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Richard Towneley and Henry Power were the first to notice the link between pressure and volume in the 17th century. Robert Boyle validated their invention through various experiments and published the final results. According to Robert Gunther and other experts, Boyle’s assistant, Robert Hooke, developed the experimental apparatus. The Law of Boyle is based on the experiments which consisted of air, which was still considered one of the four elements at the time, but Boyle disagreed. Air was still considered one of the four components, but Boyle disagreed. Boyle’s interest was most likely in comprehending air as a vital component of life; for example, Boyle wrote research on growing plants without air.
What is gas compression?
The amount of air present in our surroundings is vast, and all living things need air to survive, which contains oxygen. However, living creatures aren’t the only ones who need air. Air is also required for the effective and efficient operation of engines and equipment. Air is essential in various industries, including transportation, factories, manufacturing plants, and services. However, this air is not the same as the air we inhale.
What do you mean by “compressed air”?
Many molecules and gas components make up atmospheric air, which is the air that we are continually surrounded by. The following is an example of the composition:
Nitrogen molecules account for 78 %.
Oxygen molecules make about 20-21% of the total.
Water vapour, carbon dioxide, and other gas molecules make up 1-2 %.
These molecules spontaneously space themselves out as gases, filling the available volume in free space.
When the air is compressed, those molecules are forced closer together. The air contains the same molecules, but they take up less space (vol) than when they were free. The effects of squeezing these gas together are incredible:
Basic Theory of Compressing Air
Nitrogen and oxygen are the two main components of the air we breathe.
Although air is not a “perfect” or pure gas, the existence of oxygen and nitrogen in significant proportions causes it to behave very similarly to a “perfect” gas. Perfect gases are known to follow a few rules:
Boyle’s law (PV = k)
Charles’s law (V/T = k)
Here P, V, T, K defines:
k = constant
V = Volume
P = Pressure
T = Temperature
The combination of the Charles and Boyle’s law can be represented as:
PV/T = k
Derivation and Formula
The Law of Gay-Lussac states that the ratio of initial pressure and temperature of a gas is equal to the final pressure and temperature for a gas that is kept at constant volume in fixed mass. This formula is expressed as:
(P1/T1) = (P2/T2)
In the above equation, P1, P2, T1, and T2 represent:-
P1 – initial pressure
P2 – final pressure
T2 – final temperature
T1 – initial temperature
The above expression could be derived from pressure and temperature proportionality for the gas. Since P ∝ T for gases of fixed mass kept at constant volume:
P1/T1 = k (constant) (initial pressure/ initial temp= constant)
P2/T2 = k (constant) (final pressure/ final temp= constant)
hence, P1/T1 = P2/T2 = k
Or, P1 x T2 = P2 x T1
Gay-Lussac’s Law Examples
When a pressurised aerosol can (spray-paint can or deodorant can) is warmed, the pressure imposed by the gas on the container increases, resulting in explosions. Many pressurised containers have warning warnings stating that they must be maintained in a cold environment and kept away from fire.
Pressure cookers are another example of Gay-law Lussac’s in action. The pressure applied by the steam inside the vessel grows as the cooker heats up, and the food cooks faster due to the high temp & pressure inside the container.
Assumptions for the Ideal Gas Model
The ideal gas model makes many assumptions. The following are the details:
Gas molecules are thought of as little, rigid spheres indistinguishable from one another.
There is no energy loss in collisions or motion since all motions are frictionless & collisions are elastic.
Ideal gas follows all the Laws of Newton.
The avg distance between molecules is substantially lower than the size of the molecules.
Molecules are constantly moving in random directions and at different speeds.
Besides point-like collisions with walls, molecules do not attract or repel one another.
There are no long-range forces between the gas molecules and their environment.
Conclusion
In the above chapter, we have read about pressure and gas compression and the definition of Gay-Lussac’s Law. An ideal gas is a hypothetical gas in which gas particles travel arbitrarily with no interparticle interaction. In actuality, there is no such thing as a perfect gas. It is based on the ideal gas equation, a simple equation that we will learn more about and that can be analysed using statistical mechanics. Most gases are assumed to behave like an ideal gas under conventional pressure and temperature conditions. According to IUPAC, at constant pressure and temperature, 1 mole of an ideal gas has a potential of 22.71 litres. The formula for the Gay Lussacs law is denoted by (P1/T1) = (P2/T2).
