chapter 11_jan14 new version (shortened)

8/12/2019 Chapter 11_Jan14 New Version (Shortened)

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Chapter 11

I ntermolecular F orces,L iquids, and Solids


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CONTENTS11.1 Introduction: A Molecular Comparison

11.2 Types of Intermolecular Forces : 11.2.1 Ion-dipole forces

11.2.2 Dipole-dipole forces 11.2.3 Hydrogen bonding 11.2.4 London dispersion forces 11.2.5 Comparison between forces

11.3 Properties of Liquids 11.3.1 Viscosity and surface tension


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11.4 Phase Change

11.4.1 Energy changes accompanying phasechange 11.4.2 Cooling curve 11.4.3 Critical temperature and pressure

11.5 Vapour pressure

11.6 Phase diagram 11.6.1 Phase diagram of CO 2 11.6.2 Phase diagram of water


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11.7 Structure of solids 11.7.1 Crystalline and amorphous solids 11.7.2 Unit cell - PC, FCC & BCC structure

11.8 Bonding in Solids 11.8.1 Molecular solids 11.8.2 Covalent-network solids 11.8.3 Ionic and metallic solids


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Learning Outcomes

Able to differentiate the 4 main intermolecularforces (IMF)

Able to relate IMF of a compound to the boiling point, surface tension, viscosity and vapor pressureAble to use phase diagram in problem solvingAble to use knowledge on unit cell of an elementto calculate molar mass, density, radius etc.


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11.1 I ntroduction: A M olecular ComparisonSolid - L iqui d - Gas


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Solid - L iquid - GasState Gas Liquid Solid

Volume Assumes vol. &shape ofcontainer

Assumes shapeof the portion of

container but

not volume

Retain its ownshape &volume

Compressibility

Compressible Virtuallyincompressible

Virtuallyincompressible

Diffusion Occurs rapidly Occurs slowly Extremely low

Flow Readily Readily Does not flow


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11.2 I ntermolecular F orces

The attraction between molecules is anintermolecular force.Intermolecular forces are much weaker than

ionic or covalent bonds.When a substance melts or boils, intermolecularforces are broken.When a substance condenses, intermolecular

forces are formed.Boiling points/melting points indicate the strengthof intermolecular forces.


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11.2 Types of I ntermolecular F orces

Ion-dipole force

Dipole-dipole force

Hydrogen bonding

London-Dispersion force

+ +- -

+-+

A - H :B -


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11.2.1 I on-Dipole F orces

Exists between ions and the partial charge onthe end of polar molecules/dipoleImportant for solutions of ionic compounds in

polar liquidsE.g.: NaCl in waterMagnitude of attraction increases:

charge of the ion increases magnitude of dipole moment increase

Ionic bonding > Ion-dipole > Dipole-dipole


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I on-Dipole F orces

+ -

+

+

-

--

+

-

-+

+

+

+

+

-

--


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11.2.2 Dipole-dipole F orces

Exists between neutral polar moleculesThe partially positive end of one molecule attracts thepartially negative end of another.

Weaker than ion-dipole forces

E.g.: CO---CO

C

O

C C

C

C O

OO

O


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Dipole-dipole F orces

+

- Attractionsare greater

thanrepulsion, sothe

moleculesfeel a netattraction toeach other

-+

+

+

+

+

+

-

-

- -

-

-+


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For molecules of approximately equal mass and size , strength of attraction increases

with increasing polarityMass

(amu)Dipole

moment ( )Boiling point

(K)CH3CH2CH3 44 0.1 231

CH3OCH 3 46 1.3 248CH3CHO 44 3.9 294


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For molecules of comparable polarity ,those with smaller molecular volumes generally experience higher dipole-dipole attractive forces.


