1 Intermolecular Forces 11. 2 INTERMOLECULAR FORCES Van der Waals’ forces Hydrogen bonds...

1

Intermolecular Intermolecular ForcesForces

1111

2

INTERMOLECULAR FORCES

Van der Waals’ forces

Hydrogen bonds

Dipole-dipole forces

London Dispersion forces

3

Johannes van der Waals (1837−1923).

Fritz London (1900−1954).

4

3 types of dipoles

3 types of dipoles

Permanent dipole

Permanent dipole

Instantaneous dipole

Instantaneous dipole

Induced dipole

Induced dipole

5

Permanent dipolePermanent dipole

A permanent dipole exists in all polar molecules as a result of the difference in the electronegativity of bonded atoms.

A permanent dipole exists in all polar molecules as a result of the difference in the electronegativity of bonded atoms.

6

Instantaneous dipoleInstantaneous dipoleAn instantaneous dipole is a temporary dipole that exists as a result of fluctuation in the electron cloud.

An instantaneous dipole is a temporary dipole that exists as a result of fluctuation in the electron cloud.

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Instantaneous dipoleInstantaneous dipoleAn instantaneous dipole is a temporary dipole that exists as a result of fluctuation in the electron cloud.

An instantaneous dipole is a temporary dipole that exists as a result of fluctuation in the electron cloud.

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Induced dipoleInduced dipoleAn induced dipole is a temporary dipole that is created due to the influence of neighbouring dipole (which may be a permanent or an instantaneous dipole).

An induced dipole is a temporary dipole that is created due to the influence of neighbouring dipole (which may be a permanent or an instantaneous dipole).

Permanent dipole

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11.11.22 Van der WaalsVan der Waals

’ Forces’ Forces

10

Van der Waals’ ForcesVan der Waals’ Forces



Dipole-Dipole

Interaction

Dipole-Dipole

Interaction

Dipole-Induced Dipole

Interaction

Dipole-Induced Dipole

Interaction

Instantaneous Dipole-Induced Dipole

Interaction

Instantaneous Dipole-Induced Dipole

InteractionLondon dispersion forces

11

Dipole-dipole interactionsDipole-dipole interactions• Electrostatic interactions between polar

molecules

12

Dipole-dipole interactionsDipole-dipole interactions• In a sample containing many polar

molecules

A balance of attraction and repulsion holding the molecules together

13

Dipole-induced dipole interactionsDipole-induced dipole interactions• When a non-polar molecule approaches a

polar molecule (with a permanent dipole), a dipole will be induced in the non-polar molecule.

Dispersion forces exist among all molecules and contribute most to the overall van der Waals’ forces.

14

Polarizability : - A measure of how easily the electron cloud of an atom/molecule can be distorted to induce a dipole

Polarization

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In general, size of electron cloud

electron cloud is less controlled by positive nuclei

extent of electron cloud distortion polarizability stronger dispersion forces

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Instantaneous dipole-induced Instantaneous dipole-induced dipole interactionsdipole interactions

11.2 Van der Waals’ forces (SB p.277)

• The instantaneous dipole arises from constant movement of electrons.

• Induces dipoles in neighbouring atoms or molecules

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Instantaneous dipole-induced Instantaneous dipole-induced dipole interactionsdipole interactions

20

Evidence for the presence of London dispersion forces

1. Condensation of noble gases at low temperatures to form liquids and solids

presence of attractive forces between non-polar atoms

E.g. Xe(g) Xe(s) Hsub = -14.9 kJ mol1

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Evidence for the presence of London dispersion forces

2. The non-ideal behaviour of gases

nRTbnVV

naP

2

van der Waals’ equation

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Strength of van der Waals’ forcesStrength of van der Waals’ forces11.2 Van der Waals’ forces (SB p.279)

Much weaker than covalent bonds

Less than 10% the strength of covalent bonds

van der Waals’ radius > covalent radius

I2

23

Q.59

The electron clouds of adjacent iodine molecules would repel each other strongly until the equilibrium van der Waals’ distance is restored.

24

The strength of van der Waals’ forces can be estimated by

melting point, boiling point, enthalpy change of fusion or enthalpy change of vapourization.

Higher m.p./b.p./Hfusion/Hvap stronger van der Waals’ forces

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Strength of van der Waals’ forcesStrength of van der Waals’ forces

Depends on three factors (in decreasing order of importance) : -

1. Size of molecule

2. Surface area of molecule

3. Polarity of molecule

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Size of molecule Size of molecule

Size of electron cloud

Size of electron cloud

Molecule Boiling point (o

C)

Helium

Neon

Argon

-269

-246

-186

Fluorine

Chlorine

Bromine

-188

-34.7

58.8

Methane

Ethane

Propane

-162

-88.6

-42.2

1. Size of 1. Size of MoleculeMolecule

Polarizability Polarizability

Dispersion forces Dispersion forces

Rel. molecular mass Rel. molecular mass

Sometimes !

