The Gaseous State Chapter 10 1 Objectives 1

The Gaseous State Chapter 10 1

Objectives 1. 2. 3. 4. 5. Understand the definition of pressure. Use the definition to predict and measure pressure experimentally Describe experiments that show relationships between pressure, temperature, volume, and moles of a gas sample Use empirical gas laws to predict how change in one of the properties of a gas will affect the remaining properties. Use empirical gas laws to estimate gas densities and molecular masses. Use volume-to-mole relationships obtained using the empirical gas laws to solve stoichiometry problems involving gases. 2

Objectives 6. Understand the concept of partial pressure in mixtures of gases. 7. Use the ideal kinetic-molecular model to explain the empirical gas laws. 8. List deficiencies in the ideal gas mode 3 el that will cause real gases to deviate from behaviors predicted by the empirical gas laws. Explain how the model can be modified to account for these deficiencies. 3

Definition of Gas: large collection of particles moving at random throughout a volume that is primarily empty space. Have relatively large amount of energy. Gas pressure: due to collisions of randomly moving particles with the walls of the container. Force/unit area 5

Definition of Gases • STP: 0°C, and 1 atmosphere pressure • Elements that exist as gases at STP: hydrogen, nitrogen, oxygen, fluorine, chlorine and Noble Gases • Ionic compounds are all solids • Molecular compounds - depends on the intermolecular forces. Most are liquids and solids. Some are gaseous • http: //www. chemistry. ohiostate. edu/betha/neal. Gas. Law/fr 1. 1. html 6

Properties of Gases • Assume the volume and shape of their container • Compressible • Mix evenly and completely when confined to the same container • Lower densities than liquids and solids • Allotropes: O 2 ↔O 3

Kinetic Molecular Theory of Gases 1. Tiny particles in continuous motion ( the hotter the gas, the faster the molecules are moving) with negligible volume compared to volume of container. 2. Molecules are far apart from each other 3. Do not attract or repel each other (? ). 4. All collisions are elastic (gas does not lose energy when left alone). 5. The energy is proportional to Kelvin temperature. At a given temperature all gases have the same average KE. 8

Properties of Gases Observation Hypothesis Gases are easy to expand Gases are easy to compress Gases have densities that are 1/1000 of solid or liquid densities Gases completely fill their containers Hot gases leak through holes faster than cold gases 9

Properties of Gases Observation Gases are easy to expand Gases are easy to compress Gases have densities that are 1/1000 of solid or liquid densities Hot gases leak through holes faster than cold gases Hypothesis Gas molecules do not strongly attract each other Gas molecules don’t strongly repel each other Molecules are much farther apart in gases than in liquids and solids Gas molecules are in constant motion 10

Atmospheric Pressure Intensive or Extensive Property? 11

Pressure • Pressure is due to collisions between gas molecules and the walls of the container. Magnitude determined by: force of collisions and frequency. • Pressure: force per unit area: P =F/A • Standard temperature: 0ºC = 273. 15 K • Standard pressure: 1 atm in US; 1 bar elsewhere 12

Pressure Unit Symbol Conversions Pascal Pa 1 Pa = 1 N/m 2 Psi lb/in 2 Atmosphere Atm Bar 1 atm = 101325 Pa = 14. 7 lb/in 2 1 bar = 100000 Pa Torr 760 torr = 1 atm Millimeter mercury mm Hg 1 mm Hg = 1 torr 13

Pressure: Examples 1. How much pressure does an elephant with a mass of 2000 kg and total footprint area of 5000 cm 2 exert on the ground? 2. Estimate the total footprint area of a tyrannosaur weighing 16 000 kg. Assume it exerts the same pressure on its feet that the elephant does. 14

Pressure • Measuring pressure: • Strategy: – Relate pressure to fluid column heights • You can’t draw water higher than 34 feet by suction alone. Why? • Hypothesis: atmospheric pressure supports the fluid column • Develop the equation 15

Measuring Pressure 16

Pressure: Barometer measures atmospheric pressure as a mercury column height. 17

Pressure: Open-Manometer measures gas pressure as a difference in mercury column heights. Two types: closed manometer open manometer 18