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11.2.3 H ydrogen Bonding

Special type of intermolecular attraction.Hydrogen bonding is a special case of dipole-dipole interactions.H-bonding requires:

H bonded to a small electronegative element (mostimportant for compounds of F, O, and N)

An unshared electron pair on a nearby smallelectronegative ion or atoms (usually F, O, or N onanother molecule)


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11.2.3 H ydrogen Bonding

E.g.:

Ion-dipole > Hydrogen bonding > Dipole-dipole

OH

H

OH

H

NH

H

OH

H

H


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H ydrogen Bonding

Hydrogenbond


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H ydrogen Bonding

Hydrogen bonding is important in :

Stabilising the structure of protein

Folding of protein moleculesSurvival of aquatic in frozen lake

Ice is less dense than water Ice floats so forms an insulating layer on top of

lakes or river. Therefore, aquatic life cansurvive in winter.


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11.2.4 L ondon Dispersion F orces

What intermolecular forces exist betweennonpolar molecules in liquid and solid state ? Dipole-dipole attractions cannot exist betweennon-polar atoms or molecules Non-polar molecules do not have permanentdipoles but all nonpolar substances can be

liquefied Therefore, there must exist some kind ofattractive interactions between the particles


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Two schematic presentations of theinstantaneous dipoles on two adjacent heliumatoms, showing the electrostatic attractionbetween them

L ondon Dispersion F orces


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In a collection of helium atoms, the averagedistribution of electrons about each nucleus isspherically symmetrical

The atoms are nonpolar and posses nopermanent dipole moment


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In any atom or molecule, electrons constantlymovingThe nucleus of one molecule (or atom) attractsthe electrons of the adjacent molecule (oratom).For an instant, the electron clouds become

distorted.In that instant a dipole id formed (called ani n s t an taneous d ipo le )


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One instantaneous dipole can induce anotherinstantaneous dipole in ad adjacent molecule(or atom).

These two temporary dipoles attract eachother.The attraction is called London-dispersionforce (LDF) , or simply a dispersion force.LDF is the weakest of all intermolecular forces.LDF exist between all molecules (polar ornonpolar)


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L ondon Dispersion F orces (L DF )

What affects the strength of a dispersion force? Molecules must be very close together for

these attractive forces to occur. POLARIZABILITY is the ease with which an

electron distribution can be deformed. The larger the molecule, the more polarize it is. LDF increase as molecular weight increases. LDF depend on the shape of the molecule.


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Factors affecting London dispersion force :1. Molecular weight

increase in mol. wt. (increase in atomic radii)results in increase number of electrons the larger the molecule, the farther its electron

from the nuclei, the greater its polarizability greater polarizability, the greater the strength of

London dispersion force


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E.g. : Noble gases . As go down theperiodic table, mol. wt. increase

(atomic radii increase) , resultinggreater boiling point

Substance Mol. Wt (amu) Boiling point (K)He 4.0 4.6

Ne 20.2 27.3 Ar 39.9 87.5Kr 83.8 120.9Xe 131.3 166.1


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Factors affecting London dispersion force2. Shape of molecule

greater the surface area available

for contact, greater the dispersionforce

e.g.: straight chain molecule

>branched chain molecule

Neopentane

(bp=282.7K)

N-pentane

(bp=309.4K)


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11.2.5 Comparison between F orces

LDF are found in all substances. Their strength depends on molecular shapes

and molecular weights. Dipole-dipole forces add to the effect of LDF.

They are found only in polar molecules. Ion-dipole interactions are stronger than H-

bonding.


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H-bonding is a special case of dipole-dipole interactions.

It is the strongest of intermolecular forcesinvolving neutral species.

H-bonding is most important for H

compounds of N, O, and F.


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If ions are involved, ion-dipole (if a dipoleis present) and ionic bonding are

possible. Ion-dipole interactions are stronger than H-bonding.

Ordinary ionic or covalent bonds aremuch stronger than these interactions!



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11.3 Proper ties of L iquids

11.3.1 VISCOSITY Viscosity - resistance of liquid to flow. Liquid flows by sliding molecules over

each other. The stronger the intermolecular

forces, the higher the viscosity. Viscosity usually decreases with an

increase in temperature.


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11.3.2 SURFACE TENSION bulk molecules (those in liquid) are equally

attracted to their neighbours

surface molecules are only attractedinwards to the bulk molecules therefore surface molecules are packed

more closely than bulk molecules

Surface tension - amount of energy required toincrease the surface area of a liquid by a unitamount.

