When heating up a liquid, the temperature will rise until it begins to boil, at which point the temperature will remain constant until all the liquid boils away. One thing I’ve never understood is what property of liquid allows it to hold a constant temperature during boiling. The molecules of the liquid are moving around at all different velocities. The temperature is just the average of all that movement. How is all that seemingly random movement able to average out to a constant temperature when the liquid is boiling? What force is preventing the liquid’s temperature from continuing to rise past the boiling point as heat is continually added?
Molecules seem able to leave the liquid while it’s being heated up. Mist can often be seen rising above the liquid as heat is applied. As the liquid gets hotter, more and more mist rises up. Why doesn’t the mist just keep getting thicker and thicker and the liquid keep getting hotter and hotter as the heat is applied? Why do the molecules of the liquid seem to say to each other “Everyone! The temp is at the boiling point. No more moving around faster until every one of us has left.”
It’s easier for me to understand how a solid can keep a certain temperature when it reaches the melting point. With the crystalline structure of a solid, the force of a wiggling molecule will be transferred to the rest of the solid through the connections between the molecules and average out. But the molecules in the liquid are not connected to each other. They’re bouncing around at all varied speeds. If a rogue molecule in a boiling liquid gets some extra energy to where the temperature would rise, how is the liquid able to moderate that extra energy such that the temperature instead remains constant?
The molecules in a liquid are certainly connected to each other. It’s just that the connections are weaker and there is more wiggling going on. The energy being put into the liquid is being used to break intermolecular interactions (connections) allowing them to go into the gaseous state.
Because the boiling point is the definition of what temperature a given liquid turns into a gas at a given pressure. If the liquid is hotter than the boiling point then it becomes a gas and if a lower temperature it remains a liquid.
Then there is the latent heat of the liquid. It is the additional energy needed to vaporize the liquid.
If you add more heat to a liquid then it will boil faster (you will convert the water to steam more quickly).
You have the gist of the answer already, the temperature of the liquid is a measure of the average energy of all of the molecules, and each molecule is “bouncing” off of many others many times per second. One of your extra hot molecules, if I can use that term, is either going to undergo a collision with a more sedate molecule and transfer some of it’s energy, or it’s going to reach the surface of the liquid and have the energy to break through the surface tension and escape. In essence, the boiling point of a liquid is the temperature at which the average energy of each molecule is sufficient that the molecule is able to break through the surface when it reaches it. To boil all of the liquid away, you have to keep providing heat, but the temperature of the liquid will not rise because it is as hot as it can get. If you make it more difficult for the molecules to escape, the energy required to break through the surface tension rises and so the boiling point increases.
But you don’t need to boil a liquid for it to turn to gas. If you spill water on the floor, it dries out. At any temperature some of the molecules are always moving fast enough to turn to gas, so liquid is constantly turning to gas at all temperatures by evaporation at the surface.
So this doesn’t answer OP’s question: as you heat a liquid, why is there a temperature at which suddenly 100% of the energy input (however large) goes into phase change from liquid to gas, and the temperature of the remaining liquid cannot be increased further?
Boiling point is when the vapor pressure of the liquid equals ambient pressure. Consider the horizontal surface of a beaker of liquid. Ambient pressure pushes down. But the fastest liquid molecules are constantly escaping upward into vapor. You can think of this vapor pressure as upward pressure that reduces the net downward pressure on the surface of the liquid. As the liquid temperature increases, the average speed of all the liquid molecules increases, so there are more molecules moving fast enough to escape at the surface, and the vapor pressure increases. When the vapor pressure equals the ambient pressure, there is now no net downward pressure at all at the surface of the liquid. With no external pressure, there is nothing to stop liquid molecules turning to vapor anywhere throughout the entire volume of the liquid.
So boiling point is when vapor pressure equals ambient pressure, and thus when liquid can turn to gas not only through limited evaporation at the surface of the liquid, but to an unlimited extent (subject to continued energy input to maintain vapor pressure) anywhere throughout the entire volume of the liquid.
