On the surface of the Earth, we are exposed to the pressure exerted by the atmosphere. This is called the atmospheric or barometric pressure. Most people regard this pressure as caused by the mass of the atmosphere pressing down on them. A flaw in this argument is that the pressure remains in a bottle after it is sealed, although its contents are contained and are no longer exposed to the column of air above. The physically correct explanation is that atmospheric pressure is generated by collisions of the molecules of gas in accordance with the kinetic theory of gases. Either explanation is acceptable for the following discussion.
The pressure decreases as we move upward through the atmosphere and increases as we move down into a mine or into the sea. At the top of Mount Everest the atmospheric pressure is about 40 per cent of that at sea level. Because water is much heavier than air, the pressure changes experienced by divers over a particular depth change are much greater than those encountered by climbers or aviators as they change altitude.
Pressure is measured in a variety of units from either of two reference points. It can be expressed with respect to a vacuum, i.e. zero pressure. This reading is called an absolute pressure. The second method measures pressures above or below local pressure. These readings are called gauge pressures. At sea level, the absolute pressure is 1 atmosphere (1 ATA) and the gauge pressure is 0. These units are commonly abbreviated to ATA and ATG.
Common examples are the barometric pressure used by weather forecasters, which is an absolute pressure, and the blood pressure, which is a gauge pressure reading.
With descent in water, pressure increases. For each 10 metres of depth in sea water, the pressure increases by 1 atmosphere, starting from 1 ATA or 0 ATG at the surface. The gauge pressure remains 1 atmosphere less than the absolute pressure. For example, at 10 metres, the pressure is 2 ATA and 1 ATG. At 90 metres, the pressure is 10 ATA and 9 ATG.
Table 2.1 Pressure conversion factors (commonly used approximations shown in brackets)
Because diving involves facets of engineering and science, it is plagued with many units of pressure. These include absolute and gauge atmospheres, pascals and multiples such as the kilopascal, metres or feet of sea water, bars, pounds per square inch, torr and several other rarer units. Table 2.1 lists conversions for the more commonly used units.
Pressure and the diver’s body
Many people have difficulty in understanding why the pressure of the water does not crush the diver. The answer to this problem may be considered in two parts:
The solid and liquid parts of the body are virtually incompressible, so a pressure applied to them does not cause any change in volume and is transmitted through them. After immersion, the increased pressure pushes on the skin, which in turn pushes on the tissues underneath, and so the pressure is transferred through the body until the skin on the other side is pushed back against the water pressure. Therefore, the system remains in balance. This is in accordance with Pascal’s Principle, which states: ‘A pressure exerted anywhere in a confined incompressible fluid is transmitted equally in all directions throughout the fluid such that the pressure ratio remains the same’.
However, the effect of pressure on the gas spaces in the diver’s body is more complex. The applied pressure does not cause any problems if the pressure in the gas space is close to that of the surrounding water. There is, for example, no physical damage to a diver’s lungs if the air space was exposed to an internal pressure of 100 metres of water, provided that this pressure is balanced by the pressure exerted by surrounding water acting on the walls of the lung to balance any tendency of the lungs to expand. If the lungs were exposed to an internal pressure sufficiently more than the surrounding atmospheric tissue, they would overexpand and burst.
Water pressure and lung inflation
Immersion up to the neck in water reduces vital capacity by about 10 per cent (Figure 2.1 shows lung volumes). This is caused in part by the hydrostatic pressure of the water compressing the thorax. With immersion, there is also a loss of gravitational effects. This reduces the volume of blood in lower, mainly leg, veins and increases thoracic blood volume. This change in turn reduces the compliance of the lungs.
When a diver is using breathing equipment, pressure at the point from which the gas is inhaled can be different from the pressure at the chest. If upright in the water, a scuba diver is inhaling air released at the pressure at the level of the mouth. A snorkel diver is inhaling air from the surface, and this is at surface pressure. In both these cases, the air is at a lower pressure than the diver’s lungs. This reduces the amount of air the diver can inhale because part of the inhalation force is used in overcoming this pressure difference.
Conversely, when descending, face-down, a diver whose air is released at mouth pressure can inhale to greater than normal vital capacity but could not exhale to the normal residual volume. This is because in this orientation, the water pressure is helping to inflate the lungs.
