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Ch. 11: Liquids and Intermolecular Forces

Learning goals and key skills:

� Identify the intermolecular attractive interactions (dispersion, dipole-dipole, hydrogen bonding, ion-dipole) that exist between molecules or ions based on their composition and molecular structure and be able to compare the relative strengths of these intermolecular forces � Explain the concept of polarizability and how it relates to dispersion forces

� Explain the concepts of viscosity and surface tension in liquids� Know the names of various changes of state for a pure substance�Interpret heating curves and be able to calculate quantities related to temperature and enthalpies of phase changes

� Define critical pressure, critical temperature, vapor pressure, normal boiling point, normal melting point, critical point, and triple point� Be able to interpret and sketch phase diagrams. Explain how water’s phase diagram differs

from most other substances, and why.�Understand how the molecular arrangements characteristic or nematic, smectic, and choloesteric liquid crystals differ from ordinary liquids and from each other. Be able to recognize the features of molecules that favor formation of liquid crystalline phases.

GasesGases

In gases the particles/molecules are far apart and the intermolecular forces between them are weak.

... vs liquids and solids... vs liquids and solids

In liquids and solids this

is not true. Instead we must consider the intermolecular forces, i.e., the forces that exist between atoms and molecules.

Strength of intermolecular attractionsStrength of intermolecular attractions

States of MatterStates of Matter

The fundamental difference between states of matter is the distance between particles.

Because in the solid and liquid states particles are closer together, we refer to them as condensed phases.

The States of MatterThe States of Matter

The state a substance is in at a particular temperature and pressure depends on:

– The kinetic energy of the particles

– The strength of the attractions between the particles

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Intermolecular ForcesIntermolecular Forces

Intermolecular ForcesIntermolecular Forces

•ion-dipole

•hydrogen bonding

•dipole-dipole van der Waals

•dispersion forces (London) forces

The strength of these intermolecular forces is directly related to the melting/boiling points, enthalpy of fusion, enthalpy of vaporization, and solubility of the substances.

Dispersion forcesDispersion forces

These induced dipoles lead to intermolecular attractions. This is how nonpolar gas molecules (e.g. He, N2, etc.) can liquefy.

The ease with which the electron distribution in a molecule is distorted is called its polarizability.

London dispersion forcesLondon dispersion forces

Dispersion forces tend to increase with increasing molecular weight.

MW Boiling Pt.

(g/mol) (K)

F2 38 85

Cl2 71 239

Br2 160 332

I2 254 458

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Factors Affecting Dispersion ForcesFactors Affecting Dispersion Forces

• The shape of the molecule affects the dispersion forces: long, thin molecules (like n-pentane tend to have stronger dispersion forces than short, round ones (like neopentane).

• This is due to the increased surface areain n-pentane.

Polar Covalent Bonds and Electronegativity

Polar Covalent Bonds and Electronegativity

Although atoms often form compounds by sharing electrons, the electrons are not always shared equally.

Electronegativity is the ability of atoms in a molecule to attract electrons to itself.

Dipole-Dipole ForcesDipole-Dipole Forces Dipole-Dipole Forces

Dipole-Dipole Forces

Influence of dipole-dipole forces is seen in the boiling points of simple molecules.

MW (g/mol) Boiling Pt (K)

N2 28 77

CO 28 81

Br2 160 332

ICl 162 370

Dipole-Dipole ForcesDipole-Dipole ForcesDipole-dipole forces increase with increasing polarity (dipole moments).

MW Dipole Boiling

(g/mol) Moment, µ (D) Pt (K)

CH3CH2CH3 44 0.1 231

CH3OCH3 46 1.3 248

CH3CHO 44 2.7 294

CH3CN 41 3.9 355

Hydrogen BondingHydrogen Bonding

A special form of dipole-dipole attraction, which enhances dipole-dipole attractions.

H-bonding is strongest when X and Y are N, O, or F

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Boiling Points of Simple Hydrogen-Containing Compounds

Base-Pairing through H-BondsBase-Pairing through H-Bonds

Portion of a DNA chain

Double helix of DNA

Ion-dipole forcesIon-dipole forces

HH

water dipole

••

••

O-δδδδ

+δδδδ

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Intermolecular Forces SummaryIntermolecular Forces Summary

Ionic bonding Ions

Ion-dipoleIon charge, dipole

moment

Hydrogen bondingVery polar H-X bond and lone pair of e-

(X = N, O, F)

Dipole-dipole Dipole moment

Dispersion forces (Induced dipole-

induced dipole) onlypolarizabilityIn

crea

sing

Str

engt

h“Like dissolves like”“Like dissolves like”

Polar molecules are more likely to dissolve in a polar solvent, and

nonpolar molecules are more likely to dissolve in a nonpolar solvent.

LiquidsLiquids

• molecules close together → appreciable intermolecular forces

• liquids are almost incompressible• molecules are in constant motion• finite volume, but indefinite shape

ViscosityViscosity

The resistance of liquids to flow

Increases with stronger intermolecular forces and decreases with increasing temperature.

