Intermolecular Forces

Intermolecular Forces
11 Intermolecular Forces

INTERMOLECULAR FORCES London Dispersion forces
Hydrogen bonds Van der Waals’ forces London Dispersion forces Dipole-dipole forces

Johannes van der Waals (1837−1923).
Fritz London (1900−1954).

3 types of dipoles Permanent dipole Instantaneous dipole
Induced dipole

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

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

Induced 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

11.2 Van der Waals’ Forces

Van der Waals’ Forces Van der Waals’ forces Dipole-Dipole Interaction
Dipole- Induced Dipole Interaction Instantaneous Dipole- Induced Dipole Interaction London dispersion forces

Dipole-dipole interactions
Electrostatic interactions between polar molecules

Dipole-dipole interactions
In a sample containing many polar molecules A balance of attraction and repulsion holding the molecules together

Dipole-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.

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

In general, size of electron cloud 
 electron cloud is less controlled by positive nuclei  extent of electron cloud distortion   polarizability   stronger dispersion forces

Instantaneous dipole-induced dipole interactions
11.2 Van der Waals’ forces (SB p.277) Instantaneous dipole-induced dipole interactions The instantaneous dipole arises from constant movement of electrons. Induces dipoles in neighbouring atoms or molecules

Instantaneous dipole-induced dipole interactions

Evidence for the presence of London dispersion forces
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 = kJ mol1

van der Waals’ equation
Evidence for the presence of London dispersion forces 2. The non-ideal behaviour of gases van der Waals’ equation

Strength of van der Waals’ forces
11.2 Van der Waals’ forces (SB p.279) Strength of van der Waals’ forces Much weaker than covalent bonds Less than 10% the strength of covalent bonds van der Waals’ radius > covalent radius I2

Q.59 The electron clouds of adjacent iodine molecules would repel each other strongly until the equilibrium van der Waals’ distance is restored.

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

Strength of van der Waals’ forces
Depends on three factors (in decreasing order of importance) : - Size of molecule Surface area of molecule Polarity of molecule

Size of electron cloud 
1. Size of Molecule Sometimes ! Molecule Boiling point (oC) Helium Neon Argon -269 -246 -186 Fluorine Chlorine Bromine -188 -34.7 58.8 Methane Ethane Propane -162 -88.6 -42.2 Size of molecule  Rel. molecular mass  Size of electron cloud  Polarizability  Dispersion forces 

2. Surface area of molecule
The van der Waals’ forces also increase with the surface area of the molecule. ∵ van der Waals' forces are short-ranged forces Atoms or molecules must come close together for significant induction of dipoles.

2,2-dimethylpropane (C5H12)
Pentane (C5H12) 2,2-dimethylpropane (C5H12) Boiling point: 36.1°C Boiling point: 9.5°C Both are non-polar Same no. of electrons

rod-shaped spherical in shape larger contact area smaller contact area
2,2-dimethylpropane molecules pentane molecules larger contact area smaller contact area

Pentane (C5H12) Boiling point = 36.1C Larger contact surface area
Higher chance of forming induced dipoles stronger dispersion forces

2,2-dimethylpropane (C5H12) Boiling point = 9.5C
Smaller contact surface area lower chance of forming induced dipoles weaker dispersion forces

Polar/polar > polar/non-polar > non-polar/non-polar
3. Polarity of molecules For molecules with comparable molecular sizes and shapes, dispersion forces are approximately equal. Then, strength of van der Waals’ forces depends on the polarity of molecules involved Polar/polar > polar/non-polar > non-polar/non-polar

Dispersion forces only
Dipole-dipole forces +  +  Dispersion forces RMM = 58.0, b.p. = 50C b.p. = 0C Dispersion forces only

Other examples : - 1. Graphite layers of large surface area  strong van der Waals’ forces 2. Polyethene vs ethene (m.p. > 100C) (m.p. = 169C)

% contribution to the overall van der Waals' forces Dipole-dipole
Molecule % contribution to the overall van der Waals' forces Dipole-dipole interaction Dipole-induced dipole Instantaneous dipole- induced dipole interaction C4H10 100 HCl 15 4 81

Q.60(a) CH3Cl < CH3Br < CH3I b.p./C The strength of dispersion forces increases with molecular size/mass. Thus, b.p. increases with molecular size/mass Although chloromethane is more polar, the effect of dispersion forces outweights that of dipole-dipole forces.

