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Unit 1 (formerly Module 2)
Gases and Their Applications
Lesson 2-1
About Gases
2
Gas is one of the three main states of matter
Gas particles may be atoms or molecules, depending on the type of substance (ie, element or compound)
Gas particles have much more space between them than liquids or solids.
Gases are said to be an expanded form of matter, solids and liquids are condensed forms of matter.
3
R
General Properties of a Gas
Gases do have mass (although it is
sometimes difficult to measure).
Gases have no definite volume,
Gases have no definite shape.
Gases are compressible, meaning they can
be squeezed into smaller containers, or can
expand to fill larger containers.
– Because gases compress, the density of gases
can only be compared under specific conditions.
4
Some Important Gases Oxygen (O2): clear, breathable, supports combustion.
Ozone (O3): poisonous, unstable form of oxygen
Nitrogen (N2): clear, low activity, most abundant gas in the Earth’s atmosphere.
Hydrogen (H2): clear, lighter than air, flammable/explosive
Carbon dioxide (CO2): clear, but turns limewater cloudy. Does not support respiration but low toxicity. Heavier than air. Largely responsible for the greenhouse effect (global warming)
Sulphur dioxide (SO2): smelly gas. When it combines with oxygen and water vapour it can form H2SO4, responsible for acid rain.
5
Some Important Gases Carbon monoxide (CO): clear, colourless, but very
toxic. It destroys the ability of blood to carry oxygen. About the same density as air.
Ammonia (NH3): toxic, strong smell, refrigerant . Very soluble in water, forms a basic solution called ammonia-water (NH4OH) which is found in some cleaners.
Freon® or CFC: Non-toxic (safe-to-inhale in moderation) refrigerant used in air-conditioners & freezers. Freon may catalyze ozone breakdown. The original Freon formula is now banned, but low chlorine versions are still in use.
Methane (CH4): flammable gas, slightly lighter than air, produced by decomposition. Found in natural gas. Methane is also a “greenhouse” gas.
Helium (He): inert, lighter than air. Used in balloons and in diver’s breathing mixtures.
6
Acetylene (C2H2): AKA ethene, it is used as a fuel in
welding, lanterns and other devices.
Propane (C3H8): used as a fuel in barbecues, stoves,
lanterns and other devices.
Radon (Rn): A noble gas that is usually radioactive. It is
heavier than air, and sometimes found in poorly ventilated
basements.
Neon (Ne) and Xenon (Xe): Noble gases found in
fluorescent light tubes, and as insulators inside windows.
They glow more brightly than other gases when electrons
pass through them. Neon is slightly lighter than air, Xenon
is quite a bit heavier.
Compressed Air (78% N2, 21% O2): Not actually a pure
gas, but a gas mixture that acts much like a pure gas. It is
used by scuba divers (at shallow depths), and to run
pneumatic tools, and for producing foam materials.
7
Fun Gases
(of no real importance)
Nitrous Oxide (N2O) – AKA: Laughing gas, Happy gas, Nitro, NOS
– Once used as an 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
8
Match the gas with the problem it causes
Gas Problem
Carbon Dioxide Ozone layer depletion
CFCs Global Warming
Methane Toxic poisoning
Carbon monoxide Noxious smell
Sulfur dioxide Acid Rain
Next slide: Summary
9
Some Gases Classified by
Relative Density Low Density gases Neutral Density Gases High Density gases
“lighter than air”<25 g/mol “similar to air” 29±4 g/mol “Denser than air” (>34 g/mol)
Testable Property*:
Balloon will float in air
Balloon drops slowly
through air
Balloon drops quickly through
air
Examples:
Hydrogen (H2) 2
Helium (He) 4
Methane (CH4) 16
Ammonia (NH3) 17
Neon (Ne) 20
Hydrogen Fluoride (HF) 21
Examples:
“Cyanide“ (HCN) 27
Acetylene (C2H4) 28
Nitrogen (N2) 28
Carbon monoxide 28
Ethane (C2H6) 30
Oxygen (O2) 32
Examples:
Fluorine (F2) 38
Argon (Ar) 40
Carbon dioxide (CO2) 44
Propane (C3H8) 44
Butane (C4H10) 58
Sulphur Hexafluoride (SF6) 146
*balloon test: Fill a large, lightweight balloon with the gas, then release it from a height of about 1.8 m in a room with still air. If the gas is lighter than air the balloon will float upwards. If it is close to air, the balloon will fall very slowly. If the gas is heavier than air, the balloon will fall quickly.
Some Gases Classified by
Chemical Properties Combustible gases
(combustion /explosion)
Oxidizing (reactive) Gases
(support combustion)
Non-Reactive gases
(don’t support combustion)
Testable property:
Burning splint produces
“pop”
Testable property:
Glowing splint reignites,
burning splint grows brighter
Testable property:
Burning splint is
extinguished, glowing
splint is dimmed
Other properties:
Useful as fuels
Other properties:
Cause metals and some
other materials to corrode or
oxidize. Can improve
combustion.
Other properties:
Can be used to preserve
foods by slowing
oxidation
Examples:
Hydrogen (H2)
Methane (CH4)
Propane (C3H8)
Acetylene (C2H4)
Examples:
Oxygen (O2)
Fluorine (F2)
Chlorine (Cl2)
Nitrous Oxide (NO2)
Examples:
Carbon dioxide (CO2)
Nitrogen (N2)
Argon (Ar)
Helium (He)
Textbook Assignments
Read Chapter 1: pp. 37 to 50
Do the exercises on pages 51 and 52
– Questions # 1 to 22
12
Summary:
• Know the properties of gases
• Know the features of some important gases, esp:
• Oxygen
• Hydrogen
• Carbon dioxide
• Know the environmental problems associated with some gases, eg.
• Carbon dioxide
• CFC’s
• Sulfur dioxide
13
Chapter 2
Physical Properties of Gases
Includes:
The Kinetic Theory
“Moving, moving, moving,
Keep those atoms moving...”
The Gas Laws.
“Jumping Jack Flash, It’s a gas, gas, gas...”
14
2.1 Kinetic Theory
• Overview:
The kinetic theory of gases (AKA. kinetic-
molecular theory) tries to explain the behavior
of gases, and to a lesser extent liquids and
solids, based on the concept of moving
particles or molecules.
The Kinetic Theory of Gases (AKA: The Kinetic Molecular Theory)
• The Kinetic Theory of Gases tries to explain the
similar behaviours of different gases based on
the movement of the particles that compose
them.
• “Kinetic” refers to motion. The idea is that gas
particles* are in constant motion.
