Knowing the properties of gas is very important for anybody who wishes to use a gas or gas mix in any form. The physical properties of gases can change depending on its surrounding conditions. For instance, pressure is one of these conditions that can change the properties of a gas.

If you are working with a flammable gas, an increase of pressure can cause it to explode. If however you are working with a relatively inert gas an increase of pressure will increase the gas concentration and its pressure, whilst reducing its volume.

### An Ideal Gas

If you assume a gas is made up of individual gas atoms which do not react with each other and the gas particle distribution in an area is perfectly uniform, then you can apply the ideal gas law to calculate what effect changing pressure, temperature or concentration will have on the other factors.

### Real Gas

In reality, gases do not follow this idealism quite so perfectly, but the rule does hold up to give a very good approximation of a real gas.

### Gas Mixes

This is all very well for a pure gas, but what if you want to mix two gasses or more with varying amounts of each gas depending on what you want to do with it. How can you tell what pressure each individual gas is exerting? This is important to know in various things e.g. SCUBA diving. The partial pressure of a gas in the body can equate to the relative amount of that gas that is dissolved in the body. When a diver uses gas mixes for breathing, they must calculate the partial pressure of each gas at various pressures (depths) very carefully if they want to come back alive.

### Calculations

The total pressure of a gas can be subdivided into the pressure exhibited by each gas within the mixture. For instance if you have a bottle of gas and the pressure is 20 bars and there are 2 gasses, each in equal amounts then the partial pressure of each gas will be 10 bars (1 bar = 1 atmosphere).

There are two ways to find the partial pressure of any gas. If you know the combined partial pressures of all the other gases and know the total pressure, then the difference will be the partial pressure of the gas you want to know. The problem with the first way is that if any condition changes, then so do the values of all the other variables, so a better way must be found. This is done by modifying the ideal gas law.

1) Partial pressure gas law
pP = nRT / V

2) Partial pressures directly relate to total pressure
P = pP + pP + pP

Key:

• P = total pressure.
• pP = partial pressure of the gas.
• n = moles of gas present (equates to concentration).
• R = Avagadros constant (a number which does not change but needs to be there to make the equation work.
• T = Temperature (Celsius).
• V = Volume (litres).

### Application of these Partial Pressure Laws

I shall use the diving example to explain this law (rather than being really vague). If you fill a cylinder with gas, you know at what pressure it is at by looking at the gas gauge, you also know the concentration of each gas (you might have altered this depending on what kind of dive you are doing) and finally you know the volume of the container e.g. 15L cylinder. You now have a known gas with known values in which the temperature can be calculated (gas temp. is not very important in diving as it tends to be fairly constant underwater).

When breathing, you breathe this compressed air at atmospheric pressure; as such the pressure must drop along with the partial pressure. To calculate partial pressure from this point onwards can be done quite simply by using the second equation above. E.g. in compressed air, oxygen partial pressure is 0.21 Bar (21% of air is oxygen). If the total pressure doubles (depth), then the partial pressure will double, if the partial pressure then reaches 1.6 you die. A slight simlification, you actually get an oxygen hit which can be fatal.

As you can see, these calculations are not that hard but are very necessary. Anotherther area in which these type of equations are used is in chemical plants where attaining the exact conditions is very important in attaining the optimum reaction speed. The equations used in this example are a little more complex due to the dynamic equilibrium between reactants and products being considered; however they have exactly the same basis and are bound by the same laws.

Bibliography
All info. is from memory. Original info. will be either from my Chemistry lectures or BSAC training lectures.

Dude, I hate to be somewhat anal but you don't die instantly when the partial pressure of O2 gets to 1.6 bar.

The body is designed to take a partial pressure of about 0.21 bar. Subjecting it to higher partial pressures causes complex physiological changes to the way the body works. These changes and effects of the exposure is know as oxygen toxicity.

### What Happens When the Partial Pressure of O2 Gets Too High

There are two types of adverse reaction one, the first is caused by the brain basically going into epileptic fit when exposed to high partial pressures over 1.6 bar (the classic O2 hit ). The fit itself is not dangerous but as the victim is under water trying to hold a regulator between their teeth, the danger of losing consiousness and dropping the regulator and drowning is somewhat high.

The second is caused by caissons disease above 0.21 bar (Pulmonary oxygen toxicity). The effect of this is the lungs start to ache and then burn and respiration becomes somewhat difficult, eventually the lining of the lung itself starts to decompose.

Both of these toxic reactions will cure themselves if the body is returned to the atmospheric pressure it is designed for ( i.e. you get out of the water). The only problem is both the hits and the O2 pulmonary toxicity are not easy to predict and each individual is totally random in how resistant they are to the effects. Being fit young and healthy is no real indicator either, as it is also dependent on the individual's physiological makeup and frame of mind.

### The Influence of Nitrox

When the first nitrox tables were created by Haldane for the mine clearance divers on the D-Day landing beaches he recommended the partial pressure of O2 to be no higher than 2.0 bar. Subsequent table designs have reduced this over the years.

### Today

The modern recommendation from the RSTA is 1.4 bar and a absolute maximum of 1.6 bar (this is working on the assumption that some fool will just try and break the rules if they are written down so they recommend as low as possible). A working pressure of 1.2 bar is recommended as best practise best benefits of O2 for metabolism and the least danger of getting a 'hit'.