011: EPICC Review Week 5: 5 Essential Gas Laws for Air Medical Transport

5 Essential Gas Laws to know for Air Medical Transport Professionals

Spend much time in Aeromedical Transport and you’re bound to hear people talking about Flight Physiology and the impact of the Gas Laws.

 In this episode of the podcast I talk about the 5 most essential gas laws you need to know about as a flight paramedic or flight nurse and the effects each one may have on you and your patient.

The 5 Essential Gas Laws

Law #1: Boyle’s Law

Law #2: Dalton’s Law

Law #3: Charles Law

Law #4: Henry’s Law

Law #5: Fick’s Law

Boyle's Law.jpg

Boyle’s Law

At a constant temperature, the volume of a gas is inversely proportional to its pressure.

P1 x V1 = P2 x V2


To remember Boyle’s Law, think “B” for Balloon.  What happens to a balloon as it rises higher in the sky?  It expands.

As the balloon ascends, the outside air pressure drops.

According to Boyle’s Law, as the pressure drops, the volume of air inside the balloon increases.  

Because the balloon is capable of stretching (much like to balloon of your ET tube), the expanding volume of gas inside the balloon causes the container to expand too.

Physiological effects

  • ETT / Foley / Blakemore / Other Cuffs – Expansion of Cuff Can Cause Tissue Necrosis

    • Watch Your Cuff Pressures of consider filling your cuffs with NS.

  • Air Left Inside an IV Bag – Expansion of Air Causes Infusion Rates to Increase as Altitude Increase (Unpressurized Cabin)

    • Remove ALL Air From IV Bag Before Flight

  • Air Inside Pleur-Evac – If Not Vented Can Cause Re-Collapse of Lung

    • Vent or Place to Suction

  • Air Trapped Inside Body Cavity – Air Expansions Can Cause Problems

    • Small Pneumothorax can Tension

    • Pneumocephalus (Air Trapped in the Skull) Can Cause Neurological Damage

  • Air Trapped Inside Newly Applied Cast Can Expand and Damage Cast

    • Cast Should be Cut Along Lateral and Medial Sides and Wrapped with Elastic Bandage to Allow for Expansion

  • Air Trapped inside new dental work can be extremely painful to patients.

  • Clogged sinuses will cause severe pain during changes in altitude


Dalton’s Law

The total pressure of a mixture is equal to the sum of the partial pressures of each gas in the mixture.

P = P1 + P2 + P3…

 Put another way, the partial pressure of a gas in a mixture is equal to it’s percentage within that gas X the pressure of that gas.

The most important example of Dalton’s Law is the partial pressure oxygen in a volume of air.

We know that at Sea Level the Standard Atmospheric Pressure is 760 torr, and that air is comprised of 21% oxygen, 78% nitrogen, and 1% other gases.

If we apply Dalton’s Law to normal air at standard sea level pressure than…

The partial pressure of oxygen in air (P1) = 0.21 X 760 torr.

 P1, or the partial pressure of oxygen, then equals 160 torr at sea level.

Physiological effects

  • Explosive Decompression at High Altitudes can Result in rapid decompression

    • Partial pressure of oxygen higher inside the blood compared to the partial pressure of oxygen in the atmospheric air.

      • Causes oxygen to actively dissociate from hemoglobin leading to rapid hypoxia and loss of consciousness.

  • Crew Members May Require Oxygen At Higher Altitudes in Unpressurized Cabins.

  • Patients May Require Supplemental Oxygen Just Because You’re Flying At Altitude.

  • Patients Who Already Receiving Supplemental Oxygen to Maintain Normal Oxygenation WILL Require Higher FiO2 During Flight.

I underlined the last point above because I consistently hear that this concept is covered on both Aeromedical Exams (FP-C and CFRN).


So how do you apply this to your practice?

To determine the exact amount of oxygen a patient needs to maintain a constant level of blood oxygenation at a higher altitudes requires a bit of math….., but there is a easy method to solve their oxygenation problem.


First the Easy Method… Ready?




(Now, for those of you who are like me and want me to prove it, here you go!)

Now I’ll explain the math and give the rationale behind my simple answer above. But before I get into this, there’s a few things we need to know about changes in air pressure and converting one atmospheric pressure to another.

  1. Air pressured can be expressed in several ways: inHg, mmHg, and torr The most useful unit of measurement for our purpose is mmHg (or torr).

  2. To convert from inHg (inches of Mercury) to mmHg you multiple inHg by the constant 25.4.

  3. As you ascend, the atmospheric pressure drops by about 23-27 mmHg / 1000ft at lower altitudes and closer to 20 mmHg / 1000ft at higher altitudes (above 8000MSL).

