Intermolecular Forces,

Liquids,

and Solids

Chapter 11

• What are the 3 states of matter?• What physical properties do we know

about them?• Liquids and solids are similar

compared to gases…– Incompressible– Density doesn’t change with

temperature• Because solids and liquids are

molecules that are so close together while gases are so far apart.

• What holds them so close together?

States of Matter

Intermolecular Forces• Physical Properties of solids and

liquids is due largely to Intermolecular Forces (IMFs).

• IMFs are forces that exist between molecules within a solid or a liquid

• There are many different types of IMFs.

• By understanding the nature & strength of IMFs we can relate composition & structure to physical properties.

Types of IMFs• Strong Intermolecular Forces:

– Covalent Bonding– Ionic Bonding

• Weak Intermolecular Forces:– Dipole-Dipole– Hydrogen “bonding”– London Dispersion Forces

• During phase changes, molecules stay intact… NRG used to overcome forces

What is a dipole?Dipole – Dipole Forces:• Attraction between molecules with dipole

moments.• Maximizes + to – interactions and minimizes +

to + or – to – interactions.• 1% the strength of ionic bonds• Force strength increases as polarity increases• Not important for gases… why?

Ion – Dipole Force:• Force between Ion and dipole of a polar

molecule.• Weaker than Dipole – Dipole forces.

+- +-

+-

+

-+

-

+-

+-+-

+-

+-

Hydrogen Bonding• Special dipole-dipole attraction

– Hydrogen covalently bonded to highly electronegative elements

– Especially strong with: N, O, F (because they have high electronegativity and they are small enough to get close)

• Stronger than other dipole-dipole attractions

• Important for bonding of Water and DNA

• Effects Boiling Points

CH4

SiH4

GeH4SnH4

PH3

NH3 SbH3

AsH3

H2O

H2SH2Se

H2Te

HF

HI

HBrHCl

Boiling Points

0ºC

100

-100

200

Water

+-

+

London Dispersion Forces• Instantaneous dipoles

• Electrons move randomly, and creates a momentary nonsymmetrical distribution of charge (even in nonpolar molecules)

• Induces short-lived dipole with neighboring molecules

• Exist between ALL molecules, but are the weakest forces of attraction

Example

H H H HH H H H

+ H H H H

+ - +

London Dispersion Forces• Weak and short lived

• Last longer at Low Temp.

• Polarizabliity – ability to distort the electron cloud; “squishiness” of the molecule.– increases with the number of electrons

in a molecule… CCl4 experiences greater LDFs than CH4.

• Larger molecules = more LDF = higher melting and boiling points

IMF Summary• London Dispersion Forces• Dipole-Dipole Forces• Hydrogen Bonding

ALL 3 are sometimes referred to as

“Van der Waals Forces”• See flow chart on pg 417 to review

what types of forces are involved with different types of bonds.

Increasingstrength

The Liquid State• Many of the properties of liquids

are due to the internal attraction of atoms:– Surface Tension

– Beading

– Capillary Action

– Viscosity

• Stronger IMFs cause each of these to increase.

Surface tension

• Molecules in the middle are attracted in all directions.

Molecules at the top are only pulled inside.

Minimizes surface area.

•High IMF means Higher surface tension

Capillary Action

• Liquids spontaneously rise in a narrow tube.

• IMFs are cohesive (connect like things)

• Adhesive forces form between two separate “like” objects (polar molecules attract to polar bonds in material making up container)

Beading• If a polar substance

is placed on a non-polar surface. – There are cohesive,– But no adhesive

forces.

• And Visa Versa H2O

Wax Paper

Viscosity

• Measure of a liquid’s resistance to flow.

• Increases with IMFs

• Increases with molecular size.

Liquids… a structural model

• Strong IMFs (like solids)

• Considerable molecular motion (like gases)

Phase Changes• Phase of a substance determined by…

1. Temperature

2. Pressure

3. Strength of IMFs

• Stronger Forces = Bigger Effect– Hydrogen Bonding

– Polar Molecule (dipole – dipole)

– London Dispersion Forces

Decreasing Strength

• Melting / Freezing

• Vaporization / Condensation

• Sublimation / Deposition

We’ll go more into the energy changes, heating curves, and

phase diagrams later in this unit…

Types of Phase Changes:

Solids

• Two major types.

