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I understand that all solids and liquids are unstable to vaporisation in vacuum (because some molecules will always have the energy to escape the bonds due to the nature of thermal energy distribution).

It is said that solids and liquids reach an equilibrium at any non-zero pressure when the number of molecules leaving it become equal to the number of molecules assimilating.

My small brain thinks that this instability should be unrelated to the pressure and that such an equilibrium should not, in general, be possible. At best, the instability should be related only to the partial pressure of the vapour of the solid/liquid substance around it.

What I mean to say is that even under very large air pressure, some molecules of my table would still have the energy to leave it while none of the molecules of the air would assimilate into the table.

In other words, any liquid should boil as easily in air as it does in vacuum (unless it is liquefied nitrogen or oxygen).

Why don't all liquids kept in open just boil all the time (just as they would in vacuum)? Why is the behaviour different in vacuum and in open? Isn't keeping a liquid in open same as keeping it in vacuum since the partial pressure of the liquid in open should approach zero because the vapour would drift away quickly?

Where am I going wrong in my thinking?

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    $\begingroup$ What you're saying is basically correct. If the partial pressure of the substance in the surrounding air is equal to its equilibrium vapor pressure, it will stop evaporating. If it is less than the equilibrium vapor pressure, it will continue evaporating until its partial pressure reaches the equilibrium vapor pressure. $\endgroup$ – Chet Miller Dec 7 '17 at 13:03
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    $\begingroup$ Further to @ChesterMiller 's point, equilibrium vapor pressures of most elements at STP are really really really small, so things can 'evaporate' for billions of years without you noticing. Water is one of the obvious counter-examples, where water vapor is a quite noticeable component of the atmosphere, often not in equilibrium. $\endgroup$ – Jon Custer Dec 7 '17 at 13:45
  • $\begingroup$ @ChesterMiller If so, why don't all liquids kept in open just boil all the time (just as they would in vacuum)? Why is the behaviour different in vacuum and in open? Isn't keeping a liquid in open same as keeping it in vacuum since the partial pressure of the liquid in open should approach zero because the vapor would get drifted away quickly? $\endgroup$ – Ritesh Singh Dec 7 '17 at 17:59
  • $\begingroup$ @JonCuster I see, thanks for confirming that! There is something else that makes me wonder. I have edited my question to highlight what intrigued me more. $\endgroup$ – Ritesh Singh Dec 7 '17 at 18:04
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    $\begingroup$ To answer your questions, No. For the liquid to boil, its equilibrium vapor pressure must be equal to the total pressure of the gases in the system, and not just equal to the partial pressure of the substance in the gas phase. Bubbles can only form below the liquid surface if the pressure within the bubbles is high enough to push back the surrounding atmosphere. $\endgroup$ – Chet Miller Dec 7 '17 at 18:17
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From the comments section, it seems like you're asking why—if condensed matter such as liquids and solids tend to evaporate or sublimate into the gaseous state, which they do—why don't bubbles of the gaseous state spontaneously form inside their bulk.

There are several interrelated impediments. One is the extreme unlikelihood of so many molecules entering the gaseous state at once. A visible bubble would require many more than a billion billion molecules breaking their bonds simultaneously at roughly the same location. This isn't so unusual at the boiling point, but remember than the vapor pressure (which corresponds to the likelihood of bonds breaking) is exponentially dependent on temperature. At room temperature, you will simply never see water spontaneously form a vapor bubble, for example. Any bubble that you do see forming consists of dissolved atmospheric gases coming out of solution.

Another impediment is the work that has to be done against the atmosphere to form a vapor bubble inside the material. The molar volume of the gas phase might be a thousand times larger than that of the condensed-matter phase, for example. The resulting movement of the interface therefore has to push against the surrounding atmosphere, representing mechanical work.

Still another impediment has to do with the energy penalty associated with a new surface. It costs energy to form a surface because a surface represents potential bonds that couldn't be made. This phenomenon is the origin of surface tension. It's also the reason why we have to "superheat" materials (i.e., heat them slightly above their boiling point) to actually obtain vapor bubbles—because there's an energy barrier to forming that vapor-condensed matter interface that has to be surmounted. This factor also explains why solids spontaneously evaporate into the gaseous phase but don't spontaneously form pools of liquid on their surfaces. Liquids have a positive surface tension; gases don't.

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