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Phase‱Transitions:‱How‱does‱water‱evaporate‱below‱its‱boiling‱point?how‱boiling‱point‱of‱a‱component‱reduces‱when‱partial‱pressure‱reduces?

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Phase‱Transitions Water Chemistry Science Physics

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I will try to combine the correct aspects of the answers and comments posted so far, andelaborate a bit.

1. Evaporation. A phase transition, in this case from liquid to gas, occurs when conditions(temperature and pressure) are such there is a net reduction of free energy to take amolecule out of the liquid and put it in the gas. At an interface between a liquid and a gas,a water molecule can randomly fly out of the liquid and join the gas, or drop out of the gasand join the liquid. This gas/liquid system is at equilibrium if the change in free energy isthe same for a molecule to go in either direction. In this case, the molecules aretransferred in both directions at the same rate, and there is no net change. If you have asituation that is not in equilibrium, then molecules will be transferred from the liquid to thegas or the gas to the liquid preferentially, and this will continue until the pressure ortemperature have changed to achieve equilibrium, or until one phase is completely gone. In many common examples, equilibrium is never reached. For example, at roomtemperature, the equilibrium point for water will be at a pressure of about 20 Torr. If therelative humidity is less than 100%, then the partial pressure of water in the air is less than20 Torr. So, if you put a glass of water in such an atmosphere, the water molecules willtend to leave the liquid and join the gas. Since the atmosphere is much larger than theglass of water, this evaporation will not noticeably change the partial pressure of water inthe atmosphere, so the water will continue to evaporate until it is entirely converted intogas. On the other hand, if you put a cap on the glass, the evaporation would only continueuntil the partial pressure of water went up to 20 Torr.

OK, so when does boiling occur? The difference between boiling and evaporation is that inboiling, bubbles of gas form, within or at the surface of the liquid. The formation of thesebubbles takes extra energy. So boiling takes place when there is a net reduction in freeenergy to take a molecule out of the liquid and put in the gas while expanding a bubble. For bubble formation, first a microscopic bubble has to randomly form (nucleate), andthen it has to grow. The energy of a bubble comes from surface tension (the gas/liquidinterface has higher energy than a liquid/liquid "interface") and from the need to expandagainst pressure (the bubble has to lift the water above it). As discussed above, at roomtemperature the equilibrium point between liquid and gaseous water is at 20 Torr. If abubble were to nucleate at room temperature, gas would only be added to it up to apressure of 20 Torr, which is not enough to grow the bubble against atmospheric pressureof 760 Torr. So this does not happen, and water does not boil at room temperature andatmospheric pressure. It only evaporates. We see that for boiling to occur, the equilibriumvapor pressure of the gaseous water must be at least 760 Torr (that is, the water vaporpressure must be at least atmospheric pressure). For water at 1 atm, this point is reachedat 100 C. At this point (or just above), bubbles can grow near the surface of liquid waterbecause the pressure within the bubble is greater than the pressure trying to collapse thebubble. Boiling may still not occur however, because the bubbles must nucleate before theycan grow. Nucleation is complicated, and depends on the particular situation (e.g. waterin a smooth glass container may not boil above 100 C).

2. Passage around the critical point. The answer along these lines is also correct. At

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sufficiently high pressure and temperature the distinction between liquid and gas breaksdown, and there is no phase transition between the two, hence no evaporation or boiling.By starting as a liquid, going into the supercritical regime, then coming back into thesubcritical regime in the gas phase, you avoid boiling.

3. Sublimation. Although maybe the asker was just thinking of liquid to gas transitions,solid to gas transitions can also occur, in this case, when there is a net reduction in the freeenergy from taking a molecule out the solid and putting it into a gas. For water, this takesplace at pressure below about 5 Torr.

Updated‱12‱Jul‱•‱View‱Upvotes‱•‱Asked‱to‱answer‱by‱Yishan‱Wong

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Originally‱Answered:‱How‱does‱water‱evaporate‱below‱its‱boiling‱point?

There's a very simple piece of the puzzle missing from the other answers. Liquid wateralways has some pressure of water vapor above it ­­ the higher the temperature, the morepressure of vapor you get above the water. The "boiling point" is not some magicaltemperature where the properties of water change, it is merely the temperature at whichthe vapor pressure produced by the water is higher than local atmospheric pressure. If thewater vapor pushes out harder than the air pushes it in, the vapor is able to force its wayaway from the surface of the liquid and escape. Instead of needing to diffuse away or beblown by external wind, the vapor can flow away at a rapid rate under its own power. In aroiling boil, the vapor pressure at the hot surface is higher than the hydrostatic pressure inthe pot of water plus atmospheric pressure, so the most energetic water molecules haveenough force to form gas bubbles below the surface.

