Phase diagram of water in real life

[fog37]

Hello,

All substances that are either in the liquid or solid state have a small vapor pressure which implies that they are always slowly (solids especially) evaporating and turning into a gas (vapor).

The phase diagrams of water and carbon dioxide indicate only states (P , V) of pressure and temperature that are in thermodynamical equilibrium but in real life, both water and dry ice are not in equilibrium: water is present both as a liquid (in oceans and lakes) and as a vapor at the same conditions of pressure P and temperature T.
That said, how is the phase diagram of water useful since water is not in stable equilibrium? Is the phase diagram useful only for water in a laboratory setting?

Thanks!
fog37

 

[NFuller]

The phase diagrams of water and carbon dioxide indicate only states (P , V) of pressure and temperature that are in thermodynamical equilibrium but in real life, both water and dry ice are not in equilibrium: water is present both as a liquid (in oceans and lakes) and as a vapor at the same conditions of pressure P and temperature T.

If the system is closed, such that the water is in a sealed container, and is held at constant temperature and pressure, then the system is in equilibrium. Specifically, this could be considered an Isothermal-isobaric ensemble since N, P, and T are fixed constants.

 

[rbelli1]

water is present both as a liquid (in oceans and lakes) and as a vapor at the same conditions of pressure P and temperature T.

This answers your question

Is the phase diagram useful only for water in a laboratory setting?

When you "move" from one region to another in the phase diagram you have a change in energy but not P or T. It is at the exact same point in the diagram thus the quotes around move.

BoB

 

[fog37]

Thank you.

rbelli1, I am grasping how moving from one region of the diagram to another does not change the P and T. Each point in the diagram has a different P and V.
And as NFuller mentioned, the equilibrium is achievable within a closed container but water on Earth is not inside a closed container and that is why I am wondering when the water phase diagram is useful except of the closed container situation...

 

[Borek]

It definitely tells you when to expect water to freeze (think ice on puddles) and when to expect water to condense and whether it will condense into ice or water (think clouds), so it seems to be pretty useful outside of the lab settings.

 

[fog37]

Thanks Borek.

• Just to make sure, in any phase diagram, regardless of the single substance state (solid, liquid, gas), the temperature T is the temperature of the substance itself (measured by inserting a thermometer in it) and the pressure P is the total external pressure acting on the substance, correct? A solid, a liquid and a gas also have an internal pressure pushing towards the outside. How does that factor in? Is the internal pressure always equal to the total external pressure since we talking about stable equilibrium states?

• As far water goes, is the provided partial pressure of water in the air (mixture of gases) always equal to the saturated vapor pressure of water? I don't think so. If that was the case, RH=100% and evaporation would never be possible. At The higher the temperature the higher the water partial pressure will be and the higher the reachable saturated vapor pressure will be. The pressure in the table below should just be the partial vapor pressure (not the saturated vapor pressure).

T, °C P, torr P, atm
0 32 4.5851 0.0060
5 41 6.5450 0.0086
10 50 9.2115 0.0121
15 59 12.79 0.0168
20 68 17.54 0.0231
25 77 23.76 0.0313
30 86 31.84 0.0419
35 95 42.20 0.0555
40 10 55.3 0.0728
45 11 71.92 0.0946
50 12 92.5 0.1218​

When dynamic equilibrium is reached, the water vapor pressure is the maximum pressure the water vapor can have at that specific temperature, correct? The relative humidity is 100% when the water partial pressure is saturated, i.e. maximum. When further cooled, the airborne water vapor will condense.

• In the discussion of dew point, condensation, saturation the presence of the other gas constituents in the air seem to not have an effect, correct? Well, but when we say that water freezes at p=1 atm and T=0 C, we are including the partial pressures of all other gases in determining the freezing behavior of water, correct?