Measuring Gas Pressure Closed-manometer : the arm not connected to the gas sample is closed to the atmosphere and is under vacuum. Explain how you can read the gas pressure in the bulb. 19

Pressure: Examples 3. Calculate the difference in pressure between the top and the bottom of a vessel exactly 76 cm deep filled at 25 ºC with a) water; b) mercury (d = 13. 6 g/cm 3) (7. 43 x 103 Pa; 100. 9 x 103 Pa) 4. How high a column of air would be necessary to cause the barometer to read 76 cm of mercury, if the atmosphere were of uniform density 1. 2 kg/m 3? d. Hg = 13. 53 kg/m 3 (8. 6 km) 5. A Canadian weather report gives the atmospheric pressure as 100. 2 k. Pa. What is the pressure in atmospheres? Torr? Mm Hg?

The Gas Laws: State of Gas Property Symbol Unit Property Type atm, torr, Intensive Pa Pressure P Volume V L, cm 3 Extensive Temperature T K Intensive Moles mol extensive n 21

The Gas Laws: State of Gas • Any equation that relates P, V, T, and n for a material is called an equation of state. • Experiment shows PV = n. RT is an approximate equation of state for gases. • R is the gas law constant – Determined by measuring P, V, T, n and computing R = PV/n. T – Value depends on units chosen for P, V, T – Notice: 1 Joule = 1 N m = 1(Pa) (m 3) 22

The Gas Laws Gas laws deal with the MACROSCOPIC view of gases and we try to explain the macroscopic properties by examining the microscopic behaviors (many molecule behaviors) http: //www. phy. ntnu. edu. tw/ntnujava/viewtopic. php? t=42 http: //jersey. uoregon. edu/vlab/Piston/index. html

Prentice Hall Simulations of Gas Laws • http: //cwx. prenhall. com/bookbind/pubbook s/hillchem 3/chapter 5/deluxe. html

Boyle’s Law: Experiment Relate volume to pressure when everything else is constant. Experiment: trapped air bubble at 298 K Volume, m. L 10. 0 Pressure, torr 760. 0 20. 0 379. 6 30. 0 253. 2 40. 0 191. 0 Graphs? PV )m. L torr) 25

Boyle’s Law: Experiment Relate volume to pressure when everything else is constant. Experiment: trapped air bubble at 298 K Volume, m. L 10. 0 Pressure, torr 760. 0 PV (m. L torr) 7. 60 x 103 20. 0 379. 6 7. 59 x 103 30. 0 253. 2 7. 60 x 103 40. 0 191. 0 7. 64 x 103 Graphs? 26

Boyle’s Law: Volume/Pressure Relationship At constant n, and T, the volume of a gas decreases proportionately as its pressure increases. If the pressure is doubled, the volume is halved. 27

Boyle’s Law: Volume/Pressure Relationship What happens to the volume of the gas as the pressure increases? Mathematical Relationship? 28

Plot of Boyle’s Law V versus P V versus 1/P Type of Graphs? 29

Boyle’s Law – the volume of a fixed amount of gas at constant temperature and constant number of moles is inversely proportional to the gas pressure. MOLECULAR VIEW 30

Boyle’s Law – the volume of a fixed amount of gas at constant temperature and constant number of moles is inversely proportional to the gas pressure. MOLECULAR VIEW: Confining molecules to a smaller space increases the number (frequency) of collisions, and so increases the pressure

Charles' Law (V/T Relationships) Relate volume to temperature, everything else is constant. Experiment: He bubble trapped at 1 atm. V, m. L 40. 0 T, ºC 0. 0 T, (K) 273. 2 44. 0 25. 0 298. 0 47. 7 50. 0 323. 2 51. 3 75. 0 348. 2 55. 3 100. 0 373. 2 80. 0 273. 2 546. 3 V/T, m. L/K 32

Charles' Law (V/T Relationships) Relate volume to temperature, everything else is constant. Experiment: He bubble trapped at 1 atm. V, m. L 40. 0 T, ºC 0. 0 T, (K) 273. 2 V/T, m. L/K 0. 146 44. 0 25. 0 298. 0 0. 148 47. 7 50. 0 323. 2 0. 148 51. 3 75. 0 348. 2 0. 147 55. 3 100. 0 373. 2 0. 148 80. 0 273. 2 546. 3 0. 146 33