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The van der Waals’ forces also increase with the surface area of the molecule.The van der Waals’ forces also increase with the surface area of the molecule.

2. Surface area of 2. Surface area of moleculemolecule

∵ van der Waals' forces are short-ranged forces

Atoms or molecules must come close together for significant induction of dipoles.

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Pentane (C5H12)2,2-dimethylpropane

(C5H12)

Boiling point: 36.1°C Boiling point: 9.5°C

Both are non-polar

Same no. of

electrons

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2,2-dimethylpropane moleculespentane molecules

larger contact areasmaller contact area

rod-shaped spherical in shape

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Pentane (C5H12)Pentane (C5H12)

Larger contact surface area Higher chance of forming induced

dipoles stronger dispersion forces

Boiling point = 36.1C

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2,2-dimethylpropane (C5H12)

2,2-dimethylpropane (C5H12)

Smaller contact surface area lower chance of forming induced

dipoles weaker dispersion forces

Boiling point = 9.5C

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3. Polarity of molecules3. Polarity of moleculesFor molecules with comparable molecular

sizesand shapes, dispersion forces are

approximatelyequal.

Polar/polar > polar/non-polar > non-polar/non-polar

Then, strength of van der Waals’ forces dependson the polarity of molecules involved

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C O

H3C

H3C

CH3

CH2

H2C

H3C

RMM = 58.0,

RMM = 58.0,

+

C O

H3C

H3C

+

Dipole-

dipole forces

+ Dispersion forces

CH3

CH2

H2C

H3C

Dispersion forces only

b.p. = 50C

b.p. = 0C

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Other examples : -

1. Graphite layers of large surface area

strong van der Waals’ forces

2. Polyethene vs ethene

(m.p. > 100C) (m.p. = 169C)

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Molecule

% contribution to the overall van der Waals' forces

Dipole-dipole

interaction

Dipole-induced dipole

interaction

Instantaneous dipole-induced dipole

interaction

C4H10 0 0 100

HCl 15 4 81

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Q.60(a)

CH3Cl < CH3Br < CH3Ib.p./C -24.2 3.56 42.4

The strength of dispersion forces increases

with molecular size/mass.Thus, b.p. increases with molecular size/massAlthough chloromethane is more polar,

the effect of dispersion forces outweightsthat of dipole-dipole forces.

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Q.60(b)

H3C

H2C

CH2

H2C

CH3H3C

H2C

CHCH3

CH3

CH3

C

H3CCH3

CH3< <

9.5C 27.7C 36.1C

Less spherical Greater surface

area

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Q.60(c)

F2 Cl2 ClF CH2Cl2

F2 < ClF < Cl2 < CH2Cl2

-188C -100C -34.0C 39.6C

ClF > F2. It is because

1.ClF has a greater molecular size than F2 and thus has stronger dispersion forces than F2

2. ClF is polar and its molecules are held by both dipole-dipole forces and dispersion forces.

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Q.60(c)

F2 < ClF < Cl2 > CH2Cl2

-188C -100C -34.0C 39.6C

Cl2 > ClF. It is because

1.Cl2 has a greater molecular size than ClF and thus has stronger dispersion forces than ClF.

2.Although ClF is polar, the effect of dispersion forces outweights that of dipole-dipole forces.

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Q.60(c)

F2 < ClF < Cl2 > CH2Cl2

-188C -100C -34.0C 39.6C

CH2Cl2 > Cl2. It is because

1.CH2Cl2 has a greater molecular size than Cl2 and thus has stronger dispersion forces than Cl2.

2.CH2Cl2 is polar and its molecules are held by both dipole-dipole forces and dispersion forces.

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Q.60(d)

NO < C2H6

RMM 28.0 28.0b.p./C -151 -89

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1 pm = 0.001 nm1 nm = 109 m

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C2H6 > NO. It is because

1.C2H6 has a greater molecular size and contact surface area than NO and thus has stronger dispersion forces than NO.

2.Although NO is polar, the effect of dispersion forces outweights that of dipole-dipole forces.

NO < C2H6

RMM 28.0 28.0b.p./C -151 -89

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The melting of a solid involves the separation of molecules from a regularly packed molecular crystal.

Thus, m.p. of a solid depends on

1. The strength of van der Waals’ forces

2. Packing efficiency of molecules in the crystal lattice

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Symmetry of molecule Packing efficiency m.p.

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Q.61

H3C

H2C

CH2

H2C

CH3H3C

H2C

CHCH3

CH3

CH3

C

H3CCH3

CH3< <

-160C -136C -20Cm.p.

Increasing symmetry

Increasing packing efficiency

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H3C

H2C

CH2

H2C

CH3H3C

H2C

CHCH3

CH3

<

Greater surface area Stronger van der Waals’ for

ces

Q.61

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Molecular Molecular CrystalsCrystals

11.11.44

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Molecular crystalsMolecular crystals

A molecular crystal is a structure which consists of individual molecules packed together in a regular arrangement by

weak intermolecular forces.

