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Gases and the Atmosphere

RVCC Fall 2008 CHEM 103-1&2 – General Chemistry I. Gases and the Atmosphere. Chemistry: The Molecular Science, 3 rd Ed. by Moore, Stanitski, and Jurs. Gases. can be compressed exert pressure on whatever surrounds them expand into whatever volume is available mix completely one another

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Gases and the Atmosphere

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  1. RVCC Fall 2008 CHEM 103-1&2 – General Chemistry I Gases and the Atmosphere Chemistry: The Molecular Science, 3rd Ed. by Moore, Stanitski, and Jurs

  2. Gases • can be compressed • exert pressure on whatever surrounds them • expand into whatever volume is available • mix completely one another • can be described in terms of their temperature, pressure, the volume occupied, and the amount (number of moles) present P, V, T, n

  3. Pressure balls demo Pressure is the force exerted per unit area. = kg x m/s2 m2 = N (Newton) m2 = Pa (Pascal)

  4. Pressure • Other units (normal atmospheric pressure): • = 14.7 lb/in2 (psi) • = 101.325 kPa (N/m2) • = 760 mm Hg = 760 Torr • = 1 atm • = 1.01325 bar Practice 10.1 A TV weather person says the barometric pressure is 29.5” of mercury. What is this pressure in atm? mm Hg? kPa?

  5. Atmospheric Pressure • Atmosphere gases are drawn to the earth by gravitational forces • Mass of gases in our atmosphere = 5 x 1018 kg Acceleration due to gravity = 9.8 m/s2 F = ma = (5 x 1018 kg) (9.8 m/s2) = 5 x 1019 N • Pressure is force exerted per unit area, so the pressure exerted by the atmosphere is P = F / A surface area of earth = 5 x 1014 m2 P = (5 x 1019 N) / (5 x 1014 m2) = 1.0 x 105 N/m2 or 1.0 x 105 Pa or 100 kPa

  6. Elevation and Atmospheric Pressure compression

  7. Torricellian Barometer (1643) • At sea level, the column of mercury measure 760 mm above the surface of the mercury in the dish. • Does the diameter of the column matter?

  8. P = m x acceleration area = m x g x height volume Torricellian Barometer (1643) gravity g = 9.81 ms-2 = density x g x h = d g h • IF the liquid is water, the column will be 34 ft high! (DO)

  9. Kinetic Molecular Theory: Gases • gas molecules are much smaller than the distance between them (→ easily compressible, completely miscible) • gas molecules move in continuous, random, rapid motion (→ fill the container) • attractive and repulsive forces between gas molecules are negligible (→no intermolecular forces) • collisions between gas molecules are elastic • (→ total kinetic energy before and after remains the same, P doesn’t decrease over time) • the average kinetic energy of gas molecules is proportional to the absolute temperature (Kelvin scale)

  10. Distribution of Molecular Speeds All the molecules are moving but they don’t all have the same kinetic energy. • higher temperature • higher average speed. • Total area of curves • (# molecules) is the same.

  11. Molar Mass Effect on Molecular Speeds At the same temperature… lighter molecules move faster (on average). Diffusion –spread of gas molecules of one type through another. Effusion - escape of gas molecules from a container through a hole.

  12. Practice • For each of the following pairs of gases, state which will diffuse more rapidly under the same conditions of temperature and pressure. • CO2 and Br2 • N2 and HCl • NO2 and C3H8

  13. Practice True or False • If you increase the temperature of a gas (same volume), the pressure increases. • If you increase the volume of a gas (same temperature), the pressure will increase. • At the same temperature and volume, equal amounts of O2 and N2 exert the same pressure.

