CHAPTER 11 MOLECULAR COMPOSITION OF GASES

1. CHAPTER 11 MOLECULAR COMPOSITION OF GASES Section 11.1 Volume-Mass Relationships

2. Molar Volume of Gases • In Chapter 3, we used the mole road map to convert grams, moles, and representative particles. • Since gases are measured according to volume, we can use the molar gas volume to convert gases. • Avogadro found that the volume occupied by 1 mole of any gas at STP = 22.414 L. This is the standard molar volume of gas.

3. THE MOLE ROAD MAP ATOMS, MOLECULES, FORM. UNITS MASS Use molar mass MOLE Use Avogadro’s # Use 22.4 L VOLUME

4. CHAPTER 11 MOLECULAR COMPOSITION OF GASES Section 11.2 Ideal Gas Law

5. Ideal Gases • In this chapter we are going to assume the gases behave ideally. • Ideal gases do not really exist but it makes the math easier and is a close approximation. • In an ideal gas, particles have no volume and no attractive forces.

6. More on Ideal Gases • There are no gases for which this is true but real gases behave this way at high temperature and low pressure.

7. Why don’t Ideal Gases exist? • Molecules do take up space • There are attractive forces • otherwise there would be no liquids

8. So why do Real Gases behave like Ideal Gases? • When the molecules are far apart • The molecules do not take up as big a percentage of the space • We can ignore their volume. • This is at low pressure

9. In Conclusion….. • Real Gases behave like Ideal gases when: • When molecules are moving fast. • Collisions are harder and faster. • Molecules are not next to each other very long. • Attractive forces can’t play a role.

10. The Ideal Gas Law • We now have a new way to count moles. By measuring T, P, and V. • We aren’t restricted to STP. PV = nRT R = universal gas constant (value depends on what pressure is measured in). See pg. 342 for values! n = number of moles

11. Using the Ideal Gas Law to find Density or Molar Mass • We can use the Ideal Gas Law to find Density (D) or Molar Mass (M). • It is all based on “n”. • Remember: we use the mole road map to go from moles (n) to mass (m) using the molar mass (M). n = m/M • Remember too that Density (D) = m/V

12. Using the Ideal Gas Law to find Density or Molar Mass • By substituting for “n” into the Ideal Gas Law, we get PV = nRT PV = mRT M M = mRT PV

13. Using the Ideal Gas Law to find Density or Molar Mass • Then by substituting for “m” into the Ideal Gas Law, we get M = mRT PV M = DVRT (V’s reduce) PV PM = DRT PM = D RT

14. CHAPTER 11 MOLECULAR COMPOSITION OF GASES Section 11.3 Stoichiometry of Gases

15. Solving “Stoich” Problems with Gases • Stoich problems with gases are very much the same as the ones we learned in Chapter 9. • Be sure the equation is balanced! • Identify what you know and what you need. • Set up a proportion • The numerator is always what you know and need.

16. Solving “Stoich” Problems with Gases • For the denominator: • If moles, just use coefficients from balanced equations. • If mass, use the coefficients and the molar masses. • If volume, use the coefficients and 22.4 L. • If particles, use the coefficients and 6.02 x 1023.

17. CHAPTER 11 MOLECULAR COMPOSITION OF GASES Section 11.4 Effusion & Diffusion

18. Graham’s Law of Effusion • The rate of effusion and diffusion is inversely proportional to the square root of the molar mass of the molecules. Rate of effusion of A = MB = DB Rate of effusion of B MA DA

19. Graham’s Law of Effusion – Con’t • Bigger molecules move slower at the same temp. (by Square root) • Bigger molecules effuse and diffuse slower • Helium effuses and diffuses faster than air - escapes from balloon.

20. Diffusion • Molecules moving from areas of high concentration to low concentration. • Perfume molecules spreading across the room. • Effusion: Gas escaping through a tiny hole in a container. • Depends on the speed of the molecule.