### Re: Liquid oxygen/nitrogen stored 120/40 psi, gaseous stored 3000 psi; why/how?

Area: Chemistry
Posted By: Allan Harvey, Staff,National Institute of Standards and Technology
Date: Tue Apr 22 14:29:14 1997
Area of science: Chemistry
ID: 861499456.Ch
Message:

There's really two separate questions here: 1) Why are these substances liquid at certain conditions and gaseous at others?
2) Why are these particular conditions chosen for storage?

For the first question, you need to understand that each substance has what is known as a "vapor-pressure curve". Take water as an example. At a temperature of 100 C, it exerts a vapor pressure of one atmosphere. If the pressure is lower, it boils (becomes a vapor). If the pressure is higher, it stays a liquid. For example, in Colorado where I live the atmospheric pressure is only about 12.5 psia because of the altitude, so the boiling point (temperature needed to exert that much vapor pressure) is only about 96 C. Similarly, what a pressure cooker does is raise the boiling temperature by raising the pressure.

So for any substance, the border between gas and liquid will be a curve of pressure/temperature points. The higher the temperature, the higher the vapor pressure it exerts. This curve is why you can often liquefy a gas either by cooling it or by compressing it.

But there is one more important point, which is that the vapor pressure curve does not go on forever. Again imagine water. As you go up in temperature, the pressure gets higher. So, at the boundary between vapor and liquid, the vapor gets denser. Meanwhile, the liquid at the boundary expands with increasing temperature (it is also being compressed by the extra pressure, but its thermal expansion is bigger). Eventually, at a temperature of 647.096 K and a pressure of 22.064 MPa, the coexisting vapor and liquid become identical, and the vapor pressure curve ends. This is called the "critical point". Above the critical temperature, there is no phase transition; you can just compress and expand and continuously change the density. This condition is not truly gas or liquid; it is often called a "supercritical fluid".

All this about vapor pressures and the critical point can be found in any introductory college thermodynamics textbook, and most freshman chemistry textbooks.

Now we get to your oxygen and nitrogen. The critical point of oxygen is at about 155 K and 5 MPa. The critical point of nitrogen is at about 126 K and 3.4 MPa. The normal boiling points (temperature at which the vapor pressure is one atmosphere) are about 90 K and 77 K for oxygen and nitrogen, respectively. So at the pressures you cite, a rough estimate in my head says that your liquid oxygen would be at or below a temperature of about 110 K, and your liquid nitrogen at or below about 85 K. The key factor is that they are cold.

On the other hand, the gas cylinders are probably at room temperature. This is way above the critical temperature for both fluids, so you will not get a liquid no matter how much pressure you put on it. The gases in the cylinders are supercritical fluids, though when you get that far above the critical temperature their behavior is much like that of a normal gas unless you get to very high pressures.

Now for why they are stored at these conditions. In storing fluids, you want to pack as much as you can into a given volume. For the cold liquids, adding more pressure would not increase the density very much because liquids are nearly incompressible. But the supercritical fluid in the cylinder is very compressible - to a first approximation you would get ten times as much gas in the cylinder at 3000 psi as you would at 300 psi. So it is storage capacity that dictates that gases are often stored at high pressures. The reason the pressure is not even higher is probably safety (and, implicitly, cost), since at higher pressures you would have to use special high-pressure fittings and equipment.

Allan Harvey, aharvey@boulder.nist.gov
"Don't blame the government for what I say, or vice-versa."

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