Well, it seems our initial return to blogging suffered a slight setback. But we’ve been pretty hard at work over the past year on our new look, new website and many other new developments. And we’re happy to be returning to blogging at long last as well. We’re pleased to have you visiting, and invite you to stay tuned for much more of our latest news, happenings and thinking.
This is the third followup of the blog article: "A perfect storm in a cup of salt water?"
I woke up last night with a perfect explanation for the mysterious temperature gradient observed in a saturated salt solution. It is the recrystallization of salt at the bottom of the cup that releases the heat.
Since water molecules are constantly evaporating from the surface of the solution, a corresponding amount of ions must return to the crystal form at the same time--because a reduced amount of water in a saturated solution in the cup cannot take them any more. This most likely occurs at the bottom since the surface of the precipitate already provides a perfect ground of crystal growth. When ions adhere to the surface of a crystal, heat is released. The amount of released heat is approximately equal to half of the cohesive energy of the salt crystal (because it is a surface adhesion), which may be quite high because of the strong electrostatic attractions in the ionic crystal. The released heat transfers to the solution near the bottom and, together with the evaporative cooling effect on the surface, creates the temperature gradient we observed. The entire process runs continually across the solution because of the diffusion of water molecules and ions driven by their concentration gradients: the concentration of water/salt becomes lower/higher at the surface when water evaporates.
There are four evidences that support this theory:
- The temperature gradient disappears when we sealed the cup, because that stopped the evaporation at the surface as well as the recrystallization at the bottom.
- We observed no temperature gradient in an unsaturated solution because there is no recrystallization process.
- The temperature hiked when the sensor touched the salt deposit at the bottom.
- This temperature gradient lasts for a long time because this process will continue until all the water molecules evaporate.
The diagram above illustrates the energy cycle of a saltpan/ionic power plant combo. This design is based on a chain reaction that involves two phase changes in a salt solution to convert solar energy into electricity through the ionic potential.
PS: I found in Wikipedia the concept of solar pond that uses a large pool of saltwater to collect solar energy. I think its mechanism is different from what I discussed above. I have had no luck reproducing the effect of a solar pond in a cup yet.
I did an experiment to investigate the relationship of the salt concentration with the mysterious temperature gradient in a cup of salt water. The experiment was to measure the top-bottom temperature differences in three cups of salt water: low-concentration, medium-concentration, and saturated solution. In the saturated solution, there is a salt precipitate at the bottom of the cup. In all measurements, a fast-response temperature sensor was moved up and down in a cup. And the solutions had existed for over 100 hours to ensure that the salt was completely dissolved and the systems had reached thermal equilibrium with the environment.
The results shown in the graph above clearly indicate that the temperature gradient exists only in the saturated solution. The two unsaturated solutions exhibit no appreciable temperature gradients and measure approximately the same temperature with plain water.
The results were confirmed by an IR image shown above (from left to right: low-concentration, medium-concentration, and saturated).
This experiment suggests that there is probably no ion gradient in an unsaturated solution. An unsaturated salt solution has the same temperature everywhere and that temperature is the same as that of the plain water, whatever its concentration is. I originally expected that an unsaturated solution would have a temperature gradient more or less proportional to the salt concentration, as would a colligative property. This surprising result made me think that the prime suspect is the salt precipitate at the bottom of the cup. We know there is a lot going on on the surface of the precipitate layer. The dissolving and crystallization never cease. It is just that the two processes reach a dynamic equilibrium--the rate of dissolving becomes the same as the rate of crystallization. Sort of like what is shown below:
Let's think a bit more about the meaning of this experiment. Notice that the temperature curve of the saturated solution lies entirely between that of the ambient temperature and that of the pure water temperature in the graph. This means that the existence of the precipitate somehow weakens the evaporative cooling effect, and probably the evaporation process itself. Why would the evaporation of water at the top of the solution slow down in the presence of some precipitate at the bottom? Exactly how does this process contribute to the temperature gradient existing in the solution?
We can plausibly reason that the rate of evaporation decreases because the ions somewhat act as binding agents that hold the water molecules more tightly through the strong electrostatic attractions. This is known as the water shell effect—water molecules are attracted to an ion and form a dynamic shell around it. As a result, it is more difficult for water molecules to leave the solution to evaporate. But this picture cannot explain why there is virtually no difference between the temperature of a cup of pure water and the temperature of a cup of unsaturated salt water.
It seems the mystery is far from being uncovered. While clarifying a few things, this experiment makes the phenomenon more baffling. Stay tuned for our next investigation.
In terms of its other implications, there is one thing that we can rule out now. There is no such effect in the ocean, as sea water is not saturated.
This is the first followup of the blog article last week "A perfect storm in a cup of salt water?".
Several people including Bob Tinker, John Loosmann, and Einar Berg suggested that it was the evaporation of water that drives the observed persisting temperature gradient. It turned out that they were right. After sealing the salt cup with plastic wrap and leaving the three cups for 24 hours, their thermogram shows the sealed cup vanishes from the infrared image (see the image to the right--the sealed salt water cup is in the middle, which is invisible). This means that the temperature everywhere in the cup is the same as the ambient temperature. The infrared image also shows the baking soda cup, which has not been sealed, continues to show a temperature gradient.
