Partial pressure - please put in simple words

[Texasguy]

Got to admit, I am not totally understanding the term of "partial pressure".

Can someone in layman terms explain partial pressure.

Second, why a change in partial pressure becomes deadly?

How does lethality of other gases change with partial pressure?

Thank you.

 

[Mustard Dave]

Air is roughly 21% O2. Atmospheric pressure is about 1 bar at sea level. The partial pressure of O2 is therefore 0.21 bar and 0.79 bar of nitrogen.

Go down 10 metres and the ambient pressure goes up by 1 bar to 2 bar. 21% of 2 bar is 0.42, so we have a partial pressure of oxygen (ppO2) of 0.42 bar.

At 20 metres, it is 0.63 bar, and so on.

Our tolerance to O2 depends on partial pressure and a ppO2 of 1.4 bar is generally accepted as safe. With air this would be 56m, but richer nitrox mixes would be toxic at a shallower depth.

The effect of nitrogen narcosis increases with elevated ppN2 in the same manner.

 

[sharkbaitDAN]

Partial Pressure is the fraction of a gas in a mix multiplied by the atmospheres the gas is being used at. Equations may or may not help, but here's one anyway:

PP(gas)=Fraction(gas)*P(pressure in atm)

Let's say we have a mixture of 20% oxygen and 80% Nitrogen (0.2 O2 and 0.8 N2). At sea level (1 atm), the fraction will equal the partial pressure. At 33 feet of salt water (10m or 2atm) partial pressure of oxygen is 0.4, or 40%, and partial pressure of nitrogen is 1.6, or 160%.

You can alway take the partial pressure, and divide by atm to determine the fraction of gas in a mix.

0.4(PPo2)/2(atm)=0.2(fraction of oxygen)

If you add up all of the partial pressures in a mix, at a given depth, the total will be equal to the pressure in atm.

0.4(PPO2)+1.6(PPN2)=2

If you go to altitude, say 1/2 atm, then the partial pressure of oxygen would be 0.1, or 10%. It is the reduction in partial pressure that makes you "out of breath" at high altitudes, not a change in the fraction of oxygen, if that makes sense.

I don't really understand the mechanisms behind why high partial pressures of O2 create problems, so I'll leave that for someone more knowledgeable.

In terms of lethality, the only other gas I can think of, off the top of my head is, carbon monoxide. It binds to hemoglobin easily and can displace oxygen. This is a problem at any pressure, but is made worse when partial pressures of CO increase.

Helium also causes problems at very high pressures (HPNS), but I don't know if that's purely because of partial pressures or overall pressure, or some combination. I haven't looked in years, but the last I checked, the physiology behind HPNS was poorly understood, although mitigation measures do exist, and are well documented.

 

[fire_diver]

I don't think it can be explained any simpler than mustard said.

 

[NAM001]

Texas guy
Let me give you an analogy.

You pull up to the gas pump and put in 15 gallons. How much gas is this. The answer is 13.5 gallons of gas and 1.5 of alcohol. You can refer to these numbers as partial volums.

Now when workng with pressure teh reference is an atmosphere. or 14.7psi or atm or bar. Usint the example above (sorta) if you put 15 bar of air in a vessel then in presure terms how much oxygen is in there and how much nitrogen is there. Since O is say 20% of airthen 20% of the15 bar is pure O2 and teh 13.5 bar is the Nitrogen.

now ty this with a tank of trimix of 25% O2 ,,,,,40% Helium ,,,,,and the rest Nitrogen. The gas is being breathed at 200' or in bar it is at 7 bar

What is the partial presure's in bar at 200 feet for O2 He and N2. BTW diving says do not exceed 1.4 bar partial presure of O2.

Partial presure answers should be
7 (atm) times .25 (O2) or 1.75 bar
7 (atm) times .40 (He) or 2.8 bar
7 (atm) times .35 (N2) or 2.45 bar




This more complicated explanation is that the gas you breath is made up of all gasses. Each gas exerts a presure in proportion to the % of that gas compared to the total presureof all gasses combined.

Hope tios works for you.

 

[agilis]

As a gas (like air) is compressed it becomes more dense, meaning that there are more molecules of the gas or gasses in the same size volume. As the number of these molecules increases with additional compression a person breathing that pressurized gas inhales more of these molecules with each breath.

That's why the data provided by Mustard Dave shows a doubling of partial pressures at at 33 feet, or 2 bar (barometric units), a tripling at 66 feet (3 bar), etc. The first stage of your regulator is adjusting the pressure of the gas you are breathing to what is the equivalent of surface pressure, but not surface molecular densities because that is a factor of the physics involved.

