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Chapter 5Gases

Chemistry A Molecular Approach, 1st Ed.Nivaldo

Tro

- Roy Kennedy
- Massachusetts Bay Community College
- Wellesley Hills, MA

2008, Prentice Hall

Air Pressure Shallow Wells

- water for many homes is supplied by a well less

than 30 ft. deep with a pump at the surface - the pump removes air from the pipe, decreasing

the air pressure in the pipe - the outside air pressure then pushes the water up

the pipe - the maximum height the water will rise is related

to the amount of pressure the air exerts

Atmospheric Pressure

- pressure is the force exerted over an area
- on average, the air exerts the same pressure that

a column of water 10.3 m high would exert - 14.7 lbs./in2
- so if our pump could get a perfect vacuum, the

maximum height the column could rise is 10.3 m

Gases Pushing

- gas molecules are constantly in motion
- as they move and strike a surface, they push on

that surface - push force
- if we could measure the total amount of force

exerted by gas molecules hitting the entire

surface at any one instant, we would know the

pressure the gas is exerting - pressure force per unit area

The Effect of Gas Pressure

- the pressure exerted by a gas can cause some

amazing and startling effects - whenever there is a pressure difference, a gas

will flow from area of high pressure to low

pressure - the bigger the difference in pressure, the

stronger the flow of the gas - if there is something in the gass path, the gas

will try to push it along as the gas flows

Atmospheric Pressure Effects

- differences in air pressure result in weather and

wind patterns - the higher up in the atmosphere you climb, the

lower the atmospheric pressure is around you - at the surface the atmospheric pressure is 14.7

psi, but at 10,000 ft it is only 10.0 psi - rapid changes in atmospheric pressure may cause

your ears to pop due to an imbalance in

pressure on either side of your ear drum

Pressure Imbalance in Ear

If there is a difference in pressure across the

eardrum membrane, the membrane will be pushed out

what we commonly call a popped eardrum.

The Pressure of a Gas

- result of the constant movement of the gas

molecules and their collisions with the surfaces

around them - the pressure of a gas depends on several factors
- number of gas particles in a given volume
- volume of the container
- average speed of the gas particles

Measuring Air Pressure

- use a barometer
- column of mercury supported by air pressure
- force of the air on the surface of the mercury

balanced by the pull of gravity on the column of

mercury

Gas Pressure

- Pressure is defined as the force per unit area,

and is usually measured in Pascals, which are

N/m2. - We measure pressure in mmHg or torr.
- These units of pressure are equivalent come

from measurements using a Torricellian barometer.

P pressure, F Force mass (g), A (area)

cm2, d density (g/cm3), h height

Common Units of Pressure

Example 5.1 A high-performance bicycle tire has

a pressure of 132 psi. What is the pressure in

mmHg?

132 psi mmHg

Given Find

1 atm 14.7 psi, 1 atm 760 mmHg

Concept Plan Relationships

Solution

since mmHg are smaller than psi, the answer makes

sense

Check

Manometers

- the pressure of a gas trapped in a container can

be measured with an instrument called a manometer - manometers are U-shaped tubes, partially filled

with a liquid, connected to the gas sample on one

side and open to the air on the other - a competition is established between the pressure

of the atmosphere and the gas - the difference in the liquid levels is a measure

of the difference in pressure between the gas and

the atmosphere

Manometer

for this sample, the gas has a larger pressure

than the atmosphere, so

Boyles Law

- pressure of a gas is inversely proportional to

its volume - constant T and amount of gas
- graph P vs V is curve
- graph P vs 1/V is straight line
- as P increases, V decreases by the same factor
- P x V constant
- P1 x V1 P2 x V2

Boyles Experiment

- added Hg to a J-tube with air trapped inside
- used length of air column as a measure of volume

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Boyles Experiment, P x V

When you double the pressure on a gas, the volume

is cut in half (as long as the temperature and

amount of gas do not change)

Boyles Law and Diving

- since water is denser than air, for each 10 m you

dive below the surface, the pressure on your

lungs increases 1 atm - at 20 m the total pressure is 3 atm

if your tank contained air at 1 atm pressure you

would not be able to inhale it into your lungs

Example 5.2 A cylinder with a movable piston

has a volume of 7.25 L at 4.52 atm. What is the

volume at 1.21 atm?