A molecular crystal is a structure which consists of individual molecules packed together in a regular arrangement by

weak intermolecular forces.

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IodineIodine

A unit cell of iodine crystal showing the orientation of I2 molecules

f.c.c. structure

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Dry iceDry ice

A unit cell of dry ice (CO2)

f.c.c. structure

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Structure and bonding of Structure and bonding of fullerenesfullerenes

Fullerenes are molecules composed entirely of carbon atoms, in the form of hollow spheres or hollow tubes.

53

Buckminsterfullerene (or buckyball)The first fullerene discovered was buckminsterfullerene (C60).

Buckminsterfullerene. A soccer ball.

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R.F. Curl

H.W. Kroto R.E. SmalleyDiscovered C60 in 1985

Awarded Nobel prize for Chemistry in 1996

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BuckminsterfullerenBuckminsterfullereneeC60

icosahedron正二十面體

truncated icosahedron

Cutting at 12

vertices

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BuckminsterfullerenBuckminsterfullerenee

12 pentagons by cutting at 12 vertices.

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BuckminsterfullerenBuckminsterfullerenee

20 hexagons by cutting 20 triangular faces.

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Named after the architect Richard Buckminster Fuller

A geodesic dome

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Each carbon atom is connected to three other carbon atoms by one double covalent bond and two single covalent bonds.

Buckminsterfullerene

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Each pentagon is connected to five hexagonsEach hexagon is connected to three pentagons and three hexagons alternately.

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Graphite is planar because it is made of hexagonal rings linked together.

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In C60, pentagonal rings prevent the sheet from being planar, making it spherical.

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0.1386 nm

0.1434 nm

Why are there two types of bond in C60 ?

64

The surface of the sphere is NOT planar

2pz orbitals are NOT parallel to one another

Delocalization of es is NOT favourable

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Family of fullerenes

C28 C32 C50 C70

Some of the more stable members of the fullerene family. (a) C28 (b) C32 (c) C50 (d) C70

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Molecular structure

C60 molecules held by dispersion forces

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Substance Melting point (°C)

Graphite 3730

Diamond 3550

Buckminsterfullerene

1070

Fullerene molecules are held together by weak van der Waals’ forces.

1. Melting point1. Melting point

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GraphiteGraphiteGraphiteGraphite

DiamondDiamondDiamondDiamond

insolubleinsolublein all liquid in all liquid

solventssolvents

insolubleinsolublein all liquid in all liquid

solventssolvents

2. Solubility2. Solubility

FullerenesFullerenesFullerenesFullerenesdissolvesdissolves

in benzenein benzenedissolvesdissolves

in benzenein benzene

Giant covalent structure

Molecular structure

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Buckminsterfullerenes are relatively strong and hard compared with most other molecular solids.

buckminsterfullerenemolecule (C60)

The C60 molecules are packed closely together in solid state.

3. Strength and hardness3. Strength and hardness

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Pure buckminsterfullerene (C60) is an electrical insulator.(no delocalized electrons)

4. Electrical conductivity4. Electrical conductivity

The buckminsterfullerene with potassium atoms filling the spaces between its molecules is a superconductor. Its formula is K3C60.

buckminsterfullerene

potassium atom

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Carbon nanotube (CNT) or buckytube

First discovered by Dr. Sumio Iijima in 1991

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It is formed by carbon atoms arranged in a long cylindrical hollow tube.

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The diameter of a nanotube is in the order of a few nanometres (109 m).

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A sheet of graphite rolls up

into a tube.

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Graphite sheet

Rolls up

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The ends of CNTs are capp

ed by half of a buckminster

fullerene molecule.

77

The tensile strength of carbon nanotubes is exceptionally high due to the strong covalent bonds holding the atoms together

The strongest materials on earth.~100 times stronger than steel

Applications : clothes, sports equipments, space elevators…

Properties of nanotubes

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Carbon nanotube is an electrical conductor because of the movement of delocalized electrons along the graphite sheets.

Depending on their structures, carbon nanotubes can be semi-conducting or as electrically conductive as metals.

Properties of nanotubes

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Hydrogen Hydrogen BondingBonding

11.11.55

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Evidence of hydrogen bondingEvidence of hydrogen bondingLook at the boiling points of some simple hydrides of Group IV to VII elements (p.87).

81

B.p. as molecular size

Group 4 hydrides are non-polar, only dispersion forces exist

Dispersion forces as molecular size .

82

B.p. as molecular size (dispersion > dipole-dipole)

However, H2O, HF and NH3 have abnormally high b.p.

There exist unusually strong dipole-dipole forces (H-bond)

All are polar

83

Formation of hydrogen bondingFormation of hydrogen bondingWhen a hydrogen atom is directly bonded to a highly electronegative atom (e.g. fluorine, oxygen and nitrogen), a highly polar bond is formed.

