INTERMOLECULAR FORCES:

LIQUIDS, SOLIDS & PHASE CHANGES (Silberberg, Chapter 12)

Ideal gas

molecules act

independently

PV=nRT

Liquid

molecules “stick” to

one another

Real gas

molecules

attract/repulse one

another

nRTnbVV

anP ))((

2

2

Intermolecular interactions

Centrosymmetric formamide dimer at the crystal geometry.

Blue shows regions of the molecular surface with a partial

positive charge, red those regions with a partial negative charge

and white regions which are neutral.

Types of Intermolecular interactions (intermolecular forces)

1. London force (“dispersion force”)

Electron cloud gives instantaneous dipole moment (e- build up at one

end of molecule leaving the nucleus at the other end partially

exposed)

Molecules stick together because their partial charges attract one another

London forces act between all types of molecules (polar and non-polar).

• Strength increases with molar mass : heavier molecules have more electrons (further from the nuclei)

bigger fluctuations in electron movement (they are more polarizable)

Eg. F2 and Cl2 are gases, Br2 is a liquid, I2 a solid at room temperature.

Strength influenced by molecule shape

Rod-shaped molecules have greater London forces of

attraction because the instantaneous dipoles can get

closer

Note: “van der Waals forces” is the generic term for ALL intermolecular forces.

It is NOT interchangeable for “dispersion forces.”

2. Dipole-dipole interactions:

Polar molecules have permanent partial charges.

Polar molecules that are brought close to one another tend to orient their dipole moments so that the plus

end of one molecule faces the minus end of another. This is a dipole-dipole interaction:

•Strength depends on magnitude of bond dipoles and shape of molecule. If bond dipoles cancel one

another within a molecule, then the molecule itself has no dipole moment.

3. Hydrogen bonding:

Occurs when a hydrogen atom is bonded to a strongly electronegative atom with lone pairs (N, O or F)

R-X-H … Y-R’ (X, Y = F, O or N)

O

H H

O

H H

O

H H

O

H H

Eg. water

+

+

-

-

covalent bond

hydrogen bond

Hydrogen bonding is the strongest of the intermolecular interactions

CCH3

O

O

H

C CH3

O

O

Hacetic acid

NH H

NHH

NH Hamine

In hexagonal ice I, (natural form of ice on Earth), each water molecule is H-bonded to 4

neighbouring molecules in a tetrahedral arrangement. H-bonds are less ordered at 20°C but they still

account for the cohesiveness of liquid water where each molecule remains H-bonded to an average

of 3.5 neighbours (at that temperature water molecules exchange their positions about 1011 times per

second !). Finally, the bonds weaken at higher temperatures and water vaporizes because thermal

motion causes an uncertain random orientation of the water dipoles. Vapour is like a gas where

molecules are too distant and move too fast to be able to interact.

http://wwwarpe.snv.jussieu.fr/td_2_eng/lsh.html

Table of intermolecular forces and energies

Type of interaction

Typical

energy

(kJ mol-1)

Interacting species

Intermolecular

London (dispersion) 0.5-40 Polarizable e- clouds

(all atoms/molecules)

dipole-dipole 5-25 dipole-dipole

(polar molecules)

Hydrogen bonding 10-40 Polar bond to H – dipole

(molecules where H is bonded to

electronegative atom with lone

pairs)

Chemical bonding

Ionic 400-4000 Cation-anion

Covalent 150-1100 atoms sharing electrons in a bond

Properties of liquids; influence of intermolecular interactions:

Viscosity

• Viscosity = resistance to flow

• High viscosity slow flow rate

• Strong intermolecular forces molecules held together can’t

move past one another easily high viscosity

• Viscosity as temperature because molecules have higher Ek

O

H H

H

H

H

H

H

Hvs

But London forces can add up to quite significant degree:

CH3(CH2)4CH=CHCH2CH=CH(CH2)7COOH CH3(CH2)7CH=CH(CH2)7COOH

Linoleic acid Oleic acid

Properties of liquids:

Surface tension

In a liquid, molecules experience equal

forces from all directions except at the

surface (experience a net inward force)

Strong intermolecular forces surface tension

Water meniscus curves upward because forces

between water molecules and the oxygen atoms

and OH groups in glass are stronger than the

forces between water molecules.

Mercury meniscus curves downwards

because forces between mercury atoms

are stronger than between Hg atoms and

the glass

capillarity

intermolecular interactions

occur in 3 dimensions

int. interactions

occur across &

below surface

net vector

for attractive

forces is down

Phase transition Name Example

Gas to liquid Condensation or liquefication

Dew

Liquid to gas Vaporisation Boiling water steam

Gas to solid Condensation or deposition

Frost

Solid to gas Sublimation Evaporation of CO2

Liquid to solid Freezing water ice

Solid to liquid Melting or fusion ice water

Phase transitions

solid liquid gas

melting

freezing

vaporizing

condensing

sublimination exothermic

endothermic

Vapour Pressure

• At a given T, molecules in liquid have a distribution of speeds.

• Some move fast enough to escape liquid (enter gas phase)

• At higher T, average kinetic energy is higher so greater fraction

of molecules can escape liquid

Consider liquid in sealed container:

Some fast moving molecules escape into gas phase

These gas molecules exert pressure on the liquid surface

The pressure as more molecules enter gas phase

Some gas molecules collide with surface and stick to it, re-entering the liquid phase

Eventually, rate of molecules evaporating equals rate of molecules condensing (equilibrium is reached)

vapour pressure : partial pressure exerted by a vapour over a liquid at a fixed temperature

Vapour pressure is characteristic for a given liquid (or solid)

Function of the intermolecular forces

Vapour pressure as temperature for a given liquid

Example: Which would you expect to have a higher boiling point, p-dichlorobenzene or o-dichlorobenzene?

