 10 s 1 2 p 3 6 4 N d 5 10 32 f 7 14 Principal quantum number, n Energy 1 s s p s p s p s p df s p s p df d d d f 234567 Figure 2.4 Schematic representation of the relative energies of the electrons for the various shells and subshells. (From K. M. Ralls, T. H. Courtney, and J. Wulff, Introduction to Materials Science and Engineering, p. 22. Copyright 1976 by John Wiley & Sons, New York. Reprinted by permission of John Wiley & Sons, Inc.) © 1496T_c02_15-37 11/10/05 10:42 Page 20 REVISED PAGES