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APPLIED FIZZICS Field Notes – Edition #5
Looking at the universe through a glass of Champagne.
Dear Reader,
This is an important edition of Field Notes. In this edition, we're going to discover one of the most important principles in all of physics—and it turns out to explain far more than just Champagne.
But first, let's look at the quiz from Field Notes #4.
Last Week's Quiz
Last time, we asked about putting two freshly opened Champagne bottles back in the fridge, one completely full, and the other about half full.
The question was, the next day, which one is fizzier? That is, which one has the higher concentration of dissolved CO₂?
If you guessed the full bottle, you are correct.
You remember from last time that nearly all of the CO₂ loss in a carbonated beverage occurs from diffusion at the liquid/air interface. Because the wine level is up into the narrow neck of the bottle, the full bottle has a much smaller Champagne/air interface and loses CO₂ more slowly.
Not only that, the full bottle has a larger "reservoir" of dissolved CO₂, so the full bottle has more CO₂ to give up.
Then we asked a second question: What if you capped both bottles with airtight stoppers? Which one would be fizzier?
Even with airtight stoppers, the full bottle still wins.
When you cap the bottles, CO₂ will continue to diffuse out of solution until a new equilibrium is reached. Remember, equilibrium is the condition at which as many molecules of CO₂ leave solution as enter solution.
In the full bottle, the headspace is tiny, and the reservoir of dissolved CO₂ is enormous—there is an entire bottle of Champagne available to supply the headspace.
In the half-empty bottle, the headspace is much larger. Much more CO₂ will have to evolve out of solution to reach equilibrium, at the expense of dissolved CO₂. So the half-empty bottle will be flatter.
If you open and immediately restopper a full bottle, it will last a remarkably long time with almost undetectable loss of fizziness. The equilibrium pressure reached in the headspace will be very close to the pressure that existed before the cork was popped.
But the emptier the bottle is, the less effective the stopper is.
Our calculations show something rather surprising: if the bottle is half-empty when you recap it, it will contain about 57% less dissolved CO₂ than the full bottle when it comes to equilibrium (which generally takes a few hours).
Curious how much carbonation you're losing from your own bottle? We built a Headspace CO₂ Pressure Calculator that lets you explore exactly how much CO₂ you lose each time you pour another glass.
So every time you open a bottle, pour a glass, and put the stopper back on, the bottle settles into a new equilibrium.
But there is something subtle happening here.
The CO₂ in the headspace isn't magically keeping the CO₂ in the Champagne dissolved. In fact, every molecule of CO₂ in the headspace came from the Champagne itself.
The less Champagne that remains in the bottle, the larger the headspace, and the more CO₂ molecules it takes to fill the void. Those molecules have to come from somewhere—and the only place they can come from is the dissolved CO₂ in the wine, which is exactly where you wanted them to stay.
By the time only a glass or two remains, the stopper is doing very little to preserve the remaining carbonation.
So here's the obvious question.
If there's already a cushion of CO₂ above the Champagne, why doesn't that cushion keep the rest of the CO₂ dissolved?
The answer lies in one of the most important principles in all of science.
More than 200 years ago, the English chemist and physicist John Dalton discovered that when different gases occupy the same space, each gas behaves almost as though the others were not there. Today, we call this Dalton's Law of Partial Pressures.
That simple idea explains not only sparkling wine, but also why nitrogen gives Guinness stout its creamy head, why fish can breathe underwater, why scuba divers have to worry about the bends, and countless other phenomena.
Physics Corner
A short aside that goes one layer deeper on the science.
To make Dalton's Law visible, we built a molecular simulator that lets us watch individual gas molecules move, collide, and come to equilibrium. In physics, this molecular approach is known as kinetic theory.
What the simulator shows is that at the molecular level, the world doesn't behave according to "bulk" properties like density, pressure, and temperature. Rather, we see that bulk properties like density, pressure, and temperature are consequences of the collective action of many, many molecules.
When we look at gases this way, many seemingly complicated phenomena become surprisingly easy to understand.
Watch our molecular simulator in action below to understand why only CO₂ can keep CO₂ in solution.
Video of Applied Fizzics Molecular Simulator in action
Want to experiment with the simulator yourself? Try the Applied Fizzics Molecular Simulator.
In the simulator, I used only 40 molecules. In the real world, there are roughly 25 quintillion molecules in every cubic centimeter (that's 25 billion billion). Yet each molecule is so tiny that, on average, it travels several hundred molecular sizes before colliding with another.
If molecules were the size of baseballs, the spacing between collisions would be like baseball players throwing the ball around on the diamond.
This enormous amount of empty space is the key to understanding Dalton's Law. Although gas molecules collide with one another constantly, they spend most of their time flying through empty space. As a result, each type of gas spreads through the available volume almost as though the other gases were not there. Carbon dioxide responds only to the partial pressure of carbon dioxide—not to the combined pressure of all the gases present. That's why only CO₂ can keep CO₂ dissolved.
As always, we finish with a quiz.
This Week's Quiz
Suppose you have two identical freshly opened bottles of Champagne. A bottle of Champagne has about 25 ounces of wine in it.
From the first bottle, Bottle A, you pour two 5-ounce glasses at once, for 10 ounces total, then immediately recap it and put it back in the fridge.
From the second bottle, Bottle B, you pour one 5-ounce glass, recap it, and put it back in the fridge. The next day, you pour a second 5-ounce glass and recap it again, then put it back in the fridge.
Assume both bottles are stored under identical conditions.
Which bottle will be fizzier (i.e., have more dissolved CO₂ left) after the second pour?
- Bottle A from which two glasses were poured at once.
- Bottle B from which one glass was poured on two separate occasions.
- Both bottles will contain the same amount of dissolved CO₂.
We'll reveal the answer next time. In the meantime, you can test your intuition with our Headspace CO₂ Pressure Calculator and see if your prediction is correct.
Reveal the Answer
The correct answer is 1. Bottle A, from which two glasses were poured at once.
Opening a bottle twice vents the pressurized headspace twice. Each time, additional CO₂ must leave solution to re-establish equilibrium, so the twice-opened bottle ends up with less dissolved CO₂.
Another way to look at it is this: Bottle A has 10 ounces of headspace volume to replenish with CO₂ while coming to equilibrium. Bottle B has to replenish 5 ounces of headspace with CO₂ after the first pour, and 10 ounces after the second pour, for a total of 15 ounces, or 50% more.
Even though both bottles dispense the same amount of Champagne, minimizing the number of times you open the bottle preserves more carbonation.