2.1 2.1 2.14.0 3.5 3.0

84

Electrostatic attractions exist between this partial positive charge and the

These attractions are called hydrogen bonds

lone pair electrons on a highly electronegative atom (i.e. fluorine, oxygen or nitrogen) of another molecule.

85

hydrogen bond

86

Formation of hydrogen bonds between H2O molecules.

hydrogen bond

87

Reasons for abnormal strength of H-Reasons for abnormal strength of H-bondbond

2. H atom does not have inner electrons.

its nucleus (proton) is partially exposed dueto unequal sharing of electron.

The partial positive charge on H is so concentrated that it can come very close

to the lone pair of a small & highly electronegative atom (F, O or N)

Abnormally strong dipole-dipole forces

1. the polarity of H–X bond is great when X is F , O , or N.

88

Two essential requirements for the formation of a hydrogen bond:

• One molecule must contain at least one H atom attached to a highly electronegative atom (i.e. F, O or N).

• The other molecule must contain an F, O or N atom that provides the lone pair of electrons.

89

Identify the hydrogen atoms of the following species that are capable of forming hydrogen bonding with water molecules.

adenine

glucose

Soluble in water

90

An exceptional case : -

C H

Cl

Cl

Cl

O C

CH3

CH3

+

Due to the combined effect of the three electronegative Cl atoms, the H atom becomes sufficiently positive to form hydrogen bond

H-bond

91

Relative strength of van der Waals’ forces, Relative strength of van der Waals’ forces, hydrogen bond and covalent bondhydrogen bond and covalent bond

PhenomenonEnergy

absorbed(kJ mol-1)

Forces overcome

He(s) He(g) 0.11Van der Waals’ f

orces

H2O(s) H2O(g) 46.90Hydrogen

bonds

O2(g) 2O(g) 494.00 Covalent bonds

92

Q.65

Tendency of H-bond formation : -

C – H < S – H < Cl – H < N – H < O – H < F – HEN 0.4 0.4 0.9 0.9 1.4 1.9

No lone pair on C

N is smaller than Cl

H can come closer

93

Q.66

Substance

Relative molecular

mass

Boiling point (°C)

NH3 17 -33.3

HF 20 19.5

H2O 18 100

HF > NH3 because

H – F bond is more polar than N – H bond

94

Q.66

Substance

Relative molecular

mass

Boiling point (°C)

NH3 17 -33.3

HF 20 19.5

H2O 18 100

H2O > HF because

H2O can form H-bonds more extensively, regardless of the fact that H-F bond is more polar than H-O bond.

95

Each NH3 molecule has only ONE lone pair.

hydrogen bond

On the average, each NH3 molecule can form only ONE hydrogen bond

96

Each HF molecule has only ONE hydrogen atom. On the average, each HF molecule can form only ONE hydrogen bond

97

hydrogen bond

Each H2O molecule has TWO hydrogen atoms and TWO lone pairs. On the average, each H2O molecule can form TWO hydrogen bonds

98

The lone pairs of oxygen atom of each water molecule forms hydrogen bonds with two hydrogen atoms of nearby water molecules

Structure and bonding of iceStructure and bonding of ice

a water molecule

hydrogenbond

hydrogenbond

hydrogen atom

oxygen atom

99

The two hydrogen atoms of each water molecule also form hydrogen bonds with the lone pairs of oxygen atoms of nearby water molecules.

hydrogenbond

hydrogenbond

hydrogen atom

oxygen atom

100

2

1

3

4

Each H2O molecule is bonded tetrahedrally to four H2O molecules

101

Open : the maximum number of hydrogen bonds can be formed

Regular : all molecules are held in positions by strong hydrogen bonds

In solid ice, the tetrahedral arrangement repeats over and over again, resulting in an open and regular network structure of water molecules.

102

The oxygen atoms in the structure of ice are arranged in a hexagonal shape.

103

The hexagonal symmetry of a snowflake reflects the structure of ice.

104

In liquid state, water molecules pack together more closely and randomly.

Hydrogen bonds are continuously formed and broken. Liquid water takes the shapes of the containers

105

Properties of iceProperties of ice1. Density

Most substances have higher densities in the solid state than in the liquid state.

ice

water

liquid paraffin

solid paraffin

Solid paraffin is denser

than liquid paraffin.

Solid paraffin is denser

than liquid paraffin.

106

ice

water

liquid paraffin

solid paraffin

Ice has a lower density than liquid water!At 0°C, density of ice = 0.92 g cm−3 density of liquid water = 1.00 g cm−3

107

This allows fish and other aquatic organisms to survive.

In cold weather, ice forms a layer on the top of a pond.

Ice acts as an insulator for the water beneath.

Ice

108

ExplanationIn ice, water molecules are arranged in an orderly manner in an open network structure because of extensive formation of hydrogen bonding.Open network structure!

Open network structure!