• From the pure water phase diagram, we see that at a temperature of 20 °C and absolute pressure of 1 atm (NTP), the water vapor pressure is just 0.00231 atm. When we study the phase diagram of water, we would use 1 atm (not 0.00231 atm), and obtain that water is supposed to be a liquid (not a vapor). However, in real life, vapor seems to be the direction/phase water is going for (evaporation would evaporate all water if it wasn't for the water cycle)...

 

[Chestermiller]

In cases where the gas phase consists of a mixture of air and water vapor, the equilibrium is established when the partial pressure of water vapor is equal to the equilibrium vapor pressure of water at the liquid temperature. This applies even if the total pressure (due mostly to the air) is 1 atm.

 

[fog37]

Thank you Chestermiller.

So the other air molecules (nitrogen, oxygen, etc.) present in the air can, at most, slightly delay when dynamic equilibrium (condensation=vaporization for water) is reached due to collisions with the water vapor molecules. The partial water vapor pressure is only affected by T (same goes for all other gases in the air).

But, just to make sure, at a certain specific temperature T, the water vapor pressure would continue to increase until it reaches its saturated vapor pressure value (the highest pressure it can be for that temperature) once evaporation rate = condensation rate.
But if stating the partial pressure at a certain T means stating the saturated vapor pressure, that implies that equilibrium is always reached at a certain temperature...Is that the case?

 

[Borek]

Not sure what your question is - given enough time every system goes to equilibrium.

 

[fog37]

Sorry for the lack of clarity. Let me reformulate my question better:

The water vapor pressure (i.e. partial vapor pressure) at the temperature T=20 C is 0.02 atm (see table above). Is that tabulated vapor pressure referring to the saturated vapor pressure of water vapor?

From the phase diagram of water, the boiling temperature is 100 C when P=1atm. That pressure of 1 atm is the total air pressure (due to all gases in the air) above the liquid interface and not just the partial vapour pressure which is only 0.86 atm at that temperature. Is that correct?

 

[Chestermiller]

Sorry for the lack of clarity. Let me reformulate my question better:

The water vapor pressure (i.e. partial vapor pressure) at the temperature T=20 C is 0.02 atm (see table above). Is that tabulated vapor pressure referring to the saturated vapor pressure of water vapor?

Yes.

From the phase diagram of water, the boiling temperature is 100 C when P=1atm. That pressure of 1 atm is the total air pressure (due to all gases in the air) above the liquid interface and not just the partial vapour pressure which is only 0.86 atm at that temperature. Is that correct?

No. The vapor pressure of water at 100 C is 1 atm. When water boils at 100 C (in contact with the atmosphere), the bubbles that form within the bulk of the liquid are pure water vapor with virtually no air inside them. These bubbles are able to form because their pressure is high enough to allow the surrounding liquid water to push back the air atmosphere above. If the room is not hermetically sealed (and if you have enough liquid water), if you keep adding heat to the liquid, the air will be purged from the room, and you will eventually have only water vapor.

 

[fog37]

Thank you!

But if the pressure of 1 atm (boiling at T=100 C and P=1 atm) is solely and entirely due to water vapor, it would mean that the actual total external pressure must be higher than 1 atm since there are other gases and vapors that also contribute to the overall pressure with their partial pressures. It was said before that the variable pressure P in the phase diagram is the total external pressure and not just the water vapor pressure...

 

[Chestermiller]

Thank you!

But if the pressure of 1 atm (boiling at T=100 C and P=1 atm) is solely and entirely due to water vapor, it would mean that the actual total external pressure must be higher than 1 atm since there are other gases and vapors that also contribute to the overall pressure with their partial pressures. It was said before that the variable pressure P in the phase diagram is the total external pressure and not just the water vapor pressure...

I never said that, and, it's not correct. The pressure P on the phase diagram is the equilibrium vapor pressure of water. In cases where other gases are present (e.g., air), we can still use the phase diagram if we regard equilibrium for water to be established if the partial pressure of the water vapor in the gas mixture above the liquid is equal to the equilibrium vapor pressure. But, at 100 C, if the total pressure is 1 atm and equilibrium is present, the gas space must contain only water vapor. Therefore, if it is possible for gases to escape the room, all the air will eventually be purged from the room (provided there is still liquid water present).

