1. 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

2. 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.

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

4. 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

5. (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

6. (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

7. (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

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

9. . 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

10. 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?

11. ATOMIC STRUCTURE
Atoms are the structural unit of all engineering materials

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

13. 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

14. 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= x 1023 atoms ( Avogadro’s number NA ).

15. 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 x ( 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

16. 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

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

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

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

20. 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 J.s
c = Speed of light
λ = Wavelength of light
v=frequency of photon
Absorb
Energy
(Photon)
Emit
Energy
(Photon)
Energy levels
2-6

21. 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)

22. Ans 3:
ΔE =

23. 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

24. 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

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

26. 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

27. 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

28. 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

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

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

31. 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

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

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

34. 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

35. 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

36. 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

37. 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

38. 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

39. 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

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

41. 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

42. 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

43. 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 = KJ/mol
Melting point = 1923oC
2-21

44. 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

45. 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

46. 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

47. 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

48. 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

49. 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

50. 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)

51. 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

52. 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

53. 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