V1 7.25 L, P1 4.52 atm, P2 1.21 atm V2, L

Given Find

P1 V1 P2 V2

Concept Plan Relationships

Solution

since P and V are inversely proportional, when

the pressure decreases 4x, the volume should

increase 4x, and it does

Check

Practice A balloon is put in a bell jar and the

pressure is reduced from 782 torr to 0.500 atm.

If the volume of the balloon is now 2780 mL, what

was it originally?

A balloon is put in a bell jar and the pressure

is reduced from 782 torr to 0.500 atm. If the

volume of the balloon is now 2780 mL, what was it

originally?

V2 2780 mL, P1 762 torr, P2 0.500 atm V1, mL

Given Find

P1 V1 P2 V2 , 1 atm 760 torr

(exactly)

Concept Plan Relationships

Solution

since P and V are inversely proportional, when

the pressure decreases 2x, the volume should

increase 2x, and it does

Check

Charles Law

- volume is directly proportional to temperature
- constant P and amount of gas
- graph of V vs T is straight line
- as T increases, V also increases
- Kelvin T Celsius T 273
- V constant x T
- if T measured in Kelvin

Charles Law A Molecular View

- the pressure of gas inside and outside the

balloon are the same - at low temperatures, the gas molecules are not

moving as fast, so they dont hit the sides of

the balloon as hard therefore the volume is

small

- the pressure of gas inside and outside the

balloon are the same - at high temperatures, the gas molecules are

moving faster, so they hit the sides of the

balloon harder causing the volume to become

larger

The data fall on a straight line. If the lines

are extrapolated back to a volume of 0, they

all show the same temperature, -273.15C, called

absolute zero

Example 5.3 A gas has a volume of 2.57 L at

0.00C. What was the temperature at 2.80 L?

V1 2.57 L, V2 2.80 L, t2 0.00C t1, K and C

Given Find

Concept Plan Relationships

Solution

since T and V are directly proportional, when the

volume decreases, the temperature should

decrease, and it does

Check

Practice The temperature inside a balloon is

raised from 25.0C to 250.0C. If the volume of

cold air was 10.0 L, what is the volume of hot

air?

The temperature inside a balloon is raised from

25.0C to 250.0C. If the volume of cold air was

10.0 L, what is the volume of hot air?

V1 10.0 L, t1 25.0C L, t2 250.0C V2, L

Given Find

Concept Plan Relationships

Solution

since T and V are directly proportional, when the

temperature increases, the volume should

increase, and it does

Check

Avogadros Law

- volume directly proportional to the number of gas

molecules - V constant x n
- constant P and T
- more gas molecules larger volume
- count number of gas molecules by moles
- equal volumes of gases contain equal numbers of

molecules - the gas doesnt matter

Example 5.4 A 0.225 mol sample of He has a

volume of 4.65 L. How many moles must be added

to give 6.48 L?

V1 4.65 L, V2 6.48 L, n1 0.225 mol n2, and

added moles

Given Find

Concept Plan Relationships

Solution

since n and V are directly proportional, when the

volume increases, the moles should increase, and

it does

Check

Ideal Gas Law

Example 5.6 How many moles of gas are in a

basketball with total pressure 24.3 psi, volume

of 3.24 L at 25C?

V 3.24 L, P 24.3 psi, t 25 C, n, mol

Given Find

Concept Plan Relationships

Solution

1 mole at STP occupies 22.4 L, since there is a

much smaller volume than 22.4 L, we expect less

than 1 mole of gas

Check

Standard Conditions

- since the volume of a gas varies with pressure

and temperature, chemists have agreed on a set of

conditions to report our measurements so that

comparison is easy we call these standard

conditions - STP
- standard pressure 1 atm
- standard temperature 273 K
- 0C

Practice A gas occupies 10.0 L at 44.1 psi and

27C. What volume will it occupy at standard

conditions?

A gas occupies 10.0 L at 44.1 psi and 27C. What

volume will it occupy at standard conditions?

