1 CHEMICAL BONDS

Mar Athanasius College of

Engineering - MECH - Dr. Binu

C.Yeldose

1

METALLURGY AND MATERIAL SCIENCE

CHEMICAL BONDS

Dr. Binu C. Yeldose,Lecturer,

Department of Mechanical Engineering,Mar Athanasius College of Engineering.

KOTHAMANGALAM



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

• Most elements: Electron configuration not stable.

Element Hydrogen Helium Lithium Beryllium Boron Carbon ... Neon Sodium Magnesium Aluminum ... Argon ... Krypton

Atomic # 1 2 3 4 5 6

10 11 12 13

18 ... 36

Electron configuration 1s 1

1s 2 (stable) 1s 22s 1 1s 22s 2 1s 22s 22p 1 1s 22s 22p 2 ...

1s 22s 22p 6 (stable) 1s 22s 22p 63s 1 1s 22s 22p 63s 2 1s 22s 22p 63s 23p 1 ...

1s 22s 22p 63s 23p 6 (stable) ...

1s 22s 22p 63s 23p 63d 10 4s 246 (stable)

Adapted from Table 2.2,

Callister 6e.

SURVEY OF ELEMENTS



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Shell Designation Sub Shell No. of Electrons

K s 2

L s 2

p 6 8

M s 2p 6d 10 18

N s 2p 6d 10f 14 32



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The ionization energy is the energy required to remove asingle shell (valence) electron from an isolated neutral atom. The ionization energy is a measure of the ease with which conduction electrons can be created.

Ionization energy



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ElectronegativityElectronegativity

Is a measure of the ability of an Is a measure of the ability of an atom in a molecule atom in a molecule to attract electrons to itself.to attract electrons to itself.

ALLOYING ????

Concept proposed by Linus Pauling 1901-1994Concept proposed by Concept proposed by LinusLinus PaulingPauling 19011901--19941994



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• Columns: Similar Valence Structure

Electropositive elements:

Readily give up electronsto become + ions.

Electronegative elements:

Readily acquire electronsto become - ions.

He

N e

Ar

Kr

Xe

Rn

inert

gases

accept

1e

accept

2e

giv

e u

p 1

e

giv

e u

p 2

e

giv

e u

p 3

e

F Li Be

Metal

Nonmetal

Intermediate

H

Na Cl

Br

I

At

O

S Mg

Ca

Sr

Ba

Ra

K

Rb

Cs

Fr

Sc

Y

Se

Te

Po

Adapted from Fig. 2.6, Callister 6e.

THE PERIODIC TABLE



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Ionization energy vs. atomic numberIonization energy vs. atomic number

Ca

K

H

He

Li

B

Be C

NO

F

Ne

Na

MgAlSiP S

Cl

Ar

0

500

1000

1500

2000

2500

0 2 4 6 8 10 12 14 16 18 20

Element

Ionization energy (kJ/m

ol)



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Atomic radius vs. atomic numberAtomic radius vs. atomic number

Ca

K

H He

Li

B

Be

C N O F

Ne

Na

MgAl Si P S Cl

Ar

0

50

100

150

200

250

0 2 4 6 8 10 12 14 16 18 20

Element

Atomic Radius (pm)



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Atomic Radius

• Measures as distance from nucleus to nucleus and divided by 2.

• Unit commonly used is pm

• picometer= 10-12m

• Example: iodine atomic radius 140pm



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• ATOMIC NUMBER - N – r0

• Strength of material - correlate



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STRENGTH AND ATOMIC RADIUS?

• atomic number is the number of protons in the nucleus and also the number of electrons around the nucleus.

• as the atomic number increases, the number of electrons increases, the number of shells increases therefore atomic radius increases.

• as a result the valence shell and valence electrons become away from the nucleus reducing the attraction between outer shell and nucleus which in turn reduces the strength.

• thus, as atomic number increases, strength decreases and as atomic number decreases strength increases.



