Pressure

1. Pressure Pressure: Force applied per unit area. Barometer: A device that measures atmospheric pressure. Manometer: A device for measuring the pressure of a gas in a container.

2. Pressure Units of Pressure Pascal: (abbrev. Pa) The SI unit for pressure. 1 standard atmosphere = 1.000 atm = 760.0 mm Hg = 760.0 torr 1 standard atmosphere = 101,325 Pa = 101.325 kPa 1.000 atm = 14.69 psi

3. Pressure and Volume: Boyle’s Law Boyle’s Law: Pressure times Volume equals a constant. PV=k where k is a constant at a specific temperature for a given amount of gas. If we know the volume of a gas at a given pressure, we can predict the new volume if the pressure is changed, provided that neither the temperature nor the amount of gas is changed.

4. Pressure and Volume:Boyle’s Law Example A sample has a volume of 1.51 L at a pressure of 635 torr. Calculate the final volume of the gas if the final pressure is 785 torr.

5. Volume and Temperature:Charles’s Law Charles’s Law: Proportionality constant times temperature is equal to volume. V = bT where T is in Kelvins and b is the proportionality constant. Charles’s Law implies that the amount of gas (moles) and pressure are constant. The volume of the gas is directly proportional to temperature on the Kelvin scale.

6. Volume and Temperature:Charles’s Law Example A sample has a temperature of 28oC and a volume of 23 cm3 at 1 atm. The final temperature was found to be 18oC, assuming no change in pressure. Calculate the final volume.

7. Volume and Moles:Avogadro’s Law Avogadro’s Law: For a gas at constant temperature and pressure, the volume is directly proportional to the number of moles of gas. V = an orV= a n where V is the volume of the gas N is the number of moles a is the proportionality constant.

8. Volume and Moles:Avogadro’s Law Example 3H2(g) +N2(g)  2NH3(g) If one has 15.0 L of H2(g), what volume of N2(g) is required for a complete reaction, given that both gases are at the same temperature and pressure?

9. The combined gas law Combined Gas Law: The following equation is called the combined gas law. It holds when the amount of gas (moles) is held constant. P1V1 = P2V2 T1 T2

10. The combined gas law Example A sample has a volume of 11.0 L at a temperature of 13oC and a pressure of 0.747 atm. The sample is heated to 56oC at a final pressure of 1.18 atm. Calculate the final volume.

11. Standard (STD) molar volume Lets define the volume occupied by 1 mol of a gas under specified conditions. For 1 mol of an ideal gas at 273.15 K and 1.0 atm, the volume of the gas is 22.414 L, regardless of gas. 0oC (273.15 K) and 1.0 atm = standard temperature and pressure (STP)

12. The Ideal Gas Law Ideal gas Law: The equation for the ideal gas law is PV=nRT where R=0.08206 L atm/mol K (universal gas constant). Derived from STP and standard molar volume

13. The Ideal Gas Law Example A 1.5 mol of a sample of gas has a volume of 21.0 L at 33oC. What is the pressure of the gas.

14. Density of Gases •  = m/V • PV = nRT • n = m/MM •  PV = (m/MM)RT • m/V = PMM/RT = 

15. Dalton’s Law of Partial Pressures Partial Pressure: The pressure that the gas exerts if it were above in the container. For a mixture of gases in a container, the total pressure exerted is the sum of the partial pressures of the gases present. This can be expressed as Ptotal = P1 + P2 + P3 + . . . . . . where the subscripts refer to the individual gases. The pressures P1, P2, and P3 are the partial pressures. Important Points 1. The volume of the individual gas particles must not be very important. 2. The forces among the particles must not be very important.

16. Dalton’s Law of Partial Pressures For a mixture of ideal gases, it is the total number of moles of particles that is important, not the identity of the individual gas particles. We can calculate the partial pressure of each gas from the ideal gas law.

17. Dalton’s Law of Partial Pressures Example A 2.0 L flask contains a mixture of N2 and O2 gas 25oC. The total pressure of the mixture is 0.91 atm. The mixture is known to contain 0.050 mol of N2. Calculate the partial pressure of O2 and the number of moles of O2 present.

18. Dalton’s Law of Partial Pressures • Pressure of gas in mixture of gases = product of its mole fraction and total pressure of mixture • PA = XAPtot • Ex: Xnitrogen gas = 0.78. What’s its partial pressure at STP?

19. The Kinetic Molecular Theory of Gases Kinetic Molecular Theory: The behavior of individual particles (atoms or molecules) in a gas. Postulates of the Kinetic Molecular Theory of Gases 1. Gases consist of tiny particles (atoms or molecules). 2. These particles are continually in rapid and random motion. 3. The particles are assumed not to attract or to repel each other. 4. All gases, regardless of MM, have = average KE @ = temp.

20. The Implications of the Kinetic Molecular Theory Gas speed u2 = (3RT/MM) Maxwell’s equation = gas molecule speeds  temperature Same av. KE @ same temp BUT, different av. Speeds! Smaller MM molecules go faster

21. The Implications of the Kinetic Molecular Theory • Diffusion = random mixing of gases • Effusion = gas movt through a small opening in a container into another container with lower pressure • Graham’s Law: • Effusion rate1/effusion rate2 = (MM2/MM1) • Gas effusion used initially for 235UF6/238UF6 separation

22. Gas Stoichiometry Example Calculate the volume of H2 produced at 1.50 atm and 19oC by the reaction of 26.5 g of Zn with excess HCl. Zn(s) + 2HCl(aq) ZnCl2(aq) + H2(g)

23. Solution