Is it correct to say that there is NO such thing as an ideal gas? I was taught that, technically, a 100% ideal gas only exists in theory because:
a) the atoms are assumed to be point masses which, by definition of a point which has no thickness, exhibit only translational motion (i.e.: no chance of rotation at all)
b) all collisions between atoms are assumed to be perfectly elastic.
If you were taught about the nitty gritty on the definition of an ideal gas, is this how you were taught it? Any other stipulations that I missed? - Jinx
It’s been a little while since I studied the ideal gas equation but you’re mostly correct. There is no such thing as a perfectly ideal gas and even gasses which are generally considered ideal become non-ideal at extreams of temperature and pressure.
The one stipulation that you missed, which is I believe is the most important, is that there are attractive/repulsive forces (electrical) between the molecules in the gas.
There is an equation (the Van der Waal equation) which accounts for the non-ideal behaviour caused by both the non-zero size of the molecules and the forces between the molecules. One downside of the Van der Waal equation is that you need to know the empirical constants for the gas you are working with.
Two general rules to remember: First, in physics, nothing is really ideal. Second, physicists almost always assume that they are.
The key is that you have to know when your approximations break down. When that happens, generally what you do is move to a new set of approximations, which are more difficult to calculate, but which are more accurate in the problem at hand. You never actually drop all of your approximations, because you can’t: We do not yet have a complete model of the world, and probably never will. And, in fact, there are some cases where we know our approximations break down, but we don’t yet have anything to replace them with.
It is a matter of degree. For example, the ideal compressible gas laws for steam (a molecule) flowing through a steam turbine at 1050F and 1000 psi are “very accurate.” Plus there are tests, as previously mentioned (Van der Waal), to determine the degree of “non-ideal” behavior at specific conditions. Also, under certain conditions, you could not prove a gas wasn’t ideal. (double -)
I should mention, to those reading this thread just out of curiosity, the Ideal Gas Law, PV=nRT, is good enough for high school. And, many freshman chemisty courses will teach you to use a “fudge factor”, z, which is read off a graph. The Ideal Gas Law becomes PV=znRT. Usually, a chem prof will give some guideline that gasses where z=>0.95, then we can say it is an ideal gas. Note: the value is a percent expressed as a decimal, and the value of “z” must fall in the range from 0<z<1. Oh, I should say “z” is known as the “compressability factor”, IIRC.
True, Van der Waal’s equation is more accurate, but what a pain in the butt to use!
There are probably dozens of “equations of state” like the ideal gas equation of state. (Good old Redlich-Kwong). Many rely on empirical or other factors to get better results.
I use the ideal gas equation of state all the time not because I deal with ideal gases (the flows I deal with are very hot and probably still reacting a bit when I apply it) but rather because it gives a pretty darn good approximation . Real gas effects just make everything a little bit of a pain to calculate. Even using compressibilty annoys me to no end (curse you ethylene!)
I don’t see using equations like Van Der Waals as being that onerous, since we have these shiny computers and such sitting on our desks or carried in our purses. Even my Excel calculations have a plugin to calculate this. I think it’s only a problem if one is doing estimating or off-the-cuff calculations, in which case fine accuracy isn’t the prime factor anyhow.