Pressure and volume changes
When a diver descends, the increased pressure of the surrounding water compresses gas in the gas spaces within the diver’s body. These spaces include the lungs, middle ears, sinuses and intestines.
This is one of the many aspects of diving medicine that is concerned with the relationship between pressure change and change of gas volume. The relationship between changes in volume of a gas and the pressure applied to it is described by Boyle’s Law. This states: ‘if the temperature remains constant, the volume of a given mass of gas is inversely proportional to the absolute pressure’. This means that the absolute pressure multiplied by volume has a constant value, and this constant changes with the mass of gas considered. To a mathematician, this means that P × V = K or P1 × V1 = P2 × V2, where P and V are pressure and volume. For example, 10 litres of gas at sea level pressure (1 ATA) will be compressed to:
- 5 litres at 2 ATA (10 metres).
- 2 litres at 5 ATA (40 metres).
- 1 litre at 10 ATA (90 metres).
During ascent into the atmosphere, the reverse happens and the gas expands. This means that the 10 litres of air would expand to 20 litres at 0.5 ATA (an altitude of about 5000 metres or 18 000 feet) and to 40 litres at 0.25 ATA (an altitude of about 10 300 metres or 33 400 feet).
Gas volumes expand when pressure decreases and contract when pressure increases.
The volume of a mass of gas in a flexible container decreases with pressure or depth increase and expands during ascent or pressure reduction (Figure 2.2). It should be noted that volume changes are greatest near the surface. Conversely, gas has to be added if the volume of a container or gas space is to remain constant as the pressure is increased. The effects of this law are important in many aspects of diving medicine.
During descent, the increasing pressure in the water is transmitted through the body fluids to the tissue surrounding the gas spaces and to the gas spaces themselves. The pressure in any gas space in the body should increase to equal the surrounding pressure. In the lungs, during descent on breath-hold dives, this is accompanied by a decrease in lung volume. Air should enter cavities with rigid walls, such as the sinuses or the middle ear. If air entry does not take place to equalize pressures, then a pressure difference between the space and the surrounding tissue will develop, with the pressure in the gas space being less than in the surrounding tissue. The results are tissue distortion and damage, such as congestion, oedema or haemorrhage.
During ascent, as the pressure decreases, gas within body spaces will expand. Unless gas is vented from the space, the expanding gas will exert pressure on the surrounding tissue and will eventually damage it. Pressure changes in the middle ear can also result in rupture of the tympanic membrane.
The same volume changes with pressure occur in bubbles in tissue or blood. Again, the volume changes are greatest close to the surface. An injury caused by pressure change is called barotrauma.
Barotrauma is the general name for an injury caused by pressure change.
Respiration in water and under pressure
While breathing air underwater, the diver’s respiratory volume is about the same as it would be if he or she worked at the same rate on the surface. A consequence of this is that a cylinder that contains enough air for 100 minutes at 1 ATA would last about 50 minutes at 2 ATA (10 metres) or 20 minutes at 5 ATA (40 metres) for dives with the same energy expenditure. This is because the gas in the cylinder expands to a smaller volume when it is released against the ambient pressure at depth than it would if used at the surface. A cylinder that contains 5000 litres of gas if it is released at the sea surface would yield only 1000 litres of gas if it is released at 5 ATA, or 40 metres. A diving physician needs to keep this in mind when estimating the amount of gas needed for any task or therapy.
With depth, gas is compressed and there is an increase in density of the gas because there are more molecules in a given space. So, at depth, a diver must move a greater mass of gas with each breath. This requires greater effort and involves an increase in the work of breathing. In some situations, this can limit the capacity to do work.
The density of the breathing gas can be reduced by replacing nitrogen with a lighter gas such as helium. For example, the density of air at 1 ATA is about 1.3 kg/cubic metre. At 10 ATA, the density of air would be about 13 kg/cubic metre. The use of lighter gas helps to reduce density. For example, at 40 ATA, the density of a 1 per cent oxygen and helium mixture is 6.7 kg/cubic metre.
As the density of a gas increases, there is an increased tendency for the flow to become turbulent. This causes a further increase in the energy used in breathing. These factors can lead to fatigue of the inspiratory muscles and reduce maximum breathing capacity and the work output. To minimize this load, the body responds by using less gas for a given workload. This can result in the development of hypercapnia. Continued exposure to dense gas, as is encountered in deep dives, may cause an adaptive response.