The effects of the denser more molecular rich gas under pressure is what can cause Nitrogen narcosis and Oxygen toxicity. You are getting much more of the essence of those gasses when you breath them under pressure. This is of especially vital importance when you are breathing mixed gasses under pressure.

For the sake of clarity, be aware that there is one atmosphere (bar) at the surface, two at 33 feet, three at 66 feet, etc. Some people, for reasons I could never understand, start counting from 'one' at 33 feet. One is at the surface, where we enjoy the benefits of one atmosphere of pressure by obvious definition.

Don't let the techy language confuse you. It's really not very complex.

 

[TSandM]

The core of "partial pressure" is that gases can exist in a mixture (like air, being 21% oxygen and 79% nitrogen) but the gases behave as though they don't know the others are there. So, when you put gases into contact with a liquid, the gases dissolve into the liquid according to the pressure of each individual gas. If you have water sitting in contact with atmospheric air, the amount of oxygen that will dissolve into the liquid is the same amount that would dissolve if you put the water in a container and put only .21 ATA of pure O2 in it.

When you descend, the gas you are breathing is more dense, so any given volume contains a great many more molecules of each gas. That's like putting a gas at a higher pressure in contact with the liquid. So, at 33 feet, you have 21% oxygen, but the total pressure isn't 1 ATA any more, it's two. So the total pressure of oxygen is .42 ATA, and the amount of oxygen that will dissolve into your blood is the amount that would happen if you put pure oxygen at .42 ATA in contact with that blood. Each gas sorts itself the same way; they don't affect one another, so you can consider them as totally independent.

Because of this, you can use a shorthand for describing the concentration of dissolved oxygen (or any gas) in a liquid. You don't have to measure millimoles per liter, or nanograms per deciliter. You can just use the pressure that would generate that concentration, and say, ".42 ATA" instead, and everyone will know what that means.

This gets very important in two ways. One is that oxygen becomes toxic at high concentrations (remember, we will describe concentrations as partial pressures). Above 1.4 ATA, you start to have a risk of seizures. The absolute risk is not well-defined, because it varies between people and in any given person from day to day, but the danger of seizing underwater is so great that we simply arbitrarily set a limit beyond which it is unwise to go, at least for very long.

The other way is that nitrogen at high concentrations is an anesthetic, manifested in recreational depths by nitrogen narcosis. You have to have a moderate concentration of nitrogen to notice the effects -- but again, we describe concentrations of gases as "partial pressures", meaning the fraction of the total pressure represented by THAT gas. A lot of people like to keep their nitrogen concentrations at, or lower than that of air at 100 feet, which is about 3.2 ATA.

Did that help?

 

[boulderjohn]

I think Agilis did the right thing in getting you to think in terms of the number of molecules.

Lets say that a lung full of air at sea level has 1 million molecules. Since air is about 21% oxygen, then about 21% of those molecules are oxygen. The total pressure of the air in your lungs is 1 atmosphere, and that total pressure is made up of all the gases in the mix. That means the partial pressure of oxygen is 0.21.

Go down to 33 feet of sea water, and now you are inhaling twice as many molecules with each breath--2 million molecules. Now the total pressure of the gases is 2.0 atmospheres. Oxygen is still 21% of that, but since there are 2 atmospheres instead of 1, the partial pressure of oxygen is 0.42.

Go to 5 atmospheres of pressure, and the partial pressure of oxygen is now 5 X 0.21. or 1.05. That means that if you are breathing air at 5 atmospheres, each lungful contains more oxygen molecules than if you were breathing pure oxygen at sea level.

 

[jbark]

. . . So the total pressure of oxygen is .42 ATA, and the amount of oxygen that will dissolve into your blood is the amount that would happen if you put pure oxygen at .42 ATA in contact with that blood. . .


Did that help?

YES.

There is the key to all of this, IMO. I knew what "partial pressure" means to a chemist or physicist, and it means the same thing here. But, TSandM just taught me a very valuable lesson: why do I need to know the partial pressure and how does the partial pressure affect my body.

Thanks!
Jerry

 

[kensuf]

In simple terms, I like to explain it as a "dose" of a gas, and the dose of the gas is dependent upon the % of the gas in the mixture you are breathing multiplied by the pressure you are under. Mustard did a good example of how that goes up. I will frequently draw a box of "air" in the surface (with red O2 molecules and green N2 molecules) then the same box at 4ATA highlighting the increased molecules.

The ramifications for diving are simple, just like with medicine, there are safe and unsafe doses of a gas, and the negative impact of the doses are time dependent as well. I can take the contents of a bottle of advil in regularly spaced intervals over the course of a month, and I may be fine, but if I took a bottle of advil all at once I would have problems.