Introduction to Material Science & Engineering Atom and Bonding Acknowledgement: This slides are largely obtained from Dr.Neoh Siew Chin UniMAP on the subject Material Engineering

Describe and differentiate material science and engineering Material science is the exploration and research for basic knowledge about the internal structure, properties and processing of materials. (Ex: research on things such as, density, elasticity, conductivity, or chemical compound) Material engineering is the application of fundamental and applied knowledge of materials so that the materials can be converted into products needed or desired by society For example: if a company is looking for a high density metal object, but wants it to have low conductivity, then they would look to a material engineer for a design of a man made material that they could use.

Classify types of materials: Metallic Polymeric Ceramic Composite Electronic materials Fundamental classes Processing/ applicational classes

Classify types of materials: Metallic - inorganic substances - composed of one or more metallic elements (may also contain some non-metallic elements) - crystalline structure where atoms are arranged in an orderly manner - generally a good thermal and electrical conductors - Ex: iron, copper, aluminum, nickel, titanium non-metallic elements: carbon, nitrogen, oxygen

(2) Polymeric - mostly consist of long molecular chains or networks that are usually based on organics (carbon-containing precursors) - non-crystalline OR mixture of non-crystalline + crystalline - strength and ductility vary - poor conductors of electricity - Ex of application: DVDs

(3) Ceramic - inorganic materials ( metallic + non-metallic elements chemically bonded together ) - can be crystalline, noncrystalline or mixtures - high hardness and high-temperature strength but tend to be brittle - Advantages : light weight, high strength & hardness, good heat and wear resistance - Ex applications: furnace linings for heat treatment

(4) Composite - two or more materials (phases or constituents) integrated to form new one – the constituents keep their properties and the overall composite will have properties different than each of them) - Ex application: aerospace, avionics, automobile, civil structural and sports equipment industries - Advantages : high strength and stiffness-to-weight ratio - Disadvantage : brittleness and low fracture toughness

(5) Electronic - important for advanced engineering technology - most important electronic materials – pure silicon modified in various ways to change its electrical properties

. The application and advancement in material science and technology. Ex: smart materials; nanomaterials Materials that are able to sense external environmental stimuli (temp., stress, light, humidity) and respond to them by changing their properties (mechanical, electrical or appearance). Use in sensors, actuators

ASSIGNMENT 1 - 10% Due Date: 16/12/10 Group of 4-5 5-10 pages Synthesis and Evaluation problems (a) Name the important criteria for selecting materials to use in a protective sports helmet (b) Identify materials that would satisfy these criteria (c) Why would a solid metal helmet not be a good choice?

ATOMIC STRUCTURE Atoms are the structural unit of all engineering materials

Fundamental Concept Each atoms consist of nucleus composed of protons and neutron and surrounded by electrons n= Quantum Number

Mass and Charge of the Proton, Neutron, and Electron Mass (g) Charge (C) Proton 1.673X10-24 +1.602X10-19 Neutron 1.675X10-24 Electron 9.109X10-28 -1.602X10-19 Source: Smith, W.F. and Hashemi, J., Foundations of Materials Science and Engineering, McGraw-Hill, 2006

Atomic number, Z - Number of protons (p) Atomic number, Z - Number of protons (p). In a neutral atom the atomic number is equal to the number of electrons (Z=e). Atomic mass, A - Total mass of proton and neutron in the nucleus ( A=Z+N ). Isotope - atoms that have two or more atomic mass. Same number of proton but different number of neutron. 1 atomic mass unit (a.m.u) = 1/12 of the atomic mass of carbon 1 mole= 6.023 x 1023 atoms ( Avogadro’s number NA ).

Atomic Number and Atomic Mass Atomic Number = Number of Protons in the nucleus Unique to an element Example :- Hydrogen = 1, Uranium = 92 Relative atomic mass = Mass in grams of 6.203 x 1023 ( Avagadro Number) Atoms. Example :- Carbon has 6 Protons and 6 Neutrons. Atomic Mass = 12. One Atomic Mass unit is 1/12th of mass of carbon atom. One gram mole = Gram atomic mass of an element. Example :- One gram Mole of Carbon 12 Grams Of Carbon 6.023 x 1023 Carbon Atoms 2-3

Z X A Ex: Determine the number of electron (e), neutron (N) in fluorin atom Ans: A=p+N= Z=p=e= 9 So, N=A-Z=10 9 F 19

Periodic Table Source: Davis, M. and Davis, R., Fundamentals of Chemical Reaction Engineering, McGraw-Hill, 2003. 2-4

Example: 1 mole aluminium have mass of 26.98 g and 6.023 x 1023 atoms. 1) What is the mass in grams of 1 atom of aluminium (A=26.98g/mol) Ans:

2) How many atom of Copper (Cu) in 1 gram Of Copper ?(A=63.54g/mol) Ans:

Electron Structure of Atoms Electron rotates at definite energy levels. Energy is absorbed to move to higher energy level. Energy is emitted during transition to lower level. Energy change due to transition = ΔE = = h=Planks Constant = 6.63 x 10-34 J.s c = Speed of light λ = Wavelength of light v=frequency of photon Absorb Energy (Photon) Emit Energy (Photon) Energy levels 2-6

Example 3) Calculate the energy in joules (J) and electron volts (eV) of the photon whose wave length  is 121.6nm. (Given 1.00eV=1.60X10-19J; h= 6.63X10-34J.s)

Ans 3: ΔE =

Energy in Hydrogen Atom Hydrogen atom has one proton and one electron Energy of hydrogen atoms for different energy levels is given by (n=1,2…..)principal quantum numbers Example:- If an electron undergoes transition from n=3 state to n=2 state, the energy of photon emitted is Energy required to completely remove an electron from hydrogen atom is known as ionization energy 2-7

A hydrogen atom exists with its electron in the n= 3 state A hydrogen atom exists with its electron in the n= 3 state. The electron undergoes a transition to the n=2 state. Calculate (a) the energy of the photon emitted, (b) its frequency, and (c) its wavelength. Ans: (a) Energy of the photon emited is

b) The frequency of the photon is c) The wavelength of the photon is

Quantum Numbers of Electrons of Atoms Subsidiary Quantum Number l Represents sub energy levels (orbital). Range 0…n-1. Represented by letters s,p,d and f. Principal Quantum Number (n) Represents main energy levels. Range 1 to 7. Larger the ‘n’ higher the energy. n=1 s orbital (l=0) n=2 n=2 n=1 p Orbital (l=1) n=3 2-8

Quantum Numbers of Electrons of Atoms (Cont..) Electron spin quantum number ms. Specifies two directions of electron spin. Directions are clockwise or anticlockwise. Values are +1/2 or –1/2. Two electrons on same orbital have opposite spins. No effect on energy. Magnetic Quantum Number ml. Represents spatial orientation of single atomic orbital. Permissible values are –l to +l. Example:- if l=1, ml = -1,0,+1. I.e. 2l+1 allowed values. No effect on energy. 2-9

Electron Structure of Multielectron Atom Maximum number of electrons in each atomic shell is given by 2n2. Atomic size (radius) increases with addition of shells. Electron Configuration lists the arrangement of electrons in orbitals. Example :- 1s2 2s2 2p6 3s2 For Iron, (Z=26), Electronic configuration is 1s2 2s2 2p6 3s2 3p6 3d6 4s2 Number of Electrons Orbital letters Principal Quantum Numbers 2-10

Stable electron configuration Stable electron configurations... have complete s and p subshells tend to be unreactive.

• Most elements: Electron configuration not stable. • Why? Valence (outer) shell usually not filled completely.

Example Write the electron configuration for the following atoms by using conventional spdf notation. a) Fe atom (Z=26) and the Fe2+ and Fe3+ ions

remember !!!!! 7s 7p 7d 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d

Electrons fill up in this order The order for writing the orbitals

Electron Structure and Chemical Activity (Cont..) Electronegative elements accept electrons during chemical reaction. Some elements behave as both electronegative and electropositive. Electronegativity is the degree to which the atom attracts electrons to itself Measured on a scale of 0 to 4.1 Example :- Electronegativity of Fluorine is 4.1 Electronegativity of Sodium is 1. Na Te N O Fl Electro- negative Electro- positive K 1 W 2 H Se 3 4 2-12

Atomic and Molecular Bonds Ionic bonds :- Strong atomic bonds due to transfer of electrons Covalent bonds :- Large interactive force due to sharing of electrons Metallic bonds :- Non-directional bonds formed by sharing of electrons Permanent Dipole bonds :- Weak intermolecular bonds due to attraction between the ends of permanent dipoles. Fluctuating Dipole bonds :- Very weak electric dipole bonds due to asymmetric distribution of electron densities. 2-12

Ionic Bonding Ionic bonding is due to electrostatic (Coulombic) force of attraction between cations and anions. It can form between metallic and nonmetallic elements. Electrons are transferred from electropositive to electronegative atoms Electropositive Element Electronegative Atom Electron Transfer Electrostatic Attraction Cation +ve charge Anion -ve charge IONIC BOND 2-14

Ionic Bonding Lattice energies and melting points of ionically bonded solids are high. Lattice energy decreases when size of ion increases. Multiple bonding electrons increase lattice energy. Sodium Atom Na Chlorine Cl Sodium Ion Na+ Chlorine Ion Cl - I O N C B D

Ionic Force for Ion Pair Nucleus of one ion attracts electron of another ion. The electron clouds of ion repulse each other when they are sufficiently close. Force versus separation Distance for a pair of oppositely charged ions Figure 2.11 2-16

Ion Force for Ion Pair (Cont..) Z1,Z2 = Number of electrons removed or added during ion formation e = Electron Charge a = Interionic seperation distance ε = Permeability of free space (8.85 x 10-12c2/Nm2) (n and b are constants) 2-17