Cl

Cl

Cl

Cl

Boiling point and melting point

In open container, atmosphere exerts a pressure on the liquid

surface. As T liquid molecules move more quickly & leave

surface more often. At some temperature, KE is great enough for

bubbles of vapour to appear in the liquid (“the liquid boils”).

Vapour pressure as function of

temperature and intermolecular

forces

The boiling point is the temperature at which the vapour

pressure equals the external pressure

The normal boiling point, Tb, of a liquid is defined as the boiling

point at 1 atm external pressure

The normal freezing point, Tf, of a liquid is the temperature at

which it freezes at 1 atm pressure

(Tf is only slightly dependent on pressure)

Strong intermolecular forces low vapour pressure

high boiling point

The normal boiling point, Tb, is defined as the boiling point at

1 atm external pressure

Heat of phase transitions

1. Gas cools:

2. Gas condenses:

3. Liquid cools:

5. Solid cools:

4. Liquid freezes:

Note:

• Within a phase, a change in heat is accompanied by a change in temperature (because of change in Ek). Heat lost or

gained depends on amount of substance, molar heat capacity for that phase and change in temperature.

•During a phase change, a change in heat occurs at constant temperature (energy needed to cause change in phase)

•Melting or boiling is endothermic; Condensation or freezing is exothermic.

Phase diagrams

• summarise phases of a substance under different conditions of temperature and pressure

Phase diagram of CO2:

solid liquid

gas

triple

point

critical

point

Phase diagram of H2O:

Amorphous and crystalline silicon dioxide.

When molten silica cools quickly it becomes a

glass. The atoms (red=O; black=Si) are

arranged in a disorderly fashion.

Quartz is a crystalline form of silica, SiO2. The

atoms (red=O; black=Si) are arranged in an

orderly network.

The Solid State (Silberberg, Chapter 12)

Amorphous: atoms/molecules in random arrangement (“frozen liquid”)

Crystalline: atoms/molecules in ordered pattern. Often have flat “faces”and definite angles at edges

(formed by orderly stacks of atoms)

Class Examples Characteristics

metallic s- and d- block elements malleable, ductile, lustrous, electrically and

thermally conducting

ionic NaCl, CsCl, KNO3, CuSO4 .

5 H2O

hard, rigid, brittle,

high melting points,

those soluble in water conduct electrical current;

when molten they are electrical conductors

network

(covalent

bonding)

B, C, BN, SiO2

hard, rigid, brittle,

very high melting points, insoluble in water

molecular

(covalent

bonding;

intermolecular

interactions)

H2O (ice), S8, I2, glucose,

sucrose, naphthalene

relatively low melting and boiling points, brittle

if pure

Crystalline solids can be classified in terms of their bonding:

portion of a 3-D

lattice

lattice point

unit

cell

The crystal lattice and unit cell

• All particles in a crystal are

packed in an orderly way

• Imagine placing a point at each

repeating part of the pattern –

the points form a regular array

called a crystal lattice

• All lattice points are identical

The smallest part of the crystal that, if repeated in 3-D gives the whole crystal, is called the unit cell

Seven crystal systems (different shapes of unit cells) and fourteen kinds of unit cells exist in nature

We consider only the cubic crystal system (all sides are equal length (called a) and all angles 90)

In the cubic system there are 3 possible packing arrangements:

Simple cubic /

Primitive cubic (P)

atom at each of the 8 corners

of cube

Body-centred cubic (I)

atom at each corner

& atom at centre of cube

Face-centred cubic (F)

atom at each corner

& atom in centre of each face of

cube

1 atom

at center

Coordination number = 8 Coordination number = 12 Coordination number = 6

atom

at 8 corners

1 8 –

Atoms/unit cell = x 8

= 1

1 8 –

atom

at 8 corners

1 8 – atom

at 8 corners

1 8 –

– atom

at 6 faces

1 2

Atoms/unit cell =

( x 8) + ( x 6)=4 1 8 – 1 –

2

Atoms/unit cell =

( x 8) + 1 = 2 1 8 –

Primitive cubic Body-centred cubic Face-centred cubic

Metals

• cations (identical spheres) packed tightly together and surrounded by a “sea” of electrons

• denser than most solids

• malleable and ductile

• conduct heat and electricity

• many pack in cubic unit cells

Copper: face-centred cubic

Figure 12.28 Packing of spheres.

simple cubic

(52% packing efficiency)

body-centered cubic

(68% packing efficiency)

hexagonal

unit cell

Figure 12.28 (continued)

closest packing of first

and second layers

layer a

layer a

layer b

layer c

hexagonal

closest

packing cubic closest

packing

abab… (74%) abcabc… (74%)

expanded

side views

face-centered

unit cell

expanded view space-filling

Ionic solids

•formed from oppositely charged ions (overall crystal is neutral)

•held together by attraction of cations and anions

•hard, rigid, brittle

•conduct electricity when molten or in solution; not as solids

•anions close-pack and cations fit into interstitial spaces

Allotropes (polymorphs) of Carbon

Diamond and graphite are examples of Network solids

• atoms covalently bonded

• entire crystal is one network

• hard, brittle, high melting points

• individual molecules, held in crystal structure by

intermolecular forces only

• physical properties depend on types & strengths

of intermolecular interactions

• relatively low melting points (100 - 500C)

Molecular crystal structures of carbon

C60 Face centred cubic

Manipulating atoms.

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The future …?