 

[Borek]

Imagine a cylinder filled with water, with a movable piston resting on the water surface. External pressure is that of the atmosphere - 1 atm. You start to heat water. Once you get to 100 °C water starts to boil and the vapor pushes the cylinder up. But, as the cylinder moves pressure of water vapor never exceeds 1 atm.

In reality there is no well defined boundary between water vapor and the air, but the principle remains the same.

 

[fog37]

Ok thanks.

That makes sense. So The pressure P on the vertical axis is the saturated vapor pressure of water and at T=110 ^oC the saturated water vapor pressure must be 1 atm.

However, in real life, when we see a pot of water boiling, there is air above the water (which contains a variable amount of water vapor, on average around 1% at sea level, and 0.4% over the entire atmosphere) so the contribution from water vapor is usually very small and does not contribute with 1 atm. The measured outside pressure of 1 atm derives from all the partial pressures of the gases/vapors in the air together. that means the water is boiling at T=110 ^oC even when the water vapor pressure is much much less than 1 atm.

 

[Borek]

at T=110 oC the saturated water vapor pressure must be 1 atm.

No, it must be higher. 1 atm is at 100 °C.

Note:

• by definition boiling means water starts to evaporate not only at the surface, but also under the surface
• despite being defined as force/surface pressure is a property of a point in space

 

[Chestermiller]

Ok thanks.

That makes sense. So The pressure P on the vertical axis is the saturated vapor pressure of water and at T=110 ^oC the saturated water vapor pressure must be 1 atm.

However, in real life, when we see a pot of water boiling, there is air above the water (which contains a variable amount of water vapor, on average around 1% at sea level, and 0.4% over the entire atmosphere) so the contribution from water vapor is usually very small and does not contribute with 1 atm. The measured outside pressure of 1 atm derives from all the partial pressures of the gases/vapors in the air together. that means the water is boiling at T=110 ^oC even when the water vapor pressure is much much less than 1 atm.

That is because the water vapor in the air is not at equilibrium with the liquid water. Only at the surface between the liquid water and gas phase is the water partial pressure equal to the equilibrium vapor pressure (and, at this location the partial pressure of the oxygen and nitrogen in the air is zero). The water vapor generated by the evaporation diffuses from the surface to the bulk air, driven by the partial pressure difference between the surface and the bulk air. Gradually, by this process, the air in the room gets "humidified."

 

[fog37]

Thanks again.

1) "That is because the water vapor in the air is not at equilibrium with the liquid water". I agree.

2) "Only at the surface between the liquid water and gas phase is the water partial pressure equal to the equilibrium vapor pressure (and, at this location the partial pressure of the oxygen and nitrogen in the air is zero). " Why are the partial pressure of oxygen and nitrogen zero at the interface and only the water vapor exerts a pressure? That would explain how the only vapor pressure present at the interface is the water vapor pressure and its value is exactly 1 atm. But what happened to those other gases? Did they get pushed out of that region of space near the air-liquid interface?

3) The water vapor generated by the evaporation diffuses from the surface to the bulk air, driven by the partial pressure difference between the surface and the bulk air. Gradually, by this process, the air in the room gets "humidified." I agree.

 

[Chestermiller]

2) "Only at the surface between the liquid water and gas phase is the water partial pressure equal to the equilibrium vapor pressure (and, at this location the partial pressure of the oxygen and nitrogen in the air is zero). " Why are the partial pressure of oxygen and nitrogen zero at the interface and only the water vapor exerts a pressure? That would explain how the only vapor pressure present at the interface is the water vapor pressure and its value is exactly 1 atm. But what happened to those other gases? Did they get pushed out of that region of space near the air-liquid interface?

Yes. Even below the boiling point, the partial pressure of water vapor at the interface is equal to the equilibrium vapor pressure, so the partial pressures of the other gases total 1 atm minus the equilibrium vapor pressure of water.

