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Liquids, solids, & Liquids, solids, & intermolecular intermolecular

forcesforces

Chapter 11

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KMT meets liquidsKMT meets liquids

Ideal gas is a gas even at absolute zeroReal gas condenses to liquid at low T/high P

Attractive forces exist between real gas molecules

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Intermolecular attractionsIntermolecular attractions

Attractive forces exist between all atoms/molecules

Relative strength of attractions indicated byBoiling point (higher b.p. = stronger attractions)Vapor pressure (high v.p. = weaker attractions)∆Hvaporization (large ∆Hvap = stronger attractions)

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Instantaneous or momentary dipolesInstantaneous or momentary dipoles

e– distribution is asymmetric –– just for a momentAtom/molecule is polar –– just for a moment

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Induced dipolesInduced dipoles

Momentary dipole in one atom induces a dipole in a neighboring atom . . . which induces a dipole in another neighboring atom, and so on, causing a little ripple of dipoles

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Dispersion forceDispersion force

Taken together, instantaneous & induced dipoles create an attractive force between molecules, called the dispersion force

Each dipole is tiny, but the constant ripple of countless dipoles throughout the substance makes this the primary attractive force between molecules

Even noble gas atoms show dispersion force between atoms

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PolarizabilityPolarizability

Magnitude of dispersion force depends on polarizability

Larger e– cloud = more polarizable Dispersion force increases with increasing molar massMelting and boiling points of molecular substances

generally increase as molar mass increases

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Molar mass & boiling pointMolar mass & boiling point

For compounds of similar structure, boiling point increases as molar mass increases

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PolarizabilityPolarizability

Polarizability is greater in elongated molecules than in compact ones of similar mass

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Permanent dipolesPermanent dipoles

Polar molecules tend to arrange themselves +/– to maximize attractions

Extra ordering increases tendency to stick together in liquid state

Boiling point of a polar substance is higher than that of a nonpolar substance of similar mass.

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Nonpolar/polarNonpolar/polar

Molecules have similar masses

Permanent dipoles increase b.p.

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The van der Waals forcesThe van der Waals forces

Together, dispersion and pemanent dipole forces are known as the van der Waals forcesWhen comparing substances of comparable mass

(±10%), the presence of a permanent dipole increases boiling point significantly

When comparing substances of different molar masses, the dispersion force (related to mass) is more important than the permanent dipole

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ExamplesExamples

Which would you expect to have the highest boiling point, and why: C3H8, CO2, CH3CN

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ExamplesExamples

Which would you expect to have the highest boiling point, and why: C3H8, CO2, CH3CNmasses similar (C3H8 = 44, CO2 = 44, CH3CN = 41)

CH3CN polar = highest bp

Actual values: C3H8 = 231K, CO2 = 195K, CH3CN =

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ExamplesExamples

Arrange these in order of increasing boiling point: Ne, He, Cl2, (CH3)2CO, O2, O3

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ExamplesExamples

Arrange these in order of increasing boiling point: Ne, He, Cl2, (CH3)2CO, O2, O3

masses: Ne = 20, He = 4, Cl2 = 71, (CH3)2CO = 58, O2 = 32, O3 = 48

Ordered by mass: He, Ne, O2, O3, (CH3)2CO, Cl2

(CH3)2CO is polar & has large surface area = higher bp

Predict He, Ne, O2, O3, Cl2, (CH3)2CO

Actual values: He = 4K, Ne = 27K, O2 = 90K, O3 = 161K, Cl2 = 238K, (CH3)2CO = 329K

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Then there’s hydrogen . . .Then there’s hydrogen . . .

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Weak bond to neighboring O is a hydrogen bond

O–H bond is very polar, and atoms are very small

Dipoles are close together, so their attraction is very strong

H atom is covalently bonded to its own O and weakly bonded (dotted line) to the neighboring O

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Hydrogen bondingHydrogen bonding

Hydrogen bonding occurs only between molecules containing N–H, O–H, and F–H bonds

Hydrogen bonding is much stronger than ordinary dispersion/dipole → much higher boiling points than expected for their mass

Hydrogen bonds are not as strong as covalent bonds (15-40 kJ/mol, vs >150 kJ/mol)

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Intermolecular forcesIntermolecular forces

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Substances that are not molecularSubstances that are not molecularIonic substances

Held together by lattice energyGenerally high mp & bp

Metallic substancesMetal cations in sea of electronsGenerally high mp & bp

Network covalent solids (e.g. diamond)Melting = disrupt covalent bondsVERY high mp & bp

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VaporizationVaporization

At liquid surface, faster molecules have enough kinetic energy to escape (vaporize or evaporate)

As higher-energy molecules leave the liquid, average kinetic energy of the liquid decreases

Temperature of liquid decreases (evaporative cooling)

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VaporizationVaporization

For liquid temperature to remain constant during evaporation, liquid must absorb energy from surroundings

Amount of energy liquid must absorb to keep temperature constant during evaporation = enthalpy (heat) of vaporization (∆Hvaporization)

Vaporization is endothermic, so ∆Hvap is positive

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ExampleExample

How much energy is required to vaporize 2.35 g of diethyl ether, (C2H5)2O, at 298 K? ∆Hvap for diethyl ether at 298 K is 29.1 kJ/mol.

