2. Properties Structure Processing Electronic level (subatomic) Atomic (molecular level, chemical composition) Crystal (arrangement of atoms or ions wrt one another) Microstructure (can study with microscopes) Macrostructure (can see with naked eye) “Materials Science” “Materials Engineering”

3.  Valence electrons determine all of the following properties 1) Chemical 2) Electrical 3) Thermal 4) Optical 3

4. What promotes bonding? What types of bonds are there? Bonding  material properties What properties are inferred from bonding? (melting point, thermal expansion, elasticity) Electronic structure & bonding Based on Chapter 2 (Callister)

5.  atom – electrons – 9.11 x 10 -31 kg protons neutrons  Electric charge of proton = 1.6 x 10 -19 C  atomic number = # of protons in nucleus of atom = # of electrons of neutral species  A [=] atomic mass unit = amu = 1/12 mass of 12 C Atomic wt = wt of 6.023 x 10 23 molecules or atoms 1 amu/atom = 1g/mol C 12.011 H 1.008 etc. } 1.67 x 10 -27 kg

6. 6

7. Nucleus: Z = # protons = 1 for hydrogen to 94 for plutonium N = # neutrons Atomic mass A ≈ Z + N Adapted from Fig. 2.1, Callister 6e. electrons: n = principal quantum number n=3 2 1 n labels shells; shells are composed of sub-shells: s, p, d, f, …

8. 8

9.  Electrons have wavelike and particulate properties.  This means that electrons are in orbitals defined by a probability.  Each orbital at discrete energy level determined by quantum numbers. (like x, y, z) Quantum # Designation n = principal (energy level-shell) K, L, M, N, O (1, 2, 3, etc.7) l = subsidiary (form and number of orbitals) s, p, d, f (0, 1, 2, 3,…, n -1) m l = magnetic (Orientation of orbitals) 1, 3, 5, 7 (- l to + l )  Pauli exclusion principle m s = spin +½, -½ aut 9

10. 10 1s1s 2s2s 2p2p K-shell n = 1 L-shell n = 2 3s3s 3p3p M-shell n = 3 3d3d 4s4s 4p4p 4d4d Energy N-shell n = 4 have discrete energy states tend to occupy lowest available energy state. Electrons... Adapted from Fig. 2.4, Callister 7e.

11. 11 Why? Valence (outer) shell usually not filled completely. Most elements: Electron configuration not stable. Electron configuration (stable)... 1s1s 2 2s2s 2 2p2p 6 3s3s 2 3p3p 6 (stable)... 1s1s 2 2s2s 2 2p2p 6 3s3s 2 3p3p 6 3d3d 10 4s4s 2 4p4p 6 (stable) Atomic # 18... 36 Element 1s1s 1 1Hydrogen 1s1s 2 2Helium 1s1s 2 2s2s 1 3Lithium 1s1s 2 2s2s 2 4Beryllium 1s1s 2 2s2s 2 2p2p 1 5Boron 1s1s 2 2s2s 2 2p2p 2 6Carbon... 1s1s 2 2s2s 2 2p2p 6 (stable) 10Neon 1s1s 2 2s2s 2 2p2p 6 3s3s 1 11Sodium 1s1s 2 2s2s 2 2p2p 6 3s3s 2 12Magnesium 1s1s 2 2s2s 2 2p2p 6 3s3s 2 3p3p 1 13Aluminum... Argon... Krypton Adapted from Table 2.2, Callister 7e.

12.  Valence electrons – those in unfilled shells  Filled shells more stable  Valence electrons are most available for bonding and tend to control the chemical properties  example: C (atomic number = 6) 1s 2 2s 2 2p 2 12 valence electrons

13. ex: Fe - atomic # = aut 13 26 valence electrons Adapted from Fig. 2.4, Callister 7e. 1s1s 2s2s 2p2p K-shell n = 1 L-shell n = 2 3s3s 3p3p M-shell n = 3 3d3d 4s4s 4p4p 4d4d Energy N-shell n = 4 1s 2 2s 2 2p 6 3s 2 3p 6 3d 6 4s 2

14. aut 14 Columns: Similar Valence Structure Adapted from Fig. 2.6, Callister 7e. Electropositive elements: Readily give up electrons to become + ions. Electronegative elements: Readily acquire electrons to become - ions. give up 1e give up 2e give up 3e inert gases accept 1eaccept 2e O Se Te PoAt I Br He Ne Ar Kr Xe Rn F ClS LiBe H NaMg BaCs RaFr CaKSc SrRbY

15. aut 15 Ranges from 0.7 to 4.0, Smaller electronegativityLarger electronegativity Large values: tendency to 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.

