There are four laws, known as Gas Laws, which describe how gases behave. The four laws are Boyle’s Law, Charles’s Law, Gay-Lussac’s Law and Avogadro’s Law.
Robert Boyle, a famous English chemist, discovered in 1662 that if you pushed on a gas, its volume would decrease proportionately. For example, if you doubled the pressure on a gas (increase the pressure two times), its volume would decrease by half (decrease the volume two times). The opposite is also true. If you reduced the pressure on a gas by 3.5 times, then its volume would increase by 3.5 times. This law is an example of an inverse relationship - if one factor increases, the other factor decreases.
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Pulling up on the lid of a sealed container increases the volume and decreases the pressure. Pushing down on the lid of a sealed container decreases the volume and increases the pressure.
Boyle’s Law in Everyday Life
Here’s a story from British Airways. Back when British Airways was called British Overseas Airways Corporation (BOAC) (before 1974), female flight attendants in the airline were finding that their uniform skirts were fitting on take-off but once they reached cruising altitude, their skirts felt too tight. This tight-skirt mystery was solved using gas laws! A spokesman for BOAC used Boyle’s Law to explain what was going on. He explained that as the pressure in the cabin decreased at the higher altitude, the pressure in the flight attendants’ stomachs also decreased, thus causing the volume of their stomachs to increase (making their stomachs bulge). Since then, female flight attendants wear adjustable skirts.
The working of a syringe can also be explained using Boyle’s Law. When the plunger of a syringe is pulled out, the volume inside the barrel increases, resulting in a decrease in the pressure inside the barrel. Fluids (such as water) flow from a high pressure area to a low pressure area. This means that once the pressure inside a syringe is lower than the pressure outside the syringe, a fluid near the needle (e.g., water, medicine, etc.) will flow into the syringe.
The opposite is also true. When the plunger is pushed back in, the volume decreases and the pressure increases. Once the pressure is greater than that outside the syringe, the fluid inside the barrel will flow out.
The operation of your lungs also can be explained using Boyle’s Law. When you inhale (breathe in), your diaphragm (a large muscle below your lungs) lowers, which increases the volume inside your lungs. This makes the air pressure inside your lungs lower than the air pressure outside your lungs (and your body); therefore, the outside air is drawn into your lungs (much like the syringe). When you exhale (breathe out), your diaphragm pushes upwards, reducing the volume inside your lungs, increasing the pressure and forcing the air outwards.
A weather balloon is a special type of high altitude balloon. These balloons can reach heights of 18 to 37 km above the Earth carrying instruments for measuring atmospheric pressure, temperature and wind among other things. When weather balloons are sent up, they are only partly filled with gas (typically with helium). Why don’t they fill them completely? Short answer – because they would pop! At higher elevations, the air pressure outside the balloon is lower than the pressure of the helium inside the balloon. As Boyle’s Law states, this causes the volume inside the balloon to increase. If the balloon was already full, this increase in volume could cause the balloon’s rubber to stretch beyond its breaking point.
Commercial Diving Instructor and Scuba Instructor
Diving is an amazing activity that allows humans to explore the world under water for relatively long periods of time. Breathing is made possible through either a cylinder of compressed air that a Diver brings with them under water or through an “umbilical” (a type of hose) that is attached to an air supply at the surface.
As a Diving Instructor, I teach people who want to dive for fun and those who need to work underwater. This work can include inspecting and repairing structures such as oil rigs, boats and bridges, conducting archeological or other scientific investigations, search and recovery missions, the production of films, and much more.
Understanding the properties of gases, especially how they behave under pressure and how they interact with the human body is vitally important to my work and to all members of the dive community. Humans are used to normal air pressure (the pressure that is exerted on us by the air in the Earth’s atmosphere) but as we dive deeper under water, the pressure exerted on our bodies increases. As a result, the gases in our bodies (like the air in our sinuses and lungs) get compressed. The Diver’s breathing air is also compressed by the pressure of the surrounding water. This pressure causes nitrogen, an inert gas present in the breathing air, to dissolve in body tissues. There is a limit to how much nitrogen can be absorbed this way without causing harm and consequently we have depth and time limits when diving. When we begin to make our way back to the surface, the pressure decreases and the nitrogen doesn’t stay dissolved in body tissues any more. A very slow and controlled ascent rate allows the expanding nitrogen in our tissues to safely escape.
I also teach students how to calculate the amount of air they will need, which depends on time and the depth of the dive (and how much pressure is exerted on their body).
The best part of the job is having the ability to feel weightless like an astronaut. Seventy percent of the Earth's surface is covered in water and to a Diver there is so much to explore!