V1 10.0 L, P1 44.1 psi, t1 27 C, P2 1.00

atm, t2 0C V2, L

Given Find

Concept Plan Relationships

Solution

1 mole at STP occupies 22.4 L, since there is

more than 1 mole, we expect more than 22.4 L of

gas

Check

Molar Volume

- solving the ideal gas equation for the volume of

1 mol of gas at STP gives 22.4 L - 6.022 x 1023 molecules of gas
- notice the gas is immaterial
- we call the volume of 1 mole of gas at STP the

molar volume - it is important to recognize that one mole of

different gases have different masses, even

though they have the same volume

Molar Volume

Density at Standard Conditions

- density is the ratio of mass-to-volume
- density of a gas is generally given in g/L
- the mass of 1 mole molar mass
- the volume of 1 mole at STP 22.4 L

Gas Density

- density is directly proportional to molar mass

Example 5.7 Calculate the density of N2 at

125C and 755 mmHg

P 755 mmHg, t 125 C, dN2, g/L

Given Find

Concept Plan Relationships

Solution

since the density of N2 is 1.25 g/L at STP, we

expect the density to be lower when the

temperature is raised, and it is

Check

Molar Mass of a Gas

- one of the methods chemists use to determine the

molar mass of an unknown substance is to heat a

weighed sample until it becomes a gas, measure

the temperature, pressure, and volume, and use

the ideal gas law

Example 5.8 Calculate the molar mass of a gas

with mass 0.311 g that has a volume of 0.225 L at

55C and 886 mmHg

m0.311g, V0.225 L, P886 mmHg, t55C, molar

mass, g/mol

Given Find

m0.311g, V0.225 L, P1.1658 atm, T328 K,

molar mass, g/mol

Concept Plan Relationships

Solution

Check

the value 31.9 g/mol is reasonable

Practice - Calculate the density of a gas at 775

torr and 27C if 0.250 moles weighs 9.988 g

Calculate the density of a gas at 775 torr and

27C if 0.250 moles weighs 9.988 g

m9.988g, n0.250 mol, P775 mmHg, t27C,

density, g/L

Given Find

m9.988g, n0.250 mol, P1.0197 atm, T300. K

density, g/L

Concept Plan Relationships

Solution

Check

the value 1.65 g/L is reasonable

Mixtures of Gases

- when gases are mixed together, their molecules

behave independent of each other - all the gases in the mixture have the same volume
- all completely fill the container ? each gass

volume the volume of the container - all gases in the mixture are at the same

temperature - therefore they have the same average kinetic

energy - therefore, in certain applications, the mixture

can be thought of as one gas - even though air is a mixture, we can measure the

pressure, volume, and temperature of air as if it

were a pure substance - we can calculate the total moles of molecules in

an air sample, knowing P, V, and T, even though

they are different molecules

Partial Pressure

- the pressure of a single gas in a mixture of

gases is called its partial pressure - we can calculate the partial pressure of a gas if
- we know what fraction of the mixture it composes

and the total pressure - or, we know the number of moles of the gas in a

container of known volume and temperature - the sum of the partial pressures of all the gases

in the mixture equals the total pressure - Daltons Law of Partial Pressures
- because the gases behave independently

Composition of Dry Air

The partial pressure of each gas in a mixture can

be calculated using the ideal gas law

Example 5.9 Determine the mass of Ar in the

mixture

PHe341 mmHg, PNe112 mmHg, Ptot 662 mmHg, V

1.00 L, T298 K massAr, g

PAr 0.275 atm, V 1.00 L, T298 K massAr, g

Given Find

Concept Plan Relationships

PAr Ptot (PHe PNe)

Solution

Check

the units are correct, the value is reasonable

Practice Find the partial pressure of neon in a

mixture with total pressure 3.9 atm, volume 8.7

L, temperature 598 K, and 0.17 moles Xe.

Find the partial pressure of neon in a mixture

with total pressure 3.9 atm, volume 8.7 L,

temperature 598 K, and 0.17 moles Xe

Ptot 3.9 atm, V 8.7 L, T 598 K, Xe 0.17

mol PNe, atm

Given Find

Concept Plan Relationships

Solution

the unit is correct, the value is reasonable

Check

Mole Fraction

the fraction of the total pressure that a single

gas contributes is equal to the fraction of the

total number of moles that a single gas

contributes

the ratio of the moles of a single component to

the total number of moles in the mixture is

called the mole fraction, c for gases, volume

/ 100

the partial pressure of a gas is equal to the

mole fraction of that gas times the total pressure

Sample Problems

- A mixture of 2.50 moles neon, 1.45 moles helium,

and 2.80 moles argon has a pressure of 1.45 atm.

