Monkeymaniac

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I've got a dumb question here.

How exactly is evaporation different from boiling? So water molecules can evaporate at room temperature, or separated from other chunk of water molecules and be freed to the atmosphere as gas molecules. Same thing happens when water boils, liquid water is converted to gas.

If this liquid->gas conversion can happen even at room temperature, then why does the state change graph like the one we saw millions of times in our chemistry classes (please see the link below) show that only liquid exists until 100'C is reached and then water->gas conversino occurs only during the phase change that happens at 100'C? Could anyone please clarify this?

 

rocuronium

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Any water molecule that enters the gaseous state must gain enough energy to reach the boiling point and make the conversion. Evaporation is not an exception to this rule.

Another way to look at this is by realizing that water in the air is indeed above 100`C. Whether water achieves the required energy at the boiling point (boiling) or below it (evaporation) is irrelevant.

Does that help?
 
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Monkeymaniac

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Any water molecule that enters the gaseous state must gain enough energy to reach the boiling point and make the conversion. Evaporation is not an exception to this rule.

Another way to look at this is by realizing that water in the air is indeed above 100`C. Whether water achieves the required energy at the boiling point (boiling) or below it (evaporation) is irrelevant.

Does that help?
Is water in the air really above 100'c? If I pour some water into a bowl, bring my face on top of it, wait for some of the water molecules to evaporate, I don't think I would burn my face :). So from my understanding, at room temperature, there will be a equilibrium of some set of water molecules gaining enough kinetic energy and escaping into the air and set of gaseous water molecules in the air losing kinetic energy and getting back into the body of the liquid water. And boiling point is when this equilibrum breaks and water molecules begin to evaporate at much faster rate than condensation. Is this right?
 
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rocuronium

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Is water in the air really above 100'c?
Is it possible to have water vapour at a temperature less than that (assuming standard conditions)? What would happen if it were?
 
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Monkeymaniac

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Yes, I think it's possible to have water vapor at temperature less than that. Even at STP, the atmostphere contains some modicum amount of water in it. The air we breathe as we speak contains water vapors. My question was, from the state graph, it seems that it would be impossible to have a gaseous form of water unless the system temperature is above or equal to 100'C. In other words, how can the formation of water vapor at room temperature (evaporation) be done when water is not boiling (temperature above or equal to 100'C)? which was my original question.
 

rocuronium

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I think that you may be confusing the difference between water vapour and suspended water molecules (such as in a mist).

If I take an atomizer filled with water and spray it into the air, that water isn't immediately a gas. In order for it to become a gas, it has to gain sufficient energy, just as a molecule of water in a boiling pot would have to.

This gets back to my point about water vapour being above 100`C. If I could convince you of this fact, it might make the answer to your question (why it seems that it would be impossible to have a gaseous form of water unless the system temperature is above or equal to 100'C) more obvious.
 
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Monkeymaniac

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No I wasn't confused about that. After doing some research, I found out that the end result of evaporation and boiling is the same. That the liquid form of water molecules turn into gaseous form of water molecules. The difference is that the evaporation is a surface phenomenon, where as boiling is a volume phenomenon. That is, molecules evaporate at the surface of the water body, since the molecules experience weaker attractive force than molecules within the body, ones surrounded by more molecules. But for boiling, molecules inside the body rise up to the surface and escape to the air in gageous forms.

Now going back to the original point. The reason why the graph implies that there aren't gaseous water molecules at temperature below 100'C is because the molecules may have enough kinetic energy to rise up to the air, but since the surrounding temperature is below 100'C, it quickly comes back to the original state, in its liquid form.

Another way to look at this is by realizing that water in the air is indeed above 100`C.
The surrounding temperature doesn't have to be 100'C for a liquid water molecule to become a gaseous molecule. At room temperature, water molecules at sufrace may not get enough kinetic energy due to high surrounding temperature (say 100'C like you said) but it could get enough kinetic energy from colliding with other molecules to become a gas molecule.
 
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wanderer

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Water vapor in the air is not necessarily 100 degrees, it can certainly be less. Think what would happen if this were the case, any water at the surface would have to be at or near 100 degrees, even if the system was at a much lower temperature. (The alternative explanation would be that water suddenly rises to 100 degrees upon evaporation, and this doesn't occur.)

Now going back to the original point. The reason why the graph implies that there aren't gaseous water molecules at temperature below 100'C is because the molecules may have enough kinetic energy to rise up to the air, but since the surrounding temperature is below 100'C, it quickly comes back to the original state, in its liquid form.
Perhaps in theory, but in the real world this doesn't exactly happen. (Ever notice what happens in chem lab if you leave a container of ether open?)
My guess is that the amount of heat absorbed due to evaporation is insignificant if the surface area is small and the time of the experiment is relatively short. Also once your sample has started boiling, it doesn't matter whether x amount of water was lost to evaporation or boiling.
The surrounding temperature doesn't have to be 100'C for a liquid water molecule to become a gaseous molecule. At room temperature, water molecules at sufrace may not get enough kinetic energy due to high surrounding temperature (say 100'C like you said) but it could get enough kinetic energy from colliding with other molecules to become a gas molecule.
 

