Transcript Chapter 11 Intermolecular Forces - X

Chemistry 100 – Chapter 11
Intermolecular Forces
Intramolecular and Intermolecular
Forces
We have just been discussing the covalent
bond - the force that holds atoms together
making molecules.
 We have also talked about the ionic bond.
These are intramolecular forces.
 There are also forces that cause molecules
to attract each other. These are called
intermolecular forces.
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Intermolecular Forces
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Salt (NaCl) is a solid because of the strong
electrostatic attraction of the Na+ and the
Cl– ions: the ionic bond
Q: Why is Cl2 a gas, Br2 a liquid and I2 a
solid? A: Intermolecular forces
Q: Why is molasses “thick” while water has
low viscosity? Same answer.
Q: Why is it possible to float a needle on
water?
Intermolecular Forces (cont’d)
Q: Why is water a liquid but H2S is a
gas at 25ºC?
 Q: Why are real gases not ideal? A:
van der Waals forces (same thing,
different name)
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States of Matter
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Gas: No defined shape or volume,
compressible, rapid diffusion, flows readily
Liquid: Takes shape of container, virtually
incompressible, low diffusion, flows readily
Solid: Has its own shape and volume,
virtually incompressible, extremely slow
diffusion, does not flow
Liquids and solids are called condensed
phases because particles are close together
Intermolecular Forces
Much weaker than chemical bonds.
Covalent bonds 200 kJ/mol and more.
 Intermolecular forces less than 50
kJ/mol
 When a liquid vaporizes, the
intermolecular forces must be
overcome. But no covalent bonds are
broken.
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Ion - dipole forces
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Dissolve an ionic
compound in water
Charged ions
interact with the
dipole of water
molecules
Dipole-dipole Forces
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Interactions between the dipoles in a polar
liquid leads to a net attraction
Dipole - Induced dipole force
In a mixture of two liquids were one is
polar and the other is not, the dipole of
the polar molecule can induce a dipole in
the other.
 The energy of this force depends on the
polarizabilty of the non-polar molecule.
Larger molecules are more polarizable
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London Dispersion Forces
How do we account for the fact that
non-polar gases can be liquefied and
solidified?
 Fritz London proposed instant
dipoles. These are dipoles that result
from the random movement of the
electron cloud.
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Dispersion forces work
over very short distances
Polarizability
The ease with which a dipole can be
induced depends on the polarizability of
the molecule
 Large molecules are more polarizable easier to distort the electron cloud
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Polarizability and Boiling Points
Boiling points (K) of halogens
F2
Cl2
Br2
I2
85.1
238.6
332.0
457.6
Boiling points of Noble gases
He
Ne
Ar
Kr
4.6
27.3
87.5
120.9
Xe
166.1
Dipole or Dispersion?
Dispersion forces operate between all
molecules - polar and non-polar
 Molecules with comparable molecular
weights and shapes have approximately
equal dispersion forces.
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Any difference is due to dipole-dipole
attractions
When molecules differ widely in
molecular weight, dispersion forces tend
to be the decisive ones
Water And Ammonia Have Unusual
Boiling Points!!
The Hydrogen Bomb (err – Bond)
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A special dipole-dipole intermolecular attraction
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H in a polar bond (H-F, H-O or H-N)
an unshared electron pair on a nearby electronegative
ion or atom (generally F, O or N)
Hydrogen bonds (4 to 25 kJ/mol, or larger) are
weaker than covalent bonds but stronger than
most dipole-dipole or dispersion forces.
Molecules with H bonding
HF (can behave as if it were H2F2)
 NH3
 H2O (ice is less dense than liquid at 0ºC)
 alcohols (e.g. CH3OH)
 amines (e.g CH3NH2)
 carboxylic acids (e.g. CH3COOH)
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Summary
Liquids - viscosity
Viscosity (resistance to flow)
 Molecules that have strong intermolecular
forces cannot move very easily - more
viscous
 Viscosity decreases at higher
temperatures. The kinetic energy
overcomes the intermolecular forces
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Liquids - surface tension
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Surface tension: the energy that
must be expended to increase the
surface of a liquid
Surface tension
Liquids with strong intermolecular bonds
have high surface tensions
Water: strong H-bonds, so high surface
tension, 7.29  10–2 J/m2
Mercury: atoms held by metallic bonds, so
even higher surface tension, 4.6  10–1 J/m2
Wetting & capillary action
Water wets clean glass (spreads out)
but beads on a waxy surface
 Water climbs up a capillary tube
 Mercury does not wet glass
 Mercury level is depressed in
capillary tube
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To wet or not to wet?
