posted
I found out in the course of my work this week that saturated steam at 450 degrees F has more enthalpy (a measure of the energy stored in the system) than the same mass of saturated steam when it is at 500 degrees F. There's about 2.9 BTUs difference per pound mass. So what that means is that if you had a pound of sat. steam at 450 F, you should be able to let it do work, like move a piston to run an engine or something, and then it would be in the lower energy state of 500 F (at a much higher pressure too). This is totally counterintuitive to me, and I have been puzzling how it can be true for a couple of days.
Can any of you chemical / thermodynamical types explain to me what is going on?
My best guess so far has to do with U, the internal energy of the system, which I am told is made up of stuff like molecular rotation, springiness in molecular bonds, and so on, as well as temperature. The equation is H = U + PV, where H is enthalpy, U = internal energy, P = Pressure and V = Volume of the system. I think that the enthalpy decrease with increasing temperature at saturated conditions can't happen with ideal gases, and this weird anomaly is only true because steam isn't an ideal gas.
Am I even on the right track? Can anyone explain it to me? Feel free to verify with any steam tables. It's certainly true, odd though it seems.
posted
I don't have anything to contribute, really, but the physical fact you're talking about is fascinating, and completely counter-intuitive.
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posted
It reminds me of another weird anomaly about water, that it's less dense as a solid than as a liquid. I can't think of any other substance for which that's true, though there must be some.
The fact that ice floats is very important to life on earth, since rivers, lakes, and oceans get a blanket of ice that insulates them from the cold air, instead of ice forming on the surface then sinking immediately to the bottom, which would mean bodies of water would freeze solid in the winter.
I wonder how many other very odd properties water has?
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posted
Tatiana , maybe I am getting this all wrong (it is late ) but your numbers to say the opposite thing:
450 F : H = 1000.6 kJ/kg 500 F : H = 1134.4 kJ/kg
Also, the numbers seem to check out. Since the enthalpy is given in J/kg, it is enthalpy per unit mass. So, H = U + P/rho, where rho is the density or H = U + Pv where v is the specific volume.
So, the number from your tables give
450 F: H = 996.92*10^3 + 29.103*10^5*0.0012132 = 1000.6kJ/kg (approx)
500 F works out too. Are you talking about some other numbers? The trend for the internal energy is also consistent
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posted
That's the saturated liquid. The saturated vapor does indeed have a higher enthalpy at 450.
I have no idea what the answer to your question is, Tatiana. Looks like it's both internal energy and greater volume, though. If you multiply pressure and volume for saturated vapor, it comes out that 450 degrees has a greater PV. I suppose there's some point where the pressure rises faster than the volume--it would be interesting to plot the PV curve for various temperatures. If I had my thermo textbook at home, I could probably do it.
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posted
The boilers at the power plant I work at are classed as “supercritical.” That means that the water doesn’t actually boil. The boilers do not have drums just tubes. We run the water pressure in the boiler tubes up to over 3500 psi and at some point (nobody knows just where or when) the water changes from a liquid state to steam. And the steam is at the same density as water. The water wall pressure runs about 790 degrees F. at about 3900 psi. Then it flows through the superheat sections of the boiler and picks up even more heat before hitting the turbines.
I just mention this because it is another wild and whacky thing about water. I don’t even begin to understand all the physics involved nor the math. I just know which buttons to push and which valves to operate and when to run away.
I do know, however, why ice is less dense than water. That’s so the folks in Minnesota can go ice fishing.
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posted
So, um...if you allow 500 degree water to cool to 450 you can get more work out of it? If you, say, use 500 degree water to spin a turbine, will it drop to 450 degrees, at which point you can use its excess energy to bring it back up to 500 degrees, to then spin a turbine, etc., etc.?
Should I buy stock, or start an Internet scam?
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posted
Dasa, look where it says "Enthalpy Saturated Liquid (h)" then two lines down from that it says "Saturated Vapor (hg)". That's the line I mean. 4th line from the bottom. It's 1205.6 and 1202.5 BTU/pound mass at 450 and 500 degrees, respectively. (I apologize for using the cruddy U.S. units system, but that's what we use.)
Samuel Bush, you guys in fossil probably run into all the same steam properties as we do in nuclear. Have you ever heard of this before? I think we may run lower temperatures, though. Is that right? The steam in our secondary system doesn't get above the critical point, (where steam and water have the same properties, and it's all one fluid, like you described). The range I've picked for my example is actually higher Temp and Pressure than it is at the outlet from our steam generators. I picked 450 F because it's the peak enthalpy, and it goes downhill from there.
"I just know which buttons to push and which valves to operate and when to run away."
