Gases & Atmospheric Chemistry

Unit 5

States of MatterState PropertiesSolid •Definite shape and volume

•Are virtually incompressible•Do not flow easily

Liquid •Assume the shape of the container but have a definite volume•Are virtually incompressible•Flow readily

Gas •Assume the shape and volume of the container•Are highly compressible•Flow readily

Solid, Liquid & Gas

Forces Holding Solids Together The forces that are holding a solid

together are very strong

Forces: Ionic Covalent Some intermolecular forces in some

substances

Forces Holding Solids TogetherExplains why solids: Have a definite shape

Strong bonds holding molecules together (rigid)

Do not flow readily In order to be able to flow particles have to

slip past one another, strong bonds do not allow this

Cannot be compressed Strong bonds mean that there are few

empty spaces between the particles

Forces Holding Liquids Together The forces that are holding a liquid

together are not as strong as ionic or covalent bonds

Forces: intermolecular

Bonds hold molecules closely together but do not lock them into place

Liquids can spread out and take the shape of the container while keeping a constant volume

Because gases have NO definite shape or volume there appears to be an absence of forces between the molecules in a gas

No limit to the diffusion of gas molecules into the atmosphere (a very large container)

Gases – Lack of Forces

The Kinetic Molecular Theory Kinetic Molecular Theory = the idea

that all substances contain particles that are in constant, random motion

Particles are continually moving & colliding

Explains:1. Diffusion2. Evaporation

Diffusion Example: food colouring is added to

water will slowly spread out

Explanation from Kinetic Molecular Theory: molecules of food colouring and molecules of water are moving and colliding with each other which causes them to mix

Evaporation Example: water in an open container

slowly decreases as some of the water evaporates

Explanation from Kinetic Molecular Theory : some water molecules in the open container obtain sufficient energy from collision to escape from the liquid

3 Types of Motion A particle an exhibit 3 type of motion:

1. Vibrational = back-and-forth motion of atoms within a molecule

2. Rotational = spinning

3. Translational = straight line

3 Types of Motion

Motion in Relation to State Solid – mainly vibration due to

restriction of the strong bonds Particles stay together in a relatively

ordered state Liquid – some of all 3 types of motion

Less orderly state than solid Gas – rotate and vibrate but

translational (straight-line) motion is the most significant Most disordered state with no organization

Properties of Gases1. Gases are compressible: When pressure is

increased, the volume of a gas decreases. When pressure is decreased, the volume of a gas increases. The volume of a liquid and a solid remain constant during changes in temperature because their particles cannot move independently of one another like the gas particles can.

2. Gases expand as the temperature increases (much more than water and solid).

3. Gases have very low viscosity (they flow fast). 4. Gases have much lower densities than solids or

liquids. 5. ALL Gases are miscible (some liquids are miscible

yet some are immiscible).

Earth’s Leaky Atmosphere? Many of the gases that make up Earth’s

atmosphere and those of the other planets are slowly leaking into space.

Hot gases, especially light ones, evaporate away

chemical reactions and particle collisions eject atoms and molecules

and asteroids and comets occasionally blast out chunks of atmosphere

The Atmosphere

Pressure = force exerted on an object per unit of surface area Unit = Pa (pascal)

Atmospheric Pressure = the force per unit area exerted by air on all objects 100kPa one standard atmosphere (1atm)

Measurement of Gas Pressure

Units of PressureUnit of pressure Symbo

lInstruments that use the unit

1) Millimetres of Mercury: mm of Hg.

mmHg Blood pressure meters

2) 1 Torr torr Vacuum pumps

3) Pascal (Pa) the SI unit of pressure. 1 kPa = 1000 Pa

Pa Pressure sensors in pipelines

4) Bars: 1 bar bar Pressure sensors in scooba gear

5) Atmospheres (atm) atm Gas compressors

6) Pounds per square inch Psi Hydraulic pumps

Conversion: 1 atm = 760 mm of Hg = 101.325 kPa = 1.01325 bar = 760 torr = 14.7 psi

