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I was just going through acids on Wikipedia in my free time, and noticed this neat trend:

  1. Formic acid, Ethanoic acid, uptil Nonanoic acid, are all liquids. The first few are colourless, while the latter are yellowish and also "oily" (more viscous).
  2. Decanoic acid and all its elder brothers are white crystals (or "powders")
  3. Aromatic acids like benzoic acid or picric acid are solids.

I am unsure if there could be a reasonable logic for the colours. But, I am sure there could be a good logic for the physical state of these acids at room temperature.

To me, all these acids have hydrogen bonding common, so we can rule out that factor. Apparently, the physical state seems to be related to the "size" of the acid. The heavier acids are solids, while the lower ones are liquid. My question is: why is it so?


PS: List of all acids at one place is at the bottom of this Wikipedia page.

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    $\begingroup$ Van der waals interactions $\endgroup$ – Cyclohexanol. Feb 3 '18 at 9:09
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    $\begingroup$ Come to think of it, higher everything is solid (alkanes, alkenes, alcohols, etc.) $\endgroup$ – Ivan Neretin Feb 3 '18 at 9:37
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    $\begingroup$ On a cool day, you probably would not be asking this question. Formic acid: mp. 8 deg. C (47 deg. F); acetic acid: 17 deg. C (62 deg. F). $\endgroup$ – user55119 Feb 3 '18 at 16:01
  • $\begingroup$ @user55119 haha, sharp observation! ;) Though on a more serious note I have always hoped (but have never seen or read about) that chemical laboratories of high standard have temperature control builtin, so they maintain temperature at around 25degC throughout their room. $\endgroup$ – Gaurang Tandon Feb 3 '18 at 16:05
  • $\begingroup$ Not when it isn't working! $\endgroup$ – user55119 Feb 3 '18 at 19:32
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Acids with a larger size have greater intermolecular forces than 'smaller' acids, meaning it would take more heat energy to break those bonds in larger acids than in smaller ones. Hence at room temperature the heavier acids would be solid.

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  • $\begingroup$ Really, that's it? This was so straightforward :/ Thanks though! :) $\endgroup$ – Gaurang Tandon Feb 3 '18 at 15:03
  • $\begingroup$ @GaurangTandon Well, this question was rather trivial and I'm pretty sure similar ones were asked here many times... :( $\endgroup$ – Mithoron Feb 3 '18 at 23:03

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