Chapter 12
Intermolecular Attractions
and the properties of liquids
and Solids
12.1 Gases, liquids, and solids and intermolecular
distances
• Learning Objective:
• To understand why the physical properties of liquids and
solids depend so heavily on their chemical composition while
the properties of gases don’t
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
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Intermolecular Attractions
• Converting gas  liquid or solid
• Molecules must get closer together
• Cool or compress
• Converting liquid or solid  gas
• Requires molecules to move farther apart
• Heat or reduce pressure
• As T decreases, kinetic energy of molecules decreases
• At certain T, molecules don’t have enough energy
to break away from one another’s attraction
Intermolecular Forces
• Physical state of molecule depends on
• Average kinetic energy of particles
• Recall KE  Tave
• Intermolecular Forces
• Energy of Inter-particle attraction
Physical properties of gases, liquids and solids
determined by
• How tightly molecules are packed together
• Strength of attractions between
molecules
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12.2 Types of intermolecular Forces
• Learning Objectives:
• To learn nature and relative strength of the principal kinds of
intermolecular attractions
• To understand the factors that influence the strengths of
intermolecular attractions
• To be able to use a Lewis structure, bond polarities and
predicted molecular geometry to anticipate the nature of the
intermolecular attractions that exist in liquid and solid states of
a substance.
Intermolecular 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
Covalent Bond (strong)
Cl
H
Intermolecular attraction (weak)
Cl
H
Intermolecular Forces
Intermolecular forces are attractive forces between molecules.
Intramolecular forces hold atoms together in a molecule.
Intermolecular vs Intramolecular
•
41 kJ to vaporize 1 mole of water (inter)
•
930 kJ to break all O-H bonds in 1 mole of water (intra)
“Measure” of intermolecular force
Generally, intermolecular
forces are much weaker
than intramolecular
forces.
boiling point
melting point
DHvap
DHfus
DHsub
Electronegativity Review
Electronegativity: Measure of attractive force that one atom in a
covalent bond has for electrons of the bond
9
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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Net Dipoles
• Symmetrical molecules
• Even if they have polar bonds
• Are non-polar because bond
dipoles cancel
• Asymmetrical molecules
• Are polar because bond dipoles
do not cancel
• These molecules have
permanent, net dipoles
• Molecular dipoles
• Cause molecules to interact
• Decreased distance between
molecules increases amount of
interaction
Intermolecular Forces
•
When substance melts or boils
• Intermolecular forces are broken
• Not covalent bonds
•
•
•
Responsible for non-ideal behavior of gases
Responsible for existence of condensed states of matter
Responsible for bulk properties of matter
• Boiling points and melting points
• Reflect strength of intermolecular forces
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Three Important Types of Intermolecular Forces
1.
London dispersion forces
2.
Dipole-dipole forces
• Hydrogen bonds
3.
Ion-dipole forces
• Ion-induced dipole forces
13
London Forces
• When atoms near one another, their
valence electrons interact
• Repulsion causes electron clouds in each
to distort and polarize
• Instantaneous dipoles result from this
distortion
• Effect enhanced with increased
volume of electron cloud size
• Effect diminished by increased
distance between particles and
compact arrangement of atoms
London Forces
Instantaneous dipoleinduced dipole
attractions
•
London Forces
•
Dispersion
forces
Operate between all
molecules
•
Neutral or net
charged
•
Nonpolar or
polar
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London Dispersion Forces
•
Ease with which dipole moments can be induced and
thus London Forces depend on
1. Polarizability of electron cloud
2. Points of attraction
• Number atoms
•
Molecular shape (compact or elongated)
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Polarizability
 Plarizability: ease with which the electron cloud can be distorted
 Larger molecules often more polarizable
 less tightly held electrons
Table 12.1 Boiling Points of Halogens and Noble
Gases
Larger molecules have stronger London forces and
thus higher boiling points.
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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
How Intermolecular Forces Determine Physical
Properties
Hexane, C6H14
Propane, C3H8
BP 68.7 °C
BP –42.1 °C
 More sites (marked with *) along its chain where
attraction to other molecules can occur
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Molecular Shape
• Increased surface area available for contact = increased
London forces
• London dispersion forces between spherical molecules are
lower than chain-like molecules
• More compact molecules
• Hydrogen atoms not as free to interact with hydrogen
atoms on other molecules
• Less compact molecules
• Hydrogen atoms have more chance to interact with
hydrogen atoms on other molecules
21
Physical Origin of Shape Effect
• Small area for
interaction
• Larger area for
interaction
More compact – lower BP Less compact – higher BP
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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
• Positive to negative
+

+

• As dipole moment increases,
intermolecular force increases
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Dipole-Dipole Attractions
• Tumbling molecules
• Mixture of attractive and
repulsive dipole-dipole
forces
• Attractions (- +) are
maintained longer than
repulsions(- -)
• Get net attraction
• ~1–4% of covalent bond
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Dipole-Dipole Attractions
• Interactions between net dipoles in polar molecules
• About 1–4% as strong as a covalent bond
• Decrease as molecular distance increases
• Dipole-dipole forces increase with increasing polarity
25
Hydrogen Bonds
• Special type of dipole-dipole Interaction
• Very strong dipole-dipole attraction
• ~10% of a covalent bond
• Occurs between H and highly electronegative atom (O, N, or F)
• H—F, H—O, and H—N bonds very polar
• Electrons are drawn away from H, so high partial charges
• H only has one electron, so +H presents almost bare
proton
• –X almost full –1 charge
• Element’s small size, means high charge density
• 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
H
O
H
H
N
H
H
H
H
H
H
H
F
N
O
H
H
H
O
H
H
H
F
H
N
N
H
H
H
H
O
H
H
N
H
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Hydrogen Bonding in Water
water crystallizes into an open hexagonal form.
• 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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Example:
List all intermolecular forces for CH3CH2OH.
A. Hydrogen-bonds
B. Hydrogen-bonds, dipole-dipole attractions, London
dispersion forces
C. Dipole-dipole attractions
D. London dispersion forces
E. London dispersion forces, dipole-dipole attractions
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Example:
In the liquid state, which species has the strongest
intermolecular forces, CH4, Cl2, O2 or HF?
A. CH4
B. Cl2
C. O2
D. HF
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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
• Attractions between ion and polar molecules
 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
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Ion-Induced Dipole Attractions
• Attractions between ion and dipole it induces on neighboring
molecules
• Depends on
• Ion charge and
• Polarizability of its neighbor
• Attractions can be quite strong as ion charge is constant,
unlike instantaneous dipoles of ordinary London forces
• E.g., I– and Benzene
Summary of Intermolecular Attractions
Dipole-dipole
• Occur between neutral molecules with permanent dipoles
• About 1–4% of covalent bond
• Mid range in terms of intermolecular forces
Hydrogen bonding
• Special type of dipole-dipole interaction
• Occur when molecules contain N—H,
H—F and O—H bonds
• About 10% of a covalent bond
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Summary of Intermolecular Attractions
London dispersion
• Present in all substances
• Weakest intermolecular force
• Weak, but can add up to large net attractions
Ion-dipole
• Occur when ions interact with polar molecules
• Strongest intermolecular attraction
Ion-induced dipole
• Occur when ion induces dipole on neighboring particle
• Depend on ion charge and polarizability of its neighbor
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