  14. The Behavior of Ideal Gases P is pressure (in atm, psi, kPa, mm Hg,…) V is volume (in liters, dm3) nis moles of gas T is temperature (in Kelvin) R = is the universal gas constant (matching units)

  15. The Ideal Gas Law, R? • STP (standard temperature and pressure) 0°C (273.15 K) 1 atm (101.3kPa, 760mm Hg) • At STP: one mole of a gas occupies22.414L Solve for R @ STP using any units of P: R = 0.08206 L atm K-1 mol-1 R = 62.36 L mmHg K-1 mol-1 R = 8.314 L kPa K-1 mol-1

  16. Boyle’s Law – Pressure and Volume Boyle’s Law: The volume of a gas (at a given temperature) is inversely proportional to the applied pressure. When T and n are constant… or syringe

  17. P ↑ V ↓ P ↓ V ↑ Boyle’s Law The constant (k=nRT) doesn’t change…

  18. Boyle’s Law : pressure-volume relationship. or

  19. P V P V Boyle’s Law - Practice • A sample of chlorine gas has a volume of 1.8 L at 1.0 atm. If the pressure increases to 4.0 atm (at constant temperature), what would be the new volume? • A scuba diver exhales 3.5L of air while swimming at a depth of 20m. By the time the bubbles of air rise to the surface (1.0 atm), the total volume is 10.5L. What is the pressure at the depth of 20m? • The world record for diving without supplemental air tanks (“breath-hold diving”) is about 125 meters: a depth at which the pressure is about 12.5atm. If a diver’s lungs have a volume of 6L at the surface, what is their volume at 125m? 0.45 L 3.0atm 0.48 L

  20. Charles’ Laws Charles’ Law: The volume of a gas (at a given pressure) is directly proportional to the temperature. When P and n are constant…

  21. V ↑ If T ↑ If T ↓ V ↓ Charles’ Laws

  22. Hydrogen (H2) V is proportional to T Oxygen (O2) Gas Volume (mL) 10 20 30 40 50 All gases intersect the T-axis at the same point. Absolute zero -273.15°C -300 -200 -100 0 100 200 300 Temperature (°C) Charles’ Law

  23. Charles’ Law - Practice • An anesthesiologist delivers a gas with volume of 7.5 mL at 25oC (room temperature). What would the volume of the gas be at 37oC (body temperature)? 7.8 mL • A sample of methane gas that has a volume of 3.8 L at 5.0 oC is heated to 86.0 oC at constant pressure. Calculate its new volume. 4.9 L

  24. Avogadro’s Law Avogadro’s Law: The volume of a gas (at a given temperature and pressure) is directly proportional to the amount.

  25. 2 H2(g)+ O2(g) 2 H2O(g) 2 L 1 L 2 L The Law of Combining Volumes Joseph Gay-Lussac (1809) At constant T and P, the volumes of reacting gases are always in ratios of small whole numbers. For an ideal gas, V is proportional to number of moles.

  26. All The Gas Laws Boyle’s Law n,T are constant Charles’ Law n, P are constant Avogadro’s Law P,T are constant Pressure - Temperature n, V is constant Combined Gas Law n is constant bottles demo

  27. Practice PV = nRT • If a helium-filled balloon has a volume of 340 cm3 at 25oC and 120.0 kPa, what is its volume at STP? • The pressure and temperature of a gas are 960 mm Hg and 200.0 oC, respectively. If the volume and number of moles are held constant, what would the pressure of the gas be if the temperature is reduced to 150.0oC? • 0.56 moles of gas are held in a fixed volume and temperature with the pressure of 223kPa. After using some of the gas, the new pressure in the tank is 79.6kPa. How many moles of gas was removed? • How many moles of neon gas are there are in a cylinder with a volume of 25.0 liters at 25.0oC with a pressure of 389.9 kPa? 0.369 L 858 mm Hg 0.20 moles remain, 0.36 moles were used 3.93 moles

  28. Molar Mass (MM) Determination We can use it in the ideal gas equation: solving for molar mass

  29.      Molar Mass Determination A 5.17 gram sample of an unknown gas occupied a volume of 1.92 L at 25 oC and a pressure of 1.08 atm. Calculate its molar mass. MM= 61.1 g/mol •  •  OR

  30. Density using the Ideal Gas Law Calculate the density of ozone, O3 (MM = 48.0g/mol), at 50 oC and 1.75 atm of pressure. 3.16 g/L Start with… Since, Rearrange the equation to give…

