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Water and its interesting 3 phases...

  1. Nov 12, 2016 #1
    Hello Forum,
    I am reviewing water and its 3 phases (solid, liquid and gas).

    a) At the triple point, water can be in the three phases simultaneously. Does that mean that for a certain amount of water, say 30grams, we would find, approximately, 10 grams of liquid water, 10 grams of ice and 10 grams of gaseous water?

    b) The boiling point for water, at atmospheric pressure, is 100 Celcius. I clearly see how water exists as a liquid at temperatures below 100C. But there is also water in the gaseous state below 100C. How do we explain that? My explanation is that many water molecules that are in the liquid phase are energetic enough to escape the liquid phase and get into the gaseous state. but water is predominantly liquid below 100 C. Surely, there is still a lot of water (vapor) in the air. So water appears to exist both as a liquid and gas at the same temperature and pressure. Can other substances do that?
    I know that other materials, even when solid, can slowly evaporate (metals, for example). It seems that all three states of matter are actually taking place at the same time but some are more dominant than others at particular temperature and pressure...

    Steam is water in the gaseous state. But steam looks white while water in the gaseous state is clear. Why? Is steam just a denser cloud of gaseous water? I know that there is a difference between gas and vapor: gas is when there is no liquid coexisting with the gas. Vapor is the gaseous state when the substance is also in the liquid state around it...

    thanks!
     
  2. jcsd
  3. Nov 12, 2016 #2

    Borek

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    No. Any combination of amounts will do.

    All substances do that.

    No, typically only two phases are present at equilibrium at given P,T (with the exception of the triple point).

    What looks white is not a gaseous water, but water that have already condensed into tiny droplets (so it is a liquid).
     
  4. Nov 12, 2016 #3

    BvU

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    Hello Fog,

    Fog and visible steam are small water droplets.
    a) No. At the triple point of water any amount of each of the phases can be present.
    b) yes. Vapour-liquid equilibrium is common

    [Borek was faster ...]
     
  5. Nov 12, 2016 #4

    rbelli1

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    Last edited by a moderator: May 8, 2017
  6. Nov 12, 2016 #5

    Ygggdrasil

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    All liquids will have a vapor pressure at a specified temperature, which describes the amount of vapor present over the liquid at equilibrium. The boiling point of a substance is defined as the temperature at which the vapor pressure exceeds atmospheric pressure (hence, why water has a lower boiling point at higher elevations where the atmospheric pressure is lower). At temperatures below the boiling point, water can evaporate only from the surface of the liquid. However, when the vapor pressure exceeds the atmospheric pressure, vapor can now push against the pressure of the liquid to form bubbles of vapor throughout the liquid (hence the phenomenon of boiling).
     
    Last edited: Nov 12, 2016
  7. Nov 12, 2016 #6
    Thank you all for the good information. I really appreciate it.

    • So steam, which looks white, is a actually a small cloud (tiny water droplets). In general, a gas is defined as that state a substance has when it is above its critical temperature. As we compress the gaseous substance, it will remain a gas until the critical temperature (specific to that substance) is reached. More compression will convert the gas into a liquid.
    • A saturated vapor is similar to a gas but a vapor is usually described in the context of of a closed container that also contains the liquid form of the substance: the faster molecules in the liquid manage to escape into the gaseous form until dynamic equilibrium is reach (equal number of particles going from/to the liquid into the vapor/liquid phase). We call the vapor "saturated" when that equilibrium is reached. This saturated vapor can exert a pressure. Is the the water in the air we breath, above the seas, the oceans, the lakes, etc. a saturated water vapor even if we are not really dealing with a closed container since the Earth atmosphere does not really have a lid?
    • "The boiling point of a substance is defined as the temperature at which the vapor pressure exceeds atmospheric pressure..." The atmospheric pressure is given by the sum of all the gaseous (or vapor?) particles in the air itself (water, carbon dioxide, etc.). Each group of particles will exert a partial pressure which I would call the vapor partial pressure. Boiling takes place inside the liquid where air bubble manage to grow and become buoyant to eventually make it to the surface....
    Thanks!
     
