INTERMOLECULAR FORCESINTERMOLECULAR FORCESLIQUIDS AND SOLIDSLIQUIDS AND SOLIDS

Chapter 11

STATES OF MATTERSTATES OF MATTER

Weak attractive forces between

molecules

Intermolecular forces stronger

Strong intermolecular

forces

The state of a substance depends largely on the balance between the kinetic energies of the particles and the interparticle energies of attraction.

INTERMOLECULAR FORCESINTERMOLECULAR FORCESintramolecular force- covalent bond between atoms in a molecule

When a substance melts or boils the intermolecular forces are broken (not the covalent bonds). Many properties are governed by the strength of the intermolecular forces e.g. boiling point, melting point, vapor pressure, viscosity, etc.

intermolecular force - attraction between molecules

Strong431 kJ/mol

weak16 kJ/mol

(can also be ionic bond)

HCl boils at -85oC at 1 atm

A liquid boils when:

A solid melts when:

4 types of intermolecular forces:

• Ion-Dipole Forces

• Dipole-Dipole Forces

• London Dispersion Forces

• Hydrogen Bonding

electrostatic forces

Can you arrange them in order of increasing strength?

What is a dipole?+ -

van der Waals forces

Ion-Dipole ForcesIon-Dipole Forces

• Interaction between an ion and a dipole • Strongest of all intermolecular forces.• Ion-dipole interactions make it possible for ionic substances to

dissolve in polar solvents.

Dipole-Dipole ForcesDipole-Dipole Forces

• Interaction between two dipoles• Weaker than ion-dipole forces.• Polar molecules need to be close together.

Molecules in liquids are free to move results in both attractive and repulsive forces. overall and stronger attractive force

If two molecules have about the same mass and size, then intermolecular attractions (dipole-dipole forces) increase with increasing polarity.

London Dispersion ForcesLondon Dispersion Forces• Interaction between nonpolar atoms or molecules, but can also

occur between polar molecules• Occur only when atoms or molecules are close together • Weakest of all intermolecular forces.

Instantaneous dipoles formed!

While the electrons in the 1s orbital of helium would repel each other (and therefore tend to stay far away from each other), it does occasionally happen that they wind up on the same side of the atom.

These dipoles are temporary

Polarisability: the ease with which the distortion of the charge distribution occurs.

Greater polarisablility stronger dispersion

In general, larger molecules tend to have greater polarisability as they have more electrons which are further away from the nuclei.

ExerciseExercise

Explain why n-pentane has a higher boiling point than neopentane? Both have the molecular formula C5H12.

Hydrogen BondingHydrogen Bonding

Boiling pints

Special case of dipole-dipole forces

nonpolar

polar

Hydrogen bonds are dipole-dipole interactions between the H-atom in a polar bond (usually H-F, H-O or H-N) and an unshaired e- pair on a nearby small electronegative ion or atom (usually F, O, N)

Hydrogen bond

Small and electronegative!Why are we specifying F, O and N?

Hydrogen only has 1e- +ve nucleus is rather exposed.

Hydrogen can approach the small electronegative atom closely and interact strongly.

ExerciseExercise

Why does ice float on water?

ExerciseExercise

Which of the following molecules can hydrogen bond with itself?

CH2F2 NH3 CH3OH H3C– C–CH3

O

SUMMARYSUMMARY

PROPERTIES OF LIQUIDSPROPERTIES OF LIQUIDSViscosityViscosity The resistance of a liquid to flow

Which intermolecular forces also increase with molecular weight?

- related to the ease with which molecules can move past each other

Viscosity increases with molecular weight and decreases with higher temperature.

Surface tensionSurface tension A measure of the inward forces that must be overcome in order to expand the surface area of a liquid

Surface tension of water at 20oC:

0.0729 J/m2

Molecules in the interior are attracted equally in all directions.

Surface molecules experience a net inward force.

Molecules at the surface can pack more closely together

PHASE CHANGESPHASE CHANGESEvery phase change is accompanied by a change in energy of the system.

(fusion)Hfus

H (change in enthalpy)

Hvap Hcond

Hfreez

Hsub Hdepos

Hfus

Hvap

Hvap > Hfus

In the transition from liquid to vapour phase, the molecules must essentially sever all intermolecular interactions.

In melting, many of these interactions remain.

Hsub =

Plot of temperature versus heat added is a heating curve

e.g. Ice initially at -25oC is heated (constant P = 1 atm)

While melting, heat added is used to break intermolecular forces

While evaporating, heat added is used to break intermolecular forces

From the graph determine:

Hfus

Hvap

Recall: Hvap > Hfus

HAB

HCD

HEF

?