Bob, that's just the thing. If you let 500 degree water cool to 400, (while keeping the pressure at the correct value to maintain it as 100% saturated steam), you'd have to put work IN to get it to do that, according to the steam tables. If you got work out of it, (let it do work) then it would get hotter and have a higher pressure than it did to begin with. Maybe you could use the hotter steam to heat up some saturated steam that's down in the reasonable range of the curve, and then extract its energy in useful work, though. So hold onto your seed capital. There are still possibilities.
posted
Thanks Shigosei and Tatiana for the correction. As Shigosei said, both the internal energy and Pv contribute to the drop.
U drops by about 5kJ/kg and Pv by about 2kJ/kg (sorry I just can't figure out the Standard units ). I can kind of understand how Pv could drop, but the internal energy drop really confuses me. Are there other substances which have lower internal energy when temperature increases?
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quote:The boilers at the power plant I work at are classed as “supercritical.” That means that the water doesn’t actually boil. The boilers do not have drums just tubes. We run the water pressure in the boiler tubes up to over 3500 psi and at some point (nobody knows just where or when) the water changes from a liquid state to steam. And the steam is at the same density as water.
This isn't quite accurate because about the critical pressure, there is not phase transition therefore the liquid and gas phases don't really exist -- there is only one phase. Also, this isn't a property that is unique to water. Most if not all liquids have a critical point.
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posted
Rabbit, right. But do you understand how saturated steam can have a lower enthalpy at a higher T and P above 450 F?
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quote:So what that means is that if you had a pound of sat. steam at 450 F, you should be able to let it do work, like move a piston to run an engine or something, and then it would be in the lower energy state of 500 F (at a much higher pressure too). .
This may not be true. You have to consider both the second law and the first law to determine if work can be done by a system.
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Some of the lower H (at 500 F) comes from lower PV and some from lower U, as you point out, Shigosei and Dasa. So while you're decreasing the volume, the pressure is not increasing correspondingly, but rather, it is increasing at a lower rate than the volume decreases. Is that what you are saying? That seems to be part of the answer. And the explanation of why that would be probably has something to do with hydrogen bonds, which seem to come into most of the anomalous properties of this extremely strange and most familiar molecule.
Part comes from internal energy, which is lower. My friend Paul the chemist (whom some of you may remember used to post here as "Paul") pointed out that the entropy is also less at the higher temperature. Meaning the randomness in the system must be decreasing with higher pressure and density faster than it's increasing from higher temperature.
I wonder if you actually did the experiment would the system spontaneously settle into a state at higher T and P once it passed through the inflection point? And if it did that, could you use that somehow to your advantage?
posted
Rabbit, that's true. I think thermo problems are often assumed to be adiabatic with quasi-equilibrium conditions during changes and all that stuff, which makes it come out different than the real world.
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posted
I couldn't find anything in a quick Google, but it's possible that there is a temperature somewhere in that range where a new vibration mode becomes available. That would absorb some energy by the virial theorem, and account for some of your U.
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posted
Y'know, I read that link, and it's very interesting; a lot of those more easily observable properties are indeed bizarre when compared to other substances, but most of us learn about them by experience at a young age and then kind of take them for granted most of the time. It's not until one starts categorizing and comparing the properties to those of similar substances that one realizes how strange they are.
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posted
OK, Here is the best way I can explain this apparent anomally.
Temperature is, essentially, a measure of the average translational energy (i.e. their velocity) of the molecules in a system. Internal energy includes not only the translational energy, but also the vibrational, rotational and chemical potential energy in the system.
In gas phase, the chemical potential energy generally doesn't change over a very wide range of temperatures and pressures so we tend to forget it exists.
There is nearly always a strong correlation between translational, rotational and vibrational energies for a given kind of molecule, although the relationship is different for different kinds molecules because of the different possible rotational and vibrational modes. This is why it take more energy to heat a mole of water vapor 1 C than to heat a mole of Argon 1 C. H2O has three atoms and it is nonlinear so it has more rotational and vibrational modes than Argon. When you add energy to Ar gas, all it can do is move faster because a single atom doesn't have any rotational or vibrational modes. In contrast, when you add energy to steam, only part of it goes to increasing the translational energy and the rest goes to increasing rotations and vibrations so you have to add more energy to the steam than to Ar to get the temperature to rise 1 degree.
At least that is the story at low pressures. When the pressures get high enough, it is no longer accurate to neglect the chemical potential energy. At 500 F and 630 psi (saturation conditions), the water molecules are close enough together for there to be significant electrostatic attraction between the water molecules. If the steam is then taken to 450 F and 422 psi (also saturation conditions), the rotational, translational and vibrational energies of the molecules will decrease. That is what the temperature is telling us. Because these energies are going down, our intuition tells us that we will have to take energy out of the system to achieve the decrease in temperature. But now we have to consider that in going from saturation at 500 F to saturation at 450 F, there is a signifcant increase in the specific volume of the gas. This means that we will need to put in energy in order move the attracting molecules further apart. In other words, the chemical potential energy at 450 F and 422 psi is greater than the chemical potential energy at 500 F and 630 psi.