  8. Nov 14, 2016 #7
    Hello again,

    Water boils at ##T= 100 C## (if p = 1 atm) and it is said that the boiling (water vapor bubble formation) occurs when the saturated vapor pressure becomes equal to the external atmospheric pressure. But water bubbles form inside the liquid while the saturated vapor pressure idea seem to pertain to the vapor forming outside of the liquid, above its surface in the case of a closed container. Boiling also occurs in an open pot of water.

    How do we relate/connect the idea that saturated vapor pressure must equal the atmospheric pressure for boiling to occur inside the volume of the liquid?

    My explanation is that the atmospheric pressure ##P_{atm}## acts at the free surface of the liquid but it is also transmitted at every point inside the liquid. As more and more heat is transferred to the water from the heat source at the bottom, bubbles will start forming and growing if the internal pressure of the bubbles becomes, at least, equal to the external liquid pressure at the bubble location which is equal to ##P_{atm}+(rho\times g\times depth)##

    If that is correct, I guess the vapor pressure represents in this case the internal vapor pressure inside the bubble and that vapor, forming the spherical volume must be saturated, i.e. in dynamic equilibrium with its surrounding liquid.
     
  9. Nov 14, 2016 #8

    Borek

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    At first sight looks OK to me.
     
  10. Nov 14, 2016 #9

    Bystander

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  11. Nov 14, 2016 #10

    rbelli1

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    While substances above their critical temperature are gasses not all gasses are above that substance's critical temperature.

    BoB
     
  12. Nov 15, 2016 #11
    Thanks rbelli1.

    Soupercritical ##CO_{2}## is used ion the coffee industry for decaffeination. Supercritical fluids are considered a continuum which has both liquid and gas properties. This continuum is obtained when a gas is brought to a pressure and a temperature higher than its critical values.

    Overall, does a supercritical fluid look more like a liquid or like a gas in appearance?
     
  13. Nov 15, 2016 #12

    Borek

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    As far as I know they are gaseous enough you can't watch them - you can see a liquid, as it is a condensed phase that doesn't expand rapidly without a pressure, you can't see a gas in the same way, as it is not separated from the surroundings.
     
  14. Nov 15, 2016 #13
    So we can see things because they scatter light towards us. A piece of glass is transparent even if it is solid. Most gases, I guess, are not visible.

    As far as water in liquid and vapor phase, water exists as liquid in lakes and ocean and as a vapor in the air. So the below 100 C, water is not only in its liquid state, clearly. Sometimes Same goes for temperatures at 0 or below 0 Celcius: we don't have only the solid phase, i.e. ice, but also the vapor state, i.e. the water in the air, correct?

    The liquid/vapor phases coexist even below the boiling temperature.
     
  15. Nov 15, 2016 #14

    Borek

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    In every conditions water exists as a gas in equilibrium with something else. Sometimes it is a solid, sometimes it is a liquid, at triple point it is both.

    That the water exists as a gas only above the boiling point is a common misconception. I see people being surprised it exists as a gas at room temperature many times a year. It is funny, as from the point of view of the liquid/gas equilibrium boiling point is in no way special - it is the external pressure that makes it distinct, not properties of the equilibrium itself.
     
  16. Nov 15, 2016 #15
    Ok, thanks. As reality shows, water at temperatures between 0 and 100 C, at 1 atm, is present both as a liquid and as a gas. But the phase diagram of water shows that water is only a liquid at 1 atm and within that temperature range.

    What arguments are used to explain that the liquid/gas equilibrium exists for water?
     
  17. Nov 15, 2016 #16

    Drakkith

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    Vapor pressure?
     
  18. Nov 16, 2016 #17

    Bystander

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    The phase rule?
     
  19. Nov 16, 2016 #18

    Borek

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    Sorry, lousy wording. Phase diagram shows pressure of just water vapor - that's not equivalent to the conditions in general (external pressure). Sure, if there is not enough water to produce saturated vapor, there will be no liquid/solid, only the gas. What I meant was that if there is enough water, as it evaporates we will be moving through the phase diagram till we hit one of the equilibrium lines.
     
  20. Nov 16, 2016 #19
    Ok Borek, I need your patience.....You mention that the phase diagram shows pressure of just water vapor. I am just not there yet...