-30oC

0oC

50oC

D

C B

Awater

ice

ExerciseExercise

Calculate the enthalpy change, ΔH, when 100 g of water at 50 oC is cooled down to ice at -30 oC.

Given:Specific heat capacities: Water = 4.18 J/g K, Ice = 2.09 J/g KΔ Hfus = 6.01 kJ/mol

HAB = (4.18 J/g K)(100 g)(0 - 50 K) = -20.9 kJ

T = Tf - Ti

HBC = ΔHfreez = -ΔHfus = -(6.01 kJ/mol)(5.55 mol) = -33.4 kJ

n = (100 g)/(18.016 g/mol)

n = 5.55 mol

HCD = (2.09 J/g K)(100 g)(-30 - 0 K) = -6.27 kJ

HAD = HAB + HBC + HCD = -60.6 kJ

HAB

HCD

HBC = Hfreez

H = CmT

What is the enthalpy change when 100 g of ice at -30 oC is heated up to water at 50 oC?

• Some of the molecules on the surface of a liquid have enough energy to escape the attraction of the bulk liquid these molecules move into the gas phase.

VAPOUR PRESSUREVAPOUR PRESSURE

• As the number of molecules in the gas phase increases, some of the gas phase molecules strike the surface and return to the liquid.• After some time the pressure of the gas will be constant at the vapour pressure when liquid and vapour reach dynamic equilibrium.

Closed container

As the temperature increases, the fraction of molecules that have enough energy to escape increases.

The boiling point of a liquid is the temperature at which its vapor pressure equals atmospheric pressure.

Normal boiling point boiling point at 1 atm.

Why does water boil at a higher temperature at the coast than here in Jhb?

PHASE DIAGRAMSPHASE DIAGRAMSPhase diagrams display the state of a substance under various pressure and temperature conditions.

The solid lines show the conditions P,T conditions under which equilibrium exists between phases.

Line AD: solid -liquid interface- Each point along this line is the melting point of the substance at that pressure.

Critical point B: above this critical temperature and critical pressure the liquid and vapour are indistinguishable from each other.

Line AB: liquid-vapour interface- Each point along this line is the boiling point of the substance at that pressure.

Line AC: solid-vapour interface- Each point along this line is the sublimation point of the substance at that pressure.- Note: the substance cannot exist in the liquid state below A

Triple point A: the temperature and pressure condition at which all three states are in equilibrium.

The slope of the solid–liquid line is negative. The melting point decreases with increasing pressure.

Phase Diagram of WaterPhase Diagram of Water

Note the high critical temperature and critical pressure:

due to the strong van der Waals forces between water molecules.

WHY?

Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm.CO2 sublimes at normal pressures.

Phase Diagram of Carbon DioxidePhase Diagram of Carbon Dioxide

BONDING IN SOLIDSBONDING IN SOLIDSSolids can be:

crystalline

Particles are in highly ordered arrangement

amorphous

No particular order in the arrangement of particles.

or

The physical properties of crystalline solids (e.g. m.p., hardness) depend on the arrangement of particles and on attractive forces between particles.

There are 4 types of crystalline solids:

• Molecular• Covalent network• Ionic• Metallic

Most substances that are gases or liquids at room temperatureform molecular solids at low temperature.

Atoms or molecules are held together by intermolecular forces

dipole-dipole forces, London dispersion forces, hydrogen bonds

Molecular SolidsMolecular Solids

Because of these weak forces they are soft and have relatively low melting points (<200oC)

Benzene Toluene Phenolm.p./oC 5 -95 43 b.p./oC 80 111 182

OHCH3

Explain the m.p.’s and b.p.’s observed below:

Diamond Graphite

Which one is harder and has the higher m.p.? Explain.

Covalent-Network SolidsCovalent-Network Solids

Atoms are held together in large networks or chains by covalent bonds.

Covalent bonds much stronger than intermolecular forces

Harder solids and higher melting points than molecular solids.

Ionic SolidsIonic Solids

Ions are held together by ionic bonds

strength of the ionic bond depends on the charges of the ions

Ions pack themselves so as to maximize the attractions and minimize repulsions between the ions

Depends on relative size and charge of ions

Metallic SolidsMetallic SolidsConsist of entirely metal atoms.

Metals are not covalently bonded, but the attractions between atoms are too strong to be van der Waals forces.

Bonding due to valence electrons delocalized throughout the solid.

In general, the strength of bonding increases as the no. of electrons available for bonding increases

The m.p. for sodium is 97.5oC and for chromium is 1890oC. Explain.

SUMMARYSUMMARY