High purity gas systems: Small Leaks = Big Problems
05/01/2003
By Mike Fitzpatrick and Ken Goldstein, Ph.D.
When end users discuss high-purity gas systems, it's nearly impossible not to touch upon the considerable effort and expense required to install them. Designers, component manufacturers and installing contractors go to extraordinary lengths to minimize all potential sources of contamination in these systems.
Acceptable leak rates are often specified to not exceed 10-9 atm—cc/sec—or, one billionth of a cubic centimeter per second—at one atmosphere of differential pressure between the system interior and the external ambient.
 Mike Fitzpatrick |
If we were to bubble test a leak of this magnitude, we would have to wait several days for a visible bubble to form. Obviously, our testing solution would dry out long before that. The bottom line: Locating and repairing leaks of this magnitude can be time-consuming, expensive and frustrating.
 Ken Goldstein, Ph.D. |
Let's assume we're dealing with a non-toxic and biologically inert gas such as nitrogen. If the system is positively pressurized with respect to the ambient, the leak will occur outward. And who cares if there is a trace amount of a biologically inert, non-toxic gas in our facility? With or without leaks, there is plenty of nitrogen all around us. So, why bother with these very small leaks?
Don't forget, not all of the leakage is outward. Some of the critical leakage will be inward, from the ambient to the inside of our system. But how can this be if the system is positively pressurized?
The answer is "partial pressure." Before proceeding, let's clearly state that the positive pressure inside the system really does mean that the net flow of gas leakage will be outward. Indeed, the total pressure difference drives this net movement, since all fluids will flow from high pressure to low pressure areas. It's possible, however, for individual gas molecules to move against the tide and travel from the outside to the inside. Partial pressure differences make this happen. Let's use some simple examples to demonstrate the importance of partial pressure. First, let's explain the term "partial pressure."
Consider the air all around us. It's not a single gas, but a mixture of separate gases—all peacefully coexisting in our atmosphere. What is the pressure of our atmosphere? Let's call it one atmosphere, or around 15 psi. While not exactly correct, this is close enough for our purposes and makes the arithmetic a bit simpler.
Again, to keep things uncomplicated, let's assume that nitrogen makes up 80 percent of our atmosphere, oxygen 16 percent, carbon dioxide three percent, and water vapor one percent. Yes, there are other trace constituents in the air, but because their concentrations are low, let's ignore them. These numbers should be close enough for our purposes.
Dalton's Law tells us that the total pressure of our air (one atmosphere, or 15 psi) is actually the sum of the pressures of individual components of the air. In other words, our one atmosphere is really the sum of the pressure of the nitrogen in the air, plus the pressure of the oxygen in the air, plus the pressure of the carbon dioxide in the air, plus the pressure of the water vapor in our air. In equation form, we might write:
PT = PN2 + PO2 + PCO2 + PH2O where
PT is the total pressure of the air (one atmosphere, or 15 psi) and
PN2, POv, PCO2 and PH2O are the individual pressures of the nitrogen, oxygen, etc.
Now, this does not seem to help us much unless we can calculate the individual pressure of the air's constituents. It turns out that we can get close. For each of the constituent gases, the individual pressure of that gas is approximately equal to the product of the concentration of that gas multiplied by the total pressure. So, in our example:
- The total pressure (PT) is 15 psi and, therefore, the nitrogen pressure (PN2) is equal to
If we add these all together (12.0 + 2.4 + 0.45 + 0.15), we get 15.0 psi, or 100% of PT as we must, considering where the numbers came from. Each of these individual pressures is called the partial pressure of that gas.
Now, let's use this approach to examine a high purity gas line operating at a "high" pressure of 100 psi. And let's assume that there is only one impurity in this gas: water vapor at a concentration of 0.1 parts per million (ppm). What is the partial pressure of the water vapor in the nitrogen line? Using the same approach, we can estimate this as the concentration of water vapor (0.1 ppm, or 0.00001 percent) multiplied by the total pressure (100 psi), giving us 0.00001 psi.
Finally, let's compare the partial pressure of water vapor in the air to the partial pressure of water vapor in our nitrogen line. Recall that in air, the partial pressure of water vapor was 0.15 psi, while in our "high" pressure (100 psi) nitrogen line, it was 0.00001 psi. Which is higher? Clearly, the partial pressure of water vapor in ambient air is 15,000 times greater than the partial pressure of water vapor in our high purity line. This high-pressure differential will drive water vapor (H2O molecules) from the air into our high purity line. A little thought will quickly convince the reader that the higher the purity level of our process gas, the higher the differential driving pressure.
Because partial pressure differences will always be with us, the only solution to this apparent problem is to find and eliminate all leaks. Nothing else will do. We cannot count on the higher total pressure of our lines to keep out contaminants.
Michael Fitzpatrick is a Senior Member of the Institute of Environmental Sciences and Technology. Mike can be reached at [email protected].
Ken Goldstein is principal of Cleanroom Consultants Inc. in Phoenix, Ariz., and is a member of the CleanRooms Editorial Advisory Board. He can be reached at [email protected].