* For simplicity, I usually call the gas particles “molecules”,
although in truth, they could include atoms or ions.
2.1
Page 54
16
The Particle Model Not in text
• The Kinetic Theory is part of the Particle
Model of matter, which includes the following
concepts:
– All matter is composed of particles (ions, atoms or
molecules) which are extremely small and have a
varying space between them, depending on their
state or phase.
– Particles of matter may attract or repel each other,
and the force of attraction or repulsion depends on
the distance that separates them.
– Particles of matter are always moving.
+ +
-
17
R
Kinetic Molecular Theory
And Temperature
• The absolute temperature of a gas (Kelvins) is
directly proportional to the average kinetic
energy of its molecules.
– In other words, when it is cold, molecules move
slowly and have lower kinetic energy.
– When the temperature increases, molecules speed
up and have more kinetic energy!
18
Particle Motion
and Phases of Matter • Recall that:
• In solids, the particles (molecules) are moving
relatively slowly. They have low kinetic energy
• In liquids, molecules move faster. They have
higher kinetic energy.
• In gases, the particles move fastest, and have high
kinetic energy.
• But, as we will find out later: • Heavy particles moving slowly can have the same
kinetic energy as light particles moving faster.
2.1.1
Page 54
19
R
Kinetic Theory Model of States
Solid Particles vibrate
but don’t “flow”. Strong molecular attractions keep them in place.
Liquid Particles vibrate, move
and “flow”, but cohesion (molecular attraction) keeps them close together.
Gas Particles move freely
through container. The wide spacing means molecular attraction is negligible.
20
Kinetic Motion of Particles
• Particles (ie. Molecules) can have 3 types
of motion, giving them kinetic energy
– Vibrational kinetic energy (vibrating)
– Rotational kinetic energy (tumbling)
– Translational kinetic energy (moving)
2.1.1
Page 55
21
Kinetic Theory and Solids & Liquids
• When it is cold, molecules move slowly
• In solids, they move so slowly that they are held
in place and just vibrate (only vibrational energy)
• In liquids they move a bit faster, and can tumble
and flow, but they don’t escape from the
attraction of other molecules (more rotational
energy, along with a little bit of vibration & translation)
• In gases they move so fast that they go
everywhere in their container (more translational
energy, with a little bit of rotation & vibration).
2.1.1
Page 56
22
Plasma, the “Fourth State” (extension material)
• 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”
23
Kinetic Theory and the Ideal Gas
• As scientists tried to understand how gas
particles relate to the properties of gases,
they saw mathematical relationships that
very closely, but not perfectly, described
the behaviour of many gases.
• They have developed theories and
mathematical laws that describe a
hypothetical gas, called “ideal gas.”
2.1.3
Page 61
24
• To make the physical laws (derived from
kinetic equations from physics) work, they
had to make assumptions about how
molecules work. • Four of these assumptions are listed on page 61 of
your textbook
• Other textbooks contain additional assumptions
associated with the kinetic theory.
2.1.3
Page 61
2
2
1: mvEequationkinetic k
25
Kinetic Theory Hypotheses
about an Ideal Gas
1. The particles of an ideal gas are infinitely small, so the size is negligible compared to the volume of the container holding the gas.
2. The particles of an ideal gas are in constant motion, and move in straight lines (until they collide with other particles)
3. The particles of an ideal gas do not exert any attraction or repulsion on each other.
4. The average kinetic energy of the particles is proportional to the absolute temperature.
2.1.3
Page 61
26
No Gas is Ideal
• Some of the assumptions on the previous
page are clearly not true. • Molecules do have a size (albeit very tiny)
• Particles do exert forces on each other (slightly)
• As a result, there is no such thing as a
perfectly “ideal gas” • However, the assumptions are very good
approximations of the real particle properties.
• Real gases behave in a manner very close to
“ideal gas”, in fact so close that we can usually
assume them to be ideal for the purposes of
calculations. 27
Other “Imaginary Features”
of Ideal Gas • An ideal gas would obey the gas
laws at all conditions of
temperature and pressure
• An ideal gas would never
condense into a liquid, nor
freeze into a solid.
• At absolute zero an ideal gas
would occupy no space at all.
2.1.3
Page 61
28
Please Notice:
• Not all molecules move at exactly the same speed. The kinetic theory is based on averages of a great many molecules. – Even if the molecules are identical and at a uniform
temperature, a FEW will be faster than the average, and a FEW will be slower.
– If there are two different types of molecules, the heavier ones will be slower than the light ones – ON THE AVERAGE! – but there can still be variations. That means SOME heavy molecules may be moving as fast as the slowest of the light ones.
• Temperature is based on the average (mean) kinetic energy of sextillions of individual molecules.
29
“Slow”
molecules
The range of kinetic energies can be
represented as a sort of “bell curve.”
Maxwell’s Velocity Distribution Curve.
Increasing kinetic energy
Average
kinetic energy
Incre
asin
g #
mole
cule
s
Most molecules
mode
mean
“Average”
molecules
The mean & mode can help establish
“average” molecules
“Fast”
Molecules
30
So, Given two different gases at
the same temperature…
What is the same about them?
• The AVERAGE kinetic energy is the same. • Not the velocity of individual molecules
• Not the mass of individual molecules.
• In fact, the lighter molecules will move faster
• Ek = mv2 kinetic energy of molecules
2
So, kinetic energy depends on both the speed (v) and on the mass (m) of the molecules.
31
Distribution of Particles Around
Average Kinetic Energies.
Kinetic Energy of molecules
(proportional to velocity of molecules)
Num
ber
of
mole
cule
s
Averag
e k
ineti
c e
nerg
y o
f m
ole
cu
les
Averag
e k
ineti
c e
nerg
y o
f w
arm
er m
ole
cu
les
Faster
than
average molecules
Slower
than
average molecules
32
Averag
e k
ineti
c e
nerg
y o
f co
lder m
ole
cu
les
Kinetic Theory Trivia • 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.
• Each air molecule has about 1010 (ten billion) collisions per second
• 10 billion collisions every second means they bounce around a lot!
• The number of oxygen molecules in a classroom is about:
• 722 400 000 000 000 000 000 000 000
– that’s more than there are stars in the universe!
• The average distance air molecules travel between collisions is about 60nm.
– 0.00000006m is about the width of a virus.
33
Videos
• Kinetic Molecular Basketball
– http://www.youtube.com/watch?v=t-Iz414g-ro&NR=1
• Average Kinetic Energies
– http://www.youtube.com/watch?v=UNn_trajMFo&NR=1
• Thermo-chemistry lecture on kinetics
34
Assignments
• Read pages 53 to 61
• Do Page 62 # 1-11
Chapter 2.2
• Behaviors of Gases
– Compressibility
– Expansion
– Diffusion and Effusion
– Graham’s Law
37
• 2.2.1 Compressibility:
– Because the distances between particles in a
gas is relatively large, gases can be squeezed
into a smaller volume.
