Chemical Forums
Chemistry Forums for Students => High School Chemistry Forum => Topic started by: Shahab Mirza on July 12, 2015, 08:51:18 AM
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Hello,
I want to ask about Ideal gas , There is always written about ideal gases that " Real gases shows Ideal Behaviour at High temperature and Low Pressure" , and Ideal gases are those gases which have no inter molecular attaraction , and occupy no space .
So , According to Boyles Law . If I set Low Pressure then Volume will be high , and According to Charles Law If I set temperature to High then there will be more K.E hence more intermolecular attaraction hence mean free path and avg distance increases hence volume will be high, then How real gas behave Idealy at Low Pressure and High temperature ? how Decreasing Pressure can decrease volume ?
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You are confusing volume occupied by the gas with the volume occupied by its molecules. Volume occupied by the gas is the volume of the molecules plus the space between them.
You are also confusing collisions with attraction between the molecules. Collisions are the only kind of interaction between molecules of ideal gas, molecules of real gas also attract each other.
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, and According to Charles Law If I set temperature to High then there will be more K.E hence more intermolecular attaraction hence mean free path and avg distance increases hence volume will be high, then How real gas behave Idealy at Low Pressure and High temperature ? how Decreasing Pressure can decrease volume ?
That's the thing that's confusing -- higher temperature increases collisions in an ideal gas, or even a real gas. So, some of your conclusions are just incorrect. You mix them up in between correct statements, so it becomes a muddle. Hint: try to express them as formulas or a table of statements,instead of a paragraph, to avoid confusion.
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, and According to Charles Law If I set temperature to High then there will be more K.E hence more intermolecular attaraction hence mean free path and avg distance increases hence volume will be high, then How real gas behave Idealy at Low Pressure and High temperature ? how Decreasing Pressure can decrease volume ?
That's the thing that's confusing -- higher temperature increases collisions in an ideal gas, or even a real gas. So, some of your conclusions are just incorrect. You mix them up in between correct statements, so it becomes a muddle. Hint: try to express them as formulas or a table of statements,instead of a paragraph, to avoid confusion.
You are confusing volume occupied by the gas with the volume occupied by its molecules. Volume occupied by the gas is the volume of the molecules plus the space between them.
You are also confusing collisions with attraction between the molecules. Collisions are the only kind of interaction between molecules of ideal gas, molecules of real gas also attract each other.
Thanks for reply Sir . But then tell me that what to speak about when it is asked that how real gases behave Ideally at high temperature and low pressure.
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Ideal gas is an imaginary gas whose molecules move randomly and have no intermolecular interaction except the elastic collisions.
But in reality none of the gases are 100% percent ideal, because they interact with each other through many types of intermolecular interactions.
At high temperature and low pressure the intermolecular interactions decrease, but still they are not 100% ideal.
At high temperature, they have high kinetic energy and hence they move very fast and only collide with each other and move away fast. At low pressure, the number of molecules per volume is less and hence the chances of the molecules interacting with each other becomes less as well, which means they are close to ideal gas behavior.
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At elevated temperature and low pressure (sufficiently far from critical parameters) any real gas can be treated as the ideal one in most calculations. This is only better or worser approximation. This statement concerns also any (advanced) equations of gas state. These equations are only better aproximations, but they are not absolutely exact.