Learning objective 2.16:The student is able to explain the properties (phase, vapor pressure, viscosity, etc.) of small and large molecular compounds in terms of the strengths and types of intermolecular forces.

•Distinguish between intermolecular and intramolecular attractions

•Put a list of compounds in order of increasing melting point, boiling point, and vapor pressure

•Use and label the parts of a phase diagram

•Use the Clausius-Clapeyron equation to relate temperature to vapor pressure of a substance

•Vapor pressure•Viscosity•Surface tension•ΔH of fusion•ΔH of vaporization•ΔH of sublimation

•Sublimation•Deposition•Condensation•Evaporation•Melting•Freezing•Freezing point•Boiling point

•Polar•Nonpolar•Dipole-dipole forces•Ion-dipole forces•Hydrogen “bonding”•London dispersion forces

Define:

What are three factors determine whether a substance is a solid, a liquid, or a gas: 1.The attractive intermolecular forces between particles that tend to draw the particles together. 2.Temperature: The kinetic energies of the particles (atoms, molecules, or ions) that make up a substance. Kinetic energy tends to keep the particles moving apart. 3.Pressure: pressure is increased or decreased as the volume of a closed container changes

Solid, Liquid, or GasSolid, Liquid, or Gas

There are several types of attractive intermolecular forces:

1.Ionic2.Ion-dipole forces 3.Dipole-dipole forces 4.Hydrogen bonding5.Induced-dipole forces

a) Ion-inducedb) Dipole-induced

6.London dispersion forces

Types of Attractive Forces

All of the intermolecular forces that hold a liquid together are called cohesive forces.

Ionic BondsIonic Bonds Electrons are transferred

Electronegativity differences are generally greater than 1.7 The formation of ionic bonds is always exothermic!

• Q is the charge.• r is the distance between the centers.• If charges are opposite, E is negative

• exothermic

• Same charge, positive E, requires energy to bring them together.• endothermic

19 1 2(2.31 10 )QQ

E x J nmr

1 2QQE

r

Coulomb’s LawCoulomb’s Law

An ion-dipole force is an attractive force that results from the electrostatic attraction between an ion and a neutral molecule that has a dipole.

•Most commonly found in solutions. Especially important for solutions of ionic compounds in polar liquids. •Ion-dipole attractions become stronger as either the charge on the ion increases, or as the magnitude of the dipole of the polar molecule increases.

Ion-Dipole Forces

•Dipole-dipole forces are attractive forces between the positive end of one polar molecule and the negative end of another polar molecule.

•They are much weaker than ionic or covalent bonds and have a significant effect only when the molecules involved are close together (touching or almost touching).

Dipole-Dipole Forces

Hydrogen Bonding

Bonding between hydrogen and more electronegative neighboring atoms such as oxygen and nitrogen

Hydrogen bonding between ammonia and water

Hydrogen Bonding in DNA

N O

OH

OP

O

OH

OHN

N

NNH2

NO

OH

OP

O

OH

OH

NH

O

O

CH3

TT AA

Thymine hydrogen bonds to Adenine

Hydrogen Bonding in DNA

CC GG

NO

OH

OP

O

OH

OH

N

NH2

O

N O

OH

OP

O

OH

OHN

NH

N

NH2

O

Cytosine hydrogen bonds to Guanine

Induced dipole forces result when an ion or a dipole induces a dipole in an atom or a molecule with no dipole. These are weak forces.

Induced-Dipole Forces

Ion-Induced Dipole Forces

An ion-induced dipole attraction is a weak attraction that results when the approach of an ion induces a dipole in an atom or in a nonpolar molecule by disturbing the arrangement of electrons in the nonpolar species.

A dipole-induced dipole attraction is a weak attraction that results when a polar molecule induces a dipole in an atom or in a nonpolar molecule by disturbing the arrangement of electrons in the nonpolar species.

Dipole-Induced Dipole Forces

London Dispersion Forces

The temporary The temporary separations of charge separations of charge that lead to the London that lead to the London force attractions are force attractions are what attract one what attract one nonpolarnonpolar molecule to its molecule to its neighbors.neighbors.

Fritz London Fritz London 1900-19541900-1954

London forces increase London forces increase with the size of the with the size of the molecules.molecules.

London Dispersion Forces

London Forces in Hydrocarbons

Boiling point as a measure of intermolecular attractive forces

Relative Magnitudes of Forces

The types of bonding forces vary in The types of bonding forces vary in their strength as measured by their strength as measured by average bond energy. average bond energy.

*Hydrogen bonding (12-16 kcal/mol )Dipole-dipole interactions (2-0.5 kcal/mol)

London forces (less than 1 kcal/mol)

Strongest

Weakest

*Ion-dipole interactions

Ionic bonds

Ion induced dipole interactions

Induced Dipole-dipole interactions

 

Identify the predominant intermolecular forces present in the solids of each of the following substances:

a. dipole-dipoleb. ion-dipolec. hydrogen bondingd. London dispersione. dipole-induced dipolef. ionicg. ion-induced dipole

Identify the predominant intermolecular forces present in the solids of each of the following substances:

a. dipole-dipoleb. ion-dipolec. hydrogen bondingd. London dispersione. dipole-induced dipolef. ionicg. ion-induced dipole

Identify the intermolecular forces present between the two substances listed:

a. dipole-dipoleb. ion-dipolec. hydrogen bondingd. London dispersione. dipole-induced dipolef. ionicg. ion-induced dipole

Identify the strongest intermolecular forces present between the two substances listed:

a. dipole-dipoleb. ion-dipolec. hydrogen bondingd. London dispersione. dipole-induced dipolef. ionicg. ion-induced dipole

What Is a Liquid? No, really, what IS a liquid??!!

