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# Worked example: Calculating partial pressures

In a mixture of ideal gases, each gas behaves independently of the other gases. As a result, we can use the ideal gas law to calculate the partial pressure of each gas in the mixture. Once we know the partial pressures of all of the gases, we can sum them using Dalton's law to find the total pressure of the mixture. Created by Sal Khan.

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• If F=ma and P=F/A, then why doesn't the mass of the particle matter?
• I was wondering the same thing. But then I believe Sal answered the question a bit later at .

If I'm not mistaken, each particle has the same amount of energy as defined by the term: temperature (average kinetic energy). That means that the bigger particles are going to have a lower acceleration than the smaller particles because they all have the same amount of energy, but applied to a higher mass. Although the acceleration is lower in the bigger particles, the pressure contribution upon impact of the surface is the same as a smaller particle because its higher mass then makes up for its lack of acceleration.
• why is oxygen gas usually stored under pressure in metal cylinders?
• Oxygen gas, like all gases are stored under pressure. This is because gases are compressible and are more efficent to store large amounts in a small area. Furthermore as NIGEL 1994 said, by doing so, you'll be able to get out the gas, through diffusion , where areas of high concentration would go to low concentration to acheive an "isotonic" state. Another thing is, if you didn't pressurize the gas, It would take up a large volume, as gas particles disperse from each other. Imagine carrying an oxygen tank underwater that was the size of a humvee!
• So let me get this straight. The O2 and H2 and N2 molecules are all moving at the same velocity, but because the H2 molecules are less massive, they exert less pressure.
• No. The molecules have the same average kinetic energy, they are not all moving at the same velocity. The more massive molecules move more slowly than the lighter molecules.

If temperature and volume are constant, then the pressure that each type of gas in mixture of gases exerts is proportional to the number of molecules of each kind, not the size of the molecules. Thus, the average pressure per molecule is the same no matter the mass of the molecule.
• hey a question just striked my mind when sal said about hydrogen at ..
as we know that in most of the cases a hydrogen atom consists of an electron revolving around the proton(as there are no neutrons).. so my question is that.. IF WE TOOK AN ELECTRON NEAR A ISOLATED PROTON THEN WILL I BE ABLE TO CREATE HYDROGEN ATOM?
• That's entirely possible provided that you manage to put the electron close enough to the proton. The second condition is that the electron moves slow enough in order to give the proton a chance to catch it.
• What is partial pressure and is it specific for a particular gas?
• Partial pressures are determined by the mole fraction of the gas in the mixture, thus there are no specific values for gases. For example, if a mixture contains 1 mole of gas A and 2 moles of gas B and the total pressure is 3 atm. The partial pressure of gas B is 2 atm and 1 atm for gas A. :)
• At , Sal scrolls down to show us the three types of "R"s. My high school offers this Chemistry course in high school so I am doing a little prep work by watching Chemistry videos before starting the course. My question is, when tests come up regarding this topic, do we generally have to memorize these "R"s or are they given to us. I am asking since these look very daunting if we have to memorize these. I know I will be going to a different school but I want to know in your case what happened. Thanks.
• It is not as rigorous as it seems. Unless you have a very strict teacher, you will probably be given formula sheets. If you aren't given formula sheets, you tend to pick up alot of the more difficult topics naturally as you hear about them in class, rather than brute force memorization.
• Suppose you have a 300 mL sample of He at 32oC. Assuming the volume of the container can vary so that the gas pressure is held constant, to what temperature would you have to heat the gas to increase its volume to 475 mL?
• ideal gas law. Number of moles is constant, pressure is constant, so only two things changing is volume and temperature.

T1V1 / V2 = T2
• At , why is it that we don't care about the mass, but the amount of molecules when reguarding pressure? Because isn't hydrogen molecules so much smaller than say oxygen molecules ?
• Ideal gas equation is based on the assumption that there is no forces of attraction or repulsion between the particles. So if we have bigger molecules, they take up space, they have more density. So the molecules will be "closer" to each other such that the forces of attraction or repulsion can't be ignored. This fact is ignored when we deal with ideal gases where the size or mass of the particles are considered to be significantly less. The assumption is made so that people can grasp the concept better. That's why we don't care about the mass. Hope this helps
• I thought you have to use liters in pv=nrt. You can't use m^3 I thought?
(1 vote)
• No, depending on the type of units you use in your expression/question, you can either use cubic metre or litre.
• Will the number of molecules in the container always equal 100?
20 O2
30 H2
50 N2
----------
100moles
• Yes. Moles refers to the number of molecules so it does not matter how you change the temperature/pressure )as long as your molecules do not react).