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Summary – Atomic Structure

N -atomic number

Atomic radius

Strong/weak attraction

Strength/hardness, M.P. of element/material correlated to atomic radius

Hardness/strength is inversely proportional to

the chemical bond length.

Cohen in 1985 predicted a new form of carbon (C=6) nitride (N=7) that exhibits an extremely low compressibility and super high hardness that exceeds that of diamond Ref:- B. Bhusan



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Bond LengthBond Length• Bond length is the distance between the nuclei of two bonded

atoms. ro = 1-2 Å – primary bonds; 2-5 Å–secondary bonds.

• Bond length increases with increase with atomic number/ radius.

ro

Fmax = F / Cohesive strength of materials

ro



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ro = 1 - 2 Å – primary bonds;2 - 5 Å – secondary bonds

• 1eV = 100 KJ/mole

• The difference in ro is too small but large

variation in properties.



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Periodic Trends in Lattice Energy

Coulomb’s Law

charge A X charge B

electrostatic force αdistance2

energy = force X distance therefore

charge A X charge B

electrostatic energy αdistance

cation charge X anion charge

electrostatic energy αcation radius + anion radius

α ∆H0lattice



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• Net bonding force ( )M Nr r

A BF r ≡ −

For N>M

r – centre to centre spacing

A an B – Material constants –type of bonds

M = 2 PRIMARY BOND

= 7 Van der Waals bonds

N = 7 to 10 Metallic bonds

= 9 to 12 Covalent bonds

= 10 to 12 Ionic bonds

First term and second terms of right hand side represents the attractive and repulsive forces



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Fmax = F / Cohesive strength of materials/Theoretical strength of material.

Fmax per unit cross-sectional area A, reaches to the value of materials Young's modulus E.

Structural steel E = 200 GPa. Therefore Fmax = 200 GPa.

But practically it is much lower value = 0.6 GPa?

Presence of Imperfections

maxmax

FE

Aσ = =



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Cohesive strength

courtesy: Callister

∞



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Charge and Attractive Force Control on Effective Ionic Radii



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ro

Liquid

Solid

Gas

E0



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CHEMICAL BONDS



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Marriage

Divorce

Forming

of a bond

is like

marriage

•More stable/

equilibrium

•exothermic

The

breaking

of a bond

relates to

a divorce.

•Less stable/

want of electrons to

fill outer shell

•Endothermic



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Types of Chemical Bonding

1. Metal with nonmetal:

electron transfer and ionic bonding

2. Nonmetal with nonmetal:

electron sharing and covalent bonding

3. Metal with metal:

electron pooling and metallic bonding



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(a) Primary bonds (Strong Bonds:-100 - 1500 kJ/mole/ 1-15 eV/bond - ro = 1 - 2 Å )

• Ionic bond (NON METAL+ METAL) • Covalent bond (NON METAL+ NON METAL)• Metallic bond (METAL+ METAL)

(b) Secondary bonds (Weak Bonds:- 1- 50 kJ/mole / 0.01- 0.5 eV/bond - ro = 2 - 5 Å )

Inter molecular, gases and liquids - feeble and less stable.• Di-pole bond• Hydrogen bond• Van der Waals bond

Classification of Chemical bonds



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All bonding forces are due to electrostatic charge. Opposite charges attract, Like charges repel.

This diagram shows the attraction and repulsion

between atoms: The outer ring (e-) is the electron

cloud. The inner red ring is the nucleus.

Bonding Overview



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Dog - Bone

Bonds

• the natural attraction between dogs and bones is like the attraction between opposite chargesand atomic bonds.



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Ionic Bonding



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Characteristics Ionic BondingMetal+ non metal

Transfer of electrons - easily exchange/ transfer electrons so as

to stabilize their outer shells (i.e. become more inert gas-like)

NON DIRECTIONAL – over lapping of electrons orbitals are spherically symmetric.

magnitude of the bond equal in all directions around an ion OR BOND POSSES EQUAL STRENGTH IN ALL DIRECTIONS.

