SCUBA Diving and Gas Laws

Polly Dornette
Product Developer and Former SCUBA Instructor

© J. Robert Patrick

SCUBA diving nuts

The dry air we breathe every day is equanimous of 21% oxygen, 78% nitrogen, and <1% other gases. Its average pressure at sea level is i atm (14.7 psi). For SCUBA, this air is compressed into a SCUBA cylinder or "tank." SCUBA tanks tin exist fabricated of steel or aluminum; each of these materials has pros and cons that affect the diver'southward decision on which blazon to use.

The compressed air in the tank is delivered to the diver through a regulator, which reduces the pressure from the tank to match the ambient pressure. At the surface, ambient force per unit area is one atm and it increases by 1 atm for every 10 m in depth through which a diver descends. Note: Other gas mixes such as nitrox (an oxygen/nitrogen mixture with a greater amount of oxygen than air), heliox (a helium and oxygen mixture), and trimix (a mixture of oxygen, nitrogen, and helium) or even pure oxygen are likewise used for technical diving, but those mixes go beyond the scope of this word.

Boyle's law: P _one V _i = P _ii V _two

A key rule of SCUBA diving is to "never hold your breath." A look at Boyle'south law explains why this dominion exists. When a diver inhales air from a SCUBA tank, the air that enters the diver's lungs is at ambient pressure. If a diver inhales from the tank on the surface, the pressure in her lungs will be at 1 atm. If she inhales air from her tank at a depth of 30 g (~99 ft), the pressure in her lungs will be iv atm (30 m / x k/atm = three atm from the water plus 1 atm from the air at the surface = iv atm). Assuming the diver'due south lung volume is i L, we can consummate the left side of the equation for Boyle's police force. If a diver at 30 m has 1 L (V ₁) of air at a pressure of 4 atm (P _ane) in her lungs and ascends to the surface (P _ii) while holding her breath, the post-obit equation applies:

4 atm × 1 Fifty = 1 atm × V ₂

© John Simone

Afterward working through this instance, students often ask why free divers are able to swoop to such extreme depths. Free divers make full their lungs at the surface with air at ambient force per unit area (P ₁) then descend while property their breath. The pressure change has the opposite touch on on the volume of their lungs. A free diver diving to a depth of 30 m would have his lungs shrink to ¼ of their initial volume, which can be determined using the following equation:

one atm × 1 L = 4 atm × V _two

SCUBA instructors sometimes demonstrate this principle to their students by bringing a foam cup along on a dive. As the force per unit area increases with depth, the gas bubbles trapped in the foam decrease in book, shrinking the cup.

Boyle's law also has implications on the amount of air used from the tank with each breath. At 10 grand (2 atm) twice as many oxygen and nitrogen molecules are inhaled with each breath. Deeper dives require closer monitoring of a diver's air supply considering the diver uses his supply more speedily. Some other question students often inquire in this give-and-take is, "How is the SCUBA tank impacted by these changes in pressure?" Because the tank is a rigid container, its volume is not altered with the change in external pressure nor is the gas it contains affected.

Gay-Lussac'south law: P ₁ / T _ane = P ₂ / T ₂

In SCUBA diving, Gay-Lussac's law (sometimes referred to every bit Amontons' law of pressure-temperature) is well-nigh of import in relation to the amount of breathable air in a tank. The pressure level of an "empty" tank is depression (around 500 psi), and the temperature is equal to the ambient temperature. SCUBA tanks made out of aluminum typically have a rated fill pressure of 3,000 psi.

A SCUBA tank is a rigid container, therefore its book is held constant. When a tank is filled, additional oxygen and nitrogen molecules are added to the tank and the pressure and temperature increment. If a tank is filled rapidly to 3,000 psi (P _ane), its temperature can rise to as much as 150° F (65.half dozen° C). Since all gas laws use accented temperatures, this temperature needs to be converted.

