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Module 2.Module 2.

Structures of Engineering Materials1.Atomic Structure

Prof. Dr. Ir. Bondan T. Sofyan, M.Si.

2.Atomic Bonding

2.1. Atomic Structure


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Atomic StructureAll materials consist of elements and all elements consist of atoms. Each atoms has its own characteristics → has different properties → see Periodic Table

proton

Atomic Structure


proton

neutron e-

orbital electrons: n = principal quantum number

n=3 2 1

BOHR ATOM

Nucleus: Z = # protons

n=3 Adapted from Fig. 2.1, Callister 6e.


= 1 for hydrogen to 94 for plutoniumN = # neutrons

Atomic mass A ≈ Z + N

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• have discrete energy states• tend to occupy lowest available energy state.

Electrons...ELECTRON ENERGY STATES

py gye

asi

ng

en

erg

y

2

n=3

n=4

3s2p

3p

4s4p

3d


Inc

re

n=1

n=2

1s2s

2p

Adapted from Fig. 2.5, Callister 6e.

• have complete s and p subshells• tend to be unreactive.

Stable electron configurations...STABLE ELECTRON CONFIGURATIONS

Z Element Configuration

2 He 1s2

10 Ne 1s22s22p6

18 Ar 1s22s22p63s23p6

2 2 6 2 6 10 2 6

Adapted from Table 2.2, Callister 6e.


36 Kr 1s22s22p63s23p63d104s24p6

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• Most elements: Electron configuration not stable.Element Hydrogen Helium Lithi

Atomic # 1 2 3

Electron configuration 1s1 1s2 (stable)

SURVEY OF ELEMENTS

Lithium Beryllium Boron Carbon ... Neon Sodium Magnesium Aluminum

3 4 5 6

10 11 12 13

1s22s1 1s22s2 1s22s22p1 1s22s22p2 ... 1s22s22p6 (stable) 1s22s22p63s1 1s22s22p63s2 1s22s22p63s23p1

Adapted from Table 2.2, Callister 6e.


• Why? Valence (outer) shell usually not filled completely.

Aluminum ... Argon ... Krypton

13

18 ... 36

... 1s22s22p63s23p6 (stable) ... 1s22s22p63s23p63d104s246 (stable)

ne

rt g

ase

s e

pt

1e

ep

t 2

e

ive

up

1e

e

up

2e

3

e

Metal

• Columns: Similar Valence Structure

THE PERIODIC TABLE

He

Ne

Ar

Kr

Xe

ina

cc

ac

cgg

ive

giv

e u

p 3

F Li Be

Nonmetal

Intermediate

H

Na Cl

Br

I

O

S Mg

Ca

Sr

K

Rb

Sc

Y

Se

Te


Rn At Ba

Ra

Cs

Fr

Po

Electropositive elements:Readily give up electronsto become + ions.

Electronegative elements:Readily acquire electronsto become - ions.


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• Ranges from 0.7 to 4.0,

He H

• Large values: tendency to acquire electrons.

ELECTRONEGATIVITY

-

Ne -

Ar -

Kr -

Xe -

Rn -

F 4.0

Cl 3.0

Br 2.8

I 2.5

At 2.2

Li 1.0

Na 0.9

K 0.8

Rb 0.8

Cs 0.7

2.1

Be 1.5

Mg 1.2

Ca 1.0

Sr 1.0

Ba 0.9

Ti 1.5

Cr 1.6

Fe 1.8

Ni 1.8

Zn 1.8

As 2.0

Prof. Dr. Ir. Bondan T. Sofyan, M.Si. 7

Smaller electronegativity Larger electronegativity

Fr 0.7

Ra 0.9

Adapted from Fig. 2.7, Callister 6e. (Fig. 2.7 is adapted from Linus Pauling, The Nature of the Chemical Bond, 3rd edition, Copyright 1939 and 1940, 3rd edition. Copyright 1960 by CornellUniversity.

2.1. Atomic Bonding


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P i b d

Atoms may form two types of bond:

Primary bonds: a. Ionic bond b. Covalent bond c. Metallic bond

Secondary bonds: a. Van der Waal bond


a. Van der Waal bond b. Hydrogen bond

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Primary BondsElectrons are directly transferred from one atom to another and localised. Atoms are bonded by coulombic force. Bonds are non-

Ionic Bonding

directional.


• Occurs between + and - ions.• Requires electron transfer.• Large difference in electronegativity required.

IONIC BONDING

Na (metal) unstable

Cl (nonmetal) unstable

electron

• Example: NaCl


+ - Coulombic Attraction

Na (cation) stable

Cl (anion) stable

8

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However, more energy is released as Na+ and Cl- ions are getting closer.

The attractive energy , EA, is given by:

))((1 eZeZE −=

where ε0 is the permittivity of free space (8.85 x 10-12 F/m), Z1 and Z2 are the valences of the two ions, e is the electronic charge (1.602 x 10-19 C) and r is the separation of the ions.In convention, negative is for the energy released in forming

))((4 21

0

eZeZr

EA πε=


g gy gbonds and positive for the energy required to break bonds.