109

In this open structure, water molecules are further apart than they are in liquid water.

liquid water

melts

open structure collapses

water molecules tend to pack more closely together

ice

110

More H-bonds

Less H-bonds

More stable Open & regular

less stable Close & random

Energy is absorbed to break some of the hydrogen bonds

111

1. Melting point and boiling point

Effect of hydrogen bonding on Effect of hydrogen bonding on properties of waterproperties of water

The melting point (0°C) and boiling point (100°C) of water are much higher than expected.

A lot of energy is required to overcome the hydrogen bonds between water molecules and separate them.

112

High surface tension of water allows water striders to ‘walk’ on it.

2. Surface tension2. Surface tension

113

2. Surface tension2. Surface tensionSurface tension of molecular liquids arises from intermolecular forces.Stronger intermolecular forces leads to higher surface tension

LiquidRelative surface tension at 25C

Hexane 18.4

Methanol 22.6

Ethanol 22.8

Water 72.3

114

Water molecules at the surface are strongly attracted by neighboring molecules on the same surface.

intermolecular forces


The surface of water is like a tightly-stretched skin such that small insects can walk on it.

115

Water forms droplets rather than spreading out on leaf.


116

2. Surface tension2. Surface tensionIn a sample of water,each water molecule is attracted to neighboring water molecules in all directions and there is a balance of force.

117

2. Surface tension2. Surface tensionThere is an imbalance of force

forthe molecules at the surface.The water molecules at thesurface tend to be pulled

inwardsby other water molecules belowthe surface.As a result, water forms

dropletsrather than spreading out on

leaf.In other words, water tends toreduce its surface area by

takingthe spherical shape.

118

The high surface tension of water allows water to be transported to the top of trees by capillary action.

The tallest tree on earth

115.56m

119

3. Viscosity

The higher the viscosity of a liquid, the more slowly it flows.

The resistance of a liquid to flow.Viscosity

Viscosity arises from intermolecular forces

120

Strong hydrogen bonds hold water molecules together and do not allow them to move past one another easily.

Liquid Relative viscosity

Benzene 1

Water 15

Water has high melting and boiling points, high surface tension and high viscosity.

Surface tensionSurface tension

121

Effect of hydrogen bonding on Effect of hydrogen bonding on properties of alcoholsproperties of alcoholsConsider an ethanol molecule.

lone pairs of electrons

hydroxyl group

122

1. Boiling point1. Boiling pointEthanol molecules are held together by H-bonds. high boiling point

hydrogen bond

H-bond strengthH-bond strength

123

Alcohols vs Thiols (p.90)

Alcohol CH3OH C2H5OH C3H7OH C4H9OH

b.p.(C) 64.5 78 97 117

Thiol CH3SH C2H5SH C3H7SH C4H9SH

b.p.(C) 5.8 37 67 97

Dispersion forces : Thiol > alcohol

Boiling point : Alcohol > thiol

124

water ethanol

hydrogen bonds

2. Solubility in water2. Solubility in water

Ethanol and water are completely miscible

125

3. Viscosity3. Viscosity

Ethanol is viscous because of the presence of extensive intermolecular hydrogen bonds.

Ethanol is viscous, completely miscible with water, and has a high boiling point.

126

Viscosity as no. of OH groups per molecule

propan-1-ol

propane-1,2-diol

propane-1,2,3-triol

ViscosityViscosity

127

Explain the following.

Water is easily absorbed by tissue paper rather than forming droplets on it.

Tissue paper is composed of cellulose which is a natural polymer made of glucose molecules.

Thus, tissue paper can form extensive hydrogen bonds with water molecules.

128

Carboxylic Acids

H3C C

O

O

H

CH3C

O

O

H

Ethanoic acid exists as dimers, (CH3COOH)2, in vapour phase or in non-polar solvents

RMM = 260 = 120

H-bonds

RMM = 60

129

Q.67

Ethanoic acid molecules form H-bonds with polar solvent molecules rather than with other ethanoic molecules.

H3C C

O

O

H

O

HH

O

H

H

130

Q.67

Ethanoic acid molecules form H-bonds with polar solvent molecules rather than with other ethanoic molecules.

H3C C

O

O

H

RO

H

ROH

131

Intramolecular Hydrogen Bonding

N

O O

O

H

O

H

N

O

O

b.p. = 214C

b.p. = 279C

132

Formation of intramolecular hydrogen bonds prevents the formation of intermolecular hydrogen bonds lower boiling point

N

O O

O

H

O

H

N

O

O

b.p. = 214C

b.p. = 279C

133

Roles of Hydrogen Bonding in Biochemical SystemsProteins : polymers of amino acids

Peptide linkage

Primary structure : sequence of amino acids

134

Secondary structures : -

1. -pleated sheet

2. -helix

135

1. -pleated sheet

Intermolecular H-bonds formed between peptide linkages of adjacent protein chains

136

N C

C

O

H

C

N C

C

O

H

C

Both N and C of the N – C bond are sp2 hybridized

to facilitate delocalization of electrons

planar

137

N C

C

O

H

C

N – C bond has double bond character free rotation w.r.t. the bond axis is restricted N – H and C = O groups are held in opposite positions to facilitate the formation of inter-chain H-bonds

N C

C

O

H

C

planar

139

2. -helical structure

C

O

N

H

Intramolecular hydrogen bond

140

The molecular chains of protein can be held in position to give the -helical structure by forming intramolecular hydrogen bonds.