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Surface TensionSurface Tension

Surface tension results from the net inward force experienced by the molecules on the surface of a liquid.

Cohesive and adhesive forcesCohesive and adhesive forcesCohesive forces

- between the liquid molecules

Adhesive forces- between the liquid molecules and another substance

Properties of liquidsProperties of liquids

Viscosity – the resistance of liquids to flow; “the thickness”

Surface tension – energy required to break through the surface or to disrupt a liquid drop and make the drop spread out like a film.

Cohesive forces – forces between liquid molecules

Adhesive forces – forces between liquid molecules and another substance

Capillary action – the rise of liquids up narrow tubes and other surfaces

Phase changesPhase changes

Changes of state involve energy (at constant T)

Ice + (heat of fusion) → water

Water + (heat of vaporization) → steam

Phase changesPhase changes Energy changes associated with changes of state

Energy changes associated with changes of state

For water ∆Hfus = 6.01 kJ/mol or 334 J/g∆Hvap = 40.67 kJ/mol or 2258 J/g

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Heating/Cooling Curve for WaterHeating/Cooling Curve for Water Heating/Cooling Curve for WaterHeating/Cooling Curve for Water

The total heat is the sum of each step qtotal = ∑ qiThe total heat is the sum of each step qtotal = ∑ qi

The ice is heated qice = sicem∆T

The ice melts to water qfus = ∆Hfusm

The water is heated qwater = swaterm∆T

The water evaporates to steam qvap = ∆Hvapm

The steam is heated qsteam = ssteamm∆T

The ice is heated qice = sicem∆T

The ice melts to water qfus = ∆Hfusm

The water is heated qwater = swaterm∆T

The water evaporates to steam qvap = ∆Hvapm

The steam is heated qsteam = ssteamm∆T

Example: Heating curve problemExample: Heating curve problem

Determine the amount of heat (in kJ) required to heat

500. g of ice from -50.0 °C to steam at 200. °C.

sice = 2.09 J/gK

∆Hfus = 334 J/g

swater = 4.184 J/gK

∆Hvap = 2258 J/g

ssteam=1.84 J/gK

qice, -50.0 to 0 °C = +52 250 J

qfus, ice to water = +167 000 J

qwater, 0 to 100 °C = +209 200 J

qvap, water to steam = +1 129 000 J

qsteam, 100 to 200. °C = +92 000 J

qtotal = 1 650 000 J = 1650 kJ

Note that most of the heat is used to convert water into steam.

Example: Heating curve problemExample: Heating curve problem

Determine the amount of heat (in kJ) required to

convert 42.0 g of ethanol from -155 °C to 78 °C.

ethanol (C2H5OH) melts at -114 °C and boils at 78 °C.

The enthalpy of fusion of ethanol is 5.02 kJ/mol and

the enthalpy of vaporization is 38.56 kJ/mol. The

specific heats of solid and liquid ethanol are 0.97 J/gK

and 2.3 J/gK, respectively.

Critical Temperature: The highest temperature at which a distinct liquid phase can form

Critical Pressure: Pressure required to bring about liquefaction at this critical temperature.

Supercritical Fluid: Above this critical point, we have a supercritical fluid where the density is similar to a liquid and the viscosity is similar to a gas.

Vapor PressureVapor Pressure

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Temperature effects on the distribution of energy in a liquid

Temperature effects on the distribution of energy in a liquid

E needed to evaporate liquid

Liquids that evaporate rapidly are volatile

Vapor pressure curvesVapor pressure curves

Equilibrium Vapor Pressure & the Clausius-Clapeyron Equation

Equilibrium Vapor Pressure & the Clausius-Clapeyron Equation

• Clausius-Clapeyron

equation — used to find

heat of vaporization, ∆Hvap.

• The logarithm of the vapor

pressure P is proportional

to ∆Hvap and to 1/T.

• ln P = –(∆Hvap/RT) + C

−

∆−=

121

2 11ln

TTR

H

P

P vap

Clausius-Clapeyron EquationClausius-Clapeyron Equation

Liquid bromine has a vapor pressure of

400. torr at 41.0 °C and a normal boiling

point of 58.2 °C. Calculate the heat of

vaporization of bromine (in kJ/mol).

−

∆−=

121

2 11ln

TTR

H

P

P vap

Phase DiagramsPhase DiagramsPhase diagrams display the state of a substance at various pressures and temperatures and the places where equilibria exist between phases.

Normal Phase Diagram: CO2Normal Phase Diagram: CO2

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Phase Diagram of WaterPhase Diagram of Water

Note the high critical temperature and critical pressure.

–These are due to the strong van der Waals forces between water molecules.

Info from P-T phase diagramsInfo from P-T phase diagrams

From a (pressure-temperature) phase diagram we can find:

• the normal melting point

• the normal boiling point

• the triple point

• whether the liquid is more or less dense than the solid

• the vapor pressure curve

• the sublimation pressure curve

• the critical temperature (Tc) and critical pressure (Pc)

Liquid CrystalsLiquid Crystals

Liquid Crystals

• Rod-shaped and contain double or triple bond near the middle

• Polar groups create a dipole moment and promote alignment

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