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Pure Liquid HCl

  1. Apr 15, 2013 #1
    I had read that water is essential for any acid. When HCl is dissolved in water, it produces H+ ions which cause the litmus paper to change it colour! But if HCl gas is liquidified and then tested with liquid, what would be the results? As wel as this pure liquid HCl is corrosive or not?
     
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  3. Apr 15, 2013 #2

    Borek

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    No idea what you mean. No idea at all.
     
  4. Apr 15, 2013 #3

    AGNuke

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    I believe he wants to know the acidic properties of the HCl "liquid" (solvent form).
     
  5. Apr 15, 2013 #4

    Borek

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    Up to him to explain, not up to us to guess.
     
  6. Apr 17, 2013 #5
    I am guessing a bit here, but if you managed to keep the water away from this HCl liquid then you would not get the H+/OH- dissociation of acid/alkali chemistry so it should not be acidic as such. Without water it should be less reactive but if it came near anything with water, like flesh, it would be extremely corrosive.
     
  7. Jun 10, 2015 #6

    SRJ

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    Theoretically, pure HCl in gaseous or liquid form does not show any acidic properties, but if you try this in the lab, it will absorb water from the glass beaker, the air(moisture) and even the litmus paper that you would try to test it with. Therefore, in practice it behaves like an acid even in "pure" form.
     
  8. Jun 11, 2015 #7

    TeethWhitener

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    This is not strictly true. Pure HCl undergoes self-ionization, just like water, ammonia, HF, and a host of other pure protic liquids. This is most evident in the fact that pure HCl has a nonzero electrical conductivity (3×10-8 mho/cm, according to this source; cf. 5.5×10-6 mho/cm for pure deionized water). HF is much more extensively studied than HCl in this respect, and it is known that HF has the self-ionization equilibrium: [tex]3HF \rightleftharpoons H_{2}F^{+} + HF_{2}^{-}[/tex] I would assume the HCl equilibrium is similar. The pH range wouldn't be the same, however.
    [EDIT] To be precise, since the conductivity is lower in HCl than in water, you'd expect the obtainable pH's in pure (anhydrous) HCl to be higher (more basic) than in water (surprisingly enough). Two examples: the conductivity of pure HF is higher than pure water, and the obtainable pH's in HF are anywhere from 0 down to -20 or so (extremely acidic). On the other hand, the conductivity of pure ammonia is much lower than pure water, and the obtainable pH's in ammonia range from about 10 to 35 (extremely basic). The reason for this is that the conductivity of the liquid is related to the equilibrium constant for the self-ionization reaction. You might want to search for terms like "self-ionization" and "solvent leveling effect" to understand this better.
     
    Last edited: Jun 11, 2015
  9. Jun 11, 2015 #8

    Borek

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    While I understand what you are thinking about, this is a rather dangerous way of using the term "pH". Technically speaking pH is defined only for water solutions, and shouldn't be used in other contexts to avoid ambiguity. The part of "higher pH meaning it is more basic" is exactly what I am talking about, that's comparing apples and oranges.

    If anything, it would be better to switch to Hammett acidity function for such comparisons.
     
  10. Jun 11, 2015 #9
    I understand Borek's caution. But pH is only a -ve log of "proton" concentration and can be used for all systems. Just we cannot assume 7 is neutral, that is specifically for water; generally acids and bases are defined with that 7 in mind.

    I take TeethWhitener's "more basic" as just relative term to aqueous HCL; less acidic would have been easier, though both mean same. It is same as saying less protons. Less protons and so higher pH. So, pure HCL can have a pH above 7 and still be "acid"! Also, their reactivity will be different because in water protons stay as H3O+ but in other medium it will stay as something else (in more or less reactive form).

    Now we can understand the importance of Borek's caution.
     
  11. Jun 11, 2015 #10

    Borek

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    It even doesn't work for water - depends on the temperature and ionic strength.
     
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