At a separation, r*, the energy released equals the 1.5eV required to transfer the electrons. So at separations smaller than this there will be a net release of energy and so the bond is stable.

EEA


The variation of the attractive coulomb interaction,EA, with ionic separation, r.

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The approaching ions also suffer a repulsive force due to the interaction of the inner electron shells of the two ions.A repulsive energy, ER, is given by:

nR rBE =r

where n is an exponent ~8 and B is an empirical constant.

ER


The variation of the repulsive interaction, Ur, withionic separation, r.

ER

EA


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The overall energy of interaction, Etot, is the combination of the attractive energy, EA and the repulsive energy, ER.

ntot rBeZeZ

rE +−= ))((

41

210πε

ETOT

EW


The equilibrium spacing between the two ions, given as r0, is where the energy of interaction is a minimum.r0 = bond length

Structure of Salt (NaCl)

Schematic model Lattice


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• Predominant bonding in Ceramics

MgONaCl

MORE EXAMPLES: IONIC BONDING

He -

Ne -

Ar -

Kr -

Xe -

Rn

F 4.0

Cl 3.0

Br 2.8

I 2.5

At

Li 1.0

Na 0.9

K 0.8

Rb 0.8

Cs

H 2.1

Be 1.5

Mg 1.2

Ca 1.0

Sr 1.0

Ba

Ti 1.5

Cr 1.6

Fe 1.8

Ni 1.8

Zn 1.8

As 2.0

CsCl

g

CaF2O

3.5


Give up electrons Acquire electrons

-At

2.2Cs 0.7

Fr 0.7

a 0.9

Ra 0.9

Adapted from Fig. 2.7, Callister 6e. (Fig. 2.7 is adapted from Linus Pauling, The Nature of the Chemical Bond, 3rd edition, Copyright 1939 and 1940, 3rd edition. Copyright 1960 by CornellUniversity.

• Requires shared electrons• Example: CH4

C: has 4 valence eshared electrons from carbon atomH

CH

COVALENT BONDING

C: has 4 valence e,needs 4 more

H: has 1 valence e,needs 1 more

Electronegativitiesare comparable.

from carbon atom

shared electrons from hydrogen atoms

HH

H

C

CH4



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He -

H 2.1 SiC

C(diamond)

H2OH2

Cl2

F2

co

lum

n IV

A

EXAMPLES: COVALENT BONDING

Ne -

Ar -

Kr -

Xe -

Rn -

F 4.0

Cl 3.0

Br 2.8

I 2.5

At 2.2

Li 1.0

Na 0.9

K 0.8

Rb 0.8

Cs 0.7

Fr

2.1

Be 1.5

Mg 1.2

Ca 1.0

Sr 1.0

Ba 0.9

Ra

Ti 1.5

Cr 1.6

Fe 1.8

Ni 1.8

Zn 1.8

As 2.0

SiCC

2.5

Cl2

Si 1.8

Ga 1.6

G A

Ge 1.8

O 2.0

c

Sn 1.8Pb 1.8

Ad t d f Fi 2 7 C lli t 6 (Fi 2 7 i


• Molecules with nonmetals• Molecules with metals and nonmetals• Elemental solids (RHS of Periodic Table)• Compound solids (about column IVA)

Fr 0.7

Ra 0.9 GaAsAdapted from Fig. 2.7, Callister 6e. (Fig. 2.7 is

adapted from Linus Pauling, The Nature of the Chemical Bond, 3rd edition, Copyright 1939 and 1940, 3rd edition. Copyright 1960 by Cornell University.

Sharing electrons between adjacent atoms. Bonds are directional.

The energy released, Ua, associated with electron sharing:

a

Covalent Bonding

If the two atoms become too close together, the repulsive energy Ur is given by: b

nr raU −=

where r is the distance between the two atoms, n is ~6, and a is a constant.

g y

mr rbU =

where m is an exponent (~12) and b is a constant.

The overall covalent interaction energy Utot is given by:


mntot rb

raU +−=

energy Utot is given by:

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Covalent bonding in amolecule of methane (CH4)(CH4)


In a metal, valence electrons leave their parent atoms and combine to form an electron "gas" which freely wander around metal ions.

Metallic Bonding


The electrons are completely delocalized and provide a bonding force between the metal ions. The bonding between the atoms is non-directional.

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+ + + 1. The complete delocalization of the electrons means metal "ions" are readily i t h bl diff t

Metallic Bonding

+ + +

+ + +

interchangeable → different alloys.

2. The electrons are easy to move → metals are good electrical conductors.

3. The metal ions pack


• Primary bond for metals and their alloys


ptogether very well → metals have high densities and assume simple crystallographic structures.

