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GAS LAWS

Category: EMS Operations

Topic: Flight Physics

Level: Critical Care

Next Unit: Stressors of Flight

24 minute read

Ideal Gas Laws

The ideal gas laws are a collection of mathematical formulas that describe how gases respond to volume, pressure, and temperature changes. These laws are usually seen in first-year college physics courses but become relevant to EMS professionals when a patient is taken on an airplane or helicopter. This section will attempt to explain the basics of these laws and how they affect the health of your patients.


Boyle's Law

Boyles law is the most clinically relevant gas law for air medical providers. It describes the relationship between the pressure and the volume of a gas. The changes in atmospheric pressure as an aircraft ascends and descends make Boyle's Law relevant.
INCREASE in Volume = DECREASE in Pressure: Picture a rubber birthday balloon when it is filled up and tied off. No air can escape. However, when you leave the store and accidentally let go of the string, the balloon climbs high into the sky. As altitude increases, the volume of that gas inside it expands because the atmospheric pressure outside the balloon decreases.
DECREASE in Volume = INCREASE in Pressure: If you take that same ballon but have it filled on Pikes Peak in Colorado (14,114 ft elevation), then travel by car to Colorado Springs (6,035 ft elevation), the volume of the gas within the balloon decreases and the pressure inside the balloon increases.

Implications/Complications as a Result of Boyle's Law

  • Tension Pneumothorax may occur or worsen as you increase in altitude. You may have taken off with a patient who had a "suspected" pneumothorax and, due to the decrease in pressure expanding the volume of air in the chest, now has a significant pneumothorax that needs care.
  • Pneumocephalus is the presence of air or gas within the cranial cavity; this is most common in patients with traumatic injuries to the head. The intracranial pressure (ICP) may increase due to the expansion of the gas within the skull and can cause significant life-threatening problems for your patient.
  • Sinus Pain barosinusitis is pain due to the expansion of trapped gas in the sinuses. Generally, small air passages prevent gas from being trapped in the sinuses. Illness or injury can block these passages leading to a painful sensation of pressure when this gas expands.
  • ETT bulbs will increase in size as you gain altitude. This can lead to trauma to the trachea, especially in pediatric patients. While rupture of these bulbs is rare, a sudden air leak or difficulty ventilating the patient could be secondary to this.
  • Gastric distention Distention of the stomach can occur during ascent in intubated patients and cannot burp. This distention can make ventilation more difficult. Placing a nasogastric or orogastric tube will prevent this.
  • Colostomy bags may need to be "burped" as you climb in altitude as the volume of the gas from stomach contents expands in the bag.

Charles's Law

Charles's Law describes the volume of a gas in relation to its temperature. As temperature rises, there is a direct increase in volume. When an oxygen cylinder is left in the sun, the volume of the gas inside increases. As the temperature around the cylinder decreases, the gas within will decrease in volume and, in turn, pressure.

Implications/Complications as a Result of Charle's Law

Besides the change in the pressure of gas cylinders listed above, Charles law takes on a special significance in rotor-wing transport. Helicopters often have a strict weight limit. This limit is directly affected by the temperature. A warmer gas has a higher volume. This means it also has a lower density as the gas molecules are further apart. A decrease in gas density decreases the lift generated by the helicopter blades and lowers the aircraft's weight limit.


Dalton's Law

Dalton's Law states that the total pressure of a gas mixture is based on the pressures of each component gas. This is the foundation of the critical concept of "partial pressure."

Our atmosphere is made up of nitrogen, oxygen, carbon dioxide, and small amounts of other gases. No matter the altitude, the ratio between these gases is identical. 21% oxygen, 78% nitrogen, and 1% carbon dioxide.

At sea level, the air pressure is 1 atmosphere (atm); this is sometimes referred to in torr (760mmHg) or kilopascals (101.325 kPa). We will use kilopascals for this illustration.

The partial pressure of oxygen at sea level is approximately 21kPa because oxygen makes up 21% of the atmosphere, and the pressure at sea level is approximately 100kPa.

