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We performed an experiment where tissue paper soaked in various chemicals was wrapped around the end of a temperature probe.

The chemicals with stronger intermolecular forces had a smaller drop in temperature (from room temperature to final after most of the alcohol had evaporated off). Two of the chemicals were ethanol, with hydrogen bonding, and n-pentane, with only London dispersion forces.

My question is why?

Below are just my initial attempts to understand

My initial feeling was that molecules with stronger intermolecular forces would take more energy to evaporate and would have a larger temperature change than those with weaker bonds, but we see the opposite in the data.

I was also thinking that maybe went the other way: chemicals with stronger forces had fewer molecules evaporate, so the temperature drop was smaller. Is this correct?

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    $\begingroup$ It's a matter of kinetics, not thermodynamics. If you measured heat not temp. you should get expected results. $\endgroup$
    – Mithoron
    Commented Oct 31, 2018 at 23:18
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    $\begingroup$ The heat (enthalpy) needed vaporise 1 g of each solvent is approx 840 J for ethanol and approx 360 for the pentane. So unless you have equal or at least known quantities in your tissue your experiment can give misleading results. However, your initial idea about stronger forces needing more energy to evaporate makes sense. $\endgroup$
    – porphyrin
    Commented Nov 1, 2018 at 9:37

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