Intermolecular Forces

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Chapter 11 Intermolecular Forces
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A phase is a homogeneous part of the system in
contact with other parts of the system but
separated from them by a well-defined boundary.
2 Phases
Solid phase - ice
Liquid phase - water
11.1
3
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 boiling
point melting point DHvap DHfus DHsub
11.2
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Intermolecular forces are feeble but without
them, life as we know it would be impossible.
Water would not condense from vapor into solid or
liquid forms if its molecules didn't attract
each other. Intermolecular forces are
responsible for many properties of molecular
compounds, including crystal structures (e. g.
the shapes of snowflakes), melting points,
boiling points, heats of fusion and
vaporization, surface tension, and densities.
Intermolecular forces pin gigantic molecules
like enzymes, proteins, and DNA into the shapes
required for biological activity.
http//www.nationmaster.com/encyclopedia/ImageMyo
globin.png
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Intermolecular Forces

• 1. London Forces (Dispersion Forces)
• Dipole-Dipole Interactions
• 3. Ion-Dipole Interactions (Salt dissolving in
  solution)
• 4. Hydrogen Bonding

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Dispersion Forces
Occur between every compound and arise from the
net attractive forces amount molecules which is
produced from induced charge imbalances
Figure 10-8 Olmsted Williams
The magnitude of the Dispersion Forces is
dependent upon how easily it is to distort the
electron cloud. The larger the molecule the
greater its Dispersion Forces are.
Figure 10-9 Olmsted Williams
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The boiling point of alkanes increase with the
length of the carbon chain. Long-chain alkanes
have larger dispersion forces because of the
increased polarizability of their larger electron
cloud.
Olmsted Williams Fig 10-10
Pg 437
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How molecular shape affects the strength of the
dispersion forces The shapes of the molecules
also matter. Long thin molecules can develop
bigger temporary dipoles due to electron movement
than short fat ones containing the same numbers
of electrons. Long thin molecules can also lie
closer together - these attractions are at their
most effective if the molecules are really
close. For example, the hydrocarbon molecules
butane and 2-methylpropane both have a molecular
formula C4H10, but the atoms are arranged
differently. In butane the carbon atoms are
arranged in a single chain, but 2-methylpropane
is a shorter chain with a branch.
                                                  
                                                  
                                                  
       Butane has a higher boiling point because
the dispersion forces are greater. The molecules
are longer (and so set up bigger temporary
dipoles) and can lie closer together than the
shorter, fatter 2-methylpropane molecules.
http//www.chemguide.co.uk/atoms/bonding/vdw.html
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Polarizability
the ease with which the electron distribution in
the atom or molecule can be distorted.

• Polarizability increases with
• greater number of electrons
• more diffuse electron cloud

11.2
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Is the Molecule Polar?
We have already talked about diatomic molecules.
The more Electronegative atom will pull the
electron density of the bond Closer to itself
giving it a partial negative charge leaving the
other Atom with a partially positive charge. Thus
giving the molecule A dipole moment. But what
about molecules made up of more than two
molecules?
11
Molecules with 3 Atoms
Even though the C-O bond is polar, the bonds
cancel each other out because the molecule is
linear the dipole moments are equal and in
opposite directions.Therefore CO2 is non-polar.
CO2
The dipole moment between H-C points in the
direction of C. The dipole moment points between
C-N points in the direction of the N. Therefore
the dipole vectors are additive and HCN is polar
HCN
SO2
SO2 is a polar molecule because the S-O dipole
Moments dont cancel each other out due to the
angle
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Molecules with 4 Atoms
CCl4 is non-polar
CHCl3 is polar
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How to Determine if a Molecule Is Polar

• Draw Lewis Structure
• If all of the regions of electron density are
• bound to the same thing (CCl4 CO2 ) than the
• molecule is non-polar
• If the regions of electron density are not bound
  to
• the same thing than the molecule is polar (HCN
  SO2)

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dipole moment polar molecule
dipole moment polar molecule
no dipole moment nonpolar molecule
no dipole moment nonpolar molecule
10.2
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Dipole-Dipole Forces
Attractive forces between polar molecules
11.2
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Dipole Forces occur between molecules containing
a dipole moment. The positive end of the dipole
moment on one mole is attracted to the Negative
end of the dipole moment on a nearby molecule.
Consider 2-methyl propane (left) and acetone
(right) Both compounds are about Equal in size
and shape therby Having similar dispersion
forces, But Acetone contains an Oxygen (red) and
causes the Molecule to have a dipole Moment
allowing it to have Dipole forces and thus a
Higher boiling point
Figure 10-11
Olmsted Williams
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Ion-Dipole Forces
Attractive forces between an ion and a polar
molecule
The larger the charge the stronger the force
11.2
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Fig 10-34
Olmsted Williams
A molecular picture showing the ion-dipole
Interaction that helps a solid ionic crystal
dissolve in water. The arrows indicate
ion-dipole interactions.
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What type(s) of intermolecular forces exist
between each of the following molecules?
HBr
HBr is a polar molecule dipole-dipole forces.
There are also dispersion forces between HBr
molecules.
CH4
CH4 is nonpolar dispersion forces.
SO2
SO2 is a polar molecule dipole-dipole forces.
There are also dispersion forces between SO2
molecules.
11.2
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Hydrogen Bond
The hydrogen bond is a special dipole-dipole
interaction between they hydrogen atom in a polar
N-H, O-H, or F-H bond and an electronegative O,
N, or F atom.
A B are N, O, or F
11.2
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Fig 10-16B Pg 444
Crystals of benzoic acid contain pairs of
molecules held together head to head by
hydrogen bonds. These pairs then stack in
planes which are Held together by dispersion
forces.
Courtesy Stephen Frisch
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Intermolecular Forces

• 1. London Forces (Dispersion Forces)
• Dipole-Dipole Interactions
• 3. Ion-Dipole Interactions (Salt dissolving in
  solution)
• 4. Hydrogen Bonding

These forces affect how molecules will interact
with each other and As a general rule as the
strength of the force increases the boiling
Point of the compound increases
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Liquids and Surface Tension
Surface tension is the amount of energy required
to stretch or increase the surface of a liquid by
a unit area.
Strong intermolecular forces
High surface tension
11.3
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Properties of Liquids
Cohesion is the intermolecular attraction between
like molecules
Adhesion is an attraction between unlike molecules
11.3
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Condensation
Evaporation
11.8
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The equilibrium vapor pressure is the vapor
pressure measured when a dynamic equilibrium
exists between condensation and evaporation
A substance with a high Vapor pressure is
considered To be volitile therefore, the
lower The boiling point the higher the Vapor
pressure and the weaker The intermolecular forces
11.8
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The boiling point is the temperature at which the
(equilibrium) vapor pressure of a liquid is equal
to the external pressure.
The normal boiling point is the temperature at
which a liquid boils when the external pressure
is 1 atm.
11.8
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The critical temperature (Tc) is the temperature
above which the gas cannot be made to liquefy, no
matter how great the applied pressure.
The critical pressure (Pc) is the minimum
pressure that must be applied to bring about
liquefaction at the critical temperature.
11.8
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The melting point of a solid or the freezing
point of a liquid is the temperature at which the
solid and liquid phases coexist in equilibrium
Freezing
Melting
11.8
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11.8
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Molar heat of sublimation (DHsub) is the energy
required to sublime 1 mole of a solid.
Sublimation
Deposition
DHsub DHfus DHvap
( Hesss Law)
11.8
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A phase diagram summarizes the conditions at
which a substance exists as a solid, liquid, or
gas.
Phase Diagram of Water
11.9
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11.9