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Chapter 11“Liquids, Solids,and IntermolecularForces”

Comparisons of theStates of Matter

• The solid and liquid states have a muchhigher density than the gas state

• The solid and liquid states have similardensities– generally the solid state is a little denser

• The molecules in the solid and liquid state arein close contact with each other, while themolecules in a gas are far apart

Freedom of Motion• The molecules in a gas have complete

freedom of motion• The molecules in a solid are locked in place,

they cannot move around– though they do vibrate

• The molecules in a liquid have limitedfreedom – they can move around a littlewithin the structure of the liquid

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Kinetic - Molecular Theory

• The properties of solids, liquids, and gasescan be explained based on the kinetic energyof the molecules and the attractive forcesbetween molecules

• Kinetic energy tries to give moleculesfreedom of motion– degrees of freedom = translational, rotational,

vibrational• Attractive forces try to keep the molecules

together, kinetic energy depends only on thetemperature

Gas Structure

Gas molecules arerapidly moving inrandom straight linesand free from stickingto each other.

Explaining the Properties ofSolids

• The particles in a solid are packed closetogether and are fixed in position

• The close packing of the particles results insolids being incompressible

• The inability of the particles to move aroundresults in solids retaining their shape andvolume when placed in a new container; andprevents the particles from flowing

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Explaining the Properties ofLiquids

• They have higher densities than gasesbecause the molecules are in close contact

• They have an indefinite shape because thelimited freedom of the molecules allows themto move around enough to get to thecontainer walls

• They have a definite volume because the limiton their freedom keeps them from escapingthe rest of the molecules

Compressibility

Phase Changes

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Why are molecules attractedto each other?

• Intermolecular attractions are due to attractiveforces between opposite charges– + ion to - ion– + end of polar molecule to - end of polar molecule

• H-bonding especially strong

– even nonpolar molecules will have temporarycharges

• larger the charge = stronger attraction• longer the distance = weaker attraction

Trends in the Strength ofIntermolecular Attraction?

• The stronger the attractions between theatoms or molecules, the more energy it willtake to separate them

• Boiling a liquid requires we add enoughenergy to overcome the attractions betweenthe molecules or atoms

• The higher the normal boiling point of theliquid, the stronger the intermolecularattractive forces

Dispersion Forces

• Fluctuations in the electron distribution inatoms and molecules result in a temporarydipole

• The attractive forces caused by thesetemporary dipoles are called dispersionforces (aka London Forces)

• All molecules and atoms will have them• As a temporary dipole is established in one

molecule, it induces a dipole in all thesurrounding molecules

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Dispersion Force

Size of the Induced Dipole• The magnitude of the induced dipole depends on

several factors• Polarizability of the electrons

– volume of the electron cloud– larger molar mass = more electrons = larger

electron cloud = increased polarizability = strongerattractions

• Shape of the molecule– more surface-to-surface contact = larger induced

dipole = stronger attraction

Boiling Points of n-Alkanes

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Dipole-Dipole Attractions

• Polar molecules have a permanent dipole– because of bond polarity and shape– dipole moment

• The permanent dipole adds to the attractiveforces between the molecules– raising the boiling and melting points relative to

nonpolar molecules of similar size and shape

Effect of Dipole Moment onBoiling Points

Attractive Forces andSolubility

• Solubility depends on the attractive forcesof solute and solvent molecules– Like dissolves Like

• Polar substance dissolve in polar solvents• Nonpolar molecules dissolve in nonpolar

solvents• Many molecules have both hydrophilic and

hydrophobic parts - solubility becomescompetition between parts

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Immiscible Liquids

Hydrogen Bonding

• When a very electronegative atom is bondedto hydrogen, it strongly pulls the bondingelectrons toward it (O-H, N-H, or F-H)

• Since hydrogen has no other electrons, whenit loses the electrons, the nucleus becomesdeshielded (exposing the H proton)

• The exposed proton acts as a very strongcenter of positive charge, attracting all theelectron clouds from neighboring molecules

H-Bonding

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H-Bonding

Practice

Ion-Dipole Attraction

• In a solution, ions from an ioniccompound are attracted to the dipole ofpolar molecules

• The strength of the ion-dipole attractionis one of the main factors thatdetermines the solubility of ioniccompounds in water

