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Chapter 11Chapter 11

Fritz London 1900-1954.Studied intermolecularinduced-dipole interactions.

Johannes D. van der Waals1837-1923.*Studied intermolecular forcesin VPT relationships inliquids and gases.

Liquids, Solids, and MaterialsLiquids, Solids, and Materials

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Kinetic Energy

Attractive Intermolecular Forces

High temperaturesLow temperatures

Gases, Liquids and SolidsGases, Liquids and Solids

3

Ionic ForcesIonic Forces

++

-- -

-+

++

Ion-Ion

e.g. NaCl(s)

Ion-Dipole

e.g. NaCl(aq)

Ions form strong intermolecular forces with the polar molecule water.

Ion-ion forces are very strong and produce high boiling points and melting points.

NaCl dissolved in water

From now on we will concentrate on covalent molecules

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Intermolecular InteractionsIntermolecular InteractionsNon-Polar Molecules:

Polar Molecules:

Molecules withF-H, O-H, N-HBonds:

1. Dispersion (Disp) Forces

1. Dispersion (Disp) Forces2. Dipole-Dipole (DD) Forces

1. Dispersion (Disp) Forces2. Dipole-Dipole (DD) Forces3. Hydrogen Bonding (HB) Forces

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London Dispersion Forces London Dispersion Forces (aka van der Waals Forces)(aka van der Waals Forces)

The electrons on one atom are attracted tothe nucleus on a neighboring atom.

This creates an “instantaneous” (i.e. temporary)dipole on the first atom.

The instantaneous dipole on the first atomthen induces an instantaneous dipole on thesecond atom.

The two induced dipoles attract each other.

London Dispersion Forces are proportional to a molecule’s polarizability,which is the ease with which the electron cloud can be deformed.

The polarizability, is approximately proportional to the number of electronsin the molecule.

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Element MW # e- Tbp

[amu] [K]

He 4.0 2 4

Ne 20.2 10 27

Ar 39.9 18 87

Kr 83.9 36 121

Xe 131.3 54 166

The noble gases are spherical, non-polar atoms.

As the molecule (atom) increases in size, the boiling point increases (vapor pressure decreases).

Noble GasesNoble Gases

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London Dispersion ForcesLondon Dispersion ForcesLondon Dispersion Force # e-

London Dispersion Force MW

Tbp is a good measure of the strength of intermolecular forces.

A higher Tbp indicates stronger intermolecular forces.

Compd. MW # e- Tbp

[amu] [oC ]

F2 38 18 -188

Cl2 71 34 -34

Br2 160 70 +59

I2 254 106 +184

Compd. MW # e- Tbp

[amu] [oC ]

CH4 16 10 -161

C2H6 30 18 -88

C3H8 44 26 -42

n-C4H10 58 34 0

Disp

ersion

Fo

rces

Increasin

g

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Dispersion forces also depend upon the molecular shape.Pentane: C5H12 - MW = 72 amu

H3C

CH2

H2C

CH2

CH3

n-pentane

Tbp = 36oC

neopentane

Tbp = 9oC

H3C

C

CH3

CH3H3C

LargeContact Area(Surface Area)

SmallContact Area(Surface Area)

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Dipole-Dipole ForcesDipole-Dipole Forces

In liquids of polar molecules, oppositelycharged ends of the molecules tend toattract each other, causing partial alignment.

For molecules of roughly equal MW’s(i.e. with similar Dispersion Forces),the molecule with the higher dipole momentwill have a higher boiling point due to greaterDipole-Dipole forces.

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C

C

CH

HH

H H

H

H H

Propane

CH3CH2CH3

Acetonitrile

CH3CN

MW = 44 amu

Tbp = -42 oC

MW = 41 amu

Tbp = +82 oC

C C

H

H

H

N

0No Dipole Moment

= 3.9 Debye (D)

Large Dipole Moment

Why is the boiling point of acetonitrile so muchhigher than the boiling point of propane???

Dipole-Dipole Forces

Consider this….

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Hydrogen Bonding

F-H Bonds:

H-F

O-H Bonds:

e.g. H2O , CH3OH

N-H Bonds:

e.g. NH3, CH3NH2

O

H H

O

H H

N

HH

H

N

HH

H

H F H F

H bonding:between H and a veryelectronegativeatom

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Compd. MW # e- [amu] [D]

F-F 38 18 0

H-Cl 36 18 1.1

Tbp

[oC ]

-188

-85

H-F 20 10 1.8 +20 Why so high??

