Liquification: KMT and condensed phases
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Transcript Liquification: KMT and condensed phases
P1V1 P2V2
n1T1 n 2T2
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.
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
KE depends on mass (m) and speed (u)
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KE mu 2
2
Temperature (in K) proportional
to average molecular KE
At higher T, average speed higher
At lower T, average speed lower
At T = 0, speed = 0
(molecules stop moving)
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1
KE mu2
2
Different gases at same temperature
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
What if there are more molecules?
More collisions → higher pressure
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Molecules colliding with container → gas pressure
What if the container is smaller?
More collisions → higher pressure
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Molecules colliding with container → gas pressure
What if the molecules are moving faster?
Harder, more frequent collisions → higher pressure
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Moving molecules fill
the container
Light molecules escape faster,
heavy molecules more slowly
Large spaces between molecules allow gas to be
compressed
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?
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
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)
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.
At same temp. (same KE), molecules are gases,
liquids, solids.
This suggests that some molecules have stronger
attractions between molecules.
How do you judge the strength of the attractive
forces between molecules?
Look at the boiling point.
Low Boiling point = weak intermolecular forces
(molecules are not sticky)
High Boiling point = strong intermolecular forces
(molecules are sticky)
Molecules of similar structure:
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
Predict which would have a higher boiling
point, and why?
Na or K
F2 or Br2
As you heat a solid, the temp. increases until
the solid melts.
Temp. will remain constant until solid melts
completely.
You observe the same pattern when a gas is
cooled until it changes into a solid
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
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)
Heat of Fusion- change in energy when a solid
substance melts
Heat of Vaporization- change in energy when
liquid substance vaporizes (evaporates)
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?
A quick look into the past…
At any given temp. molecules have the same avg. KE
Light molecules = fast
Heavy molecules = slow
Having the same avg. KE means some molecules are
moving faster than avg., and some slower.
Consider if you will…
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
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
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.
Changing temp. changes equilibrium, and vapor
pressure remains constant
at any given temp.
Imagine this…
Open container and heat it.
KE increases, more molecules evaporate
Bubbles of water form throughout liquid, rise to
surface and break… boiling begins
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