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?