Most students know they can convert a Celsius temperature to an absolute temperature of Kelvin by adding 273. All the same, they are not probable to be aware that they tin can add 460 to a Fahrenheit temperature to convert it to a Rankine temperature, which is based on the Fahrenheit calibration but with zero representing absolute nada. As the tank cools to ambient temperature (T _ii) after the rapid make full, the gas pressure in the tank will also decrease. Assuming the ambient temperature is 70° F (21° C), the post-obit equations can be used to make up one's mind the pressure at the lower temperature:

Charles's constabulary: V _one / T ₁ = Five _two / T ₂

Charles's police force is seldom relevant to diver safety; however, the implications of this law are responsible for an interesting miracle for divers using dry suits. A dry suit is a watertight garment worn by divers (typically over warm habiliment) that serves to go on the diver warm past trapping a layer of air between the diver and the suit. Dry suits are ordinarily worn in common cold air and/or water temperatures.

During the dive, divers can add and remove air from their dry suits through their regulators. This allows them to adjust for changes in their suits' gas volumes due to force per unit area changes during assent and descent. If the air temperature is colder than the h2o temperature when the divers emerge at the end of the dive, they tin can become "vacuum sealed" in their suits due to the decrease in their suits' gas volumes. Defined can add air to the suits from their tanks, or unzip their suits, to release the "squeeze."

Dalton's law: P _Full = P ₁ + P _ii + P ₃ . . .

Besides known as Dalton's constabulary of fractional pressures, this law states that the total pressure of a gas mixture is equal to the sum of the fractional pressures of its component gases. Every bit mentioned earlier, dry out air is a mixture composed of 21% oxygen and 78% nitrogen. Both of these gases can accept negative impacts on a diver at high pressures. Low partial pressures of oxygen are also dangerous but are just an issue for technical diving, which is beyond the scope of this word.

Oxygen can become toxic to a diver when the partial pressure of the oxygen breathed is above one.six atm. Symptoms of oxygen toxicity can include changes in vision, dizziness/vertigo, and seizures, all of which can be problematic for a diver and can lead to death. To calculate at what depth a diver might begin to feel symptoms of oxygen toxicity when diving with compressed air, nosotros need to first calculate at what air pressure level would the partial pressure of oxygen exist equal to 1.6 atm or greater.

At 1 atm of full pressure level for air, oxygen would have a partial pressure level of 0.21 atm. Therefore, the full force per unit area of the air would be vii.6 atm (1.6/0.21 atm) for the fractional pressure level of oxygen to be at i.6 atm or greater. Remember that for each 10 k of depth the pressure increases past 1 atm, only the pressure at the surface is 1 atm, and then the partial pressure of oxygen in air would be one.six atm at 66 m (216 ft).

Nitrogen narcosis can event from a diver's exposure to high partial pressures of nitrogen during her dive. Symptoms of nitrogen narcosis about closely resemble those of alcohol intoxication. These symptoms appear more gradually than those of oxygen toxicity simply also increase with depth.

Henry'southward police

Henry's law states that the concentration of a gas dissolved in a liquid at a given temperature is direct proportional to the fractional force per unit area of the gas above the liquid. The implication of this police force for SCUBA diving is that as depth increases (and therefore pressure level) the corporeality of a gas dissolved in the diver's blood volition also increase. Oxygen is consumed by the body'south physiological processes, only nitrogen is physiologically inert. The longer that a diver remains at depth, the more nitrogen is dissolved in his blood.

During long dives a considerable amount of nitrogen tin can be dissolved in the diver's bloodstream. When the diver ascends the partial pressure of nitrogen drops, and due to Henry'southward law the dissolved nitrogen begins to come out of solution. Nitrogen bubbles class in the diver's bloodstream, which can lead to decompression sickness (DCS).

The symptoms of DCS and their severity depend on where in the diver's body the bubbles migrate and tin range from soreness in the joints or blisters under the skin to death. Treatment for DCS typically involves several sessions in a hyperbaric oxygen chamber. In their preparation, divers are taught to stay inside dive time and depth limits to minimize their run a risk of DCS and to ascend slowly from every dive.

Related products

For further exploration of the gas laws, nosotros recommend the following products:

• Spouting Cylinder (detail #752529)
• Inquiries in Science®: Expanding on the Gas Laws Kit (item #251205)
• Carolina STEM Challenge®: Cartesian Divers Kit (item #750024)