 

[russ_watters]

One point I'm not sure if you have clear about how boiling works:

Boiling happens where heat is applied. In a pot of water on an electric range, that's at the bottom of the pot, where the pot touches the heating element. So the water that is actually boiling is the water touching the hot bottom of the pot. This water, clearly, only has water around it; no other gases. The pressure at the bottom of the pot is atmospheric pressure plus the weight of the water above, which makes the boiling temperature just slightly above 100C. The bubbles form, then rise to the surface (and cool to 100C).

That's what differentiates boiling from evaporating: it happens at the bottom of the pot, not at the surface. So it doesn't matter what the composition of the air above is; it just matters how fast you add heat. Turn off the burner and the water will still be at 100C, but it will evaporate from the surface at a rate dependent on the composition of the air.

 

[fog37]

Thanks russ_water.

My understanding of boiling is the following: boiling is vaporization throughout the entire volume while evaporation is vaporization only happening at the surface.
As you mention, the atmospheric pressure at the water surface is transmitted to all points inside the liquid water. Position inside the water that are deeper in the pot experience a higher pressure due to the added hydrostatic pressure contribution. Boiling take places when bubbles of water vapor manage to form (with the help of condensation nuclei). A bubble forms when the bubble vapor internal pressure is equal to the pressure around them in the liquid (atmospheric+hydrostatic). For the bubble to form, the rates of evaporation and condensation must be also be equal hence the bubble internal pressure is the saturated vapor pressure. Being less dense, these bubbles rise and pop and a convective cell is form.

So the bubbles only really form at the bottom of the pot (where there is more heat) near the range and not through the entire volume. So I guess it is incorrect to say that boiling is vaporization occurring through the entire volume...

My entire thread started from my being uncertain on how to interpret the pressure P in a phase diagram (total external pressure or just saturated vapor pressure of the substance in the diagram).

 

[Borek]

That's what differentiates boiling from evaporating: it happens at the bottom of the pot

I am afraid that idea only adds to the confusion. By definition boiling happens everywhere in the bulk of the liquid.

Yes, in practice, because there is a (gravitational) pressure gradient inside of the liquid, because the heat transfer is slow, because surfaces that are in the contact with the liquid provide nucleations sites helping to overcome the activation energy required to create bubbles, evaporation inside of the liquid takes place mainly on the pot walls and especially on the heater surface. But technically that's not how boiling is defined.

If you have ever observed water heated in a glass beaker, if there are any solid impurities inside of the liquid (think boiling stone small enough to be floating) you can clearly see bubbles forming on them inside of the liquid. That's because glass surface is too smooth for the bubbles to start.

 

[russ_watters]

I am afraid that idea only adds to the confusion. By definition boiling happens everywhere in the bulk of the liquid.

I disagree with "happens everywhere", but if you mean can happen anywhere [depending on other factors], then i agree.

The wiki actually describes four separate types of boiling. I was describing what is seen in a specific situation, but did not mean to imply it covered every situation of boiling.

 

[Borek]

I disagree with "happens everywhere", but if you mean can happen anywhere [depending on other factors], then i agree.

Then we are talking about the same thing. Definition of boiling assumes an ideal model and ignores activation energy required for the bubble creation.

 

[zbikraw]

In practice, differences between various states defined as equilibrium ones may be small in comparison with measurement errors and deviations of real systems from ideal equilibrium ones. For using phase diagrams measured and calculated for pure water, we must remove from experimental sample any traces of another substances (the "impurities"), mainly atmospheric gases and water-soluble substances. In the presence of such substances we must consider solution equilibria.
In principle, equilibrium in many-phase systems is attainable only at any phase surface. This leads to many equilibria inside any phase "bulk", different from surface ones. In solution thermodynamics, the composition of surface phases may be different from "bulk" ones. Such a situation is exploited in separation and purification processes, for example zone melting.
Surface pheses in general have higher energy than "bulk" ones, so they more easily surpass activation energy for phase change.
There is another complication, rarely mentiond in thermodynamics, namely the consequences of a Curie rule (one of the most fundamental in science). For the sake of present discussion, it states that the symmetry of present (or future) state is the result of symmetry of past (or present) one. This means that equlibrium can be attained only between phases of the same symmetry, in practice it is sufficient to have some common symmetries. So it is impossible to obtain equilibria between gases, liquids and solids. As the real life equilibria exist in phase boundaries ("surface phases"), we can obtain different phases with inexpected symmetries. This is mainly behind the real research, but frequently results in impossibility of applying the simple and well proven rules in a given situation. In practice we met situations characterized by unexpected variability of phase transfer kinetics, of course we mention only the sluggishness of some processes.

