What are the deviations of the ideal gas accounted for in the equations of state of real gass?

Answer 1

An ideal gas, by definition, follows the Ideal Gas Law, which states PV=nRT. Any behavior for which that equation does not hold is considered non-ideal. What then are the causes of non-ideal behavior? The Ideal Gas Law doesn't work for many gases (in other words, many gas are not actually ideal) because the Gas Law makes two assumptions, that in certain conditions break down. Assumption #1 is that there are no interactions between atoms/molecules in the gas phase. In this model, there are no attractive or repulsive forces between two neighboring atoms/molecules in the gas phase. This is not always correct, and especially at very low temperatures, gases tend to condense, and so attractive forces between them start to be significant. Attractive forces tend to make the measured pressure lower than it is predicted to be. Assumption #2 is that the volume of the container holding the gas is infinitely larger than the volume taken up by the gas molecules themselves. In other words, it assumes that molecules have zero volume, which is of course not true. This assumption breaks down significantly at very high pressures, where the volume taken up by the gas is significant compared to the volume of the container. To correct for this, the molecular volume taken up by the gas is subtracted from the volume of the empty container. Therefore, there are significant deviations from the Ideal Gas Law at high pressures or very low temperatures. The actual amount of deviation depends on the molecules individual properties. H2 gas or He gas are both very "ideal" gases under most conditions. However, H2O, with strong intermolecular attractive forces, or SO2 (a fairly large molecule also with strong intermolecular forces) do not obey the Ideal Gas Law under most conditions.

Answer 2

Ideal gas is only a term used to represent when we are assuming the behavior of a gas fits the principles of an ideal gas, so in reality all gases are real gases and when we use the term ideal gas when we are doing calculations . The difference between the two is that when looking at ideal gases the equations used assume that 1) gases occupy a negligible fraction of the total amount of gas and 2) The interaction between the individual gas molecules with each other is negligible. These two assumptions cause a difference between real and ideal gases, since gases in real life do in fact take up space and have intermolecular interaction.

Answer 3

actually the real gases deviate from the ideal behaviou under two conditions

• at high pressure
• at low temperature

and actually under these conditions the two of the postulates get disproved for real gas

i.e one that there are no intermolecular forces between the molecules of an ideal gas but these forces become very significant for the real gases at high pressure and low temperature .. and make the gas to deviate from the ideality

and at high pressure the volume of the gas molecules become very significant disobeying the postulate that the volume is negligilble as compared to the volume of the container ... so the decrase in volume of the real gases is not same for real gases as that for ideal gas...

Answer 4

Real gases deviate from ideal gases because the weight of their molecules are not negligible, the force of attraction between their molecules are also not negligible... Also real gases obey the vander waals equation of state while ideal gases obey dalton's/kinetic equation of state

Answer 5

Gases will deviate from ideal behavior when any intermolecular bonding happens. For instance, water vapor in high concentrations will form hydrogen bonds, pulling the affected molecules together. This will cause a net decrease in pressure will everything else equal. Also, during molecular collisions, there is a possibility of an electron moving from one molecule to the next, charging both. This would produce as slight electromagnetic repulsion, increasing pressure with all else equal (ceteris paribus).

Answer 6

There are (1)attraction forces and (2) the fact that the molecules actually take up some space, which at high pressures and low temperatures can make a large difference in volume from what the ideal gas law predicts.

Answer 7

Low temperatures and high pressures cause gases to deviate from ideal gas behavior.

Answer 8

It explodes

:) :)