  31. Practice • What is the density of helium at 27°C and 1.25 atm? 0.203 g/L • What is the density of air at sea level (755 mm Hg and 0°C) compared to the top of Mt. Everest (210 mm Hg and 0°C). Assume the MM of air to be 29.0 g/mol. sea level, 1.29 g/L Mt. Everest, 0.358 g/L

  32. Practice PV = nRT According to the Ideal Gas Law, • Equivalent amounts of different gases at the same P and T, have the same density. • Density increases with pressure. • Molar mass increases with temperature. • For a fixed amount of gas, the quantity PV/T is a constant. • When T and n are kept constant, doubling the volume doubles the pressure. F T F T F

  33. 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 KClO3 in a test tube. How many liters of oxygen can you collect at 298 K and 1.02 atm? 0.360 L 0.360 L

  34. Practice • Automobile air bags inflate during a crash by the rapid generation of nitrogen gas from sodium azide according to the following reaction 2NaN3 (s) 2Na (s) + 3 N2 (g) How many grams of sodium azide are needed to fill a 60.0L bag to a pressure of 1.20 atm and a temperature of 15°C with N2? 132g NaN3

  35. Practice Ammonium nitrate, NH4NO3, is an explosive that undergoes the decomposition: 2NH4NO3(s)  4H2O(g) + O2(g) + 2N2(g) If 10.0 g of NH4NO3 explodes, how many total liters of gas are generated at 0.00 °C and 1.00 atm? 0.437 mol of gas (total) 9.79 L

  36. Partial Pressures of Gas Mixtures * green house gas † environmental importance What we call atmospheric pressure is the sum of the pressures exerted by all these individual gases.

  37. Partial Pressures of Gas Mixtures

  38. 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.

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

  40. 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.

  41. Practice A mixture of 10.0 g of hydrogen and 13.0 g of nitrogen in placed in a 4.0 L reaction vessel at 25ºC. a. Calculate the total pressure. b. Calculate the mole fraction of each gas. c. Calculate the partial pressure of each gas. Part a.

  42. A mixture of 10.0 g of hydrogen and 13.0 g of nitrogen in placed in a 4.0 L reaction vessel at 25ºC. a. Calculate the total pressure. b. Calculate the mole fraction of each gas. c. Calculate the partial pressure of each gas. Part b. Part c.

  43. A sample of KClO3 is heated and decomposes to produce O2 gas. The gas is collected by water displacement at 25°C (VPH2O=23.8 mm Hg @ 25°C). The total volume of the collected gas is 229 mL at a pressure of 754 mm Hg. How many moles of oxygen formed? Practice Hint: The gas collected is a mixture so use Dalton’s Law to calculate the pressure of oxygen then the ideal gas law to find the number of moles oxygen. PT = PO2 + PH2O 9.0 mmol O2

  44. Urban Air Pollution – Photochemical Smog

  45. Urban Air Pollution – Photochemical Smog Los Angeles, CA [NO] builds during the rush hour. Secondary pollutants NO2 and O3 build up later.

  46. Chlorofluorocarbons Noon UV radiation measured in Lauder, New Zealand. 1 Dobson ≡ 10-m layer of O3 at STP

  47. Sulfur Dioxide A major contributor to smog and acid rain. U.S. coal contains 1 to 4% pyrite (FeS2). 4 FeS2(s) + 11 O2(g)  2 Fe2O3(s) + 8 SO2(g) • Atmospheric SO2 reacts with O2 (and O3): • 2 SO2(g) + O2(g)  2 SO3(g) • Forming sulfuric acid • SO3(g) + H2O(l)  H2SO4(aq)

  48. Atmospheric CO2 and Global Warming • CO2 is continually produced and destroyed • Removal by: • photosynthesis • washing from the atmosphere CO2(g) + 2 H2O(l) → H3O+(aq) + HCO3-(aq) HCO3-(aq) + H2O(l) → H3O+(aq) + CO32-(aq) • Production by: • combustion • animal respiration. Global CO2 levels are rising

  49. Atmospheric CO2 and Global Warming

  50. Atmospheric CO2 and Global Warming • CO2 (Mauna Loa, HI) • +19.3% from 1959 to 2004

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