Bond energies : 600 – 1500 kJ/mol (6-15 eV / atom)

High melting temperature, hard, brittle, electrically and thermally insulative.

Density???????????- high



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• Both ionically and covalent bonded materials are poor conductors, because electrons are not to leave their host atoms.

• The delocalized electrons of metals move easily along a potential gradient. Therefore metallic boned materils are good conductors



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Ionic bonds: One big greedy thief dog

• Ionic bonding is like one

big greedy dog steeling the

other dog's bone.

• The bone represents the

electron that is up for

grabs.

• When the big dog gains an

electron he becomes

negatively charged and the

little dog who lost the

electron becomes

positively charged.



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Properties of Ionic Compounds

• Crystalline structure.

• A regular repeating arrangement of ions in

the solid.

• Ions are strongly bonded.

• Structure is rigid.

• High melting points- because of strong

forces between ions.



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Conductivity

• Conducting electricity is allowing charges

to move.

• In a solid, the ions are locked in place.

• Ionic solids are insulators.

• When melted, the ions can move around.

• Melted ionic compounds conduct.

• Dissolved in water they conduct.



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• Predominant bonding in Ceramics

Give up electrons Acquire electrons

He

-

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

0.7

H

2.1

Be

1.5

Mg

1.2

Ca

1.0

Sr

1.0

Ba

0.9

Ra

0.9

Ti

1.5

Cr

1.6

Fe

1.8

Ni

1.8

Zn

1.8

As

2.0

Cs Cl

MgO

CaF 2

NaCl

O

3.5

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.

EXAMPLES: IONIC BONDING



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Covalent Bond



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Covalent Bond

covalent bond- a bond formed when atoms share

electrons equally

Two atoms share one or more pairs of outer-shell electrons.



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Characteristics Covalent Bonding

Non Metal+ Non metal

Sharing of electrons

strongest of all chemical bondsstrongest of all chemical bonds

Sharing – net decrease in P.E - good overlap of orbital- bring shared electrons close both the nuclei-Overlapping orbitalsare directionally oriented and not spherically symmetric.

DIRECTIONAL – magnitude of the bond is not equal in all directions .

Bond energies : 100 – 1000 kJ/mol (1-10 eV / atom)

High melting temperature, hard, brittle, electrically and thermally insulative

Density???????????- low



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Covalent Bonds: Dogs of equal strength

Think of two or more dogs with equal attraction to the bones.

The dogs (atoms) are identical, so the dogs share the pairs of available bones evenly.

Since one dog does not have more of the bone than the other dog, the charge is evenly distributed.



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Oxygen AtomOxygen Atom Oxygen AtomOxygen Atom

Oxygen Molecule (OOxygen Molecule (O22))

Oxygen O 8 1s22s22p4



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

He

-

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

0.7

H

2.1

Be

1.5

Mg

1.2

Ca

1.0

Sr

1.0

Ba

0.9

Ra

0.9

Ti

1.5Cr

1.6

Fe

1.8Ni

1.8

Zn

1.8

As

2.0

SiC

C(diamond)

H2O

C

2.5

H2

Cl2

F2

Si

1.8

Ga

1.6

GaAs

Ge

1.8

O

2.0

colu

mn IV

A

Sn

1.8

Pb

1.8

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 Cornell University.

EXAMPLES: COVALENT BONDING



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Density???????????- high –ionic bond

Density???????????- low-covalent bond



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Metallic Bonding



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Characteristics Metallic BondingMetal+ metal

Delocalized electrons /electron cloud/sea of electrons

NON DIRECTIONAL – over lapping of electrons orbitalsare spherically symmetric.

magnitude of the bond equal in all directions around an ion OR BOND POSSES EQUAL STRENGTH IN ALL DIRECTIONS

Bonding energy: wide range: -E (Hg) = 0.7eV/atom, E (W) = 8.8.eV/atom

Metals are lustrousDuctile, electrically and thermally conductive

Density???????????- high



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• The free electron theory of metals, which started with Drude (1902) and Lorentz(1916).