Charles’ Law: Volume/Temperature Relationships At constant n and P, the volume of a gas increases proportionately as its absolute temperature increases, If the absolute temperature doubles, the volume is doubled. K = ºC + 273 34

Charles’ Law A plot of V versus T for a gas sample. What type of graph? Equation? 35

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Charles' Law Kinetic Interpretation of Charles's Law? Why higher pressure? Equation? Frequency and force of collision… 37

Charles’ Law The volume of the gas is directly proportional to its Kelvin temperature, when everything else is constant. MOLECULAR VIEW

Charles’ Law The volume of the gas is directly proportional to its Kelvin temperature, when everything else is constant. MOLECULAR VIEW Raising temperature increases the number of collisions and force of collisions (KE increases) with container wall. If the walls are flexible, they will be pushed back and the gas expands. 39

Charles’ Law Assume that you have a sample of gas at 350 K in a sealed container, as represented in (a). Which of the drawings (b) – (d) represents the gas after the temperature is lowered from 350 K to 150 K 40

Gay Lussac’s Law The pressure of the gas is directly proportional to its Kelvin temperature, when everything else is constant. Molecular View; 41

Gay Lussac’s Law The pressure of the gas is directly proportional to its Kelvin temperature, when everything else is constant. Molecular View; Raising the temperature increases the number of collisions and the kinetic energy of the molecules. More collisions with greater energy (force) means higher pressure. 42

Combined Gas laws

Avogadro’s Law: Relates n to Volume 44

Volume of Real Gases at STP 45

Avogadro’s Law: Relates n to Volume At constant T and P, the volume of a gas is directly proportional to moles of gas. Molar volume is almost independent of the type of gas. Samples of two gases with the same V, P, T contain the same number of molecules. MOLECULAR VIEW 46

Avogadro’s Law: Relates n to Volume At constant T and P, the volume of a gas is directly proportional to moles of gas. Molar volume is almost independent of the type of gas. Samples of two gases with the same V, P, T contain the same number of molecules (moles). MOLECULAR VIEW Type of gas does not influence distance between molecules too much. 47

Avogadro’s Law: Example 6 Show the approximate level of the movable piston in drawings (a) and (b) after the indicated changes have been made to the initial gas sample. 48

Avogadro’s Law: Answer to Example 6 49

Example 7 Show the approximate level of the movable piston in drawings (a), (b), and (c ) after the indicated changes. 50

Gas Laws: Examples 8. A balloon indoors, where the temperature is 27. 0 ºC, has a volume of 2. 00 L. What will be its volume outdoors, where the temperature is -27. 0 ºC? (Assume no change in pressure) [ 1. 67 L] 9. A sample of nitrogen occupies a volume of 2. 50 L at 120 ºC and 1. 00 atm. Pressure. To which of the following approximate temperatures should the gas be heated in order to double its volume while maintaining a constant pressure? -240 ºC - 60. 0 ºC -12. 0 ºC 30. 0 ºC [30. 0 ºC] 51

Gas Laws: Examples 10. Calculate the volume occupied by 4. 11 g of methane gas at STP. [5. 74 x 103 L] 11. What is the mass of propane, C 3 H 8, in a 50. 0 L container of the gas at STP? 52

Ideal Gas Law PV = n. RT Gas Constant R = 0. 082057 (L atm)/(mol K) 53

Examples 12. Sulfur hexafluoride, SF 6 is a colorless, odorless, very unreactive gas. Calculate the pressure (in atm) exerted by 1. 82 moles of the gas in a steel container of volume 5. 43 L at 69. 5 ºC. (9. 42 atm) 13. Calculate the volume (in liters) occupied by 7. 40 g of CO 2 at STP. ( 3. 77 L) 54

Gas Laws: Examples 14. A gas initially at 4. 0 L, 1. 2 atm, and 66 º undergoes a change so that its final volume an temperature become 1. 7 L and 42 º C. What is its final pressure? Assume the number of moles remains unchanged. 15. A certain container holds 6. 00 g of CO 2 at 150. 0 ºC and 100. k. Pa pressure. How many grams of CO 2 will it hold at 30. 0 ºC and the same pressure?