    A phase diagram plots p versus T (the internal pressure of the substance versus the temperature of the substance) for a particular substance. Let's talk about water. At each point ##(p,T)## on the diagram, water is either a solid (ice, of various types), a gas, a liquid or a supercritical fluid. A supercritical fluid is a gaseous state at a temperature and pressure above its critical point ##(p_{c},T_{c}) ## (it can effuse through solids like a gas, and dissolve materials like a liquid). We cannot liquefy a gas that is at or above its critical temperature ##T_{c}##. But there is also the critical pressure ##p_{c}## which is defined as the smallest pressure needed to liquefy a gas at its critical temperature...This seems contradictory since it was said that there is no sufficiently large pressure that can liquefy a gas at or above ##T_{c}##...

    A gas that is below its critical temperature ##T_{c}## is called a vapor. Does it need to also be below critical pressure ##p_{c}##? In general, we look at what happens at p= 1 atm, i.e. the gas internal pressure is 1 atm, the same as the outside atmospheric pressure. When the substance is in the atmosphere, its pressure is simply 1 atm unless we mechanically compress it or expand it.

    That said, I am still not clear on how I can explain, using the the phase diagram of water, why water can be in its liquid state and also in its vapor state at the same time (not mixed together). For p =1 atm and 0<T<100, the diagram shows that water is only a liquid.

    Let's see if I reach some enlightment...
     
  21. Nov 16, 2016 #20

    Borek

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    There is some misunderstanding, perhaps because of my lousy wording again.

    What I meant was that in the air you may have a pressure of 1 atm but it is not necessarily the pressure that should be found on the phase diagram to find out what state the water is in. What is important is not the total pressure, but just the partial pressure of the water. Or, in other words, phase diagram can be used to predict state of the system that consists of only water (which is rarely what we are dealing with). In a real world even if the P,T parameters seem to suggest water should be present only as a liquid (say, 50 °C, 1 atm), the partial pressure of water is much lower.

    Yes, and in a closed system with nothing else, pressurized to 1 atm, there will be no gas.
     
  22. Nov 16, 2016 #21
    Thanks Borek. You are very clear and I am making progress. It is just taking me a while.

    The air is a mixture of many elements. Water is about 0-4%. So, in the phase diagram p versus T, is p the vapor pressure, not the (total) external pressure, and T the internal temperature of the substance under consideration? I don't think p is the internal pressure of the substance....

    As you mention, to understand how water is both liquid and vapor in real life at ordinary T and at 1 atm, we need to worry about the partial pressure of water vapor which is much smaller than 1 atm. That said, what if we looked at the phase diagram at ordinary temperature and at that much lower than 1 atm pressure? Would the phase diagram predict both phases? I don't think so.
     
  23. Nov 16, 2016 #22

    Ygggdrasil

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    If you spill some water at room temperature, 1 atm pressure, and normal humidity does the puddle just stay there or does it eventually evaporate?

    The puddle eventually evaporates reflecting the fact that liquid water is not thermodynamically favored when the partial pressure of water in the air is lower than its vapor pressure at a specified temperature. So, under normal conditions, you should expect any water you have in an open container to evaporate away. It just takes a long time to reach equilibrium because the water can only evaporate on the surface. When the vapor pressure exceeds atmospheric temperature, boiling can occur, which increases the speed with which liquid water can convert to water vapor, letting the system reach equilibrium more quickly.
     
  24. Nov 16, 2016 #23
    Thanks Ygggdrasil. your inputs are helpful.

    Just for my sake and for those like me that are trying to learn, this is what I think I know with confidence and correctly:

    a) Closed container initially empty (vacuum): a liquid is introduced, fast surface molecules escape the liquid phase into the vapor phase. This process is called evaporation (or surface vaporization) and happens at any temperature. The vapor exerts a pressure on the container walls. As time goes on and evaporation continues, there will be less an less liquid and more vapor. The vapor pressure will increase until dynamics equilibrium is reached (rate of molecules condensing back to the liquid phase = rate of molecules going into vapor phase). At that point the vapor is called "saturated" and its pressure on the container's walls and on the liquid itself is the saturated vapor pressure.
    If we repeated the same experiment with a very very larger container, the entire liquid will probably evaporate into the vapor phase but the final saturated vapor pressure will be the same as in the first example. If we repeated the experiment but the liquid was at a higher temperature, evaporation would be faster and the final saturated vapor pressure would be larger.