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Metallic BondsMellow dogs with plenty of bones to go around

• Think of a room full of puppies who have plenty of bones to go around and are not possessive of any one particular bone. This allows the electrons to move through the substance with little restriction. The model is often described as the "kernels of atoms in a sea of electrons."



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Metallic Bond

Metallic Bonding

Bonding between atoms with

low electronegativity. ie 1,2

or 3 valence electrons,

therefore there are many

vacancies in valence shell.



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Sea of Electrons

+ + + +

+ + + +

+ + + +

• Electrons are free to move through the

solid.

• Metals conduct electricity.



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• Metals have low ionization energieslow ionization energies, thus they do not have a

tight hold on their valence electrons.

• These outer electrons easily move around, as they do not

"belong" to any one atom, but are part of the whole metal crystal.

• Metals are good conductors of heat and electricity. This is directly due to the mobility of the electrons.

-Metals are malleable (can be flattened) and ductile (can be drawn into wires) because of the way the metal cations and electrons can "flow" around each other, without breaking the crystal structure.



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- Metals are lustrous. This is due to the uniform way that the

valence electrons of the metal absorb and re-emit light energy.

• The free electrons oscillate in the alternating electric field of the incident light beam, absorbing energy at all wavelengths and so making the metal opaque.

• In turn the oscillating electrons emit waves and in this way produce the reflected beam.



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PROPERTIES FROM BONDING



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Deep well

Shallow well

r0

r5

r4



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Deep wellShallow well

r0

r5r4

Weakly bonded solid Strongly bonded solid

Mean inter atomic distance

Bond energy



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Shallow well

Deep well



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Young’s modulus is a numerical constant, named for the

18th century English physician and physicist Thomas Young,

that describes the elastic properties of a solid

undergoing tension or compression in only one direction.



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Elastic Modulus - Metal Cutting



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TABLE 2.2 Mechanical Properties of Various Materials at Room Temperature

Metals (Wrought) E (GPa) Y (MPa) UTS (MPa)

Elongation

in 50 mm

(%)Aluminum and its alloysCopper and its alloys

Lead and its alloys

Magnesium and its alloysMolybdenum and its alloys

Nickel and its alloys

SteelsTitanium and its alloys

Tungsten and its alloys

69–79105–150

14

41–45330–360

180–214

190–20080–130

350–400

35–55076–1100

14

130–30580–2070

105–1200

205–1725344–1380

550–690

90–600140–1310

20–55

240–38090–2340

345–1450

415–1750415–1450

620–760

45–465–3

50–9

21–540–30

60–5

65–225–7

0

Nonmetallic materialsCeramics

DiamondGlass and porcelain

Rubbers

ThermoplasticsThermoplastics, reinforced

Thermosets

Boron fibersCarbon fibers

Glass fibers

Kevlar fibers

70–1000

820–105070-80

0.01–0.1

1.4–3.42–50

3.5–17

380275–415

73–85

62–117

—

——

—

——

—

——

—

—

140–2600

—140

—

7–8020–120

35–170

35002000–3000

3500–4600

2800

0

——

—

1000–510–1

0

00

0

0

Note: In the upper table the lowest values for E, Y, and UTS and the highest values for elongation are for pure metals.

Multiply gigapascals (GPa) by 145,000 to obtain pounds per square in. (psi), megapascals (MPa) by 145 to obtain psi.



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Problems in machining Titanium

Compared to high strength steels, titanium, due to its unique physical and chemical properties, poses the following problems:

• Lower thermal conductivity of Ti hinders quick dissipation of the heat caused by machining, leading to increased wear of cutting tools.