Interionic Force - Example Force of attraction between Na+ and Cl- ions Z1 = +1 for Na+, Z2 = -1 for Cl- e = 1.60 x 10-19 C , ε0 = 8.85 x 10-12 C2/Nm2 a0 = Sum of Radii of Na+ and Cl- ions = 0.095 nm + 0.181 nm = 2.76 x 10-10 m Cl- Na+ a0 2-18

Interionic Energies for Ion Pairs Net potential energy for a pair of oppositely charged ions = Enet is minimum when ions are at equilibrium seperation distance a0 Attraction Energy Repulsion Energy Energy Released Energy Absorbed 2-19

Ion Arrangements in Ionic Solids Ionic bonds are Non Directional Geometric arrangements are present in solids to maintain electric neutrality. Example:- in NaCl, six Cl- ions pack around central Na+ Ions As the ratio of cation to anion radius decreases, fewer anion surround central cation. Ionic packing In NaCl and CsCl Figure 2.13 CsCl NaCl 2-20

Bonding Energies Lattice energies and melting points of ionically bonded solids are high. Lattice energy decreases when size of ion increases. Multiple bonding electrons increase lattice energy. Example :- NaCl Lattice energy = 766 KJ/mol Melting point = 801oC CsCl Lattice energy = 649 KJ/mol Melting Point = 646oC BaO Lattice energy = 3127 KJ/mol Melting point = 1923oC 2-21

Overlapping Electron Clouds Covalent Bonding In Covalent bonding, outer s and p electrons are shared between two atoms to obtain noble gas configuration. Takes place between elements with small differences in electronegativity and close by in periodic table. In Hydrogen, a bond is formed between 2 atoms by sharing their 1s1 electrons H + H H H 1s1 Electrons Electron Pair Hydrogen Molecule Overlapping Electron Clouds

Covalent Bonding - Examples In case of F2, O2 and N2, covalent bonding is formed by sharing p electrons Fluorine gas (Outer orbital – 2s2 2p5) share one p electron to attain noble gas configuration. Oxygen (Outer orbital - 2s2 2p4) atoms share two p electrons Nitrogen (Outer orbital - 2s2 2p3) atoms share three p electrons H H F F F + F F F H Bond Energy=160KJ/mol O + O O O O = O Bond Energy=28KJ/mol N N N N N + N Bond Energy=54KJ/mol 2-23

Covalent Bonding in Benzene Chemical composition of Benzene is C6H6. The Carbon atoms are arranged in hexagonal ring. Single and double bonds alternate between the atoms C H Structure of Benzene Simplified Notations

Metallic Bonding Positive Ion Atoms in metals are closely packed in crystal structure. Loosely bounded valence electrons are attracted towards nucleus of other atoms. Electrons spread out among atoms forming electron clouds. These free electrons are reason for electric conductivity and thermal conductivity Since outer electrons are shared by many atoms, metallic bonds are Non-directional Positive Ion Valence electron charge cloud

Metallic Bonds (Cont..) Overall energy of individual atoms are lowered by metallic bonds Minimum energy between atoms exist at equilibrium distance a0 Fewer the number of valence electrons involved, more metallic the bond is. Example:- Na Bonding energy 108KJ/mol, Melting temperature 97.7oC Higher the number of valence electrons involved, higher is the bonding energy. Example:- Ca Bonding energy 177KJ/mol, Melting temperature 851oC 2-29

Secondary Atomic Bonding Secondary bonds are due to attractions of electric dipoles in atoms or molecules. Dipoles are created when positive and negative charge centers exist. Bonding result from the columbic attraction between positive end of one dipole and the negative region of an adjacent one Sometimes called Van der Waals bond There two types of bonds permanent and fluctuating. Dipole moment=μ =q.d q= Electric charge d = separation distance +q d

Fluctuating Dipole Bond Weak secondary bonds in noble gasses. Dipoles are created due to asymmetrical distribution of electron charges. Electron cloud charge changes with time. Symmetrical distribution of electron charge Asymmetrical Distribution (Changes with time)

Permanent Dipole Bond CH4 CH3Cl Secondary bond created by the attraction of molecules that have permanent dipole. That is, each molecule has positive and negative charge center separated by distance. Symmetrical Arrangement Of 4 C-H bonds CH4 No Dipole moment CH3Cl Asymmetrical Tetrahedral arrangement Creates Dipole

Hydrogen Bond Hydrogen bonds are Dipole-Dipole interaction between polar bonds containing hydrogen atom. Special type of intermolecular permanent dipole attraction that occur between hydrogen atom bonded to a highly electronegative element (F, O, N or Cl) and another atom of a highly electronegative element. 105 0 O H Hydrogen Bond

Mixed Bonding Chemical bonding of atoms or ions can involve more than one type of primary bond and can also involve secondary dipole bond. Ionic - covalent Metallic – covalent Metallic – ionic Ionic – covalent - metallic