zbikraw

 

[Ygggdrasil]

My entire thread started from my being uncertain on how to interpret the pressure P in a phase diagram (total external pressure or just saturated vapor pressure of the substance in the diagram).

The P in the phase diagram refers to the partial pressure of water vapor. Along the curve between the condensed phase (liquid/solid) and vapor phase, the values of (P, T) give the saturated vapor pressure of water at the given temperature.

When the partial pressure of water vapor is below the saturated vapor pressure (at the given temperature), evaporation will be favored and you will have a net transfer of material from the solid/liquid phase to the gas phase over time. When the atmospheric pressure is below the saturated vapor pressure (at a given temperature), boiling will occur which will significantly speed the rate of mass transfer from the liquid phase to gas phase.

You are correct that we observe many systems that are out of equilibrium. Weather can occur only because the atmosphere is not at equilibrium.

 

[fog37]

Hello everyone.

• In an older post, Ygggdrasil explains that the term "vapor pressure" does not actually mean pressure of the vapor. Vapor pressure always means and implies saturated vapor pressure, i.e. the partial pressure of a vapor over a liquid at equilibrium given a specific temperature. That seems clear.

• If the water temperature is T= 25 Celcius (room temperature), the partial pressure of water in the open atmosphere at saturation gets called vapor pressure and is about 24 torr. The overall atmospheric pressure is 1 atm =760 torr . Going back to the phase diagram of water and how we can use it to correctly predict the state of water based on T and P: the pressure P in the phase diagram always represents the saturated partial pressure of vapor and not the total external pressure applied to the substance (by a cylinder or by the joint partial pressure of all the gases in the mixture). For example, at room temperature T =25 C, the atmospheric pressure is 1 atm and water vapor only contributes 8 torr with its partial pressure to the total pressure of 760 torr (due to all partial pressures of the gases together). P= 8 torr is lower than the vapor pressure (24 torr) of water at room temperature For example, from the water phase diagram, the normal boiling point is indicated to be pressure P = 760 torr and T=100 C. But if P on the vertical axis is the saturated vapor pressure, that would imply that the saturated, equilibrium vapor pressure of water is 760 torr at boiling. I don't think that is the case since 1 atm is the entire pressure of the air mixture...

• Any point on the phase diagram of water indicates a state (T and P) that is thermodynamically in equilibrium. That means that the substance (water) will not be going through a state transformation at those condition of P and T. For example, let's pick P=500 torr and T= 60 C on the diagram. The diagram tells us that water must be a liquid. There is not mention of the fact that water, in real life, is always slowly evaporating away. So the diagram implicitly works is we assume situations like RU=100% in the open atmosphere or liquid water in the presence of its saturated vapor in a close container...

Thanks for any correction.

Fog37

 

[Chestermiller]

Hello everyone.

• In an older post, Ygggdrasil explains that the term "vapor pressure" does not actually mean pressure of the vapor. Vapor pressure always means and implies saturated vapor pressure, i.e. the partial pressure of a vapor over a liquid at equilibrium given a specific temperature. That seems clear.