Proper ties of L iquids


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Explaining Surface Tension

Copyright Houghton Mifflin Company. All rights reserved 11-16

11_16

Surface


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SURFACE TENSION (CONT..)

Strong intermolecular forces cause higher

surface tension. Water has a high surface tension (H-

bonding)

Hg(l) has even higher surface tension (thereare very strong metallic bonds between Hgatoms)



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SURFACE TENSION (CONT..)

Cohesive forces Intermolecular forces that bind molecules to

one another Adhesive forces

Intermolecular forces that bind molecules toa surface



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To illustrate this:Meniscus - the shape of the liquid surface

if adhesive forces > than cohesive forces - the liquid is attracted to its container, meniscus is U-shaped (e.g. water in glass)

if cohesive forces > than adhesive forces - meniscus curved downwards (e.g. Hg in glass)

Capillary action the rise of liquids up very narrowtubes.The liquid climbs until adhesive and cohesive forcesare balanced by gravity.


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Liquid Levels in Capillaries: Capillary Rise

Copyright Houghton Mifflin Company. All rights reserved 11-19A

11_19a


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11.4 Phase Changes

A phase change occurs when the matter istransformed from one physical state to another

During phase change temperature remains constant Sublimation : Solid Gas Vaporization : Liquid Gas Melting or Fusion : Solid Liquid Deposition : Gas Solid Condensation : Gas Liquid Freezing : Liquid Solid


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During phase change, the energy change of thesystem are as given below

Sublimation DHSUB > 0 (Endothermic) Vaporisation DHVAP > 0 (Endothermic) Melting OR Fusion DHFUS > 0 (Endothermic)

Deposition DHDEP < 0 (Exothermic) Condensation DHCON < 0 (Exothermic) Freezing DHFREEZ < 0 (Exothermic)

11.4.1 Energy changes accompanying phase changes

Phase Changes


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The amount of energy required to cause a phasechange increases as the strength of the intermolecularforces increases

Generally heat of fusion is less than heat ofvaporisation

Less energy is needed to allow particles to movepast one another than to separate them totally

Steam can cause severe burns, when it comes in

contact with skin because condensation (exothermic)

11.4.1 Energy changesaccompanying phase changes

11 4 2 Cooling curve


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+

11.4.2 Cooling curve

Cooling curve for the conversion of gaseous water to ice.


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+

11.4.2 Cooling curve

Supercooling:When a liquid iscooled below itsfreezing pointand it stillremains a liquid.


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11.4.3 Cri tical Temperatur e andPressure

Gas can be liquefied by compressing them at a suitable T As temperature increases, gas become more difficult to liquefy

because of the increasing kinetic energy

For every substance, there exist a temperature above which the

gas cannot be liquefied, regardless of the pressure

CRITICAL TEMPERATURE (T C) - the highest temperature atwhich a substance can exist as a liquid (using pressure)

CRITICAL PRESSURE (P C) - the pressure required to bringabout liquefaction at T C

The greater the intermolecular forces, the easier it is to liquefy asubstance, thus the higher the T C of the substance.


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11.5 Vapour Pressure

What is vapor pressure ? We place a quantity of ethanol ( b.p. 78.3C @ 1 atm)

in an evacuated container some of the molecules on the surface of the liquid have

enough energy to escape the attraction of the bulkliquid

these molecules moves into the gaseous phase as the number of molecules in the gas phase

increases, some of the gas phase molecules struck thesurface and returns to the liquid


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Measurement of the Vapour Pressure of Water



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It will reach a condition where the rate at which moleculesreturn to the liquid is equal to the rate they escape from theliquid.

At this condition ; the number of molecules in the gas phase reaches a

steady state pressure of the vapor remains constant dynamic equilibrium is attained between the rate of

condensation and rate of evaporation pressure of the vapor remains constant (vapor pressure)

Vapour Pressure


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Vapour pressure of a liquid isdefined as the pressure exerted by its

vapor when the liquid and the vaporstate are in dynamic equilibrium(Dynamic equilibrium is a condition inwhich two opposing processes occursimultaneously at equal rates .)

