PLT104

1. Introduction to Material Engineering PLT104

2. THE MARS ROVERS; SPIRIT AND OPPORTUNITY Spirit and Opportunity are made up of materials such as * Metals * Ceramics * Composites * Polymers * Semiconductors

3. WHAT ARE MATERIALS? • Materials may be defined as substance of which something is composed or made. • Materials are obtained from earth crust and atmosphere. • Examples :- • Silicon and Iron constitute 27.72 and 5.00 percentage of weight of earths crust respectively. • Nitrogen and Oxygen constitute 78.08 and 20.95 percentage of dry air by volume respectively.

4. WHY THE STUDY OF MATERIALS ARE IMPORTANT? • Production and processing of materials constitute a large part of our economy. • Engineers choose materials to suite design. • New materials might be needed for some new applications. • Example:- High temperature resistant materials. • Space station and Mars Rovers should sustain conditions in space. * High speed, low temperature, strong but light. • Modification of properties might be needed for some applications. • Example :- Heat treatment to modify properties.

5. MATERIAL SCIENCE AND ENGINEERING? • Materials science deals with basic knowledge about the internal structure, properties and processing of materials. • Materials engineering deals with the application of knowledge gained by materials science to convert materials to products. Materials Science and Engineering Materials Science Materials Engineering Applied Knowledge of Materials Basic Knowledge of Materials Resultant Knowledge of Structure and Properties

6. TYPES OF MATERIALS • Metallic Materials • Composed of one or more metallic elements. Example:- Iron, Copper, Aluminum. • Metallic element may combine with nonmetallic elements. Example:- Silicon Carbide, Iron Oxide. • Inorganic and have crystalline structure. • Good thermal and electric conductors. Metals and Alloys Ferrous Eg: Steel, Cast Iron Nonferrous Eg:Copper Aluminum

7. TYPES OF MATERIALS • Polymeric (Plastic) Materials • Organic giant molecules and mostly noncrystalline. • Some are mixtures of crystalline and noncrystalline regions. • Poor conductors of electricity and hence used as insulators. • Strength and ductility vary greatly. • Low densities and decomposition temperatures. Examples :- Poly vinyl Chloride (PVC), Polyester. Applications:- Appliances, DVDs, Fabrics etc.

8. TYPES OF MATERIALS • Ceramic Materials • Metallic and nonmetallic elements are chemically bonded together. • Inorganic but can be either crystalline, noncrystalline or mixture of both. • High hardness, strength and wear resistance. • Very good insulator. Hence used for furnace lining for heat treating and melting metals. • Also used in space shuttle to insulate it during exit and reentry into atmosphere. • Other applications : Abrasives, construction materials, utensils etc. • Example:- Porcelain, Glass, Silicon nitride.

9. TYPES OF MATERIALS • Composite Materials • Mixture of two or more materials. • Consists of a filler material and a binding material. • Materials only bond, will not dissolve in each other. • Mainly two types :- • Fibrous: Fibers in a matrix • Particulate: Particles in a matrix • Matrix can be metals, ceramic or polymer Examples :- • Fiber Glass ( Reinforcing material in a polyester or epoxy matrix) • Concrete ( Gravels or steel rods reinforced in cement and sand) Applications:- Aircraft wings and engine, construction.

10. TYPES OF MATERIALS • Electronic Materials • Not Major by volume but very important. • Silicon is a common electronic material. • Its electrical characteristics are changed by adding impurities. Examples:- Silicon chips, transistors Applications :- Computers, Integrated Circuits, Satellites etc.

11. COMPETITION AMONG MATERIALS Example:- • Materials compete with each other to exist in new market • Over a period of time usage of different materials changes depending on cost and performance. • New, cheaper or better materials replace the old materials when there is a breakthrough in technology Figure 1.14 Predictions and use of materials in US automobiles.

12. CASE STUDY: MATERIAL SELECTION Steel and alloys Wood Carbon fiber Reinforced plastic Aluminum alloys Ti and Mg alloys • Problem: Select suitable material for bicycle frame and fork. Low cost but Heavy. Less Corrosion resistance Light and strong. But Cannot be shaped Very light and strong. No corrosion. Very expensive Light, moderately Strong. Corrosion Resistance. expensive Slightly better Than Al alloys. But much expensive Cost important? Select steel Properties important? Select CFRP

13. ATOMIC STRUCTURE } 1.67 x 10-27 kg Fundamental concept • Atom – electrons – 9.11 x 10-31 kg protonsneutrons • Atomic number = # of protons in nucleus of atom = # of electrons of neutral species • Isotopes - atoms of some elements have two or more different atomic masses(the number of protons is the same for all atoms of a given element, the number of neutrons (N) may be variable.) • A [=] atomic mass unit = amu = 1/12 mass of 12C1 mole = 6.023 x 1023 molecules or atoms1 amu/atom = 1g/mol • Quantum mechanical principle - A set of principles and laws • that govern systems of atomic and subatomic entities

14. ATOMIC MODEL Bohr atomic model - Electrons are assumed to revolve around the atomic nucleus in discrete orbital and the position of any particular electron is more or less well defined in terms of its orbital. An electron may change energy, but in doing so it must make a quantum jump either to an allowed higher energy (with absorption of energy) or to a lower energy (with emission of energy).