Correct

• If the water temperature is T= 25 Celcius (room temperature), the partial pressure of water in the open atmosphere at saturation gets called vapor pressure and is about 24 torr. The overall atmospheric pressure is 1 atm =760 torr . Going back to the phase diagram of water and how we can use it to correctly predict the state of water based on T and P: the pressure P in the phase diagram always represents the saturated partial pressure of vapor and not the total external pressure applied to the substance (by a cylinder or by the joint partial pressure of all the gases in the mixture). For example, at room temperature T =25 C, the atmospheric pressure is 1 atm and water vapor only contributes 8 torr with its partial pressure to the total pressure of 760 torr (due to all partial pressures of the gases together). P= 8 torr is lower than the vapor pressure (24 torr) of water at room temperature

This is correct. If the pressure of water vapor in the bulk of the gas phase is at 8 torr, the system is not at equilibrium, and water vapor is continuing to evaporate. The relative humidity is 33.3%.

• For example, from the water phase diagram, the normal boiling point is indicated to be pressure P = 760 torr and T=100 C. But if P on the vertical axis is the saturated vapor pressure, that would imply that the saturated, equilibrium vapor pressure of water is 760 torr at boiling. I don't think that is the case since 1 atm is the entire pressure of the air mixture...

This is correct. This is not a system in equilibrium. The bubbles that form within the bulk of the liquid water below the surface are saturated with water vapor at 760 torr. But the air above the boiling water is not saturated with water vapor. And typically, the gas phase temperature is below 100 C. If the system were in equilibrium, there would have to be no air present, and the entire gas phase would have to be water vapor at 100 C. It could also work if the total pressure were greater than 760 torr.

• Any point on the phase diagram of water indicates a state (T and P) that is thermodynamically in equilibrium. That means that the substance (water) will not be going through a state transformation at those condition of P and T. For example, let's pick P=500 torr and T= 60 C on the diagram. The diagram tells us that water must be a liquid. There is not mention of the fact that water, in real life, is always slowly evaporating away. So the diagram implicitly works is we assume situations like RU=100% in the open atmosphere or liquid water in the presence of its saturated vapor in a close container...

Fog37

or liquid water with no water vapor in a closed container.

 

[Borek]

In an older post, Ygggdrasil explains that the term "vapor pressure" does not actually mean pressure of the vapor. Vapor pressure always means and implies saturated vapor pressure, i.e. the partial pressure of a vapor over a liquid at equilibrium given a specific temperature. That seems clear.

Without added context this is incorrect. If you put some water in a tank with head space above, if there is enough water you will get to the point when there are both liquid and gas and the system is in a dynamic equilibrium, same amount of water condenses and evaporates in any given period of time. But if there were not enough water, you will end without a liquid, there will be only vapor with a pressure below that of the saturated one, which clearly falsifies the statement "vapor pressure always means and implies saturated vapor pressure".

 

[Ygggdrasil]

Without added context this is incorrect. If you put some water in a tank with head space above, if there is enough water you will get to the point when there are both liquid and gas and the system is in a dynamic equilibrium, same amount of water condenses and evaporates in any given period of time. But if there were not enough water, you will end without a liquid, there will be only vapor with a pressure below that of the saturated one, which clearly falsifies the statement "vapor pressure always means and implies saturated vapor pressure".

My point was that the term vapor pressure has a specific definition in chemistry: "Vapor pressure or equilibrium vapor pressure is defined as the pressure exerted by a vapor in thermodynamic equilibrium with its condensed phases (solid or liquid) at a given temperature in a closed system." (https://en.wikipedia.org/wiki/Vapor_pressure).

In this and other threads, posters have the tendency to use the term "vapor pressure" when instead they should be saying "partial pressure of water vapor." In your example, you are correct that the partial pressure of water vapor will be below the vapor pressure of water at the particular temperature, but you are not interpreting the term "vapor pressure" correctly (likely an issue of how terms translate between English and Polish).

 

[Ygggdrasil]

Any point on the phase diagram of water indicates a state (T and P) that is thermodynamically in equilibrium. That means that the substance (water) will not be going through a state transformation at those condition of P and T. For example, let's pick P=500 torr and T= 60 C on the diagram. The diagram tells us that water must be a liquid. There is not mention of the fact that water, in real life, is always slowly evaporating away. So the diagram implicitly works is we assume situations like RU=100% in the open atmosphere or liquid water in the presence of its saturated vapor in a close container...