Boiling is just normal evaporation that is able to occur much more rapidly because theevaporation rate is no longer limited by local air pressure. Lower air pressure means alower boiling point. In Denver, it takes longer to boil an egg than in Houston. In a pressurecooker, water boils at a higher temperature because the container prevents the vapor fromescaping. And in a vacuum, water will boil away into a cloud of ice crystals almostinstantaneously, because the boiling point is actually below the freezing point.

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Originally‱Answered:‱How‱does‱water‱evaporate‱below‱its‱boiling‱point?

All substances are composed of molecules. For a given body—say, a glass of water—thetemperature of the whole is really just the average molecular energy of the molecules itcomprises. But the average is only an average, and at a given moment, the speeds ofindividual molecules vary, with a generally bell­shaped distribution. That means some aretravelling much slower, and others much faster. Since the thermal energy is in the form ofrandom molecular movements, with everything smashing and crashing around, for anygiven molecule, the speed with which it travels changes with each new collision.

Now water is held together not only by the simple weight of gravity holding it down, but bya collection of weak bonds called hydrogen bonds. This is because of the uniquely "bent"shape of the water molecule.

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Source The effect is that a hydrogen atom from one molecule has a tendency to "stick" to theoxygen atom of an adjacent molecule. This is what makes water remain in a liquid state ata much higher temperature than you would expect by its molecular weight alone.

Now what happens is that at the surface of the liquid, a given molecule might get enoughenergy to simply break free of the hydrogen bonds that keep it in the liquid. When it doesthis, it becomes a vapor, and it starts behaving as a gas—the molecules with which it nowinteracts are traveling with a much higher velocity, and they do not collide at speeds lowenough (in general) for hydrogen bonding to resume. The exception to that last point is inthe formation of fog, clouds, and ultimately rain.

All the while, however, there are other water molecules in the gas phase (by phase here, Imean the designation of either liquid, solid, gas, etc.) directly above the surface of theliquid phase. For the converse reason, the distribution of speeds in the gas phase willinclude a few molecules with such a *low* speed, that when it collides with the surface ofthe liquid, it loses the energy it would need to remain as a vapor, and succumbs to thehydrogen bonds that make it now a part of the liquid.

The net effect here is that the liquid will gain some, and it will lose some. The net directiondepends on whether the air is saturated with H2O or not.

Usually, it is not saturated, so you expect a puddle to dry up over time, even if thetemperatre of the water never gets near boiling. But the drier the air is—that is to say,more arid, less humid—the less the loss of water molecules from liquid to gas will be offsetby transfers from gas to liquid. So after a rainstorm, the air is close to saturated withwater, and puddles dry very slowly. This is also why spilled water will dry much quicker inPhonix or Las Vegas, even if it's not a hot day. And it's the reason why blowing on thingsdries it out—you're forcing the saturated vapor at the surface away, and pushing in airthat's much drier and can remove more moisture. The converse can also happen, but we don't normally notice it in the liquid itself. If youhave a cup of iced beverage with an open top, water from the air around it will crash intothe surface, lose the energy needed to escape, and increase the overal amount of liquid inthe cup. But it's a samll amount generally not noticed. What we *do* tend to notice iswhen this happens to the cold, dry exterior of the cup itself. This condenses onto the cupan makes the outside of it wet; that water on the outside of the cup is from the air around

30/11/2015 (2) How does water evaporate below its boiling point? ­ Quora

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it condensing, rather than from the beverage in the cup itself. You can also have a situation where the rate of change from liquid to gas is exactly thesame as the rate of change from gas to liquid. This is called equilibrium, and there's no netchange in the liquid volume. The best example of this would be on a closed bottle or jar. The air in the container soon becomes saturated, and so long as the jar remains closed andthere's no great change in temperature for the jar, the liquid volume remains fixed.

Written‱31‱Dec‱2012‱•‱View‱Upvotes

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Originally‱Answered:‱Why‱does‱water‱become‱steam‱before‱100°C?

Others answers have already addressed the issue of difference between boiling andevaporation which explains the vapors(which is not steam) of water at temperatures below100C.

