A doubt from Kinetic Theory of Gases

In summary, the conversation discusses the average translation kinetic energy for ideal gases and how it relates to the degrees of freedom of the gas molecules. It is established that an ideal gas has three degrees of freedom due to its lack of structure, and more complex molecules can have additional degrees of freedom for energy storage. The concept of an ideal gas is well established and typically refers to gases with no volume, perfectly elastic collisions, and no internal degrees of freedom. The conversation also mentions the ability to edit posts on the forum.
  • #1
vijayram
26
1
I have read Average translation kinetic energy is 1/2RT per degree of freedom and Average translation kinetic energy for an ideal gases is 3/2RT.How? Does it imply f=3 for all ideal gases?
 
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  • #2
There are three dimensions in our universe, so translation always has 3 degrees of freedom. Even if you constrain motion in one dimension: if the constraint is not too strong (or if you ignore quantum effects), it will have motion in that dimension.
 
  • #3
Yes, exactly. The ideal gas molecule has no structure, so it can't rotate or vibrate. All it can do it move in three dimensions. f=3. Atomic gasses (at least the ones that don't make dimers) approximate this case and differ from the ideal primarily only in that they take up space reducing the available volume. More complicated molecules can vibrate and rotate and each independent motion is another degree of freedom where energy can be stored. The gas has more capacity to store energy as a function of temperature.
 
  • #4
well this is weird. I keep trying to quote mfb's comment above about an ideal gas with f>3, and the app keeps quoting an entirely different post, one that doesn't even show up in the thread (on my browser anyway)

In any case, regarding the comment that an ideal gas can have degrees of freedom >3:

No volume, only perfectly elastic collisions, and no internal degrees of freedom is the definition of an ideal gas. I wouldn't be surprised if someone somewhere postulated and "an ideal diatomic gas" or similar, but in almost all contexts the meaning of "ideal gas" is extremely well established.
 
  • #5
mike.Albert99 said:
Yes, exactly. The ideal gas molecule has no structure, so it can't rotate or vibrate. All it can do it move in three dimensions. f=3. Atomic gasses (at least the ones that don't make dimers) approximate this case and differ from the ideal primarily only in that they take up space reducing the available volume. More complicated molecules can vibrate and rotate and each independent motion is another degree of freedom where energy can be stored. The gas has more capacity to store energy as a function of temperature.

I know, replying to myself is goofy ...

I just wanted to correct myself in that the atoms in an atomic gasses also have to have negligible chemical interactions in order to be close to ideal. That means noble gasses for the most part.
 
  • #6
I misunderstood the first post, then edited my post. I guess you opened the thread when the old text was there, but if you quote it the forum loads the current text. Ignore the old text.
mike.Albert99 said:
I just wanted to correct myself
You can edit your posts.
 
  • #7
mfb said:
You can edit your posts.

Oh, that would be useful. I don't see a way to do it here (accessing by browser) do I need the app?
 
  • #8
I don't know if it is possible via the app, it is certainly possible via the browser (bottom left of a post), but there is some time limit to it.
 

1. What is the Kinetic Theory of Gases?

The Kinetic Theory of Gases is a scientific model that explains the behavior of gases based on the motion of their individual molecules. It states that gas molecules are in constant, random motion and that the temperature of a gas is directly proportional to the average kinetic energy of its molecules.

2. How do gas molecules behave according to the Kinetic Theory of Gases?

According to the Kinetic Theory of Gases, gas molecules are in constant motion and collide with each other and the walls of their container. These collisions create pressure and cause the gas to expand and fill its container. Gas molecules also have different speeds and directions of motion, leading to their random motion.

3. What are the assumptions of the Kinetic Theory of Gases?

The Kinetic Theory of Gases is based on the following assumptions: 1) Gas particles are considered to be point masses with no volume, 2) Gas molecules are in constant, random motion, 3) Collisions between gas molecules are perfectly elastic, meaning there is no loss of energy during collisions, and 4) Gas molecules do not exert attractive or repulsive forces on each other.

4. How does temperature affect the behavior of gases according to the Kinetic Theory?

The Kinetic Theory of Gases states that as the temperature of a gas increases, the average kinetic energy of its molecules also increases. This means that the gas molecules will move faster and collide more frequently, resulting in an increase in pressure and volume.

5. What is the relationship between pressure, volume, and temperature in the Kinetic Theory of Gases?

The Kinetic Theory of Gases explains the relationship between pressure, volume, and temperature using the ideal gas law: PV = nRT. This equation states that the pressure and volume of a gas are inversely proportional to each other, while the pressure and temperature are directly proportional to each other. This relationship is known as Boyle's Law and Charles' Law, respectively, and can be explained by the behavior of gas molecules according to the Kinetic Theory.

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