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## Class 11 Physics (India)

Created by Ryan Scott Patton.

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• Where does .041mol come from @ •  A great question! It took me some time to find the answer and like the most of us watching the video I just took it to notice and didn't think about it. but dividing 1 through 22,4 gives a slightly different number of moles,
But let me explain it what i found out:

He is saying starting @ "... if you filled up four balloons to exactly one liter at 25 degrees celsius..."
and the key point is the one liter at 25°C and it took me some browsing on google looking for how much volume takes one mole of ideal gas at 25°C.
- We know that one mole of ideal gas at STP (Standard Temperature and Pressure- which means 0°C and 1 atmosphere of pressure) takes 22,4 liters
But how much of volume increase would it be if we heat up our balloons from 0°C to room temperature or more exactly 25°C (presuming that we are filling the four balloons somewhere at the garage or some simple room and not inside a pressure chamber where you can change the atmospheric pressure, so our pressure would stay 1 atmosphere.
And finally I found this at wikipedia over the ideal gases:
"The molar volume of an ideal gas at one atmosphere of pressure is

22,414 dm3/mol at 0°C
24,465 dm3/mol at 25 °C"

dm3 equals to liters so by rearranging we can figure out that if we would have 1 mol of gas at one atmosphere and 25°C it would make our balloon 24,465 liters big, but here we have a 1 liter big balloon so it contains 1 / 24,465 moles which is 0,0408 rounded up to....

yes exactly here it is where it comes from the 0,041 moles

References:
http://en.wikipedia.org/wiki/Molar_volume
• this may be a dumb question but im asking anyways..
Different gas molecules have different sizes, then why do the different gases at equal volume have same number of molecules ....??
To make it more clear take an example of hydrogen and methane ...
How can same number of molecules of methane fit in 1 litre of container as by hydrogen when methane molecule is obviously bigger than hydrogen molecule ?
plz answer fast cause my exams in a week • How Did He Get That .97 in the problem as n2 • but what about ,force=mass *acc., so, molecules having more mass should have more pressure as compared to oters having less mass or less volume? • boi someone explain more simple I don't get it • Okay so, I'm using this video to revise for my Chemistry end-of-year, and I wanted to ask a bit of a stupid question: Where is s.t.p. applicable? Usually, questions I've encountered about molar gas volume involve compounds at r.t.p., so where is s.t.p. used? Is s.t.p. a better measure than r.t.p.? • S.T.P is used when the gas is present at these conditions. In which when gas is at 1 atm and 273K then the gas will also be 22.4 L per mole. Usually, within problems, it is clearly stated when gas is at S.T.P. This is usually used within combined gas laws, ideal gas law, as well as problems relating to specific laws (i.e. Boyle's law, Avogadro's law, etc.). Between S.T.P and R.T.P, usually, S.T.P is mainly used within gas law problems especially.
• Thank You for explaining this better than my teacher • 22.5 liters of something in the gas phase at STP = 1 mol does not make any sense because:

1) Molar mass is different between 1 gas and another and the size of the molecules is also different(Like CH4 vs O2 for example with CH4 being larger than O2)

2) What if you lower the pressure of something so that it becomes a gas and then raise it to STP? The molar volume while still in the gas phase would be different because of pressure initially being lowered.

3) every gas has a different melting point and boiling point so that N2 vs O2 will have one of them(The O2 I think) moving around really fast and the other a little slower and that would effect the molar volume. • Consider this:
22.4 L of something in the gas phase at STP = 1 mol is correct because
1) We don't take into consideration molar mass of the gas, we look at the NUMBER OF MOLES. Now for any gas, 1 mole means the exact same thing. If its O2, then in mole O2 there are
6.022* 10^23 O2 molecules. In 1 mole of CH4, there are 6.022* 10^23 CH4 molecules and so on. That doesn't change.
Also, one of the assumptions of the ideal gas equation is that we consider each and any molecule or atom (if the gas is momoatomic) a point mass. We consider its atomic/molecular volume negligible. I think this one in one of the videos... :)

2) Since Volume is inversely proportional to Pressure (Boyle's Law) once you take away the pressure, the volume (which decreased when pressure increased) should return to 22.4 L again (if n = 1 mole). The two work together.

3) You are right about boiling point. But if you remember, we standardised the Temperature. Temperature is basically the average kinetic energy of a gas (they have a video on this too). So if temperature of both gases is same, their kinetic energy is same, so their molecules will have equal amount of energy - so that will not affect molar volume.  