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I noticed, and perhaps many others have too, that the application of hand sanitizer (mainly ethanol), leaves one's hands feeling rather chilly after application.

What is responsible for this phenomena? Is it the high heat of vaporization of hand sanitizer? However, explanation doesn't hold water; water has a standard heat of vaporization of 40.65 kJ/mol while ethanol has a heat of vaporization of 38.56 kJ/mol.

Could it be the low boiling point of ethanol? Hand sanitizer disappears (vaporizes) within seconds upon rubbing the hands together. Water, however, does not.

Additionally, how does one square a high heat of vaporization with a low boiling point? If it takes a lot of energy to vaporize something, then how can that something have a low boiling point?

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    $\begingroup$ I noticed the same thing with acetone. I think it's evaporation related, but that's all I can say. $\endgroup$ – user137 Sep 3 '14 at 19:36
  • $\begingroup$ My best shot at an answer to your last paragraph is here: chemistry.stackexchange.com/questions/32504/… $\endgroup$ – hBy2Py Jun 6 '15 at 22:15
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As user137 noted, this is from evaporation of the alcohol. Evaporation of a liquid takes the enthalpy of vaporization away with it, making you cooler. Just like how sweating helps keep you cool in the summer, and how swamp coolers (evaporative coolers) keep houses cool in the southwest.

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Hand sanitizer leaves your hands feeling cool because the particles in the gel which posses the most amount of energy are able to evaporate from the gel and off your hands. This results in only the particles with low energy levels to be left on your hand which therefore reduces the hands overall temperature.

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The above explanations are not incorrect, but they don't address the focus of the original question: the comparison of alcohol with water. The questioner was, in effect, saying they knew about "Evaporative Cooling", but had rejected that explanation due to the higher enthalpy value for water verses alcohol.

I believe the answer to this aspect of the question has to due with the quantity of alcohol vs. water that evaporates per unit time. That is, although more energy is removed from the skin per gram of water evaporated, it takes a lot longer for each gram of water to evaporate compared to each gram of alcohol. That is, alcohol evaporates so much faster that it ends up being more efficient at removing heat.

The temperature we "feel" is often deceptive because it's usually a difference we detect, having to do with a lot of other things related to our nerves and their recent history, so there could be other factors related to alcohol's stronger interactions with hydrophobic molecules on our skin (especially oils), or the "drying out" effect that alcohol has after it evaporates from our skin.

But none of this rationalization means that much from a scientific perspective unless actual tests are made. So, I would suggest different solvents be applied to skin and/or suitable skin-like surfaces and the rate of temperature tracked with one of those IR thermometers. Alcohols of different chain lengths, butanol, propanol, ethanol, methanol (warning, poisonous!) would be perfect, but it'd also be interesting to try things like acetone or even oils. The results of those tests would be the real answer.

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  • $\begingroup$ Nice answer. You might find it useful, however, to edit the post to insert some paragraphs/formatting, the single block of text is a little hard to read as it stands. $\endgroup$ – NotEvans. Jul 21 '17 at 11:28

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