What the internal energy numbers are telling us is that it takes more energy to move the molecules further away from each other than is released when the molecules slow down.
Water is somewhat unusual in this respect. Most of the gases we deal with commonly are nonpolar and so they tend to repel rather than attract until they are forced very close together. The strongly polar water molecules would be expect to attract each other over a much longer distance than typical chemicals that we deal with in gas phase.
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quote:Originally posted by King of Men: I couldn't find anything in a quick Google, but it's possible that there is a temperature somewhere in that range where a new vibration mode becomes available. That would absorb some energy by the virial theorem, and account for some of your U.
I considered this possibility but what we are dealing with here is sort of a negative specific heat. I way sort of becaue specific heats are general defined at either constant pressure or constant volume and never at constant %saturation which makes analyzing the system all in all a bit wierd. I can't see any way that you could get a negative specific heat by adding a vibrational mode, I think you have to add in a chemical potential well to get this result.
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posted
Bingo, Dr. Rabbit! That makes total sense. In a sense we are having to stretch the gas out to get it into the bigger volume, when changing it from 500 to 450 degrees. The energy has to come from somewhere, and so the gas cools down. That makes sense to me. You r0xx0rz!
So if we were heating the gas, and had an electronic "100% saturation maintainer" that would move the piston to whatever volume would maintain conditions at exactly what was required to stay at 100% saturated steam, when the system got to 450 F, would it quickly slide down the curve past the inflection point, and back up to whatever temperature has the same enthalpy as it does at 450 F? In a sense, it would be sort of an explosion (implosion?) that would rapidly take the temperature of the system to a much higher level. (I guess once it gets above the critical point, our saturation-ator would not know what to do and would just leave things the same volume.)
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quote:when the system got to 450 F, would it quickly slide down the curve past the inflection point, and back up to whatever temperature has the same enthalpy as it does at 450 F? In a sense, it would be sort of an explosion (implosion?) that would rapidly take the temperature of the system to a much higher level.
Once again, not necessarily. We have only considered the first law. To know whether or not a process will occur spontaneously we need to look at the second law. The Gibbs free energy will tell you whether or not a constant pressure process will occur spontaneously. Helmholtz free energy will tell you whether or not a constant volum process will occur spontaneously, but there isn't a standard free energy equation for constant %saturation processes.
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posted
Okay, the condition that we maintain constant 100% saturation is sort of odd, I understand, because that's not a common situation in real world boilers and things, and also, it's probably not easy to study. But my next question is, does the fact that steam acts this way mean that there IS a hill of some kind to be slid down, when saturated steam at 450 degrees will spontaneously increase in temperature? Perhaps at constant pressure or constant volume? It seems that there would be.
Glad you're enjoying it, AJ. Only on hatrack! Also, it's sort of interesting to me that so many of the principle participants in this thread are female. That is not a common situation in my real life, either. Posts: 6246 | Registered: Aug 2004
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posted
It's sort of bizarre that hatrack has such a prominant contingent of female engineers. I wonder what that means.
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I'm back in my dorm now, and I once again have access to my thermo textbook. I was looking at the tables in the back, and while it looks like some materials don't exhibit this peak in the range of temperatures and pressures given in the book, saturated propane vapor appears to have a peak at around 75 degrees C and refrigerant 134a has a peak at around 80 C. Which makes me wonder if the ammonia and refrigerant 22 tables simply don't go up high enough to show the peak. Does every substance have this property, do you think? If not, why not? I doubt propane is polar, so that can't be the only explanation.
Tatiana, I was thinking along the same lines earlier. It's odd to think that there might be a gas that would heat up spontaneously like that.
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posted
Yes, it is! I have been realizing something about steam that seems to make it make more sense to me. As the pressures and temperatures at saturation rise toward the critical point, the properties of sat. steam become more like the properties of sat. water, because of course, at the critical point they become indistinguishable. Water naturally has a whole lot less enthalpy than the same mass of steam at the same T and P, because of the latent heat of vaporization and all that. Maybe the reason steam's enthalpy begins to decrease as the P and T rise is because to bridge that gap between the enthalpy of water and that of steam, both substances properties trend more toward the average of the two. In other words, as you approach the critical point, it seems natural to realize that water becomes more like steam, and steam more like water, until they blur together into one indistinguishable thing. To become more like water, steam's enthalpy per unit mass has to decrease.
Viewed that way, it no longer seems like such a strange anomaly, and I begin to think most substances would act similarly. I still do wonder whether (and believe it must) at some point the system would rather violently implode suddenly, resulting in a system of greatly reduced volume, and much higher T and P. It seems odd to think it would, but I think that has to be what these steam table numbers are pointing to.
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