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Despite the fact that all bodies that come into contact with each other exert pressure on each other, pressure is a physical magnitude that we tend to associate much more frequently with gases than with solid bodies.
In physics, pressure is defined as force per unit area, and is given by the ratio F/A. This means that, to modify the pressure, we only have to modify either the force, or the area on which the force is applied. For example, if we wanted to increase the pressure we exert on, say, the surface of a table, we could increase the force (for example, by adding more weight or pressing more on the table), we could decrease the area on which we apply the force ( for example applying the force with the tip of a nail instead of with our hand), or both things at the same time.
However, how can we increase the pressure exerted by a gas? Furthermore, how is it that gases, being so ethereal and shapeless, are able to exert pressure on the walls of the containers that contain them? Understanding these aspects of one of the most important properties of gases is extremely important, as it allows us to understand many phenomena that we can observe on a daily basis, from the inflation of automobile tires to the explosion of a sealed can when heated. too much, or even the behavior of the weather.
For this reason, in this article we will explore some basic aspects of the pressure of gases, as well as the three different ways in which we can increase the pressure of a gas.
How do gases exert pressure?
Anyone who has ever attended a procession or a sporting event, such as a football game where a giant flag was unfurled over a crowd of people, will immediately understand how gases exert pressure.
Gases are substances made up of individual particles that move independently and randomly in all directions. When the gas is contained within a closed container, these particles will inevitably collide frequently with the walls of the container. Each collision of a gas particle against the walls of the container is like a hand pushing the flag from below.
The point is that, due to the immense number of particles that can be in any sample of a gas, these collisions occur with a very high frequency, generating an almost constant force that pushes the surface of the container. This is similar to the multiple pushes given to the flag from below by spectators, which does not let the flag fall, but rather keeps it in an almost constant state of tension, as if it were inflated from below.
Factors Affecting the Pressure of a Gas and the Ideal Gas Law
Gases are the simplest systems that chemistry studies. In fact, an ideally behaving gas is fully characterized by only a handful of variables which are the number of moles (n), the volume (V), the temperature (T) and, of course, the pressure ( Q). These four variables (called state functions) define the state of a sample of any gas, which means that if we know them, we all know about the gas and can predict its behavior in different situations.
Despite being four, in reality we only need to know 3 of them, since we can find the fourth one by means of the ideal gas equation of state, also known as the ideal gas law, which is given by :
This means that the pressure of a gas is determined by the values of the other three variables, ie the number of moles, the temperature and the volume, and this relationship can be obtained by isolating P from the ideal gas law, as shown shows below:
How to increase the pressure of a gas
As can be seen in the above equation, the pressure is directly proportional to the number of moles and to the temperature, but inversely proportional to the volume. This means that there are three different ways to increase the pressure, and these are:
Increasing the number of moles of the gas
The fact that the pressure is directly proportional to the number of moles means that the greater the number of moles, the greater the pressure. This implies that one way to increase the pressure is by injecting a greater quantity of the gas into the container that contains it. An example of this is when we inflate the tire or rubber of a car, motorcycle or bicycle, or when we inflate a basketball.
What the pump does is introduce more gas particles into the container. But why does this increase the pressure? To understand it better, we must remember how gases exert pressure. The pressure of the gas is the consequence of the multiple collisions between the gas particles and the walls of the container. If we introduce more particles of the gas, the frequency with which these particles collide with the surface will increase, and therefore the pressure will increase.
increasing the temperature
Pressure is also proportional to temperature. Therefore, as the temperature increases, the pressure will also increase. An everyday situation in which we can see this phenomenon in action is when we overheat a sealed can and it bursts due to the increase in pressure inside.
To understand why temperature affects pressure, we must consider what temperature itself is. Temperature is a measure of the average kinetic energy of the particles that make up a substance. Therefore, changing the temperature implies changing the kinetic energy of the particles. Since they cannot change their mass, then they will necessarily change the speed with which they move.
As the gas particles move faster, two things happen:
- On the one hand, the frequency with which the particles collide with the walls increases, since each particle takes less time to get from one wall to the other. This has the same effect as before increasing the number of particles.
- On top of this, by moving faster, each particle transfers a greater amount of kinetic energy to the wall during the collision, which is another way of saying it hits harder. As more force implies more pressure, then the latter increases.
In summary, the increase in temperature increases the pressure because it causes an increase in the number of collisions and also in the force of each collision.
reducing the volume
Unlike temperature and number of moles, the relationship between pressure and volume is inverse. This means that the lower the volume, the higher the pressure. Therefore, the last way to increase the pressure is to decrease the volume.
Here again the effect has two causes. The first is that, as the volume decreases, the path that each particle must take to get from one wall of the container to the other decreases, so the net frequency of collisions increases. Furthermore, the volume reduction is generally accompanied by a reduction in the surface area exposed to the gas. Remembering the original definition of pressure, as area decreases, pressure increases.
References
Atkins, P., & dePaula, J. (2014). Atkins’ Physical Chemistry (rev. ed.). Oxford, United Kingdom: Oxford University Press.
Brown, T. (2021). Chemistry: The Central Science (11th ed.). London, England: Pearson Education.
Chang, R., Manzo, Á. R., Lopez, PS, & Herranz, ZR (2020). Chemistry (10th ed.). New York City, NY: MCGRAW-HILL.
CK-12 Foundation. (2020, May 18). Factors Affecting Gas Pressure. Retrieved from https://www.ck12.org/chemistry/factors-affecting-gas-pressure/lesson/Factors-Affecting-Gas-Pressure-CHEM/
Flowers, P. (2018, October 19). Relating Pressure, Volume, Amount, and Temperature: The Ideal Gas Law – Chemistry: Atoms First 2e. Retrieved from https://opentextbc.ca/chemistryatomfirst2eopenstax/chapter/relating-pressure-volume-amount-and-temperature-the-ideal-gas-law/
Socratic. (2014, May 26). What causes gas pressure? Retrieved from https://socratic.org/questions/what-causes-gas-pressure