    b) Boiling is the same as evaporation but happens throughout the volume of the liquid once the saturated vapor pressure inside a bubble of water vapor inside the liquid is equal to or larger than the atmospheric pressure+the hydrostatic pressure. The vapor bubble does not get squashed and climbs to the surface by buoyancy. Both evaporation and boiling are endothermic, i.e. they require thermal energy from the outside: when puddle evaporates it get energy from the surrounding air. In the case of boiling, an external source of heat is needed. Otherwise the temperature would decrease due to the the fast molecules leaving the liquid phase;

    b) The air is composed of many elements, mainly oxygen and nitrogen (). Water is 0-4%. CO_2 is about 0.03%, etc. Each element exerts its own partial pressure as if it was existing independently. If the atmospheric pressure is 1 atm, then the sum of all the partial pressures is 1 atm;

    c) In the open atmosphere, it is said that both liquids and solids have a vapor pressure in the sense that they evaporate, I guess: molecules go from the liquid phase of from the solid phase into the vapor phase. The vapor pressure of solid is very very small. What pressure are we referring to in this case? The pressure exerted by what? By a vapor that is inside the solid or the liquid?

    d) Phase diagrams: I used to think about the phase diagram in the wrong way, I guess. For example, the phase diagram of ##CO_{2}## would easily predict how at p= 1atm and ambient temperature dry ice would sublimate. But now that I am trying to understand how water truly behaves in a container open to the atmosphere, I am confused. Could you confirm what the pressure p and the temperature T pertain to? Is p the vapor pressure of the substance due to the vapor inside or outside the substance?

    Thanks for the immense patience.
     
  25. Nov 16, 2016 #24

    Ygggdrasil

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    A note on terminology. Vapor pressure does not mean pressure of the vapor. Vapor pressure is a specific, precisely defined, scientific term referring to the partial pressure of vapor over a liquid at equilibrium given a specific temperature. The vapor pressure of water at 25°C is ~24 torr. If the partial pressure of water in the atmosphere is less than 24 torr then water will tend to evaporate until pH2O rises to 24 torr. If the partial pressure of water in the atmosphere is greater than 24 torr, then vapor will condense and precipitate from the atmosphere until pH2O drops to 24 torr.

    To rephrase the second part of your statement in more precise language: The partial pressure of water in the container (not the vapor pressure!) will increase until dynamic equilibrium is reached. At this point the vapor is saturated. The partial pressure of water in the atmosphere at saturation is equal to 24 torr, the vapor pressure of water at room temperature.

    Regarding the first part of the statement and why condensed phases have vapor pressures: In a liquid (or solid), the fast molecules at the surface of the liquid will escape into the vapor phase. Similarly, slow molecules in the vapor phase that collide with the liquid phase will stick to the liquid and condense. The rate of evaporation from the liquid is fairly constant (for a given temperature). However, the rate at which vapor returns to the liquid phase depends on the number of molecules in the vapor phase (the more molecules there are, the more often one will collide with liquid-air interface). Thus, as the partial pressure of water in the atmosphere increases, the rate of condensation increases. At some point the partial pressure of water in the atmosphere will give a rate of condensation equal to the intrinsic rate of evaporation from the surface, at which point equilibrium is reached and there is no net exchange of material between the liquid and vapor phases. The partial pressure of vapor in the atmosphere that leads to this situation is called the vapor pressure.

    The vapor pressure is related to the thermodynamics of the liquid to gas transition, and is essentially an equilibrium constant. Water molecule have a lower chemical potential energy (enthalpy) in the liquid phase (because of interactions between water molecules in the liquid phase) but a higher entropy in the gas phase (molecules occupy a larger volume and have more freedom of movement). At lower temperatures, systems will prefer to minimize their potential energy, and at higher temperatures, systems will prefer to maximize their entropy. At some intermediate temperature, the trade off between minimizing enthalpy and maximizing entropy will be equal and this is when equilibrium occurs.
     
    Last edited: Nov 16, 2016
  26. Nov 17, 2016 #25

    rbelli1

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    You should take close note of the point Ygggdrasil made that in the equilibrium state the phase transition has not stopped. Equilibrium occurs when the transitions occur at equal rates.

    BoB
     
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