Less spherical Greater surface area
Q.60(b) < 9.5C C 36.1C Less spherical Greater surface area

F2 < ClF < Cl2 < CH2Cl2
Q.60(c) F2 Cl2 ClF CH2Cl2 F2 < ClF < Cl2 < CH2Cl2 -188C C C C ClF > F2. It is because 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.

F2 < ClF < Cl2 > CH2Cl2
Q.60(c) F2 < ClF < Cl2 > CH2Cl2 -188C C C C Cl2 > ClF. It is because Cl2 has a greater molecular size than ClF and thus has stronger dispersion forces than ClF. Although ClF is polar, the effect of dispersion forces outweights that of dipole-dipole forces.

F2 < ClF < Cl2 > CH2Cl2
Q.60(c) F2 < ClF < Cl2 > CH2Cl2 -188C C C C CH2Cl2 > Cl2. It is because CH2Cl2 has a greater molecular size than Cl2 and thus has stronger dispersion forces than Cl2. CH2Cl2 is polar and its molecules are held by both dipole-dipole forces and dispersion forces.

Q.60(d) NO < C2H6 RMM b.p./C

1 pm = nm 1 nm = 109 m

NO < C2H6 b.p./C RMM C2H6 > NO. It is because C2H6 has a greater molecular size and contact surface area than NO and thus has stronger dispersion forces than NO. Although NO is polar, the effect of dispersion forces outweights that of dipole-dipole forces.

The melting of a solid involves the separation of molecules from a regularly packed molecular crystal. Thus, m.p. of a solid depends on The strength of van der Waals’ forces Packing efficiency of molecules in the crystal lattice

Symmetry of molecule  Packing efficiency  m.p. 

Increasing symmetry Increasing packing efficiency
Q.61 < < m.p. -160C C C Increasing symmetry Increasing packing efficiency

Greater surface area Stronger van der Waals’ forces
Q.61 < Greater surface area Stronger van der Waals’ forces

11.4 Molecular Crystals

Molecular crystals A molecular crystal is a structure which consists of individual molecules packed together in a regular arrangement by weak intermolecular forces.

A unit cell of iodine crystal showing the orientation of I2 molecules
f.c.c. structure Iodine A unit cell of iodine crystal showing the orientation of I2 molecules

A unit cell of dry ice (CO2)
f.c.c. structure Dry ice A unit cell of dry ice (CO2)

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

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

R.F. Curl H.W. Kroto R.E. Smalley Discovered C60 in 1985 Awarded Nobel prize for Chemistry in 1996

Buckminsterfullerene
icosahedron正二十面體 truncated icosahedron Cutting at vertices

12 pentagons by cutting at 12 vertices.

20 hexagons by cutting 20 triangular faces.

Named after the architect Richard Buckminster Fuller
A geodesic dome

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

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

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

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

nm nm Why are there two types of bond in C60 ?

The surface of the sphere is NOT planar
 2pz orbitals are NOT parallel to one another  Delocalization of  es is NOT favourable

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

Molecular structure C60 molecules held by dispersion forces

1. Melting point Fullerene molecules are held together by weak van der Waals’ forces. Substance Melting point (°C) Graphite 3730 Diamond 3550 Buckminsterfullerene 1070

Giant covalent structure
2. Solubility Graphite Diamond insoluble in all liquid solvents Giant covalent structure Fullerenes dissolves in benzene Molecular structure

3. Strength and hardness Buckminsterfullerenes are relatively strong and hard compared with most other molecular solids. buckminsterfullerene molecule (C60) The C60 molecules are packed closely together in solid state.