• Lower modulus of elasticity leads to high spring back, causing Ti parts to move away from the cutting tool.

• Lower hardness and high chemical reactivity of Ti lead to galling with the cutting tool.



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Metallic

Covalent

8013.3640NaClIonic

6603.4324Al

15384.2406Fe

- 390.768Hg

14104.7450Si

28005.21000MgO

34108.8849W

> 35507.4713C (Diam.)

M.T.

(oC)eV/atomkJ/molsubstanceBond



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Hydrogen

- 1890.087.7ArVan der

Waals

- 780.3635NH3

- 1010.3231Cl2

00.5251H2O

M.T.

(oC)eV/atomkJ/molsubstanceBond



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BOND ENERGIES AND BOND LENGTHS

0.13680C=C

0.154370C-C

0.18340C-Cl

0.14450C-F

0.14360C-O

0.15305C-N

0.11435C-H

0.12890

0.18535C=O

Bond length

nm

Bond energy

kJ/molBond

C C≡



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• Single bond: one pair of electrons are shared between two atoms

• Double bond: two pairs of electrons are shared between two atoms

• Triple bond: three pairs of electrons are shared between two atoms

Bond energy: the amount of energy required to break a bond holding two atoms

together.

triple bond > double bond > single bond

Bond length: the distance separating the nuclei

single bond > double bond > triple bond



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SUMMARY-BONDING -PROPERTIES

• HARDNESS/STRENGTH –Atomic radius• Tensile Vs Compressive force - more required?• MP• DIRECTIONAL / NON-DIRECTIONAL• DENSITY• CONDUCTIVITY• E- Young's modulus, MP, CTE.• α – CTE• DUCTILITY/BRITTLE ??• Metals are lustrous and opaque



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REVIEW QUESTIONS1. Correlate the strength of an element with atomic radius.

2. Explain the maximum number of electrons permitted in K, L, M, N shells.

3. Correlate the bond length with coefficient of thermal expansion.

4. Correlate the bond length with young's modulus of a material.

5. Correlate the bond length with melting point of a material.

6. What is the significance of cohesive strength of a material.

7. To separate two atoms, tensile or compressive force more required? Explain why?

8. Why ionic and covalent bonded material exhibit bad conductors of heat and

electricity ?

9. Why metallic bonds are opaque?

10.Why metallic bonded material are ductile?

11.Why ionic and covalent bonded material are hard and brittle?

12.What is directional and non-directional bonds?

13.Why covalent bonded material are less denser than metallic and ionic bonded

material?

14.What kind of bonding you expect in the following materials: NaCl, Cadmium

telluride, Bronze, SO2, RbI, FeC, C6H6, InAs, UH3, CaS, BN, GdO, GdTe.

15. Explain the need of highly pressurized injection of diesel into the cylinder.



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Secondary Bonds



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+-



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Vander Waals Forces

• Random electron movements creates

dipoles.

• Dipoles induces dipoles.

• Process is repeated.

• Oppositely charged dipoles attract.



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Dipole-dipole attraction

. More electronegative atom in a covalent

bond.

. Permanent dipoles.

. Attraction between dipoles.



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Dipole-dipole attraction-Factors

• More the number of electronegative

atoms, higher attraction

• More electronegative atoms, higher

attraction



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Why ice float on water



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Hydrogen Bond

. Hydrogen in a covalent bond with nitrogen, oxygen or fluorine.

. Permanent dipoles.

. Attraction between dipoles.

. Strongest intermolecular force.