STP & SATP STP = Standard Temperature &

Pressure Exactly 0°C (273K) 1atm or101.325kPa

SATP = Standard Ambient Temperature and Pressure exactly 25°C (298K) 100kPa

Gases Moving Gases naturally move from areas of

high pressure to low pressure, because there is empty space to move into

Examples of Spray Cans: whipped cream, hair spray, paint

a propellant forces the product out

Gas Law – Boyle’s LawRelationship: Pressure & Volume

As pressure on a gas increases, the volume of the gas decreases

Gas Law – Boyle’s LawRelationship: Pressure & Volume

Boyle’s Law = as the pressure on a gas increase, the volume of the gas decreases proportionally

p1v1 = p2v2

Provided that the temperature and amount of gas are constant

The volume and pressure of a gas are inversely proportional

Kelvin Temperature Scale Absolute Zero = believed to be the

lowest possible temperature = -273ºC

Kelvin Temperature Scale = a temperature scale with zero kelvin (0 K) at absolute zero and the same size divisions as the Celsius temperature scale

Converting Celsius to Kelvin To convert degrees Celsius to Kelvin

add 273

T(K) = t (ºC) + 273

Absolute Zero As temperature decreases the

volume of a gas decreases (or the pressure drops, if you keep the volume the same)

We can deduce how cold you would have to make the gas, in order for the volume to be zero (-273°C or 0K)

A gas cannot have a zero volume therefore absolute zero is an unattainable limit

Relationship: Volume & Temp.

Charles’s LawRelationship: Volume & Temp. As temperature increases volume

increases

Charles’s LawRelationship: Volume & Temp.

Charles’s Law = the volume of a gas varies directly with its temperature in kelvin, if the pressure and the amount of gas are constant

v = kT

v = volume (L) T = Temperature in Kelvin (K) k = constant (slope of the straight line in the

graph)

Charles’s LawRelationship: Volume & Temp. Charles’s Law can be written comparing

any two sets of volumes and temperatures:

k = v1/T1 and k = v2/T2

Therefore:

v1/T1 = v2/T2 (Charles’s Law)

Practical Applications Should you throw an aerosol can into

a fire? What could happen?

When should your automobile tire pressure be checked?

Gay-Lussac’s LawRelationship: Pressure & Temp.

Gay-Lussac’s LawRelationship: Pressure & Temp.

Gay-Lussac’s LawRelationship: Pressure & Temp. Pressure & Temperature Law =

the pressure exerted by a gas varies directly with the absolute temperature if the volume and the amount of gas remain constant

p1/T1 = p2/T2

The Combined Gas Law Combined Gas Law = the product of

the pressure and volume of a gas sample is proportional to its absolute temperature in kelvin

p1v1/T1 = p2v2/T2

Gas Laws Summary T(K) = t (ºC) + 273

Boyle’s Law: p1v1 = p2v2

Charles’s Law: v1/T1 = v2/T2

Gay-Lussac’s Law: p1/T1 = p2/T2

Combined Gas Law: p1v1/T1 = p2v2/T2

No change? What happens if you don't change the

conditions of a gas, but just want to find out what a gas is like when it's sitting in a container, not doing much? 

The gas laws we’ve looked at so far won't help you much, because they are equations which depend on making a change and comparing the conditions before the change and after the change to make determinations about what the gas is like.

The Ideal Gas Law The ideal gas law is an equation of state, which means that:

you can use the basic properties of the gas to find out more about it without having to change it in any way. 