• Amorphous- those with much disorder in their structure.

• Crystalline- have a regular arrangement of components in their structure.

Crystals• Lattice- a three dimensional grid

that describes the locations of the pieces in a crystalline solid.

• Unit Cell-The smallest repeating unit in of the lattice.

• Three common shapes:– Cubic– Body Centered Cubic– Face Centered Cubic

Cubic

Body-Centered Cubic

Face-Centered Cubic

SolidsAmorphous Solids:

• Components “frozen in place” and lacking orderly arrangement.

Ex: Glass and plastic

We’ll focus more crystalline solids…

Crystalline Solids:

• Ionic Solids have ions at lattice points (Salt)

• Molecular Solids have molecules (Sugar)

• Metallic Solids (Cu, Fe, Al, Pt…)

• Covalent Network Solids (Diamonds, Quartz)

The book drones on about

• Using diffraction patterns to identify crystal structures.

• Talks about metals and the closest packing model.

• It is interesting, but trivial.• We need to focus on metallic bonding.• Why do metal atoms stay together.• How their bonding effects their

properties.

Metallic bonding

1s

2s2p

3s

3pFilled Molecular Orbitals

Empty Molecular Orbitals

Magnesium Atoms

Filled Molecular OrbitalsEmpty Molecular Orbitals

The 1s, 2s, and 2p electrons are close to nucleus, so they are not able to move around.

1s

2s2p

3s

3p

Magnesium Atoms

Filled Molecular OrbitalsEmpty Molecular Orbitals

1s

2s2p

3s

3p

Magnesium Atoms

The 3s and 3p orbitals overlap and form molecular orbitals.

Filled Molecular OrbitalsEmpty Molecular Orbitals

1s

2s2p

3s

3p

Magnesium Atoms

Electrons in these energy level can travel freely throughout the crystal.

Filled Molecular OrbitalsEmpty Molecular Orbitals

1s

2s2p

3s

3p

Magnesium Atoms

This makes metals: -conductors -Malleable because the bonds are flexible.

Carbon- A Special Atomic Solid

• There are three types of solid carbon.

• Amorphous- coal uninteresting.

• Diamond- hardest natural substance on earth, insulates both heat and electricity.

• Graphite- slippery, conducts electricity.

• How the atoms in these network solids are connected explains why.

Diamond- each Carbon is sp3hybridized, connected to four other carbons.• Carbon atoms are

locked into tetrahedral shape.

• Strong bonds give the huge molecule its hardness.

Why is it an insulator?

Empty MOs

Filled MOs

EThe space between orbitals make it impossible for electrons to move around

• Each carbon is connected to three other carbons and sp2 hybridized.

• The molecule is flat with 120º angles in fused 6 member rings.

• The bonds extend above and below the plane.

Graphite is different.

This bond overlap forms a huge bonding network.• Electrons are free to move through out

these delocalized orbitals.

• The layers slide by each other.

Molecular solids.

• Molecules occupy the corners of the lattices.

• Different molecules have different forces between them.

• These forces depend on the size of the molecule.

• They also depend on the strength and nature of dipole moments.

Those without dipoles or H bonding.

• Most are gases at 25ºC.

• The only forces are London Dispersion Forces.

• These depend on size of atom.

• Large molecules (such as I2 ) can be solids even without dipoles.

• Example: the Halogens

The Halogens

F

Cl

Br

I

Boiling points

85 K

239 K

333 K

458 K

State at 20 oC

Gas

Gas

Liquid

Solid

Those with dipoles.• Dipole-dipole forces are generally

stronger than L.D.F.• Hydrogen bonding is stronger than

Dipole-dipole forces.• No matter how strong the intermolecular

force, it is always much, much weaker than the forces in bonds.

• Stronger forces lead to higher melting and freezing points.

