1. OFB Chapter 5 The Gaseous State
5-1 The Chemistry of Gases
5-2 Pressure and Boyle’s Law
5-3 Temperature and Charles’s Law
5-4 The Ideal Gas Law
5-5 Chemical Calculations for Gases
5-6 Mixtures of Gases
5-7 Real Gases
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2. Chapter 5 The Gaseous State
Examples / Exercises
5-1 thru 5-13
Problems
8, 18, 26, 34, 38, 46, 68
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3. OFB Chapter 5 The Gaseous State
Early discoveries of gases formed by chemical reactions:
2 HgO(s)  2 Hg(l) + O2(g)
Lavoisier used this to establish the conservation of mass theory
heat
heat
Marble: CaCO3(s)  CaO(s) + CO2(g)
NH4Cl(s)  HCl(g) + NH3(g)
heat
Nitroglycerin: 4 C3H5(NO3)3(l) 
6 N2(g) + 12 CO2(g) + O2(g) + 10 H2O(g)
CaCO3(s) + HCl(aq) 
CaCl2(aq) + H2O(g) + CO2(g)
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4. Pressure and Boyle’s Law
force (F) = mass * acceleration = Newton (N) = kg m s-2
acceleration (a) = velocity per unit of time [m s-2]
mass (m) = quantity of matter [kg]
area (A) = m2
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5. Pressure and Boyle’s Law
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6. Pressure and Boyle’s Law
P = gdh
g = acceleration of gravity at the surface of the Earth
= m s-2
d = density of the liquid = for Hg at 0ºC
= g cm-3 = kg m-3
h = height of mercury in the column
= 76 cm = 760 mm = 0.76 m
P = gdh = ( m s-2)( kg m-3)
(0.76 m)
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7. A pressure of 101.325 kPa is need to raise the column of
Hg 760 mm or 76 cm
Called standard pressure
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8. Boyle’s Law: The Effect of Pressure on Gas Volume
The product of the pressure and volume, PV, of a sample of gas is a constant at a constant temperature:
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9. STP = standard temperature and pressure = 0oC and 1 atm
Boyle’s Law: The Effect of Pressure on Gas Volume
STP
For 1 mole of any gas
(i.e., 32.0 g of O2; 28.0 g N2; 2.02 g H2),
STP = standard temperature and pressure = 0oC and 1 atm
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10. Boyle’s Law: The Effect of Pressure on Gas Volume
Exercise 5-3
The long cylinder of a bicycle pump has a volume of 1131 cm3 and is filled with air at a pressure of 1.02 atm. The outlet valve is sealed shut, and the pump handle is pushed down until the volume of the air is 517 cm3. The temperature of the air trapped inside does not change. Compute the pressure inside the pump.
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11. Temperature and Charles’ Law
Charles’ Law: The Effect of Temperature on Gas Volume
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12. Charles’ Law: The Effect of Temperature on Gas Volume
Absolute Temperature
V = Vo ( )
t oC
Kelvin temperature scale
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13. Charles’ Law: The Effect of Temperature on Gas Volume
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14. Exercise 5-4
The gas in a gas thermometer that has been placed in a furnace has a volume that is 2.56 times larger than the volume that it occupies at 100oC. Determine the temperature in the furnace (in degrees Celsius).
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15. (at a fixed temperature)
P1V1 = P2V2
(at a fixed temperature)
Boyle’s Law
Charles’ Law
V1 / V2 = T1 / T2
(at a fixed pressure)
(at a fixed pressure and temperature)
Avogadro
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16. The Ideal Gas Law
V  nTP-1
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17. The Ideal Gas Law
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18. Exercise 5-5
At one point during its ascent, a weather balloon filled with helium at a volume of 1.0  104 L at 1.00 atm and 30oC reaches an altitude at which the temperature is -10oC yet the volume is unchanged. Compare the pressure at that point.
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19. ? The Ideal Gas Law R universal gas constant
for fixed V, P, and T, the number of n is fixed as well, and independent of the particular gas studied
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20. PV = nRT
ideal gas law
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21. Exercise 5-6
What mass of Hydrogen gas is needed to fill a weather balloon to a volume of 10,000 L at 1.00 atm and 30C?
Strategy
1.) use PV = nRT
2.)
3.)
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22. Exercise 5-6
What mass of Hydrogen gas is needed to fill a weather balloon to a volume of 10,000 L at 1.00 atm and 30C?
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23. Gas Density and Molar Mass
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24. Exercise 5-7 Gas Density and Molar Mass
Calculate the density of gaseous hydrogen at a pressure of 1.32 atm and a temperature of -45oC.
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25. Gas Density and Molar Mass
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26. Exercise 5-8
Fluorocarbons are compounds of fluorine and carbon. A g sample of a gaseous fluorocarbon contains 7.94 g of carbon and g of fluorine and occupies 7.40 L at STP (P = 1.00 atm and T = K). Determine the approximate molar mass of the fluorocarbon and give its molecular formula.
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27. Exercise 5-8
Fluorocarbons are compounds of fluorine and carbon. A g sample of a gaseous fluorocarbon contains 7.94 g of carbon and g of fluorine and occupies 7.40 L at STP (P = 1.00 atm and T = K). Determine the approximate molar mass of the fluorocarbon and give its molecular formula.
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28. Why use Volume for gases in chemical reaction calculations?
Chemical Calculations for Gases
Why use Volume for gases in chemical reaction calculations?
The volume of a gas is easier to measure than the mass of a gas.
Exercise 5-9
Ethylene burns in oxygen:
C2H4(g) + 3 O2(g)  2 CO2(g) + 2H2O(g)
A volume of 3.51 L of C2H4(g) at a temperature of 25oC and a pressure of 4.63 atm reacts completely with O2(g). The water vapor is collected at a temperature of 130oC and a pressure of atm. Calculate the volume of the water vapor.
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29. C2H4(g) + 3 O2(g)  2 CO2(g) + 2H2O(g)
Exercise 5-9
Ethylene burns in oxygen:
C2H4(g) + 3 O2(g)  2 CO2(g) + 2H2O(g)
A volume of 3.51 L of C2H4(g) at a temperature of 25oC and a pressure of 4.63 atm reacts completely with O2(g). The water vapor is collected at a temperature of 130oC and a pressure of atm. Calculate the volume of the water vapor.
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30. C2H4(g) + 3 O2(g)  2 CO2(g) + 2H2O(g)
Exercise 5-9
Ethylene burns in oxygen:
C2H4(g) + 3 O2(g)  2 CO2(g) + 2H2O(g)
A volume of 3.51 L of C2H4(g) at a temperature of 25oC and a pressure of 4.63 atm reacts completely with O2(g). The water vapor is collected at a temperature of 130oC and a pressure of atm. Calculate the volume of the water vapor.
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31. Dalton’s Law of Partial Pressures
Mixtures of Gases
Dalton’s Law of Partial Pressures
The total pressure of a mixture of gases equals the sum of the partial pressures of the individual gases.
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32. Mole Fractions and Partial Pressures
The mole fraction of a component in a mixture is define as the number of moles of the components that are in the mixture divided by the total number of moles present.
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33. Exercise 5-11
A solid hydrocarbon is burned in air in a closed container, producing a mixture of gases having a total pressure of 3.34 atm. Analysis of the mixture shows it to contain g of water vapor, g of carbon dioxide, g of oxygen, g of nitrogen, and no other gases. Calculate the mole fraction and partial pressure of carbon dioxide in this mixture.
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34. Section 5-7 Kinetic Theory of Gases Section 5-8 Real Gases
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