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In order to understand how respiratory diseases and their respective oxygenation treatments work in the human body, an in-depth understanding of the fundamentals of hyperbaric oxygenation treatment (HBOT) is essential for professionals working in the healthcare field.
General information
In order to study the physical basis of the foundation of hyperbaric oxygenation treatment, it is necessary to know the following general characteristics of gases: the constants, pressure and density, the composition of air and the characteristics of oxygen, nitrogen and carbon dioxide.
The relevant information is based on Boyle’s and Mariotte’s gas law, the universal gas equation, Dalton’s law, Henry’s law and Fick’s law of diffusion. Also, it is necessary to know the principle of the adiabatic decompression law, Avogadro’s number and the units of pressure.
Basic concepts
- Molecular weight: 1 mole of a substance (atoms, ions, molecules or formula units) is 2 Avogadro numbers.
- Avogadro’s number: (6.022 x 1023) is approximately the number of particles (atoms, ions, molecules or formula units) contained in 1 mole of an oxygen, which has a molecular weight of 32 g, defined as the molecular weight of a substance.
Characteristics of gases
Avogadro’s Law
Avogadro’s law postulates that equal volumes of gases at the same temperature and pressure contain equal numbers of molecules. At standard conditions (0 °C, 1.013 bar) the volume of any gas is 22.42 l/mol.
Avogadro (1811) and Ampère (1814) independently proposed the existence of gaseous molecules formed by two or more equal atoms. However, both authors reached the same conclusion: in a chemical reaction, a molecule of a reactant must react with one or more molecules of another reactant. This gives rise to one or more molecules of the product. However, a molecule cannot react with a non-integer number of molecules since the minimum unit of a reactant is the molecule.
Therefore, there must be a simple integer relationship between the molecules of the reactants, and between these molecules with those of the product.
Boyle and Mariotte’s Law
First described independently by Mr. Robert Boyle (1627-1691) and Edme Mariotte (1620-1684). It is also called the ‘Boyle-Mariotte law’. Its equation is pxV= constant at constant temperature.
The product of pressure (p) and volume (V) in a quantity of gas at equal temperature (T) remains constant. For a confined quantity of a gas in two different states at constant temperature the equation changes: p1V2 = psV2.
- Hyperbaric relevance: within hyperbaric chambers, any gas volume confined in the human body and in equipment (medical) is subject to this law. In organs or gas-filled spaces with rigid walls, this effect must be taken into account during compression and decompression at higher pressures. This is most important between 1 bar and 1.5 bar (100 kPa – 150 kPa) where pressure changes cause the greatest relative volume changes.
Dalton’s Law
First described by John Dalton (1766-1844) in 1801. This gas law is also called Dalton’s partial pressure law, which states that the total pressure exerted by a gaseous mixture is equal to the sum of the pressures that would be applied by the gases if each were present and occupied the total volume.
Each gas in the mixture acts as if it were alone, so the pressures of each gas forming a mixture add up. The partial pressure of a gas (p1) equals the product of the total pressure of the gas mixture (Pt) and the fraction of the gas (F1). Equation: P1=Pt x F1.
- Practical relevance: gases that are not toxic when inhaled at ambient pressure, in a certain percentage of a gaseous mixture (vol. %) may become toxic when inhaled at elevated total pressure. It is the partial pressure, not the percentage in a gaseous mixture, that causes toxicity. In addition, if the partial pressure of oxygen is increased in the air mixture, it will exert higher partial pressure and consequently higher total pressure at the level of the pulmonary alveolus.
Henry’s Law
This law, first formulated by William Henry (1775-1836) in 1803, states that the mass of a gas (C) dissolving in a defined volume of liquid is directly proportional to the pressure of the gas (P), provided that the gas does not react with the solvent. Equation: αP/C = constant at constant temperature.
- P=partial pressure of oxygen.
- C=gas concentration in liquid.
- A=Bunsen solubility coefficient (specific for gases and liquids).
As a basic principle, the solubility of gases is greater in cold liquids.
- Practical relevance: The pressure-dependent solubility of inert gases (e.g. nitrogen) in body fluids and tissues is crucial for the development of decompression sickness (DCS). This is due to the supersaturation of tissues in relation to the reduced environmental pressure after exposure. The solubility of nitrogen in water is much lower than the solubility of oxygen. It constitutes one of the causes of decompression sickness of the diver. On the other hand, as the pressure inside the hyperbaric chamber increases, the pressure in the hematoalveolar space produced increases the amount of oxygen diluted in the plasma and tissue fluids. Thus, the higher the pressure, the greater the amount of gas diluted in the liquid, at constant temperature.
Fick’s Law
Fick’s diffusion laws were derived by Adolf Fick in 1858. The first law is used in steady state diffusion and results in the formula below (which indicates the rate of diffusion of a gas across a membrane): Diffusion rate= k x A x ΔP.
- K=constant (determined by experiment, gas and temperature).
- A=surface area where diffusion takes place.
- ΔP=difference between the partial pressures between the two sides of the membrane.
- D=distance over which diffusion takes place.
- Practical relevance: At various locations in the human body, the partial pressures of dissolved gases (such as oxygen or nitrogen) depend on diffusion. Variables for gas diffusion such as the size of the diffusion area, the thickness of the diffusion barrier (or distance) and the difference between the gas partial pressures are considered. If the difference is greater, the diffusion rate is greater between both sides of the membrane (for the same membrane size and the same distance).
Atmospheric air
Atmospheric air is a mixture of different gases (21% oxygen, 79% nitrogen, including 1% argon). The CO2 fraction is practically imperceptible, being slightly more important in expired air (approximately 4 %).
Water vapor is a highly variable component of air. At higher temperatures, the air may contain larger amounts of water vapor. The unit ‘% relative to humidity’ is temperature dependent. Also, like all other gases in the air mixture, water vapor produces a gas pressure (pH2O) at 37 °C and 100 % relative to humidity (= 100 % saturation with water vapor) pH2O equals 47 mmHg.
Finally, the different physical laws are applied in the foundation of hyperbaric oxygen therapy both in the desired effect and in the considerations of adverse events. Their understanding makes it possible to distinguish the different effects that can be obtained from the combination of increased atmospheric pressure and increased partial pressure of inspired oxygen.
Respiratory diseases
Currently there is a resurgence in the use of hyperbaric oxygenation treatment (HBOT) as an auxiliary tool in different medical specialties. TECH Technological University offers different high quality postgraduate courses such as the Master’s Degree in Respiratory Physiotherapy in Rehabilitation Medicine and the Master’s Degree in Improvement of Intensive Care Medicine, in which the topics related to the different respiratory diseases and their respective studies are deepened.
The Master’s Degree in Hyperbaric Medicine teaches how the creation of new generation hyperbaric chambers, more accessible to use, cost, and installation in public and private health institutions, has made different professionals incorporate this tool in their daily practice.