What are the partial pressure of all the gases in

this system?

PA XAPTotal

Sample Problems

A mixture of 2.50 moles neon, 1.45 moles helium,

and 2.80 moles argon has a pressure of 1.45 atm.

What are the partial pressure of all the gases in

this system?

2.50 mol Ne 1.45 mol He 2.8 mol Ar 6.75

mole total

Sample Problems

A mixture of 2.50 moles neon, 1.45 moles helium,

and 2.80 moles argon has a pressure of 1.45 atm.

What are the partial pressure of all the gases in

this system?

Sample Problems

A mixture of 2.50 moles neon, 1.45 moles helium,

and 2.80 moles argon has a pressure of 1.45 atm.

What are the partial pressure of all the gases in

this system?

Check Your Answer PNe PHe PAr .537 atm .311

atm .601 atm 1.45 atm total

Mountain Climbing Partial Pressure

- our bodies are adapted to breathe O2 at a partial

pressure of 0.21 atm - Sherpa, people native to the Himalaya mountains,

are adapted to the much lower partial pressure of

oxygen in their air - partial pressures of O2 lower than 0.1 atm will

lead to hypoxia - unconsciousness or death
- climbers of Mt Everest carry O2 in cylinders to

prevent hypoxia - on top of Mt Everest, Pair 0.311 atm, so PO2

0.065 atm

Deep Sea Divers Partial Pressure

- its also possible to have too much O2, a

condition called oxygen toxicity - PO2 gt 1.4 atm
- oxygen toxicity can lead to muscle spasms, tunnel

vision, and convulsions - its also possible to have too much N2, a

condition called nitrogen narcosis - also known as Rapture of the Deep
- when diving deep, the pressure of the air divers

breathe increases so the partial pressure of

the oxygen increases - at a depth of 55 m the partial pressure of O2 is

1.4 atm - divers that go below 50 m use a mixture of He and

O2 called heliox that contains a lower percentage

of O2 than air

Partial Pressure Diving

Ex 5.10 Find the mole fractions and partial

pressures in a 12.5 L tank with 24.2 g He and

4.32 g O2 at 298 K

mHe 24.2 g, mO2 43.2 g V 12.5 L, T 298

K cHe, cO2, PHe, atm, PO2, atm, Ptotal, atm

Given Find

nHe 6.05 mol, nO2 0.135 mol V 12.5 L, T

298 K cHe0.97817, cO20.021827, PHe, atm, PO2,

atm, Ptotal, atm

Concept Plan Relationships

Solution

Collecting Gases

- gases are often collected by having them displace

water from a container - the problem is that since water evaporates, there

is also water vapor in the collected gas - the partial pressure of the water vapor, called

the vapor pressure, depends only on the

temperature - so you can use a table to find out the partial

pressure of the water vapor in the gas you

collect - if you collect a gas sample with a total pressure

of 758.2 mmHg at 25C, the partial pressure of

the water vapor will be 23.78 mmHg so the

partial pressure of the dry gas will be 734.4

mmHg - Table 5.4

Vapor Pressure of Water

Collecting Gas by Water Displacement

Ex 5.11 1.02 L of O2 collected over water at

293 K with a total pressure of 755.2 mmHg. Find

mass O2.

V1.02 L, P755.2 mmHg, T293 K mass O2, g

Given Find

V1.02 L, PO2737.65 mmHg, T293 K mass O2, g

Concept Plan Relationships

Solution

Practice 0.12 moles of H2 is collected over

water in a 10.0 L container at 323 K. Find the

total pressure.

0.12 moles of H2 is collected over water in a

10.0 L container at 323 K. Find the total

pressure.