– Compressibility makes it possible to store
large amounts of a gas compressed into small
tanks
• 2.2.2 Expansion:
– Gases will expand to fill any container they
occupy, due to the random motion of the
molecules.
2.2.3 Diffusion
Diffusion is the tendency for molecules to
move from areas of high concentration to
areas of lower concentration, until the
concentration is uniform. They do this
because of the random motion of the
molecules.
Effusion is the same process, but with the
molecules passing through a small hole or
barrier
Next slide: 38
Rate of Diffusion or Effusion
It has long been
known that lighter
molecules tend to
diffuse faster than
heavy ones, since
their average
velocity is higher,
but how much
faster?
39
heavy particle
light particle
Graham’s Law
Thomas Graham (c. 1840)
studied effusion (a type of
diffusion through a small hole)
and proposed the following law:
“The rate of diffusion of a gas is
inversely proportional to the
square root of its molar mass.”
In other words, light gas particles
will diffuse faster than heavy gas
molecules, and there is a math
formula to calculate how much
faster.
Next slide: Example
1
2
2
1
M
M
v
v
40
Where: v1= rate of gas 1
v2= rate of gas 2
M1= molar mass of gas1
M2=molar mass of gas 2
Internet demo of effusion
Graham’s Law Version #1, based on Effusion Rate
• The relationship between the rate of effusion or diffusion
and the molar masses is:
1
2
2
1
M
M
v
v
Where: v1 is the rate of diffusion of gas 1, in any appropriate rate units*
v2 is the rate of diffusion of gas 2, in the same units as gas 1
M1 is the molar mass of gas 1
M2 is the molar mass of gas 2
*Rate units must be an amount over a time for effusion (eg: mL/s or L/min), or a distance over a time for diffusion (eg: cm/min or mm/s)
Note: See the inversion of the 1 and 2 in the 2nd ratio!
Thomas Graham (1805-1869)
• Graham derived his law by treating
gases as ideal, and applying the kinetic
energy formula to them.
• Ek = ½ mv2
• All gases have the same kinetic energy
at the same temperature,
• Therefore, mv2 for the first gas = mv2
for the second gas: m1v12 = m2v2
2.
• A bit of algebra then gave him his
famous law.
And in my spare time I
invented dialysis, which has saved the
lives of thousands of
kidney patients
Graham’s Law Version #2, Based on Effusion Time
• Sometimes it’s easier to measure the time it takes for a
gas to effuse completely, rather than the rate. Graham’s
law can be changed for this, but the relationship between
time and molar mass is direct as the square root:
2
1
2
1
M
M
t
t
Where: t1 is the time it takes for the first gas to effuse completely.
t2 is the time it takes for an equal volume of the 2nd gas to effuse
M1 is the molar mass of the first gas
M2 is the molar mass of the second gas.
Note: In this variant law, the relationship is not inverted!
Example of Graham’s Law: How much faster does He diffuse than N2?
Nitrogen (N2) has a molar
mass of 28.0 g/mol
Helium (He) has a molar
mass of 4.0 g/mol
The difference between
their diffusion rates is:
Notice the reversal of
order!
So helium diffuses 2.6
times faster than nitrogen
He
N
N
He
M
M
v
v 2
2
6.22
3.5
/4
/28
molg
molg
MN2=2x14.0=28 g/mol
Next slide: 2.3 Pressure of Gases
MHe=1x4.0=4 g/mol
44
Assignments
Read pages 63 to 67
Do Questions 1 to 10 on
page 68
Chapter 2.3
• Pressure of Gases
– What is Pressure
– Atmospheric Pressure
– Measuring Pressure
100 km < 0.003 kPa
40 km 1 kPa
20 km 6 kPa
10 km 25 kPa
5 km 55 kPa
0 km 101 kPa
Mt Everest 31 kPa
46
Highest Jet 4 kPa
Edge of Space
X15 (1963) Spaceship 1 (2006) Outer Space (immeasurable)
Mr. Smith
Pressure
• Pressure is the force exerted by a gas on
a surface. • The surface that we measure the pressure on is
usually the inside of the gas’s container.
• Pressure and the Kinetic Theory • Gas pressure is caused by billions of particles
moving randomly, and striking the sides of the
container.
• Pressure Formula:
Pressure = force divided by area
A
FP
47
Atmospheric Pressure
• This is the force of a 100 km high
column of air pushing down on us.
• Standard atmospheric pressure is • 1.00 atm (atmosphere), or
• 101.3 kPa (kilopascals), or
• 760 Torr (mmHg), or
• 14.7 psi (pounds per square inch)
• Pressure varies with: • Altitude. (lower at high altitude)
• Weather conditions. (lower on cloudy days)
48
Pressure conversions
)(
)(
)(
)(
wantedunitsSP
wantedunitsP
givenunitsSP
givenunitsP
Example 1: convert 540 mmHg to kilopascals
kPa
P
mmHg
mmHg
3.101760
540=72.0 kPa
Example 2: convert 155 kPa to atmospheres
atm
P
kPa
kPa
00.13.101
155=1.53 atm
SP 1.00 atm
760 mmHg
760 Torr
101.3 kPa
14.7 psi
1013 mB
29.9 inHg
Divide
Measuring Pressure
• Barometer: measures atmospheric
pressure.
– Two types:
• Mercury Barometer
• Aneroid Barometer
• Manometer: measures pressure in a
container (AKA. Pressure guage) • Dial Type: Similar to an aneroid barometer
• U-Tube: Similar to a mercury barometer
• Piston type: used in “tire guage” 50
• A tube at least 800 mm long is filled with
mercury (the densest liquid) and inverted
over a dish that contains mercury.
• The mercury column will fall until the air
pressure can support the mercury.
• On a sunny day at sea level, the air
pressure will support a column of mercury
760 mm high.
• The column will rise and fall slightly as the
weather changes.
• Mercury barometers are very accurate,
but have lost popularity due to the toxicity
of mercury.
the Mercury Barometer
51
The Aneroid Barometer
• In an aneroid barometer,
a chamber containing a
partial vacuum will
expand and contract in
response to changes in
air pressure
• A system of levers and
springs converts this into
the movement of a dial.
• Manometers work much like
barometers, but instead of
measuring atmospheric
pressure, they measure the
pressure difference between
the inside and outside of a
container.
