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Gases & Atmospheric Chemistry Unit 5. States of Matter StateProperties Solid Definite shape and volume Are virtually incompressible Do not flow easily.

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Presentation on theme: "Gases & Atmospheric Chemistry Unit 5. States of Matter StateProperties Solid Definite shape and volume Are virtually incompressible Do not flow easily."— Presentation transcript:

1 Gases & Atmospheric Chemistry Unit 5

2 States of Matter StateProperties Solid 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

3 Solid, Liquid & Gas

4 Forces Holding Solids Together  The forces that are holding a solid together are very strong  Forces:  Ionic  Covalent  Some intermolecular forces in some substances

5 Forces Holding Solids Together Explains 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

6 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

7  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

8 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. Diffusion 2. Evaporation

9 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

10 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

11 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

12 3 Types of Motion

13 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


15 Properties of Gases 1. 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).

16 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

17 The Atmosphere

18  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

19 Units of Pressure Unit of pressureSymbolInstruments that use the unit 1)Millimetres of Mercury: mm of Hg. mmHgBlood pressure meters 2) 1 TorrtorrVacuum pumps 3) Pascal (Pa) the SI unit of pressure. 1 kPa = 1000 Pa PaPressure sensors in pipelines 4) Bars: 1 barbarPressure sensors in scooba gear 5) Atmospheres (atm)atmGas compressors 6) Pounds per square inchPsiHydraulic pumps Conversion: 1 atm = 760 mm of Hg = 101.325 kPa = 1.01325 bar = 760 torr = 14.7 psi

20 STP & SATP  STP = Standard Temperature & Pressure  Exactly 0°C (273K)  1atm or101.325kPa  SATP = Standard Ambient Temperature and Pressure  exactly 25°C (298K)  100kPa

21 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

22 Gas Law – Boyle’s Law Relationship: Pressure & Volume  As pressure on a gas increases, the volume of the gas decreases

23 Pressure and Volume Relationship  As pressure increases volume decreases

24 Gas Law – Boyle’s Law Relationship: Pressure & Volume  Boyle’s Law = as the pressure on a gas increase, the volume of the gas decreases proportionally p 1 v 1 = p 2 v 2  Provided that the temperature and amount of gas are constant  The volume and pressure of a gas are inversely proportional


26 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


28 Converting Celsius to Kelvin  To convert degrees Celsius to Kelvin add 273 T(K) = t (ºC) + 273

29 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

30 Relationship: Volume & Temp.

31 Charles’s Law Relationship: Volume & Temp.  As temperature increases volume increases

32 As temperature increases volume increases

33 Charles’s Law Relationship: 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)

34 Charles’s Law Relationship: Volume & Temp.  Charles’s Law can be written comparing any two sets of volumes and temperatures:  k = v 1 /T 1 and k = v 2 /T 2  Therefore: v 1 /T 1 = v 2 /T 2 (Charles’s Law)

35 Practical Applications  Should you throw an aerosol can into a fire? What could happen?  When should your automobile tire pressure be checked?

36 Gay-Lussac’s Law Relationship: Pressure & Temp.


38  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 p 1 /T 1 = p 2 /T 2

39 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 p 1 v 1 /T 1 = p 2 v 2 /T 2

40 Gas Laws Summary  T(K) = t (ºC) + 273  Boyle’s Law: p 1 v 1 = p 2 v 2  Charles’s Law: v 1 /T 1 = v 2 /T 2  Gay-Lussac’s Law: p 1 /T 1 = p 2 /T 2  Combined Gas Law: p 1 v 1 /T 1 = p 2 v 2 /T 2

41 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.

42 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

43 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

44 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

45 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

46 The Ideal Gas Law  P = pressure in kPa  V = volume in Liters  n = number of moles of gas  R = gas constant  Depends on the units of p, T and v  T = temperature in kelvin

47 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 kPaL/molK R = 0.08206 atm L /mol K

48 The Ideal Gas Law  At STP, 1mol of an ideal gas would occupy a volume of 22.4 L

49 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

50 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

51 Dalton’s Law of Partial Pressures

52 P total = P 1 + P 2 + P 3 ….

53  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

54 Percentage Composition of Air

55 Altitude vs Pressure

56 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

57 Partial Pressure Application

58  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

59 Airplanes & Altitude  Humans organs have evolved 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

60 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

61 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

62 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.

63 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

64 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

65 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 STP 1 mole of any gas = 24.8 L at SATP v1/n1 = v2/n2





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