Now, we know that that isn't exactly true, that individual molecules of course do take up volume. But this is a reasonable assumption, because generally speaking, it might be a very, very infinitesimally small fraction of the total volume of the space that they are bouncing around in. And so these are the two assumptions we make when we talk about ideal gasses.
That's why we're using the word ideal. In future videos we'll talk about non-ideal behavior. But it allows us to make some simplifications that approximate a lot of the world. So let's think about how we can describe ideal gasses. We can think about the volume of the container that they are in.
We could imagine the pressure that they would exert on say the inside of the container. That's how I visualize it. Although, that pressure would be the same at any point inside of the container. We can think about the temperature. And we wanna do it in absolute scale, so we generally measure temperature in kelvin.
And then we could also think about just how much of that gas we have. And we can measure that in terms of number of moles.
And so that's what this lowercase n is. So let's think about how these four things can relate to each other. So let's just always put volume on the left-hand side.
How does volume relate to pressure? Well, what I imagine is, if I have a balloon like this and I have some gas in the balloon, if I try to decrease the volume by making it a smaller balloon without letting out any other air or without changing the temperature, so I'm not changing T and n, what's going to happen to the pressure?
Well, that gas is going to, per square inch or per square area, exert more and more force. It gets harder and harder for me to squeeze that balloon. So as volume goes down, pressure goes up. Or likewise, if I were to make the container bigger, not changing, once again, the temperature or the number of moles I have inside of the container, it's going to lower the pressure.
So it looks like volume and pressure move inversely with each other. So what we could say is that volume is proportional to one over pressure, the inverse of pressure.
How do you know which ideal gas constant to use? What is the ideal gas constant for butane? Why is ideal gas law in kelvin? Why is the ideal gas constant important?
What is the di-electric constant? What volume L will 0. How do you calculate the molar mass of a gas? When should I use the ideal gas law and not the combined gas law? All collisions, both between the molecules themselves, and between the molecules and the walls of the container, are perfectly elastic. That means that there is no loss of kinetic energy during the collision. And then two absolutely key assumptions, because these are the two most important ways in which real gases differ from ideal gases:.
The volume occupied by the molecules themselves is entirely negligible relative to the volume of the container. On the whole, this is an easy equation to remember and use. The problems lie almost entirely in the units. I am assuming below that you are working in strict SI units as you will be if you are doing a UK-based exam, for example.
Pressure is measured in pascals, Pa - sometimes expressed as newtons per square metre, N m These mean exactly the same thing. Be careful if you are given pressures in kPa kilopascals. For example, kPa is Pa. You must make that conversion before you use the ideal gas equation.
This is the most likely place for you to go wrong when you use this equation. That's because the SI unit of volume is the cubic metre, m 3 - not cm 3 or dm 3.
So if you are inserting values of volume into the equation, you first have to convert them into cubic metres. Similarly, if you are working out a volume using the equation, remember to covert the answer in cubic metres into dm 3 or cm 3 if you need to - this time by multiplying by a or a million.
If you get this wrong, you are going to end up with a silly answer, out by a factor of a thousand or a million. So it is usually fairly obvious if you have done something wrong, and you can check back again. This is easy, of course - it is just a number. You already know that you work it out by dividing the mass in grams by the mass of one mole in grams. I don't recommend that you remember the ideal gas equation in this form, but you must be confident that you can convert it into this form.
A value for R will be given you if you need it, or you can look it up in a data source. The SI value for R is 8. Note: You may come across other values for this with different units. A commonly used one in the past was The units tell you that the volume would be in cubic centimetres and the pressure in atmospheres.
Unfortunately the units in the SI version aren't so obviously helpful.
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