Remeber that the pressure is the partial pressure of water vapor, not the total atmospheric pressure. If the partial pressure of water is 500 torr at 60°C it is well above its vapor pressure (~150 torr). In this case, water vapor will precipitate until the amount of water vapor has a partial pressure of ~ 150 torr (100% relative humidity).

Under normal conditions (25°C, 8 torr water vapor), you will see that the phase diagram shows that vapor is the preferred phase. This matches with expectation as evaporation is preferred. If you spill a little water on your kitchen table it will slowly evaporate over time under those conditions. If you leave a wet towel hanging in the bathroom it will eventually dry out. However, if you have just taken a hot shower and the 25°C room suddenly has a partial pressure of water above the vapor pressure of water at that temperature (24 torr), water vapor will begin to condense from the air onto other surfaces in the bathroom. Under these conditions, your wet towels will not begin to dry out until the saturated air in the bathroom mixes with dry air from the rest of your home to lower the partial pressure of water to below 24 torr.

 

[fog37]

Sorry for being highly repetitive, but how should we then interpret what happens, for example, in the water phase diagram at T =100 C and P = 760 torr? Most would say that water boils at T = 100 C when the pressure P is 1 atm and implicitly think it that the pressure P is the total atmospheric pressure applied on the water since in everyday life that is what happens when water boils. But that does not seem correct from the standpoint of the phase diagram of water...

That leads to think that using the phase diagram of water or carbon dioxide to predict how those two substance behave in everyday life, under the atmospheric pressure, in the open atmosphere, is not good practice...

 

[Ygggdrasil]

Sorry for being highly repetitive, but how should we then interpret what happens, for example, in the water phase diagram at T =100 C and P = 760 torr? Most would say that water boils at T = 100 C when the pressure P is 1 atm and implicitly think it that the pressure P is the total atmospheric pressure applied on the water since in everyday life that is what happens when water boils. But that does not seem correct from the standpoint of the phase diagram of water...

Evaporation occurs when the partial pressure < vapor pressure. (e.g. in a typical room is 25°C, 8 torr water vapor, so water will tend to evaporate).
Boiling occurs when the vapor pressure > atmospheric pressure.

Boiling implies evaporation but evaporation does not imply boiling.

 

[fog37]

I believe I understand, thanks to your help, the concepts you are mentioning: evaporation always occurs when the vapor pressure is less than the equilibrium vapor pressure and boiling when the equilibrium vapor pressure is larger than atmospheric pressure.

So, in the graph below, the coordinates P=1 atm and T=100 C indicate the boiling point (the normal boiling point since any other point sitting along that same line is also a boiling point) and that means the water vapor pressure P has to equal to 1 atm for boiling to occur, i.e. the vapor pressure on its own must be equal to the the same pressure that the entire mixture of gases composing air exerts on the free surface of the water and inside the liquid by pascal principle. It remains that the vertical axis is always and only the equilibrium, saturated, vapor pressure of water no matter if we are talking about the blue (vapor), orange(solid) or green(liquid) area in the diagram...

 

[Borek]

My point was that the term vapor pressure has a specific definition in chemistry: "Vapor pressure or equilibrium vapor pressure is defined as the pressure exerted by a vapor in thermodynamic equilibrium with its condensed phases

OK, I see. Seems like English nomenclature is quite convoluted. There are three things:

1. Equilibrium (saturated vapor pressure).
2. Non-equilibrium vapor pressure.
3. Partial pressure.

Calling the first one "vapor pressure" (when it already has two unambiguous names - "equilibrium vp" and "saturated vp"), and the second "partial" (when "partial pressure" typically means "pressure exerted by one of the components of the mixture") seems to be designed to make things confusing. So, when humidity is 50% do we refer to it as "partial partial pressure of water vapor?"