• Like barometers they come in
mercury and aneroid types.
There is also a cheaper
“piston” type used in tire
gauges, but not in science.
U-tube manometer Pressure gauge
(mercury manometer) (aneroid) You Tube manometer
Manometers (Pressure Gauges)
Tire gauge
(piston manometer)
Reading U-tube manometers • When reading a mercury U-
tube manometer, you
measure the difference in
the heights of the two
columns of mercury.
• If the tube is “closed” then
the height (h) is the gas
pressure in mmHg.
P(mmHg)=h(mmHg)
• If the tube is “open” and h is
positive (the pressure you
are measuring is greater
than the atmosphere) then
you must add atmospheric
pressure in mmHg.
Pgas(mmHg) = Patm(mmHg)+h(mm)
Must be in
mmHg, not cm or kPa!
Atm. pressure
After you finish, you can convert your answer to kPa, or atm. Or whatever.
Manometer Examples on a day when the air pressure is 763mmHg (101.7 kPa)
Closed tube: Pgas(mm Hg)=h (mm Hg)
Pgas = h = 4 cm = 40 mm Hg Pgas =
kPakPa
Hgmm
Hgmm3.53.101
760
40
Open: Pgas(mmHg)=P atm(mmHg) +h (mmHg)
Pgas = 763 + 60mm Hg =823 mm Hg
Pgas =
kPakPaHgmm
Hgmm7.1093.101
760
823
Open: Pgas(mmHg)=P atm(mmHg) -h (mmHg)
Pgas = 763 - 60mm Hg =703 mm Hg
Pgas =
kPakPa
mmHg
mmHg7.933.101
760
703
4 cm
6
9
Assignments
• Read pages 69 to 73.
• Do Page 74, Questions 1 to 4.
Chapter 2.4
• The Simple Gas Laws – Boyle’s Law Relates volume & pressure
– Charles’ Law Relates volume & temperature
– Gay-Lussac’s Law Relates pressure & temperature
– Avogadro’s Law Relates to the number of moles
• Other Simple Laws that are a Gas: – Clarke’s Laws Relates possible and impossible
– Murphy’s Law Anything that can go wrong will
– Cole’s Law Relates thinly sliced cabbage
to vinegar.
57
Clarke’s Laws of the impossible*
Clarke’s 1st Law: If an elderly and respected
science teacher (like me) tells you that
something is possible, he is probably right. If he
tells you something is impossible, he’s almost
certainly wrong.
Clarke’s 2nd Law: The only way to find the limits
to what is possible is to go beyond them.
Clarkes 3rd Law: Any sufficiently advanced
technology is indistinguishable from magic.
*these are slightly paraphrased, I quote them from memory. They were
developed by science fiction writer Arthur C. Clarke
Lesson 2.4.1
Boyle’s Law Robert Boyle (1662)
“For a given mass of gas at a
constant temperature, the volume
varies inversely with pressure.”
For Pressure and Volume
VP
1
59
Robert Boyle
Born: 25 January 1627
Lismore, County Waterford, Ireland
Died 31 December 1691 (aged 64)
London, England
Fields: Physics, chemistry; Known for
Boyle's Law. Considered to be the
founder of modern chemistry
Influences: Robert Carew, Galileo
Galilei, Otto von Guericke, Francis
Bacon
Influenced: Dalton, Lavoisier, Charles,
Gay-Lussack, Avogadro.
Notable awards: Fellow of the Royal
Society
60
Pressure Gas pressure is the force placed on the sides of a
container by the gas it holds
Pressure is caused by the collision of trillions of gas particles against the sides of the container
Pressure can be measured many ways Standard Pressure
Atmospheres (atm) 1 atm
Kilopascals (kPa)or(N/m2) 101.3 kPa = 101.3 N/m2
Millibars (mB) 1013 mB
Torr (torr) or mm mercury 760 torr = 760 mmHg
Centimetres of mercury 76 cmHg
Inches of mercury (inHg) 29.9 inHg (USA only)
Pounds per sq. in (psi) 14.7 psi (USA only)
61
Next slide: Air in Syringe
Example of Boyle’s Law:
Air trapped in a syringe
If some air is left in a syringe, and the needle removed and sealed, you can measure the amount of force needed to compress the gas to a smaller volume.
Next slide: Inside syringe 62
Inside the syringe…
The harder you press, the smaller the volume of air becomes. Increasing the pressure makes the volume smaller!
The original pressure was low, the volume was large. The new pressure is higher, so the volume is small. Click Here for an internet demo using
psi (pounds per square inch) instead of kilopascals (1kPa=0.145psi)
Next slide: PV
low
high
63
This means that:
As the volume of a contained gas decreases, the pressure increases
As the volume of a contained gas increases, the pressure decreases
This assumes that:
no more gas enters or leaves the container, and
that the temperature remains constant.
The mathematical formula for this is given on the next slide
Next slide: Example 64
Boyle’s Law Relating Pressure and Volume of a Contained Gas
• By changing the shape of a gas container, such
as a piston cylinder, you can compress or
expand the gas. This will change the pressure
as follows:
2211 VPVP
Where: P1 is the pressure* of the gas before the container changes shape.
P2 is the pressure after, in the same units as P1.
V1 is the volume of the gas before the container changes, in L or mL
V2 is the volume of the gas after, in the same units as V1
*appropriate pressure units include: kPa, mmHg, atm. Usable, but inappropriate units include psi, inHg.
Example 1
You have 30 mL of air in a syringe at 100 kPa.
If you squeeze the syringe so that the air
occupies only 10 mL, what will the pressure
inside the syringe be?
P1 V1 = P2 V2, so..
100 kPa 30 mL = ? kPa 10 mL
3000 mL·kPa 10 mL = 300 kPa
The pressure inside the syringe will be 300 kPa
Next slide: Graph of Boyle’s Law 66
Graph of Boyle’s Law The Pressure-Volume Relationship
Pressure (kPa)
Volu
me (
L)
100 200 300 400 500 600 700 800
1
2
3
4
5
6
7
8
Boyle’s Law produces an inverse relationship graph.
100 x 8 = 800
200 x 4 = 800
400 x 2 = 800
800 x 1 = 800
P(kpa) x V(L)
Next slide: Real Life Data
300 x 2.66 = 800
500 x 1.6 = 800
600 x 1.33 = 800
700 x 1.14 = 800
67
Example 2: Real Life Data
2 4 6 8 10 12 14 16 18
5
10
15
20
25
30
35
40
In an experiment Mr. Taylor and Tracy put weights onto a syringe of air.
