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The temperature at which water boils depends on pressure. You can demonstrate this by dramatically lowering the pressure on a water-filled plastic syringe at room temperature.
Fill a syringe a quarter full of water. Try to keep out as much air as possible.
Hold the syringe in one hand and cover the tip with your finger. With your other hand, pull on the plunger.
Notice that, as you pull on the plunger, it pulls back on you. Notice also that an empty space—a bubble of air—appears inside the syringe.
Allow the plunger to slide slowly back into the syringe. Look again to see if there’s still an air bubble in the syringe.
Keep your finger on the tip and pull out the plunger again. Release the plunger suddenly. Notice that it snaps back quickly.
Pull on the plunger a third time. Notice that, this time, bubbles form inside the water, and the water appears to be boiling.
Not only does the water in your syringe appear to be boiling, it is boiling.
Living as we do at typical atmospheric pressures, we tend to think that water has to be hot to boil. But the transition from liquid to gas can occur not just as the result of increased temperature, but also as the result of decreased pressure.
Pulling on the plunger reduces the pressure on the gases inside the syringe by increasing the volume—a relationship given by Boyle’s Law: For a gas in an enclosed space at a constant temperature, volume and pressure vary inversely. In other words, doubling the volume halves the pressure.
Tap water has air dissolved in it. When you reduce the pressure in the syringe by pulling out the plunger, the dissolved air comes out of solution and forms an air pocket at the tip of the syringe. When you slowly allow the plunger to slide back in, the air that has come out of solution stays out of solution. That’s why there may seem to be more air in the syringe than when you started.
But something else happens when you pull back on the syringe: Under reduced pressure, water changes from liquid to gas. Bubbles of water vapor form inside the liquid, and the water in the syringe boils at room temperature.
In a clean liquid, it’s not easy for small bubbles to form, or nucleate. However, when you pull out the plunger and allow it to snap back in, you create many tiny “seed” bubbles throughout the water. The next time you pull back the plunger, boiling happens more easily, thanks to the nucleation sites provided by these seed bubbles.
One way to understand the importance of seed bubbles is to remember the last time you tried to blow up a balloon. At first, a balloon can be difficult to get started. But once started, it’s easier to inflate. The same rule applies to bubbles of water vapor forming in water. If you reduce the pressure and temperature even further, either by using a vacuum pump or by traveling to the surface of Mars, it’s possible to have water both boil and freeze at the same time. At this combination of low temperature and pressure—0.01 °C and 0.006 atmospheres, also known as the triple point of water—all three phases of water can exist at the same time.
Try this experiment with carbonated water.
Carbonated water has carbon dioxide dissolved in it. When it’s in an unopened bottle, this carbon dioxide is under pressure and remains in solution. When you open the bottle and lower the pressure, the carbon dioxide comes out of solution, forming the bubbles you see when you pour it into a glass. Sometimes you can see the bubbles rise in long, vertical lines, originating from a scratch or pit in the wall or bottom of the glass.
With carbonated water in the syringe, the carbon dioxide comes out of solution rapidly when you pull back on the plunger, and even more bubbles form in the water. By tapping the sides of the syringe, you can force all the bubbles together into one larger bubble.
Take the pressure off pickling.
Discover the relationship between temperature and volume of a given amount of gas.
Make a reverse Cartesian diver.
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Attribution: Exploratorium Teacher Institute