Methane? how much will it expand if we vapourize it

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When liquid methane at -165 degrees Celsius and 1 atm is vaporized, it expands significantly, with estimates suggesting a volume increase of around 600 times. To determine the exact expansion, one must use the ideal gas law (PV = nRT) and compare the densities of liquid and gaseous methane. The density of liquid methane is crucial for calculating the number of moles present in a given volume, which can then be used to find the volume of methane gas at standard conditions. The discussion emphasizes the importance of temperature and pressure in these calculations, particularly noting that methane's behavior as an ideal gas can guide the volume estimation. Ultimately, the expansion of methane when vaporized depends on the specific conditions of the system.
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If we have 1m3 of methane liquid at 1 atm and -165 degrees, how much will it expand if we vapourize it?
I have heard someone say that it expand with 600 times.
 
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That's intreguing isn't it? How would you go about checking that?
What is the property that relates how much of something you have to the volume it occupies?
Is there something special about the temperature of -165degrees (units?)
 


All gases have approximately the same volume per molecule. Use the ideal gas law: PV = nRT.
You'll need to look up the density of liquid methane, then just compare densities. Of course, your answer depends on the temperature and pressure of the gas.
 


Simon Bridge said:
That's intreguing isn't it? How would you go about checking that?
What is the property that relates how much of something you have to the volume it occupies?
Is there something special about the temperature of -165degrees (units?)

If you want to have methane in liquid at 1 bar, the temperature needs to be below -160 degrees. Thats why a said -165 deg.
So if we for example have 1 liter(1dm3) of methane(liquid), how much volume will it uccupy if it goes from liquid-> gas ?
 


Khashishi said:
All gases have approximately the same volume per molecule. Use the ideal gas law: PV = nRT.
You'll need to look up the density of liquid methane, then just compare densities. Of course, your answer depends on the temperature and pressure of the gas.

But if I have 1dm3 of methane in liquid state ( 1 bar, -165 degrees), how much volume will it occupy if we vapourise it?
 


charlie95 said:
If you want to have methane in liquid at 1 bar, the temperature needs to be below -160 degrees. Thats why a said -165 deg.
So if we for example have 1 liter(1dm3) of methane(liquid), how much volume will it uccupy if it goes from liquid-> gas ?
You can answer your own question when you answer the others in my reply. hint: try the second one.

You seem to be interested in comparing the volume of liquid methane at (or near) it's boiling point at 1 bar with that of methane gas, at the same temperature at 1 bar.

If you read carefully you will find that Khashishi has actually already answered them for you.
Please understand that we do not usually just provide answers for you in PF - rather we try to empower you to find the answers yourself.
 
If I read carefully: PV=nRT... and compare density. Where can I find a table that show me the density at different temperatures?

I just want to know if you spill 1 liter of methane on the ground( or liquified natural gas) how much will it expand when it goes from liquid to gas? I know that I should try to find the answer myself, but I wouldn't ask if I hadnt tried..!
 
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Look up the density of liquid methane on google. Use this to calculate the weight of the methane in 1 cubic meter. Calculate the number of moles of methane for this weight of methane. Use the ideal gas law to calculate the volume of that number of moles of methane at 1 atm and its atmospheric boiling point -161.5C.
 
  • #10
charlie95 said:
I just want to know if you spill 1 liter of methane on the ground( or liquified natural gas) how much will it expand when it goes from liquid to gas? I know that I should try to find the answer myself, but I wouldn't ask if I hadnt tried..!

You know that a gas does not have a proper volume. It expands as much as is allowed to by the container enclosing the gas.
So if you just spill it on the ground, in the open, then I am afraid there is no specific answer to your question. It does not expand to a specific value of the volume.

Assuming that you do it in a closed chamber instead, then the vapors will occupy the whole volume of the chamber and the answer will depend on the volume of this chamber.
Of course, if the volume is limited in this way, evaporation will reach an equilibrium when the vapor pressure is reached in the chamber.
 
  • #11
charlie95 said:
If I read carefully: PV=nRT... and compare density. Where can I find a table that show me the density at different temperatures?

I just want to know if you spill 1 liter of methane on the ground( or liquified natural gas) how much will it expand when it goes from liquid to gas? I know that I should try to find the answer myself, but I wouldn't ask if I hadnt tried..!
Did you try googling "density of methane" and such things?
You need to say what you have tried. From the information suppied to that point, it looked like you may not have realized about density. There are tables of densities for all kinds of substances online.

Of course, if you just spill it on the floor, the gas will eventually expand to fill the whole room, mixing with the air. I suspect that the situation you need is where the liquid is in a cylinder under a piston. The piston is compressed to maintain a constant pressure, The liquid is heated at it's boiling point at constant temperature until it has completely changed state, What is the new volume?
Right?

Or do you just want to know how far a spill will spread?
 
  • #12
Simon Bridge said:
Did you try googling "density of methane" and such things?
You need to say what you have tried. From the information suppied to that point, it looked like you may not have realized about density. There are tables of densities for all kinds of substances online.

Of course, if you just spill it on the floor, the gas will eventually expand to fill the whole room, mixing with the air. I suspect that the situation you need is where the liquid is in a cylinder under a piston. The piston is compressed to maintain a constant pressure, The liquid is heated at it's boiling point at constant temperature until it has completely changed state, What is the new volume?
Right?

Or do you just want to know how far a spill will spread?

-------The liquid is heated at it's boiling point at constant temperature until it has completely changed state, What is the new volume?----- YES.
 
  • #13
charlie95 said:
-------The liquid is heated at it's boiling point at constant temperature until it has completely changed state, What is the new volume?----- YES.

So how do I do it ?
 
  • #14
charlie95 said:
-------The liquid is heated at it's boiling point at constant temperature until it has completely changed state, What is the new volume?----- YES.

The volume of the container in which the gas is, well, contained.
A gas has no proper volume to talk about.
 
  • #15
charlie95
At STP ( standard temperature and pressure ) 1 mole of an ideal gas will occupy a volume of 22.4 litres, which comes from the ideal gas law.

To find the volume at different temperatures or pressures you would the same deal gas law. that has already been discussed.

so, assuming that methane behaves as an ideal gas, you want to find the volume of a container that would hold methane gas if it changed state from a liquid.
find the mass of 1 cubic meter of methane and convert that to number of moles of methane, and use PV = nRT.
 
  • #16
256bits said:
charlie95
At STP ( standard temperature and pressure ) 1 mole of an ideal gas will occupy a volume of 22.4 litres, which comes from the ideal gas law.

To find the volume at different temperatures or pressures you would the same deal gas law. that has already been discussed.

so, assuming that methane behaves as an ideal gas, you want to find the volume of a container that would hold methane gas if it changed state from a liquid.
find the mass of 1 cubic meter of methane and convert that to number of moles of methane, and use PV = nRT.

I got the Volume t be 624m3.. Is this correct?
V=(26430mol*8,314J/K*mol*288K)/atmos pressure = 624 m3

I assumed that the temperature was 15 degrees.
 
  • #17
charlie95 said:
I got the Volume t be 624m3.. Is this correct?
V=(26430mol*8,314J/K*mol*288K)/atmos pressure = 624 m3

I assumed that the temperature was 15 degrees.

But this does not correspond to your proposed situation, with liquid spilled on the floor. It may give some idea about the order of magnitude, maybe.

Even if you evaporate methane in the atmosphere, the partial pressure of the methane is not necessarily equal to the atmospheric pressure. What constraint will make the vapors of methane to have atmospheric pressure?
 

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