Doodl3s

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Water vapor is CERTAINLY BELOW 100 degrees Celsius...

Whats happening in evaporation is that some random water molecules have enough energy to break free from the liquid state. Again, water does NOT need to be 100 degrees to do this. 100 Degrees is only the point at which water's vapor pressure becomes 1 atm. Thats it!

But from 0 degrees to 100 degrees, water has a varying vapor pressure from 0 to 1 atm.... So, here's when evaporation occurs:

When the partial pressure of water vapor above water is LESS than the vapor pressure of water at that temperature, water will evaporate until the the partial pressure of the water vapor is equal to its vapor pressure....

So, the higher the temperature of water, the higher its vapor pressure and therefore the water wants the partial pressure of water vapor above it to equal the vapor pressure...

this is why if you leave a pot of water out, it will continue to evaporate. What's happening is, water keeps evaporating to raise the partial pressure of water vapor above it. However, in an open room, this water vapor diffuses away from above the solution and the partial pressure of water vapor above the liquid keeps dropping below its water vapor...

THERE! :D

EDIT: TOo long didnt read? Evaporation occurs when the partial pressure of water vapor above the liquid is less than its vapor pressure.
 

rocuronium

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Water vapor is CERTAINLY BELOW 100 degrees Celsius...

Whats happening in evaporation is that some random water molecules have enough energy to break free from the liquid state. Again, water does NOT need to be 100 degrees to do this. 100 Degrees is only the point at which water's vapor pressure becomes 1 atm. Thats it!

But from 0 degrees to 100 degrees, water has a varying vapor pressure from 0 to 1 atm.... So, here's when evaporation occurs:

When the partial pressure of water vapor above water is LESS than the vapor pressure of water at that temperature, water will evaporate until the the partial pressure of the water vapor is equal to its vapor pressure....

So, the higher the temperature of water, the higher its vapor pressure and therefore the water wants the partial pressure of water vapor above it to equal the vapor pressure...

this is why if you leave a pot of water out, it will continue to evaporate. What's happening is, water keeps evaporating to raise the partial pressure of water vapor above it. However, in an open room, this water vapor diffuses away from above the solution and the partial pressure of water vapor above the liquid keeps dropping below its water vapor...
This explanation seems to make sense to me, however I still am confused on a few points.

If we define the temperature of a gas as the average kinetic energy of its molecules, how is it that when water evaporates (gains enough kinetic energy to escape) the temperature of those molecules does not increase? Do the escaped molecules not have a higher kinetic energy than those left behind in the liquid? Note that I never did say that the water (or system) had to be 100 degrees - I think that may be people putting words in my mouth.

I had this conversation with my physics professor a few years back. It seems that either I (clearly) didn't understand what he was saying, or he wasn't right. I'm fine with it either way :D
 
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rocuronium

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I came across this in an article on Wikipedia. It seems to support my previous point.

"In terms of intermolecular interactions, the boiling point represents the point at which the liquid molecules possess enough thermal energy to overcome the various intermolecular attractions binding the molecules as liquid (eg. dipole-dipole attraction, instantaneous-dipole induced-dipole attractions, and hydrogen bonds) and therefore incur a phase change into the next phase (gas)."
 

Doodl3s

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This explanation seems to make sense to me, however I still am confused on a few points.

If we define the temperature of a gas as the average kinetic energy of its molecules, how is it that when water evaporates (gains enough kinetic energy to escape) the temperature of those molecules does not increase? Do the escaped molecules not have a higher kinetic energy than those left behind in the liquid? Note that I never did say that the water (or system) had to be 100 degrees - I think that may be people putting words in my mouth.

I had this conversation with my physics professor a few years back. It seems that either I (clearly) didn't understand what he was saying, or he wasn't right. I'm fine with it either way :D
The average kinetic energy of the molecule does not need to change (although it may). The energy it absorbs is just enough to break the intermolecular bonds. Remember, energy always required to break bonds. So, the molecule may still vibrate with the same KE/temp.... but the energy went directly to overcoming bond strength.

At the boiling point, ALL the molecules have enough energy to form into a gas. This is why at boiling point you see bubbles of water vapor rise from seemingly random points throughout the volume of the liquid. Molecules on the surface have MUCH less bonds to overcome. (only molecules below them are pulling them down, as opposed to molecules all around them as in a volume of water.)