15. ATOMIC STRUCTURE • Valence electrons determine all of the following properties • Chemical • Electrical • Thermal • Optical

16. ELECTRON ENERGY STATES Bohr model was eventually found to have some significant limitations because of its inability to explain several phenomena involving electrons. A resolution was reached with a wave-mechanical model, in which the electron is considered to exhibit both wave-like and particle-like characteristics (a) The first three electron energy states for the Bohr hydrogen atom. (b) Electron energy states for the first three shells of the wave-mechanical hydrogen atom.

17. ELECTRONIC STRUCTURE • 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.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) ml = magnetic 1, 3, 5, 7 (-l to +l) ms = spin ½, -½

18. ELECTRONS IN ATOMS Comparison of the • Bohr • Wave-mechanical atom models in terms of electron distribution

19. ELECTRON ENERGY STATES The relative energies of the electrons for the various shells and subshells.

20. 4d N-shell n = 4 4p 3d 4s 3p M-shell n = 3 Energy 3s 2p L-shell n = 2 2s 1s K-shell n = 1 ELECTRON ENERGY STATES Electrons... • have discrete energy states • tend to occupy lowest available energy state.

21. QUANTUM NUMBERS The number of available electron states in some of the electron shells and subshells

22. Element Atomic # Electron configuration Hydrogen 1 1s 1 Helium 2 (stable) 1s 2 Lithium 3 1s 2 2s 1 Beryllium 4 1s 2 2s 2 Boron 5 1s 2 2s 2 2p 1 1s 2 2s 2 2p 2 Carbon 6 ... ... Neon 10 1s 2 2s 2 2p 6 (stable) Sodium 11 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 3s 2 Magnesium 12 1s 2 2s 2 2p 6 3s 2 3p 1 Aluminum 13 ... ... 1s 2 2s 2 2p 6 3s 2 3p 6 (stable) Argon 18 ... ... ... Krypton 36 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 (stable) SURVEY OF ELEMENTS • Most elements: Electron configuration not stable.

23. valence electrons ELECTRON CONFIGURATION Valance electron are extremely important, as will be seen, they participate in the bonding between atoms to form atomic and molecular aggregates. Many of the physical and chemical properties of solids are based on these valence electrons. • Valence electrons – those in unfilled shells • Filled shells - stable valence electron shell • Valence electrons are most available for bonding and tend to control the chemical properties • example: C (atomic number = 6) 1s22s2 2p2

24. 1s2 2s2 2p6 3s2 3p6 3d6 4s2 valence electrons 4d N-shell n = 4 4p 3d 4s 3p M-shell n = 3 Energy 3s 2p L-shell n = 2 2s 1s K-shell n = 1 ELECTRON CONFIGURATION 26 ex: Fe - atomic # =

25. ELECTRON CONFIGURATION The electrons fill up the lowest possible energy states in the electron shells and subshells The filled and lowest unfilled energy states for a sodium atom.

26. THE PERIODIC TABLE Periodic table – Element classification according to electron configuration The periodic table of the elements. The numbers in parentheses are the atomic weights of the most stable or common isotopes.

27. inert gases give up 1e give up 2e accept 2e accept 1e give up 3e H He Li Be O F Ne Na Mg S Cl Ar K Ca Sc Se Br Kr Rb Sr Y Te I Xe Cs Ba Po At Rn Fr Ra THE PERIODIC TABLE • 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.

28. ELECTRONEGATIVITY • Ranges from 0.7 to 4.0 • Large values: tendency to acquire electrons. Smaller electronegativity Larger electronegativity

29. IONIC BONDING Ionic bond – metal + nonmetal donates accepts electrons electrons Dissimilar electronegativities ex: MgO Mg 1s2 2s2 2p63s2 O 1s2 2s2 2p4 Mg2+1s2 2s2 2p6O2- 1s2 2s2 2p6

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

31. IONIC BONDING Ionic bonding in sodium chloride (NaCl)

32. IONIC BONDING Bonding Energies and Melting Temperatures for Various Substances

33. NaCl MgO CaF 2 CsCl Give up electrons Acquire electrons EXAMPLE: IONIC BONDING • Predominant bonding in Ceramics

34. shared electrons H from carbon atom CH 4 H H C shared electrons from hydrogen H atoms COVALENT BONDING • Similar electronegativity share electrons • Bonds determined by valence – s & p orbitals dominate bonding • Example: CH4 C: has 4 valence e-, needs 4 more H: has 1 valence e-, needs 1 more Electronegativities are comparable.

