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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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standard temp and pressure stp
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?
using pv nrt to find density
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?
using pv nrt to find density1
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.
using pv nrt to find mw
Using PV=nRT to find MW
  • This is similar to the above:
dalton s law of partial pressures
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:
dalton s law of partial pressures1
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?
rms speed of gas particles
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:
rms speed of gas particles1
RMS Speed of Gas Particles
  • What does this mean?
  • That heavy gases move slower than light ones!
graham s law of effusion
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.
graham s law of effusion1
Graham’s Law of Effusion
  • More importantly, we can compare the rates of 2 gases:
graham s law of effusion2
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.
gas stoichiometry
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?
deviations from the ideal gas law
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
deviations from the ideal gas law1
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!
deviations from the ideal gas law2
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.
deviations from the ideal gas law3
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):