GAS MIXTURES

2-12 GAS MIXTURES

Dalton’s Law

Expressing Small Quantities of Pressure

Calculating Surface Equivalent Value

Gas Diffusion

Humidity

Gases in Liquids

Solubility

Henry’s Law

Gas Tension

Gas Absorption

Gas Solubility

If a diver used only one gas for all underwater work, at all depths, then the general gas law would suffice for most of his necessary calculations. However, to accommodate use of a single gas, oxygen would have to be chosen because it is the only one that provides life support. But 100 percent oxygen can be dangerous to a diver as depth and breathing time increase. Divers usually breathe gases in a mixture, either air (21 percent oxygen, 78 percent nitrogen, 1 percent other gases) or oxygen with one of the inert gases serving as a diluent for the oxygen. The human body has a wide range of reactions to various gases under different conditions of pressure and for this reason another gas law is required to help compute the differences between breathing at the surface and breathing under pressure.

Dalton’s law states: “The total pressure exerted by a mixture of gases is equal to the sum of the pressures of each of the different gases making up the mixture, with each gas acting as if it alone was present and occupied the total volume.”

In a gas mixture, the portion of the total pressure contributed by a single gas is called the partial pressure (pp) of that gas. An easily understood example is that of a container at atmospheric pressure (14.7 psi). If the container were filled with oxygen alone, the partial pressure of the oxygen would be one atmosphere. If the same container at 1 atm were filled with dry air, the partial pressures of all the constituent gases would contribute to the total partial pressure, as shown in Table 2-3.

If the same container was filled with air to 2,000 psi (137 ata), the partial pressures of the various components would reflect the increased pressure in the same proportion as their percentage of the gas, as illustrated in Table 2-4.

Table 2-3. Partial Pressure at 1 ata.

Table 2-4. Partial Pressure at 137 ata.

The formula for expressing Dalton’s law is:

P_TOTAL

pp_A

pp_B

pp_C

...

Where: A, B, and C are gases and

Another method of arriving at the same conclusion is to use the T formula. When using the T formula, there can be only one unknown value. Then it is merely a case of multiplying across, or dividing up to solve for the unknown value.The T formula is illustrated as:

partial pressure

________________________________________________

atmosphere(s) absolute / % volume (in decimal form)

Sample Problem 1. Use the T formula to calculate oxygen partial pressure given 10 ata and 16 percent oxygen.

1. Fill in the known values:

		pp
		__________
		10/.16

2. Multiply the pressure by the volume to solve for the oxygen partial pressure (pp):

		1.6ppO₂
		__________
		10/.16

The oxygen partial pressure is 1.6.

Sample Problem 2. What happens to the breathing mixture at the operating depth of 130 fsw (4.93 ata)? The air compressor on the ship is taking in air at the surface, at normal pressure and normal mixture, and sending it to the diver at pressure sufficient to provide the necessary balance. The composition of air is not changed, but the quantity being delivered to the diver is five times what he was breathing on the surface. More molecules of oxygen, nitrogen, and carbon dioxide are all compressed into the same volume at the higher pressure. Use Dalton’s law to determine the partial pressures at depth.

1. Calculate the oxygen partial pressure at depth.

ppO₂	=	.21(surface)	X	4.93 ata

	=	1.03 ata

2. Calculate the nitrogen partial pressure at depth.

ppN₂	=	.79(surface)	X	4.93 ata

	=	3.89 ata

3. Calculate the carbon dioxide partial pressure at depth.

ppCO₂	=	.0003(surface)	X	4.93 ata

	=	.0014 ata

Expressing partial pressures of gases in atmospheres absolute (ata) is the most common method employed in large quantities of pressure. Partial pressures of less than 0.1 atmosphere are usually expressed in millimeters of mercury (mmHg). At the surface, atmospheric pressure is equal to 1 ata or 14.7 psia or 760 mmHg. The formula used to calculate the ppCO₂ at 130 fsw in millimeters of mercury is:

0.03

760mmHg

ppCO₂

------------

4.93 ata

------------

100

1ata

1.12mmHg

From the previous calculations, it is apparent that the diver is breathing more molecules of oxygen breathing air at 130 fsw than he would be if using 100 percent oxygen at the surface. He is also inspiring five times as many carbon dioxide molecules as he would breathing normal air on the surface. If the surface air were contaminated with 2 percent (0.02 ata) carbon dioxide, a level that could be readily accommodated by a normal person at one ata, the partial pressure at depth would be dangerously high—0.0986 ata (0.02 x 4.93 ata). This partial pressure is commonly referred to as a surface equivalent value (sev) of 10 percent carbon dioxide. The formula for calculating the surface equivalent value is:

		pp at depth (in ata) X 100%
sev	=	---------------------------------------------
		1 ata

		0.0986 ata
	=	___________	X	100%
		1 ata

	=	9.86% CO₂

Another physical effect of partial pressures and kinetic activity is that of gas diffusion. Gas diffusion is the process of intermingling or mixing of gas molecules. If two gases are placed together in a container, they will eventually mix completely even though one gas may be heavier. The mixing occurs as a result of constant molecular motion.

An individual gas will move through a permeable membrane (a solid that permits molecular transmission) depending upon the partial pressure of the gas on each side of the membrane. If the partial pressure is higher on one side, the gas molecules will diffuse through the membrane from the higher to the lower partial pressure side until the partial pressure on sides of the membrane are equal. Molecules are actually passing through the membrane at all times in both directions due to kinetic activity, but more will move from the side of higher concentration to the side of lower concentration.

