# Feel the Pressure! What happens when you shake a can of soda and then open it? What is happening to the gases in the can of soda?

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Science of Diving: How Science Helps Navy Divers Stay Safe Science Topic: Physiology and Physics

Feel the Pressure! What happens when you shake a can of soda and then open it? What is happening to the gases in the can of soda?

Breathe Deep Are gases in soda the same that you breathe? What gases do scuba divers breathe when they go underwater? Like gases in soda, gases breathed by divers are under pressure.

Navy Divers in Action

I’d Like to be Under the Sea
How would you breathe underwater?

SCUBA Self-Contained Underwater Breathing Apparatus

Is Diving Safe? SCUBA is safe, like motorcycling or sky-diving If you understand the physics and physiology diving is safe

Gas Laws and Diving Essential Question: “What conditions lead to perils facing divers including the bends and embolism?”

How much does air weigh at sea level?
Review! How much does air weigh at sea level? What is the weight of water at 33 feet (10 meters)? What is the pressure on a divers body at 33 feet underwater? Pressure increases linearly with depth (not exponentially). Pressure = psi (4 x 14.68) at 99 feet: 3 x for each 33 feet of water + 1 x for the 1 atmosphere of pressure from air pressure of the atmosphere.

Around the Bends What are the bends? How is decompression sickness treated?

Bends and embolism The bends result when dissolved gases (mostly nitrogen) absorbed as air is breathed under pressure are released as gas back into the bloodstream. Embolism results when gases in the lung expand during ascent and rupture lung tissue, causing air bubbles to enter the bloodstream. These conditions can happen if a diver ascends too quickly.

Boyle: volume and pressure Charles: temperature and volume
Law Makers Boyle: volume and pressure Charles: temperature and volume Outline basic concepts. Do not introduce equations at this stage.

Cartesian Diver Activity
Fill medicine dropper for neutral buoyancy Fill soda bottle to just below top Place dropper in soda bottle Tighten bottle cap Squeeze bottle Note observations Students should observe dropper descending when bottle is squeezed and ascending when squeezing is released. Boyle’s Law – When a gas is at constant temperature, there is an inverse relationship between the volume and pressure of the gas. As the volume decreases, pressure increases and as the volume increases, pressure decreases.

Add small amount of water to soda can Place can on hot plate
Pressure Can Activity Add small amount of water to soda can Place can on hot plate When water in can boils remove can Quickly invert can and place in bowl of water Note your observations Students will see the can rapidly collapse. As air in the can heats up, its pressure falls. When air cannot escape due to the can being inverted in water, the greater air pressure on the outside causes the can to collapse. Charles’ Law – When a gas is at constant pressure, there is a positive relationship between the volume of a given mass of the gas increases and its temperature. As the temperature increases its volume increases and vice versa.

The Gas Laws Essential Questions:
What are the principles of Boyle’s Law and Charles’ Law? How is the ideal gas law used to calculate changes in volume, pressure and temperature when one or the other variables is held constant? Students will see the can rapidly collapse. As air in the can heats up, its pressure falls. When air cannot escape due to the can being inverted in water, the greater air pressure on the outside causes the can to collapse.

The Equations Boyle’s Law: PV = k Charles’ Law: V1/T1 = V2/T2
Ideal Gas Law: PV = nRT n = number of moles of gas R = gas constant ( L·atm/mol·K) T = temperature in Kelvin

Boyle’s Law: What does it look like?
Pressure Pressure Click so that only the last line remains. This represents Boyle’s Law. Volume Volume

Charles’ Law: What does it look like?
Volume Click so that only the last line remains. This represents Charles’ Law. Temperature Temperature

Should divers care about the gas laws?
Essential question: How do the gas laws relate to safe diving practices?

Why should divers care about Boyle’s Law?
Divers who know Boyle’s Law know if pressure decreases, volume increases. A diver breathes gases under pressure. If a diver ascends too quickly pressure goes down so gases dissolved in the bloodstream increase in volume. Remember what happens when you shake a soda can? Yup. Gases start to fizz in the diver’s bloodstream. Not good.

Why should divers care about Charles’ Law?
Air in a scuba tank is pressurized to about 3000 psi. What happens if the tank is heated, for example in a car on a summer’s day? Yup. The tank could explode. Not good. If a diver is close to a bends or embolism situation and is hypothermic, what could happen if the diver is quickly warmed up?

What can divers do to avoid the bends and embolism?
Safe Diving! What can divers do to avoid the bends and embolism? What is a safety stop? What is a decompression stop? Photo credit: ©Roger Harris, JunglePhotos.com

Dive Tables What are dive tables?

Charles’ Law pressure the bends diving scuba atmospheres hyperbaric chamber embolism depth Boyle’s Law ascent nitrogen volume dissolved gases temperature

Why are gas laws important to divers?

Graph of Dive Table Data
Bottom time at 70 feet (21.3 m)

Misinterpreting Dive Computers
Dive computers help divers avoid risky situations Algorithms use depth and time data to estimate the likelihood of decompression sickness What happens if a diver cannot use the dive computer properly?

How Does Volume Decrease With Depth?
Assume a pressure increase of 1.0 atm for every 10 m of depth below the surface Use Boyle’s Law to calculate the volume of a gas at pressure every 10 m of depth to 100 m Chart your data on graph paper

Volume Decreases With Depth

Depth vs. Volume vs. Pressure

Example Problem If a diver runs out of air at 15 meters underwater, and immediately ascends to the surface, what are the risks of an embolism? Assumptions: Constant temperature Lung capacity = 6 liters Pressure increases 1.0 atmosphere for every 10.0 m of depth below the surface.

Solution Let P1, V1 and T1 = pressure, volume, and temperature respectively of the air in the diver’s lungs when the last underwater breath is taken. Let P2, V2 and T2 = pressure, volume, and temperature respectively of the air in the diver’s lungs at the surface. Since temperature is constant, use Boyle’s Law to calculate the increase in volume of the air in the diver’s lungs: P1 x V1=P2 x V2 Rearranging, V2 = (P1 x V1)/P2 P1 (pressure at depth) = 2.5 atmospheres, since = 15/10 x 1 = = 2.5 (Note that 1 is added for the 1 atmosphere of pressure at the surface.) V1 (initial breath underwater) = 6 L P2 (pressure at surface) = 1 atm Therefore, V2 = (2.5 x 6)/1 = 15 L The volume of air is 15 liters at the surface. Since lung capacity is 6 liters, embolism is inevitable.

Enriched Air (Nitrox) How might divers overcome the problem of absorbing nitrogen which then turns to bubbles if the ascent is too fast? Breathing pure oxygen under pressure carries a significant risk of oxygen toxicity. Divers use enriched air (nitrox) which has a higher percentage of oxygen than air, but with a lower risk of oxygen toxicity.

Enriched Air (Nitrox)

Essential Questions What conditions lead to perils facing divers including the bends, embolism and oxygen toxicity? How do the gas laws relate to safe diving practices? What are the principles of Boyle’s Law and Charles’ Law? How is the ideal gas law used to calculate changes in volume, pressure and temperature when one or the other variables is held constant?

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