4. Electrical conductivity
Pure buckminsterfullerene (C60) is an electrical insulator.(no delocalized electrons) The buckminsterfullerene with potassium atoms filling the spaces between its molecules is a superconductor. Its formula is K3C60. buckminsterfullerene potassium atom

Carbon nanotube (CNT) or buckytube
First discovered by Dr. Sumio Iijima in 1991

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

The diameter of a nanotube is in the order of a few nanometres (109 m).

A sheet of graphite rolls up into a tube.

Graphite sheet Rolls up

The ends of CNTs are capped by half of a buckminsterfullerene molecule.

Properties of nanotubes
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
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.

11.5 Hydrogen Bonding

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

B.p.  as molecular size  Group 4 hydrides are non-polar, only dispersion forces exist Dispersion forces  as molecular size .

All are polar 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)

Formation of hydrogen bonding
When 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 4.0 3.5 3.0

Electrostatic attractions exist between this partial positive charge and the
lone pair electrons on a highly electronegative atom (i.e. fluorine, oxygen or nitrogen) of another molecule. These attractions are called hydrogen bonds

hydrogen bond

hydrogen bond Formation of hydrogen bonds between H2O molecules.

Reasons for abnormal strength of H-bond
1. the polarity of H–X bond is great when X is F , O , or N. 2. H atom does not have inner electrons.  its nucleus (proton) is partially exposed due to 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

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.

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

An exceptional case : - H-bond
 + H-bond   Due to the combined effect of the three electronegative Cl atoms, the H atom becomes sufficiently positive to form hydrogen bond

Relative strength of van der Waals’ forces, hydrogen bond and covalent bond
Phenomenon Energy absorbed (kJ mol-1) Forces overcome He(s)  He(g) 0.11 Van der Waals’ forces H2O(s)  H2O(g) 46.90 Hydrogen bonds O2(g)  2O(g) 494.00 Covalent bonds

Tendency of H-bond formation : -
Q.65 Tendency of H-bond formation : - EN C – H < S – H < Cl – H < N – H < O – H < F – H No lone pair on C N is smaller than Cl H can come closer

Relative molecular mass
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

Relative molecular mass
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.

hydrogen bond Each NH3 molecule has only ONE lone pair.  On the average, each NH3 molecule can form only ONE hydrogen bond

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

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

Structure and bonding of ice
The lone pairs of oxygen atom of each water molecule forms hydrogen bonds with two hydrogen atoms of nearby water molecules a water molecule hydrogen bond hydrogen bond hydrogen atom oxygen atom

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

Each H2O molecule is bonded tetrahedrally to four H2O molecules
1 4 2 3

In solid ice, the tetrahedral arrangement repeats over and over again, resulting in an open and regular network structure of water molecules. Open : the maximum number of hydrogen bonds can be formed Regular : all molecules are held in positions by strong hydrogen bonds

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

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

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

Solid paraffin is denser than liquid paraffin.
Properties of ice 1. Density Most substances have higher densities in the solid state than in the liquid state. Solid paraffin is denser than liquid paraffin. 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 ice water liquid paraffin solid paraffin

Ice In cold weather, ice forms a layer on the top of a pond.
Ice acts as an insulator for the water beneath. This allows fish and other aquatic organisms to survive.

Open network structure!
Explanation In 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 structure collapses
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

Energy is absorbed to break some of the hydrogen bonds
Less H-bonds  less stable  Close & random More H-bonds  More stable  Open & regular

Effect of hydrogen bonding on properties of water
1. Melting point and boiling point 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.

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

Relative surface tension at 25C
Surface tension of molecular liquids arises from intermolecular forces. Stronger intermolecular forces leads to higher surface tension Liquid Relative surface tension at 25C Hexane 18.4 Methanol 22.6 Ethanol 22.8 Water 72.3

intermolecular forces
2. Surface tension intermolecular forces Water molecules at the surface are strongly attracted by neighboring molecules on the same surface. The surface of water is like a tightly-stretched skin such that small insects can walk on it.

Water forms droplets rather than spreading out on leaf.
2. Surface tension Water forms droplets rather than spreading out on leaf.

