Transcript Liquification: KMT and condensed phases

P1V1 P2V2

n1T1 n 2T2
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or
PV  nRT
R = 0.08206 atm L/mol K
Ideal gas equation predicts gas behavior
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Ideal gas law describes what gases do, but not why.
Kinetic Molecular Theory of Gases (KMT): model
that explains gas behavior.
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developed in mid-1800s
based on concept of an ideal or perfect gas
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Tiny particles in constant, random, straight-line
motion
Molecules collide w/ each other & w/ walls of
container
Gas molecules are points; gas volume is empty
space between molecules
Molecules independent of each other (no attractive
or repulsive forces between them).
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Moving molecule has kinetic energy
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KE depends on mass (m) and speed (u)
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KE  mu 2
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Temperature (in K) proportional
to average molecular KE
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At higher T, average speed higher
At lower T, average speed lower
At T = 0, speed = 0
(molecules stop moving)
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KE  mu2
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Different gases at same temperature
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All have same average KE (same temperature)
Heavier gases are slower; lighter gases are faster
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Molecules colliding with container → gas pressure
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What if there are more molecules?
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More collisions → higher pressure
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Molecules colliding with container → gas pressure
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What if the container is smaller?
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More collisions → higher pressure
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Molecules colliding with container → gas pressure
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What if the molecules are moving faster?
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Harder, more frequent collisions → higher pressure
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Moving molecules fill
the container
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Light molecules escape faster,
heavy molecules more slowly
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Large spaces between molecules allow gas to be
compressed
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Ideal gas remains a gas when cooled, even to 0 K
Real gases condense to liquid state when cooled
Pressure
(atm)
0
0
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Ideal gas pressure
decreases steadily &
becomes zero at
absolute zero
Real gas pressure decreases
abruptly to zero when gas
condenses to liquid
Temperature
(K)
How do we explain condensation?
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KMT ignores attractions between gas
molecules
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Gas molecules are too far apart & too fast for
attractions to act
BUT . . . attractive forces do
exist between all molecules!
At low enough T, attractions overcome kinetic
energy & molecules stick together to form a
liquid
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For every substance there are 2 opposing
tendencies:
Kinetic energy of the molecules, which tend to make
them move apart from each other (gas-like)
 Attraction between molecules, which tends to make
them stick together (liquid-like)
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At any given temp., kinetic energy is the same
for all molecules, so the attractive forces
between molecules determines whether
something is a liquid, solid, or gas.
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At same temp. (same KE), molecules are gases,
liquids, solids.
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This suggests that some molecules have stronger
attractions between molecules.
How do you judge the strength of the attractive
forces between molecules?
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Look at the boiling point.
Low Boiling point = weak intermolecular forces
(molecules are not sticky)
High Boiling point = strong intermolecular forces
(molecules are sticky)
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Molecules of similar structure:
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Boiling points increase as molar mass increases
Suggests that intermolecular attractions increase as
molecular mass increase
Formula Molar
Mass
BP
Formula Molar
Mass
BP
F2
38 g/mol
85 K
CH4
Cl2
71 g/mol
239 K C2H6
30 g/mol 184 K
Br2
160 g/mol
332 K C3H8
44 g/mol 231 K
I2
254 g/mol
457 K C4H10
58 g/mol 273 K
16 g/mol 111 K
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Predict which would have a higher boiling
point, and why?
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Na or K
F2 or Br2
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As you heat a solid, the temp. increases until
the solid melts.
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Temp. will remain constant until solid melts
completely.
You observe the same pattern when a gas is
cooled until it changes into a solid
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Temp. will remain constant until gas changes into
solid
Temperature
Add energy
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Temperature
(g)
boiling
condensing
boiling
point
(l)
melting/freezing
point
melting
freezing
(s)
Add energy
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Melting & boiling are ENDOTHERMIC
Freezing & condensing are EXOTHERMIC
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Changing temperature changes KE (#1, 3, 5)
Changing state changes potential energy (#2, 4)
energy to melt or freeze = heat of fusion (∆Hfusion)
 energy to vaporize or condense = heat of vaporization
(∆Hvaporization)
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The O–H bond in water is very
polar, and the atoms are very
small
 The dipoles are close together,
so their attraction is very
strong
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An H atom is covalently bonded (red-white) to its
own O and weakly bonded (dotted line) to the
neighboring O
This weak bond to a neighboring O is called a
hydrogen bond
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Hydrogen bonding occurs only between
molecules containing N–H, O–H, and F–H
bonds
Hydrogen bonding is much stronger than
ordinary intermolecular attractions ⇒ very
high boiling points for their mass
Hydrogen bonds are not as strong as covalent
bonds (15-40 kJ/mol, vs >150 kJ/mol)
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Heat of Fusion- change in energy when a solid
substance melts
Heat of Vaporization- change in energy when
liquid substance vaporizes (evaporates)
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Given that the heat of fusion (ice) is 6.0 kj/mol,
how much energy is needed to melt an ice cube
with a mass of 42 g?
Given that the heat of vaporization (water) is
41 kj/mol, how much energy is need to convert
24 g of water into steam?
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A quick look into the past…
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At any given temp. molecules have the same avg. KE
 Light molecules = fast
 Heavy molecules = slow
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Having the same avg. KE means some molecules are
moving faster than avg., and some slower.
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Consider if you will…
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A container of water at 25oC (room temp.)
All the water molecules have the same avg. KE
Water is not hot, but a few fast molecules will leave
and become water vapor.
Vapor molecules exert pressure = vapor pressure
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Now, consider this…
Water molecules can escape from open container,
but not a closed container.
 Energy is transferred by collisions w/ other vapor
molecules.
 Some vapor molecules will be slow enough to return
to liquid.
 Eventually, rate of evaporation = rate of condensing
 This is called DYNAMIC EQUILIBRIUM
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Molecules continue to evaporate and condense.
# of vapor molecules and vapor pressure are
constant.
A liquid in a closed container has a constant
equilibrium vapor pressure.
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Changing temp. changes equilibrium, and vapor
pressure remains constant
at any given temp.
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Imagine this…
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Open container and heat it.
KE increases, more molecules evaporate
Bubbles of water form throughout liquid, rise to
surface and break… boiling begins
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As temp. increases, vapor pressure increases
A liquid boils when the vapor pressure matches
the external pressure.
In an open container, external pressure is
atmospheric pressure