V10.0 L, nH20.12 mol, T323 K Ptotal, atm

Given Find

Concept Plan Relationships

Solution

Reactions Involving Gases

- the principles of reaction stoichiometry from

Chapter 4 can be combined with the gas laws for

reactions involving gases - in reactions of gases, the amount of a gas is

often given as a volume - instead of moles
- as weve seen, must state pressure and

temperature - the ideal gas law allows us to convert from the

volume of the gas to moles then we can use the

coefficients in the equation as a mole ratio - when gases are at STP, use 1 mol 22.4 L

P, V, T of Gas A

mole A

mole B

P, V, T of Gas B

Ex 5.12 What volume of H2 is needed to make

35.7 g of CH3OH at 738 mmHg and 355 K?CO(g) 2

H2(g) ? CH3OH(g)

mCH3OH 37.5g, P738 mmHg, T355 K VH2, L

Given Find

nH2 2.2284 mol, P0.97105 atm, T355 K VH2, L

Concept Plan Relationships

Solution

Ex 5.13 How many grams of H2O form when 1.24 L

H2 reacts completely with O2 at STP?O2(g) 2

H2(g) ? 2 H2O(g)

VH2 1.24 L, P1.00 atm, T273 K massH2O, g

Given Find

Concept Plan Relationships

H2O 18.02 g/mol, 1 mol 22.4 L _at_ STP 2 mol H2O

2 mol H2

Solution

Practice What volume of O2 at 0.750 atm and 313

K is generated by the thermolysis of 10.0 g of

HgO?2 HgO(s) ? 2 Hg(l) O2(g)(MMHgO 216.59

g/mol)

What volume of O2 at 0.750 atm and 313 K is

generated by the thermolysis of 10.0 g of HgO?2

HgO(s) ? 2 Hg(l) O2(g)

mHgO 10.0g, P0.750 atm, T313 K VO2, L

Given Find

nO2 0.023085 mol, P0.750 atm, T313 K VO2, L

Concept Plan Relationships

Solution

Properties of Gases

- expand to completely fill their container
- take the shape of their container
- low density
- much less than solid or liquid state
- compressible
- mixtures of gases are always homogeneous
- fluid

Kinetic Molecular Theory

- the particles of the gas (either atoms or

molecules) are constantly moving - the attraction between particles is negligible
- when the moving particles hit another particle or

the container, they do not stick but they bounce

off and continue moving in another direction - like billiard balls

Kinetic Molecular Theory

- there is a lot of empty space between the

particles - compared to the size of the particles
- the average kinetic energy of the particles is

directly proportional to the Kelvin temperature - as you raise the temperature of the gas, the

average speed of the particles increases - but dont be fooled into thinking all the

particles are moving at the same speed!!

Gas Properties Explained Indefinite Shape and

Indefinite Volume

Because the gas molecules have enough

kinetic energy to overcome attractions, they keep

moving around and spreading out until they fill

the container.

As a result, gases take the shape and the volume

of the container they are in.

Gas Properties Explained - Compressibility

Because there is a lot of unoccupied space in the

structure of a gas, the gas molecules can be

squeezed closer together

Gas Properties Explained Low Density

Because there is a lot of unoccupied space in the

structure of a gas, gases do not have a lot of

mass in a given volume, the result is they have

low density

Density Pressure

- result of the constant movement of the gas

molecules and their collisions with the surfaces

around them - when more molecules are added, more molecules hit

the container at any one instant, resulting in

higher pressure - also higher density

Gas Laws Explained - Boyles Law

- Boyles Law says that the volume of a gas is

inversely proportional to the pressure - decreasing the volume forces the molecules into a

smaller space - more molecules will collide with the container at

any one instant, increasing the pressure

Gas Laws Explained - Charless Law

- Charless Law says that the volume of a gas is

directly proportional to the absolute temperature - increasing the temperature increases their

average speed, causing them to hit the wall

harder and more frequently - on average
- in order to keep the pressure constant, the

volume must then increase

Gas Laws ExplainedAvogadros Law

- Avogadros Law says that the volume of a gas is

directly proportional to the number of gas

molecules - increasing the number of gas molecules causes

more of them to hit the wall at the same time - in order to keep the pressure constant, the

volume must then increase

Gas Laws Explained Daltons Law of Partial

Pressures

- Daltons Law says that the total pressure of a

mixture of gases is the sum of the partial

pressures - kinetic-molecular theory says that the gas

molecules are negligibly small and dont interact - therefore the molecules behave independent of