Temperature and volume changes
Charles’ Law states: ‘If the pressure is constant, the volume of a mass of gas is proportional to the absolute temperature’.
The absolute temperature (A°) is always 273° more than the centigrade temperature. A more useful expression of the law is as follows:
Where V1 is the volume of a mass of gas at temperatures T1°A and V2 is its volume after the temperature has changed to T2°A.
This law has much less relevance to diving medicine than Boyle’s Law. However, it should be remembered when considering gas volumes and how they may change.
Boyle’s and Charles’ Laws may be combined and used if temperature and pressure both change – from P1 and T1 to P2 and T2 with a volume change from V1 to V2. The combined laws can be expressed as the universal gas equation:
A temperature-pressure problem that often causes discord can be used to illustrate the use of this equation. This is the effect of temperature on the pressure in a gas cylinder.
A diver may ask to have the compressed air cylinder filled to 200 ATA. The gas compressor heats the gas so the cylinder may be charged with gas at 47°C. When the diver gets in the water at 7°C, the diver may find that he or she has only 175 ATA in the cylinder. In this case V1 = V2 because the cylinder is rigid and the pressure falls as the gas cools.
So the reduced pressure is a result of temperature change, not a leaking valve or fraud by the air supplier.
Partial pressures in gas mixtures
Dalton’s Law states: ‘the total pressure exerted by a mixture of gases is the sum of the pressures that would be exerted by each of the gases if it alone occupied the total volume’. The pressure of each constituent in a mixture is called the partial pressure (Figure 2.3). In air, which is approximately 80 per cent nitrogen and 20 per cent oxygen, the total pressure at sea level (1 ATA) is the sum of the partial pressures of nitrogen, 0.8 ATA, and oxygen, 0.2 ATA. At 2 ATA (10 metres) these partial pressures will rise to 1.6 and 0.4 ATA, respectively.
The partial pressures of breathing gases can be manipulated to the diver’s advantage. For example, the composition of the gas breathed may be modified to reduce the chance of decompression sickness (DCS) by decreasing the percentage of inert gas in the mixture.
Undesirable effects can also occur. Air from an industrial area may contain more than 0.3 per cent carbon dioxide and 0.002 per cent carbon monoxide. If incorporated in compressed breathing gas and delivered at high partial pressures, both constituents could be toxic unless measures were taken to remove these contaminants before use.
It may be necessary to combine Boyle’s and Dalton’s Laws in calculations. For example, it may be decided that a diver should be given a mixture with a partial pressure of 0.8 ATA oxygen and 1.2 ATA nitrogen in a recompression chamber pressurized to 2 ATA. If oxygen and air are the only gases available, the gas laws can be used to calculate how to prepare a cylinder charged with the right gas mixture.
The mixture will need to be 40 per cent oxygen and 60 per cent nitrogen (Dalton’s Law). If the gas is to be prepared in a cylinder charged to 200 ATA, it should contain 120 ATA of nitrogen (60 per cent of 200). If this is to be obtained from compressed air (assumed to be 80 per cent nitrogen in this exercise), it will be necessary to put 150 ATA of compressed air into the cylinder (30 ATA of oxygen and 120 ATA of nitrogen) with 50 ATA of oxygen.
This simple mixing process cannot be used as successfully with helium mixtures. At high pressures, helium does not follow the predictions of Boyle’s Law accurately. It is less compressible than the ideal gas described by Boyle’s Law. Mixing can be conducted with allowance for this or by putting a calculated weight of each gas in the cylinder.
Solution of gases in liquids
Henry’s Law states: ‘at a constant temperature, the amount of a gas that will dissolve in a liquid is proportional to the partial pressure of the gas over the liquid’. This law implies that an equilibrium is established with each gas passing into and out of any solution in contact with it (Figure 2.4). At sea level (1 ATA), an individual’s body tissues contain about 1 litre of gaseous nitrogen in solution. If the diver dived to 10 metres and breathed air at 2 ATA, more gas would dissolve and he or she would eventually reach equilibrium again and have twice as much nitrogen in solution in the body. The time taken for any inert gas to reach a new equilibrium depends on the solubility of the gas in the tissues and the rate of gas supplied to each tissue.
When the total pressure, or the partial pressure of a particular gas, is reduced, gas must pass out of solution. If a rapid total pressure drop occurs, a tissue may contain more gas than it can hold in solution. In this situation, bubbles may form and may cause DCS.