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

2. Surface tension There is an imbalance of force for
the molecules at the surface. The water molecules at the surface tend to be pulled inwards by other water molecules below the surface. As a result, water forms droplets rather than spreading out on leaf. In other words, water tends to reduce its surface area by taking the spherical shape.

The tallest tree on earth 115.56m
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

3. Viscosity Viscosity The resistance of a liquid to flow.
The higher the viscosity of a liquid, the more slowly it flows. Viscosity arises from intermolecular forces

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 tension

Effect of hydrogen bonding on properties of alcohols
Consider an ethanol molecule. hydroxyl group lone pairs of electrons

1. Boiling point Ethanol molecules are held together by H-bonds.
 high boiling point hydrogen bond H-bond strength

Dispersion forces : Thiol > alcohol
Alcohols vs Thiols (p.90) Alcohol CH3OH C2H5OH C3H7OH C4H9OH b.p.(C) 64.5 78 97 117 Thiol CH3SH C2H5SH C3H7SH C4H9SH 5.8 37 67 Dispersion forces : Thiol > alcohol Boiling point : Alcohol > thiol

2. Solubility in water hydrogen bonds water ethanol
Ethanol and water are completely miscible

3. 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.

no. of OH groups per molecule 
propan-1-ol propane-1,2-diol propane-1,2,3-triol Viscosity  as no. of OH groups per molecule  Viscosity

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.

Carboxylic Acids H-bonds RMM = 60
Ethanoic acid exists as dimers, (CH3COOH)2, in vapour phase or in non-polar solvents RMM = 260 = 120

Q.67 Ethanoic acid molecules form H-bonds with polar solvent molecules rather than with other ethanoic molecules.

Intramolecular Hydrogen Bonding
b.p. = 214C b.p. = 279C

Formation of intramolecular hydrogen bonds prevents the formation of intermolecular hydrogen bonds  lower boiling point b.p. = 214C b.p. = 279C

Roles of Hydrogen Bonding in Biochemical Systems
Proteins : polymers of amino acids Primary structure : sequence of amino acids Peptide linkage

Secondary structures : -
-pleated sheet -helix

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

planar Both N and C of the N – C bond are sp2 hybridized
to facilitate delocalization of  electrons

planar 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

2. -helical structure Intramolecular hydrogen bond

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

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

Tertiary structure 3-D arrangements of secondary structures Myoglobin

Quaternary structure 3-D arrangements of tertiary structures
Haemoglobin

Hydrogen bonding in DNA
DNA (DeoxyribonNuclei Acid) carries genetic information hydrogen bonds

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

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

Cytosine Guanine

Thymine Adenine

Sequence of bases = genetic code CAGACTTGCAAT… Or GTCTGAACGTTA…

Without hydrogen bond, life becomes impossible
hydrogen bonds

Without hydrogen bond, life becomes impossible

It allows oxygen to dissolve in water
Without hydrogen bond, life becomes impossible +  It allows oxygen to dissolve in water

Change of states and intermolecular forces
3 different states: solid, liquid and gas Change of states involves breaking or forming of intermolecular forces of the molecular substances

Phase Diagram A 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

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.

A system having two phases in the same liquid state
water Phase boundary oil

A system having three phases in two states
water Phase boundaries oil glass

A system having four phases in three states
water Phase boundaries oil glass air

Phase Diagram of Carbon Dioxide
Three regions in each of which only one phase is stable T / C P / atm Liquid Solid Vapour

The three regions meet at three lines, along which two phases coexist in equilibrium. T / C P / atm Liquid Solid Vapour

AT is the sublimation curve T / C P / atm Liquid Solid A T Vapour

CO2(s) CO2(g) sublimation T / C P / atm Liquid Solid A T Vapour

AT shows the variation of sublimation temperature of carbon dioxide with external pressure T / C P / atm Liquid Solid A T Vapour

TB is the melting curve positive slope (most common) T / C P / atm B Liquid Solid A T Vapour