Evaporation and boiling are different things. They work differently. Evaporation is not slow motion boiling. I do not think the OP was asking about evaporation.
I would be inclined to say that the difference between boiling and evaporation is very slight.
The linked article only just glosses over a critical point - that these are systems with various equilibriums in play.
At the surface of a body of water there is a constant stream of molecules with enough energy to break free doing just that, and other molecules that find themselves banged around and heading back into the body of water. If you let things sit contained in still air this will reach an equilibrium. If there is a way for the water molecules to move far away from the body of water they will take energy with them, and the water will cool. Which is why blowing on a water cools it, you are using a mechanism to take the water molecules away so they can’t recombine with the liquid water, so they carry away the energy they needed to break free of the surface. The water is evaporating.
The OP asked about why water won’t go above boiling point when it is boiling. The key is that it takes non-zero energy to get a water molecule away from the surface. The difference between boiling and evaporation is only that there isn’t enough pressure inside the body of water to stop molecules making that break, whereas at the surface of a body of water there is an equilibrium of molecules breaking free and returning. When water starts to boil it can create internal bubbles, and these bubbles provide new surfaces that water molecules can break free into. There is still an exchange of molecules at the surface, but with enough energy on tap, there is energy to allow more molecules to enter that bubble than drop out of it. That is all there is to boiling - enough energy to support the creation of internal bubbles.
The question about constant temperature falls out from here. It takes energy to break free of the surface. Every molecule that does so takes that energy out of the liquid water. As those molecules escape from the water they take that energy away forever. So you end up with an equilibrium. More energy in goes into more molecules with enough energy to escape, and the more more escape. To a good approximation the water can’t get any hotter as the escaping molecules are taking every bit of energy added. To allow water to continue boiling this energy must be replenished - otherwise it near instantly stops boiling.
Boiling is just the edge case of evaporation. It looks different because it demands continued input of energy to maintain, whereas evaporation just cools the liquid if there is no other input of energy.
This brings up a few evil edge cases. If you add more pressure, the energy needed in any molecule to stay free goes up. You can reach a point where under huge amounts of pressure the water contain a lot of energy. If there is a near instantaneous removal of that containing pressure a lot of water molecules will suddenly find themselves with enough energy to break free. The results can be catastrophic. This is a BLEVE - boiling liquid explosive event.
I think maybe the disconnect is that it’s easy to imagine that liquid, being heated in a container, is a discrete entity; unless you think about what’s happening to (and between) individual molecules in the liquid, it’s possible to wonder why you can’t just suddenly dump a load of heat into the liquid and trick the liquid, as a bulk object, into momentarily getting hotter than its own boiling point, before it turns to gas; as though you can sneak up on the boiling point and over-run it before the evaporation catches up or something.
(I don’t think that myself, but I think that’s sort of what’s happening sometimes when people try to intuit the behaviour)
Both involve liquid turning to gas. So I was just noting that your definition of boiling…
…was not correct.
I don’t think it’s useful to say that evaporation and boiling “work differently” as though they are unrelated. Certainly something qualitatively different happens at boiling point, and in order to understand what that is (and to answer OP’s question) you need the concept of vapor pressure.
I’m not sure exactly what you mean. The concept of latent heat is relevant to any phase change, including evaporation below boiling point. It explains how we cool ourselves through the evaporation of sweat.
The concept of latent heat is not sufficient to explain the process of boiling, but it is obviously relevant to where the energy is going when you continue put energy into a liquid at boiling point. The temperature does not rise further, but the liquid does not all instantly vaporize.
Isn’t that what this thread is all about? Not explaining the whole process of boiling, but specifically why the temperature remains constant as water does so?
Anyway, that’s what I’d have said. Latent heat of vaporization.