2. -helical structure

141

Both - and - structures were first suggested by Linus Pauling.

142

Tertiary structure

3-D arrangements of secondary structures

Myoglobin

143

Quaternary structure

3-D arrangements of tertiary structures

Haemoglobin

144

hydrogen bonds

DNA (DeoxyribonNuclei Acid) carries genetic information

Hydrogen bonding in DNAHydrogen bonding in DNA

145

hydrogen bonds

Effect of hydrogen bonding on DNAEffect of hydrogen bonding on DNAThe presence of intermolecular H-bonds helps maintain the double helical shape of DNA molecules.

146

The double helical structure is maintained by intermolecular hydrogen bonds formed between specific base pairs

Effect of hydrogen bonding on DNAEffect of hydrogen bonding on DNA

147


Cytosine Guanine

148


Thymine Adenine

149


Sequence of bases = genetic codeCAGACTTGCAAT…GTCTGAACGTTA…

Or

150

hydrogen bonds

Without hydrogen bond, life becomes impossible

151


152


O O

H

O

H+

+

It allows oxygen to dissolve in water

153

Change of states and Change of states and intermolecular forcesintermolecular forces

• 3 different states: solid, liquid and gas

•Change of states involves breaking or forming of intermolecular forces of

the molecular substances

154

Phase DiagramA phase diagram is a graph summarizing the conditions of pressure and temperature under which the different phases of a substance are stable. Phase state

E.g. C in the same state may have different phases

Graphite, diamond, C60

155

A phase is any homogeneous and physically distinct part of a system which is separated from other parts of the system by a definite physical boundary known as the phase boundary.

156

water

Phase boundary

oil

A system having two phases in the same liquid state

157

A system having three phases in two states

water

Phase boundaries

oilglass

158

A system having four phases in three states

water

Phase boundaries

oilglass

air

159

T / C

P / atm

Three regions in each of which only one phase is stable

Solid

Liquid

Vapour

Phase Diagram of Carbon Dioxide

160

T / C

P / atm

The three regions meet at three lines, along which two phases coexist in equilibrium.

LiquidSoli

d

Vapour


161

T / C

P / atm

AT is the sublimation curve

Solid

Liquid

Vapour


A

T

162

T / C

P / atm

Solid

Liquid

Vapour


A

T

CO2(s) CO2(g)sublimation

163

T / C

P / atm

Solid

Liquid

Vapour


A

T

AT shows the variation of sublimation temperature of carbon dioxide with external pressure

164

T / C

P / atm

TB is the melting curve

Solid

Liquid

Vapour


A

T

B positive slope

(most common)

165

T / C

P / atm

Solid

Liquid

Vapour


A

T

B

CO2(s) CO2(l)melting

166

T / C

P / atm

Solid

Liquid

Vapour


A

T

B

TB shows the variation of melting temperature of carbon dioxide with external pressure

167

T / C

P / atm

TC is the boiling curve

Solid

Liquid

Vapour


A

T

BC

168

T / C

P / atm

Solid

Liquid

Vapour


A

T

BC

CO2(l) CO2(g)boiling

169

T / C

P / atm

Solid

Liquid

Vapour


A

T

BC

TC shows the variation of boiling temperature of carbon dioxide with external pressure

170

T / C

P / atm

Solid

Liquid

VapourA

T

Q.62

(a) Condensation by T

171

T / C

P / atm

Solid

Liquid

VapourA

T

Q.62

(a) Condensation by P

172

T / C

P / atm

Solid

Liquid

VapourA

T

Q.62

(b) Boiling by T

173

T / C

P / atm

Solid

Liquid

VapourA

T

Q.62

(b) Boiling by P

174

T / C

P / atm

Solid

Liquid

VapourA

T

Q.62

(c) Freezing by T

175

T / C

P / atm

Solid

Liquid

VapourA

T

Q.62

(c) Freezing by P

176

T / C

P / atm

Solid

Liquid

VapourA

T

Q.62

(d) Melting by T

177

T / C

P / atm

Solid

Liquid

VapourA

T

Q.62

(d) Melting by P

178

T / C

P / atm

Solid

Liquid

VapourA

T

Q.62

(e) Sublimation by T

179

T / C

P / atm

Solid

Liquid

VapourA

T

Q.62

(e) Sublimation by P

180

T / C

P / atm

Solid

Liquid

VapourA

T

BC

T is the triple point where all three phases coexist in equilibrium.