If you ascend to 20,000 feet, the pressure of the atmosphere is approximately half, or 50kPa. However, the partial pressure is now only 10.5kPa despite the same percentage of oxygen! Oxygen is still 21% of the atmosphere.

This point is vital because it illustrates that the partial pressure of oxygen is essential, not the percentage of it in the gas we are breathing. 100% oxygen at 20,000 feet is not the same as 100% oxygen at sea level!

Implications/Complications as a result of Dalton's Law

Dalton's Law is used to calculate a patient's expected (PO2 vs. their actual PO2 at a given altitude on arterial and venous blood gas. Many modern portable machines do this automatically when the altitude is entered.

Dalton's Law also explains why a patient with an oxygen saturation of 95% on the ground may see a decrease at altitude in a fixed-wing or rotor-wing flight. Keep in mind that fixed-wing aircraft generally pressurize their cabins to an equivalent of 7,000 feet.


Fick's Law

Fick's Law is one of the most critical gas laws in clinical medicine. However, unlike Boyle's Law, the factors that go into it are less noticeable.

Fick's Law describes the diffusion of a gas across a membrane. For our purposes, this is most important when considering the diffusion of oxygen and carbon dioxide across the alveolar-capillary membrane.

The partial pressure of the gas in question, the membrane's surface area, and the membrane's thickness are the three key factors in Fick's Law. Increases in partial pressure and surface area speed diffusion, increases in membrane thickness dramatically slow diffusion.

Implications/Complications as a Result of Fick's Law

  • Partial pressure: As you climb in altitude and the partial pressure of oxygen decreases, less of it can diffuse across the membranes into the lungs. This may result in hypoxia.
  • Surface area: If pus from pneumonia or water from pulmonary edema blocks the alveoli, the surface area for gas exchange decreases. This may result in hypoxia regardless of the partial pressure of oxygen.
  • Surface area: The alveoli can be destroyed by conditions like COPD and some toxic exposures. These patients have a chronically decreased area for exchanging oxygen and carbon dioxide. This is the reason why carbon dioxide accumulates in patients with COPD.
  • Treatment: Fick's Law explains why conditions such as pneumonia, ARDS, and pulmonary edema can improve with oxygen. Giving supplemental oxygen raises the partial pressure of oxygen. This may help overcome decreased surface area or conditions that increase membrane thickness.

Henry's Law

Henry's Law explains that, for a liquid with a gas dissolved in it, if the pressure of that gas surrounding the liquid decreases, the amount of that gas that can stay dissolved in that fluid decreases.

The opening of a soda can is an excellent example of Henry's Law. Cans are pressurized with CO2, and that CO2 dissolves into the fluid. When the can opens, the pressure is released, and the dissolved CO2 comes out of the fluid in bubbles.

Implications/Complications as a result of Henry's Law

Decompression sickness, also known as The Bends, occurs when a diver surfaces too quickly from a significant depth. As the diver descends underwater, nitrogen within the lungs is dissolved into the bloodstream due to increased pressure from the air he is breathing via his SCUBA gear. As the diver begins to ascend to the surface, the pressure decreases in the lungs, and nitrogen begins to move back into the alveoli. However, if the diver ascends too quickly, the nitrogen begins to form bubbles within the veins and the surrounding tissues as the pressure around the diver rapidly decreases. This results in severe pain at best and ischemia of critical organs at worst.


Gay-Lussac's Law

Gay-Lussac's Law states that the pressure of a gas is related to the temperature of that gas.

This is most commonly seen with aircraft at high altitudes. The extremely low pressure of the atmosphere at significant heights leads to extremely low temperatures.

Implications/Complications as a result of Gay-Lussac's Law

The application of this Law is straightforward. As an aircraft increases in altitude, the temperature within will decrease. This difference is more pronounced in rotor-wing aircraft as fixed-wing aircraft are constantly pressurized to around 7,000 feet. Patients will be more prone to hypothermia in-flight situations and should be monitored closely.