Vapour Pressure


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Liquids boil when the external pressure at theliquid surface equals the vapour pressure.

The normal boiling point is the boiling point at 760mmHg (1 atm).The temperature of the boiling point increases asthe external pressure increases.Two ways to get a liquid to boil:

Increase temperature or decrease pressure.

Vapour Pressure and BoilingPoint


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Under appropriate condition, equilibrium can existbetween:

liquid - gas solid - liquid solid - gas

We can summarise these conditions as phase

diagram

Phase diagram - a plot of pressure vs. temperaturesummarizing all equilibria between phases.

11.6 Phase Diagram


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(Explanation for Figure ) A-B = vapor pressure curve, represent equilibriumbetween liquid and gas phaseVapor pressure curve ends at the critical point

A = triple point (temp. & pressure at which all threephases are in equilibrium

A-D = change in melting point with increasing pressureThe solid is denser than liquid. An increase inpressure will result in more compact solid phase i.e.higher temperature are required to melt a solid in highpressure

Phase Diagram


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11.6.1 Phase Diagrams for CarbonDioxide

Notice that the solid-liquid equilibrium(melting point) line ofCO 2 follows the normal

behaviour i.e. its meltingpoint increases withincreasing temperature.

For solid CO 2 to exist asliquid the pressure mustexceed 5.11 atm

CO 2 does not melt, butsublimes at 1 atm


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11.6.2 Phase Diagram for Water

The melting point ofwater decreases with

increasing pressure Liquid form is more

compact than its solidform


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The melting point and hardness depends onarrangement of particlesattraction forces between them

We can classify solids according to types offorces (4 types)

11.8 Bonding in Solids


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Bonding in Solids

Molecular (formed from atoms or molecules) usually soft with low melting point and conductivity

Covalent network (formed from atoms) very hard with very high melting point and poor conductor

Ionic (formed from ions) hard, brittle, high melting point and poor conductivity

Metallic (formed from metal atoms) soft or hard with high melting point, good conductivity

and ductile)


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Intermolecular forces

Dipole-dipole, London dispersion and H-bonds.

Weak intermolecular forces give rise to lowmelting point.Efficient packing of molecules is important(since they are not regular spheres)

E.g. : Benzene (m.p 5C & b.p 80C)Toluene (m.p -95C & b.p 111C)

11.8.1 Molecular Solids


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E.g. : Benzene VS Toluene

Benzene

m.p 5C, b.p 80CSymmetrical planarmolecule packedefficiently in solidform - high meltingpointIMF in liquid statelower than toluene

Toluene

m.p -95C, b.p 111CLower symmetry,prevents from packingefficientlyIntermolecular forcesdepend on closecontact - not effectiveIMF in liquid state largerthan benzene

Molecular Solids


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Atoms are held together in large networks or chain bycovalent bondsCovalent bonds are stronger than intermolecular forcesE.g. Diamond & Graphite

Diamond Each carbon atom is bonded to 4 other carbon atom

Graphite The carbon atoms are arranged in layers of inter-

connected hexagonal rings The layers are held by dispersion forces Readily slide past one another

11.8.2 Covalent-NetworkSolids


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Structures of Diamond and Graphite



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Ions are held together by ionic bondsThe strength of ionic solids depends on the charges of theions

NaCl ---> Na + Cl - ---> has m.p. of 801C MgO ----> Mg 2+ O 2- ----> has m.p. of 2852C

The structure adopted by ionic solids depends largely onthe charges and relative sizes of the ions

E.g.: NaCl (Na + has CN of 6 and surrounded by 6Cl) whereas CsCl (Cs is surrounded by 8 Cl - atoms)

increase in CN results in increase in structure

11.8.3 Ionic Solids


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M etall ic Solids

Hardness and melting point depends on strength ofbonding

Bonding strength increases with increase inbonding electrons

e.g. Na (one valence electron- m.pt= 97.5C) Cr (6 valence electrons - m.pt = 1890c)

Mobility of electron explains why electrons aregood conductor of heat and electricity


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END of CH PTER

chapter 11_jan14 new version (shortened)

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