5.1 atm

56.4C

181

Solid

Liquid

Vapour

Triple point

vapour pressure above solid

= vapour pressure above liquid

182

T / C

P / atm

Solid

Liquid

VapourA

T

BC

Dry ice sublimes when heated at 1atm

5.1 atm

56.4C

1 atm

183

T / C

P / atm

Solid

Liquid

VapourA

T

BC

Dry ice is so called because it never melts (goes wet) at normal pressure

5.1 atm

56.4C

1 atm

184

T / C

P / atm

Solid

Liquid

VapourA

T

BC

At P > 5.1 atm, dry ice melts to give liquid CO2 when heated

5.1 atm

56.4C

10 atm

Triple point video

185

T / C

P / atm

Solid

Liquid

VapourA

T

BCPc=

73atm

Tc= 31C

TC curve terminates at C beyond which the boundary between liquid and vapour disappears

Critical point

186

T / C

P / atm

Solid

Liquid

VapourA

T

BCPc=

73atm

Tc= 31C

Above Tc, the vapour cannot be condensed no matter how high the external pressure is

Gas

187

T / C

P / atm

Solid

Liquid

VapourA

T

BCPc=

73atm

Tc= 31C

Gas

Supercritical fluid

As dense as a liquidAs mobile as a gas

188

T / C

P / atm

Solid

Liquid

VapourA

T

BCPc=

73atm

Tc= 31C

Gas

Supercritical fluid

Decaffeination using supercritical CO2

189

T / C

P / atm

Solid

Liquid

VapourA

T

BCPc=

73atm

Tc= 31C

Gas

Supercritical fluid

In winter, T < Tc CO2 is in liquid phase

Q.63

190

T / C

P / atm

Solid

Liquid

VapourA

T

BCPc=

73atm

Tc= 31C

Gas

Supercritical fluid

In summer, T > Tc CO2 is in gas phase

Q.63

191

T / C

P / atm

Solid

Liquid

VapourA

T

BCPc=

73atm

Tc= 31C

Gas

Supercritical fluid

Q.64

1 atm

So T

As P , CO2(l) CO2(g)H>0

and CO2(g) CO2(s)

192

T / C

P / atm

Solid

Liquid

Vapour

Gas

A

B

C

T

Supercritical fluid

Phase Diagram of Water Negative

slope

m.p. when external P

(very rare)

193

T / C

P / atm

Solid

Liquid

Vapour

Gas

A

B

C

T

Supercritical fluid

Phase Diagram of Water

H2O(s) H2O(l)

P

194

T / C

P / atm

Solid

Liquid

Vapour

Gas

A

B

C

T

Supercritical fluid


H2O(s) H2O(l)

P

195

H2O(s) H2O(l)P

Ice melts below 0C when an extremely high pressure is applied to it.

This results in a decrease in friction between contact surfaces and makes possible

ice-skating and

the movement of tremendously massive glaciers.

196

T / C

P / atm

Solid

Liquid

Vapour

Gas

A

B

C

T

Supercritical fluid


0.006 atm

0.01C

1 atm

0C 100C

197

T / C

P / atm

Solid

Liquid

Vapour

Gas

A

B

C

T

Supercritical fluid


Pc=217.2atm

Tc=374C

198

Gas He H2 Ne N2 O2 CO2 NH3 H2O

Tc(K) 4.2 33.3 44.5 126 154 304 405 647

Critical temperature and the strength of intermolecular forces

Higher Tc

stronger intermolecular forces

greater deviation from ideal gas behaviour

non-polar

Polar with H bond

199

Q.68

Cu(NH3)4SO4NH3 does not exist.

Reason : -

Unlike H2O, each NH3 has one lone pair and three H atoms.

Thus, NH3 cannot form hydrogen bonds in the same way as H2O in the crystal lattice.

200

The END

201

How is the enthalpy of vaporization related to intermolecular forces of a simple molecular

substance like neon?

Back

The enthalpy of vaporization of a substance is the energy needed to vaporize one mole of the substance at its boiling point. Consider a substance like neon, which consists of single atoms, Neon liquefies when the temperature is lowered to –246 oC at 1 atm. The enthalpy of vaporization of the liquid at this temperature is 1.77 kJ mol-1. Some of this energy is needed to push back the atmosphere when the vapour forms. The remaining energy must be supplied to overcome the intermolecular attractions. Because each molecule in a liquid is surrounded by several neighbouring molecules, this remaining energy is some multiple of a single molecule-molecule interaction. Typically, this multiple is about 5.

Answer


202


(a)Comment on the relative strength of van der Waals’ forces in solid, liquid and gaseous bromine.

(a) The relative strength of van der Waals’ forces decreases in the order:

Solid bromine > liquid bromine > gaseous bromine

The van der Waals’ forces are highly dependent on the distance between adjacent molecules. It decreases exponentially with the separation between the molecules. Going from solid to liquid and then to gaseous state, the separation between molecules increases, so the van der Waals’ forces become weaker and weaker.

Answer

203


(b) Plastics are substances which have very strong van der Waals’ forces. Explain why the van der Waals’ forces are so strong in plastics.

