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Q & A: Boiling with ice

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Most recent answer: 10/22/2007
Q:
My teacher recently conducted an expieriment that has me a little puzzled. He heated a flask with water inside about half full until the water was at a hard boil. He then removed the heat which allowed the boil to stop and then put a rubber stopper on top of the flask. After that, he took a peice of ice and rubbed it around the top of the flask causing the water inside the flask to begin to boil. Ofcoarse, this boil was no where near the intensity of the boil caused by the heat, but still to my amazment, it worked! I am so confused, could you please explain to me how this occured?
- Shawn (age 14)
Pace High School, Pace, FL, USA
A:
Hi Shawn,

That's a really cute demonstration! And, as you say, it can be quite puzzling. It's a pity your teacher left his students puzzled about it rather than explaining it fully. I can give you my guess of what went on in that experiment.

When the water is boiling with the stopper out, liquid water is being converted to its vapor phase, steam, which fills up the top half of the flask. If the flask is then taken off of the heater (and then you have to wait for it to stop boiling otherwise the stopper will be blown out the top) and stoppered, then water at a temperature just under the boiling point is in the flask with water vapor above it, all at approximately 100 degrees Celsius. Rubbing ice around the top of the flask cools the vapor down and condenses it back into liquid water on the sides of the flask. Doing this reduces the pressure of the water vapor over the liquid, creating a partial vacuum.

The vapor pressure of water (that is, the pressure of water vapor which is in equilibrium with liquid water) depends quite a lot on the temperature. For water at 100 degrees celsius, it is quite high (about an atmosphere). If the vapor pressure of the water is higher than the partial pressure of water vapor over the liquid, then the liquid will evaporate until it comes into equilibrium. If the vapor pressure is higher than the total pressure of all the gas over the liquid, then the water will boil, because the evaporation may spontaneously happen anywhere in the liquid. You actually need the vapor pressure to be higher than the pressure in the water at the depth at which the bubbles form in this case.

You can even get water to boil at room temperature by putting it into a sealed glass jar and pumping the air out with a vacuum pump. Water takes 540 calories per gram to evaporate (regardless of the pressure), and this evaporation cools the liquid water down. In a demonstration I saw once with water in a vacuum jar, the water boiled until it froze.

Tom J.

(published on 10/22/2007)

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