16. donates accepts electrons electrons Dissimilar electronegativities ex: MgOMg 1s 2 2s 2 2p 6 3s 2 O 1s 2 2s 2 2p 4 [Ne] 3s 2 aut 16 Mg 2+ 1s 2 2s 2 2p 6 O 2- 1s 2 2s 2 2p 6 [Ne] [Ne]

17. aut 17 Occurs between + and - ions. Requires electron transfer. Large difference in electronegativity required. Example: NaCl Na (metal) unstable Cl (nonmetal) unstable electron + - Coulombic Attraction Na (cation) stable Cl (anion) stable

18. 18

19. 19  Energy – minimum energy most stable  Energy balance of attractive and repulsive terms Attractive energy E A Net energy E N Repulsive energy E R Interatomic separation r r A n r B E N = E A + E R =  Adapted from Fig. 2.8(b), Callister 7e. aut

20. 20 Predominant bonding in Ceramics 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. Give up electronsAcquire electrons NaCl MgO CaF 2 CsCl

21. aut 21 C: has 4 valence e -, needs 4 more H: has 1 valence e -, needs 1 more Electronegativities are comparable. Adapted from Fig. 2.10, Callister 7e. similar electronegativity  share electrons bonds determined by valence – s & p orbitals dominate bonding Example: CH 4 shared electrons from carbon atom shared electrons from hydrogen atoms H H H H C CH 4

22. Molecules with nonmetals Molecules with metals and nonmetals Elemental solids (RHS of Periodic Table) Compound solids (about column IVA) Adapted from Fig. 2.7, Callister 6e. (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. 3.5 Prevalent in ceramics and polymers

23. Arises from a sea of donated valence electrons (1, 2, or 3 from each atom). Primary bond for metals and their alloys  high electrical conductivity. Why? What about ionic/covalent? Adapted from Fig. 2.11, Callister 6e.

24.  Much easier to deform materials with metallic than with ionic bonding. Why? Sliding atom planes over each other (deformation) very unfavorable energetically in ionic solids!  metals are ductile & ceramics (ionic) are brittle

25.  Metallic Bond -- delocalized as electron cloud  Ionic-Covalent Mixed Bonding % ionic character = where X A & X B are Pauling electronegativities aut 25 %)100( x Ex: MgOX Mg = 1.3 X O = 3.5

26. aut 26 Arises from interaction between dipoles Permanent dipoles-molecule induced Fluctuating dipoles -general case: -ex: liquid HCl -ex: polymer Adapted from Fig. 2.13, Callister 7e. Adapted from Fig. 2.14, Callister 7e. asymmetric electron clouds +-+- secondary bonding HHHH H 2 H 2 secondary bonding ex: liquid H 2 H Cl H secondary bonding secondary bonding +-+- secondary bonding

27. aut 27 Type Ionic Covalent Metallic Secondary Bond Energy Large! Variable large-Diamond small-Bismuth Variable large-Tungsten small-Mercury smallest Comments Nondirectional (ceramics) Directional (semiconductors, ceramics polymer chains) Nondirectional (metals) Directional inter-chain (polymer) inter-molecular

28. aut 28 Bond length, r Bond energy, E o Melting Temperature, T m T m is larger if E o is larger. r o r Energy r larger T m smaller T m EoEo = “bond energy” Energy r o r unstretched length

29. aut 29 Coefficient of thermal expansion,   ~ symmetry at r o  is larger if E o is smaller. =  (T 2 -T 1 )  L L o coeff. thermal expansion  L length,L o unheated, T 1 heated, T 2 r o r smaller  larger  Energy unstretched length EoEo EoEo

30. Why does thermal expansion occur?  As atoms vibrate, easy to move apart, but difficult to move closer  As temperature increases, average bond length increases. Bond Energy vs. Distance Curve is asymmetric

31. Elastic modulus, E E ~ curvature at r o E is larger if curvature is larger. E similar to spring constant

32. aut 32 Ceramics (Ionic & covalent bonding): Metals (Metallic bonding): Polymers (Covalent & Secondary): Large bond energy large T m large E small  Variable bond energy moderate T m moderate E moderate  Directional Properties Secondary bonding dominates small T m small E large  secondary bonding