Home Podcasts THMG075 – Boyle, Henry, Charles, and Dalton

THMG075 – Boyle, Henry, Charles, and Dalton


In this episode, we explore the four laws of gasses – Boyle’s, Henry’s, Charles’s, and Dalton’s.

Complete Show Notes

3:20 Gas Laws and Why They Matter

  • These are 4 of the most common gas laws in chemistry, although they fall under the heading of kinetic molecular theory
  • Very important because we need to know how the gasses we interact with on a daily basis work

5:50 Boyle’s Law

  • States that there is a direct correlation between the pressure and volume of a container
  • This assumes a fixed temperature (usually 77 degrees F or 68 degrees F) and a known amount of moles, which is the measurement of atoms in a sample
  • Law shows that when a gas is pressurized, the volume decreases – we see this in hydraulics, because air can be squeezed (unlike liquids)
  • Formula: V = 1/P

7:50 Henry’s Law

  • Covers how gas is dissolved into a solution – i.e. soda and carbon dioxide
  • In this example, bubbles are forced into a solution – once the head pressure matches the vapor pressure of the liquid, the bubbling stops – until you come in to open the vessel
  • With the drop in head pressure, the bubbles come out of the solution to release that lost head pressure
  • We have to consider three different factors when dealing with this simple law:
    • The pressure of the gas surrounding the solution
    • The solubility of the gas in a solvent
    • The temperature of the fluid
  • Formula: V = P x solubility

11:00 Charles’s Law

  • Correlation between volume and temperature
  • Make sure the moles of the product and the pressure are fixed and won’t increase
  • States that volume and temperature are directly connected – if one goes up, so does the other
  • Real-life example: We have a pressurized propane cylinder that’s in a fire – if we didn’t have this stupid law, we’d never have a BLEVE
  • Formula: V = temperature x pressure

12:55 Dalton’s Law

  • Deals with when we have a mixture of different gasses with different properties
  • When we have this mixture, the pressure of the vessel is the sum of all of the pressures of the individual gasses
  • We care about this law because it determines how we choose our PPE
  • Real-life example:
    • If we have a liquid with a low vapor pressure – and all of the other gasses and liquids have a pressure of less than the ambient pressure – there’s a low probability of saturation of the atmosphere
    • This means we don’t have to worry as much about a vapor hazard – we should still be wearing our Level A, though
    • On the other hand, if we have high vapor pressure (and therefore high partial pressure), we probably have a high saturation point in the air
    • This poses a respiratory risk, which means we should strongly consider wearing our Level A
    • Formula: gas 1 pressure + gas 2 pressure + gas 3 pressure, etc.

16:15 Pressure, Temperature, and You

  • We can’t move from one physical state to another without accounting for critical temperature and critical pressure
  • The application of pressure forces gasses into liquids – with enough pressure, this will happen even if the temperature isn’t low enough for it to happen on its own
  • The molecules in the compressed liquid still have the kinetic energy to be a gas – and will spring back into that state if allowed – but for now, the pressure keeps the gas in check
  • The compacting of a gas either by cooling or pressurizing is a way to pack more of it into a cylinder – shown by expansion factors
  • It’s crucial that we understand this fact, because we’re likely to encounter it on-scene

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