Evaporative cooling | Water, acids, and bases | Biology | Khan Academy
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- čas přidán 30. 06. 2015
- Evaporative cooling. Why sweating cools you down.
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TL:DR - When your body is hot, it releases sweat. Sweat absorbs the heat energy from your skin to turn into vapour. The transfer of heat from skin to sweat cools the skin.
tl;dr using my own words: your body is hot > it releases sweat (water molecules) > sweat (water molecules) absorb the heat and use it to break hydrogen bonds ('the water on ur skin evaporates') rather than raise the temperature of the molecules. this whole process is called "evaporative cooling" which helps stabilize temperatures.
"sweating helps dissipate our excess body heat"
clapped but it makes sense to me at least
I've been thinking of "Wet Bulb Temperature" for about 3 hours... I know this is not the purpose of the video.. but It helped me answering all my questions. Thank you !!!
Love your explanation of evaporative cooling! It helped me understand completely the concept, thank you.
Thanks so much, I have understood what I did not know
Wow you are good at this. Thank you!
I really like his voice
It's very calming😊
thank you ❤️
Cool
Thankyou sir. You explained it giving detailed information that how are water molecules getting that heat energy at the first pkace
Thank you sir
When our body is warm it releases swet. And swet absorbs heat from our skin to make its vapor due to this the overay system temperature decreases
If Any one can give the opinion on this system vs missing system, which is better and why?
Great video! It might be stupid questions, but:
1) does the specific heat of substances enter somewhere here?
2) If so, how does it influence the cooling effect? (like, substances with low specific heat being better at cooling -- I'm thinking about alcohol here, which appears to cool our skin faster than water, but at the same time, water being able to absord more water I think would lead to it being better at cooling)
3) Reading about, for example, plant transpiration, a book (Principles of terrestrial ecosystem ecology, by Chapin III) said it takes 580 times more energy (2.45 MJ kg^-1) to vaporize 1g of water at 20°C than to increase its temperature by 1°C, so I wonder: how does water vaporizes (cooling the plant in the process, in this case) instead of just heating?
water has a high heat capacity so it can absorb a lot of the heat --> cools you down since there's less heat
Cool🤜🏻👍
So would I cool down more if I wipe the sweat away, and then produce new sweat that can absorb more energy then the already heated up sweat that I just wiped away?
As long as I can produce more sweat [having a readily available source of water], I guess it would be better to wipe it away?
OR is it good to leave it on, because it's already so hot that the molecules will be more likely to escape than fresh sweat? Only problem I would think is that it's not going to absorb much new heat from the skin because it's already just as hot as the skin so there is almost an equilibrium?...
GPT-4 ANSWERING:
The cooling effect of sweat primarily comes from its evaporation from the skin, not from the sweat itself being cool when it is produced. Here's how it works:
1. Evaporative Cooling: When sweat evaporates from your skin, it absorbs heat energy from your body to transition from a liquid to a gas phase. This process, known as evaporative cooling, dissipates the absorbed heat into the environment, thereby cooling your body down
2. Wiping Sweat Away: If you wipe away the sweat before it has a chance to evaporate, you interrupt this cooling process. The energy that would have been absorbed by the evaporation process remains in your body, so you'll need to produce more sweat to achieve the same cooling effect.
3. Sweat Production: Your body will continue to produce sweat as long as it senses a need to cool down. However, wiping sweat away may lead to a need for increased sweat production, as each new batch of sweat will provide a new opportunity for evaporative cooling
4. Sweat Saturation and Evaporation: Fresh sweat will absorb energy from the body as it evaporates, regardless of whether it is replacing previously wiped-away sweat or adding to sweat that has remained on the skin. The rate of evaporation, and thus the rate of cooling, may be affected by external factors such as the relative humidity of the surrounding air.
The equilibrium you're referring to, where the sweat and the skin are at a similar temperature, doesn't hinder the evaporative cooling process. Evaporation will continue to absorb heat from your body as long as there's liquid sweat present to evaporate.
In summary, it's more beneficial to allow sweat to evaporate naturally rather than wiping it away, as the cooling effect comes from the evaporation process. If you have a ready supply of water and your body can continue to produce sweat, you'll still achieve cooling through evaporation, but wiping away sweat may simply lead to a need for increased sweat production to achieve the same level of cooling.
Tq
The reason that at room temperature (even in a warm room of lets say 37 degr. C) some water vaporize is not ONLY due to the unbalanced distribution of kinetic energy among all water molecules.
Let's assume that it is all that's happening. Think of a vessel with some water in it (a dish would be nice, as a wider area of shallow water would evaporate faster, but it's a though experiment so it doesn't matter what vessel exactly ... just that its top is opened)
Some molecules get enough energy, break their bonds and evaporate, leaving the rest of the water a little cooler (that energy they had has been used to convert these molecules in hotter waver vapors).
(Water vapors would rise in the air, until they hit the ceiling, in which case they would transfer a bit of their energy to the ceiling eventually enough to cool down and get in liquid state again (ie condensate), while the heat in the ceiling escapes in the environment. But that's also no important for this mind experiment!)
