MSE - Lecture 4,5.pptx

Materials Science and Engineering Dr. Ranjit Kumar Department of Chemical Engineering Email: [email protected]  Atomic structure Nucleus Neutron Quarks Gluons Electron Protons BOHR ATO...

Materials Science and Engineering Dr. Ranjit Kumar Department of Chemical Engineering Email: [email protected]  Atomic structure Nucleus Neutron Quarks Gluons Electron Protons BOHR ATOM orbital electrons: n = principal quantum number n=3 2 1 Nucleus: Z = # protons = 1 for hydrogen to 94 for plutonium 12 Atomic mass A ≈ Z + N N = # neutrons  Electronic structure Valence electrons determine all of the following properties: Chemical Electrical Thermal Optical 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) 1, 3, 5, 7 (-l to +l) 3 ml = magnetic ms = spin ½, -½  Electronic structure Principal Shell No. Number of electrons quantum designatio Subshells of Per subshell Per shell no. n state s 1 K s 1 2 2 2 L s 1 2 8 p 3 6 3 M s 1 2 18 p 3 6 f 5 10 d f d p 4 N s 1 2 32 d s p 3 6 p d s 5 10 d 7 14 p s f p s s 4 1 2 3 4 5  Atomic Structure Electron energy state Occupy the atomic no. at the lowest energy state to higher energy state Low K shell L shell M shell N shell No. of electron can occupy each orbital: High s = 2; p = 6; d = 10; f = 14  Electron energy states Electrons... have discrete energy states tend to occupy lowest available energy state. 4d 4p N-shell n = 4 3d 4s 3p Energy M-shell n = 3 3s 2p 2s L-shell n = 2 1s K-shell n = 1 6  Electronic configuration 1s2 2s2 2p6 3s2 3p6 3d 6 4s2 ex: Fe - atomic # = 26 valence electrons 4d 4p N-shell n = 4 3d 4s 3p Energy M-shell n = 3 3s 2p L-shell n = 2 2s 1s K-shell n = 1 7  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 2 2s 1 Beryllium 4 1s 2 2s 2 2 2p 1 Boron 5 1s 2 2s Carbon 6 1s 2 2s 2 2p 2...... Neon 10 1s 2 2s 2 2p 6 (stable) Sodium 11 1s 2 2s 2 2p 6 3s 1 Magnesium 12 1s 2 2s 2 2p 6 3s 2 Aluminum 13 1s 2 2s 2 2p 6 3s 2 3p 1...... Argon 18 1s 2 2s 2 2p 6 3s 2 3p 6 (stable)......... Krypton 36 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 (stable) Why? Valence (outer) shell usually not filled completely.  The periodic table Columns: Similar Valence Structure inert gases give up 1e give up 2e accept 2e accept 1e p 3e H He Li O F Ne Be give u Na Mg S Cl Ar Se Br Kr K Ca Sc Te I Xe Rb Sr Y Po At Rn Cs Ba Fr Ra Electropositive elements: Electronegative elements: Readily give up electrons Readily acquire electrons to become + ions. to become - ions.  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 10 Bond, 3rd edition, Copyright 1939 and 1940, 3rd edition. Copyright 1960 by Cornell University.  Bonding forces and energies Attractive force, FA Repulsive force : Dominates at r Repulsive force, FR small distance + Attractive force, FA Attractive force : Dominates at Attraction E Fdr r r r EN FN dr FA dr  larger distance Force F FR dr 0    Interatomic separation r At equilibrium : Both are equal Repulsion  E A E R Repulsive force, FR (a) Low energy : The element is Net force, FN - Repulsive energy ER more stable Interatomic separation r Net energy EN (a) The dependence of repulsive, attractive, and 11 Attractive energy EA net forces on interatomic separation for two isolated atoms. (b) (b) The dependence of repulsive, attractive, and  Properties from bonding Bond length, r Melting Temperature, Tm Energy r Bond energy, Eo r o r Energy smaller Tm unstretched length r larger Tm o r Eo = Tm is larger if Eo is larger. “bond energy” 12  Properties from bonding: thermal expansion coefficient Coefficient of thermal expansion,  coeff. thermal expansion length, Lo unheated, T1 L L =  (T -T ) 2 1 heated, T2 L o ~ symmetry at ro Energy unstretched length r o r is larger if Eo is smaller. larger  Eo Eo 13 smaller  Properties from bonding: modulus E F = kx r 14  Types of bonding: ionic 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 stable Coulombic 15 Attraction  Examples of ionic bonding Predominant bonding in Ceramics NaCl MgO CaF 2 CsC l 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 16 Bond, 3rd edition, Copyright 1939 and 1940, 3rd edition. Copyright 1960 by Cornell University.  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 shared electrons from H CH4 carbon atom H: has 1 valence e-, needs 1 more H C H Electronegativities shared electrons are comparable. H from hydrogen atoms 17  Metallic bonding Ions in a sea of electrons Attraction between free electrons and metal ions 18  Ionic-covalent mixed bonding % Ionic character = {1 − exp[−(0.25)(XA − XB)2]} × 100 where XA & XB are Pauling electronegativities XMg = 1.3 Example: MgO XO = 3.5 % ionic character {1 – exp[- (0.25) (3.5 – 1.3)2]} x (100%) 70.2% ionic 19  Secondary Bonding Secondary bonds, or van der Waals (physical) bonds, are weak in comparison to the primary or chemical bonds; bonding energies range between about 4 and 30 kJ/mol. Secondary bonding exists between virtually all atoms or molecules, but its presence may be obscured if any of the three primary bonding types is present. Secondary bonding is evidenced for the inert gases, which have stable electron structures. Secondary (or intermolecular) bonds are possible between atoms or groups of atoms, which themselves are joined together by primary (or intramolecular) ionic or covalent bonds. Van der Waals forces are of two types  Secondary bonding Arises from interaction between dipoles Fluctuating dipoles example: liquid H2 asymmetric electron clouds H2 H2 + - + - H H H H secondary secondary bonding bonding Permanent dipoles-molecule induced secondary -general case: + - + - bonding secondary -example: liquid HCl H Cl bonding H Cl se con -example: polymer dary bo n d in secondary bonding 21 g Bonding Energy and Melting Points of various substances  Summary Type Bond Energy Comments Ionic Large! Non-directional (ceramics) Variable Directional Covalent Diamond (large) (semiconductors, ceramics, Bismuth (small) polymer chains) Variable Metallic Non-directional (metals) Tungsten (large) Mercury (small) Directional Secondary Smallest Interchain (polymer) Intermolecular Ceramics Large bond energy (Ionic & covalent bonding) Large Tm and E, small  Metals Variable bond energy (Metallic bonding) Moderate Tm, E, and  Polymers Directional (Covalent & secondary) properties, Secondary bonding dominates Small Tm and E, large  3 0