CO2(s) CO2(l) melting T / C P / atm B Liquid Solid A T Vapour

TB shows the variation of melting temperature of carbon dioxide with external pressure T / C P / atm B Liquid Solid A T Vapour

TC is the boiling curve T / C P / atm B C Liquid Solid A T Vapour

CO2(l) CO2(g) boiling T / C P / atm B C Liquid Solid A T Vapour

TC shows the variation of boiling temperature of carbon dioxide with external pressure T / C P / atm B C Liquid Solid A T Vapour

Q.62 (a) Condensation by  T T / C P / atm Liquid Solid A T Vapour

Q.62 (a) Condensation by  P T / C P / atm Liquid Solid A T Vapour

Q.62 (b) Boiling by  T T / C P / atm Liquid Solid A T Vapour

Q.62 (b) Boiling by  P T / C P / atm Liquid Solid A T Vapour

Q.62 (c) Freezing by  T T / C P / atm Liquid Solid A T Vapour

Q.62 (c) Freezing by  P T / C P / atm Liquid Solid A T Vapour

Q.62 (d) Melting by  T T / C P / atm Liquid Solid A T Vapour

Q.62 (d) Melting by  P T / C P / atm Liquid Solid A T Vapour

Q.62 (e) Sublimation by  T T / C P / atm Liquid Solid A T Vapour

Q.62 (e) Sublimation by  P T / C P / atm Liquid Solid A T Vapour

Liquid Solid Vapour B C T A
T is the triple point where all three phases coexist in equilibrium. T / C P / atm B C Liquid Solid 5.1 atm 56.4C A T Vapour

Vapour Triple point Solid Liquid vapour pressure above solid
= vapour pressure above liquid

Liquid Solid Vapour B C T A Dry ice sublimes when heated at 1atm
T / C P / atm B C Liquid Solid 5.1 atm 56.4C A T Vapour 1 atm

Liquid Solid Vapour B C T A
Dry ice is so called because it never melts (goes wet) at normal pressure T / C P / atm B C Liquid Solid 5.1 atm 56.4C A T Vapour 1 atm

Liquid Solid Vapour B C Triple point video T A
At P > 5.1 atm, dry ice melts to give liquid CO2 when heated T / C P / atm B C 10 atm Liquid Solid Triple point video 5.1 atm 56.4C A T Vapour

TC curve terminates at C beyond which the boundary between liquid and vapour disappears
P / atm B C Pc= 73atm Tc= 31C Critical point Liquid Solid A T Vapour

Liquid Solid Gas Vapour
Above Tc, the vapour cannot be condensed no matter how high the external pressure is T / C P / atm B C Pc= 73atm Tc= 31C Liquid Solid Gas A T Vapour

Liquid Solid Gas Vapour As dense as a liquid Supercritical fluid
As mobile as a gas T / C P / atm B C Pc= 73atm Liquid Solid Gas A T Vapour Tc= 31C

Decaffeination using supercritical CO2
Supercritical fluid Decaffeination using supercritical CO2 T / C P / atm B C Pc= 73atm Liquid Solid Gas A T Vapour Tc= 31C

Liquid Solid Gas Vapour Supercritical fluid Q.63 In winter, T < Tc
CO2 is in liquid phase T / C P / atm B C Pc= 73atm Liquid Solid Gas A T Vapour Tc= 31C

Liquid Solid Gas Vapour Supercritical fluid Q.63 In summer, T > Tc
CO2 is in gas phase T / C P / atm B C Pc= 73atm Liquid Solid Gas A T Vapour Tc= 31C

Liquid Solid Gas Vapour Supercritical fluid Q.64
H>0 Supercritical fluid Q.64 As P , CO2(l)  CO2(g) So T  and CO2(g)  CO2(s) T / C P / atm B C Pc= 73atm Liquid Solid Gas A T Vapour 1 atm Tc= 31C

Phase Diagram of Water Liquid Solid Gas Vapour Supercritical fluid
T / C P / atm Solid Liquid Vapour Gas A B C T Supercritical fluid Phase Diagram of Water Negative slope (very rare) m.p.  when external P 

T / C P / atm Solid Liquid Vapour Gas A B C T Supercritical fluid Phase Diagram of Water H2O(s) H2O(l) P 

T / C P / atm Solid Liquid Vapour Gas A B C T Supercritical fluid Phase Diagram of Water H2O(s) H2O(l) P 

H2O(s) H2O(l) P  Ice melts below 0C 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.

