It is all about science – in particular, the three laws of physics known as the Universal Gas Laws: Boyle’s, Dalton’s and Henry’s laws. Linked at all times together, they become Universal. Alter or tweak any one and the other two adjust to keep pace.
Let’s look at the three laws, one by one:
‘At a constant temperature the volume of a gas will vary inversely with the absolute pressure’.
Creatures that live at or above sea level do not notice the weight of the atmosphere upon them because they reach atmospheric equilibrium at birth. Their bodily airspaces – lungs, sinus cavities, middle ear – equalize to the surrounding or ambient atmospheric pressure. Bodily fluids reach equilibrium with gases during gestation. Only when altitude changes occur rapidly is the weight (pressure) of gas obvious.
Example 1: When a flexible balloon is taken beneath the surface of the ocean, the volume (size) decreases proportionately to the depth. The gas inside the balloon is compressed by the weight of the seawater pressing on the surface of the balloon. The gas becomes more dense as the pressure increases around it. The pressure within the balloon equals the ambient pressure of the seawater surrounding it.
Example 2: Gas volumes expand when pressure drops. During decompression (or surfacing from depth), gas in any air cavity will expand and unless vented (equalized) will cause damage to the surrounding tissue. The same volume changes with pressure occur in bubbles in tissue or blood.
During compression in the hyperbaric oxygen therapy chamber, the pressure exerted throughout the body to tissues surrounding gas spaces increases.
The pressure inside the gas spaces must also be elevated to equal the surrounding pressure. Failure to do so would result in squeezes or barotrauma, the name generally given to any injury caused by pressure changes.
‘The total pressure of a mixture of gases is equal to the sum of the partial pressure of each gas, with each acting as if it were the only gas occupying the total volume’.
Dalton’s Law helps us to understand how gas is exchanged at various locations around the body. Whether present in air or dissolved in a solution, gas depends on differences of quantities known as partial pressures.
At sea level, air exerts a downward force of 760 mm Hg. Since air is a mixture of gases Oxygen (02 78%), nitrogen (N2 21%), carbon dioxide and inert gases. Therefore the partial pressure of oxygen is 160 mm Hg. This figure is arrived at by multiplying the percentage of oxygen (by volume) by atmospheric pressure or,0.21 X 760 = 160. This is the portion of atmospheric pressure that is contributed to by Oxygen thus the term partial pressure. This partial pressure is symbolized by P 02. The partial pressure of carbon dioxide (C02) is only 0,23 mm Hg. It’s partial pressure is symbolized by P C02.
Example: During respiration water vapour combines with air and makes up 47 mm Hg of the total alveolar pressure. The carbon dioxide partial pressure during alveolar respiration (PA C02) is 38 – 40 mm Hg. Inspiration of 100% of oxygen 02 at sea level, the partial pressure of alveolar nitrogen (PA N2) is negligible.
According to Daltons’ Law, the partial pressure of oxygen, carbon dioxide and water vapour are added to equal the total pressure.
Therefore, if 100% oxygen is breathed at sea level, the partial pressure of alveolar oxygen (PA 02) is 673 mm Hg.
When gas is breathed under increased pressure, the partial pressure of the gas is increased proportionately to the ambient pressure. Substances diffuse down gradients, so elevating partial pressures make the gradient steeper and the diffusion process more dynamic and rapid.
Example: During Hyperbaric oxygen therapy with a patient breathing 100% 02 at 3 ATA, the PA 02 (partial pressure of alveolar oxygen) rises to 2193 mm Hg. That is 21 times more than breathing air at sea level. Carbon dioxide (C02) and water vapour remain constant because they are made by the body.
Hyperbaric oxygen therapy dissolves oxygen into the tissues bypassing the need for the oxygen haemoglobin system.
‘The volume of a gas dissolved in body fluids or a liquid is directly proportional to the partial pressure of the dissolved gas.’
This law implies an equilibrium in which equal amounts of each gas are passing into and out of any liquid in contact with it. The time taken to reach equilibrium depends on the solubility of the gas in the tissues and the rate gas supplied to each tissue.
Gas moves out of solution when the total pressure or partial pressure of a gas is reduced. If this reduction occurs rapidly, then bubble formation may result because the fluid (or tissue) contains more gas than it can hold in solution. Note that it takes time for gases to move in and out of solution.
Example: Carbonated beverages. In essence, being mostly fluids, a human mimics the effect of soft drink. The body is the flavoured syrup, the container is the chamber, the gas is oxygen. Over a :90 minute treatment period, oxygen goes into solution in all fluids of the body.