Hydrogen Bond

H F

+ -

H F

+ -

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Hydrogen BondingHydrogen BondingCompd. MW # e- [amu]

H2S 34 18

H2O 18 10

Tbp

[oC ]

-60

+100

PH3 34 18

NH3 17 10

-88

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CH3CH2OH HOCH2CH2OH CH3OCH3 CH3CH2CH3

MW=74 MW=62 MW=46 MW=44

Disp Disp Disp Disp

DD DD DD

HB

LowestBP

HighestBP

Ethylene glycol -a viscous liquid;

keeps radiator fluidfrom boiling over

in your car(198°)

Propane -a gas used

for charcoalGrills(- 42°)

Ether -a veryvolatileLiquid(35°)

Ethyl alcohol -a liquid with slightly lowerbp than water

(78°)

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Vapor PressureVapor Pressure

Vap. Press. (H2O)

liquid water

H2OH2O

H2O

H2O

H2OH2O

H2O

H2O

H2O

H2OH2OH2O

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Vapor Pressure Rises with Temperature

Temperature

Va

por

Pre

ssur

e

At higher temperatures, more molecules have sufficient energy to escape from the liquid.

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The Boiling PointThe Boiling PointThe boiling point of a liquid is the temperature at whichits vapor pressure is equal to the external pressure.

The “normal” boiling point is the temperature at whichthe vapor pressure of the liquid is equal to 760 torr (1 atm.)

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The Critical TemperatureThe Critical TemperatureH2O

T Vap. Press.[oC] [atm]

25 0.03

100 1

150 5

250 40

300 100

374 218

375

Tc = Critical Temperature

Highest temperature at which substance can be liquified.

Pc = Critical Pressure

Pressure required to liquifysubstance at Tc

Tc = 374 oCPc = 218 atm

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Substance Tc Pc

Helium -268 oC 2.3 atm

Oxygen -119 50

Ethane 32 48

Propane 97 42

Freon (CCl2F2) 112 40

Water 374 218

The Critical TemperatureThe Critical Temperature

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Phase DiagramsPhase Diagrams1. Why can’t you ice skate in Red Lake, Minnesota in February?

2. Why does carbon dioxide sublime, whereas water first melts and then vaporizes?

toSt. Paul

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Phase Diagram of HPhase Diagram of H22OO

Temperature (oC)

Pre

ssu

re (

atm

)

1

0 100

Solid Liquid Vapor

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A

D

B

C

Temperature (oC)

Pre

ssu

re (

atm

) Solid Liquid Vapor

AB: Liquid-Vapor Equilibrium

AC: Solid-Liquid Equilibrium

AD: Solid-Vapor Equilibrium

A: Triple Point Solid-Liquid-Vapor

B: Critical Point

0.01

4.6torr

Phase Diagram of HPhase Diagram of H22OO

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Pressure dependence of TPressure dependence of Tbpbp

Temperature (oC)

Pre

ssu

re (

atm

)

Liquid Vapor

100

1

2

120

Liquid Vapor

d(liq) >> d(vap)

V(liq) << V(vap)

Increased pressure shiftsequilibrium in directionof lower volume.

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Pressure dependence of TPressure dependence of Tmpmp

in Hin H22OO

Temperature (oC)

Pre

ssu

re (

atm

)

0

1

120

-20

Solid (ice) Liquid

d(ice) < d(liq)

V(ice) > V(liq)

Increased pressure shiftsequilibrium in directionof lower volume.

Solid Liquid

Application toice skating

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P1

Temperature (oC)

Pre

ssu

re (

atm

)

Solid Liquid Vapor

P2

Melting point decreases with Pressure

Melting + Vaporization

Sublimation

4.6torr

0.01

Phase Diagram of HPhase Diagram of H22OO

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Phase Diagram of “Normal” SubstancesPhase Diagram of “Normal” Substances

Temperature (oC)

Pre

ssu

re (

atm

)

P1

Solid Liquid Vapor

P2

Melting point increases with Pressure

Melting + Vaporization

Sublimation

CO25.1

-56

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Pressure dependence of TPressure dependence of Tmpmp

in “normal” substancesin “normal” substances

Temperature (oC)

Pre

ssu

re (

atm

)

1

100 Solid Liquid

d(sol) > d(liq)

V(sol) < V(liq)

Increased pressure shiftsequilibrium in directionof lower volume.