Phase Diagram of Water Liquid Solid Gas Vapour Supercritical fluid B C
T / C P / atm Solid Liquid Vapour Gas A B C T Supercritical fluid Phase Diagram of Water 1 atm 0C 100C 0.006 atm 0.01C

Phase Diagram of Water Liquid Solid Gas Vapour Supercritical fluid B C
T / C P / atm Solid Liquid Vapour Gas A B C T Supercritical fluid Phase Diagram of Water Pc=217.2atm Tc=374C

Critical temperature and the strength of intermolecular forces
Gas He H2 Ne N2 O2 CO2 NH3 H2O Tc(K) 4.2 33.3 44.5 126 154 304 405 647 Polar with H bond non-polar Higher Tc  stronger intermolecular forces  greater deviation from ideal gas behaviour

Q.68 Cu(NH3)4SO4NH3 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.

The END

11.2 Van der Waals’ forces (SB p.280)
Back Let's Think 1 How is the enthalpy of vaporization related to intermolecular forces of a simple molecular substance like neon? Answer 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.

Check Point 11-2 Comment on the relative strength of van der Waals’ forces in solid, liquid and gaseous bromine. Answer 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.

Check Point 11-2 (b) Plastics are substances which have very strong van der Waals’ forces. Explain why the van der Waals’ forces are so strong in plastics. Answer (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.

Check Point 11-2 Answer Back
11.2 Van der Waals’ forces (SB p.280) Back Check Point 11-2 Arrange the following substances in an increasing order of boiling point: (i) N2, O2, Cl2, Ne (ii) H2, Br2, He Answer (c) (i) Ne < N2 < O2 < Cl2 (ii) He < H2 < Br2

Check Point 11-3 Answer Back
11.3 Van der Waals’ radii (SB p.284) Back Check Point 11-3 What is the consequence of two molecules approaching each other at a distance less than the sum of their van der Waals’ radii? Answer 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.

Relative molecular mass Hydrogen fluoride (HF)
11.5 Hydrogen bonding (SB p.291) Example 11-5A The relative molecular masses and boiling points of five compounds are given below: 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

11.5 Hydrogen bonding (SB p.291)
Example 11-5A Ammonia, hydrogen fluoride and water have similar relative molecular masses, yet their boiling points are different. Explain why. Answer (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.

Back Example 11-5A Ethanol and methanol have similar structures, yet their boiling points are different. Explain why. Answer (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.

Why it takes much longer time to boil an egg on a mountain peak?
11.5 Hydrogen bonding (SB p.293) Let's Think 2 Why it takes much longer time to boil an egg on a mountain peak? Answer 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. Back

Example 11-5B 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

Example 11-5B 11.5 Hydrogen bonding (SB p.296)
(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

Example 11-5B (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

Example 11-5B 11.5 Hydrogen bonding (SB p.296)
(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.

Example 11-5B Cyclohexane (C6H12) is insoluble in water whereas glucose (C6H12O6) is miscible with water in all proportions. Answer

Example 11-5B Back 11.5 Hydrogen bonding (SB p.296)
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

Check Point 11-5 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

Check Point 11-5 Back 11.5 Hydrogen bonding (SB p.297)
Bond broken: hydrogen bond Bonds broken: metallic bond and covalent bond Bond formed: ionic bond Bond broken: covalent bond Bonds formed: covalent bond and hydrogen bond Bonds broken: covalent bond, metallic bond and hydrogen bond Bonds formed: ionic bond and covalent bond Bonds broken: covalent bond and hydrogen bond

Intermolecular Forces

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Intermolecular Forces

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