The physiological effects of the solubility of gases are also relevant in nitrogen narcosis and oxygen toxicity.
It should be noted that each gas has a different solubility and the amount of any gas that will dissolve in a liquid depends on the liquid. For example, carbon dioxide is very soluble in water compared with other common gases. Beer aerated with compressed air instead of carbon dioxide would have far fewer bubbles. Nitrogen is more soluble in fats and oils than in aqueous solutions.
Henry’s Law is also time dependent. It takes time for gases to enter and leave solution or form bubbles. If this was not so, champagne would go flat as soon as the cork was popped.
At depth, a diver breathing air absorbs nitrogen in accord with Henry’s Law. The amount depends on depth and time of exposure. When the diver surfaces, the excess nitrogen must pass from the body. If it is eliminated from solution through the lungs, there will not be any complications. In some cases, the nitrogen comes out of solution in the blood or tissues, thus forming bubbles that may lead to DCS.
Gas movement in body tissues
Gas transfer from the lungs to the tissues is dependent on the cardiovascular circulation, and the gas supplied to a portion of tissue depends on the blood perfusing it. In a permeable substance such as body tissues, gas molecules can migrate by diffusion. That is, gas molecules dissolve in the tissue fluids and tend to move from areas of high to low partial pressure until the partial pressure of the dissolved gas is uniform. This can take hours. It is the dissolved gas pressures that tend to equilibrate, not the number of gas molecules. If a gas is twice as soluble in one tissue compared with another, then twice as many molecules will be in the first tissue to produce the same partial pressure in the tissue. This information can be estimated from the solubility coefficients of the gas in the components of the tissue.
The rate of gas movement between two points depends on several factors. The difference in partial pressure and the distance between the two points may be combined into a concentration gradient. The other major factor is the permeability of the tissue, an expression of the ease of gas movement. A large partial pressure between two points that are close together (a steep gradient) and a greater permeability both increase the rate of gas transfer.
Metabolic gas exchange
In divers, gas exchange mechanisms are basically the same as at normal pressure. Oxygen diffuses down a concentration gradient from the lungs to the tissues. The carbon dioxide gradient is normally in the opposite direction. The exchange of inert gases becomes important and there are changes in the finer details of metabolic gas exchange.
With increasing depth, there is an increase in the partial pressures of the constituents of the breathing mixture in accordance with Dalton’s Law. This causes higher alveolar pressures and arterial pressures of the inhaled gases.
Elevated pressures of oxygen facilitate oxygen transport, but they may interfere with the elimination of carbon dioxide in two ways: first, by the depression of respiration induced by high arterial oxygen tensions; and second, by direct interference with the transport of carbon dioxide. When the inspired oxygen partial pressure is elevated, there is an increase in oxygen transport in solution in the plasma (Henry’s Law).
When one is inhaling oxygen at a partial pressure above 3 ATA, the total oxygen requirement may be carried in solution. If this happens, the haemoglobin may be still saturated with oxygen in the venous blood, and this can prevent the transport of carbon dioxide in the form of carbaminohaemoglobin.
The result is an increased tissue carbon dioxide level. In some situations, there may also be an increase in the inspired carbon dioxide pressure. Causes include contamination of the breathing gas supply, the external dead space of the equipment, inadequate ventilation or failure of the absorbent system.
There is a tendency for experienced divers to be less sensitive to elevated carbon dioxide partial pressures. This reduces the total ventilation requirement during working dives. Elevated arterial carbon dioxide levels increase susceptibility to oxygen toxicity, DCS and inert gas narcosis. For these reasons, it is desirable to control the factors that cause carbon dioxide retention.
Diving is associated with a tendency to retain carbon dioxide.
Inert gas exchange
The topic if inert gas exchange is considered in the chapters on DCS. Therefore, to avoid duplication, the topic is not considered in detail here. As indicated earlier, increased total pressure is usually accompanied by an increase in nitrogen (and/or other inert gas) pressure (Dalton’s Law). This causes gas transfer to the body tissues. When pressure is reduced at the end of the dive, the transfer is reversed. If there is an excess of gas, then it can come out of solution as bubbles. These bubbles are the cause of DCS. If bubbles do occur, they are also subject to the same physical laws. Their size decreases if the pressure is increased, and gas enters or leaves them depending on the concentration gradients of gases.