Now, we have to explain why the pure water cup shows a cool infrared signature. So I added a sealed pure water cup. The thermogram to the right shows that the sealed pure water cup vanishes in the infrared image (the sealed water cup is on the right of the thermogram), whereas the open pure water cup shows a cooler image for the part filled with water. Interestingly enough, the upper part of the cup that does not have water contact is constantly almost 1 warmer than the filled part. This temperature gradient is clearly shown in the infrared image below. Why is the temperature gradient across the water line on the surface of the plastic cup so sharp?
Now go back to the evaporative updraft force. At this point, we only know that cutting off evaporation shuts down the energy loop. We still do not know how what happens under the water line in the salt water cup. The following graph clearly shows that this effect exists in not only a salt solution, but also a baking soda solution and a sugar solution.
I was bothered by an experiment I did recently about the temperature distribution in a cup of salt solution. I added a few spoons of table salt and baking soda in two cups of water to create two saturated solutions. Then I left them sit there for a few days, along with a cup of plain water. When I came back and aimed my infrared camera at them, I saw something quite puzzling: in the two cups of solution, the bottoms were always about 0.5 warmer than the tops (see the IR image above)! In contrast, a cup of plain water did not show this temperature difference--the temperature was the same everywhere just as expected.
Exactly what kind of chemical force sets up this temperature gradient? We all know that warmer water should rise and colder water should sink, and eventually the convection stops and the temperature becomes the same everywhere. But this is apparently not true in the presence of salt solute. I feel this has to do with gravity. It must be gravity that causes a concentration gradient of the solute, which in turn results in the temperature gradient. But I am not sure how exactly this happens. I have no idea what energy source feeds this temperature gradient. Don't forget that the cup material tends to eliminate it through heat conduction and the air through convection. There must be an invisible hand that counters all these thermodynamic forces. This seems pretty amazing to me.
To make sure that this is not an effect of infrared radiation, I confirmed the result by sticking a sensitive temperature probe into the solution and moved it up an down for a few times. The image below is the 60-second result recorded by the temperature probe, which clearly agrees with the IR image.
This is an example that, once again, shows the power of infrared imaging. I would not have noticed there was such a temperature gradient in a solution without my infrared camera. The infrared camera, in just one simple shot, captured the salient and subtle details that reveal very complex physics, which I still do not understand.
What is the significance of this result rather than a tempest in a teacup? Might the temperature gradient be used to generate a voltage gradient, which in turn generates electricity? In other words, might this be some kind of battery that is a 100% clean energy source?
The ocean is a gigantic solution of salt. Half Celsius of temperature difference in the ocean translates into an enormous amount of energy. Might there be such an effect in the ocean?
1) Evaporation is a driving force
2) The temperature gradient exists only in a saturated solution
3) Mystery solved?
4) Visualizing vapor pressure depression
5) Salinity gradient vs. temperature gradient
6) An evidence from an ice cube
Infrared (IR) imaging is a technique for seeing heat based on detecting thermal radiation (mostly IR) an object emits. It used to be a very expensive tool only affordable to guys in military and secret services where money is not a problem.
You can now buy "lower"-grade IR cameras with $1,500-$2,500, which are pretty cool (thank you for lowering down the prices, Flir and Fluke!). There is a vast market for this technology. Engineers and technicians buy them primarily for checking heat flow in building, electrical systems, and mechanical systems. Companies also use them to do quality assurance and safety monitoring.
I have been digging the educational potential of IR imaging lately. I feel that the tool can be very useful in education. Compare it with a microscope. Both can be used to see something invisible. In the case of a microscope, it is things that are too small to be seen. In the case of an IR camera, it is things that our eyes cannot detect. It is obvious that students need a microscope to see small things. But perhaps we can also rationalize the need for an IR camera in the classroom? What are the most important things that IR imaging can teach?
Obviously there is heat transfer. I have recently written a paper about this. But I don't want to just do the evident ones. So I have been thinking about how to broaden its applications. A direction I am taking now is its applications in chemistry, where heat is a central concept. You probably still remember that your high school chemistry teacher always wanted you to remember how much heat is released or absorbed in a chemical reaction. If a reaction produces a dramatic effect, such as a bang or a flash or a flame, then you probably were impressed. What about those reactions that mostly go unnoticed unless some sensitive methods are used to show them? For instance, most biochemical processes are pretty "calm." How does one "see" or "hear" them?
I have done an experiment that uses an IR camera to show evaporation and condensation, as mentioned in the paper. The above IR thermogram shows what happened when a piece of paper was placed on top of a cup of water. The paper did not fully cover the cup. What we see from this IR image is a cooler area that shows the evaporation process of the water in the cup and a warmer area that shows the condensation process of the water on the other side of the paper.
Last week, I did another experiment to prove that it can also be used to visualize dissolving. This experiment is introduced in a short article. The image to the right shows the thermograms of three cups: pure water, table salt solution, and baking soda solution, shortly after table salt and baking soda were added to two of the cups originally filled with pure water.
I am hoping to devise more chemistry experiments to prove the versatility of this powerful tool in making mysterious things in chemistry visible. I intuitively feel that this tool, which is essentially a bundle of thousands of IR thermometers, may be able to release students from tedious lab procedures and make chemistry experiments easier to conduct and fun to look at.