Gas Laws Summary Changing variables Variables Relationship held constant Law P, V n, T P 1 V 1 = P 2 =V 2 Boyle’s Law V, T n, P V/T = k Charle’s Law P, T n, V P/T = k Gay-Lussac’s n, V P, T V/n = k Avogadro’s P, V, T n PV/T = k Combined P, V, T, n none PV/(n. T) = R Ideal Gas Law 56

Gas Density and Molar Mass Purple M&M Do Red Too Or Michael Mo do the right thing 57

Density and Molar Mass: Examples 16. Calculate the density of methane gas, CH 4, in grams per liter, at 25 ºC and 0. 978 atm. [0. 641 g CH 4/L] 17. Under what pressure must O 2(g) be maintained at 25 ºC to have density of 1. 50 g/L? [1. 15 atm] 18. The density of a gaseous organic compound is 3. 38 g/L at 40. 0 ºC and 1. 97 atm. What is its molar mass? [44. 1 g/mol] 19. A gaseous compound is 78. 14% boron, 21. 86% hydrogen. At 27. 0 º C, 74. 3 m. L of the gas exerted a pressure of 1. 12 atm. If the mass of the gas was 0. 0934 g, what is its molecular formula? [B 2 H 6] 58

Stoichiometry Involving Gases Use regular Stoichiometry techniques, except that for non STP conditions, and 22. 4 L/mole for STP conditions. 59

Stoichiometry: The Law of Combining Volumes Involving Gases When gases measured at the same temperature and pressure allowed to react, the volumes of gaseous reactants and products are in small whole-number ratios. 60

Stoichiometry: The Law of Combining Volumes Involving Gases (Avogadro’s Explanation) When the gases are measured at the same temperature and pressure, each of the identical flasks contains the same number of molecules. 61

Examples (Stoichiometry) 20. How many liters of O 2(g) are consumed for every 10. 0 L of CO 2(g) produced in the combustion of liquid pentane, C 5 H 12, if each gas is measured at STP? [16. 0 L O 2] 21. Given the reaction C 6 H 12 O 6(s) + O 2(g) → 6 CO 2(g) + 6 H 2 O(g), calculate the volume of CO 2 produced at 37. 0 ºC and 1. 00 atm when 5. 60 g of glucose is used up in the reaction. [4. 75 L] 22. A 2. 14 L- sample of hydrogen chloride gas at 2. 61 atm and 28. 0 ºC is completely dissolved in 668 m. L of water to form hydrochloric acid solution. Calculate the molarity of the acid solution. [0. 338 M]

Dalton’s Law of Partial Pressure Assume that you have a mixture of He (4 amu) and Xe ( 131 amu) at 300 K. Which of the drawings best represents the mixture (blue= He; green = Xe)? 63

Dalton’s Law of Partial Pressure 1. What is the partial pressure of each gas – red, yellow, and green – if the total pressure inside the following container is 600 mm Hg? 2. What is the volume of each gas inside the container, if the total volume of this vessel is 1. 0 L? 64

Dalton’s Law of Partial Pressure 65

Dalton’s Law of Partial Pressures • Mole fraction: moles of component per mole of mixture • Avogadro’s Law: mole fraction = volume fraction for ideal gas Examples: 1. 2 L of He gas is mixed withy 3 L of Ne gas. What is the mole fraction of each component? 2. Air is approximately 79% N 2 and 21 %O 2 by mass. What is the mole fraction of O 2 in the air? 66

Dalton’s Law of Partial Pressures Partial Pressure – the pressure of an individual gas component in a mixture: PA Examples: 1. One mole of air contains 0. 79 moles of nitrogen and 0. 21 moles of oxygen. Compute the partial pressure of these gases at a total pressure of 1. 0 atm and at a total pressure of 3. 0 atm (about the pressure experienced by a diver under 66 ft of seawater). 2. What is the mole fraction of water in the headspace of a soda bottle, if the gas is at 2. 0 atm and 25 ºC is 23. 756 torr? 67

