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Gases Part 2

Gases Part 2. Standard Temp and Pressure, STP. For gases, chemists have defined a standard set of conditions: standard temperature and pressure or STP. STP is defined as 1.00 atm pressure and 0°C or 273.15K.

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Gases Part 2

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  1. Gases Part 2

  2. Standard Temp and Pressure, STP • For gases, chemists have defined a standard set of conditions: standard temperature and pressure or STP. • STP is defined as 1.00 atm pressure and 0°C or 273.15K. • If a 1.00 mol sample of ANY gas is at STP, then the volume which this sample occupies is 22.414L. • This means that at STP, we have another conversion factor: 1mol = 22.414L • Problem: A 12.37L sample of gas is at STP. How many moles are in this sample?

  3. Using PV=nRT to find density • Density of gases is given as mass/volume or g/L. • But in PV = nRT, this is very close to n/V or mol/volume. • Then we have: • But we really want g/L, not mol/L. • How can we get from n in mol to g?

  4. Using PV=nRT to find density • We know that: n = g/MW where MW is the molar mass (really mol. mass) • Now we substitute in n = g/MW in the above to get: • Problem: Find the density of helium at STP and at 30°C and 1.00atm.

  5. Using PV=nRT to find MW • This is similar to the above:

  6. Dalton’s Law of Partial Pressures • When we have a mixture of 2 or more gases, they act essentially independently of each other. • This means that they each exert their own pressure, or partial pressure. • Therefore, the total pressure of the gas mixture is equal to the sum of the partial pressures of the individual gases in the mixture. • This is stated thus:

  7. Dalton’s Law of Partial Pressures • We also use mol fractions, Xi: • If the partial pressure of water vapor is 23.756 torr at 25°C, what is the mol fraction of water if atmospheric pressure is 765 torr?

  8. RMS Speed of Gas Particles • For gas particles, we talk about the root-mean square speed (RMS speed) of particles instead of the average speed:

  9. RMS Speed of Gas Particles • What does this mean? • That heavy gases move slower than light ones!

  10. Graham’s Law of Effusion • Effusion is when gas particles escape through pinholes. • Diffusion is when gas particles mix throughout a container. • The speed or rate of effusion is related to the molar mass or molecular weight as seen above.

  11. Graham’s Law of Effusion • More importantly, we can compare the rates of 2 gases:

  12. Graham’s Law of Effusion • Again, this tells us what we already knew (or would have guessed intuitively): lighter gases effuse faster. • However, this difference in the rate of effusion is actually used to separate gases with different molecular weights.

  13. Gas Stoichiometry • We can use PV = nRT or the fact that at STP 1 mol = 22.414 L to solve gas stoichiometry problems. • Let’s look at the rxn of hydrogen gas with oxygen gas to produce liquid water: • How many g of water can be produced from 5.72 L of hydrogen gas at STP? • If 17.9 g of water is produced at 25°C and 1.00 atm, how many liters of oxygen were consumed?

  14. Deviations from the Ideal Gas Law • The Ideal Gas Law, PV = nRT is based on some assumptions. 1) Gases have negligible volume! • Wrong! Particularly for high pressures, the volume of gas particles may take up as much as 10% of the total volume of the container. • This means that gas particles exert a greater pressure, or Preal > Pideal

  15. Deviations from the Ideal Gas Law • Next wrong assumption: 2) Gas particles have no interactions! • Wrong! They do interact to some extent. • When they do interact, the pressure decreases, or Preal < Pideal • So these 2 factors tend to cancel each other out, and for pressures below around 4 atm, they pretty much do!

  16. Deviations from the Ideal Gas Law • So if the pressure is below 4 atm, we ignore deviations from ideal gas behavior. • For intermediate pressures the gas particle interactions are more important, so Preal < Pideal. • For high pressures the particles are closer and closer together, so the volume effect is much greater, or Preal > Pideal.

  17. Deviations from the Ideal Gas Law • These 2 deviation may be seen in the Van der Waals equation (don’t need to memorize or use):

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