Because it's an equation of state, it allows us to not only find out what the pressure, volume, and temperature are, but also to find out how much gas is present in the first place

The Ideal Gas Ideal Gas = a hypothesized gas

composed of particles that have zero size, travel in straight lines, and have no attraction to each other (zero intermolecular force)

An imaginary model of a gas that obeys all the gas laws perfectly under all conditions

The Ideal Gas We make these assumptions because:

a)  They make the equations a whole lot simpler than they would be otherwise, and

b)  Because these assumptions don’t cause too much deviation from the ways that actual gases behave

The Ideal Gas Law Ideal Gas Law = the product of the

pressure and volume of a gas is directly proportional to the amount and the kelvin temperature of the gas

Pv = nRT

The Ideal Gas LawP = pressure in kPaV = volume in Litersn = number of moles of gasR = gas constant

Depends on the units of p, T and v

T = temperature in kelvin

Pv = nRT R = Gas constant = the constant of

variation, R, that relates the pressure in kilopascals volume in liters, amount in moles and temperature in kelvins of an ideal gas

R= 8.31 kPa•L/mol•K

R = 0.08206 atm • L /mol • K

The Ideal Gas LawAt STP, 1mol of an ideal gas would occupy a volume of 22.4 L

Summary: Properties of an Ideal Gas

V-T and p-T graphs are perfectly straight lines

Gas does not condense to a liquid when cooled

Gas volume = 0 at absolute zero pv = nRT Gas particles are point size (volume of

particle = 0 ) Gas particles do not attract each other

Mixtures of Gases Partial pressure = the pressure, p, a

gas in a mixture would exert if it were the only gas present in the same volume and at the same temperature

Dalton’s Law of Partial Pressures = the total of a mixture of non-reacting gases is equal to the sum of the partial pressures of the individual gases

Dalton’s Law of Partial Pressures

Dalton’s Law of Partial Pressures

Ptotal = P1 + P2 + P3….

Explaining Dalton’s Law of Partial Pressures with Kinetic Molecular Theory: The pressure of a gas is caused by the

collisions of molecules with the walls of the container

Gas molecules act independently of each other

Therefore the total pressure (total of the collisions with the walls) is the sum of the individual pressures (collisions of only one kind of particle) of each gas present.

Dalton’s Law of Partial Pressures

Percentage Composition of Air

Altitude vs Pressure

Partial Pressure Application At high altitude the percentage of

oxygen in air may still be normal (21%) BUT the partial pressure of oxygen may be sufficiently low that the human system cannot function very well

Partial Pressure Application

Partial Pressure Application Sea level: 101kPa

21% of 101kPa = 21kPa Top of a mountain: 33kPa

21% of 33kPa = 7kPa

Most people require about 10kPa in order to survive

Airplanes & Altitude Humans are designed to live at 1atm Airplanes function best at higher

altitudes where the pressure is much lower

This problem was originally discovered when the first pilots reached an altitude at which they lost consciousness

Airplanes & Altitude At first the problem was solved by filling

tanks with pressurized oxygen and inhaling the gas through rubber tubes

Later form-fitting face masks made oxygen delivery more reliable

Then openings were sealed to prevent air from escaping, windows were reduced in size and strengthened, and the cabin inside became a pressure capsule - like a big aluminum can

Airplanes – Pressurized Cabins Airplanes fly at 35,000 feet, while

the pressurization system maintains the cabin at the pressure you would experience at 7,000 feet, sea level

If an airplane is not pressurized? Passengers would suffocate at around 20,000 feet

What would happen if you opened the emergency door of a commercial plane during a flight?

The doors of an airplane cannot be opened in flight, because they are held closed by the air pressure inside the cabin.

The doors are designed such that they must move inward before they can move outwards when they are being opened, and they cannot move inwards when the airplane is pressurized because the air pressure presses the doors against their frames with a pressure of several tons.

Thus, there is no danger of anyone opening a door in flight.

Deep Sea Diving Pressure increases the

lower you go in water

Divers need pressurized air tanks

The tank containing compressed air is attached to a regulator that releases the air at the same pressure as the underwater surroundings

Deep Sea Diving If a diver ascends to normal pressure

too quickly or while holding their breath…

Boyle’s law… … the pressure decrease – the

volume of air increases Lungs could rupture

Avogadro's Theory Avogadro's Theory – equal volumes of

gases at the same temperature and pressure contain equal numbers of molecules

1 mole of any gas = 22.4 L at STP1 mole of any gas = 24.8 L at SATP

v1/n1 = v2/n2