1. CHAPTER 2: BONDING AND PROPERTIES What promotes bonding? What types of bonds are there? How does bonding affect material properties? Much of a material’s behavior can be explained by the phenomena in this chapter. 1

2. Order: Short vs Long Range 2

3. 3 Atomic Structure atom – electrons – 9.11 x 10 -31 kg protons neutrons 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 6.022 x 10 23 molecules or atoms 1 amu/atom = 1g/mol C 12.011 H 1.008 } 1.67 x 10 -27 kg

4. 4 Atomic Structure Valence electrons determine the following properties: 1)Chemical 2)Electrical 3)Thermal 4)Optical

5. 5 Electronic Structure Electrons have wavelike and particulate properties. –This means that electrons are in orbitals defined by a probability. –Each orbital at discrete energy levels is determined by quantum numbers. Quantum # Designation n = principal (energy level-shell)K, L, M, N, O (1, 2, 3, etc.) l = subsidiary (orbitals)s, p, d, f (0, 1, 2, 3,…, n -1) m l = magnetic1, 3, 5, 7 (- l to + l ) m s = spin½, -½

7. 7 Electron Energy States 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...

9. 9 Why? Valence (outer) shell usually not filled completely. Most elements: Electron configuration not stable. SURVEY OF ELEMENTS 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 & Rethwisch 3e.

10. 10 Electron Configurations 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 valence electrons

11. 11 The Periodic Table Columns: Similar Valence Structure 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 1e - accept 2e - O Se Te PoAt I Br He Ne Ar Kr Xe Rn F ClS LiBe H NaMg BaCs RaFr CaKSc SrRbY

12. 12 Electronic Configurations ex: Fe - atomic # = 26 valence electrons 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. 14 Ranges from 0.7 to 4.0, Smaller electronegativityLarger electronegativity Large values: tendency to acquire electrons. Adapted from Fig. 2.7, Callister & Rethwisch 3e. (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. Electronegativity

15. 15 Bonding Energy 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 = 

16. 16 Bond length, r Bond energy, E o Melting Temperature, T m T m is larger if E o is larger. Properties From Bonding: T m r o r Energy r larger T m smaller T m EoEo = “bond energy” Energy r o r unstretched length

17. 17 Coefficient of thermal expansion,   ~ symmetric at r o  is larger if E o is smaller. Properties From Bonding :  =  (T 2 -T 1 )  L L o coeff. thermal expansion  L length,L o unheated, T 1 heated, T 2 r o r larger  smaller  Energy unstretched length EoEo EoEo

18. Primary Bonding Ionic Covalent Metallic 18 Bonding involves the valence electrons. Bonding occurs due to the tendency of the atoms to assume stable electron structures (completely filled outer shells)

19. 19 Occurs between + and - ions. Requires electron transfer. Large difference in electronegativity required. Example: NaCl Ionic Bonding Na (metal) unstable Cl (nonmetal) unstable electron + - Coulombic Attraction Na (cation) stable Cl (anion) stable

21. 21 Ionic bond – metal + nonmetal 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 Mg 2+ 1s 2 2s 2 2p 6 O 2- 1s 2 2s 2 2p 6 [Ne] [Ne]

22. 22 Predominant bonding in Ceramics Adapted from Fig. 2.7, Callister & Rethwisch 3e. (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. Examples: Ionic Bonding Give up electronsAcquire electrons NaCl MgO CaF 2 CsCl

23. 23 C: has 4 valence e -, needs 4 more H: has 1 valence e -, needs 1 more Electronegativities are comparable. Covalent Bonding similar electronegativity  share electrons bonds determined by valence – s & p orbitals dominate bonding Example: polymers, GaAs, InSb, SiC, CH 4 shared electrons from carbon atom shared electrons from hydrogen atoms H H H H C CH 4

24. 24 Ionic-Covalent Mixed Bonding % ionic character = where X A & X B are Pauling electronegativities Ex: MgOX Mg = 1.3 X O = 3.5 %)100( x

25. Metallic Bonding Metallic bonds have up to 3 valence electrons that are not bound to a specific atom. They drift throughout the metal forming a “sea of electrons” or “electron cloud”. The nonvalence electrons and nuclei for the “ion cores”. The free electrons act as a “glue” to hold the ion cores together. These are good conductors of heat and charge (electricity).

26. Secondary Bonding (van der Waals) Interaction between dipoles; dipoles are a separation of charge (+/-). Weaker forces (10kJ/mol) than primary bonding, yet these bonds still influence physical properties. Secondary bonding exists in virtually all atoms and molecules, but their presence may be obscured by primary bonding. 26

27. 27 Permanent dipoles-molecule induced Fluctuating dipoles -general case: -ex: liquid HCl -ex: polymer SECONDARY BONDING 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

29. Bonding energies between 600 and 1500 kJ/mol (or 3 to 8 eV/atom) are considered to be relatively large and will have correspondingly high (large) melting points.

30. 30 Type Ionic Coulombic force Covalent Metallic Secondary Van der Waals 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 Summary: Bonding

31. 31 Ceramics (Ionic & covalent bonding): Large bond energy large T m large E small  Metals (Metallic bonding): Variable bond energy moderate T m moderate E moderate  Summary: Primary Bonds Polymers (Covalent & Secondary): Directional Properties Secondary bonding dominates small T m small E large  secondary bonding