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APPLIED FIZZICS · Field Notes · Edition No. 6
A whimsical look at the universe through a glass of Champagne...or your favorite cocktail.
August 4, 2026 · Seattle, Washington
Dear Reader,
This week's Field Note answers one of the most practical questions in carbonation science: How do you keep carbonated beverages fizzy after opening them? As it turns out, the same physics also explains how beverages are force carbonated.
First, let's look at last week's quiz.
Last Week's Quiz
Question: Two identical bottles of Champagne are opened, Bottle A and Bottle B. From Bottle A, two glasses are poured at once, and the bottle is stoppered and returned to the fridge. From Bottle B, one glass is poured the first day, and the bottle is stoppered and refrigerated. Then next day, a second glass is poured, and then stoppered and refrigerated again. Which bottle is fizzier after two days?
The correct answer is Bottle A, from which two glasses were poured at once.
Each time the bottle is opened, the pressurized CO₂ in the headspace escapes. After the bottle is recapped, additional CO₂ must leave solution to rebuild that pressure, so the twice-opened bottle ends up with less dissolved CO₂.
We can quantify this. Assume the bottle holds 25 ounces of Champagne, and each glass holds 5 ounces. Bottle A has only one equilibration event. After two glasses are poured, it must replenish 10 ounces of headspace with CO₂ from the remaining 15 ounces of Champagne.
Bottle B has two equilibration events. After the first pour, the bottle must replenish 5 ounces of headspace, and after the second glass another 10 ounces—for a total of 15 ounces. So bottle B loses 50% more headspace gas from the multiple pourings.
You can check this with our Bottle Headspace Pressure calculator. Bottle A drops from an initial pressure of 65 psi to 44.3 psi; Bottle B drops from 65 psi to 54.4 psi on the first pour, to 37.1 psi on the second. That's about a 35% greater loss of CO₂. with the double pour.
The physics is simple: fewer openings mean less CO₂ lost.
Last week, we looked at what happens when different species of gas molecules occupy the same space. By treating each molecule individually, we saw that only CO₂ molecules can influence the pressure of CO₂. That simple idea—Dalton's Law of Partial Pressures—explains why only CO₂ can keep CO₂ in solution.
Try the simulator below. Click the "+" buttons to add molecules, then click Open Door to let the gases mix.
Want a larger view? Open the full-screen simulator.
But we quietly skipped over something rather important:
Champagne isn't a gas. Most of its CO₂ is dissolved in a liquid. So how do gas molecules above the wine interact with CO₂ molecules dissolved below the surface?
The answer turns out to be remarkably similar to our two-chamber Dalton's Law demonstration.
You can't "keep" CO₂ molecules in solution with any gas in the headspace except CO₂. In fact, even high-pressure CO2 in the headspace doesn't "keep" CO₂ molecules in solution. CO₂ molecules in solution will still randomly escape. But if the partial pressure of CO₂ in the headspace is sufficient, the rate at which CO₂ molecules escape will be balanced, on average, by the rate at which other CO₂ molecules dive back into the liquid.
CO₂ pressure doesn't hold CO₂ molecules in solution. It balances the traffic.
And that's something no other gas can do. Even 1000 psi of argon in the headspace won't reduce the rate at which CO₂ molecules leave the liquid by one little bit.
And that's exactly how the Perlage System works. Every time you open a bottle of Champagne, the pressurized CO₂ in the headspace escapes. Perlage simply replaces that lost CO₂, restoring the headspace to its original pressure of about 65 psi at serving temperature. Once the original CO₂ pressure is restored, the same equilibrium between dissolved CO₂ in the Champagne and gaseous CO₂ in the headspace is restored as well.
In a very real sense, when repressurized with Perlage, the bottle doesn't "know" it's been opened.
Physics Corner
A short aside that goes one layer deeper on the science.
Last week, our molecular simulator consisted of two chambers filled with different gases. This week, we made one simple change.
Instead of filling both chambers with gas, we filled the lower half of the chamber with liquid, while the upper half became the bottle's headspace. The vertical dividing wall became the horizontal liquid-gas interface.
The physics is surprisingly simple. Molecules in the liquid are allowed to escape into the headspace if they have enough energy. Molecules in the headspace, meanwhile, have a chance to dissolve back into the liquid whenever they collide with the surface at a low enough speed for the weak intermolecular (van der Waals) forces at the liquid surface to draw them into the liquid.
Each molecule follows only these simple rules. There is no equation telling the simulator where equilibrium should occur. The simulator simply applies the same microscopic rules to every molecular interaction. Equilibrium emerges naturally from the combined effect of billions of those individual interactions.
Watch this short video to see what is happening at the molecular level at equilibrium.
Video of Applied Fizzics Molecular Simulator in action
You can experiment with this yourself at the Henry's Law Solubility Simulator.
Watching these microscopic events unfold helps explain why Henry's Law isn't really a "law" that molecules obey. Instead, it is the inevitable statistical behavior that emerges from enormous numbers of randomly moving molecules.
This Week's Quiz
Question: You vigorously shake a can of beer and hand it to your friend as a joke. The beer spews all over him when he opens it. Which statement is most accurate?
- A. Shaking releases more CO₂ from solution inside the can, increasing the pressure and causing it to spew when opened.
- B. The pressure in the can stays the same, but shaking creates countless tiny CO₂ bubbles throughout the liquid that act as nucleation sites when the can is opened.
- C. You're a real jerk.
Bonus Question: Will tapping on the top of the can with your fingernail prevent spewing, and if so, why?
Make your prediction, then reveal the answer below.
Reveal the Answer
The correct answer is B (and C...you're kind of a jerk).
Shaking the can does not release CO₂ and does not increase the pressure in the can. The mechanical agitation of shaking simply creates enormous numbers of little bubbles in the body of the liquid.
When the top is popped and the headspace CO₂ rushes out, the liquid is no longer held under pressure. Then, each of those tiny bubbles suddenly expands as the pressure around it drops, just as the ideal gas law predicts. This swells the whole volume of the liquid and causes it to shoot out the top.
Further, each of those bubbles is a nucleation site, and CO₂ diffuses into them and makes them even bigger as they are on their way out of the can.
As for whether tapping the can will prevent spewing, stay tuned. We're going to do a "Myth Busters"-style video on this next week, to see if this urban legend is true.
Cheers,
Evan Wallace
President, Applied Fizzics Inc.
Makers of The Perlage System®
Want more Champagne science, carbonation experiments, and behind-the-scenes product stories? Visit the Applied Fizzics Field Notes archive .
