1 / 30

Figure 5.14: A gas whose density is greater than that of air.

Figure 5.14: A gas whose density is greater than that of air. Figure 5.15: Finding the vapor density of a substance. (Example 5.8). Sample boils Excess sample escapes When liquid is gone, submerge flask in ice water, Vapor condenses to solid Determine mass Determine density. D= m/v.

glenna
Download Presentation

Figure 5.14: A gas whose density is greater than that of air.

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript


  1. Figure 5.14: A gas whose density is greater than that of air.

  2. Figure 5.15: Finding the vapor density of a substance. (Example 5.8) Sample boils Excess sample escapes When liquid is gone, submerge flask in ice water, Vapor condenses to solid Determine mass Determine density D= m/v

  3. Density Determination • If we look again at our derivation of the molecular mass equation, we can solve for m/V, which represents density.

  4. A Problem to Consider • Calculate the density of ozone, O3 (Mm = 48.0g/mol), at 50 oC and 1.75 atm of pressure.

  5. or Molecular Weight Determination • In Chapter 3 we showed the relationship between moles and mass.

  6. Substitute into Ideal Gas Law: PV = nRT n = m/Mm Mm = molecular mass PV= mRT Mm PMm = mRT =dRT V

  7. If we solve this equation for the molecular mass, we obtain Molecular Weight Determination • If we substitute this in the ideal gas equation, we obtain

  8. A Problem to Consider • A 15.5 gram sample of an unknown gas occupied a volume of 5.75 L at 25 oC and a pressure of 1.08 atm. Calculate its molecular mass.

  9. Stoichiometry!

  10. Stoichiometry Problems Involving Gas Volumes • Consider the following reaction, which is often used to generate small quantities of oxygen. • Suppose you heat 0.0100 mol of potassium chlorate, KClO3, in a test tube. How many liters of oxygen can you produce at 298 K and 1.02 atm?

  11. Stoichiometry Problems Involving Gas Volumes • First we must determine the number of moles of oxygen produced by the reaction.

  12. Stoichiometry Problems Involving Gas Volumes • Now we can use the ideal gas equation to calculate the volume of oxygen under the conditions given. O2

  13. Partial Pressures of Gas Mixtures • The composition of a gas mixture is often described in terms of its mole fraction. • Themole fraction, c , of a component gas is the fraction of moles of that component in the total moles of gas mixture.

  14. Partial Pressures of Gas Mixtures • The partial pressure of a component gas, “A”, is then defined as • Applying this concept to the ideal gas equation, we find that each gas can be treated independently.

  15. Partial Pressures of Gas Mixtures • Dalton’s Law of Partial Pressures: the sum of all the pressures of all the different gases in a mixture equals the total pressure of the mixture. (Figure 5.16)

  16. Figure 5.17: An illustration of Dalton’s law of partial pressures before mixing.

  17. Figure 5.17: An illustration of Dalton’s law of partial pressures after mixing.

  18. A Problem to Consider • Given a mixture of gases in the atmosphere at 760 torr, what is the partial pressure of N2 (c = 0 .7808) at 25 oC?

  19. A Problem to Consider • Suppose a 156 mL sample of H2 gas was collected over water at 19 oC and 769 mm Hg. What is the mass of H2 collected? • Table 5.6 lists the vapor pressure of water at 19 oC as 16.5 mm Hg.

  20. A Problem to Consider • Now we can use the ideal gas equation, along with the partial pressure of the hydrogen, to determine its mass.

  21. A Problem to Consider • From the ideal gas law, PV = nRT, you have • Next,convert moles of H2 to grams of H2.

  22. Collecting Gases “Over Water” • A useful application of partial pressures arises when you collect gases over water. (see Figure 5.17) • As gas bubbles through the water, the gas becomes saturated with water vapor. • The partial pressure of the water in this “mixture” depends only on the temperature. (see Table 5.6)

  23. Figure 5-18 Collection of a gas over water.

  24. A Problem to Consider • Suppose a 156 mL sample of H2 gas was collected over water at 19 oC and 769 mm Hg. What is the mass of H2 collected? • First, we must find the partial pressure of the dry H2.

  25. Conceptual Problem 5.27

  26. Figure 5.27: The hydrogen fountain.Photo courtesy of American Color. Return to Slide 44

  27. Figure 5.26: Model of gaseous effusion. Return to Slide 45

More Related