Andrew mason ideal gas law and adiabatic expansion problem

In summary: However, at high temperatures and high pressures, the kinetic energy of particles becomes significant compared to intermolecular interactions, causing real gases to also approximate ideal behavior. Which approach is correct depends on the conditions of the gas being studied.
  • #1
sunnypic143
4
0
Q1.

ADIABATIC FREE EXPANSION(unrestricted and free)
in vacuum gas a volume of gas is released
expected expansion proceeds withouut change in internal energy therefore temperature is contant (joule's law) . there is no pressure acting on this system threfore no work is done.
now out of 3 thermodynamic variables 2 (p & t) are constant with only v changing
what causes this change ? ie what factor accounts for this change of volume

Q2.
I've read about joule thompson effect , where releasing gases thru a narrow orifice into vacuum causes rapid fall in temperature leading to the condensation of the gas.(used in liquefaction of air)
Isnt this case similar to adiabatic expansion of gases in vaccum? If so there ought not be a temperature change.
So what causes the condensation of gases?

Q3.
Ideal gas laws
Most books define conditions for ideal behaviour as
--> no intermolecular interactions
--> negligibility of molecular dimensions(ie diameter of moleculke<<< inter molecular distances)

So we would expect real gases to approximate to ideal gases at high temperature and low pressures


However, I came across a book where it says another condition in the list
--> the kinetic energy of particles >>> inter molecular interactions
And example cited is dense gases with high temperatures
ie real gases approximate to ideal behaviour at high temperatures and high pressures

So which approach is correct? there is a lot of misconception on this subject at the place where i live and this is a question frequently asked in exams please help.
 
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  • #2
sunnypic143 said:
Q1.

ADIABATIC FREE EXPANSION(unrestricted and free)
in vacuum gas a volume of gas is released
expected expansion proceeds withouut change in internal energy therefore temperature is contant (joule's law) . there is no pressure acting on this system threfore no work is done.
now out of 3 thermodynamic variables 2 (p & t) are constant with only v changing
what causes this change ? ie what factor accounts for this change of volume
The gas on one side of the barrier is at pressure, and the other side of the barrier has vacuum. So, when the barrier is removed, there is a pressure driving force for gas to flow from the high pressure side to the low pressure side.

Q2.
Ive read about joule thompson effect , where releasing gases thru a narrow orifice into vacuum causes rapid fall in temperature leading to the condensation of the gas.(used in liquefaction of air)
Isnt this case similar to adiabatic expansion of gases in vaccum? If so there ought not be a temperature change.
So what causes the condensation of gases?
This is very different from free expansion. In the high pressure chamber, the gas remaining has done work to force the gas ahead of it through the orifice. So its internal energy per unit mass has decreased.
Q3.
Ideal gas laws
Most books define conditions for ideal behaviour as
--> no intermolecular interactions
--> negligibility of molecular dimensions(ie diameter of moleculke<<< inter molecular distances)

So we would expect real gases to approximate to ideal gases at high temperature and low pressures However, I came across a book where it says another condition in the list
--> the kinetic energy of particles >>> inter molecular interactions
And example cited is dense gases with high temperatures
ie real gases approximate to ideal behaviour at high temperatures and high pressures

So which approach is correct? there is a lot of misconception on this subject at the place where i live and this is a question frequently asked in exams please help.
Real gases approximate ideal behavior at high temperatures and low pressures.
 

FAQ: Andrew mason ideal gas law and adiabatic expansion problem

What is the ideal gas law?

The ideal gas law is a fundamental equation in thermodynamics that describes the relationship between pressure, volume, temperature, and the number of moles of a gas. It is typically written as PV = nRT, where P is pressure, V is volume, n is the number of moles, R is the gas constant, and T is the temperature in Kelvin.

How does the ideal gas law apply to Andrew Mason's adiabatic expansion problem?

In Andrew Mason's adiabatic expansion problem, the ideal gas law can be used to determine how changes in pressure, volume, and temperature affect each other. This is because the problem involves a gas undergoing a change in volume without any heat transfer, which is known as an adiabatic process. The ideal gas law can be used to calculate the final temperature and pressure of the gas after the expansion.

What is an adiabatic process?

An adiabatic process is a thermodynamic process in which there is no transfer of heat between a system and its surroundings. This means that the change in internal energy of the system is equal to the work done on or by the system. In Andrew Mason's problem, the adiabatic process refers to the expansion of the gas without any heat transfer.

How does adiabatic expansion affect the temperature of a gas?

Adiabatic expansion causes the temperature of a gas to decrease. This is because the gas is doing work on its surroundings and therefore losing internal energy, which results in a decrease in temperature. This can be seen in Andrew Mason's problem, where the gas expands and its temperature decreases as a result.

What are the limitations of the ideal gas law?

The ideal gas law is a simplified equation that only applies to ideal gases, which do not exist in the real world. Real gases have intermolecular forces and occupy some volume, which can affect their behavior. Additionally, the ideal gas law assumes that the gas is in thermal equilibrium, which may not always be the case. It also does not take into account phase changes or chemical reactions. Therefore, the ideal gas law is limited in its application to real-world scenarios.

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