each other, each gas contributing its own

collisions to the container with the same average

kinetic energy - since the average kinetic energy is the same, the

total pressure of the collisions is the same

Daltons Law Pressure

- since the gas molecules are not sticking

together, each gas molecule contributes its own

force to the total force on the side

Deriving the Ideal Gas Law from Kinetic-Molecular

Theory

- pressure Forcetotal/Area
- Ftotal F1 collision x number of collisions
- in a particular time interval
- F1 collision mass x 2(velocity)/time interval
- no. of collisions is proportional to the number

of particles within the distance (velocity x time

interval) from the wall - Ftotal a massvelocity2 x Area x no.

molecules/Volume - Pressure a mv2 x n/V
- Temperature a mv2
- P a Tn/V, ? PVnRT

Calculating Gas Pressure

Molecular Velocities

- all the gas molecules in a sample can travel at

different speeds - however, the distribution of speeds follows a

pattern called a Boltzman distribution - we talk about the average velocity of the

molecules, but there are different ways to take

this kind of average - the method of choice for our average velocity is

called the root-mean-square method, where the rms

average velocity, urms, is the square root of the

average of the sum of the squares of all the

molecule velocities

Boltzman Distribution

Boltzmann animation

1000 meters/sec 3281 feet/sec speed of sound

at sea level 1116 feet per second or 768 mph or

about one mile in five seconds

Kinetic Energy and Molecular Velocities

- average kinetic energy of the gas molecules

depends on the average mass and velocity - KE ½mv2
- gases in the same container have the same

temperature, the same average kinetic energy - if they have different masses, the only way for

them to have the same kinetic energy is to have

different average velocities - lighter particles will have a faster average

velocity than more massive particles

Molecular Speed vs. Molar Mass

- in order to have the same average kinetic energy,

heavier molecules must have a slower average speed

Temperature and Molecular Velocities

- _
- KEavg ½NAmu2
- NA is Avogadros number
- KEavg 1.5RT
- R is the gas constant in energy units, 8.314

J/molK - 1 J 1 kgm2/s2
- equating and solving we get
- NAmass molar mass in kg/mol

- as temperature increases, the average velocity

increases

Temperature vs. Molecular Speed

- as the absolute temperature increases, the

average velocity increases - the distribution function spreads out,

resulting in more molecules with faster speeds

Ex 5.14 Calculate the rms velocity of O2 at 25C

O2, t 25C urms

Given Find

Concept Plan Relationships

Solution

Mean Free Path

- molecules in a gas travel in straight lines until

they collide with another molecule or the

container - the average distance a molecule travels between

collisions is called the mean free path - mean free path decreases as the pressure

increases

Diffusion and Effusion

- the process of a collection of molecules

spreading out from high concentration to low

concentration is called diffusion - the process by which a collection of molecules

escapes through a small hole into a vacuum is

called effusion - both the rates of diffusion and effusion of a gas

are related to its rms average velocity - for gases at the same temperature, this means

that the rate of gas movement is inversely

proportional to the square root of the molar mass

Effusion

Grahams Law of Effusion

- for two different gases at the same temperature,

the ratio of their rates of effusion is given by

the following equation

Ex 5.15 Calculate the molar mass of a gas that

effuses at a rate 0.462 times N2

Given Find

MM, g/mol

Concept Plan Relationships

Solution

Ideal vs. Real Gases

- Real gases often do not behave like ideal gases

at high pressure or low temperature - Ideal gas laws assume
- no attractions between gas molecules
- gas molecules do not take up space
- based on the kinetic-molecular theory
- at low temperatures and high pressures these

assumptions are not valid

The Effect of Molecular Volume

- at high pressure, the amount of space occupied by

the molecules is a significant amount of the

total volume - the molecular volume makes the real volume larger

than the ideal gas law would predict - van der Waals modified the ideal gas equation to

account for the molecular volume - b is called a van der Waals constant and is

different for every gas because their molecules

are different sizes

Real Gas Behavior

- because real molecules take up space, the molar

volume of a real gas is larger than predicted by

the ideal gas law at high pressures

The Effect of Intermolecular Attractions

- at low temperature, the attractions between the

molecules is significant - the intermolecular attractions makes the real

pressure less than the ideal gas law would

predict - van der Waals modified the ideal gas equation to

account for the intermolecular attractions - a is called a van der Waals constant and is

