Significance of Boyle’s Law

Boyle’s law establishes the crucial fact that gases are particularly compressible. The same number of molecules occupy less space when a given amount of gas is compressed by increased pressure.

As a result, the gas’s density rises due to an increase in mass per unit volume. For instance, the air is dense at sea level, but as height increases, both density and pressure decrease. For instance, Mount Everest has a low air pressure of only 0.5 atm. As a result, there is no longer enough oxygen in the air to support regular breathing. This results in symptoms that are typically associated with altitude sickness, such as overall unease, a sluggish sensation, headaches, etc. 

In order to prepare their bodies for the low oxygen pressure, mountain climbers and jawans in high altitudes like Ladakh either undergo lengthy training or carry oxygen cylinders for emergencies.
Similar to this, the cabins of jet aircraft that fly at extremely high altitudes (about 10,000 m) are artificially kept at normal pressure to provide adequate oxygen for breathing. In case of pressure drops, they also have emergency oxygen supplies.

The English chemist Robert Boyle (1627–1691), widely regarded as one of the pioneers of the modern experimental science of chemistry, is commonly credited with this development. He found that increasing the pressure of a sample of contained gas by two times while holding its temperature constant reduced the gas volume by half. According to Boyle’s law, a gas’s volume changes inversely with pressure when the temperature is held constant. This is an illustration of an inverted relationship. The second variable drops when one variable rises in value.