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Ion-Dipole Attraction

Liquid Properties & Structures• Surface tension is a property of liquids that results

from the tendency of liquids to minimize their surfacearea

• In order to minimize their surface area, liquids formdrops that are spherical

• The layer of molecules on the surface behavedifferently than the interior– because the cohesive forces on the surface molecules have

a net pull into the liquid interior• The surface layer acts like an elastic skin

Liquid Properties & Structures

• The stronger the intermolecularattractive forces, the higher the surfacetension will be

• Raising the temperature of a liquidreduces its surface tension

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Liquid Properties & Structures

• Viscosity is the resistance of a liquid toflow

• Larger intermolecular attractions =larger viscosity

• Higher temperature = lower viscosity

Liquid Properties & Structures

• Capillary action is the ability of a liquid toflow up a thin tube against the influence ofgravity– the narrower the tube, the higher the liquid rises

• Capillary action is the result of the two forcesworking in conjunction, the cohesive andadhesive forces– cohesive forces attract the molecules together– adhesive forces attract the molecules on the edge

to the tube’s surface

Liquid Properties & Structures• The Meniscus is the curving of the

liquid surface in a thin tube is due tothe competition between adhesive andcohesive forces

• The meniscus of water is concave in aglass tube because its adhesion to theglass is stronger than its cohesion foritself

• The meniscus of mercury is convex ina glass tube because its cohesion foritself is stronger than its adhesion forthe glass

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Liquid Properties & Structures• Molecules in the liquid are constantly in motion the

average kinetic energy is proportional to thetemperature

• Some molecules have more kinetic energy than theaverage if these molecules are at the surface, theymay have enough energy to overcome the attractiveforces– therefore – the larger the surface area, the faster the

rate of evaporation• This will allow them to escape the liquid and

become a vapor

Liquid Properties & Structures• Some molecules of the vapor will lose

energy through molecular collisions• The result will be that some of the

molecules will get captured back intothe liquid when they collide with it

• Also some may stick and gathertogether to form droplets of liquid

• This process is called condensation

Effect of Intermolecular Attractionon Evaporation and Condensation

• The less attractive forces between molecules,the less energy they will need to vaporize

• Liquids that evaporate easily are said to bevolatile

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Energetics of Vaporization

• When the high energy molecules are lostfrom the liquid, it lowers the average kineticenergy

• If energy is not drawn back into the liquid, itstemperature will decrease – therefore,vaporization is an endothermic process andcondensation is an exothermic process

Heat of Vaporization• The amount of heat energy required to

vaporize one mole of the liquid is called theHeat of Vaporization, ΔHvap

• Always endothermic, therefore ΔHvap is +• Somewhat temperature dependent• ΔHcondensation = -ΔHvaporization

Example

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Dynamic Equilibrium

• In a closed container, once the rates ofvaporization and condensation are equal, thetotal amount of vapor and liquid will notchange

• evaporation and condensation are stilloccurring, but there is no net gain or loss oreither vapor or liquid

• when two opposite processes reach the samerate so that there is no gain or loss ofmaterial, we call it a dynamic equilibrium

Dynamic Equilibrium

Vapor Pressure

• The pressure exerted by the vaporwhen it is in dynamic equilibrium with itsliquid is called the vapor pressure

• The weaker the attractive forcesbetween the molecules, the moremolecules will be in the vapor– the weaker the attractive forces, the higher

the vapor pressure and the higher thevapor pressure, the more volatile the liquid

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Vapor-Liquid Dynamic Equilibrium• If the volume of the chamber is increased, that will

decrease the pressure of the vapor inside• Eventually enough liquid evaporates so that the rates

of the condensation increases to the point where it isonce again as fast as evaporation– equilibrium is reestablished

• At this point, the vapor pressure will be the same as itwas before

Dynamic Equilibrium

• A system in dynamic equilibrium canrespond to changes in the conditions

• When conditions change, the systemshifts its position to relieve or reducethe effects of the change

Vapor Pressure vs.Temperature

• Increasing the temperature increasesthe number of molecules able to escapethe liquid

• The net result is that as the temperatureincreases, the vapor pressure increases

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Vapor Pressure Curves

Temperature vs Vapor Pressure

0

100

200

300

400

500

600

700

800

900

1000

0 10 20 30 40 50 60 70 80 90 100 110 120 130 140 150Temperature, °C

Vap

or

Pre

ssu

re, m

mH

g

w ater

TiCl4

chloroform

ether

ethanol

acetone

Boiling Point• When the temperature of a

liquid reaches a point where itsvapor pressure is the same asthe external pressure, vaporbubbles can form anywhere inthe liquid