Water is special

• Each molecule has two polar O-H bonds.

• Each molecule has two lone pair on its oxygen.

• Each oxygen can interact with 4 hydrogen atoms.

HO

H

Water is special

HO

H

HO

H

HO

H

• This gives water an especially high melting and boiling point.

Ionic Solids• The extremes in dipole dipole forces-

atoms are actually held together by opposite charges = IONIC BONDS

• Huge melting and boiling points.• Atoms are locked in lattice so hard and

brittle.• Every electron is accounted for so they

are poor conductors-good insulators.

Vapor Pressure• Vaporization - change from

liquid to gas at boiling point.

• Evaporation - change from liquid to gas below boiling point

• Heat of Vaporization (Hvap ) -

the energy required to

vaporize 1 mole at 1 atm.

• Vaporization is an endothermic process - it requires heat.

• Energy is required to overcome intermolecular forces.

• Responsible for cool earth.

• Why we sweat.

Condensation• Change from gas to liquid.• Achieves a dynamic equilibrium with

vaporization in a closed system.• What is a closed system?• A closed system means matter

can’t go in or out. • What is a “dynamic

equilibrium?”

Dynamic Equilibrium• When first sealed the molecules gradually

escape the surface of the liquid• As the molecules build up above the liquid

some condense back to a liquid.• As time goes by the rate of

vaporization remains constant but the rate of condensation increases because there are more molecules to condense.• Equilibrium is reached when rate of condensation equals rate of vaporization

Rate of Vaporization = Rate of Condensation

• Molecules are constantly changing phase “Dynamic”

• The total amount of liquid and vapor remains constant “Equilibrium”

Dynamic equilibrium

Vapor pressure• The pressure above the liquid at

equilibrium.

• Liquids with high vapor pressures evaporate easily. They are called volatile.

• Decreases with increasing intermolecular forces. – Bigger molecules (bigger LDF)– More polar molecules (dipole-dipole)

Vapor pressure• Increases with increasing

temperature.

• Easily measured in a barometer.

Dish of Hg

Patm=

760 torr

A barometer will hold a column of mercury 760 mm high at one atm.

If we inject a volatile liquid in the barometer it will rise to the top of the mercury.

There it will vaporize and push the column of mercury down.

Water

Dish of Hg

736 mm Hg

Water Vapor

• The mercury is pushed down by the vapor pressure.

• Patm = PHg + Pvap

• Patm - PHg = Pvap

• 760 - 736 = 24 torr

Temperature Effect

Kinetic energy

# of

mol

ecu

les

T1

Energy needed to overcome intermolecular forces

Kinetic energy

# of

mol

ecu

les

T1

Energy needed to overcome intermolecular forcesT1

T2

• At higher temperature more molecules have enough energy - higher vapor pressure.

Energy needed to overcome intermolecular forces

Changes of state• The graph of temperature versus

heat applied is called a heating curve.

• Heat of Fusion represents the transition from Solid to Liquid

• Heat of Vaporization represents the transition from Liquid to Gas

-40

-20

0

20

40

60

80

100

120

140

0 40 120 220 760 800

Heating Curve for Water (axis not to scale)

IceWater and Ice

Water

Water and Steam

Steam

-40

-20

0

20

40

60

80

100

120

140

0 40 120 220 760 800

Heating Curve for Water (axis not to scale)

Heat of Fusion

Heat ofVaporization

Slope is Heat Capacity

• Copy picture of Heating Curve for Water

1. In which region(s) of the graph is temperature constant?

2. Describe how the amount of energy required to melt a given mass of ice compares to the energy required to vaporize the same mass of water? Explain.

3. Which region of the graph represents the coexistence of solid and liquid? Liquid and Vapor?

HEAT IN CHANGES OF STATE

4. When a material condenses or freezes is energy released or absorbed? Explain your reasoning.

5. When a material evaporates or melts, is energy released or absorbed? Explain your reasoning.

Energy (as heat) during Changes of State

• Compounds absorb energy (heat) when breaking bonds.