At the beginning, Mr. Taylor calculated the equivalent of 4 kgf of atmospheric pressure were exerted on the syringe.
0+4= 4kg : 29 mL (116)
2+4= 6kg : 20 mL (120)
4+4=8kg : 15 mL (120)
6+4=10kg: 12 mL (120)
8+4=12kg: 10.5 mL (126)
Next slide: Boyle’s Law Experiment or skip to: Lesson 2.3 Charles’ Law: 68
Summary: Boyle’s law
• The volume of a gas is
inversely proportional to its
pressure
• Formula: P1V1=P2V2
• Graph: Boyle’s law is usually
represented by an inverse
relationship graph (a curve)
Volu
me (
L)
Pressure (kPa)
VP
1
P1V1=P2V2
69
70
Assignments on Boyle’s Law
• Read pages 75 to 79
• Do questions 1 to 10 on page 97
Lesson 2.4.2
Charles’ Law
The Relationship between Temperature
and Volume.
“Volume varies directly with Temperature”
Next slide: Jacques Charles
TV
72
Jacques Charles (1787)
“The volume of a fixed mass of gas is directly proportional to its temperature (in kelvins) if the pressure on the gas is kept constant” This assumes that the container can expand, so that the pressure of the gas will not rise.
Next slide: The Mathematical formula for this law
Born: November 12, 1746 (1746-11-12) Beaugency, Orléanais
Died: April 7, 1823 (1823-04-08) (aged 76), Paris
Nationality: France
Fields: physics, mathematics, hot air ballooning
Institutions: Conservatoire des Arts et Métiers
Charles’ Law Relating Volume and Temperature of a Gas
• If you place a gas in an expandable container,
such as a piston or balloon, as you heat the gas
its volume will increase, as you cool it the
volume will decrease.
2
2
1
1
T
V
T
V
Where: T1 is Temperature of the gas before it is heated, in kelvins.
T2 is Temperature of the gas after it is heated, in kelvins
V1 is the volume of the gas before it was heated, in L or mL
V2 is the volume of the gas after it was heated, in the same units.
Charles Law Evidence
Charles used cylinders and pistons to study and graph the expansion of gases in response to heat.
See the next two slides for diagrams of his apparatus and graphs.
Lord Kelvin (William Thompson) used one of Charles’ graphs to discover the value of absolute zero.
Next slide: Diagram of Cylinder & Piston 75
Charles Law Example
Piston
Cylinder
Trapped Gas
Next slide: Graph of Charles’ Law
Click Here for a simulated internet experiment
76
Graph of Charles Law
0°C 100°C 200°C
150°C 50°C 250°C
1L
2L
3L
4L
5L
6L
-250°C -200°C -150°C -100°C -50°C
-273.15°C
Expansion of most real gases
273°C
Next slide: Example
Liquid state
Solid state
Charles discovered
the direct relationship
Lord Kelvin
traced it back to absolute
zero.
-273.15°C is called absolute zero. It is the coldest possible temperature. At absolute zero, molecules stop moving and even vibrating. Since temperature is based on the average kinetic energy of molecules, temperature cannot be said to exist if there is no kinetic energy (movement)
William Thompson, The right honourable
Lord Kelvin 1st Baron of Largs
1824-1907
Born in Belfast Ireland Died in Largs, Scotland
Worked at the University of Glasgow
Experimented in Thermodynamics.
"There is nothing new to
be discovered in physics
now. All that remains is
more and more precise
measurement"
Kelvin’s Scale
To convert from Celsius to Kelvin, simply add 273 to the Celsius temperature. To convert back, subtract 273
Note: Temperature readings are always assumed to have at least 3 significant digits. For example, 6°C is assumed to mean 279 K with 3 sig.fig., even though the data only showed 1 sig.fig.
In 1848 Lord Kelvin suggested using a temperature scale based on absolute zero, but with divisions exactly the same as the Celsius scale. For many years this was called the “absolute temperature scale” but long after his death it was renamed to honour Lord Kelvin
Example
If 2 Litres of gas at 27°C are heated in a cylinder, and the piston is allowed to rise so that pressure is kept constant, how much space will the gas take up at 327°C?
Convert temperatures to kelvins: 27°C =300k, 327°C = 600k
Use Charles’ Law: (see below)
Answer: 4 Litres
K
Litresx
K
Litresso
T
V
T
V
600300
2:,
2
2
1
1
Next slide: Lesson 2.4 Gay Lussac’s Law
Summary: Charles’ law
• The volume of a gas is
directly proportional to its
temperature
• Formula:
• Graph: Charles’ law is
usually represented by a
direct relationship graph
(straight line)
• Video1
2
2
1
1
T
V
T
V
Absolute zero
0°C=273K Temp
Volu
me (
L)
TV
Charles’ Law Assignments
• Read pages 80 to 84
• Do questions 11 to 21 on pages 97 and 98
Charles’ Law Worksheet
1. The temperature inside my fridge is about 4˚C, If I place a
balloon in my fridge that initially has a temperature of 22˚C
and a volume of 0.50 litres, what will be the volume of the
balloon when it is fully cooled? (for simplicity, we will
assume the pressure in the balloon remains the same)
Data:
T1=22˚C
T2=4˚C
V1=0.50 L
To find:
V2= unknown
Temperatures must be converted to kelvin
=295K
=277K
2
2
1
1
T
V
T
V So:
V2=V1 x T2 ÷ T1
V2=0.5L x 277K
295K
V2=0.469 L
The balloon will have a volume of 0.47 litres
divide
84
2. A man heats a balloon in the oven. If the balloon has
an initial volume of 0.40 L and a temperature of
20.0°C, what will the volume of the balloon be if he
heats it to 250°C.
85
Data
V1= 0.40L
T1= 20°C
T2= 250°C
V2= ?
Convert temperatures to kelvin
20+273= 293K, 250+273=523k
=293 K
=523 K
Use Charles’ Law
K
V
K
L
T
V
T
V
523293
4.0... 2
2
2
1
1
0.40L x 523 K ÷ 293 K = 0.7139L
0.7139L
Answer: The balloon’s volume will be 0.71 litres
3. On hot days you may have noticed that potato chip bags
seem to inflate. If I have a 250 mL bag at a temperature
of 19.0°C and I leave it in my car at a temperature of
60.0°C, what will the new volume of the bag be?
Answer: The bag will have a volume of 285mL
Data:
V1=250 mL
T1= 19.0°C
T2=60.0°C
V2= ?