. Intermolecular forces much weaker than normal

covalent bonds



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HYDROGEN BOND(Water)

H

O

Each hydrogen has 1 valence electron

Each hydrogen wants 1 more

The oxygen has 6 valence electrons

The oxygen wants 2 more

They share to make each other happy



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Water

• Put the pieces together

• The first hydrogen is happy

• The oxygen still wants one more

H O



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Water

• The second hydrogen attaches

• Every atom has full energy levels

H OH



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Hydrogen Bonding

H2O molecule

Tetrahedron



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Ice structure

Hydrogen bondingO

H H



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Ice, like all solids, has a wellIce, like all solids, has a well--defined structure; each water defined structure; each water molecule is surrounded by four neighboring Hmolecule is surrounded by four neighboring H22Os.Os. two of these are hydrogen-bonded to the oxygen atom on the central H2O molecule, and each of the two hydrogen atoms is similarly bonded to another neighboring H2O.



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When ice melts to form liquid water, the uniform three-dimensional tetrahedral organization of the solid breaks down as thermal motions disrupt, distort, and occasionally break hydrogen bonds. The methods used to determine the positions of

molecules in a solid do not work with liquids, so there is no unambiguous way

of determining the detailed structure of water. The illustration here is probably

typical of the arrangement of neighbors around any particular H2O molecule,

but very little is known about the extent to which an arrangement like this gets

propagated to more distant molecules



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Hydrogen Bonds (H-bonds)

Hydrogen bonds of water



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This computer-generated nanoscale view of liquid water isfrom the lab of Gene Stanley of Boston University .The oxygen atoms are red, the hydrogen atoms white.

Recent work from Richard SayKally's

laboratory

shows that the hydrogen bonds in liquid

water break and re-form so rapidly

(often in distorted configurations) that

the liquid can be regarded as a

continuous network of hydrogen-bonded

molecules.



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Ice floats because it is about 9% less dense than liquid water



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• Ice floats because it is about 9% less densethan liquid water.

• lakes and rivers freeze from top to bottom, allowing

fish to survive even when the surface of a lake has

frozen over. If ice sank, the water would be displaced

to the top and exposed to the colder temperature,

forcing rivers and lakes to fill with ice and freeze solid

= Acts as a insulator (eg. Igloo).



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WHY WATER HAS HIGH SPECIFIC HEAT



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• Low molecular mass - so more moles per kg

• It is higher than most other liquids or high heat carrying away capacity or J/Kg

• This is mainly because a given mass of water contains more molecules (and therefore more degrees of freedom in which to store energy) than the same mass of other liquids.



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• molecular weight, weight of a molecule of a substance expressed in atomic mass units(amu). The molecular weight may be calculated from the molecular formula of the substance; it is the sum of the atomic weights of the atoms making up the molecule.



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88.8101 % 115.99943OxygenO

11.1899 % 21.007947HydrogenH

Mass percentNumber of

atoms

Atomic

weightElementSymbol

Elemental composition of H2O:

The molecular weight of water is thus

(2 × 1.008)+(1×15.999) = 2.016+15.999 = 18.015 amu. or g/mol



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Elemental composition of NH3:

17.7554 % 31.007947HydrogenH

82.2446 % 114.00672NitrogenN

Mass percentNumber of atomsAtomic weightElementSymbol

Molar mass (molecular weight) of NH3 is 17.0306 g/mol



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Directional

inter-chain (polymer), inter-molecular

smallestSecondary

Nondirectional

(metals)

Variable

large-Tungsten

small-MercuryMetallic

Directional (semiconductors,

ceramics, polymer chains)

Variable

large-Diamond

small-Bismuth

Covalent

Nondirectional (ceramics)Large!Ionic

CommentsBond EnergyType

SUMMARY: BONDING



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Strong covalent bonding forces withinmolecules

Weak intermolecular forces betweenmolecules

Strong forces within molecules and weak forces between them.

ATOMIZATION ??

Injection - Diesel



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Elastic properties determined by binding between atoms. HOW?

Plastic properties determined by microstructure.



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Binding of atoms

microstructure

WHY?

Resistance to plastic deformation – HARDNESS

Resistance to elastic deformation – STIFFNESS

Resistance to crack propagation



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ENGINEERING STRESS

TRUE STRESS

Documents

1 CHEMICAL BONDS