Intermolecular Attractions and the Properties of Liquids and Solids

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Chapter 12 Intermolecular Forces• Important differences between gases,

solids, and liquids:– Gases

• Expand to fill their container – Liquids

• Retain volume, but not shape– Solids

• Retain volume and shape

At room temperature, some are solid, others are liquid, others are gaseous.

Why?

• Physical Properties of Gases, Liquids and Solids determined by – How tightly molecules are packed together– Strength of attractions between molecules

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Inter vs. Intra-Molecular Forces• Intramolecular forces

– Covalent bonds within molecule – Strong – Hbond (HCl) = 431 kJ/mol

• Intermolecular forces – Attraction forces between molecules– Weak– Hvaporization (HCl) = 16 kJ/mol

Cl H Cl H

Covalent Bond (strong) Intermolecular attraction (weak)

• When substance melts or boils– Intermolecular forces are broken– Not covalent bonds

• Responsible for existence of condensed states of matter

• Responsible for bulk properties of matter– Boiling Points and Melting Points

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Electronegativity Review

Electronegativity: Measure of attractive force that one atom in a covalent bond has for electrons of the bond

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Bond Dipoles• Two atoms with different electronegativity

values share electrons unequally• Electron density is uneven

– Higher charge concentration around more electronegative atom

• Bond dipoles – Indicated with delta (δ) notation– Indicates partial charge has arisen

H F

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Three Important Types of Intermolecular Forces

1. Dipole-dipole forces– Hydrogen bonds

2. London dispersion forces3. Ion-dipole forces

– Ion-induced dipole forces

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I. Dipole-dipole Attractions• Occur only between polar

molecules– Possess dipole moments

• Molecules need to be close together

• Polar molecules tend to align their partial charges– + to –

• As dipole moment , intermolecular force

+ +

+ +

+ +

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I. Dipole-dipole Attractions• Tumbling molecules

– Mixture of attractive and repulsive dipole-dipole forces

– Attractions (- -) greater than repulsions(- -)

– Get net attraction – ~ 1% of covalent bond

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Hydrogen Bonds

• Special type of Dipole-Dipole Interaction– Very strong dipole-dipole attraction – ~40 kJ/mol

• Occurs between H and highly electronegative atom (O, N, or F)– H—F, H—O, and H—N bonds very polar

• Positive end of one can get very close to negative end of another

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Examples of Hydrogen Bonding

H O

H

H O

H

H O

H

H N

H

H

H F H O

H

H F H N

H

H

H N

H

H

H N

H

H

H N

H

H

H O

H

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Effects of Hydrogen Bonding

• Boiling points of H compounds of elements of Groups IVA, VA, VIA, and VIIA.

• Boiling points of molecules with H bonding are higher than expected.

• Don’t follow rule that BP as MM (London forces )

Boili

ng P

oint

(°C)

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Hydrogen Bonding in Water

• Responsible for expansion of water as it freezes • Hydrogen bonding produces strong attractions in liquid • Hydrogen bonding (dotted lines) between

water molecules in ice form tetrahedral configuration

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II. London Dispersion Forces• Intermolecular forces between

nonpolar molecules• Two neutral molecules (atoms) can

affect each other– Nucleus of 1 molecule (atom) attracts

e’s of adjacent molecule (atom)– Electron cloud distorts– Temporary or instantaneous dipole

forms– One instantaneous dipole can induce

another in adjacent molecule (atom)– Results in net attractive force

e

e

2+

e

e

Electrostaticattraction

He atom 1 He atom 2

2+

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London Dispersion Forces• Instantaneous dipole-induced dipole

attractions – London Dispersion Forces– London forces– Dispersion forces

• Decrease as 1/d6 (d = distance between molecules)

• Effect enhanced with increased particle mass• Operate between all molecules

– Neutral or net charged– Nonpolar or polar

London Forces as MM More e, less tightly held

London Forces as electron cloud volume (size)

Larger molecules have stronger London forces and thus higher boiling points.

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2. Number of Atoms in Molecule

• London forces depend on number atoms in molecule• Boiling point of hydrocarbons demonstrates this trend

Formula BP at 1 atm, C Formula BP at 1 atm, C

CH4 161.5 C5H12 36.1

C2H6 88.6 C6H14 68.7

C3H8 42.1 : :

C4H10 0.5 C22H46 327

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III. Ion-dipole Attractions• Attractions between ion and charged end of

polar molecules– Attractions can be quite strong as ions have full

charges

(a) Negative ends of water dipoles surround cation (b) Positive ends of water dipoles surround anion

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Ex. Ion-dipole Attractions AlCl3·6H2O

Positive charge of Al3+ ion attracts partial negative charges – on O of water molecules

Ion-dipole attractions hold water molecules to metal ion in hydrate Water molecules are found at

vertices of octahedron around aluminum ion

• Attractions between ion and polar molecules

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Using Intermolecular Forces• Often can predict physical properties (like BP and

MP) by comparing strengths of intermolecular attractions– Ion-Dipole– Hydrogen Bonding– Dipole-Dipole– London Dispersion Forces

• Larger, longer, heavier molecules have stronger IMFs• Smaller, more compact, lighter molecules have

weaker IMFs

Weakest

Strongest

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Phase Changes

• Changes of physical state – Deal with motion of molecules

• As temperature changes– Matter will undergo phase changes

• Liquid Gas– Evaporation– As heat H2O, forms steam or water vapor– Requires energy or source of heat to occur

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Phase Changes • Solid Gas

– Sublimation– Ice cubes in freezer, leave in long enough disappear– Endothermic

• Gas Liquid– Cooling or Condensation– Dew is H2O vapor condensing onto cooler ground– Exothermic

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Phase ChangesEn

ergy

of S

yste

m

Gas

Solid

Liquid

Meltingor Fusion

Vaporization Condensation

Freezing

SublimationDeposition

Exothermic, releases heat Endothermic, absorbs heat

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Rate of Evaporation• Depends on

– Temperature– Surface area– Strength of

intermolecular attractions

• Molecules that escape from liquid have larger than average KE’s

• When they leave– Average KE of

remaining molecules is less

– T lower

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Effect of Temperature on Evaporation Rate

• For given liquid– Rate of evaporation per

unit surface area as T • Why?

– At higher T, total fraction of molecules with KE large enough to escape is larger

– Result: rate of evaporation is larger

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Kinetic Energy Distribution in Two Different Liquids

• Smaller IMF’s• Lower KE required to

escape liquid• A evaporates faster

• Larger IMF’s• Higher KE required to

escape liquid• B evaporates slower

A B

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Vapor Pressure Diagram

T-t curves

Supercooling

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Phase Diagram of Water