Engineering Materials: Lecture 3

1. Engineering Materials: Lecture #3 ENGR 151

2. Quiz Announcement Quiz #1 This Thursday, Feb. 18 Open note/book Atomic Structure, Section 2.1

3. Types of Atomic Bonding Three types of PRIMARY bonds: IONIC, COVALENT, METALLIC. Nature of the bond depends on the electron structures of the bonding atoms AND type of bond depends on the tendency of atoms to assume stable electron structures (completely filling outermost electron shell)

4. Ionic Bonding Always found in compounds that have metallic and nonmetallic elements (elements at the horizontal extremities of the periodic table). Metallic elements easily give up valence electrons to nonmetallic atoms. In this process, all atoms acquire stable or inert gas configurations and electrical charges (becoming ions)

5. Ionic Bonding The attractive bonding forces within ionic bonds are COULOMBIC (positive and negative ions attracting one another). The ion of an element has a different atomic size than that of an atom of the element. (Appendix 2)

6. Ionic Bonding The coulombic attraction is a function of the interatomic distance : Fc = (-K/a2) K = k0(Z1q)(Z2q) a = interatomic distance (meters) ko = proportionality constant (9 x 109 V�m/C) Z1 = valence of ion Z2 = valence of ion q = electronic charge (1.602 x 10-19 C)

7. Ionic Bonding Ionic bonding is considered nondirectional (magnitude of bond is equal in all directions around the ion) Therefore, for an ionic material to be stable, all positive ions must be surrounded by negative charged ions, and vice versa.

8. Coordination number CN is the number of adjacent ions surrounding a reference ion (or atom) CN depends on the sizes of surrounding ions Radius Ratio: r/R r = radius of smaller ion R = radius of larger ion

9. CN and force Exercises Derive the minimum radius ratio for threefold coordination Practice Problem 2.5: Calculate the coulombic force of attraction between the Mg2+ - O2- ion pair. What is the repulsive force in this case?

10. Ionic Bonding Energies Ionic materials are hard, brittle, electrically and thermally insulative Bonding energies ranging from 600 and 1500kJ are relatively large, as reflected in high melting temperatures. Sodium Chloride (NaCl, bonding energy of 640 kJ/mol, melts @ 801 degrees Celsius) Magnesium Oxide (MgO, bonding energy 1000 kJ/mol, melts @ 2800 degrees Celsius)

11. Covalent Bonding Stable electron configurations are assumed by the sharing of electrons between adjacent atoms. Atoms contribute at least one electron to the bond, and the shared electrons are considered to belong to both atoms Covalent bonds are directional (between specific atoms)

12. Covalent Bonding -Examples include: Hydrogen (H2), Chloride (Cl2), Flouride (F2), Water (H20), diamond, silicon, germanium, elemental solids located on right hand side of the periodic table, Gallium Arsenide (GaAs), indium antimonide (InSb), and silicon carbide (SiC).

13. Covalent Bonding The number of covalent bonds for an atom is determined by the number of valence electrons. For N' valence electrons, an atom can covalently bond with at most 8 - N' other atoms.

14. Problem Using Periodic Table, determine the number of covalent bonds that are possible for atoms of the following elements: silicon, bromine, nitrogen, and sulfur

15. Answer Silicon: 8 � 4 = 4 Bromine: 8 � 7 = 1 Nitrogen: 8 � 5 = 3 Sulfur: 8 � 6 = 2

16. Covalent Bonding Energies Can be strong or weak. (diamond, very high melting temperature, 3550 degrees C; bismuth, 270 degrees C)

17. Metallic Bonding Metallic materials have one, two, or at most three valence electrons which are not bound to any atom within the solid and can drift throughout the metal Creates a �sea of electrons� (electrons that belong to the entire metal) Remaining nonvalence electrons and their atomic nuclei are called ION CORES (have net positive charge equal in magnitude to total valence charge per atom)

18. Metallic Bonding Free electrons shield ion cores from repulsive forces and act as glue to hold ion cores together Metallic bond is nondirectional Free electrons allow metal to be good conductors of heat and electricity

19. METALLIC BONDING ENERGIES Can be strong or weak (Mercury, 68 kJ/mol, -39 degrees Celsius; Tungsten, 850 kJ/mol, 3410 degrees Celsius)

20. SECONDARY BONDING Very weak bonds (10 kJ/mol) Exist between all atoms and molecules but obscured if other primary bonding is occurring Found in: -inert gases because of their stable electron structures -between molecules that are covalently bonded -arise from dipoles (separation of positive and negative portions of an atom or molecule).

21. Secondary Bonding

22. Fluctuating Induced Dipole Bonds Atom structure is not symmetric and positively charged nucleus creates dipole with electron cloud on opposite side of atom.

23. Polar Molecule-Induced Dipole Bonds polar molecules: asymmetrical arrangements of positively and negatively charged regions within molecules. (Figure 2.23, pg. 49)

24. Permanent Dipole Bonds Hydrogen Bonding: a type of secondary bonding occurring in molecules that have hydrogen Adjacent polar molecules attracting each other

25. Case History How do we know what materials look like on the atomic level? Can we see them? YES! X-ray diffraction Transmission electron microscopy (TEM): electron beam that passes through the specimen (1,000,000x) Scanning tunneling microscopy (STM): electrons flow down from tip of tiny probe Atomic force microscopy (AFM): height data taking from tiny probe

26. Mountain or Gold Cluster?

27. STM image of diamond

28. AFM � carbon nanotubes

29. The Solid State Crystalline v. amorphous Crystalline: atoms are situated in a repeating array over large atomic distances (repetitive three-dimensional pattern) Amorphous: materials that do not crystallize Unit cells, lattices Single crystal v. polycrystalline Grain boundaries

30. The Nature of Metals Metals bound by �sea of electrons� Movement Dislocations: linear defect in which atoms are misaligned Strengthening mechanisms: Solid solution (alloy with impurity atoms) Precipitation (formation of small particles in original material) Dispersion (fine particles of hard material) Quench (rapid cooling)

31. The Nature of Ceramics Metals combined with nonmetals Ionic and covalent bonding Mechanically brittle Electrically insulating

32. The Nature of Polymers Long-chain repeating molecules Carbon backbone Intra- and intermolecular forces Covalent v. Van der Waals forces Example: PE

33. The Nature of Composites Combination of two or more materials Natural example: Wood (cellulose & lignin) Plywood (laminar composite) Synthetic: Pearlitic steel: ferrite & cementite Properties: high strength, low weight, crack-sensitive