The explanation for the boiling point of water is that at 100C, the vapour pressure becomes greater than atmospheric pressure. But say you had a a jar of water sealed in argon at 1atm, which is larger than the vapour pressure of water, wouldn't some water evaporate? I'm guessing yes, because there are no water molecules in the argon initially, so wouldn't some randomly leave the liquid, making the partial pressure on average non-zero?
Another reason I can think of it being based on partial pressure is that an open jar of water in a windless dark, dry room will still eventually evaporate. So why is it that at 100C there is a sudden dramatic change? Isn't the vapour pressure at e.g. 95C more than the partial pressure of water in the atmosphere as well?
Best Answer
Yes, but there will not be boiling, only evaporation from the surface.
Yes.
At this temperature the vapor pressure inside the bubble begins to be high enough to overcome the pressure in the liquid, so the bubble survives and rises.
Yes, and the water will evaporate to the atmosphere, but only on the surface. There will be no boiling, since the pressure in the liquid is higher than the pressure in the bubble could be, so any small bubble is crushed and the vapor turns back into liquid.