2. DO NOW: 1. List three properties of gases. 2. What is pressure? Gas pressure? 3. What is a barometer?

3. A. Gases 1. State of Matter with HIGHEST amount of energy 2. Relatively few substances exist as gases under typical conditions (Ex: N 2, O 2, CO 2, Noble gases) 3. Life on earth exists in a gaseous immersion (atmosphere) 4. Properties a. Constant and random motion b. Elastic collisions c. Spreads evenly throughout container d. Mixes completely with other gases e. Easily compressible (Mostly empty space between particles)

4. B. 5.1: Pressure 5. Force exerted per unit area a. Gravity, Atmosphere 6. Gas Pressure a. Force exerted by gas particles on objects around it b. Caused by constant collision of gas particles against surfaces c. Proof? Wind, Size of Balloon indoor vs. outdoor. 7. Factors affecting Gas Pressure a. volume, temperature, quantity

5. 7. Measuring Gas Pressure a. Barometer-instrument used to measure gas pressure i. Function: 1. Column of Mercury placed in Open Container of Mercury. 2. Increase in Atmospheric Pressure Pushes Down on Hg, pushing column of Hg up, indicating HIGH atmospheric pressure. 3. Decrease in Atmospheric Pressure results in drop in Hg column, indicating LOW atmospheric pressure 4. Invented by Evangelista Torricelli (Italian, 1608-1647) http://www.flashscience.com/motion/barometer.htm

6. 8. Units of Pressure a. mm Hg (millimeters of Mercury). Most common pressure tools use column of Mercury. b. torr (in honor of Torricelli) c. atm (Standard, atmosphere) d. Pa (Pascal) or kPa (kilopascal) e. Relationship 1 atm = 760 mm Hg = 760 torr = 101.325 kPa f. Example: The pressure of in a container is measured at 927 torr. Convert to atm, mm Hg and Pa.

7. B. 5.2: The Gas Laws 1. Mathematical Laws that relate pressure of gases to temperature and/or volume 2. Derived from careful measurements of relevant gas properties

8. Boyle’s Law 1. Relationship between pressure and volume at constant temperature. 2. As pressure exerted on an object increases, volume decreases (inverse relationship) 3. Formula PV = k (constant at a given temperature) P 1 V 1 = P 2 V 2 (Common equation) animation

9. Examples 1. An aerosol can contains 400. mL of compressed gas at 5.20 atm pressure. When all the gas is sprayed into a large plastic bag, the bag inflates to a volume of 2.14 L. What is the pressure inside the bag? Assume constant temperature 2. A balloon contains 30.0 L of He gas at 103 kPa. What is the volume in mL of the balloon if the pressure is changed to 900. mm Hg? HW: Boyle’s Law Worksheet

10. A sample of an ideal gas is cooled from 50.0 °C to 25.0 °C in a sealed container of constant volume. Which of the following values for the gas will decrease? I. The average molecular mass of the gas II. The average distance between the molecules III. The average speed of the molecules (A) I only (B) II only (C) III only (D) I and III (E) II and III

11. When a sample of oxygen gas in a closed container of constant volume is heated until its absolute temperature is doubled, which of the following is also doubled? (A) The density of the gas (B) The pressure of the gas (C) The average velocity of the gas molecules (D) The number of molecules per cm 3 (E) The potential energy of the molecules

12. Gay-Lussac’s Law If the volume of a definite amount of dry gas remains constant, then the pressure varies directly with the Kelvin temperature. As T increases, gas molecules move more rapidly causing more collisions with the container walls, increasing P. DIRECT PROPORTION P = constant T P 1 = P 2 T 1 T 2 If V is constant! Temperature must be in Kelvins! T (Kelvin) = T (° Celsius) + 273

13. Charles’ Law 1. Relates temperature and volume at constant pressure. a. Direct LINEAR relationship b. As temperature increases volume increases c. NOTE: temperature extrapolates to O K for ALL gases at the same point. INFERENCE??? d. Equation: V = bT(where b is some constant specific to gas) V1=V2 T1T2 ALL GAS LAW PROBLEMS MUST BE SOLVED WITH KELVIN TEMPERATURE UNITS.

14. Example 1. A 600. mL sample of Nitrogen is heated from 37 o C to 77 o C, at constant pressure. Determine volume. 2. What is the volume of a 1.4 L sample of gas when temperature changes from 400 K to 78 o F

15. Avogadro’s Law 1. At constant temperature and pressure, equal volumes of gases contain the SAME number of particles. -- for a gas at constant temperature and pressure, the volume is directly proportional to the number of moles of gas. 2. Equation: V= an(where a is a proportionality constant, and n is moles of gas) V1=V2 n1n2

16. Example 1. Suppose you have a 12.2 L sample containing 0.50 mol Oxygen gas (O2) at a pressure of 1 atm and a temperature of 25oC. If all of this O2 was converted to Ozone, O3, at the same temperature and pressure, what would be the volume of the ozone? HW: Worksheet AND textbook page 233 Q # 33 and 34, page 236 Q # 96 (challenge problem)

17. Combined Gas Law 1. Generally, when conditions change, multiple factors are affected. 2. Helps determine unknown value related to P, V and T at same time. 3. Formula: 4. Example: A partially flexible container is filled with 2.80 L of He as at STP. It is heated until volume is 4.35 L and pressure is 1.45 atm. What is the new temp in Celcius?

