Gases & Atmospheric Chemistry

1. Gases & Atmospheric Chemistry Unit 5

2. States of Matter

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. Gases – Lack of Forces • 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)

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: • Diffusion • 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: • Vibrational = back-and-forth motion of atoms within a molecule • Rotational = spinning • 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 • 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. • Gases expand as the temperature increases (much more than water and solid). • Gases have very low viscosity (they flow fast). • Gases have much lower densities than solids or liquids. • 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. Measurement of Gas Pressure • 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)

19. Units of Pressure 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 LawRelationship: Pressure & Volume • As pressure on a gas increases, the volume of the gas decreases

23. 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

25. 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

27. Converting Celsius to Kelvin • To convert degrees Celsius to Kelvin add 273 T(K) = t (ºC) + 273

28. 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

29. Relationship: Volume & Temp.

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

31. 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)

32. 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)

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

34. Gay-Lussac’s LawRelationship: Pressure & Temp.

35. Gay-Lussac’s LawRelationship: Pressure & Temp.

36. 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

37. 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

38. 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

39. 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.

40. 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

41. 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

42. 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

43. The Ideal Gas Law • IdealGas 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

44. 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

45. 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

46. The Ideal Gas Law • At STP, 1mol of an ideal gas would occupy a volume of 22.4 L

47. 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

48. 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

49. Dalton’s Law of Partial Pressures

50. Dalton’s Law of Partial Pressures Ptotal = P1 + P2 + P3….