This bottle contains both liquid bromine [Br2(l), the darker phase at the bottom of the bottle] and gaseous bromine [Br2(g), the lighter phase above the liquid]. The circles show microscopic views of both liquid bromine and gaseous bromine.

A liquid is a state of matter in which a sample of matter:•is made up of very small particles (atoms, molecules, and/or ions). •flows and can change its shape. •is not easily compressible and maintains a relatively fixed volume.

The particles that make up a liquid: •are close together with no regular arrangement, •vibrate, move about, and slide past each other.

What Is a Liquid?

More Properties of a Liquid

Surface Tension: The resistance to an increase in its surface area (polar molecules, liquid metals).

Capillary Action: Spontaneous rising of a liquid in a narrow tube.

Even More Properties of a Liquid

Viscosity: Resistance to flow

High viscosityHigh viscosity is is an an

indication of indication of strong strong

intermolecularintermolecular forces forces

Microscopic view of a liquid.

Microscopic view after evaporation.

When a liquid is heated sufficiently or when the pressure on the liquid is decreased sufficiently, the forces of attraction between molecules do not prevent them from moving apart, and the liquid evaporates to a gas.

•Example: The sweat on the outside of a cold glass evaporates when the glass warms. •Example: Gaseous carbon dioxide is produced when the valve on a CO2 fire extinguisher is opened and the pressure is reduced.

EvaporationEvaporation is the change of a liquid to a gas.

Condensation Condensation is the change from a vapor to a condensed state (solid or liquid).

When a gas is cooled sufficiently or, in many cases, when the pressure on the gas is increased sufficiently, the forces of attraction between molecules prevent them from moving apart, and the gas condenses to either a liquid or a solid.

•Example: Water vapor condenses and forms liquid water (sweat) on the outside of a cold glass or can. •Example: Liquid carbon dioxide forms at the high pressure inside a CO2 fire extinguisher.

Microscopic view of a gas.

Microscopic view after condensation.

The vapor pressure of a liquid is the equilibrium pressure of a vapor above its liquid (or solid)The vapor pressure of a liquid varies with its temperature, as the following graph shows for water. The line on the graph shows the boiling temperature for water.

Vapor Pressure

As the temperature of a liquid or solid increases its vapor pressure also increases. Conversely, vapor pressure decreases as the temperature decreases.

•Types of Molecules: the types of molecules that make up a solid or liquid determine its vapor pressure. If the intermolecular forces between molecules are:

Factors That Affect Vapor Pressure

substancevapor

pressure at 25oC

diethyl ether 0.7 atm

bromine 0.3 atm

ethyl alcohol 0.08 atm

water 0.03 atm•Surface Area: the surface area of the solid or liquid in contact with the gas has no effect on the vapor pressure.

Temperature Dependence of Vapor Pressures

• The vapor pressure above the liquid varies exponentially with changes in the temperature.

• The Clausius-Clapeyron equation shows how the vapor pressure and temperature are related. It can be written as:

CTR

HP vap

1ln

Clausius – Clapeyron Equation

• A straight line plot results when ln P vs. 1/T is plotted and has a slope of Hvap/R.

• Clausius – Clapeyron equation is true for any two pairs of points.

122@

1@ 11ln

TTR

H

P

P vap

Tvap

Tvap

CTR

HP vapvap

1ln

Write the equation for each and combine to get:

Using the Clausius – Clapeyron Equation

• Boiling point - the temperature at which the vapor pressure of a liquid is equal to the pressure of the external atmosphere.

• Normal boiling point - the temperature at which the vapor pressure of a liquid is equal to atmospheric pressure (1 atm).

E.g. Determine normal boiling point of chloroform if its heat of vaporization is 31.4 kJ/mol and it has a vapor pressure of 190.0 mmHg at 25.0°C.E.g.2. The normal boiling point of benzene is 80.1°C; at 26.1°C it has a vapor pressure of 100.0 mmHg. What is the heat of vaporization?

334 K

33.0 kJ/mol

Phase Transitions• Melting: change of a solid to a

liquid.• Freezing: change a liquid to a

solid.• Vaporization: change of a liquid

to a gas. • Condensation: change of a gas to

a liquid. • Sublimation : Change of solid to

gas• Deposition: Change of a gas to a

solid.

H2O(s) H2O(l)

H2O(l) H2O(s)

H2O(l) H2O(g)

H2O(g) H2O(l)

H2O(s) H2O(g)

H2O(g) H2O(s)

Water phase changes

Temperature remains constant during a phase change.

Energy

Energy of Heat and Phase Change

• Heat of vaporization: heat needed for the vaporization of a liquid.

H2O(l) H2O(g) H = 40.7 kJ• Heat of fusion: heat needed for

the melting of a solid.

H2O(s) H2O(l) H = 6.02 kJ• Temperature does not change

during the change from one phase to another.

E.g. Start with a solution consisting of 50.0 g of H2O(s) and 50.0 g of H2O(l) at 0°C. Determine the heat required to heat this mixture to 100.0°C and evaporate half of the water. 130 kJ

Phase Diagrams

• Triple point- Temp. and press. where all three phases co-exist in equilibrium.

• Critical temp.- Temp. where substance must always be gas, no matter what pressure.

• Critical pressure- vapor pressure at critical temp.

• Critical point- point where system is at its critical pressure and temp.

Phase changes by Name

Water

Carbon dioxide

Carbon

Sulfur