Body tissues are permeable membranes. The rate of gas diffusion, which is related to the difference in partial pressures, is an important consideration in determining the uptake and elimination of gases in calculating decompression tables.

Humidity is the amount of water vapor in gaseous atmospheres. Like other gases, water vapor behaves in accordance with the gas laws. However, unlike other gases encountered in diving, water vapor condenses to its liquid state at temperatures normally encountered by man.

Humidity is related to the vapor pressure of water, and the maximum partial pressure of water vapor in the gas is governed entirely by the temperature of the gas. As the gas temperature increases, more molecules of water can be maintained in the gas until a new equilibrium condition and higher maximum partial pressure are established. As a gas cools, water vapor in the gas condenses until a lower partial pressure condition exists regardless of the total pressure of the gas. The temperature at which a gas is saturated with water vapor is called the dewpoint.

In proper concentrations, water vapor in a diver’s breathing gas can be beneficial to the diver. Water vapor moistens body tissues, thus keeping the diver comfortable. As a condensing liquid, however, water vapor can freeze and block air passageways in hoses and equipment, fog a diver’s faceplate, and corrode his equipment.

When a gas comes in contact with a liquid, a portion of the gas molecules enters into solution with the liquid. The gas is said to be dissolved in the liquid. Solubility is vitally important because significant amounts of gases are dissolved in body tissues at the pressures encountered in diving.

Some gases are more soluble (capable of being dissolved) than others, and some liquids and substances are better solvents (capable of dissolving another substance) than others. For example, nitrogen is five times more soluble in fat than it is in water.

Apart from the individual characteristics of the various gases and liquids, temperature and pressure greatly affect the quantity of gas that will be absorbed. Because a diver is always operating under unusual conditions of pressure, understanding this factor is particularly important.

Henry’s law states: “The amount of any given gas that will dissolve in a liquid at a given temperature is directly proportional to the partial pressure of that gas.” Because a large percentage of the human body is water, the law simply states that as one dives deeper and deeper, more gas will dissolve in the body tissues and that upon ascent, the dissolved gas must be released.

When a gas-free liquid is first exposed to a gas, quantities of gas molecules rush to enter the solution, pushed along by the partial pressure of the gas. As the molecules enter the liquid, they add to a state of gas tension. Gas tension is a way of identifying the partial pressure of that gas in the liquid.

The difference between the gas tension and the partial pressure of the gas outside the liquid is called the pressure gradient. The pressure gradient indicates the rate at which the gas enters or leaves the solution.

At sea level, the body tissues are equilibrated with dissolved nitrogen at a partial pressure equal to the partial pressure of nitrogen in the lungs. Upon exposure to altitude or increased pressure in diving, the partial pressure of nitrogen in the lungs changes and tissues either lose or gain nitrogen to reach a new equilibrium with the nitrogen pressure in the lungs. Taking up nitrogen in tissues is called absorption or uptake. Giving up nitrogen from tissues is termed elimination or offgassing. In air diving, nitrogen absorption occurs when a diver is exposed to an increased nitrogen partial pressure. As pressure decreases, the nitrogen is eliminated. This is true for any inert gas breathed.

Absorption consists of several phases, including transfer of inert gas from the lungs to the blood and then from the blood to the various tissues as it flows through the body. The gradient for gas transfer is the partial pressure difference of the gas between the lungs and blood and between the blood and the tissues.

The volume of blood flowing through tissues is small compared to the mass of the tissue, but over a period of time the gas delivered to the tissue causes it to become equilibrated with the gas carried in the blood. As the number of gas molecules in the liquid increases, the tension increases until it reaches a value equal to the partial pressure. When the tension equals the partial pressure, the liquid is saturated with the gas and the pressure gradient is zero. Unless the temperature or pressure changes, the only molecules of gas to enter or leave the liquid are those which may, in random fashion, change places without altering the balance.

The rate of equilibration with the blood gas depends upon the volume of blood flow and the respective capacities of blood and tissues to absorb dissolved gas. For example, fatty tissues hold significantly more gas than watery tissues and will thus take longer to absorb or eliminate excess inert gas.

P_TOTAL X % Vol_A

------------------------

The solubility of gases is affected by temperature—the lower the temperature, the higher the solubility. As the temperature of a solution increases, some of the dissolved gas leaves the solution. The bubbles rising in a pan of water being heated (long before it boils) are bubbles of dissolved gas coming out of solution.

The gases in a diver’s breathing mixture are dissolved into his body in proportion to the partial pressure of each gas in the mixture. Because of the varied solubility of different gases, the quantity of a particular gas that becomes dissolved is also governed by the length of time the diver is breathing the gas at the increased pressure. If the diver breathes the gas long enough, his body will become saturated.

The dissolved gas in a diver’s body, regardless of quantity, depth, or pressure, remains in solution as long as the pressure is maintained. However, as the diver ascends, more and more of the dissolved gas comes out of solution. If his ascent rate is controlled (i.e., through the use of the decompression tables), the dissolved gas is carried to the lungs and exhaled before it accumulates to form significant bubbles in the tissues. If, on the other hand, he ascends suddenly and the pressure is reduced at a rate higher than the body can accommodate, bubbles may form, disrupt body tissues and systems, and produce decompression sickness.