different for every gas because their molecules

are different sizes

Real Gas Behavior

- because real molecules attract each other, the

molar volume of a real gas is smaller than

predicted by the ideal gas law at low temperatures

Van der Waals Equation

- combining the equations to account for molecular

volume and intermolecular attractions we get the

following equation - used for real gases
- a and b are called van der Waal constants and are

different for each gas

Real Gases

- a plot of PV/RT vs. P for 1 mole of a gas shows

the difference between real and ideal gases - it reveals a curve that shows the PV/RT ratio for

a real gas is generally lower than ideality for

low pressures meaning the most important

factor is the intermolecular attractions - it reveals a curve that shows the PV/RT ratio for

a real gas is generally higher than ideality for

high pressures meaning the most important

factor is the molecular volume

PV/RT Plots

Structure of the Atmosphere

- the atmosphere shows several layers, each with

its own characteristics - the troposphere is the layer closest to the

earths surface - circular mixing due to thermal currents weather
- the stratosphere is the next layer up
- less air mixing
- the boundary between the troposphere and

stratosphere is called the tropopause - the ozone layer is located in the stratosphere

Air Pollution

- air pollution is materials added to the

atmosphere that would not be present in the air

without, or are increased by, mans activities - though many of the pollutant gases have natural

sources as well - pollution added to the troposphere has a direct

effect on human health and the materials we use

because we come in contact with it - and the air mixing in the troposphere means that

we all get a smell of it! - pollution added to the stratosphere may have

indirect effects on human health caused by

depletion of ozone - and the lack of mixing and weather in the

stratosphere means that pollutants last longer

before washing out

Pollutant Gases, SOx

- SO2 and SO3, oxides of sulfur, come from coal

combustion in power plants and metal refining - as well as volcanoes
- lung and eye irritants
- major contributor to acid rain
- 2 SO2 O2 2 H2O ? 2 H2SO4
- SO3 H2O ? H2SO4

Pollutant Gases, NOx

- NO and NO2, oxides of nitrogen, come from burning

of fossil fuels in cars, trucks, and power plants - as well as lightning storms
- NO2 causes the brown haze seen in some cities
- lung and eye irritants
- strong oxidizers
- major contributor to acid rain
- 4 NO 3 O2 2 H2O ? 4 HNO3
- 4 NO2 O2 2 H2O ? 4 HNO3

Pollutant Gases, CO

- CO comes from incomplete burning of fossil fuels

in cars, trucks, and power plants - adheres to hemoglobin in your red blood cells,

depleting your ability to acquire O2 - at high levels can cause sensory impairment,

stupor, unconsciousness, or death

Pollutant Gases, O3

- ozone pollution comes from other pollutant gases

reacting in the presence of sunlight - as well as lightning storms
- known as photochemical smog and ground-level

ozone - O3 is present in the brown haze seen in some

cities - lung and eye irritants
- strong oxidizer

Major Pollutant Levels

- government regulation has resulted in a decrease

in the emission levels for most major pollutants

Stratospheric Ozone

- ozone occurs naturally in the stratosphere
- stratospheric ozone protects the surface of the

earth from over-exposure to UV light from the sun - O3(g) UV light ? O2(g) O(g)
- normally the reverse reaction occurs quickly, but

the energy is not UV light - O2(g) O(g) ? O3(g)

Ozone Depletion

- chlorofluorocarbons became popular as aerosol

propellants and refrigerants in the 1960s - CFCs pass through the tropopause into the

stratosphere - there CFCs can be decomposed by UV light,

releasing Cl atoms - CF2Cl2 UV light ? CF2Cl Cl
- Cl atoms catalyze O3 decomposition and removes O

atoms so that O3 cannot be regenerated - NO2 also catalyzes O3 destruction
- Cl O3 ? ClO O2
- O3 UV light ? O2 O
- ClO O ? O2 Cl

Ozone Holes

- satellite data over the past 3 decades reveals a

marked drop in ozone concentration over certain

regions