[工学]材料科学基础英文版课件2

Chapter 1 Atomic Structure and Interatomic Bonding 11. Introduction Some of the important properties of solid materials depend on geometrical atomic arrangements. The properties of materials are controllable and can be tailored to the needs of a given application by controlling their structure and composition. We can examine and describe the structure of materials at different levels.2 Subatomic level: Electronic structure of individual atoms that defines interaction among atoms (interatomic bonding). Atomic level: Arrangement of atoms in materials. Nanostructure: the structure of material at a length-scale of 100nm. Microstructure: the structure of material at a length-scale of 10 to 1000nm. Macrostructure: the structure of a material at a macroscopic level where the length-scale is 1000,000 nm. 3Length-scales Angstrom = 1 = 1/10,000,000,000 meter = 10-10 m Nanometer = 10 nm = 1/1,000,000,000 meter = 10-9 m Micrometer = 1m = 1/1,000,000 meter = 10-6 m Millimeter = 1mm = 1/1,000 meter = 10-3 m Interatomic distance a few A human hair is 50 m Elongated bumps that make up the data track on CD are 0.5 m wide, minimum 0.83 m long, and 125 nm high 45 This micrograph, which represents the surface of a gold specimen, was taken with a atomic force microscope (AFM). Individual atoms for this (111) crystallographic surface plane are resolved. 6 Amorphous: lack a long-range ordering of atoms or ions. Crystalline: exhibit periodic arrangements of atoms or ions. The long-range atomic order is in the form of atoms or ions arranged in a three dimensional pattern that repeats over much larger distances (from 100 nm to up to few cm). Short-range atomic arrangements: the atoms of ions show a particular order only over relatively short distances. 72. Atomic Structure Atomic structure influences how atoms are bonded together. An understanding of this helps categorize material as metals, semiconductors, ceramics, or polymers. Atoms = nucleus (protons and neutrons) + electrons8 Charges: Electrons and protons have negative and positive charges of the same magnitude, 1.6 10-19 Coulombs. Neutrons are electrically neutral. Masses: Protons and Neutrons have the same mass, 1.67 10-27 kg. Mass of an electron is much smaller, 9.11 10-31 kg and can be neglected in calculation of atomic mass. The atomic mass (A) = mass of protons + mass of neutrons # protons gives chemical identification of the element # protons = atomic number (Z) # neutrons defines isotope number9 There are 92 naturally occurring elements, each identified by the atomic number (number of protons) and atomic weight (which includes the weight of the neutrons) and represents an average over the various isotopes that may exist). The atomic weight has units of grams per mole. A mole is the amount of material that corresponds to the atomic weight. A mole is the amount of matter that has a mass in grams equal to the atomic mass in amu of the atoms (A mole of carbon has a mass of 12 grams). The number of atoms in a mole is called the Avogadro number, Nav = 6.023 1023.10THE PERIODIC TABLE113. Electronic in AtomsBohr atomic model Orbital electron12 Quantum numbers are the numbers in an atom that assign electrons to discrete energy levels. Principal quantum number n: is assigned integral values 1,2,3,4,5that refer to the quantum shell to which the electron belongs. Orbital (Azimuthal) quantum number l: describe the energy levels in each quantum shell. l=0,1,2,n-1. Magnetic quantum number ml: describes the number of energy levels for each orbital quantum number. l, +l Spin quantum number ms: assigned values +1/2 and -1/2 and reflects the different electronic spins The number of possible energy levels is determined by the first three quantum numbers. 