GASES - Mrs. Isernhagen's Science Page · Kinetic Theory for ideal gases. 1. Particles of a gas are infinitely small. Explains effusion & compressibility. 2. Particles of a gas are
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GASES
Unit 1
Chapter 1
Motion of particles:
In solids the particles:
are moving relatively slowly.
have low kinetic energy
In liquids the particles:
molecules move faster.
have higher kinetic energy.
In gases, the particles:
move fastest,
have high kinetic energy.
Kinetic Theory Model of States
SolidParticles vibrate but
don’t “flow”. Strong
molecular
attractions keep
them in place.
LiquidParticles vibrate, rotate,
tumble and “flow”, but
cohesion (molecular
attraction) keeps them
close together.
GasParticles move freely through
container. The wide spacing
means molecular attraction is
negligible.
4
Particles can have 3 types of motion:
Vibrational kinetic energy (vibrating)
Rotational kinetic energy (tumbling)
Translational kinetic energy (flying around)
5
Kinetic Theory of s, l & g. When it is cold, molecules move slowly
In solids, they move so slowly that they are held in
place (only vibrational energy)
In liquids they move a bit faster, they can tumble
and flow, but they don’t escape from the
intermolecular attraction with other molecules
(mostly rotational energy, some vibration & translation)
In gases they move so quickly that can overcome
the intermolecular attractions and leave the
container (more translational energy, with a little bit
of rotation & vibration).
6
Plasma, the “Fourth State”
When strongly heated, or exposed to high
voltage or radiation, gas atoms may lose some
of their electrons. As they capture new
electrons, the atoms emit light—they glow. This
glowing, gas-like substance is called “plasma”
7
Properties of gases:Gases:
can be atoms or molecules
I Have No Bright Or Clever Friends
I2 H2 N2 Br2 O2 Cl2 F2 are all diatomic gases,
have mass,
no definite volume or shape,
are compressible & can expand.
Therefore properties of gasses can only be compared
under specific conditions.
Gas particles are very spaced out!
Find the properties of some common
gases: Finish for homework!
Read pages 38-43 from your textbook.
For N2, O2, CO2, radon & methane gas:
List their:
Abundance
General use (Do we breath it? Do plants use it etc)
Technological applications
For O3 how is it a useful and harmful gas?
Fun Gases (of no real importance)
Nitrous Oxide (N2O) AKA: Laughing gas, Happy gas, Nitro, NOS
Uses
anaesthetic in dentist offices, this sweet-smelling gas reduces pain sensitivity and causes euphoric sensations.
It is an excellent oxidizer, reigniting a glowing splint much like oxygen would.
It is used in racing where it is injected into the carburetor to temporarily increase an engine’s horsepower.
Sulfur Hexafluoride
One of the densest gases in common use. Fun with Sulfur hexafluoride
10
Matchthe gas with the problem it causes
Gas Problem
Carbon Dioxide Ozone layer depletion
CFCs Climate Change
Methane Toxic poisoning
Carbon monoxide Noxious smell
Sulfur dioxide Acid Rain
Last class: We talked about the motion of molecules:
Vibrational
Rotational
Translational
Gasses have mostly translational motion.
As you increase temperature their motion also
increases.
Increase in translational movement
= increase in velocity
Pressure in gases: A force applied over a unit of area.
For a gas, pressure results from gas molecules
colliding with the wall of its container.
Measured in Pascals (Pa) or kiloPascals (kPa)
Adding a gas:
Adds more gas molecules
More collisions
Increased net pressure
Ex. Double # of molecules =
double pressure
If container is not
strong enough
walls can rupture
Removed a Gas:
Removes gas molecules
Less collisions
Pressure decreasesIf container is not
strong enough
walls can collapse
Change Size of Container:
Decrease container size
Decreases space for molecules to move
Increases collisions
Increases pressure
Change Size of Container:
Increase container size
Increases space for molecules to move
Decreases collisions
Decreases pressure
Heating a Gas:
Gas molecules absorb heat
Molecules move more rapidly
Increase collisions
Increase pressure
Cooling a Gas:
Gas molecules release heat
Molecules move more slowly
Decrease collisions
Decrease pressure
Gases exert a pressure as they collide with the
walls of containers. The total pressure is
dependent on magnitude & quantity of collisions.
Concentration:
Add more gas Pressure
Remove gas Pressure
Container size:
decrease Pressure
increase Pressure
Temperature:
Increase Temp Pressure
Decrease Temp Pressure
Kinetic Molecular Theory (K.M.T):
1. A gas is composed of particles
2. Gas particles move rapidly & are in constant
random motion
3. All collisions are perfectly elastic
4. Kinetic energy is proportional to
temperature
Textbook questions
Do Page 62 # 3,4,5,6,
To be done for next class
Kinetic energy & temperature: As temperature increases molecules move faster
& have a greater KE.
