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Ideal Gas Law Multiple Gases Partial Pressures using Mole Fractions (Example 4)
You may have a mixture of gases together, which will also have multiple chemical species present.
It will have a total pressure, but that total pressure will be the sum of all the partial pressures (this is called Dalton’s Law).
You can find the partial pressures by finding the mole fraction of the species and multiplying by the total pressure.
Here’s a practice example we created:
Example problem: A gas mixture contains 20g of O\(_2\) and 30 g of CO\(_2\). What is the partial pressure of CO\(_2\) in the gas mixture?
Convert the masses into moles by dividing by each species’s molar mass.
You should keep all decimals on your calculator, for our guide we will just round so it doesn’t look so messy.
\[\frac{{30g}}{{44g/mol}} = 0.682mol\]
\[\frac{{20g}}{{32g/mol}} = 0.625mol\]
Find the total amount of moles:
\[0.682mol + 0.625mol = 1.307mol\]
Then find the mole fraction of CO\(_2\) by dividing the moles of CO\(_2\) by the total moles in the mixture:
\[{x_{C{O_2}}} = \frac{{0.628mol}}{{1.307mol}} = 0.522\]
Remember, the mole fraction is unitless.
Don’t use the mass fraction by mistake! (The mass fraction is the mass of each species divided by the total mass)
The partial pressure is the mole fraction of CO\(_2\) multiplied by the total pressure.
\[{P_{C{O_2}}} = (0.522)(3450kPa)\]
\[{P_{C{O_2}}} = 1800kPa\]
Try another Ideal Gas Law problem with combining multiple gases together!