Convert temperatures to kelvin
19+273= 292K, 60+273=333K
=292 K
=333 K
K
V
K
mL
T
V
T
V
333292
250... 2
2
2
1
1
Use Charles’ Law
250mL x 333 K ÷ 292 K = 285.10mL
285.10 mL
4. The volume of air in my lungs will be 2.35
litres Be sure to show your known information
Change the temperature to Kelvins and show them.
Show the formula you used and your calculations
State the answer clearly.
5.
6. The temperature is 279.7 K, which corresponds to 6.70 C. A
jacket or sweater would be appropriate clothing for this
weather.
Although only the answers are shown here, in order to get
full marks you need to show all steps of the solution!
Gay-Lussac’s Law
For Temperature-Pressure changes.
“Pressure varies directly with Temperature”
Lesson 2.4.3
Next slide:’
TP
88
Joseph Gay-Lussac (1802)
“The pressure of a gas is directly proportional to the temperature (in kelvins) if the volume is kept constant.”
Next slide:’ 89
Born 6 December 1778
Saint-Léonard-de-Noblat
Died 9 May 1850 @ Saint-Léonard-de-Noblat
Nationality: French
Fields: Chemistry
Known for Gay-Lussac's law
Gay-Lussac’s Law Relating Pressure and Temperature of a Gas
2
2
1
1
T
P
T
P
Where: P1 is the pressure* of the gas before the temperature change.
P2 is the pressure after the temperature change, in the same units.
T1 is the temperature of the gas before it changes, in kelvins.
T2 is the temperature of the gas after it changes, in kelvins.
*appropriate pressure units include: kPa, mmHg, atm.
Gay-Lussac’s Law
As the gas in a sealed
container that cannot
expand is heated, the
pressure increases.
For calculations, you
must use Kelvin
temperatures:
K= C+273
pressure
91
92
Example A sealed can contains 310 mL of air at
room temperature (20 C) and an internal
pressure of 100 kPa. If the can is heated to 606 C what will the internal pressure
be?
K
x
K
kPa
879293
100
2
2
1
1
T
P
T
P
x = 87900 ÷ 293
x = 300 Next slide: T vs P graph
Data:
P1= 100kPa
V1=310 mL
T1=20˚C
P2=unknown
T2=606˚C
˚Celsius must be converted to kelvins
20˚C = 293 K 606˚C = 879 K
Answer: the pressure will be 300 kPa
Remove irrelevant fact
=293K =879K
divide
Formula:
Temperature & Pressure Graph
The graph of temperature in Kelvin vs.
pressure in kilopascals is a straight line.
Like the temperature vs. volume graph, it
can also be used to find the value of
absolute zero.
93
Graph of Pressure-Temperature Relationship (Gay-Lussac’s Law)
Temperature (K)
Pre
ssure
(kPa)
273K Next slide:’ 94
Summary: Gay-Lussac’s law
• The pressure of a gas is
directly proportional to its
temperature
• Formula:
• Graph: Gay-Lussac’s law is
usually represented by an
direct relationship graph
(straight line)
2
2
1
1
T
P
T
P
Absolute zero
0°C=273K Temp
Pre
ssure
Assignment on Gay-Lussac’s Law
• Read pages 85 to 87
• Answer questions #22 to 30 on page 98
Avogadro’s Law For amount of gas.
“The volume of a gas is directly related to the number of moles of gas”
Lesson 2.4.4
Next slide: Lorenzo Romano Amedeo Carlo Avogadro di Quaregna
nV
97
Lorenzo Romano Amedeo Carlo
Avogadro di Quaregna
“Equal volumes of gas at the same temperature and pressure contain the same number of moles of particles.” Amedeo Avogadro
Born: August 9, 1776
Turin, Italy
Died: July 9, 1856
Field: Physics
University of Turin
Known for Avogadro’s hypothesis, Avogadro’s number.
You already know most of the facts that relate to Avogadro’s Law:
– That a mole contains a certain number of particles (6.02 x 1023)
– That a mole of gas at standard temperature and pressure will occupy 22.4 Litres (24.5 at SATP)
The only new thing here, is how changing the amount of gas present will affect pressure or volume.
– Increasing the amount of gas present will increase the volume of a gas (if it can expand),
– Increasing the amount of gas present will increase the pressure of a gas (if it is unable to expand).
99
It’s mostly common sense…
If you pump more gas into a
balloon, and allow it to expand
freely, the volume of the balloon
will increase.
If you pump more gas into a
container that can’t expand, then
the pressure inside the container
will increase.
100
Avogadro’s Laws Relating Moles of Gas to Volume or Pressure
2
2
1
1
n
V
n
V
2
2
1
1
n
P
n
Por
Where: V1 = volume before, in appropriate volume units.
V2 = volume after, in the same volume units
P1=pressure before, in appropriate pressure units.
P2=pressure after, in the same pressure units.
n1 = #moles before
n2 = #moles after 101
Assignments on Avogadro’s Law
• Read pages 88 to 92
• Do Questions 31 to 36 on page 98
102
Lesson 2.45
Standard Conditions and
Molar Volume
Next slide: 103
Standard Temperature & Pressure
(STP)
• Since the volume of a gas can change with
pressure and temperature, gases must be
compared at a specific temperature and
pressure. The long-held standard for
comparing gases is called Standard
Temperature and Pressure (STP)
• Standard Temperature =0°C = 273 K
• Standard Pressure =101.3 kPa
Ambient Temperature
• Some chemists prefer to compare gases at 25°C rather
than 0°C. At zero it is freezing, a temperature difficult to
maintain inside the lab. This alternate set of conditions
is known as Standard Ambient Temperature and
Pressure (SATP). Although not widely used, you should
be aware of it, and always watch carefully in case a
question uses AMBIENT temperature instead of
STANDARD temperature.
• Ambient Temperature = 25°C = 298 K
• Standard Pressure = 101.3 kPa
Molar Volume
• The volume of 1 mole of an ideal gas
depends on the conditions:
– At STP one mole of an ideal gas has a
volume of 22.4 litres
– AT SATP one mole of an ideal gas has a
volume of 24.5 litres
• Since all common gases are very near
ideal at these temperatures, we can use
these as standard molar volumes for ANY
common gas.
Comparison Standard and Ambient Conditions
Standard Temperature &
Pressure
(STP)
Ambient Temperature &
Pressure
(SATP)
Pressure 101.3 kPa 101.3 kPa Temperature °C 0 °C 25 °C Temperature K 273.15 K 298.15 K
Molar Volume 22.4 L/mol 24.5 L/mol
Assignments on Molar Volume
• Read pages 92 to 96
• Do Questions 37 to 52 on page 98
108
Lesson 2.5 & 2.6
The General Gas Law and the
Ideal Gas Law
Next slide: 109
The Combined or General
Gas Law • The general (or combined) gas law replaces
the four simple gas laws. It puts together: • Boyle’s Law
• Charles’ Law
• Gay-Lussac’s Law
• Avogadro’s Law
• Advantages of the Combined Gas Law: • It is easier to remember one law than four.
