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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

Inter vs. Intra-Molecular Forces
• Intra molecular forces
– Covalent bonds within molecule
– Strong
– 
H bond
(HCl) = 431 kJ/mol
• Inter molecular forces
– Attraction forces between molecules
– Weak
– 
H vaporization
(HCl) = 16 kJ/mol
Covalent Bond (strong) Intermolecular attraction (weak)
Cl H Cl H
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• 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 : Measure of attractive force that one atom in a covalent bond has for electrons of the bond

• Two atoms with different electronegativity values share electrons unequally
• Electron density is uneven
– Higher charge concentration around more electronegative atom
• Bond dipoles
H F
– Indicated with delta (δ) notation    
– Indicates partial charge has arisen
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1. Dipole-dipole forces
– Hydrogen bonds
2. London dispersion forces
3. Ion-dipole forces
– Ion-induced dipole forces

I.
• 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.
• 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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• 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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H F H O
H O H O
H
H H
H
H N H O
H
H H
H H
H F H N
H N H N
H
H
H H
H O H N
H H
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• 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

)

• 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.
• Intermolecular forces between
Electrostatic attraction nonpolar molecules e
 e

• Two neutral molecules (atoms) can
2+ affect each other e
 e

– Nucleus of 1 molecule (atom) attracts e

’s of adjacent molecule (atom) He atom 1
2+
He atom 2
– Electron cloud distorts
– Temporary or instantaneous dipole forms
– One instantaneous dipole can induce another in adjacent molecule (atom)
– Results in net attractive force
       
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• Instantaneous dipole-induced dipole attractions
– London Dispersion Forces
– London forces
– Dispersion forces
• Decrease as 1/d 6 (d = distance between molecules)
• Effect enhanced with increased particle mass
• Operate between all molecules
– Neutral or net charged
– Nonpolar or polar
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 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.

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
CH
C
2
H
4
6

161.5

88.6
C
C
5
6
H
H
12
14
36.1
68.7
C
C
4
3
H
H
8
10

42.1

0.5
C
22
:
H
46
:
327
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III.
• 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

•
Ex. Ion-dipole Attractions AlCl
3
·6H
2
Attractions between ion and polar molecules
O
 Positive charge of Al 3+ 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
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• Often can predict physical properties (like BP and
MP) by comparing strengths of intermolecular attractions
– Ion-Dipole
Strongest
– Hydrogen Bonding
– Dipole-Dipole
– London Dispersion Forces
Weakest
• Larger, longer, heavier molecules have stronger IMFs
• Smaller, more compact, lighter molecules have weaker IMFs
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• Changes of physical state
– Deal with motion of molecules
• As temperature changes
– Matter will undergo phase changes
• Liquid

Gas
– Evaporation
– As heat H
2
O, forms steam or water vapor
– Requires energy or source of heat to occur
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• Solid

Gas
– Sublimation
– Ice cubes in freezer, leave in long enough disappear
– Endothermic
• Gas

Liquid
– Cooling or Condensation
– Dew is H
2
O vapor condensing onto cooler ground
– Exothermic
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Gas
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Vaporization Condensation
Sublimation
Liquid
Melting or Fusion
Freezing
Solid

Exothermic, releases heat

Endothermic, absorbs heat
Deposition

• 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
A B
• Smaller IMF’s
• Lower KE required to escape liquid
• A evaporates faster
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• Larger IMF’s
• Higher KE required to escape liquid
• B evaporates slower

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