2019 Fall
Introduction to Materials Science and Engineering
09. 10. 2019 Eun Soo Park
Office: 33‐313 Telephone: 880‐7221 Email: [email protected] Office hours: by appointment 1 Materials Science and Engineering
Processing •Sintering • Heat treatment • Thin Film Structure • Melt process • Mechanical • Atomic • Crystal • Microstructure • Macrostructure Properties • Mechanical • Electrical • Magnetic Theory & • Thermal
Design • Optical 2 CHAPTER 2: BONDING AND PROPERTIES
ISSUES TO ADDRESS...
• What promotes bonding?
• What types of bonds are there?
• What properties are inferred from bonding?
3 Contents for previous class Atomic Structure 2.2 Fundamental concepts
2.3 Electrons in atoms a. atomic models Bohr’s model + Wave-mechanical model 전자는 파동성과 입자성을 동시에 갖는다 가정= 전자구름 b. Quantum #s
c. Electron configurations
2.4 Periodic table 모든 원소는 주기율표 상의 전자 배위에 의해 분류 4 : 특성의 규칙적인 변화 양상 확인 가능 2.3 Electrons in atoms a. atomic models Bohr vs. wave mechanical model
Wave Bohr mechanical model model
Bohr Wave model mechanical model
Electron position is described by Bohr energy levels to be separated a probability distribution into electron subshells or electron cloud described by quantum numbers 5 2.3 Electrons in atoms b. quantum numbers Meaning of quantum numbers
n determines the size l determines the shape
ml determines the orientation
6 2.3 Electrons in atoms c. Electronic configurations Electronic Configurations
ex: Fe - atomic # = 26 1s2 2s2 2p6 3s2 3p6 3d 6 4s2
4d 4p N-shell n = 4 valence electrons 3d 4s
Energy 3p M-shell n = 3 3s Adapted from Fig. 2.4, Callister 7e. 2p L-shell n = 2 2s
1s K-shell n = 1 7 1s1 1s2 2s1 2s2 1s 2p1 . 2s 2p . . 3s 3p 3d 2p6 3s1 4s 4p 4d 4f 3s2 3p 4s 3d 4p
8 SURVEY OF ELEMENTS • Most elements: Electron configuration not stable. Element Atomic # Electron configuration Hydrogen 1 1s 1 Helium 2 1s 2 (stable) Lithium 3 1s 22s 1 Beryllium 4 1s 22s2 Boron 5 1s 22s 22p 1 Adapted from Table 2.2, Carbon 6 1s 22s 22p 2 Callister 7e...... Neon 10 1s 22s 22p 6 (stable) Sodium 11 1s 22s 22p 63s 1 Magnesium 12 1s 22s 22p 63s 2 Aluminum 13 1s 22s 22p 63s 23p 1 ...... Argon 18 1s 22s 22p 63s 23p 6 (stable) ...... Krypton 36 1s 22s 22p 63s 23p 63d 10 4s 24p 6 (stable)
• Why? Valence (outer) shell usually not filled completely. 9 Chapter 2.4 모든 원소는 주기율표 상의 전자 배위에 의해 분류
The Periodic Table Contents for today’s class
Atomic Bonding in Solids
- Primary interatomic bonds
a. Ionic / b. covalent / c. metallic
- Secondary bonds
a. Van der Waals / b. Hydrogen
- Properties from bonding
11 Two fundamental types of bonding:
primary bonds: strong atom-to-atom secondary bonds: much weaker. It is the attractions produced by changes in attraction due to overall “electric fields”, electron position of the valence e– . often resulting from electron transfer in Example : covalent atom between two primary bonds. Example: intramolecular hydrogen atoms bond between H2 molecules gas e H + e + H + + H H e e
H Very weak intermolecular Hydrogen attraction Hydrogen Molecule + + + + Molecule
Highest Probability density of two electrons between atoms forms very strong intramolecular covalent bond 12 Chapter 2.6 Primary interatomic bonds
a. Atomic bonding When atoms of far- & near-closed shell structure are brought together ...
e e e e e e e e e e 3+ 9+ 3+ 9+ Li F Li e F e e e e e e e e e e e e e
Atoms of far-closed shell structure & near-closed one tend to lose & gain electrons, respectively
Electronegativity by L. Pauling
Excess charge induced by the transfer of electrons are compensated by the presence of ions of opposite sign Ionic bonding 14 The Periodic Table • Columns: Similar Valence Structure inert gases give up 1e accept 2e accept 1e H give up 2e He Li Be O F Ne
Na Mg give up 3e Adapted from S Cl Ar Fig. 2.6, K Ca Sc Se Br Kr Callister 7e. Rb Sr Y Te I Xe Cs Ba Po At Rn Fr Ra
Electropositive elements: Electronegative elements: Readily give up electrons Readily acquire electrons to become + ions. to become - ions. 15 Electronegativity • Ranges from 0.7 to 4.0, • Large values: tendency to acquire electrons.
