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

Video

• Simple gas laws

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