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Gases

Gases. Gas Stoichiometry #2. How many liters of ammonia can be produced when 12 liters of hydrogen react with an excess of nitrogen?. 3 H 2 (g) + N 2 (g)  2NH 3 (g). 12 L H 2. 2. L NH 3. = L NH 3. 8.0. 3. L H 2. Gas Stoichiometry #3.

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Gases

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  1. Gases

  2. Gas Stoichiometry #2 How many liters of ammonia can be produced when 12 liters of hydrogen react with an excess of nitrogen? 3 H2(g) + N2(g)  2NH3(g) 12 L H2 2 L NH3 = L NH3 8.0 3 L H2

  3. Gas Stoichiometry #3 How many liters of oxygen gas, at STP, can be collected from the complete decomposition of 50.0 grams of potassium chlorate? 2 KClO3(s)  2 KCl(s) + 3 O2(g) 50.0 g KClO3 1 mol KClO3 3 mol O2 22.4 L O2 122.55 g KClO3 2 mol KClO3 1 mol O2 = L O2 13.7

  4. Gas Stoichiometry #4 How many liters of oxygen gas, at 37.0C and 0.930 atmospheres, can be collected from the complete decomposition of 50.0 grams of potassium chlorate? 2 KClO3(s)  2 KCl(s) + 3 O2(g) 50.0 g KClO3 1 mol KClO3 3 mol O2 = “n” mol O2 0.612 mol O2 122.55 g KClO3 2 mol KClO3 = 16.7 L

  5. Therefore: • For each P = nRT/V • Ptotal = n1RT + n2RT + n3RT +... V VV • When gases are in the same container R, T and V are the same. • PTotal = (n1+ n2 + n3+...)RT V • PTotal = (nTotal)RT V

  6. Mol Fraction • Ratio of moles of a given component in a mixture to the total moles. •  = chi = mol1/mol T =P1/PT • Therefore: P1 =  PT • Partial Pressure = mol fraction x total pressure

  7. For a mixture of gases in a container, PTotal = P1 + P2 + P3 + . . . Dalton’s Law of Partial Pressures This is particularly useful in calculating the pressure of gases collected over water.

  8. Collecting Gas Over Water • When a gas is collected over water it is a mixture of gas & water vapor. • PT = Pgas + Pwater vapor • To find the pressure of the “dry” gas, subtract out the water vapor.

  9. Kinetic Energy of Gas Particles At the same conditions of temperature, all gases have the same average kinetic energy.

  10. The Meaning of Temperature • Kelvin temperature is an index of the random motions of gas particles (higher T means greater motion.)

  11. Diffusion Diffusion: describes the mixing of gases. Therateof diffusion is the rate of gas mixing.

  12. Effusion: describes the passage of gas into an evacuated chamber. Effusion

  13. Graham’s LawRates of Effusion and Diffusion Effusion: Diffusion:

  14. Diffusion • The spreading of a gas through a room. • It is slow considering molecules move at 100’s of meters per second. • Collisions with other molecules slow down diffusions. • Best estimate for the rate is Graham’s Law.

  15. Why do real gases not behave like ideal gases? • Real gases have attractive forces • Real gases have finite volume

  16. HOW TO MAKE A REAL GAS BEHAVE LIKE AN IDEAL GAS • Decrease the pressure • Increase the temperature

  17. Real Gases Must correct ideal gas behavior when at high pressure(smaller volume) and low temperature(attractive forces become important). ­ ­ corrected pressure corrected volume Videal Pideal

  18. van der Waals Constants for Various Gases

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