The Dynamics of Equilibrium (PLA 17)

Ben Lear

Alright, welcome back to The Honors Element! I'm Ben, here with Morgan—how's it going?

Morgan Vincent

Hey, Ben! I’m good! I'm ready as ever to get into this week’s chemistry adventure.

Ben Lear

Yeah, I think today’s topic is one that’s sort of deceptively simple but super foundational: chemical equilibrium. You know, if you’re like most students, you kinda picture nothing happening when a system’s at equilibrium. Like, it just stops. But it’s actually the total opposite, right?

Morgan Vincent

Exactly! I always tell students, equilibrium is dynamic, not static. It’s not about reactions shutting down. It’s about the forward and reverse reactions happening simultaneously at exactly the same speed. Imagine a busy two-way street: cars are constantly driving both ways, but if the number going east equals the number going west over time, the street looks sort of steady from a distance, even though a ton is happening up close.

Ben Lear

Great analogy! And one of the classic textbook examples that really drives the point home is that neat dinitrogen tetroxide and nitrogen dioxide reaction. Basically, dinitrogen tetroxide is colorless, but when it breaks up into nitrogen dioxide, you get this reddish-brown gas, which is basically the color of smog in LA. If you seal up some dinitrogen tetroxide and heat it, you start to see that brown color appear as nitrogen dioxide forms. But it never goes fully brown because, as some nitrogen dioxide forms, some of it recombines back into dinitrogen tetroxide. They reach this stalemate where both are present, and neither concentration changes anymore. Visually, the color stabilizes at this in-between shade.

Morgan Vincent

That’s such a striking visual. And if you change the temperature or pressure, that balance will shift. But the idea is, at equilibrium, both forward and reverse reactions are still very much happening, just at equal rates. So, no net change in concentrations.

Ben Lear

Yeah, so, big point to remember: equilibrium does not mean go toward a balance until nothing is happening. It’s more like everything is happening equally but in both directions.

Morgan Vincent

Okay, so now that we’ve got this core idea, let’s talk about how you can actually see and measure equilibrium in practice. One of my favorites is the cobalt two chloride system. It’s super dramatic in color changes, so it’s not just some abstract math on paper. If you take cobalt two chloride and dissolve it in pure water, you get this lovely pale pink color from the hexaaqua complex. But, if you dissolve cobalt two chloride in concentrated hydrochloric acid instead, you swap those water molecules out for chloride ions, and boom! The whole thing turns this vivid deep blue. That blue is from the cobalt tetrachloride complex.

Ben Lear

Right, so what’s wild is if you mix somewhere in between, say add some HCl to the pink solution, but not enough to completely convert everything, you get lavender. That’s not because those molecules have turned purple, but because you’ve got a mix: some of the pink hexaaqua species and some blue tetrachloro complex, both hanging out in the same solution. That color is your signal you’ve hit equilibrium between the two forms, at least under those conditions.

Morgan Vincent

And if you approach this from the other direction, it still works. If you start with the blue cobalt tetrachloride solution and start adding water, you shift things back towards pink. Wait long enough, and you’ll hit that same lavender shade. The same equilibrium composition, no matter whether you started with mostly pink or mostly blue. That’s such a key takeaway: equilibrium is independent of the path you took to get there. It’s all about the current conditions: mainly concentrations, temperature, and pressure, not where you started. I know sometimes it feels counterintuitive, but it’s really powerful.

Ben Lear

Yeah, and if you plot the concentrations over time, whether you started from reactants or products, those concentration curves will eventually flatten out at the same points. That’s equilibrium. I always find students are surprised by how identical those end states are, even after very different journeys to get there.

Morgan Vincent

And here’s the thing. The moment you put those cobalt complexes together, they’re already moving toward equilibrium. It’s not like they sit there, think about it, and then flip a switch. The reactions are happening continuously and instantly. And it’s not just when you’re mixing two liquids or gases. This constant push-and-pull toward balance happens between different phases too, like a liquid and its vapor

Ben Lear

Exactly. Think about a sealed water bottle. You’ve got liquid water inside, but if you’ve ever looked closely, there’s also water vapor above the liquid. Molecules are constantly leaving the liquid to become gas, and just as constantly, gas molecules are re-entering the liquid.

Morgan Vincent

If you start with some water and lots of air inside, liquid molecules will evaporate until equilibrium is reached. Different paths, same end point. So it’s not that one side ‘wins’ and the other stops. At equilibrium, both directions are happening, evaporation and condensation, at equal rates. That’s why the amount of liquid water and the humidity in the bottle stay steady.

Ben Lear

So, this gets us to quantifying all of this, right? We’ve danced around it, but to actually predict where this balance point falls, chemists use something called the equilibrium constant. Depending on your system, you’ll see it as K C when using concentrations, or K P for partial pressures. It’s just this ratio, the concentrations of products to reactants, each raised to their stoichiometric coefficients, all measured at equilibrium. It’s called the law of mass action, and, wild enough, no matter what concentrations you start with, as long as you hit equilibrium, that ratio is always the same under a set of conditions.

Morgan Vincent

That is so important. The value of the equilibrium constant is unique for a reaction but will change if you mess with the temperature, or sometimes pressure for gas reactions. If you know K and your starting amounts, you can actually work out exactly what your equilibrated mixture will look like. For example, in food chemistry, when we want to keep food from spoiling, we pay a lot of attention to water activity, which is connected to equilibrium ideas: by lowering water activity, you can shift equilibrium to prevent microbial growth. Managing equilibria is a practical skill way beyond the classroom.

Ben Lear

Oh, totally. Think of making ammonia or sulfuric acid industrially. You’ve gotta manipulate temperature, pressure, all that, to push the equilibrium toward more product. So, understanding how K changes lets you optimize yields, save energy, even make products safer. Honestly, working with equilibrium is one of those places where chemistry feels like you have a superpower for tweaking real systems.

Morgan Vincent

I love that. It’s like you’re never really stuck with a reaction’s natural outcome. You can mess around with the conditions to get what you want. That constant interplay between theory and application is just so chemistry, right?

Ben Lear

Absolutely. And I guess that’s a good place to wrap for today. We’ll get a bit deeper next time, maybe dig into how equilibrium ties into energy and thermodynamics, which is its own rabbit hole. Morgan, always a pleasure.

Morgan Vincent

Thanks, Ben. And thanks to everyone for joining us again—catch The Honors Element next episode, and don’t forget to bring your best chemistry questions. Bye!

Ben Lear

See you next time!