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In a closed system, water vapor would have nowhere to go and atmospheric pressure would not be a factor. It would reach dynamic equilibrium with only water vapor.

But in an open system, I think that the water vapor would be forced back down into the liquid state by the atmospheric pressure until it is heated to the boiling point, at which the water vapor would equal or exceed the atmospheric pressure allowing water vapor particles to escape.

Is my understanding correct?

It just doesn’t seem to be how it works since water vapor can evaporate and escape even before it reaches the boiling point. Or instead, would only fast enough particles be able to overcome and escape the atmospheric pressure and the boiling point would only be when practically all of the water particles have reached that speed

I need help understanding in-depth the process of evaporation and vaporization on a particulate level. Starting from very basic chemistry might help.

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    $\begingroup$ Have you ever washed clothes and let them dry in open air? Have you even observed water evaporating at room temperature? Have you ever observed sea, lakes or rivers boiling or being nearly as hot as boiling water ?( aside of geothermal activities) // Liquid evaporation occurs at any temperature, just its rate changes with temperature. // en.wikipedia.org/wiki/Evaporation // Remember that a particular molecule does not have temperature. $\endgroup$
    – Poutnik
    Jul 13 at 6:16
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    $\begingroup$ In contrary to some other Q/A or forum sites, answers on CH SE site are figuratively paid by the user's own effort. When you ask, you are supposed to search and thoroughly think about the topic and to provide explicit compact summary of partial answers or at least ideas or thoughts you have got until then. Effort not shown may be considered as effort not done and such a question may be closed. $\endgroup$
    – Poutnik
    Jul 13 at 6:19
  • $\begingroup$ Your second explanation is correct. I copy it : *only fast enough particles are able to overcome and escape the atmospheric pressure * $\endgroup$
    – Maurice
    Jul 13 at 8:44
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    $\begingroup$ @Maurice Rather: Only fast enough particles are able to overcome the liquid cohesive forces. $\endgroup$
    – Poutnik
    Jul 13 at 9:27
  • $\begingroup$ @Poutnik. OK for your modification. Of course. I wanted to be positive about OP's question. $\endgroup$
    – Maurice
    Jul 13 at 17:01
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Liquid evaporation occurs at any temperature, just its rate changes with temperature. You must have seen washed clothes getting dry in open air, water evaporating at room temperature. OTOH, I guess you have not observed boiling sea, lakes or rivers to obtain clouds - putting aside geothermal activities.

There are few things for you to understand:

Molecules do not have temperature. For each temperature, there is a wide distribution of molecular energies and speeds. Similarly, the same molecule energy/speed can belong to a wide range of substance temperatures.

Molecules do not deal with pressure. Well, for usual evaporating context. If a molecule would leave liquid or not does not depend on pressure(*), but on liquid intermolecular forces, the molecule energy and velocity. When a molecule leaves liquid, it moves mostly freely with occasional collisions. ( 10 billions per second as the typical value is still occasional for them )

Only fast enough molecules are able to overcome the liquid cohesive forces to escape to a gaseous phase. At any temperature, there are always some molecules with enough energy to leave. But for low temperatures, it would be very tiny amount.

OTOH, even at the boiling point, only minor fraction of molecules has enough energy to escape liquid, as there is not enough energy in liquid for all of them.(**) This should not be surprising, as there is needed about 5.5 times more energy to boil water out, compared to energy for heating the same amount of water from $\pu{0^{\circ} C}$ to $\pu{100^{\circ} C}.$


(*) In fact, increasing of total pressure very sligthly increases the liquid saturated vapor pressure, but the reason for it is out of the question scope.

(**) There are assumed conditions near to thermal equilibrium of molecular energies at given temperature.

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    $\begingroup$ @buckthorn There is a large set of statements that are generally true, with lists of extreme conditions, when they are not, but which are not useful to be mentioned at given level. :-) I have meant it for conditions not far from thermal equilibrium of molecular energy distribution. $\endgroup$
    – Poutnik
    Jul 13 at 15:34
  • $\begingroup$ Agreed. Perhaps I should have taken an example that is more on topic, but thought I'd make a point that boiling is a non-equilibrium process. $\endgroup$
    – Buck Thorn
    Jul 13 at 19:49
  • $\begingroup$ Surely boiling is not an equilibrium process, but energy distribution is still close to the equilibrium one. (In fact,all observed processes are just more or less close to equilibrium.) $\endgroup$
    – Poutnik
    Jul 13 at 20:18
  • $\begingroup$ Never mind, I see your point. If the constraint is that the temperature be at the boiling point. $\endgroup$
    – Buck Thorn
    Jul 13 at 20:33

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