• It can handle changing more than one variable at a
time (eg. Changing both temperature and
pressure)
110
= General Gas Law
The General Gas Law Relating all the Simple Laws Together
22
22
11
11
Tn
VP
Tn
VP
Where: P1 P2 are the pressure of the gas before and after changes.
V1, V2 are the volume of the gas before and after changes.
T1 T2 are the temperatures, in kelvins
n 1, n2 is the number of moles of the gas.
The neat thing about the General gas law is that it can replace the three original gas laws.
Just cross out or cover the parts that don’t change, and you have the other laws:
22
22
11
11
Tn
VP
Tn
VP
Most of the time, the number of moles stays the same, so you can remove moles from the equation.
If the temperature is constant, then you have Boyle’s law.
If, instead, pressure remains constant, you have Charles’ Law
And finally, if the volume stays constant, then you have Gay-Lussac’s Law
112
The Ideal Gas Law
The Ideal Gas Law is derived from the General Gas Law in several mathematical steps.
First, start with the general gas law, including P, V, T, and the amount of gas in moles (n) .
Next slide:
22
22
11
11
Tn
VP
Tn
VP
Remember Standard Temperature & Pressure
(STP)
Standard Temperature is 0 C or more to
the point, 273K (@SATP = 25 C = 298K)
Standard Pressure is 101.3 kPa (one
atmospheric pressure at sea level)
At STP one mole of an ideal gas occupies
exactly 22.4 Litres (@SATP = 24.5 L)
The Ideal Gas Law: Calculating the
Ideal Gas Constant.
We are going to
calculate a new constant
by substituting in values
for P2, V2, T2 and n2
At STP we know all the
conditions of the gas.
Substitute and solve to
give us a constant
Kmol
LkPa
Tn
VP
2731
4.223.101
11
11
molKkPaLTn
VP/31.8
11
11
Next slide: R-- The Ideal Gas Constant
22
22
11
11
Tn
VP
Tn
VP
The Ideal Gas Constant is the proportionality constant that makes the ideal gas law work
The Ideal Gas Constant has the symbol R
R=8.31 L· kPa / K·mol
The Ideal Gas constant is 8.31 litre-
kilopascals per kelvin-mole.
Next slide: Ideal Gas Formula
So, if
Then, by a bit of algebra: P1V1=n1RT1
Since we are only using one set of
subscripts here, we might as well remove
them: PV=nRT
RTn
VP
11
11
The Ideal Gas Law Relating Conditions to the Ideal Gas Constant
nRTPVWhere: P=Pressure, in kPa
V=Volume, in Litres
n= number of moles, in mol
R= Ideal Gas constant, 8.31 LkPa/Kmol
T = Temperature, in kelvins
The Ideal gas law is best to use when you
don’t need a “before and after” situation.
Just one set of data (one volume, one
pressure, one temperature, one amount of
gas)
If you know three of the data, you can find
the missing one.
Sample Problem
8.0 g of oxygen gas is at a pressure of 2.0x102 kPa (ie: 200 Kpa w. 2 sig fig) and a temperature of 15 C. How many litres of oxygen are there? Formula: PV = nRT
Variables: P=200 kPa
V=? (our unknown)= x
n= 8.0g 32 g/mol =0.25 mol
R=8.31 L·kPa/K·mol (ideal gas constant)
T= 15 C + 273 = 288K
200 x = (0.25)(8.31)(288) , therefore
x= (0.25)(8.31)(288) 200=2.99 L
There are 3.0 L of oxygen (rounded to 2 S.D.)
Sample problem
molmolg
g
M
mn 25.0
/0.32
0.8
8.0 g of oxygen gas is at a pressure of 2.0x102 kPa (ie: 200KPa) and a temperature of 15°C. How many litres of oxygen are there? (assume 2 significant digits)
Data:
P=200 kPa
V=unknown = X
n= not given
R=8.31 L·kPa/K·mol
T= 15°C + 273 = 288K
---
m (O2) = 8g
M (O2) = 32.0 g/mol
0.25 mol
Temperature has been converted to kelvins
Calculate the value of n using the mole formula:
nRTPV200 x = (0.25)(8.31)(288) , therefore
x= (0.25)(8.31)(288) ÷ 200=2.99 L
There are 3.0 L of oxygen (rounded to 2 S.D.)
Sample Problem • 8.0 g of oxygen gas is at a pressure of 2.0x102 kPa (ie:
200KPa) and a temperature of 15°C. How many litres
of oxygen are there? (give answer to 2 significant digits)
Data:
P = 200 kPa
R = 8.31 L·kPa/K·mol
T = 15+273 = 288K
m(O2)= 8.0 g
M(O2)= 32.0 g/mol
n =
To find:
V
molmolg
g25.0
/32
8
Formula:
Work:
nRTPV
KmolK
kPaLmolVkPa 28831.825.0200
kPa
KmolK
kPaLmol
V200
28831.825.0
LV 99.2
Next slide: Ideal vs. Real
Ideal vs. Real Gases The gas laws were worked out by assuming that gases are ideal, that
is, that they obey the gas laws at all temperatures and pressures. In
reality gases will condense or solidify at low temperatures and/or high
pressures, at which point they stop behaving like gases. Also,
attraction forces between molecules may cause a gas’ behavior to vary
slightly from ideal.
A gas is ideal if its particles are extremely small (true for most gases),
the distance between particles is relatively large (true for most gases
near room temperature) and there are no forces of attraction between
the particles (not always true)
At the temperatures where a substance is a gas, it follows
the gas laws closely, but not always perfectly.
For our calculations, unless we are told otherwise, we will
assume that a gas is behaving ideally. The results will be
accurate enough for our purposes! Next slide: Summary
Testing if a gas is ideal
If you know all the important properties of a
gas (its volume, pressure, temperature in
kelvin, and the number of moles) substitute
them into the ideal gas law, but don’t put in the
value of R. Instead, calculate to see if the
value of R is close to 8.31, if so, the gas is
ideal, or very nearly so. If the calculated value
of R is quite different from 8.31 then the gas is
far from ideal.
Example
A sample of gas contains 1 mole of particles and occupies 25L., its pressure 100 kPa is and its temperature is 27 C. Is the gas ideal?