• Solids absorb energy to melt into liquids. (ice cubes melting)

• Liquids absorb energy to evaporate into gases.

• Energy goes to changing state (breaking bonds), there is no temperature change!

• Compounds release energy when forming bonds.

• Gases release energy to condense into a liquid.

• Liquids release energy to freeze into a solid.

• Energy goes to changing state (making bonds), there is no temperature change!

Equations for Heating Curve

• No Phase Change:

Q = m(▲T)Cp

• Latent Heat of Fusion: (Solid Liquid)

Q = m ▲Hfus

• Latent Heat of Vaporization: (Liquid Gas)

Q= m ▲Hvap

Melting Point

• Melting point is determined by the vapor pressure of the solid and the liquid.

• At the melting point…• the vapor pressure of the solid =

vapor pressure of the liquid

Solid Water

Liquid Water

Water Vapor Vapor

Solid Water

Liquid Water

Water Vapor Vapor

• If the vapor pressure of the solid is higher than that of the liquid the solid will release molecules to achieve equilibrium.

Solid Water

Liquid Water

Water Vapor Vapor

• While the molecules of water condense to a liquid.

• This can only happen if the temperature is above the freezing point since solid is turning to liquid.

Solid Water

Liquid Water

Water Vapor Vapor

• If the vapor pressure of the liquid is higher than that of the solid, the liquid will release molecules to achieve equilibrium.

Solid Water

Liquid Water

Water Vapor Vapor

Solid Water

Liquid Water

Water Vapor Vapor

• While the molecules condense to a solid.

• The temperature must be below the freezing point since the liquid is turning to a solid.

Solid Water

Liquid Water

Water Vapor Vapor

• If the vapor pressure of the solid and liquid are equal, the solid and liquid are vaporizing and condensing at the same rate. The Melting point.

Solid Water

Liquid Water

Water Vapor Vapor

Boiling Point• Reached when the vapor pressure

equals the external pressure.

• Normal boiling point is the boiling point at 1 atm pressure.

• Super heating - Heating above the boiling point.

• Supercooling - Cooling below the freezing point.

Using Vapor Pressure Data

• Clausius-Clapeyron equation is used to determine the vapor pressure of a liquid at a different temp. or to find the heat of vaporization of a substance.

• ln(Pvap,T1/PvapT2) = (Hvap/R)[(1/T2) – (1/T1)]

• The graph of ln P vs 1/T is linear with a negative slope.

• Determine the boiling point of water in Breckenridge, Colorado, where the atmospheric pressure is 520 torr. For water, Hvap = 40.7 kJ/mol.

• P1 = 520 Torr; P2 = 760 Torr

• R = 8.3145 J/K x mol

• T2 = 373 K; T1 = ?

• T1 = 362 K or 89 oC

Phase Diagrams.• A plot of temperature versus

pressure for a closed system, with lines to indicate where there is a phase change.

Temperature

SolidLiquid

Gas

1 Atm

AA

BB

CCD

D D

Pre

ssur

e

D

Dissecting our Phase Diagram• A = Sublimation / Deposition• B = Vaporization / Condensation• C = Melting / Freezing• D = Liquid moving up past the critical

point and transforming into a gas

Phase Diagram terminology• Triple Point – all three phases exist in

equilibrium in a “closed system”• Critical Temperature – temp. above which

the substance cannot exist as a liquid regardless of how great the pressure.

• Critical Pressure – pressure required to produce liquefaction at the critical temperature.

• Critical Point – Point defined by the critical temp. and pressure (For Water: 374oC and 218 atm)

SolidLiquid

Gas

Triple Point

Critical Point

Temperature

Pre

ssur

e

SolidLiquid

Gas

• This is the phase diagram for water.

• The density of liquid water is higher than solid water.

Temperature

Pre

ssur

e

Solid Liquid

Gas

1 Atm

• This is the phase diagram for CO2

• The solid is more dense than the liquid

• The solid sublimes at 1 atm.

Temperature

Pre

ssur

e

• How does the melting/freezing line differ for the phase diagram of Water and Carbon Dioxide?

• What may cause that difference?