• This phenomenon is what iscalled boiling and thetemperature required to havethe vapor pressure = externalpressure is the boiling point

Heating Curve of a Liquid• As you heat a liquid, its

temperature increaseslinearly until it reaches theboiling point

• Once the temperaturereaches the boiling point, allthe added heat goes intoboiling the liquid – thetemperature stays constant

• Once all the liquid has beenturned into gas, thetemperature can again startto rise

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Clausius-Clapeyron Equation2-Point Form

• The equation below can be used with just twomeasurements of vapor pressure andtemperature (remember: the vapor pressureat the normal boiling point is 760 torr)

!!"

#$$%

&'

('=!!

"

#$$%

&

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vap

1

2

T

1

T

1

R

H

P

Pln

Example

Supercritical Fluid

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The Critical Point

• The temperature required to produce asupercritical fluid is called the criticaltemperature

• The pressure at the critical temperature iscalled the critical pressure

• At the critical temperature or highertemperatures, the gas cannot be condensedto a liquid, no matter how high the pressuregets

Sublimation and Deposition• Molecules in the solid have thermal energy

that allows them to vibrate• Surface molecules with sufficient energy may

break free from the surface and become agas – this process is called sublimation

• The capturing of vapor molecules into a solidis called deposition

• The solid and vapor phases exist in dynamicequilibrium in a closed container

Sublimation

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Melting = Fusion

• As a solid is heated, its temperature rises andthe molecules vibrate more vigorously

• Once the temperature reaches the meltingpoint, the molecules have sufficient energy toovercome some of the attractions that holdthem in position and the solid melts (or fuses)

• The opposite of melting is freezing

Heating Curve of a Solid

• As you heat a solid, itstemperature increaseslinearly until it reaches themelting point

• Once the temperaturereaches the melting point, allthe added heat goes intomelting the solid – thetemperature stays constant

• Once all the solid has beenturned into liquid, thetemperature can again startto rise

Energetics of Melting

• When the high energy molecules are lostfrom the solid, it lowers the average kineticenergy

• If energy is not drawn back into the solid itstemperature will decrease – therefore, meltingis an endothermic process and freezing is anexothermic process

• Melting requires input of energy to overcomethe attractions between molecules

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Heat of Fusion

• The amount of heat energy required to meltone mole of the solid is called the Heat ofFusion, ΔHfus

• Always endothermic, therefore ΔHfus is +• Somewhat temperature dependent• ΔHcrystallization = -ΔHfusion

• Generally much less than ΔHvap

• ΔHsublimation = ΔHfusion + ΔHvaporization

Heats of Fusion andVaporization

Heating Curve of Water

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Calculation:

Phase Diagrams• Describe the different states and state changes that

occur at various temperature - pressure conditions• Areas represent states• Lines represent state changes

– liquid/gas line is vapor pressure curve– both states exist simultaneously– critical point is the furthest point on the vapor pressure curve

• Triple point is the temperature/pressure conditionwhere all three states exist simultaneously

• For most substances, freezing point increases aspressure increases

Phase Diagrams

Pres

sure

Temperature

vaporization

condensation

criticalpoint

triplepoint

Solid Liquid

Gas

1 atmnormal

melting pt.normal

boiling pt.

Fusion Curve

Vapor PressureCurve

SublimationCurve

melting

freezing

sublimation

deposition

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Phase Diagrams

Phase Diagram of Water

Temperature

Pres

sure

criticalpoint

374.1°C217.7 atm

triplepoint

Ice Water

Steam

1 atm

normalboiling pt.

100°C

normalmelting pt.

0°C

0.01°C0.006 atm

Phase Diagram of CO2

Pres

sure

Temperature

criticalpoint

31.0°C72.9 atm

triplepoint

Solid Liquid

Gas1 atm

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Phase Diagram of CO2

Water – An ExtraordinarySubstance

• water is a liquid at room temperature

• water is an excellent solvent – dissolving many ionicand polar molecular substances– even many small nonpolar molecules have solubility in water

• water has a very high specific heat for a molecularsubstance

• water expands when it freezes– at a pressure of 1 atm– about 9%– making ice less dense than liquid water

Note

• Omit– Sections 11.10, 11.11, 11.12 and 11.13

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