Question

I know this type of question has been addressed before, but I believe mine is slightly different and hints at a potential misunderstanding I've had for a long time.

I understand the concept of equilibrium vapor pressure, that molecules with high enough energies escape the surface of a liquid until the rate of evaporation = rate of condensation. I also understand that for a given temperature, a substance has a characteristic saturation vapor pressure. But one thing bothers me: when considering a system open to the atmosphere, does it even make sense to talk about a vapor pressure developing? Consider, for example, water vapor at 95 C. Obviously, evaporation still occurs at the surface, but since its saturation vapor pressure is less than the surrounding atmospheric pressure, can we even consider a vapor pressure to develop, or is this a term that holds meaning only within the context of closed systems? (I suppose, however, that it does make sense to speak of the vapor pressure in an open system once the boiling point is reached since there is quite literally a vapor pressure building up in bubbles that can push back the atmosphere...)

Answer 1

Main Answer

In an open system, the concept of vapor pressure still holds meaning below the boiling point, as molecules with high enough energies escape the surface of the liquid until the rate of evaporation equals the rate of condensation. However, since the saturation vapor pressure is less than the surrounding atmospheric pressure at temperatures below the boiling point, it may not be practical to consider a separate vapor pressure developing in an open system. Once the boiling point is reached, there is a buildup of vapor pressure in bubbles that can push back against the atmosphere.

Step-by-step explanation

When considering a system open to the atmosphere, it may seem contradictory to speak of a vapor pressure developing since the surrounding atmospheric pressure already exists.

However, below the boiling point, molecules with high enough energies still escape the surface of the liquid until the rate of evaporation equals the rate of condensation.

This equilibrium vapor pressure is characteristic of a given substance at a given temperature and is known as its saturation vapor pressure.

In an open system, this saturation vapor pressure can still be considered as long as it is less than or equal to the surrounding atmospheric pressure.

This means that some molecules will continue to evaporate from the surface of the liquid and join the surrounding atmosphere until they reach their equilibrium state.

However, since the saturation vapor pressure is less than or equal to the atmospheric pressure, it may not be practical to consider a separate vapor pressure developing in an open system.

Instead, we can think of molecules continuously escaping and joining the surrounding atmosphere until they reach their equilibrium state.

Once the boiling point is reached, however, there is a significant difference. At this temperature, the saturation vapor pressure becomes greater than or equal to the atmospheric pressure, and bubbles begin to form within the liquid.

These bubbles contain vapor at a higher concentration than in the surrounding atmosphere, creating a buildup of vapor pressure that can push back against the atmosphere. This phenomenon is known as boiling and results in a significant increase in mass transfer from liquid to vapor phase.

In summary, below the boiling point, it may not be practical to consider a separate vapor pressure developing in an open system due to the lower saturation vapor pressures. However, once the boiling point is reached, there is a significant buildup of vapor pressure within bubbles that can push back against the atmosphere and result in boiling.