Dalton’s Law of Partial Pressures Ptotal = P 1 + P 2 + P 3 +……. Pn P 1 = x 1 PT 68

Dalton’s Law of Partial Pressures Dalton’s Law: The total pressure of a mixture of gases is just the sum of the pressures that each gas would exert if it were present alone. MOLECULAR VIEW Molecules of a gas do not attract or repel each other. The distances between particles are very large, therefore each particular gas occupies the entire container and adds its pressure to the total pressure in the container. 69

Dalton’s Law: Examples 23. A mixture containing 0. 538 mol He, 0. 315 mol of Ne, and 0. 103 mol of Ar is confined in a 7. 00 L vessel at 25 ºC. A) Calculate the partial pressure of each of the gasses in the mixture. B) Calculate the total pressure of the mixture. [P of He 1. 88 atm; P of Ne 1. 10 atm; P of Ar 0. 360 atm; P total 3. 34 atm] 24. The partial pressure of nitrogen in air is 592 torr. Air pressure is 752 torr, what is the mole fraction of nitrogen? [7. 87 x 10 -1] 70

Dalton’s Laws: Examples 25. What is the partial pressure of nitrogen if the container holding the air is compressed to 5. 25 atm? [4. 13 atm] 26. Ca(s) + H 2 O(l) →Ca(OH)2 + H 2(g) was collected over water. The volume of gas at 30. 0 ºC and P= 988 mm Hg is 641 m. L. What is the mass (in grams) of the H 2 gas obtained? The pressure of water at 30. 0 ºC is 31. 82 mm Hg. [0. 0653 g] 71

Dalton’s Laws: Additional Problems 27. A gaseous mixture made from 6. 00 g of oxygen and 9. 00 g of methane is placed in a 15. 0 – L vessel at 0. 00°C What is the partial pressure of each gas, and what is the total pressure in the vessel? [0. 281 atm O 2; 0. 841 CH 4; 1. 122 atm total] 28. A study of the effects of certain gases on plant growth requires a synthetic atmosphere composed of 1. 5 mol percent of CO 2, 18. 0 mol percent O 2; and 80. 5 mol percent of Ar. (a) calculate the partial pressure of O 2 in the mixture if the total pressure of the atmosphere is to be 745 torr. (b) If this atmosphere is to be held in a 120 –L space at 295 K, how many moles of O 2 are needed? [PO 2 = 134 torr; n. O 2 = 0. 872 mol] 72

Dalton’s law of Partial Pressure 29. The apparatus shown consists of three bulbs connected by stopcocks. What is the pressure inside the system when the stopcocks are opened? Assume that the lines connecting the bulbs have zero volume and that the temperature remains constants. [PCO 2 = 0. 710 atm; PH 2 = 0. 191 atm; P Ar = 0. 511 atm; PT = 1. 412 atm] 73

Example 30 4. 00 L CH 4 1. 50 L N 2 2. 70 atm . 58 atm 3. 50 L O 2 . 752 atm When these valves are opened, what is the partial pressure of each gas and the total pressure in the assembly? [P of CH 4 = 1. 2 atm; P of N 2 = 0. 097 atm; P of O 2 = 0. 292 atm; P total : add all the pressures]

Kinetic Molecular Theory of Gases 1. Gas particles are in continuous motion ( the hotter the gas, the faster the molecules are moving) with negligible volume compared to volume of container. 2. Molecules are far apart from each other 3. Do not attract or repulse each other (? ). 4. All collisions are elastic (gas does not lose energy when left alone). 5. The energy is proportional to Kelvin temperature. At a given temperature all gases have the same average KE. 75

Properties of Gases Observation Hypothesis Gases are easy to expand Gas molecules do not strongly attract each other Gases are compressible Particles have small volumes compared to continer. Lots of empty space Gases are easy to compress Gas molecules don’t strongly repel each other Gases have densities that are 1/1000 of solid or liquid densities Molecules are much farther apart in gases than in liquids and solids Hot gases leak through holes faster than cold gases Gas molecules are in constant motion 76