Smaller electronegativity Larger electronegativity Adapted from Fig. 2.7, Callister 7e. (Fig. 2.7 is adapted from Linus Pauling, The Nature of the Chemical Bond, 3rd edition, Copyright 1939 and 1940, 3rd edition. Copyright 1960 by Cornell University.
16 Ionic bond = metal + nonmetal
Accepts Donates electrons electrons
Dissimilar electronegativities ex: MgO Mg 1s2 2s2 2p6 3s2 O 1s2 2s2 2p4 [Ne] 3s2
Mg2+ 1s2 2s2 2p6 O2- 1s2 2s2 2p6 [Ne] [Ne]
17 Ionic Bonding • Occurs between + and – ions • Requires electron transfer • Large difference in electronegativity required • Example: NaCl
Na (metal) Cl (nonmetal) unstable unstable electron
Na (cation) + - Cl (anion) stable Coulombic stable Attraction
18 Examples: Ionic Bonding • Predominant bonding in Ceramics NaCl MgO CaF2 CsCl
Give up electrons Acquire electrons Adapted from Fig. 2.7, Callister 7e. (Fig. 2.7 is adapted from Linus Pauling, The Nature of the Chemical Bond, 3rd edition, Copyright 1939 and 1940, 3rd edition. Copyright 1960 by Cornell University.
19 20 21 • When voltage is applied to an ionic material, entire ions must move to cause a current to flow. Ion movement is slow and the electrical conductivity is poor
22 Ionic Crystal
23 b. COVALENT Covalent Bonding • similar electronegativity share electrons • bonds determined by valence – s & p orbitals dominate bonding Shared electrons • Example: CH4 H from carbon atom CH4 C: has 4 valence e-, needs 4 more H C H H: has 1 valence e-, needs 1 more
H Shared electrons Electronegativities are comparable. from hydrogen atoms
25 Adapted from Fig. 2.10, Callister 7e. Silicon
26 Covalent Bonding Si
vs.
• The tetrahedral structure of silica (Si02), which contains covalent bonds between silicon and oxygen atoms 27 Covalent Bonding • Electron sharing
Cl2
diamond 2- SiO4
directional Example : Covalent Bonding
H2O H 2 F2 C(diamond) H He - 2.1 SiC column IVA Cl2 Li Be C O F Ne 1.0 1.5 2.5 2.0 4.0 - Na Mg Si Cl Ar 0.9 1.2 1.8 3.0 - K Ca Ti Cr Fe Ni Zn Ga Ge As Br Kr 0.8 1.0 1.5 1.6 1.8 1.8 1.8 1.6 1.8 2.0 2.8 - Rb Sr Sn I Xe 0.8 1.0 1.8 2.5 - Cs Ba Pb At Rn 0.7 0.9 1.8 2.2 - Fr Ra 0.7 0.9 GaAs • molecules with nonmetals • molecules with metals and nonmetals • elemental solids (RHS of Periodic Table) • compound solids (about column IVA)
29 Covalent Bonding • Bond energy curve
• Strong directional nature of bonding • Wide range of hardness & melting point ex. High (Diamond) or low (Bismuth) melting point • Low electrical conductivities at low temperatures when specimens are pure 30 Ionic vs. Covalent Bonding
• many compounds-partially ionic and partially covalent • degree of bond type - electronegativity • a large difference in electronegativity largely ionic • similar electronegativity largely covalent
% Ionic character {1- exp[-(0.25)(X A X B )]}100 31 c. METALLIC • delocalized electron Metallic Bonding
• Arises from a sea of donated valence electrons (1, 2, or 3 from each atom) • Primary bond for metals and their alloys
Free electrons act as a “glue” to hold the ion core 33 Metallic Bonding
• The metallic bond forms when atoms give up their valence electrons, which then form an electron sea. • The positively charged atom cores are bonded by mutual attraction to the negatively charged electrons.
34 Metallic Bonding
• When voltage is applied to a metal, the electrons in the electron sea can easily move and carry a current.
35 Metallic Bonding
• mechanical property – What do you expect from oxides and metals?