Solid Liquid

Tmpo Tmp

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Phase TransitionsPhase Transitions

Solid

Gas

Liquid

Melting Freezing

Vaporization

En

thal

py

Sublimation Deposition

Condensation

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Energy Changes of Phase TransitionsEnergy Changes of Phase Transitions

Solid

Liquid

Melting Freezing

En

thal

py

Melting (Fusion)

Hfus = Hliq - Hsol

= 6.01 kJ/mol (for H2O)

ENDOTHERMIC

Freezing (Crystallization)

Hcrys = Hsol - Hliq

= -6.01 kJ/mol (for H2O)

EXOTHERMIC

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Gas

Liquid

Vaporization

En

thal

py

Condensation

Vaporization

Hvap = Hgas - Hliq

= 40.7 kJ/mol (for H2O)

ENDOTHERMIC

Condensation

Hcond = Hliq - Hgas

= -40.7 kJ/mol (for H2O)

EXOTHERMIC

Energy Changes of Phase TransitionsEnergy Changes of Phase Transitions

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Heating CurvesHeating Curves

Heat Added (Joules)

Tem

per

atu

re (

oC

)

Tmp

Tbp

1. Heating solid

1

2. Melting solid to liquid2 3. Heating liquid

3

4. Vaporizing liquid to gas

4

5. Heating gas

5

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Quantitative considerations of heating curves

Hfus= 6.0 kJ/mol

Hvap= 40.7 kJ/mol

Tmp= 0 oC

Tbp = 100 oC

Heat capacities:

Cs(sol)= 2.09 J/g-oC

Cs(liq)= 4.18 J/g-oC

Cs(gas)= 1.84 J/g-oC

For water:

To heat 18 g of H2O from -40o to 140oC:

18 x 2.09 x 40 = 1.5 kJ (heating the ice)

1.0 x 6.01 = 6.0 kJ (melting the ice)

18 x 4.18 x 100 = 7.5 kJ (heating the water)

1.0 x 40.7 = 40.7 kJ (boiling the water)

18 x 1.84 x 40 = 1.3 kJ (heating the steam)

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Properties of LiquidsProperties of LiquidsViscosityViscosity

Viscosity is the resistance of a liquid to flowing.

High viscosity liquids (e.g. molasses, motor oil) flow slowly.

Low viscosity liquids (e.g. water, gasoline) flow quickly.

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Surface tension is a measure of the strength of intermolecular attractions which pull on molecules at the surface of a liquid.

It is because of its high surface tension thatwater tends to “bead” up on a waxy surface.

That’s because a sphere gives the minimum ratioof surface area to volume.

The high surface tension of water also causes molecules onthe surface to pack very closely together. This is why someinsects can “walk on water”.

Properties of LiquidsProperties of LiquidsSurface TensionSurface Tension

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Structure of SolidsStructure of Solids

Crystalline Solid

Ordered arrangement of atoms(or molecules) in 3-dimensionalstructure, called a lattice

e.g. Quartz (SiO2)

Amorphous Solid

Irregular (disordered)arrangement of atoms(or molecules)

e.g. Silica Glass (SiO2)

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The Crystalline LatticeThe Crystalline Lattice

Lattice: Three Dimensional array of points representing the centers of each atom in the crystal.

Lattice Point: Point corresponding to atom center

Unit Cell: Parallelopiped corresponding to minimum unit which can be used to replicate crystal

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Bonding in SolidsBonding in Solids

There are four classifications of solids, dependingon the type of bonds that are present.

• Covalent-Network Solids

• Ionic Solids

• Metallic Solids

• Molecular Solids

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Covalent-Network SolidsCovalent-Network SolidsForm of particles: Atoms connected in network of covalent bonds

Forces between particles: Covalent bonds

Properties: Very Hard Very high melting point Usually poor thermal and electrical conductivity

Examples: Diamond (C), Quartz (SiO2)

Diamond

Each carbon is connected to 4 others by a covalent bond

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Ionic SolidsIonic Solids

Form of particles: Positive and negative ions

Forces between particles: Electrostatic attractions

Properties: Hard and Brittle High melting point Poor thermal and electrical conductivity

Examples: All typical salts. e.g. NaCl, Ca(NO3)3, MgBr2

++

-- -

-+

++

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Metallic SolidsMetallic SolidsForm of particles: Atoms

Forces between particles: Metallic Bonds (due to delocalized valence electrons)

Properties: Soft to very hard Low to very high melting point Excellent thermal and electrical conductivity Malleable and Ductile

Examples: All metals. e.g. Cu, Fe, Sn, Au, Ag

Bonding due to delocalized valenceelectrons (shown in blue)

Strength of bonding varies betweendifferent metals, resulting in widerange of physical properties

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Molecular SolidsMolecular SolidsForm of particles: Atoms or molecules

Forces between particles: Dispersion Dipole-Dipole (if molecules are polar) Hydrogen Bonds (if O-H, N-H, F-H)Properties: Fairly soft Moderately low melting point (usually <200 oC) Poor thermal and electrical conductivity

Examples: Argon , CH4, CO2, C6H12O6 (sucrose), H2O