Gases are an important part of our daily life. Their properties and characteristics allow us to explore the ways in which they interact.
Today we will look at the characteristics of gases, their behavior according to Ideal Gas Theory, and deviation from this theory.
It is easier to visualize and make predictions for gases in a more ideal manner. In basic level chemistry courses, this simplified understanding of gases is crucial to building a foundation of knowledge regarding interactions with gases. Let us look at ideal gases.
In an ideal and theoretical world of chemistry, gases would behave perfectly, without losing energy and sticking to our easy calculations. Alas, this is not the case. But looking at the criteria for ideal gases can help us contrast them with real gases.
1. Gas Particles are in continuous, rapid, random motion
2. There are no attractive forces between particles
3. The gas particles do not take up volume
4. Collisions between particles and the container are perfectly elastic
5. Temperature of the gas is proportional to the average kinetic energy of the molecules.
Many of these assumptions are almost met by real gases, however the differences can lead to great variations in our calculations. Lets look at a few key contrasts:
Gas Particles do, in fact, have intermolecular forces between particles. this makes their motions less rapid and increases its deviation from ideal gas. In fact, due to the intermolecular forces, the pressure that the gases exert on the container walls are less, as the molecules impact with the walls less often than is predicted with ideal gas law. Furthermore, the gas particles do take up a significant amount of volume, which tends to mess up our calculations. Finally, collisions for molecules with each other and the wall aren't fully elastic. Some of the energy that the molecules have is inevitably lost during the collisions.
It is also important to note that real gases act most like ideal gases at high temperatures and low pressures. Given two gases, and asked to select the one that will behave in a more ideal manner, choose the one with the higher temperature and lower pressure.
Now that we have contrasted ideal gases with real gases, we can explore the relationships between the levels of variables for gases.
The ideal gas law, PV = nRT* roughly gives a good approximation of the behaviour of gases under ideal conditions (High Pressure, Low Temperture. The variables from left to right in the equation represent; pressure (in atm, torr, or mmHg), volume, moles of gas, ideal gas constant (different values based on context), and temperature (convert to Kelvins).
From this law we can see many different relationships between the variables
Pressure vs Volume (Boyle's Law): Pressure of an ideal gas is inversely proportional to its volume (holding other variables constant) Explanation: Greater volumes decrease the frequency of collisions, thus decreasing the pressure
Pressure vs Temperature (Gay-Lussac's Law): Pressure of an ideal gas is directly proportional its absolute temperature (holding other variables constant) Explanation: Greater temperature increases the momentum and frequency of collisions for gas molecules, thus increasing pressure
Pressure vs Moles: Pressure of an ideal gas is directly proportional to the amount of moles of gas the container has (holding other variables constant) Explanation: Greater amount of moles make collisions between molecules and container wall more frequent, thus increasing pressure
Volume vs Moles (Avogadro's Law): Volume that a gas takes up is directly proportional to the amount of moles of gas that a container has (holding other variables constant). Explanation: Increasing the amount of moles would require a greater volume to keep the pressure constant, thus increasing volume
Volume vs Temperature (Charle's Law): Volume tat a gas takes up is directly proportional to the temperature of a gas (holding other variables constant). Explanation: Increasing temperature while keeping pressure and moles constant would require greater volume, thus increasing volume.
This isnt all! Look at the upcoming article to find out more about gases. There we will explore even more Kinetic Molecular Theory. For now, happy studying!
*note that there is an alternative equation that allows us to make better predictions about real gases. The Van der Waals equation has two more factors, a and b, which are different for each gas and are experimentally determined.
Van der Waals equation: [P + (an^2)/V^2)][V - nb) = nRT
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