35. x ( 100 %) COVALENT BONDING • Ionic-Covalent Mixed Bonding % ionic character = where XA & XB are Pauling electronegativities Ex: MgO XMg = 1.3XO = 3.5

36. METALLIC BONDING • Metallic bonding – found in metals and alloys • Metallic materials have one, two, or at most three valence electrons. These atom not bound to any particular atom in the solid and are more or less free to drift throughout the entire metal • They may be thought of as belonging to the metal as a whole or forming a “sea of electrons” or an “electron cloud.” • The remaining nonvalence electrons and atomic nuclei form what are called ion cores, which possess a net positive charge equal in magnitude to the total valence electron charge per atom

37. METALLIC BONDING The free electrons shield the positively charged ion cores from mutually repulsive electrostatic forces, which they would otherwise exert upon one another; consequently the metallic bond is nondirectional in character. Schematic illustration of metallic bonding

38. Secondary Bonding or Van Der Waals Bonding ex: liquid H 2 asymmetric electron H H 2 2 clouds + - + - H H H H secondary secondary bonding bonding + - + - Cl Cl H H secondary bonding Arises from interaction between dipoles • Fluctuating dipoles • Permanent dipoles-molecule induced secondary -general case: bonding secondary -ex: liquid HCl bonding -ex: polymer secondary bonding

39. MOLECULES Molecules Hydrogen bonding in hydrogen fluoride (HF)

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

41. Summary: Primary Bonds secondary bonding SUMMARY: PRIMARY BONDS Ceramics Large bond energy large Tm large E small a (Ionic & covalent bonding): Metals Variable bond energy moderate Tm moderate E moderate a (Metallic bonding): Polymers Directional Properties Secondary bonding dominates small Tm small E large a (Covalent & Secondary):

42. Energy typical neighbor bond length typical neighbor r bond energy • Dense, ordered packing Energy typical neighbor bond length r typical neighbor bond energy ENERGY AND PACKING • Non dense, random packing Dense, ordered packed structures tend to have lower energies.

43. CRYSTAL STRUCTURES- MATERIALS AND PACKING Crystalline materials... • atoms pack in periodic, 3D arrays • typical of: -metals -many ceramics -some polymers crystalline SiO2 Noncrystalline materials... Si Oxygen • atoms have no periodic packing • occurs for: -complex structures -rapid cooling "Amorphous" = Noncrystalline noncrystalline SiO2

44. FUNDAMENTAL CONCEPTS • For the face centered cubic crystal structure • a hard sphere unit cell representation • a reduced-sphere unit cell • an aggregate of many atoms

45. METALLIC CRYSTAL STRUCTURES • How can we stack metal atoms to minimize empty space? 2-dimensions vs. Now stack these 2-D layers to make 3-D structures

46. METALLIC CRYSTAL STRUCTURES • Tend to be densely packed. • Reasons for dense packing: - Typically, only one element is present, so all atomic radii are the same. - Metallic bonding is not directional. - Nearest neighbor distances tend to be small in order to lower bond energy. - Electron cloud shields cores from each other • Have the simplest crystal structures.

47. SIMPLE CUBIC STRUCTURE • Rare due to low packing denisty (only Po has this structure) • Close-packed directions are cube edges. • Coordination # = 6 (# nearest neighbors)

48. volume atoms atom 4 a 3 unit cell p (0.5a) 3 R=0.5a volume close-packed directions unit cell contains 8 x 1/8 = 1 atom/unit cell ATOMIC PACKING FACTOR (APF) Volume of atoms in unit cell* APF = Volume of unit cell *assume hard spheres • APF for a simple cubic structure = 0.52 1 APF = 3 a

49. BODY CENTERED CUBIC STRUCTURE (BCC) • Atoms touch each other along cube diagonals. --Note: All atoms are identical; the center atom is shaded differently only for ease of viewing. ex: Cr, W, Fe (), Tantalum, Molybdenum • Coordination # = 8 2 atoms/unit cell: 1 center + 8 corners x 1/8

50. a 3 a 2 Close-packed directions: R 3 a length = 4R = a atoms volume 4 3 p ( 3 a/4 ) 2 unit cell atom 3 APF = volume 3 a unit cell ATOMIC PACKING FACTOR: BCC • APF for a body-centered cubic structure = 0.68 a