[Vadevalor]

Q1:why are some gases more ideal than another under the same
conditions and both gases of the same amount?

Q2:Why is the PV against P graph of a real gas like a tick?why does it decrease then increase?if higher pressure the real gas is less ideal then why not just a straight line up from the origin? Why PV decrease :0

[Plato's Trousers]

(Original post by Vadevalor)
Q1:why are some gases more ideal than another under the same
conditions and both gases of the same amount?

Q2:Why is the PV against P graph of a real gas like a tick?why does it decrease then increase?if higher pressure the real gas is less ideal then why not just a straight line up from the origin? Why PV decrease :0

well, what do you know about what "ideal" means in the context of a gas?

[Vadevalor]

I know that they have no intermolecular attn and volume is negligible compared to that of container.
And real gases are most ideal under high temp and low pressure...


This was posted from The Student Room's iPhone/iPad A

[Plato's Trousers]

(Original post by Vadevalor)
I know that they have no intermolecular attn and volume is negligible compared to that of container.
And real gases are most ideal under high temp and low pressure...


This was posted from The Student Room's iPhone/iPad A

Ok. And do you know what it is that makes gases depart from ideal behaviour? If you do, that will basically give you the answer to Q1

[Vadevalor]

Uhh..if i knew the answer i wouldnt be asking i know how a gas can be more ideal under some conditions,but couldnt figure out how to compare between gases their ideality.


This was posted from The Student Room's iPhone/iPad A

[charco]

(Original post by Vadevalor)
Uhh..if i knew the answer i wouldnt be asking i know how a gas can be more ideal under some conditions,but couldnt figure out how to compare between gases their ideality.


This was posted from The Student Room's iPhone/iPad A

One of the assumptions is that gas particles themselves do not occupy any volume, in other words each particle has all of the available volume to move around in.

What is going to happen when the pressure is very high (in other words there are many particles in a small volume)?

Can each particle move around in the whole volume?

Are gas particles themselves all the same volumes for different gases?




This wasn't posted from an iPad, iPod, iPud or any other new-fangled and largely overpriced device ...

[Alchem]

(Original post by Vadevalor)
Q1:why are some gases more ideal than another under the same
conditions and both gases of the same amount?

Q2:Why is the PV against P graph of a real gas like a tick?why does it decrease then increase?if higher pressure the real gas is less ideal then why not just a straight line up from the origin? Why PV decrease :0

Question 1:

Plato's Trousers has already given you a hint about Q1. As you yourself have said, intermolecular attractions don't exist between the molecules of an ideal gas . Hence, a non-ideal gas will be one in which intermolecular attractions do exist. To answer your question, a gas with weak intermolecular attractive foreces will behave more ideally than a gas with strong intermolecular attractive forces. Example: N₂ vs HF.
Another factor that matters is the molecular volume (volume of one molecule of gas).

Question 2:

In a real gas, intermolecular attractive forces are present. However, under low pressures, since volume is large and the gas molecules are far apart, the attractive forces will be negligible due to large intermolecular distances, and the gas will behave like an ideal gas. So PV will be nearly constant (PV = nRT), i.e. the plot will be a near straight line parallel to X-axis.
Under moderately high pressures,volume will be much smaller, and intermolecular distances will be smaller too on an average. As a result, intermolecular attractive forces will become stronger. Under such conditions, as one applies slightly more pressure, the volume would decrease faster than before, aided by the stronger intermolecular attractive forces. [This is a key point. I am not sure I've been able to explain it clearly enough]. In other words, as P is increased, V would decrease faster than before, i.e. faster than an ideal gas. This means that PV would start decreasing (PV will be < nRT).

Under very high pressures, volume occupied by the gas will become really small, so small that it will become close to the total volume occupied by the molecules of the gas (remember that a gas molecule does have a finite volume). In this scenario, the molecules will practically start encroaching into one another's territory, and intermolecular repulsive forces will set in. As a result, when P is further increased, V will decrease much more slowly than an ideal gas, due to the resistance offered by intermolecular repulsive forces. So, PV will start increasing again.(P increasing, V decreasing only slightly, so product PV increases)

This explains the shape of a PV against P graph.

Sorry for a long-winded answer.

[Alchem]

What is going to happen when the pressure is very high (in other words there are many particles in a small volume)?

Can each particle move around in the whole volume?

Are gas particles themselves all the same volumes for different gases?




1. When the pressure is very high, the volume occupied by the gas sample will be more than that predicted by the Ideal Gas Equation . [ V > nRT/P]

2. No, the volume available to a molecule of the gas becomes significantly smaller than the volume available to the gas sample (I mean, the collection of all the molecules). In fact, there is an alternate Gas Equation, called the van der Waal's Equation, which explicitly takes this fact into account.

3. Molecules of different gases may certainly have different volumes. For example, a molecule of hydrogen (H₂) will have a much smaller volume than a molecule of dichloromethane - CH₂Cl₂ , which is also a gas at room temperature.

Incidentally, the smallest unit of a gas is called a molecule.

[Vadevalor]

Alchem..your explanation of the graph is awesome..couldnt get that kind of information elsewhere haha awesome!
Nonetheless,thanks others for dropping hints here and there though i'm quite dimwitted to get it.
Here's a summary for self and others' references.
Ans for Q1: different non-ideal gases have different types and strength of intermolecular forces of attn btwn the gas molecules (stronger int molecular forces btwn them= less ideal)

For Q2 :Oh so when too much pressure the volume will not decrease as much since the "limit" is kinda reached! But when moderate pressure limit havent reach
-So PV of real gas will be more than nRT of ideal gas (the horizontal line) under high pressure,PV of real gas will be less than nRT of ideal gas(the hori line) when moderately high pressure

Now i'm gonna challenge myself to think of why for PV against V graph since it's the same shape


This was posted from The Student Room's iPhone/iPad A

[Alchem]

(Original post by Vadevalor)
Alchem..your explanation of the graph is awesome..couldnt get that kind of information elsewhere haha awesome!
Nonetheless,thanks others for dropping hints here and there though i'm quite dimwitted to get it.
Here's a summary for self and others' references.
Ans for Q1: different non-ideal gases have different types and strength of intermolecular forces of attn btwn the gas molecules (stronger int molecular forces btwn them= less ideal)

For Q2 :Oh so when too much pressure the volume will not decrease as much since the "limit" is kinda reached! But when moderate pressure limit havent reach
-So PV of real gas will be more than nRT of ideal gas (the horizontal line) under high pressure,PV of real gas will be less than nRT of ideal gas(the hori line) when moderately high pressure

Now i'm gonna challenge myself to think of why for PV against V graph since it's the same shape


This was posted from The Student Room's iPhone/iPad A

Thanks Vadevalor. I'm glad you liked the explanation.