Convert to kelvins: 27 C+273=300K
PV=nRT (ideal gas law formula)
100kPa25L=1molR300K, so…
R=100kPa25L (300K1mol)
R=8.33 kPaL /Kmol expected value: 8.31 kPaL /Kmol
So the gas is not perfectly ideal, but it is very close to an ideal gas,
It varies from ideal by only 0.24% %24.0%100
31.8
)31.833.8(
Gas Laws Overview
• When using gas laws, remember that temperatures are given in Kelvins (K) – Based on absolute zero: –273 C
• The three original gas laws can be combined, and also merged with Avogadro’s mole concept to give us the Combined Gas Law.
• Rearranging the Combined Gas Law and doing a bit of algebra produces the Ideal Gas Law.
• Substituting in the STP conditions we can find the Ideal Gas Constant.
• “Ideal gases” are gases that obey the gas laws at all temperatures and pressures. In reality, no gas is perfectly ideal, but most are very close.
Gas Laws: Summary
Simple gas laws
– Boyle’s Law:
– Charles’ Law:
– Gay-Lussac’s Law:
– Combined gas law:
– Ideal gas law:
– The ideal gas constant:
22
22
11
11
Tn
VP
Tn
VP
nRTPV
VP
1
2
2
1
1
T
V
T
V
2211 VPVP 2
2
1
1
T
P
T
P
TV TP
R=8.31 Lkpa/Kmol
Assignments on the Ideal Gas
Law • Read pages 100 to 104in textbook
• Do Exercises p. 104 #1 to 16
Lesson 2.7
Stoichiometry of Gases
Stoichiometry of Gases
• When using stoichiometry with gases it is
important to remember Avogadro’s
hypothesis: that equal volumes of gas
under the same conditions of temperature
and pressure contain equal number of
particles
– Ie. At same pressure and temperature
• Same volumes have same # moles
• Volumes are proportional to numbers of moles
Simple question 1 • How many litres of hydrogen will react with
3 Litres of oxygen to form water if both
gases are at the same pressure and
temperature?
• 2 H2 + O2 2 H2O
• 2 : 1 : 2 molar ratio
• ?L : 3L : ?L volume ratio
• 6L : 3L : XX proportion solution
Answer: Six litres of hydrogen gas will react. In theory, this reaction would produce 6L of water vapour, but because the reaction is highly exothermic, the temperature would go up, so the result for water would be meaningless (hence the XX)
Assignments on the
Stoichiometry of Gases • Read pages 108 to 109 in textbook
• Do Exercises p. 110 #1 to 9
Lesson 2.8
Dalton’s Law of partial pressures
John Dalton
Born 6 September 1766
Eaglesfield,
Cumberland, England
Died 27 July 1844
Manchester, England
Notable students James Prescott Joule
Known for Atomic Theory, Law of
Multiple Proportions,
Dalton's Law of Partial
Pressures, Daltonism
Influences John Gough
Besides being the founder of modern atomic theory, John Dalton experimented on gases. He was the first to reasonably estimate the composition of the atmosphere at 21% oxygen, 79% Nitrogen
Partial Pressure
‚ Many gases are mixtures, ‚ eg. Air is 78% nitrogen, 21% Oxygen, 1% other gases
‚ Each gas in a mixture contributes a partial
pressure towards the total gas pressure.
‚ The total pressure exerted by a mixture of
gases is equal to the sum of the partial pressures
of the individual gases in the mixture.
‚ 101.3 kPa (Pair) = 79.1 kPa (N2)+ 21.2 kPa (O2) + 1.0 kPa(Other)
Next slide:
Kinetic Theory Connection
• Hypothesis 3 of the kinetic theory states
that gas particles do not attract or repel
each other.
• Dalton established that each type of gas in
a mixture behaved independently of the
other gases.
• The pressure of each gas contributes
towards the total pressure of the mixture.
Dalton’s Law The Law of Partial Pressures of Gases
Where: PT is the total pressure of mixed gases
P1 is the pressure of the 1st gas
P2 is the pressure of the 2nd gas
etc...
...21 PPPT
Variant of Dalton’s Law (used for finding partial pressure of a gas in a mixture)
T
T
AA P
n
nP
Where: PA=Pressure of gas A
nA = moles of gas A
nT= total moles of all gases
PT= Total Pressure of all gases
Uses of Dalton’s Law
In the 1960s NASA used the law of partial pressures to reduce the launch weight of their spacecraft. Instead of using air at 101 kPa, they used pure oxygen at 20kPa.
Breathing low-pressure pure oxygen gave the astronauts just as much “partial pressure” of oxygen as in normal air.
Lower pressure spacecraft reduced the chances of explosive decompression, and it also meant their spacecraft didn’t have to be as strong or heavy as those of the Russians (who used normal air).. This is one of the main reasons the Americans beat the Russians to the moon.
Carelessness with pure oxygen, however, lead to the first major tragedy of the American space program…
At 20 kPa, pure oxygen is very safe to handle, but at 101 kPa pure oxygen makes everything around it extremely flammable, and capable of burning five times faster than normal.
On January 27, 1967, during a pre-launch training exercise, the spacecraft Apollo-1 caught fire. The fire spread instantly, and the crew died before they could open the hatch.
Gus Grissom, Ed White, Roger Chaffee
Crew of Apollo 1
Exercises :
• Page 113 in new textbook, # 1 to 8
Extra practice (if you haven’t already started):
• Study guide: pp 2.12 to 2.17 # 1 to 22
– There is an answer key in the back for these
– Do these on your own as review
Summary:
• Dalton’s Law: The total pressure of a gas mixture is the sum of the partial pressures of each gas. PT = P1 + P2 + …
• Graham’s Law: light molecules diffuse faster than heavy ones
• Avogadro’s hypothesis – A mole of gas occupies 22.4L at STP and
contains 6.02x1023 particles
1
2
2
1
M
M
Rate
Rate
Summary of Kinetic Theory
• Hypotheses (re. Behaviour of gas molecules):
1. Gases are made of molecules moving randomly
2. Gas molecules are tiny with lots of space between.
3. They have elastic collisions (no lost energy).
4. Molecules don’t attract or repel each other (much)
• Results: • The kinetic energy of molecules is related to their
temperature (hot molecules have more kinetic energy because they move faster)
– Kinetic theory is based on averages of many molecules (graphed on the Maxwell distribution “bell” curve)
– Pressure is caused by the collision of molecules with the sides of their containers.
– Hotter gases and compressed gases have more collisions, therefore greater pressure.
Energy of a particle:
KE = ½ mV 2
Pressure is the result of particles colliding with the container walls.
P = F /A
Gases are made of particles
Particles move randomly!
Pressure
• The end of module 2