Loss of CO2 pressure

[mookie1010]

So I've been through a few 5lb tanks of co2 lately, and at first I assumed it was a leak in a keg or 2 that I had. I had one hooked up to the two kegs I have on tap now and it was maintaining pressure fine for several weeks. A couple of days ago I realized that I had somehow shimmied the regulator loose while cleaning that fridge last week. I didn't notice until 2 nights ago that it had completely drained empty. So I exchanged it for a new one yesterday.

Here's my problem / question - I exchanged it over lunch so it sat in my car all afternoon at approx 55 degrees. I got it home, attached the regulator and turned it on - it jumped right to 700PSI in the tank and 5 PSI delivered to the kegs for dispensing.

I keep the tank in the fridge with the kegs and over the course of the night (proabably over 4 hours or so) it dipped to 600PSI in the tank. There is no audible or visible leak, all gaskets are less than 4 months old and were repalced by me. All connections have been sprayed with water to check visibly for leaks. I'm 98% certain that it's not leaking, but it sure seems like there's a leak somewhere.

Is the change in temp responsible for the change in pressure? I turned it off last night so I wouldn't lose any more pressure in the event that I have a leak I haven't found and because I don't want to take the 20 mile drive to get another tank if I can avoid it.

Ideas? Garcias in advance

[Travisty]

Is the change in temp responsible for the change in pressure?

Yup. No worries, it's normal.

[TapItGood]

I see the same thing when I refill my tank. When it's warmer it shows a fuller tank. As it cools the level drops. That leads me to believe the gas compresses as it cools. Any rocket scientists want to elaborate?

[Travisty]

Gas pressure is directly related to temperature. Basically the colder the temperature the more the CO2 will want to stay in liquid form and the warmer the temerature, the more the CO2 will want to be in gaseous form. It's kinda like water in a pressure cooker. The pressure shown on the regulator is not really good for telling exactly how much CO2 is left since at a given temperature the pressure should remain the same until pretty much all the liquid CO2 has evaporated. Then the pressure will drop quickly. So, if the pressure starts dropping and the tank is at a constant temperature, you know the tank is near empty.

[ajdelange]

If you were to place an evacualted CO2 bottle into a constant temperature water bath at temperature T and start filling it with CO2 the gas would flow into the cylinder and the pressure increase proportional to the amount of gas pumped in up to a pressure given by

Pvap = 14.7*exp( -24.03761*ln(T) -7062.404/T +166.3861 +3.368548e-5*T*T)/101.325 -14.7 psig

where T is the temperature in Kelvins (add 273.15 to °C to get K). Thereafter the pressure will increase no further and the gas that flows into the cylinder will condense into liquid. At 0° F this pressure is 303 psig, at 10°F it is 362 psig, at 20 °F 426 psig and so on. At room temperature, 70°F, it is 823 psig and at 87.8 °F it is 993.15 psig. If the temperature is above 87.8 °F, which is called the critical temperature for CO2, condensation will no longer take place and gas which flows into the cylinder will thus continue to increase the pressure until the bottle bursts. Thus as long as the temperature is below 87.8 and there is enough gas to have raised the pressure to the vapor pressure (that is what the pressure calculated by the formula is called) for the particular temperature the pressure in the tank will depend only on the temperature independent of the amount of gas actually in the bottle. Take it outside on an 80 ° day and the gauge will read 917 psig. Put it in the freezer at 0 ° and the gauge will read 362.

As you draw off gas the liquid vaporizes. Once all the liquid is gone the pressure read on the gauge depends on the amount of gas left in the bottle and the temperature according to the gas law P = nRT/V where n is the amount of CO2 (number of moles) in the volume, V, T is the temperature (Kelvins) and R the universal gas constant. Above 87.8 the pressure again depends on the temperature and pressure again according to the gas law. Actually in neither case is the gas law strictly adhered to - CO2 is not an ideal gas.

Vapor pressure is simply the pressure of the gaseous form of a liquid over the liquid. Water has vapor pressure which is only a function of temperature. At 100 °C the vapor pressure of water is 1 atmosphere so that when water reaches that temperature it begins to boil. At 90°C the vapor pressure of water is lower than 1 atmosphere and, if you take 90 ° water far enough up a mountain to reach that lower vapor pressure it will boil at 90 °C.

[Bugeater]

Ouch! My brain got a cramp reading that. But yeah, that agrees with what I learned in school several decades ago.

The temperature increase during compression can be significant. When I used to scuba dive, we always had to put the tanks in water to dissipate the heat as we filled them. But then we were also going up to 3000 psi.

Wayne

[Kbar]

Charles's Law.

Keep it simple.............

[Kbar]

Edit.

Ideal gas law - Isochoric process - Constant Volume, p2 = p1(T2/T1)

Gay-Lussac's law