Not all molecules are moving at the same speed.
The KE of moving objects is expressed by:
This shows that the KE of molecules is
dependent on both their mass & velocity.
2
2
1mvEk
“Slow”
molecules
The range of kinetic energies can be
represented as a “bell curve.”
Maxwell’s Velocity Distribution Curve.
Increasing kinetic energyAverage
kinetic energy
Incre
asin
g #
mole
cule
s
Most molecules
mo
de
me
an
“Average”
molecules
The mean & mode can
help establish “average”
molecules
“Fast”
Molecules
Distribution of Particles Around Average
Kinetic Energies.
Kinetic Energy of molecules
(proportional to velocity of molecules)
Num
be
r o
f m
ole
cu
les
Ave
rag
e k
ine
tic
en
erg
y o
f m
ole
cu
les
Ave
rag
e k
ine
tic
en
erg
y o
f w
arm
er
mo
lec
ule
s
Faster
than
average molecules
Slower
than
average moleculesAve
rag
e k
ine
tic
en
erg
y o
f c
old
er
mo
lec
ule
sConclusion:
As temperature increases.
Curve broadens.
Average KE increases.
Two different gases at the same temperature will:
Have the same AVERAGE EK
Lighter molecules will move faster.
Heavier molecules will move slower.
Fun Fact
The average speed of oxygen molecules at 20°C is 1656km/h.
At that speed an oxygen molecule could travel from Montreal to Vancouver in three hours…If it travelled in a straight line.
= ½ mv2
Observing gases As scientists observed gases, they saw
mathematical relationships that very closely, but
not perfectly, described the behaviour of many
gases.
They have developed theories & mathematical
laws that describe a hypothetical gas, called
“ideal gas.”
It is an approximation that helps us model and
predict the behavior of real gases.
Kinetic Theory for ideal gases.1. Particles of a gas are infinitely small.
Explains effusion & compressibility.
2. Particles of a gas are in constant motion, and move in straight lines.
Until they run into another particle or wall.
Explains diffusion.
3. The particles do not attract or repel each other.Explains why gases expand to fill a space.
4. The average kinetic energy of the particles are proportional to the absolute temperature.
Explains observed changes in pressure.
28
Fun fact Each air molecule has about ten billion collisions per second
Thomas Graham (1805-1869) And in my spare time
I invented dialysis,
which has saved the
lives of thousands of
kidney patients
m = mass (kg)
v = velocity (m/s)
Graham studied the speed of diffusion &
effusion.
Diffusion is when gas molecules spread throughout
a container until they are evenly distributed
Effusion is when gas molecules pass through tiny
opening in container.
He derived his law from Ek = ½ mv2
we will use Ek= ½Mv2
M = molar mass (g/mol)
Write this!
Graham’s Law
Rate of diffusion of a gas is inversely related to the square root of its molar mass.
The equation shows the ratio of Gas 1’s speed to Gas 2’s speed.
v = velocity M = molar mass
(Leave a space for variants)
1
2
M
M
2
1
v
v
1
2
2
1
M
M
v
v Same
as
Ex. 1
Determine the relative rate of diffusion for
krypton and bromine.
1.381
Ans: Kr diffuses 1.381 times faster than Br2.
Kr
Br
Br
Kr
M
M
v
v2
2
1
2
2
1
M
M
v
v
g/mol83.80
g/mol159.80
Relative rate means find the ratio “v1/v2”!
C. Johannesson
Ex. 2Oxygen gas molecules have an average speed of 12.3 m/s at a given temp and pressure. What is the average speed of hydrogen molecules under the same conditions?
2
2
2
2
H
O
O
H
M
M
v
v
g/mol 2.02
g/mol32.00
m/s 12.3
vH2
3.980m/s 12.3
vH2
m/s49.0 vH 2
1
2
2
1
M
M
v
v
Derive an equation that allows
you to solve for.
Distance, if time is kept constant
Time, if distance is kept constant
Graham’s Law Version #2, Effusion Time
It can be easier to measure the time it takes for a gas to effuse
completely, rather than the speed.
2
1
2
1
M
M
t
t
Add the equations!
It can be useful to know the distance that a gas would spread
to in a certain amount of time.
1
2
2
1
M
M
d
d
C. Johannesson
Ex. 3
An unknown gas diffuses 4.0 times faster than O2.
Find its molar mass.
XM
g/mol32.00 16
X
O
O
X
M
M
v
v2
2
XM
g/mol32.00 4.0
2
2
16
g/mol32.00 M X g/mol2.0
The ratio “v1/v2” is 4.0.
Square both
sides to get rid
of the square
root sign.1
2
2
1
M
M
v
v
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