0
$\begingroup$

This may be a silly question, but when I was reading my chemistry textbook, there was this following statement:

"A liquid boils when its vapor pressure equals the external pressure acting on the liquid surface. At this point, bubbles of vapor form within the liquid. The temperature at which a given liquid boils increases with increasing external pressure."

But what would happen if that liquid was sealed in a vacuum, or at least submitted to very low pressure? Would the boiling point be a lot lower?

$\endgroup$
1
  • 1
    $\begingroup$ Yes, with lower pressure the boiling point lowers too. // Think about it. Water boils at a lower temperature in the mountains than it does at sea level because there is less air pressure in the mountains. $\endgroup$ – MaxW Mar 27 '20 at 18:52
6
$\begingroup$

"Sealed in a vacuum" is an oxymoron, a contradiction in terms, much like "frozen with fire". If you seal a liquid in a flask containing nothing else but vacuum, then a part of the liquid will quickly evaporate and fill the flask with vapor, so it would no longer be a vacuum.

The said vapor will exert some pressure, depending on the temperature and the nature of the liquid. For example, with water at room temperature, it will be a lot lower than the normal atmospheric pressure. Indeed, the boiling point of water at that pressure will be a lot lower than the usual $100\;^\circ\rm C$. Guess what? It will be the same as ambient temperature. That's how equilibrium works.

Now what if we just keep the liquid at vacuum by pumping out the vapor as soon as it forms? Well, the simplified answer is: the liquid will evaporate gradually, bit by bit, until nothing is left. There will be no more liquid, and consequently no more boiling point. That's what happens with the boiling point in a vacuum.


† In fact, it is more complicated. As the liquid evaporates, the rest of it gets progressively colder. At some point, it will get cold enough to freeze, and then you'll be dealing with ice. The molecules will still be leaving its surface, but slower and slower, and the complete sublimation might well take forever, or even longer. That's why the space rocks are still there after all these years.

$\endgroup$
1
$\begingroup$

Generally speaking, liquid boils at the temperature, at which its saturated vapour pressure is equal the external pressure.

With external pressure going down, boiling temperature goes down as well.

When external pressure goes down too much, the boiling point may meat the freezing point of the liquid and the liquid freezes. There is no boiling any more, but sublimation. (That is the way of "dry ice" production from liquified carbon dioxide.)

For very low vapour and external pressures, there may be needed little liquid overheating, as the vapour pressure must deal with external pressure AND surface tension, acting against bubble forming.

$\endgroup$

Not the answer you're looking for? Browse other questions tagged or ask your own question.