Question

According to the phase diagram, water is supposed to be a liquid at 1 atmosphere and room temperature. Why then did the water disappear from my glass after I left it on the table overnight? Shouldn't the water have stayed in liquid state?

Answer 1

No the water you have kept is not at room temperature than if it evaporated :) did you measure the temperature :)

Answer 2

evaporation happens, even at low temperatures. slowly, all the water in the glass will evaporate without changing temp

Answer 3

it was at room temperature but it was a dry night. @JFraser: good point, but wasn't it supposed to be in liquid state only? why would there be evaporation when the water in pure liquid state?

Answer 4

in any sample of liquid, there will always be a small number of molecules that possess the energy needed to leave the liquid phase and enter the gas phase. That's what evaporation is. it takes a long time to notice any change, and there's no temperature change needed, but it's all evaporation.

Answer 5

i agree with what you are saying, it was evaporation. but why does the phase diagram say it's a pure liquid and not a liquid-gas equilibrium, as implied by your description of evaporation?

Answer 6

the water absorbs heat energy from the surroundings until it gathers enough energy to evaporate.....as JFraser said, evaporation can occur even at very low temperatures but just that it would take a very long time.......

Answer 7

@kausarsalley : you are right, but evaporation implies phase co-existence, which the phase diagram denies.

when the TA points to the phase diagram at 1 atm and room temp and say that indicates water is a pure liquid, something is wrong.

specifically, a *open* glass of water cannot be described by the point at 1 atm and 25 deg Celsius. the situation is a little more complicated.

Answer 8

even at room temperature, some water molecules are able to escape from the pool.

Answer 9

@nincompoop yes, but still not answering my question, why does the phase diagram say, "no, it's a pure liquid".

Answer 10

we can calculate if evaporation is possible using q=mCΔT
where q is your heat, m is your mass and change in temp

Answer 11

how much water?

Answer 12

you have to realize that even at freezing point, evaporation still happens. now this may seem a bit odd, but it is a fact of nature.

Answer 13

@nincompoop i understand what you are saying, but we are still ignoring the info on the phase diagram. can the phase diagram be used to give a reasonable explanation for what happened?

Answer 14

I don't see where your phase diagram is.
I am not an expert at this yet, but I am also thinking along the possibility that because what you have is an open-system and that the temperature in the air fluctuates as the water absorb heat. also thinking how much would it require to break a hydrogen bond, that would cause the evaporation to occur as well.

Answer 15

post your data and we'll see if we can have a reasonable explanation based on calculation and varying factors.

Answer 16

@nincompoop no need to calculate, and no need for data, except for a phase diagram for water, which is in every gen-chem textbook

Answer 17

@nincompoop good point, temperature over the water might drop a little because of heat absorption, but the room is a heat reservoir and i think the temperature will go back to its initial value. so the evaporation occurs at constant, room temp.

Answer 18

yeah, your question made me go through my old lectures to see if this question was ever raised at all…

there are a few good chemistry students here like @abb0t and @agreene that might have better explanation.

Answer 19

So, I'll be honest, I didnt read all the responses, but they seemed to be on the right track.

Also, I doubt you had a sample of pure H2O in your glass. The other issues that could contribute would be changes in ATM (since I'm assuming your room isnt back pressured to 1 ATM constant), and of course the good ole' Copenhagen interpretation of QM

Answer 20

i agree with @agreene

Answer 21

i don't mean to be disagreeable. while the previous comments raise good points, but they are NOT hitting the phase diagram where it hurts, in my silly opinion anyway.

Answer 22

equilibrium implies a closed container, where the vapor can re-enter the liquid phase at the same rate as the liquid evaporates. The glass on the counter isn't a closed container.

Is it possible to think of the term "pure liquid" more as a "pure SUBSTANCE that HAPPENS to be a liquid"?

Answer 23

@JFraser: good point, it's more like a pure substance in liquid state, and there's definitely stuff dissolved. incidentally, the solutes should lower the vapor pressure, slowing down (instead of speeding up) the evaporation rate (the forward rate in the water(l) <-> water(g) equilibrium.

good point: i agree, equilibrium implies that the vapor is condensing at the same rate as the liquid evaporating. this seems to suggest that the glass must be covered, and covering it would have allowed an equilibrium to establish itself, and i wouldn't have lost the water to the atmosphere.

in light of this line of thought, let me rephrase the question: how does the water somehow know that there is covering? if i had placed the glass of water in a sealed container that's so large that it reaches to the sky, would the water have stayed? How is it different from covering the glass in the usual way?

Answer 24

What?

Answer 25

the earth's atmosphere is essentially the same thing.