Also note that Boiling point Temperature of any liquid is a function of pressure too. Theboiling point temperature decreases if surrounding pressure is low. So at high altitudes(orany location where pressure is less than 1 atm) water boils at lower temperature.Reversehappens if surrounding pressure is high.

Also boiling point temperature of a liquid can be increased by adding a non­volatile solutein the liquid. So if you have sea water,you can decrease its boiling point by purifying it.

Hope you know that Boiling point(BP) is the temperature at which vapor pressure of theliquid is equal to the surrounding pressure(which if 1 atm then its called normal BP­100Cin case of pure water).There are other dimensions to it if want to go in bit depth:­>A mixture of two volatile ideal liquids(which follow Raoult's law) has a range ofboiling points which range from BP of less volatile to BP of more volatile liquid.Depending on the composition of the mixture you can have a corresponding BP of themixture.

>If the mixture behaves Non­ideally then there are other possibilities1. Positive deviations from Raoult's LawThe vapor pressure of this mixture is always higher than you would expect from an idealmixture.The fact that the vapor pressure is higher than ideal in these mixtures means thatmolecules are breaking away more easily than they do in the pure liquids. That is becausethe intermolecular forces between molecules of A and B are less than they are in the pureliquids.You can see this when you mix the liquids. Less heat is evolved when the new attractionsare set up than was absorbed to break the original ones. Heat will therefore be absorbedwhen the liquids mix. The enthalpy change of mixing is endothermic. The classic exampleof a mixture of this kind is ethanol and water.

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30/11/2015 (2) How does water evaporate below its boiling point? ­ Quora

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If the deviations are quite large,then a maximum boiling point AZEOTROPE (Amixture of two or more liquids whose proportions cannot be altered by simple distillation)is formed. Mixture is then known as a constant boiling mixture or an azeotropic mixture oran azeotrope. The point where the two curves touch in the figure below gives theazeotropic composition.

2.Negative deviations from Raoult's LawThe vapor pressures which are less than would be expected by Raoult's Law.These are cases where the molecules break away from the mixture less easily than they dofrom the pure liquids. New stronger forces must exist in the mixture than in the originalliquids. You can recognize this happening because heat is evolved when you mix theliquids ­ more heat is given out when the new stronger bonds are made than was used inbreaking the original weaker ones.

If the negative deviations are quite large,then a minimum boiling point Azeotrope isformed as shown below.

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>At last the above is only for binary mixtures.You can imagine the complexity in a realfluid with all sorts of solutes,surrounding pressure and of course multicomponents.

Note­There are n number of reasons for your answer.Choose yours according to what suitsyour depth of knowledge.

Written‱11‱Feb‱•‱View‱Upvotes‱•‱Asked‱to‱answer‱by‱Malan‱Balasubramaniam

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Originally‱Answered:‱How‱does‱water‱evaporate‱below‱its‱boiling‱point?

A fundamental rule of Physics is that when temperature is above 0K (­273,15 °C , ­459,67°F) there is some movement in there. For liquids it means molecules are constantly ejectedfrom the liquid's surface to the surrounding vapour environment. So, next to liquid water,there is liquid vapour, actually there is a "concentration" of water molecule in the air,called vapor pressure, which is the equilibrium (the best state). (File:Water vapor pressuregraph.jpg ). The hotter, the more molecule can be in the air.

If you're in a close room like your shower, water will evaporates until it reaches this vapourpressure concentration and then it won't evaporate anymore and actually it willcondensate (turn liquid) on surfaces like mirrors or walls as part of the dynamicequilibrium. When you open the door, vapour pressure will immediately lower as thevolume expands. Then water drop will evaporate to try to reach vapour pressure. Most ofthe time volume is too big (big house or even the earth atmosphere if you open the windowof your house) so water drops will totally evaporate without being able to reach vapourpressure.

Boiling point only means that there is so much agitation (keep in mind temperature =agitation) liquid water can't exist and turns into vapour even inside a volume of liquidvapour (it forms bubbles of water vapour)

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Vaporization of a liquid is a phase transition from the liquid phase to vapor. Evaporation of a liquid is a phenomenon that occurs from the surface of a liquid into agaseous phase that is not saturated with the evaporating substance. Water becoming steam at 100 degree centigrade is vaporization, while water in ponds getsevaporated.A substance vaporizes only at a particular temperature at a particular pressure (called

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boiling point), while a liquid can evaporate at any temperature.

Difference Between Evaporation and Vaporization

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