(b) A large size of a molecule of plastics indicates that it has a large electron cloud which is more easily polarized. Therefore, the molecule of plastics is more likely induced to form an instantaneous dipole. Moreover, the molecule of plastics has an extensive surface area. These make plastics have very strong van der Waals’ forces between the molecules.

Answer

204


(c)Arrange the following substances in an increasing order of boiling point:

(i) N2, O2, Cl2, Ne

(ii) H2, Br2, He(c) (i) Ne < N2 < O2 < Cl2

(ii) He < H2 < Br2

Answer

Back

205

What is the consequence of two molecules approaching each other at a distance less than the sum of their van der Waals’ radii?

The electron clouds of the two molecules will repel each other, and the distance between the two molecules will increase until the repulsion is just balanced by the attraction.

Answer

Back11.3 Van der Waals’ radii (SB p.284)

206

The relative molecular masses and boiling points of five compounds are given below:

11.5 Hydrogen bonding (SB p.291)

Compound Relative molecular

mass

Boiling point (oC)

Ammonia (NH3) 17 -33.4

Ethanol (C2H5OH) 46 78

Hydrogen fluoride (HF)

20 19.5

Methanol (CH3OH) 32 66

Water (H2O) 18 100

207

(a)Ammonia, hydrogen fluoride and water have similar relative molecular masses, yet their boiling points are different. Explain why.


(a) H2O can form 2 hydrogen bonds per molecule while NH3 and HF can only form 1 hydrogen bond per molecule. Thus, the boiling point of water is higher than those of NH3 and HF. Besides, as F is more electronegative than N, the intermolecular hydrogen bond formed between HF molecules is stronger than that between NH3 molecules.

Answer

208

(b)Ethanol and methanol have similar structures, yet their boiling points are different. Explain why.

Back


(b) For molecules with similar structures, their boiling points depend on their relative molecular masses. As the relative molecular mass of ethanol is greater than that of methanol, the boiling point of ethanol is higher.

Answer

209


Why it takes much longer time to boil an egg on a mountain peak?

Back

The boiling point of water decreases with decreasing pressure. Although water boils easily at mountain peak, the cooking of an egg takes longer time. It is because the amount of heat delivered to the egg is proportional to the temperature of water.

Answer

210

(a)The formation of a hydrogen bond between two molecules RAH and R’B may be represented as:

R A H · · · · · · · B R’

(i) Suggest possible elements for A and B. What are their common features?

(ii) In which of the following ranges would you expect the strength of hydrogen bonds to lie?

0.1 – 10 kJ mol-1

10 – 50 kJ mol-1

100 – 400 kJ mol-1


Answer

211


(a) (i) A and B can be nitrogen, oxygen or fluorine. All of them are highly electronegative atoms, thus they form highly polar

molecules, resulting in the formation of hydrogen bonds.

(ii) 10 – 50 kJ mol-1

212

(b) Benzoic acid has an apparent relative molecular mass of 244 in hexane, but only 122 in aqueous solution. With the aid of diagrams, explain this phenomenon.


Answer

213


(b) The relative molecular mass of benzoic acid (C6H5COOH) is 122. In hexane, benzoic acid molecules form dimers with hydrogen bondings between the molecules.

However, in water, the benzoic acid molecules form hydrogen bonds with the water molecules.

214

(c) Cyclohexane (C6H12) is insoluble in water whereas glucose (C6H12O6) is miscible with water in all proportions.


Answer

215


(c) Cyclohexane is non-polar, and there are only weak van der Waals’ forces holding the molecules together. Thus, cyclohexane molecules do not form hydrogen bonds with water. On the other hand, glucose can form hydrogen bonds with water molecules via its OH groups. Therefore, glucose is soluble in water but cyclohexane is not.

Cyclohexane Glucose

Back

216


Name the types of bonding or intermolecular forces that are broken and formed in the following processes.

• H2O(s) H2O(g)

• 2Mg(s) + O2(g) 2MgO(s)

• H2(g) + F2(g) 2HF(g)

• 2Na(s) + 2H2O(l) 2NaOH(aq) + H2(g)

• CH3CH2OH(l) + 3O2(g) 2CO2(g) + 3H2O(l)Answer

217


Back

(a) Bond broken: hydrogen bond

(b) Bonds broken: metallic bond and covalent bond

Bond formed: ionic bond

(c) Bond broken: covalent bond

Bonds formed: covalent bond and hydrogen bond

(d) Bonds broken: covalent bond, metallic bond and hydrogen bond

Bonds formed: ionic bond and covalent bond

(e) Bonds broken: covalent bond and hydrogen bond

Bonds formed: covalent bond and hydrogen bond

1 Intermolecular Forces 11. 2 INTERMOLECULAR FORCES Van der Waals’ forces Hydrogen bonds...

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Transcript of 1 Intermolecular Forces 11. 2 INTERMOLECULAR FORCES Van der Waals’ forces Hydrogen bonds...