Arises from interaction between dipoles• Fluctuating dipoles van der waals

H2 H2ex: liquid H2asymmetric electron

clouds

SECONDARY BONDING

• Permanent dipoles-molecule induced hydrogen bond

+ - secondary bonding + -

HH HH

H2 H2

secondary bonding

clouds

+ - + -secondary

bonding

-general case:


Adapted from Fig. 2.14,Callister 6e.


bonding

H Cl H Clsecondary bonding

secondary bonding

g

-ex: liquid HCl

-ex: polymer

Adapted from Fig. 2.14,Callister 6e.

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Secondary Bonds

A dipole may be formed when the electrical symmetry for some atoms or molecules is instantaneously distorted. Two opposite dipoles may attract each and form a eak bond

Van der Waal Bonding

attract each and form a weak bond.

Bonds in the condensed halogen molecules, such as liquid and solid forms of Cl2, Br2 and I2.


The energy of the Van der Waals bond

mntot rb

raU +−=

rr

where r is the distance between the two atoms, n is an exponent ~6, m ~12 and a and b are constants.

Van der Waals bond:


• the weakest bond• forces fluctuate with time • non-directional

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When a single hydrogen electron is shared with dissimilar atoms in covalent bonding, a net positive charge is displaced towards the hydrogen atom. The positively

Hydrogen Bonds

charged hydrogen can form a hydrogen bond with the negative end of a neighboring molecule.


• Approximately 30 times weaker than a normal covalent bond, because only one of the contributing atoms is supplying

Hydrogen Bonds

contributing atoms is supplying electrons to it.

• Relatively easily broken.

• Directional.

• Such bonds may also exist in many polymers linking discrete


a y po y e s g d sc etechain molecules together.

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Interatomic ForcesThe bonding force, F, varies as a function of the separation, r, between atoms.

drdUF =

F(r)


The variation of force between two atoms, F, as a function of atomic separation

Some engineering materials have mixed bonds.• In ceramic: ionic/covalent mixed bonds. e.g. SiO2.• In polymer: covalent/secondary mixed bond.

Mixed Bonds


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• Two atoms of similar electronegativity form either a metallic bond or covalent bond.

• Two atoms of different electronegativity form partial ionic bond. The ionic character increases with the difference in electronegativity.

For a AB compound:

e.g. NaCl:|ENa-ECl|= |0.9-3.0|=2.1

hi hl i i


→ highly ionic

e.g SiC: |ESi-EC|= |1.8-2.5|=0.7→ highly covalent

• Bond length, r

F F

• Melting Temperature, Tm

Energy (r)

PROPERTIES FROM BONDING: TM

• Bond energy, Eo

r

Energy (r)

unstretched length

r

larger T

smaller Tm

ro


Eo=

“bond energy”

ro r

unstretched length larger Tm

Tm is larger if Eo is larger.

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Bond Strength and Melting PointBonding Type Substance Bonding Energy

(kJ/mol)Melting

Temperature (oC)Ionic NaCl 640 801

MgO 1000 2800C l t Si 450 1410Covalent Si 450 1410

C (diamond) 713 >3550Metallic Hg 68 -39

Al 324 660Fe 406 1538W 849 3410

Van der Waals Ar 7 7 189


Van der Waals Ar 7.7 -189Cl2 31 -101

Hydrogen NH3 35 -78H2O 51 0

• Elastic modulus, E cross sectional area Ao

ΔL

length, Lo

undeformed ΔL F E

Elastic modulus

PROPERTIES FROM BONDING: E

• E ~ curvature at ro

ΔL

F deformed

ΔL F Ao

= E Lo

Energy

unstretched length


r

larger Elastic Modulus

smaller Elastic Modulus

ro E is larger if Eo is larger.

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• Coefficient of thermal expansion, α

ΔL

length, Lo

unheated, T1 ΔL

coeff. thermal expansion

PROPERTIES FROM BONDING: α

• α ~ symmetry at ro

ΔL

heated, T2 = α (T2-T1) ΔL

Lo

Energy


α is larger if Eo is smaller.r

smaller α

larger α

ro

TypeIonic

C l t

Bond Energy

Large!

Variablel Di d

Comments

Nondirectional (ceramics)

Directionali d t i

SUMMARY: BONDING

Covalent

Metallic

Secondary

large-Diamondsmall-Bismuth

Variablelarge-Tungstensmall-Mercury

smallest

semiconductors, ceramicspolymer chains)

Nondirectional (metals)

Directionalinter-chain (polymer)


Secondary smallest inter hain (polymer)

inter-molecular

• There is an equilibrium spacing between two atoms determined by an attractive force associated with electrostatic attraction and a repulsive force associated with the interaction of inner shell electrons.

• The nature of bonding has a strong effect on the properties of materials.

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Ceramics(Ionic & covalent bonding):

Large bond energylarge Tm

large Esmall α

SUMMARY: BONDING

Metals(Metallic bonding):

Polymers

Variable bond energymoderate Tm

moderate Emoderate α

Directional Properties


(Covalent & Secondary):

secondary bonding

Secondary bonding dominatessmall Tsmall Elarge α

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