    • (These are approximations, but are close enough for what we’re doing.  I will generally just us 25mmHg/1000ft to ensure my calculations are at least what the patient needs).

  4. You need to know the formula for determining the required FiO2, so here it is…

(Starting FiO2 x Starting Air Pressure) / Air Pressure at Altitude (either real or approx) = FiO2 Requirements at Altitude

(F1 x P1 / P2 = F2)

So Now What?!?


Air Pressure is most often reported in either inHg (i.e. 29.92 – Standard Air Pressure at Sea Level) or mmHg (i.e. 760mm Hg Standard Air Pressure at Sea Level, This is the same as 760 torr)


As I said before, to convert from inHg to mmHg, which is more useful for our calculations, multiple inches of Hg by 25.4.


Example: 29.92 inHg x 25.4 = 760 mm Hg (Standard Air Pressure at Sea Level). See how their the same?


Take a look at this altimeter from a Agusta 106 Power.

Notice the pressure in the lower left is in inHg.


Using the formula above we get…


30.21 X 25.4 = 767 MMHG OR TORR


767 mmHg is the Atmospheric Pressure at ground level.  This is important to know because it gives us a starting point for determining the patients required FiO2 once we get to altitude during the flight.


Hint! If you’re flying in an unpressurized cabin, figure out the air pressure once you land at the sending facility or scene and write it down in mmHg.


Now, lets take, for example, a patient who’s vented on 35% FiO2 that we take out of an ER. During a typical flight our max altitude is around 2000ft AGL.


If I wanted to titrate my patients FiO2 to an FiO2 necessary for them to maintain a constant oxygenation level before leaving the referring facility, I could take the information I already know…


Air pressure at the sending facility = 767 torr

The patients current FiO2 demands = .35


And the anticipated air pressure once we reach cruise altitude of 9000ft = 767-(25x2) = 717 mmHg

…and calculate what FiO2 my patient will require once we reach altitude.

 P1 = 767

F1 = 0.35

P2 = 717

(767 x 0.35) / 717 = F2 = 37%

What Do You Notice?


Right!  From 35% to 37%. Not a big difference. That’s why I said earlier, just turn it up a bit.


Now I will caution you, that’s not always the case.  Let’s look at another EXTREME example.


Say you pick up a patient from sea level (We’ll use standard atmospheric pressure to keep it easy), and during your flight you’ll have to cross a high ridge where your pilot will have to climb to 10,000 ft MSL.


What can you expect your patient’s FiO2 requirements to be if their requiring 35% FiO2 when you pick them up?


760 x 0.35 / 510 = 52%


That’s a 17% increase.  Still not that much more.


For test purpose, you’ll likely get a questions that will ask you what FiO2 your patient will need if you’re going from one altitude to another.  If you remember that a few thousand foot altitude change requires a few additional % of FiO2, and 1 additional LMP approximately equals 4% you’ll be fine.


Unless you’re making HUGE changes in altitude, you’re patient shouldn’t need more than a few more liters per minute.


Charles Law

C = Celsius (or temp)

Law stating that the volume of an ideal gas at constant pressure is directly proportional to the absolute temperature.

As temperature goes up, volume will go up.

Practical Application

  • Air Bubbles in IV Lines, or air in IV bag, will expand when taken from cool environment to warm environment.

  • Effect on Density Altitude - As temp goes up, the air volume will expand resulting in worse aircraft performance.


Henry's law

The amount of dissolved gas is proportional to its partial pressure in the gas phase.

Pour a regular beer hard and what happens?

It bubbles over because there’s less CO2 in the atmosphere (0.04 %) keeping the CO2 in the beer.


Nitro Beer Example

Pour a Nitro beer hard and it doesn’t bubble over.


Because there’s more nitrogen in the atmosphere (78%( which keeps more nitrogen in the beer.

Physiological Effects

  • Rapid (Explosive) Decompression

  • Decompression Sickness

Screen Shot 2019-01-10 at 3.21.40 PM.png

Fick’s Law

The net diffusion rate of a gas across a fluid membrane is proportional to the difference in partial pressure, proportional to the area of the membrane, and inversely proportional to the thickness of the membrane.

Practical Implications

  • Increasing the FiO2, thus increasing the partial pressure of alveolar O2 will increase the rate of both oxygen and CO2 diffusion.

  • Increasing alveolar pressure (CPAP / PEEP) will increase the net partial pressure within the alveoli thus increasing the partial pressure of O2 and increasing the rate of diffusion.

  • Increasing alveolar pressure (CPAP / PEEP) will keep the alveoli open, increasing the surface area and thinning the wall, thus increasing the rate of diffusion.

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I hope you enjoy this episode! Until next time remember...

"Education is good, but excellence through collaboration is much better!"

Stay Safe, and Live Well!


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