Teachers wish students would always behave ideal, but in reality life happens and it gets messy.
Teachers wish students would always behave ideal, but in reality life happens and it gets messy (0:32). The same is true for gases: There is a difference between real and ideal gases (1:12). The Ideal Gas Law assumes that particles don’t experience intermolecular forces nor that they take up volume (1:34). But real gases do and this leads to deviations from the Ideal Gas Law, especially under high pressure and close to condensing (2:00). Under low temperatures, intermolecular forces become more and more significant (2:41), leading to a lower pressure of a real gas compared to an ideal gas (3:30). This is enhanced by the strength of intermolecular forces and therefore their polarity (4:36). Because particles do have volume, the usable space is less in a real gas than an ideal gas (5:12). Therefore at high pressures the volume of a real gas is larger than the volume of ideal gases (5:32).
Which molecule would have the least deviation from the Ideal Gas Law?
A. CH4 B. Ne C. H2O D. Cl2
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Hi and welcome to the APsolute Recap: Chemistry Edition. Today’s episode will recap Deviations from the Ideal Gas law.
Lets Zoom out:
Unit 3 - Intermolecular Forces and Properties
Topic - 3.6 Deviation from Ideal Gas Law
Big idea - Structure and Properties
In a teacher’s ideal world, students would always be well behaved, do exactly as they are told and what they are told. Hand in all their assignments on time, read and answer all the questions completely, spend sufficient amounts of time independently reviewing the day’s materials and prepare for the next class. In reality, life happens and it gets messy. Students interact with one another, get distracted with other things than AP Chemistry that take up space in their lives. And it’s the same for gases! Ideally, the gas molecules do not interact with one another and they do not have volume. But we don’t live in an ideal world, but a real world. And so do gases. So let’s take a closer look at how their behavior deviates from the Ideal Gas Law and why.
Let’s zoom in:
The foundation of the Ideal Gas Law is the assumption that gas particles don’t experience attractive or repulsive forces between the particles, but only elastic collisions, and that the gas molecules do not take up volume themselves. These assumptions let us predict the behavior of gases more easily by using the ideal gas law, PV = nRT, which is a great tool for estimations.
But, as well all know, nothing is ever ideal. Real gases do experience intermolecular forces and they have volume themselves and therefore take up space. The ideal gas law is a helpful tool for an approximation and works especially well under low pressure and high temperatures. Under these conditions, the intermolecular forces between the particles and the volume of particles themselves don’t carry much weight. But we also need to be aware of the deviations from the ideal gas law and how real gases behave under different conditions, like when the temperature is close to condensing or the pressure is high.
Let’s start with intermolecular forces. As mentioned before in ideal gases, we neglect the intermolecular forces, especially at high temperatures and in large volumes. But as we lower volume and temperature, the conditions are getting close to a phase change from gas to liquid, aka condensation, and the intermolecular forces of attraction and repulsion become more and more significant: particles experience forces of attraction between each other and therefore they deviate from the straight lines we assume they move in. This results in particles moving towards one another instead of the wall of the container. Even when they are close to the wall of the container, the particles are drawn away from the wall due to intermolecular forces and will hit the wall with less force. Since we use pressure as a measurement of collision of particles with the container, the pressure in a real gas will therefore be lower than in an ideal gas at lower temperatures.
The magnitude of deviation is connected to the strength of intermolecular force and related to the polarity of the molecule. The stronger the intermolecular force the particles are experiencing, the greater the deviation is. For example: the hydrogen molecule is experiencing only very weak London Dispersion forces. Therefore, hydrogen won’t deviate as much from the ideal gas law as, let’s say Ammonia, NH3, which is experiencing hydrogen bonding. If you need a recap of intermolecular forces, check out episode 17!
The second component that leads to a deviation is the volume of particles. In an ideal gas, we assume that atoms have no volume and the entire volume of the container is available. BUT: atoms do have volume and so, the usable space is less. This especially carries weight under high pressures, because at higher pressures the particles are squeezed closer together and their own volume occupies a larger fraction of the container. This means that at high pressures real gases cannot be compressed as much as ideal gases and therefore the volume of real gases is larger than the volume of ideal gases.
To recap:
The ideal gas law assumes that there are no intermolecular forces and atoms do not have volume themselves. It is a great tool for estimations, but in conditions close to condensation and under high pressure, real gases deviate from ideal gases. Two components distinguish real and ideal gases: real gases experience intermolecular forces which lead to a decreased pressure in comparison with ideal gases. This increases with increasing polarity of molecules and is especially significant close to condensation. At high pressures, the volume occupied by real gases becomes significant and leads to real gases having a larger volume than ideal gases.
Coming up next on the APsolute RecAP Chemistry Edition: Beer-Lambert Law
Today’s Question of the day is about Deviations from the Ideal Gas Law.
Which molecule would have the least deviation from the Ideal Gas Law?