Properties of Gases Observation Gases undergo elastic collisions: when gas is left alone at constnat temperature, it does not liquefy or vaporize (no energy exchange) Hypothesis Gas molecules are like billiard balls – do not stick to each other (do not attract, do not repel) Hot gases leak through holes Gas molecules are in faster than cold gases constant motion 77

Kinetic Molecular Theory of Gases • Ideal gas limitations: • Gases can be liquefied if cooled enough. • Real gas molecules do attract one another to some extent otherwise the particles would not condense to form a liquid. 78

Maxwell Distribution Curves • Average Kinetic Energy at a given temperature is constant for a gas sample • But, the speeds of the molecules vary – (during to collisions with each other and with the walls of the container) • Physics: momentum is conserved • (playing pool) 79

Maxwell’s Distribution Curves http: //jersey. uoreg on. edu/vlab/Ballo on/index. html 80

Gas Laws: Maxwell’s Distribution Curves Molecules in a gas move at different speeds. 377 m/s 900 m/s compare 1500 m/s The Maxwell Distribution Curves show many molecules are moving at a particular speed. The distribution shifts to higher speeds at higher temperatures. 81

MKT of Gases: Equations • KE = ½ m(urms)2 • Average KE = (3/2) RT • Maxwell equation for the root mean square velocity: • Urms = The Urms is not the same as the mean (average ) speed. The difference is small. 82

Average Molecular speed § Average molecular kinetic energy depends only on temperature for ideal gases. § Therefore: § Higher temperature = higher root-mean- square speed (RMS), rms § Higher molecular weight (molar mass) = lower urms speed (same temperature) 83

Average Root Mean Square: Examples 31. Calculate the Urms speed, urms, of an N 2 molecule at 25ºC. (5. 15 x 102 m/s) 32. Calculate the urms speed of helium atoms 25ºC. (1. 36 x 103 m/s) 33. Calculate the Urms speed of chlorine atoms at 25ºC. (323 m/s) 84

Average Speed of Some Molecules 85

Diffusion and Effusion (a) Diffusion: mixing of gas molecules by random motion under conditions where molecular collisions occur. (Ib) Effusion: the escape of a gas through a pinhole without molecular collisions 86

Diffusion and Effusion HCl and NH 3: What will happen? 87

Graham’s Law of Diffusion • Under the same conditions of temperature and pressure, the rate of diffusion of gas molecules are inversely proportional to the square root of their molecular masses. 88

Graham’s Law of Diffusion 34. It has taken 192 seconds for 1. 4 L of an unknown gas to effuse through a porous wall and 84 seconds for the same volume of N 2 gas to effuse at the same temperature and pressure. What is the molar mass of the unknown gas? (146 g/mol) 35. In a given period of time, 0. 21 moles of a gas of MM = 26 gmol-1 effuses. How many moles of HCN would effuse in the same period of time? 36. Calculate and compare the urms of Nitrogen gas at 35 o. C and 299 K. 89

Real Gases Problems with the Kinetic Molecular Theory of "Ideal" Gases: 1. Gas particles have volume (they are not point masses). The volume becomes important under certain conditions. 2. When gas particles are close to each other, they attract each other. 90

For 1 mole of gas: PV = n. RT equation when rearranged: Plot of (PV)/(RT) for 1 mole of gas The value for the equation is not always equal to 1 Corrections to the Ideal Gas Equation is needed 91

Factors that Affect Ideality Deviation from ideal behavior as a function of temperature for nitrogen gas: 92

Factors that Affect Ideality of Gases • Interactions between the molecules (intermolecular forces): important at low temperatures and small free volume • Actual volumes of the molecules: important at high pressures and small free volume. • Free volume: the space in the container that is not occupied by the molecules. 93

Factors Affecting Ideality of Gases: low temperatures and small free volumes Distance between molecules is related to gas concentration: • At high concentration (high P, low V): • Molecules are closer (higher concentration) = stronger intermolecular attractions = deviation from ideality • Repulsion make pressure higher than expected by decreasing free volume • Attractions make pressure lower than expected by breaking molecular collisions (plastic collisions) 94