18. 5.3: The IDEAL GAS Law 1. Ideal gas? a. Defined as a gas that strictly obeys Boyle’s Law b. Empirical Equation c. Gases that obey certain laws and express “expected behavior” are said to behave “ideally” d. Hypothetical 2. Real Gases (actual gases) a. Approach ideal gas behavior under certain conditions b. Errors in calculations c. OK to treat most gases using ideal gas law

19. An ideal gas is one for which both the volume of the molecules and forces between the molecules are so small that they have insignificant effect on its P, V and T behavior Low Pressure (under 1 atm) High Temperature Small mass Small particle volume (size) Non-polar

20. 3. Based on 3 Laws a. Boyle’s:V = k / P(at constant T and n) b. Charles’:V = bT(at constant P and n) c. Avogadro’s:V = an(at constant T and P) d. k, b and a (proportionality constants) are combined to form “R” i. Value of R = ____________________________ known as ______________________________ e. Derivation f. Ideal Gas Law Equation: ________________

21. Values of R (gas constant) in Different Units

22. Example 1 A sample of hydrogen gas has a volume of 8.56 L at a temperature of O o C and a pressure of 1.5 atm. Calculate the moles of H2 molecules present in this gas sample.

23. Example 2 Suppose we have a sample of ammonia gas with a volume of 7.0 mL at a pressure of 1.68 atm. The gas is compressed to a volume of 2.7 mL at a constant temperature. Use the ideal gas law to calculate the final pressure

24. Example 3 A sample of methane gas that has a volume of 3.8 L at 5 o C is heated to 80 o C at constant pressure. Calculate its new volume

25. Example 4 A sample of diborane gas (B 2 H 6 ), a substance that bursts into flame when exposed to air, has a pressure of 345 torr at a temperature of -15 o C and a volume of 3.48 L. If conditions are changed so that the temperature is 36 o C and the pressure is 468 torr, what will be the new volume of this sample.

26. Example 5 A sample containing 0.35 mol argon gas at a temperature of 13 0 C and a pressure of 568 torr is heated to 56 o C and a pressure of 897 torr. Calculate the change in volume that occurs.

27. HW Page 234 Questions 45, 46, 47, 48 Page 236 Questions 97, 98

28. Practice Which gas would have the largest number of molecules per mole? 1. CO2 2. CH4 3. He 4. O2 5. Xe 6. H2 7. All are equal

29. In rigid containers of equal mass and volume at same temperature, which gas would have the largest numer of molecules? 1. CO2 2. CH4 3. He 4. O2 5. Xe 6. H2 7. All are equal

30. If you have a rigid container with a mixture of gases (in some known ratio), and you removed one of the gases, what, if any, will be the effect on Pressure? If Pressure is affected, which factor(s) can be used to determine the result?

32. A 2.00 L flask contains 3.00 g of CO2 and 0.10 g of He at a temperature of 17.0oC. What the partial pressures of each gas and the total pressure in the flask? Plan

33. A 2.00 L flask contains 3.00 g of CO2 and 0.10 g of He at a temperature of 17.0oC. What the partial pressures of each gas and the total pressure in the flask? Plan Temperature? Individual moles Individual pressures Equation? Total pressure

38. Derivations of the Ideal Gas Law Density Determination Molar Mass Determination Can be derived from PV = nRT if mass of gas used is known. Useful to identify unknown sample of gas Derivation PV = nRT MM = g / mol (n) n = g/MM D = m (g) / V (L)

39. Ammonia Density Problem Calculate the density of ammonia gas (NH 3 ) in grams per liter at 752 mm Hg and 55 o C Plan?? Conversions Pressure Temperature Approach 1: Solve for n/V then convert n to m Approach 2: Solve for n first, then solve for d

40. Determining Molar Mass from mass, P, V and T Problem: A sample of natural gas is collected at 25.0 o C in a 250.0 mL flask. If the sample had a mass of 0.118 g at a pressure of 550.0 torr, what is the molecular mass of the gas? Approach 1: Use the ideal gas law to calculate n, then calculate the molar mass Approach 2: Solve for density. Rearrange ideal gas law to solve for molar mass.

41. Problem: A volatile liquid is placed ina 590.0 mL flask and heated until all of the liquid is gone. Only vapor fills the flask at a temperature of 100.0 o C and a pressure of 736 mm Hg. If the mass of the empty flask was 148.375 g and the mass of the flask containing the vapor is 149.457 g, what is the molar mass of the liquid?