13 The complete set of quantum numbers for each of the 11 electrons in sodium.3s1 electron 11 n=3, l=0, ml=0, ms=+1/2 or -1/2electron 10 n=2, l=1, ml=+1, ms= -1/2 electron 9 n=2, l=1, ml=+1, ms=+1/2 2p6 electron 8 n=2, l=1, ml=0, ms= -1/2electron 7 n=2, l=1, ml=0, ms=+1/2electron 6 n=2, l=1, ml=-1, ms= -1/2electron 5 n=2, l=1, ml=-1, ms=+1/22s2 electron 4 n=2, l=0, ml=0, ms= -1/2electron 3 n=2, l=0, ml=0, ms=+1/21s2 electron 2 n=1, l=0, ml=0, ms= -1/2electron 1 n=1, l=0, ml=0, ms=+1/214The pattern used to assign electrons to energy levelsl=0(s)l=1(p)l=2(d)l=3(f)l=4(g)l=5(h)n=1(K) 2n=2(L) 2 6n=3(M) 2 6 10n=4(N) 2 6 10 14n=5(O) 2 6 10 14 18n=6(P) 2 6 10 14 18 22Note: 2,6,10,14,refer to the number of electrons in the energy level.15This sequence emphasizes the relative energy levels of the shells so that the outer, higher energy and more asymmetric d levels may fill after the inner s level of the next shell and more asymmetric levels may fill after the inner p and even the d level of the next shell. 16 have complete s and p subshells tend to be unreactive.Stable electron configurations.Stable Electron Configurations17 Valence: The number of electrons in an atom that participate in bonding or chemical reactions. Usually, the valence is the number of electrons in the outer s and p energy levels. Mg: 1s22s22p6 valence=2 Al: 1s22s22p6 valence=3 Ge: 1s22s22p63s23p63d10 valence=4 If an atom has a valence of zero, the element is inert (non-reactive). 1s22s22p6 Valence also depends on the immediate environment surrounding the atom or the neighboring atoms available for bonding. For Example, P2O3, PH3.3s23s23p14s24p23s23p618 Electronegativity is the quantitative description of an atoms desire to gain or lose an electron. It is a function of the number of electrons in an atoms valence shell, and the distance of the shell from the nucleus. For example, chlorine, with 7 valence electrons, is very eager to gain an electron to fill its outer shell, while sodium will easily give up its 1 valence electron.4. The Periodic Table19The Periodic Table20Electronegativity increases as you go right and up on the periodic table. 21 Columns: Similar Valence StructureElectropositive elements:Readily give up electronsto become + ions.Electronegative elements:Readily acquire electronsto become - ions.The Periodic TableHe Ne Ar Kr Xe Rn inert gasesaccept 1eaccept 2egive up1egive up2egive up3eF Li Be Metal Nonmetal Intermediate H Na Cl Br I At O S Mg Ca Sr Ba Ra K Rb Cs Fr Sc Y Se Te Po 225. Atomic Bonding in solid The attractive force varies inversely with the square of the distance between the atoms for all of the various types of bonds. This force pulls the atoms together with a greater force as they get closer together. A repulsive force arises when the electron clouds of negative charge meet. This force increases much faster with distance, with an exponent in the range 6-9 that depends on the particular atom and its electron shells. When the sum of these forces is zero, the distance between the atoms is at the equilibrium value 23 Both the attractive and repulsive forces increase as the atoms are brought closer together, and the sum of the two is zero at the equilibrium point. The slope of the force vs. distance curve at the equilibrium point defines the force needed to pull the atoms slightly apart and is the slope of the stress vs. strain curve. This slope is the modulus of elasticity.24 The Bond energy vs. distance curve is the integral of the bond force curve. The lowest energy is the equilibrium point. 25 The depth of the minimum is the total bond strength which reflects the energy required to pull the atoms completely apart. The deeper the potential minimum, the higher the melting temperature. The bond energy curve is asymmetric. This is why most materials expand when heated. Because of the asymmetry of the bond energy curve, the average distance between atoms increases with temperature. The narrower the potential minimum, the lower the expansion coefficient. 26Atomic Bonding in solid There are four principal kinds of bonds that form between atoms: Metallic bond Ionic bond Covalent bond Van der Waals bondStrong BondWeak BondThe forces between atoms are electrostatic dependent directly on the electrons that surround the atoms. The different bond types are characterized by how electrons are shar