Effect of Intermolecular Attractions Orange molecules attract purple molecules. Therefore: purple molecule exert less force when it collides with the wall. No attractive forces = more force 95

Real Gases: Effect of Pressure At high pressures • Intermolecular distances between molecules decrease • Attractive forces start to play a role • Stickiness factor • Measured pressure is less than expected Correction for lower pressure 96

Real Gases: Effect of Volume should go to zero, but it does not. (a) At low pressure, the gas occupies the entire container and its volume is insignificant compared to the volume of the container. (b) At high pressure, the volume of a real gas is somewhat larger than the ideal value for an ideal gas as gas molecules take up space. 97

Correction due to volume: (V – nb) V = volume of the container n = number of moles B = volume of a mole of particles Correction for volume 98

Real Gases: Corrections • Constant needed to correct intermolecular attractive forces (make it larger) • Constant needed to correct for volume of individual gas molecules (make it smaller) The constants are characteristic properties of the substances: depend on the make-up and geometry of the substance 99

Van der Waals Constants for Common Gases Compound a (L 2 -atm/mol 2) b (L/mol) He 0. 03412 0. 02370 Ne 0. 2107 0. 01709 H 2 0. 2444 0. 02661 Ar 1. 345 0. 03219 O 2 1. 360 0. 03803 N 2 1. 390 0. 03913 CO 1. 485 0. 03985 CH 4 2. 253 0. 04278 CO 2 3. 592 0. 04267 NH 3 4. 170 0. 03707 100

Real Gases: Comparison Ideal gas Real gas Always Only at low pressures Molecular volume Zero Small, but not zero Molecular attraction Zero Small Molecular repulsion Zero small Obey PV=n. RT 101

Real Gases Large deviation form ideality: Large intermolecular attractive forces (IMF) Large Molar Mass (and subsequently volume) Real Conditions: high pressures low volumes Ideal Conditions: Low pressures (atmospheric and up to ≈ 50 atm High temperatures 102

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Factors Affecting Ideality of Gases – Tug-of-war between these two effects causes the following: • Repulsion win at very high pressure • Attractions win at moderate pressure • Neither attractions nor repulsions are important at low pressure. 105

22. 41 L atm PV PV versus P at Constant T (1 mole of Gas) CO 2 P (at constant T) 106

• V of a real gas > V of an ideal gas because V of gas molecules is significant when P is high. Ideal Gas Equation assumes that the individual gas molecules have no volume. 107

Boyle’s Law 108

Plots of Charles’ Law A plot of V versus T for a gas sample. What type of graph? 109

Kinetic Molecular Theory of Gases 1. Gases are composed of tiny atoms or molecules (particles) whose size is negligible compared to the average distance between them. This means that the volume of the individual particles in a gas can be assumed to be negligible (close to zero). 2. The particles move randomly in straight lines in all directions and at various speeds. 3. The forces of attraction or repulsion between two particles in a gas are very weak or negligible (close to zero), except when they collide. 4. When particles collide with one another, the collisions are elastic (no kinetic energy is lost). The collisions with the walls of the container create the gas pressure. 5. The average kinetic energy of a molecule is proportional to the Kelvin temperature and all calculations should be carried out with temperatures converted to K. 110

Notes: Kinetic Molecular Theory of Gases a. The observation that gases are compressible agrees with the assumption that gas particles have a small volume compared to the container. b. Elastic collisions agree with the observation that gases when left alone in a container do not seem to lose energy and do not spontaneously convert to the liquid. c. The assumptions have limitations. For example, gases can be liquefied if cooled enough. This means real gas molecules do attract one another to some extent otherwise the particles would never stick to one another in order to condense to form a liquid. 111

Measuring Pressure Open-ended Manometer Mercury Barometer Pressure? 112

Maxwell- Boltzmann Velocity (energy) Distribution Plot of Probability (fraction of molecules with given speed) versus root mean square velocity of the molecules. 113

Maxwell Distribution Curve • Variation in particle speeds for hydrogen gas at 273 K urms The vertical line on the graph represents the root-meansquare-speed (urms). The root-mean-square-speed is the square root of the averages of the squares of the speeds of all the particles in a gas sample at a particular 114 temperature.