The issue of vapor phase dew point in salt solutions
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Recently, while studying the phase equilibrium of urea solutions, a very big doubt has arisen. In fact, it can be simply studied using a sodium chloride solution system. For example, in a sodium chloride solution of a certain concentration, at normal pressure, the boiling point is raised due to the presence of the salt solution; thus, the boiling point of the solution at normal pressure might become 106 degrees. At this point, equilibrium is achieved between the gas and liquid phases. According to my understanding, the temperature of the gas phase should be 106 degrees, which means that the dew point of the gas phase should also be 106 degrees. However, using software simulation, since the gas phase consists only of water vapor at this point, the resulting dew point is 100 degrees at atmospheric pressure. Now a contradiction arises; this issue has been studied for a long time, and I hope some expert can help resolve my confusion.
This post was last edited by qugd on 2019-1-9 at 11:13. In multi-component systems, it is necessary to consider the composition of the phases. Water evaporates from the saltwater solution as a result of phase equilibrium between water molecules at the liquid surface of the saltwater and the freshly evolved water vapor; it is a two-component system, and its boiling point increases. However, the gas phase formed above by the evaporated water vapor is a single-component system (assuming that sodium chloride does not evaporate with the water); in this case, its dew point is the temperature at which the water vapor condenses, namely 100 degrees. The gas phase formed when water vapor leaves the liquid surface is a single-component system. Two-component systems are different from single-component systems because the chemical potentials differ. These are my personal opinions.
Hello, thank you for the reply. I also think the only explanation now is that the water has left the liquid surface and entered another system; in this case, the dew point of water vapor needs to be taken into account. It can no longer be considered as being in phase equilibrium with the salt solution, but rather phase equilibrium should be considered between pure water vapor and liquid water (this should be assumed even if condensation does not occur). But another issue is that, for example, in an evaporator (flash tank), a gas phase pipe is connected; so how can one determine the boundary point between these two phases of equilibrium – namely the equilibrium between steam and brine on one hand, and steam and liquid water on the other? Can it be understood that within the evaporator, there is equilibrium between steam and brine, while once the vapor enters the pipe, there is equilibrium between steam and liquid water? This interpretation seems a bit forced; after all, the evaporator and the pipes belong to the same system, so the pressure and temperature should be the same. The main issue is still not knowing how to determine under what conditions the gas leaving will form condensate; I hope someone can help clarify this.
There is another question: if the dew point of the vapor after evaporation remains at 100 degrees, but according to phase equilibrium the temperature of both the gas and liquid phases should be 106 degrees, then does the vapor become superheated steam?
Original poster, has your question from back then been resolved? I ask the same question after N years, and I couldn’t find a suitable answer in any books either. I hope some expert can help answer it; I would be very grateful!
The temperature at which a drop of liquid just vaporizes is the boiling point of the liquid, while the temperature at which a drop of gas just condenses is the dew point of the gas. Under the same conditions, the boiling point and dew point are generally different from each other.