Mixing It Up: Understanding Partial Pressures and Dalton’s Law (PLA 12)

Ben Lear

Alright, welcome back to The Honors Element, everyone. I’m Ben Lear, and as always I’m joined by Morgan Vincent. Today we’re getting into the world of gases, partial pressures and Dalton’s law. Basically, what happens when you mix more than one gas together in a container and want to understand how each one is behaving, and, well, why we care in the first place.

Morgan Vincent

Yep, and honestly this is one of those topics that sounds a little abstract at first, but it’s everywhere: soda cans, scuba diving, even the air we’re breathing. So let's start with the basics. Ben, what do we actually mean by “partial pressure”?

Ben Lear

Partial pressure is simply the pressure that a single gas in a mixture would exert if it were the only gas in the entire container. So if you picture a jar filled with oxygen, nitrogen, and I don't know, maybe a splash of carbon dioxide, each gas acts as if the others aren’t even there. And that’s the beauty of how gases behave. Thanks to kinetic molecular theory, they’re zipping around and barely interacting with each other. Each one only cares about bouncing off the walls, not the other particles.

Morgan Vincent

And that’s a big deal. If gas particles stuck together or kept influencing each other, Dalton’s law would fall apart. I like to picture it as a silent disco: everyone’s on the same dance floor, but each person has noise-cancelling headphones and their own music. No one cares what the others are doing, and yet the room is still full of energy. The key idea is that in a gas mixture, each type of particle contributes its own pressure, and the total pressure is just the sum of all those independent contributions.

Ben Lear

Exactly. Sometimes students ask, does temperature or volume change things? The answer is, each gas in a mix shares the same temperature and volume, but contributes its own pressure based on the number of moles present in that system. So, if you’re looking at air in this room, the nitrogen, the oxygen, that tiny bit of carbon dioxide, all of those are acting independently as far as pressure goes. We call their contributions partial pressures.

Morgan Vincent

And this is where Dalton’s law comes in, which we’ll get to in a second. But first, without this idea of partial pressures, it’d be nearly impossible to make sense of what’s happening in a gas mixture.

Ben Lear

So now that we’ve got partial pressures in mind, here’s the punchline: Dalton’s law of partial pressures. Basically, Dalton figured out experimentally that the total pressure in a mixture is just the sum of all those individual partial pressures. So, P total equals P one plus P two plus P three, you get the idea. It’s deceptively simple, but hugely useful.

Morgan Vincent

Yeah, and you can write it out using as many you pressures as you need! If you’ve got more gases, just keep adding. And what’s nice is that it works as long as we’re dealing with gases at low or moderate pressure, so, basically, almost every intro chemistry problem you’ll see. If you’ve got, say, a mixture of hydrogen at 2.3 atmospheres and nitrogen at 0.7 atmospheres, your total is straightforward: 3 atmospheres. Done.

Ben Lear

But let’s take it a step further, because sometimes they’ll ask for the partial pressure based on something called “mole fraction.” So, mole fraction is just a fancy way to say, “What fraction of all molecules in there is from a certain gas?” It’s the number of moles of your gas divided by the total number of moles in the mix. You times that mole fraction by the total pressure, and boom. That’s your partial pressure for that gas.

Morgan Vincent

For real chemistry problems, that’s super helpful. If you know the mole fraction, let’s call it x sub A for gas A. The partial pressure is x sub A times the total pressure. So if nitrogen is 78% of air, its mole fraction is 0.78, and its partial pressure is 0.78 times that atmospheric pressure. There’s a great real-world case with nitrogen dioxide and dinitrogen tetroxide, actually. Did you want to walk through that, Ben?

Ben Lear

Yeah! So, if you have a flask with a bunch of NO₂ at high temperature and then you cool it down. The particles won’t always stay as single molecules—some of them start pairing up into N₂O₄. That means you suddenly have fewer individual particles flying around and colliding with the walls of the container, so the total pressure you measure goes down. And what’s important to see is that Dalton’s law still works in this situation. Even though some molecules are flying solo as NO₂ and others have paired up as N₂O₄, it’s really just the same particles swapping identity. Pressure-wise, each form still acts independently, and the total pressure is simply the sum of their contributions. So honestly, the more you practice thinking through the reactions and the subsequent math, the more it’ll just click.

Morgan Vincent

Agreed! But now for the fun part. Let’s bring this into everyday life and the lab. One of the classic chemistry set-ups is collecting a gas over water. You spout gas from a reaction through water into an upside-down tube, that’s the pneumatic trough. But here’s the thing: the gas you collect is mixed with water vapor, so the total pressure you read includes both the target gas and the water. To find out how much pressure is just your gas, you use Dalton’s law to subtract the water vapor pressure from the total, and you’ve got your answer.

Ben Lear

Yeah, that’s one of those details students often want to skip, but it’s crucial. Chemists will check tables for water vapor pressure at a certain temperature, because that value changes if it’s, say, 20°C versus 40°C. If you collect oxygen at 754 torr and the vapor pressure of water is 20 torr, it’s just 754 minus 20, and that gives you the partial pressure of the oxygen. Simple, but, you know, sort of easy to forget if you’re in a rush.

Morgan Vincent

And this is also what’s going on when you open a can of soda, remember because I love food chemistry! All those bubbles? That’s CO₂ gas coming out of solution, because it was dissolved in the liquid under high pressure. The second you open the can, you release the pressure, and, poof, CO₂ comes out—making fizz and foam. That pressure was set so high in that sealed can and as soon as you break the seal, Dalton’s law kicks in and the partial pressure of CO₂ drops, letting the dissolved gas escape.

Ben Lear

Honestly, it’s the same thing if you’ve ever opened a warm soda and it explodes everywhere. That’s because CO₂ is less soluble as temperature increases, so more is ready to come out. Fun, or, well, annoying at times, but it’s a perfect, messy demonstration of partial pressures in your fridge.

Morgan Vincent

Right, and beer, sparkling water, champagne—it’s all about controlling and understanding gas pressures. So, whether you’re in the lab measuring gases over water or just trying not to spill your soda, Dalton’s law has you covered.

Ben Lear

Alright, that’s gonna wrap it up for today! Keep looking for these ideas in the world around you, because chemistry is kinda hiding everywhere if you know where to look. Morgan, thanks for the great discussion as always!

Morgan Vincent

Thanks, Ben—always a pleasure! And thanks to everyone for tuning in. We’ll catch you all on the next episode of The Honors Element. Take care!