Partial Pressure - Video Tutorials & Practice Problems
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Partial Pressure (PGas) is the pressure exerted by an individual gas within a container.
Partial Pressure of Gases
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concept
Partial Pressure
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Partial pressure is the pressure exerted by an individual gas within a mixture. So think of it as the gas's individual pressure. We're gonna say in a container of unreacting gases, total pressure of the container is the sum of partial pressures of each gas. Now this is known as the law of partial pressures. So basically, the total pressure inside of a container comes from adding up all the pressures of each individual gas. So total pressure would equal the pressure of gas 1, plus gas 2, plus gas 3, and so on, if there are additional gases. So just remember, the total pressure that a container is experiencing is contributed by each of the individual gases within it.
In a container of unreacting gases, total pressure of the container is the sum of the partial pressures of each gas.
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example
Partial Pressure Example 1
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Here we're told that a sample of neon gas exerts a pressure of 1.85 atmospheres inside a cylinder. Some nitrogen gas is also present at a pressure of 500 Torr. What is the total pressure inside the cylinder? So remember, we just learned about the law of partial pressures, which tells us that the total pressure felt inside of a container, or in this case a cylinder, comes from adding up the partial pressures of each gas present. So in the container we have neon gas, and we also have nitrogen gas. The total pressure is when you add their partial pressures together. Now the issue is we don't have the same units for these gases. Neon is in atmospheres, but nitrogen is in torres. Since atmospheres is a standard unit we usually use for pressure, let's convert the tor into atmospheres. So we're gonna have 500 torr, and remember that for every 1 atmosphere that's 700 and 62. So when we do that we're gonna get as our atmosphere is 0.65789 atmospheres. Take that and plug it in, and when we do that, we're gonna get a total pressure of 2.5078 9 atmospheres. Within our question, 1.85 has 3 sig figs, 500 only has 1 sig fig. Here, if we wanna buy 1 sig fig, this would round up to 3, which is a pretty big round there in terms of our value. So it's just better to go let's go with the 3 sig figs in this 1.85. Again the question isn't asking for number of sig figs in your final answer, we're doing this as continual continuous practice in terms of determining sig figs. Again, better to go with 3 sig figs. I know it's not the least number of sig figs but going from 2.5 to 3 3 atmospheres is such a big increase. Better just to go with 3 sig figs and then we have 2.51 atmospheres at the end. So now that we've seen this question, let's move on to a next video.
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concept
Partial Pressure
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So we know at this point that the total pressure felt within a container is a result of adding up all the partial pressures of the gases present. Now if we can focus on one of these gases and we know its moles, its temperature, and its volume, we can also find its partial pressure. Now we're going to say here if you assume that the gases behave ideally, then their partial pressures can be calculated from the ideal gas law. We're gonna say here that the pressure of that gas that I'm focusing on, so let's call it gas 1, we can find its partial pressure if we know its moles, so moles 1, r is our gas constant, times the temperature of the container, divided by the volume of the container. So here we're using the ideal gas law to just focus in on 1 gas, and from it determine its partial pressure.