Competition between two tendencies
 Cohesive forces; intermolecular forces
that bind similar molecules together.
Keeps liquid as a bead
 Adhesive forces: Intermolecular forces
between molecules of a liquid and those of
a surface. Makes liquid spread out
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Are we all wet??
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Water on clean glass
adhesive forces > cohesive forces
explains the upward meniscus of water
Water on polished table
cohesive forces win because H2O molecules
are not attracted to wax
Mercury on glass - cohesive forces win
because Hg molecules are not attracted to
glass
Explains depression in capillary tube and
downward meniscus
Phase changes
Phase changes
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Enthalpy of fusion or heat of fusion
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for water Hfus = 6.01 kJ/mol.
Enthalpy of vaporization or heat of
vaporization
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for water Hvap = 40.67 kJ/mol.
cooling effect of evaporation
refrigeration (Not Freon-12)
steam burn generally severe
Heating Curve
Supercooling, superheating
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A liquid cooled below its freezing point is
said to be supercooled
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requires very clean conditions
the molecules are moving slowly but have not
organized themselves into the solid form
A liquid heated above its boiling point is
said to be superheated
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a danger with heating water in microwave oven
Critical T and P
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A gas can be liquefied by cooling
A gas can be liquefied by increasing pressure but
only if the temperature is below the compound’s
critical temperature
substance
critical temp K and C
ammonia
405.6 (133)
carbon dioxide 304
(31)
argon
150.9 (-122)
Critical pressure: pressure needed to liquefy gas
at critical temperature
Substance at Tc and Pc - supercritical fluid
Vapour Pressure
Every liquid in a closed container gives off
vapour until a certain pressure is reached
- the liquid’s vapour pressure.
 The vapour pressure of a liquid increases
with increase in temperature.
 We can explain these facts using the
kinetic theory
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Volatile
A liquid in an open container will
evaporate
 As vapour moves away, the liquid releases
more molecules into the vapour phase to
try to build up to the correct vapour
pressure
 Liquids with high vapour pressure
evaporate more quickly - they are volatile
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Boiling point: the temperature at
which vp = 760 torr
Boiling
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A liquid boils when the vp equals the atmospheric
pressure
Normal boiling point temp -> when vp is 760 torr.
We list normal bp values in textbooks
Actual boiling point -> then the liquid has a vp
equal to the external atmospheric pressure
Water boils at temperature lower than 100ºC atop
mountains - it never reaches 100ºC
Water boils at higher temp in pressure cooker
Phase diagrams
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A graphical way to show the equilibria
between different phases of a substance
Thing to look for:
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critical point
triple point ( three phases)
how bp (and mp) varies with pressure
Triple point is not pressure dependent - the
vapour has to be at the critical pressure!
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useful for thermometer calibration
Structure of Solids
Crystalline: atoms, ions, or molecules are
well ordered. Have a well-defined melting
point. Often the solid has regular shapes.
 Amorphous: no order to the particles.
Examples are glass and rubber. Have no
defined mp; they soften over a range of
temperatures (important for glass
blowing)
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Unit cell: crystal lattice
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In a brick wall there
is a repeating
pattern, as there is
with most wallpaper
In a crystalline solid
there is a repeating
pattern - the unit
cell. The unit cell
repeats to make the
crystal lattice
The seven unit cells
Three Cubic lattices
Close Packing
Another way of looking at it
Diamonds are a …..
Some ionic solidslattice decided by size & charge
From sea to shining ...
array of metal ions in a sea of electrons
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A sea of valence
electrons
Electrons not tightly
held but can move
Explains electrical
conduction
Also explains optical
properties - most
metals “shine”