Oxides
brittle
Metals
ductile
36 Primary Bonding • Metallic Bond -- delocalized as electron cloud
• Ionic-Covalent Mixed Bonding
𝑋 𝑋 1 𝑒𝑥𝑝 100% % ionic character = 4
where XA & XB are Pauling electronegativities
Ex: MgO XMg = 1.3, XO = 3.5
. . %Ionic Character = 1 𝑒𝑥𝑝 100% = 70.2% ionic
37 Chapter 2.7 Secondary bonding Van der Waals Hydrogen SECONDARY BONDING Arises from interaction between dipoles • Fluctuating dipoles asymmetric electron ex: liquid H2 clouds H2 H2 +- +- H H H H secondary secondary bonding Adapted from Fig. 2.13, Callister 7e. bonding • Permanent dipoles-molecule induced -general case: +- secondary +- bonding Adapted from Fig. 2.14, Callister 7e. -ex: liquid HCl secondary H Cl bonding H Cl
-ex: polymer secondary bonding
39 a. Van der Waals bonding (dipole bonding)
e Although electrons have tendency of being separated as far as possible due to e-e He2+ repulsion, electrons are constantly in motion e It follows that electrons could get close enough to induce a “electric dipole moment” at atomistic level 2+ This tendency is expected to be more He significant as the number of electrons increases e e Temporal bonding due to the induced electric dipole
van der Waals bonding 40 Van der Waals Bonding induced dipole permanent dipole (polar molecule)
41 b. Hydrogen bonding When one of the components of covalent bonding is hydrogen ...
e e e e e e e e + H e 9+ 9+ F e H+ e F e e e e e e e e e
Since hydrogen atom has only one electron, there is no electron left for the formation of closed shell Bare proton is exposed without being shielded by electrons ... Strong ionic character develops locally about hydrogen atom ...
... Strong bonding develops locally ... Hydrogen Bonding
H H H H H H O S N H2OH2S NH3 H
• Strongest secondary bonding • Positively charged Hydrogen ion forms a bridge between two negatively charged ions 43 Hydrogen Bonding Ice H2O
open structure lower density Van der Waals and Hydrogen bonding
Polyethylene Nylon-6.6 & Kevlar
45 Materials-Bonding Classification
46 Bonding compared
47 Chapter 2.5 Bonding force and energies Bonding Energy : potential well concept • Energy – minimum energy most stable – Energy balance of attractive and repulsive terms A B E = E + E = N A R r r n
Repulsive energy ER
Interatomic separation r
Net energy EN
Attractive energy EA
48 Adapted from Fig. 2.8(b), Callister 7e. Properties From Bonding: Tm • Bond length, r • Melting Temperature, Tm Energy r
• Bond energy, Eo r o r Energy smaller Tm
unstretched length r larger T o r m E = o Tm is larger if Eo is larger. “bond energy”
49 Properties from Bonding: Thermal Expansion Thermal expansion asymmetric nature of the energy well Broad well (generally more asymmetric) larger expansion
50 Properties From Bonding : • Coefficient of thermal expansion,
length, Lo coeff. thermal expansion unheated, T1 L L = (T2 -T1) heated, T2 Lo
• ~ symmetry at ro Energy
unstretched length r o r is larger if Eo is smaller. Larger Eo
Eo Smaller 51 Properties From Bonding :
Energy
unstretched length r o r is larger if E is smaller. Larger E1
Eo smaller 52 MATERIALS AND PACKING
Crystalline materials... • atoms pack in periodic, 3D arrays • typical of: -metals -many ceramics
-some polymers crystalline SiO2 Adapted from Fig. 3.18(a), Callister 6e. Si Oxygen Noncrystalline materials... • atoms have no periodic packing • occurs for: -complex structures -rapid cooling
"Amorphous" = Noncrystalline noncrystalline SiO2 Adapted from Fig. 3.18(b), Callister 6e. 53 ENERGY AND PACKING • Non dense, random packing Energy typical neighbor bond length
typical neighbor r bond energy
• Dense, regular packing Energy typical neighbor bond length
typical neighbor r bond energy Dense, regular-packed structures tend to have lower energy. 54 Contents for today’s class SUMMARY: BONDING
Type Bond Energy Comments Ionic Large! Nondirectional (ceramics)
Variable Directional Covalent large-Diamond semiconductors, ceramics small-Bismuth polymer chains)
Variable Metallic large-Tungsten Nondirectional (metals) small-Mercury Directional Secondary smallest inter-chain (polymer) inter-molecular 55 Contents for today’s class Summary: Properties from Bonds Ceramics Large bond energy (Ionic & covalent bonding): large Tm large E small Metals Variable bond energy (Metallic bonding): moderate Tm moderate E moderate
Polymers Directional Properties (Covalent & Secondary): Secondary bonding dominates small Tm small E large
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