If that evaporation process continues, then eventually you'll have less liquid water in that vessel, but it would be cooler.
Which even if it's not cool enough to stop evaporating completely it would progressively slow down that evaporation (with the average temp. of the remaining water dropping).
So you can extrapolate until the last few molecules of water ... they bump into each other somehow, randomly passing more kinetic energy to one of them, which breaks its bonds and evaporates ... and then? ... eventually you'll have a molecule (or a few) that can't produce high enough kinetic energy by random disbalance between them ...
But we know that if left long enough (and that's not too long) the whole water will evaporate at room temperature. (** sure if only one or few molecules are left we can't normally see them, at least not without some very advanced instruments). But if you run the math I'm sure that the chance for another molecule to get high enough energy to escape will drop to negligible values, long before there are only few molecules left. Sure the evaporation might continue even then, but it would be that much slower.
If you run the math it should result in full evaporation times much longer than what's necessary for a few mm shallow water in a wide dish to evaporate (from my experience in a day or two usually - depends on the amount of water, surface of evaporation and room temperature of course ... and air humidity too!).
So just statistical unbalanced distributions can't explain the whole process!
From what I remember from school, they thought us that dry air even if relatively cold would still "pull" some water molecules from the body of water turning it into vapor, and that vapor would steal some energy, while making the rest of the water cooler. Not sure in details how that air dryness pulls water to evaporate much lower than boiling point (even for single molecules), but I'm quite sure that it does happen.
The air dryness is acting like a catalyst.
You just gave Richard Feynman's thumbnail description of this, but I think that it likely ignores thermodynamics, enthalpy, and specific heat of vaporization so as to provide a neat, intuitive, and only partially correct answer. Work is done and energy is used to break any bond, hydrogen bonds or otherwise, and you totally step over that.
I wish you had someone with a PhD in the subject making these videos. Because then I would be more likely to think believing them was not actively misinforming me.
Late reply but for those who are just now reading this, ignore MrShysterme's oblivious comment. This is a simplified video, sure, but there is no misinformation here. This is a video intended for biology students and hence avoids the complicity of a fully thermodynamic explanation. By the way, enthalpy and specific heat of vaporization are all a part of thermodynamics, indicating to me that MrShysterme does not know much of what he is talking about. Enjoy the video for what it is people! Sincerely, a chemical engineer.
@@Lrobert152 I'm not even a university student and I was about to call him out for that lmao.
I like these videos, but I agree with MrShysterme, there seems to be something missing here. Why when looking at an evaporative cooler, does both the evap cooler cool down, but also the air cool down? According to the video explanation I would think the air would heat up because of the fast water molecules escaping into the air, but the air cools down and does not heat up. If anyone can tell me what is going on I would appreciate it. Until then I will do a youtube hunt.
Here's an idea that I'm sure Spock would agree is only logical:
- Take for example your left arm. It has certain temperature, and in total certain energy. Now cut at the shoulder and throw it away. Great, now you've dumped all that heat and energy away from you ... additional bonus is that you would get even cooler with the blood loss
;)
It's not that difficult to dump some energy. The difficult part is to stop it from coming back ... if you don't want it back that is - if you want to heat yourself then the opposite is difficult ;) - obviously that depends on the environment.
SUMMARISED BY GPT-4:
Sweating as a Cooling Mechanism:
When our body gets warm, it sweats to prevent overheating. This cooling effect is achieved through a process called evaporative cooling.
Evaporative Cooling Explained:
1. When we sweat, our skin is covered in tiny beads of water (sweat).
2. On a microscopic level, these water molecules have varying kinetic energies. Some have high kinetic energy, while others have lower kinetic energy.
3. The temperature of a substance is determined by the average kinetic energy of its molecules.
4. Water molecules are held together by hydrogen bonds. However, some water molecules, especially those with high kinetic energy and near the surface, can break these bonds and evaporate into the atmosphere.
5. As the high kinetic energy molecules evaporate, the average kinetic energy of the remaining molecules decreases, leading to a decrease in temperature.
How It Cools the Body:
1. The skin is made up of molecules that vibrate due to their kinetic energy.
2. When high kinetic energy water molecules from sweat evaporate, the remaining water molecules on the skin have a lower average kinetic energy (lower temperature).
3. The warmer molecules of the skin then transfer their heat to these cooler water molecules, raising their kinetic energy.
4. Some of these newly energised water molecules then evaporate, continuing the cycle and leading to further cooling of the skin.
In essence, the heat from our body is used to vaporise the sweat, and as the high-energy molecules of sweat evaporate, it leaves the skin cooler. This process effectively dissipates the body's heat, allowing us to cool down.
temperature of water lesser in 1) evaporative cooling by wood wool used in EV air cooler 2) wet surface. other parameter put constant
Grade 7 lessons
Not sure who your target audience is, but I found this to be a frustratingly incomplete and oversimplified explanation. It does not explain how evaporative coolers cool both the water and the air into which the higher-energy water molecules escape.