2.35g ×1mol

74.123g=0.0317mol ×

29.1kJ1mol

=0.923kJ

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Liquid-vapor equilibriumLiquid-vapor equilibrium

When rate of vaporization = rate of condensation in a closed sysem, system has reached equilibrium

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Vapor PressureVapor Pressure

Pressure exerted by vapor in dynamic equilibrium w its liquid = vapor pressure of that liquid

Vapor pressure depends only on type of liquid & temperature

As long as both phases are present, amount of liquid in container does not affect vapor pressure

Liquids with high vapor pressure at room temperature are volatile (evaporate easily)

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Vapor Vapor pressure pressure curvescurves

Vapor Vapor pressure pressure always always

increases as increases as temperature temperature

increasesincreases

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Vapor pressure and boilingVapor pressure and boiling In open container, evaporation occurs only at surface As temperature increases, evaporation increases

At some point, evaporation begins to occur throughout the liquid instead of just at the surface: boiling!

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Vapor pressureVapor pressure & boiling & boiling

Vapor bubbles form throughout liquid

Bubbles rise to surface, burst, release vapor

All energy is used to convert liquid to vapor, so temperature remains constant while liquid boils

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Boiling pointBoiling point

Boiling begins when the liquid’s vapor pressure matches the external pressure of the atmosphere

The temperature at which this occurs is the boiling point

When the external atmospheric pressure = 1 atm, the boiling point is called the normal boiling point

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The critical pointThe critical point

Liquid heated in a rigid sealed container does not boil Vapor pressure and vapor density increase Liquid density decreases Vapor & liquid densities become equal & meniscus disappears

This point is called the critical point

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The critical pointThe critical point

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Vapor pressure and temperatureVapor pressure and temperature

Clausius-Clapeyron equation shows relationship between vapor pressure and temperature

lnP2

P1

=∆Hvaporization

R1T1

−1T2

⎛

⎝⎜⎞

⎠⎟

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Clausius-Clapeyron Clausius-Clapeyron equationequation

P (vapor pressure) can be in any unitR must be 8.3145 J/mol K∆Hvaporization is usually given in kJ/mol but must be

converted to J/mol to agree with RT is in Kelvins (duh)

lnP2

P1

=∆Hvaporization

R1T1

−1T2

⎛

⎝⎜⎞

⎠⎟

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ExampleExample

The vapor pressure of methanol is 100 mm Hg at 21.2 °C. What is its vapor pressure at 25.0 °C? ∆Hvap for methanol is 38.0 kJ/mol.

lnP2

100mmHg=

38.0 ×103 Jmol

8.3145 JmolK

1294.35K

−1

298.15K⎛⎝⎜

⎞⎠⎟

lnP2

100mmHg=0.19789

P2

100mmHg=e0.19789 =1.22

P2 =122mmHg

lnP2

P1

=∆Hvaporization

R1T1

−1T2

⎛

⎝⎜⎞

⎠⎟

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ExampleExample

The normal boiling point of isooctane is 99.2 °C and its ∆Hvap is 35.76 kJ/mol. What is the vapor pressure of isooctane at 25.0 °C?

lnP2

P1

=∆Hvaporization

R1T1

−1T2

⎛

⎝⎜⎞

⎠⎟

lnP2

760mmHg=

35.76 ×103 Jmol

8.3145 JmolK

1372.35K

−1

298.15K⎛⎝⎜

⎞⎠⎟

lnP2

760mmHg=−2.8746

P2

760mmHg=e−2.8746 =0.056438

P2 =42.9mmHg

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Clausius-Clapeyron Clausius-Clapeyron equationequation

Plot of ln P vs 1/T gives straight line w slope –∆Hvap/R

lnP2

P1

=∆Hvaporization

R1T1

−1T2

⎛

⎝⎜⎞

⎠⎟

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Changes of stateChanges of stateLiquid ↔ gas

Vaporization/boiling and condensation

Solid ↔ liquidMelting (fusion) and freezing

Solid ↔ gasSublimation and deposition

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Heating curveHeating curve

Add energy

Temperature

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Heating curveHeating curve

Add energy

Temperature

(s)

meltingfreezing

boilingcondensing

(l)

(g)

melting/freezing point

boiling point

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Part of a cooling curve for waterPart of a cooling curve for water The dotted line shows

supercooling The water remains liquid

below 0 °C At the bottom of the dotted

line, crystallization begins Crystallization releases

energy; temperature returns to freezing temperature

Temperature remains constant until freezing is completed

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Phase diagramPhase diagram

A graphical representation of the conditions of temperature & pressure under which various phases of a substance exist

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A phase diagram for iodineA phase diagram for iodine

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A phase diagram for carbon dioxideA phase diagram for carbon dioxide

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A phase diagram for waterA phase diagram for water

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Types of solidsTypes of solids

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Molecular substancesMolecular substances

Molecular solids held together by Dispersion Dipole Hydrogen bonding

Relatively low mp & bp For molecules of similar structure,

boiling point increases as molar mass increases

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Ionic substancesIonic substances

Ions held together by lattice forces Coulomb’s law:

Attraction of oppositely charged ions increases with increased charge and/or decreased ion size

Which has a higher mp, NaF or MgO? NaF mp 993 °C, MgO mp 2852 °C

NaCl or KI? NaCl mp 801 °C, KI mp 681 °C

F =kq1q2

r2

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Atomic substancesAtomic substances

Noble gas atoms held together only by dispersion forces

Metals atoms held together by metal cations in sea of electrons

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Atomic substancesAtomic substances

Atoms in network covalent solid held together by covalent bonds

Examples: C (subl 3652 °C), SiC (subl 2700 °C)