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example
Partial Pressure Example 2
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If 12 grams of helium and 20 grams of oxygen are placed inside a 5 liter cylinder at 30 degrees Celsius, what is the partial pressure of the helium gas? Alright. So they're giving us information on 2 gases. They're giving us information on helium and oxygen. But realize they're only asking for information in terms of partial pressure for the helium. With the helium, I have its grams, and from that I can determine its moles. I have the volume of the container, and I have the temperature of the container. With this information I can find the partial pressure of helium gas by utilizing the ideal gas law. So we're gonna say here pressure of helium equals moles of helium times r times t divided by v. We don't even need to look at the grams of oxygen because the question again is only asking about the partial pressure of helium. Alright. So let's take the 12 grams of helium, we look on the periodic table, you'll see that the atomic mass of helium is approximately 4.003 grams helium for every 1 mole of helium. Grams here cancel out and I'll have my moles as 2.998 moles of helium. So take that, plug it into the formula. So 2.998 moles of helium. Multiply by my gas law or my gas constant, 0.08206 liters times atmospheres over moles times k. Remember temperature must be in Kelvin. So the 30 degrees Celsius, I'm gonna add 273.15 to it, and that gives me 303.15 Kelvin. Then we take the volume, which is 5 liters, and we just plug it in. Look at the units. Liters cancel out with liters, kelvins cancel out with kelvins, moles cancel out with moles. And at the end what we have left is atmospheres. So we plug that in and we'll get 14.9149 atmospheres. If we look at the sig figs within our question, we have 3 sig figs, 3 sig figs, 2 sig figs, and 1 sig fig. Here if we wanted 1 sig fig it would round down to 10 atmospheres. Again that's such a big deviation from our actual number. So let's go with a number that makes more sense because we don't wanna round so much. We're gonna say our answer here is 14 0.9 atmospheres. Again we're constantly trying to remember significant figures play a role in a lot of our questions. Doesn't quite fit. Going from 14.9 to 10 is such a that doesn't quite fit. Going from 14.9 to 10 is such a big difference. So here, we're just gonna go with 3 sig figs. 14.9 atmospheres is more reasonable. It's not a big deviation from our original answer. So just remember, if we have the moles, the temperature and the volume, we can find the partial pressure of a gas by using the ideal gas law formula.
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concept
Partial Pressure
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Under Dalton's law, we can use the mole fraction of a gas to find the partial pressure of that gas here, we're gonna say, Dalton's law says that the partial pressure of a gas one equals the mole fraction of that gas, one times the total pressure. So if you don't quite remember how to calculate mole Fraction, I suggest you go back and take a look at my topic videos on mole fractions before proceeding further under Dalton's law.
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example
Partial Pressure Example 3
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Here in this example question it says, a container has 16.7 grams, o 2, 8.1 grams, h 2, and 35.2 grams, n 2, and contains a total pressure of 0.83 atmospheres. We're asked to calculate the mole fraction of o 2 and its partial pressure. Alright. To determine the partial pressure of o 2, so pressure of o 2 equals the mole fraction of o 2 times the pressure total. Right now we already know what the total pressure is, it is 0.83 atmospheres. So to determine partial pressure we first have to find the mole fraction of 0 2. Remember mole fraction of o 2 will equal the moles of o 2 divided by total moles of all the together. So we're gonna take here, we're gonna say we have 16.7 grams, o 2, 8.1 grams, h 2, and 35.2 grams, n 2. We're gonna convert each one of these grams into moles. So we look on the periodic table for the atomic masses of oxygen, hydrogen, and nitrogen. Here, 1 mole of o 2, 2 oxygens comes out to 32 grams, 1 mole of h 2, 2 hydrogens when you add up their atomic masses is 2 point 016 grams, and then you have 2 nitrogens, so 1 mole of n 2 is 28.02 grams. Here all of our grams cancel out and we'll have the moles for each one of these gases or each one of these, yeah, each one of these gases. So that's gonna come out to 0.5219 moles of 024.0179 molesh2 and 1.2562 moles of n 2. So take those and plug them in. So we have 0.5 219 moles of 2. On the bottom we have the collective moles of everyone. So here we're just adding them all together. K. And then when you work that out, you'll get your mole fraction for 02, which comes out to be 0.0900. Take that mole fraction and plug it in here. So here's our mole fraction of o 2, remember mole fraction is a unitless number. So then 0.0900 times 0.83 atmospheres comes out to 0.075 atmospheres. Here our answer has 2 significant figures because our lowest number of sig figs are 8 8.1 and 0.83. They both have 2 sig figs. So just remember, utilizing Dalton's law, we can use the mole fraction of any gas times the pressure toll to find the partial pressure of that particular gas.
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Problem
Problem
A sample of 3.51 g argon and an unknown amount of oxygen are mixed in a container at room temperature. The partial pressure of argon was calculated as 71.0 torr and the partial pressure of oxygen as 188 torr. What is the mass of the oxygen within the container?
A
4.27 g
B
6.18 g
C
7.44 g
D
9.16 g
E
15.2 g
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Problem
Problem
A gas mixture contains 72.8% chlorine and 27.2% neon by mass. What is the partial pressure of neon in the mixture if the total pressure is recorded as 809 mmHg?