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Dalton’s Law of Partial Pressure

Dalton’s Law of Partial Pressure. Dalton’s Law of Partial Pressures. For a mixture of gases in a container, P Total = P 1 + P 2 + P 3 + . . . The total pressure exerted is the sum of the pressures that each gas would exert if it were alone . SEE REFERENCE SHEET. Derivation from IGL.

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Dalton’s Law of Partial Pressure

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  1. Dalton’s Law of Partial Pressure

  2. Dalton’s Law of Partial Pressures • For a mixture of gases in a container, PTotal = P1 + P2 + P3 + . . . • The total pressure exerted is the sum of the pressures that each gas would exert if it were alone. • SEE REFERENCE SHEET.

  3. Derivation from IGL

  4. Therefore… • For a mixture of ideal gases, the TOTAL NUMBER OF MOLES OF PARTICLES is important • Not the identity of the gas or composition of the particles involved • The fact that pressure isn’t affected by identity tells us two things about ideal gases… • Volume of individual particles isn’t important • Forces among the particles aren’t important

  5. 2.00 atm • 9.00 L • 3.00 atm • 3.00 L Example 1 Consider the following apparatus containing helium in both sides at 45°C. Initially the valve is closed. • After the valve is opened, what is the pressure of the helium gas?

  6. Example 2 • Some solid CaCO3 was sealed into a 250mL flask along with argon gas at a pressure of 610 torr at 22oC. The flask was heated and some of the CaCO3 decomposed, giving off CO2(g). • CaCO3(s)  CaO(s) + CO2(g) • After the flask had cooled to 22oC, the pressure was 840 torr. How many moles of CO2 were produced?

  7. Example 3 • A 40.0L tank contains a gas at 1820 torr and 23oC. What would be the pressure in the tank if 0.500 mole of gas were added to the cylinder without changing the temperature?

  8. Example 4 27.4 L of oxygen gas at 25.0°C and 1.30 atm, and 8.50 L of helium gas at 25.0°C and 2.00 atm were pumped into a tank with a volume of 5.81 L at 25°C. • Calculate the new partial pressure of oxygen. 6.13 atm • Calculate the new partial pressure of helium. 2.93 atm • Calculate the new total pressure of both gases. 9.06 atm

  9. Example 5 A study of the effects of certain gases on plant growth requires a synthetic atmosphere composed of 1.5% CO2, 18.0% O2, and 80.5% Ar. • Calculate the partial pressure of O2 in the mixture if the total pressure of the atmosphere is 745 torr. (b) If the atmosphere is to be held in a 121 L space at 295K, how many moles of O2 are needed?

  10. Example 6 • Ammonia, NH3 (g), and hydrogen chloride, HCl (g), react to form solid ammonium chloride, NH4Cl (s): • NH3 (g) + HCl (g) → NH4Cl (s) • Two 2.00L flasks at 25ºC are connected by a stopcock. On flask contains 5.00g NH3 (g) and the other contains 5.00g HCl (g). When the stopcock is opened, the gases react until one is completely consumed. (a) Which gas will remain in the system after the reaction is complete? (b) What will be the final pressure of the system after the reaction is complete? (Neglect the volume of the ammonium chloride formed).

  11. Mole fraction (χ)—ratio of the number of moles of a given component in a mixture to the total number of moles in the mixture. • χ = n1/nTOTAL • And also, χ = P1/PTOTAL • So, mole fraction of each component of mixture of ideal gases is directly related to its partial pressure. • This equation can be rearranged to P1 = χ1 x PTOTAL • SEE REFERENCE SHEET.

  12. Example 7 • Gaseous iodine pentafluoride, IF5, can be prepared by the reaction of solid iodine and gaseious fluorine: • I2 (s) + 5 F2 (g) → 2IF5 (g) • A 5.00L flask containing 10.0 g I2 is charged with 10.0 g F2, and the reaction proceeds until one of the reagents is completely consumed. After the reaction is complete, the temperature of the flask is 125ºC. (a) What is the partial pressure of IF5 in the flask? (b) What is the mole fraction of IF5 in the flask?

  13. Example 8 • The partial pressure of oxygen was observed to be 156 torr in air with a total atmospheric pressure of 743 torr. Calculate the mole fraction of O2 present.

  14. Example 9 • The mole fraction of nitrogen in the air is 0.7808. Calculate the partial pressure of N2 in air when the atmospheric pressure is 760. torr.

  15. Partial Pressures • When one collects a gas over water, there is water vapor mixed in with the gas. The vapor pressure of water depends on temperature. • To find only the pressure of the desired gas, one must subtract the vapor pressure of water from the total pressure.

  16. Example 10 • A sample of KClO3is decomposed according to the following equation. 2KClO3 (s) 2KCl (s) + 3O2 (g) • The oxygen produced was collected by displacement of water at 22oC at a total pressure of 754 torr. The volume of gas collected was 0.650L, and the vapor pressure of water at 22oC is 21 torr. • (a) What is the partial pressure of oxygen in the gas collected? • (b) How many grams of KClO3 were decomposed?

  17. Example 11 • The volume of a gas mixed with water vapor at 32.0oC at 742 torr is 1350mL. What would be the volume of the gas at 0.0oC and 760 torr if all the water vapor were removed? The pressure of water vapor at 32oC is 36 torr.

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