1. Zumdahl • Zumdahl • DeCoste
World of
CHEMISTRY

2. Chapter 13
Gases

3. 13.1 Pressure
Objectives:
To learn about atmospheric pressure and the what in which barometers work.
To leant the various units of pressure.
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4. Uniformly fill the container Easily compressed
Properties of Gases
Uniformly fill the container
Easily compressed
Mixes completely with any other gas.
Exerts pressure on its surroundings.
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5. Figure 13.2: A glass tube is filled with mercury and inverted in a dish of mercury at sea level.
A Barometer measures atmospheric pressure.
The pressure exerted by the atmospheric gases on the mercury in the dish, allows the tube to maintain 760mm of mercury .
How will altitude affect atmospheric pressure?
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6. Torr and mm Hg are used interchangeably by chemists.
Units of Pressure
Mm Hg
millimeters of mercury is often called the torr
Torr and mm Hg are used interchangeably by chemists.
Atmospheric pressure = atm
1.00 atm = 760 mm Hg = 760 torr
SI unit = pascal (Pa)
1.00 atm = 101,325 Pa (105 pascals)
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7. Figure 13.3: Gas pressure = atmospheric pressure – h.
A device called a monometer is used for measuring the pressure of a gas in a container. The pressure of the gas is equal to h.
Gas pressure equals atmospheric pressure - h
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8. Figure 13.3: Gas pressure = atmospheric pressure + h.
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9. Pounds per square inch – psi is the unit used by engineering science.
1.00 atm = psi
Pressure Unit Conversions: (pg 402)
The pressure of the air in a tire is measured to be 28psi, what would the pressure be in atm, torr and pascals?
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10. 13.2 Pressure and Volume: Boyle’s Law
Objectives:
To understand the law that relates the pressure and volume of a gas
To do calculations involving this law
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11. Robert Boyle (1627 – 1691)
Boyle studied the relationship between the pressure of the trapped gas and its volume.
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12. Figure 13.4: A J-tube similar to the one used by Boyle.
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13. Table 13. 1 http://www. classzone. com/cz/books/woc_07/book_home. htm
What happens as the pressure increases?
For Boyles Law to hold true, the temperature and the amount of gas must be held constant.
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14. Figure 13.5: A plot of P versus V from Boyle’s data.
Boyles Law
P x V = k
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15. Figure 13.6: Illustration of Boyle’s law.
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16. If we know the volume of a gas at a given pressure, we can predict the new volume if the pressure is changed.
P1V1 = k P2V2 = k
P1V1 = P2V2
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17. Calculating Volume Using Boyle’s Law
Consider a 1.5 –L sample of CCl2F2 (Freon -12) at a pressure of 56 torr at a constant temp. If the pressure is changed to 150 torr at constant pressure.
Will the volume of the gas increase or decrease?
What will be the new volume?
More practice pg
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18. 13.3 Volume and Temperature: Charles Law
Objectives:
To learn about absolute zero
To learn about the law relating the volume and temperature of a sample of gas at constant moles and pressure and to do calculations involving that law
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19. Figure 13.7: Plots of V (L) versus T (°C) for several gases.
Volume of a given gas (at constant pressure) increases with the temperature of the gas.
That is, it exhibits a linear relationship.
-273°C = the lowest temperature that matter can reach.
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20. Figure 13.8: Plots of V versus T using the Kelvin scale for temperature.
Absolute Zero
The direct proportionality between volume and temperature in kelvins:
V=bT
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21. V=bT
V T = b V1/T1 = V2/T2
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22. 13.4 Volumes and Moles: Avagadro’s Law
Objectives:
To understand the law relating the volume and the number of moles of a sample of gas at constant temperature and pressure and to do calculations involving this law.
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23. What is the relationship between volume of a gas and the number of molecules present ?
Volume is proportional to the number of moles in a sample of gas (providing pressure and temperature are constant)
Volume = a n
a= constant
n= number of moles
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24. Figure 13.9: The relationship between volume V and number of moles n.
For a gas at constant temp and pressure, the volume is directly proportional to the number of moles in the gas.
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25. Focus Questions 13.1 – 13.4
Mercury is a very toxic substance. Why is it used in barometers and manometer instead of water?
What is the SI unit for pressure? What is the unit commonly used in chemistry for pressure? Why aren’t they the same?
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26. 13.5 The Ideal Gas Law Objectives:
To understand the ideal gas law and use it in calculations.
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27. The Ideal Gas Law
The equation that shows how the volume of a gas depends on the pressure, temperature and the number of moles of gas present is combined:
PV = nRT
Where R is the constant
(atm∙L / K∙mol)
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28. 13.6 Dalton’s Law of Partial Pressure
Objectives:
To understand the relationship between the partial and total pressure and total pressure of a gas mixture, and to use this relationship in calculations.
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29. The partial pressure of a gas is the pressure that the gas would exert if it were alone in a container
The total pressure exerted is the sum of the partial pressures exerted if it where alone in the container.
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30. Figure 13.10: When two gases are present, the total pressure is the sum of the partial pressures of the gases.
Ptotal = P1 +P2 + P3 + ……
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31. Figure 13.11: The total pressure of a mixture of gases depends on the number of moles of gas particles present.
Refer to combined gas laws using laws of partial pressure, pg 421
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32. Figure 13.12: The production of oxygen by thermal decomposition.
A mixture of gases are collected by the displacement of water. The gas released is a mixture.
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33. Focus Questions 13.5 – 13.6
When solving gas problems, you can use the ideal gas equation even when some of the variables remain the same. Explain how you do this.
At what temperature will 6.21 g of oxygen gas exert a pressure of 5.00 atm in a
10.0-L container.
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34. 3. The pressure of a gas is affected by the number of particles and not by the kind of gas it is. What doe this fact tell us about ideal gases?
4. A tank contains a mixture of 3.0 mol N2, 2.0 moles O2 and 1.0 mol CO2 at 25° C and a total pressure of 10.0 atm. Calculate the partial pressure (in torr) of each gas in the mixture.
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35. In figure 13. 12 oxygen gas is being collected over water
In figure oxygen gas is being collected over water. Assume that this is being done at 25°C. Write and equation for it, and find the pressure of oxygen gas (assume the pressure in the bottle is PT)
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36. 13.7 Laws and Models: A Review
Objectives:
To understand the relationship between laws and models (theories)
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37. 13.8 The kinetic Molecular Theory of Gases
Objectives:
To understand the basic postulates of the kinetic molecular theory
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38. Postulates of KMT.
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39. 13.9 The implications of the Kinetic Molecular Theory
Objectives:
To understand the term temperature
To learn how kinetic molecular theory explains the gas laes
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40. Figure 13. 13: (a) A gas confined in a cylinder with a movable piston
Figure 13.13: (a) A gas confined in a cylinder with a movable piston. (b) The temperature of the gas is increased at constant pressure P ext.
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41. 13.10 Real Gases Objectives: To describe the properties of real gases.
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42. Figure 13.14: A gas sample is compressed.
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43. 13.11 Gas Stoichiometry Objectives:
To understand the molar volume of an ideal gas
To learn the definition of STP
To use these concepts and the ideal gas equation.
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44. Focus Questions 13.7 – 13.11
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