That’s telling you where the energy is going, but it doesn’t explain why.
If you heat a liquid below boiling point, part of the energy input goes to raise the temperature of the liquid, part goes to effecting phase change (latent heat of vaporization).
When you reach boiling point, if you continue heating now all of the energy goes into effecting phase change. OP’s question is why that should be so. Why can’t you raise the temperature further?
The explanation requires an understanding of vapor pressure. The (only) thing that defines boiling point is when vapor pressure equals ambient pressure.
The discussion of the vapor pressure has made it much clearer. It’s the atmosphere squeezing on the liquid that keeps it in the liquid state. With enough energy, the liquid gets to a point where the vapor can overcome that pressure. It’s clearer now that at the boiling point, any additional energy will go into creating the vapor state since the atmosphere is no longer able to exert enough pressure to prevent bubbles from forming. I always knew that the boiling point varied with pressure, but I don’t think I realized how much the atmosphere played a part. I guess the liquid is always trying to turn into a gas, but the squeezing force of the atmosphere prevents that from happening below a certain temperature.
It’s easier for me to understand the melting point since that seems to be a constant and independent value. It seems to be an inherent property of the material and is not dependent on external factors like the atmosphere.
Yes, I know all this since I teach general chemistry. You are correct about vapor pressure, etc. However, it looks like the OP was asking why temperature doesn’t rise when you continue to heat at the boiling point of a liquid at atmospheric pressure. The answer is demonstrated in this heating curve of water:
As you add heat energy, the temperature doesn’t rise because you are putting in energy to break intermolecular interactions throughout the liquid such that the liquid disappears. As @moes_lotion states, the average energy of all the liquid molecules is sufficient to escape the liquid state. The name for this is the latent heat of vaporization (11.3 Phase Change and Latent Heat - Physics | OpenStax)
This is different than evaporation where there is always a dynamic equilibrium between the gaseous and liquid states on the surface of the liquid at a particular temperature. You keep losing liquid because the vapor escapes into the surroundings. However, in a closed system, once you pressure reaches vapor pressure, you would no longer have an increase evaporation at a fixed temperature. It reaches an equilibrium between the gas and liquid states.
The process of breaking intermolecular interactions and molecular movement is the same in both boiling and evaporation. It’s just that in boiling, it’s happening all at once and that requires the latent heat of vaporization to do it.
No, this is not a good explanation of what boiling is and how it is different from just evaporation.
This is getting repetitive, but the concept of latent heat is just as relevant to evaporation. It explains how we cool ourselves by sweating.
Again:
When we heat a liquid below boing point, PART of the energy goes to effecting phase change (latent heat of vaporization), PART of the energy goes to raising the temperature.
But if we continue to heat at boiling point, ALL of the energy goes to effecting phase change, and the temperature does not rise further. OP asked - WHY is that so? And understanding that requires the concept of vapor pressure.
I never said that the concept of latent heat isn’t relevant to evaporation. In fact, I said the exact same mechanism is occurring on the molecular level with evaporation as it does with boiling.
But both enthalpy of vaporization (latent heat of vaporization) and vapor pressure are dependent on the attractive forces between molecules or particles. That’s why water with its hydrogen bonding network has a higher heat of vaporization, a higher boiling point, and lower vapor pressure (at a given temperature) than methane which interacts with each other through dispersion forces only. As you add heat, attractive forces between molecules are overcome and they become a gas, increasing the vapor pressure. Once the vapor pressure equals the atmospheric pressure, the whole system reaches the boiling point.
So attractive forces explain both heat of vaporization and vapor pressure. In fact, the OP (@filmore) was on the right track in his last paragraph except he was only applying it to solids. If he applied his same scenario to liquids, he would understand why the temperature doesn’t increase at the boiling point. Liquid molecules are indeed “connected”. It’s just that the connections are different in solids. Solids have a more stable and ordered network of connections with a higher density (except water) than liquids.
