At the surface of the Earth the air pressure is around 14.7 psia. I’ve heard the pressure can be explained by looking at a 1" X 1" (square) column of air that extends all the way to the upper atmosphere, and the column of air weighs 14.7 pounds. “And that’s where air pressure comes from.”
I get that. It makes sense.
Now let’s say I have a very rigid jar with a very rigid lid. I remove the lid for a few seconds, and then reinstall the lid. I think we can all agree the pressure inside the jar is 14.7 psia. But there’s no column of air that leads out of the jar and up to the upper atmosphere. So where is the pressure coming from? I would assume the answer would be, “Molecules moving around and banging on everything due to temperature.” I’m good with that. But if I were to (again) remove the lid the story suddenly changes back to the “column of air” explanation.
The column of air above you compresses the air, to the point where the air molecules are moving at an average speed that generate 14.7 psi of pressure.
Air pressure really comes from the combined effect of gas molecules bumping into surfaces. That’s the micro-level model, and it explains everything about gas pressure.
Unfortunately it’s a bit clunky, so often we just talk about the average result of all those molecules and call it “gas”.
Gas pressure comes from confining gas. Instead of your jar, let’s pretend we have a tube with one blocked end and a piston. Put it in space and the gas will expand and the piston will be pushed out. Put it in a different volume of gas and the piston will move inwards or outwards until the pressure inside is equal to the pressure outside.
We can use this model of confined gas to explain atmospheric pressure, and if you do the scenario with the jar won’t be weird. But if you skip the general model and jump straight to the specific idea of atmospheric pressure you have, as you already know, a recipe for confusion.
Or to put it another way, this is wrong. There’s now “changing back”.
With the lid off, the volume within the perimeter of your jar is confined by the walls of the jar and the air pressure at the opening, when you put the lid on it is instead confined by the walls and the lid. But the fundamental idea is that such a volume of gas, at that temperature, is always at that pressure.
Any random collection of air at sea level, whether a on open jar, a jar you just put a lid on, the air in your open fist, the air in the bubble you just blew… it is in every direction 14.7psi (or close, depending on weather, temperature, etc…
As naita says, whether you confine it or not, the air at any point you choose to measure has a pressure of 14.7psi.
But why is all the air that pressure - because all the air around it is. Why is all the air around it at that pressure? Because gravity is pulling down the air. Air may not weigh much, but it has weight, and over a hundred miles or two, that weight matters. think of all those air molecules bouncing off each other and surfaces. Some bounce upward. Even at the molecular level, something bouncing up slows, and as it comes back down it accelerates (until it hits another air molecule, imparting that velocity to it.) Some molecules get hot (achieve high enough velocity) that they can escape the earth entirely - but very very few, and most of those courtesy of collisions with solar particles.
So air is no different than being on the bottom of a football or rugby pile-on - the guy at the bottom is going to feel the weight of three or four layers of bodies bouncing onto him, the guy in the middle one, and the guy on top is simply contributing his body weight. If they were gel sacs instead of muscle the effect would be more pronounced…
Imagine placing a piece of cling wrap on top of the jar. Clearly there is an equilibrium of pressure on both sides of the membrane. The air molecules inside the jar are banging into the membrane with the same speed (on average) as the molecules on the outside, the same number are banging into the membrane on each side - and the membrane doesn’t move.
Now walk up a hill. The column of air above the jar is now less deep, and the pressure less, and you find that the number of molecules of air banging against the membrane are a tiny bit lower. But the number of air molecules in the jar have not changed, and so they remain banging on the membrane, and the membrane is forced upwards. This happens until a new equilibrium is found where the drop in pressure inside the jar (due to the air now having more room to bang around in) decreases, and the number of molecules banging on the membrane from within the jar again matches the number banging on it from the outside. You will find the pressure inside the jar and out side the jar are now the same. Measure and calibrate the bulge in the membrane and you have just built an altimeter. (Or wait for the weather systems to change, and the air pressure to change, and you have a barometer.)
However try the same trick with a solid lid, and all you find is that the lid doesn’t move, and if you move up the hill the pressure in the jar doesn’t change. Eventually, as you continue to climb, the difference in pressure (from the air molecules banging on the lid from the inside versus the much fewer of them banging on the lid from the outside) might pop the lid off.
If you take this analysis far enough and apply a couple of conservation laws (energy, momentum) and the principle of equipartition of energy, you can derive the gas laws for an ideal gas. PV = nRT.
As noted above, gravity is the causative agent here … if I might reword the OP a lil’ bit … air pressure at the surface is 14.7 pounds[sub]force[/sub] per square inch … we add up all the weight[sub]force[/sub] of the air molecules and atoms* in that 1 inch square column … this definition is very useful in meteorology, however once we start talking about closed jars and pistons, maybe the other definition is better suited (i.e. “Pressure is the amount of force applied perpendicular to the surface of an object per unit area.”) …
If you carry the sealed jar up a tall mountain where the air pressure is lower, when you open it, you’ll hear a rush of escaping air as the contained pressure inside equalises with the external pressure of the surrounding air
If you do the same without walking up the mountain, there is no rush of air, because, whilst the jar was containing the air, it wasn’t constraining it to stay pressurised any more or less than it would be it the air was free.
That is, at the original location, the pressures on the inside and outside surfaces of the sealed jar are the same - so it’s not that any switch between column and jar takes place, it’s actually just that the jar is sort of irrelevant to the question of pressure until there is a difference.
The story doesn’t change because the two explanations are the same explanation. The “column of air” imagery is just a way of conceptually quantifying why the surface atmospheric pressure averages just short of about 15 psi. What actually happens is that the weight of all that air compresses the lower atmosphere, creating that pressure at the surface, and that compression pressure exists in all directions. That’s why it doesn’t matter whether the jar is upright, or lying on its side, or upside down. When you put an airtight lid on it, the air remains compressed to the same degree. Molecules moving around due to temperature is kind of beside the point – that’s what temperature is – the salient fact is the higher density of molecules.
One might wonder how this works with water pressure in the ocean, and the answer is: exactly the same way. Water isn’t really incompressible, it just has a very low compressibility. Pressure increases with depth very rapidly and the volume of water does change, just very little. But under the tremendous forces in the deep ocean, it’s more significant. At a depth of 4000 m, water is compressed about 1.8% compared to surface volume at the same temperature, but it takes incredible pressure to do that.