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amateurawesomeness
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Might just try that then. I'm ordering some more sodium bisulphate tonight. What amounts should I use?
EDIT:Nevermind I found a video which shows the volumes pretty well.
[Edited on 17-4-2013 by amateurawesomeness]
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subsecret
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If you live in the USA (Or Canada, I assume), it might be wise to buy a gallon of ~8-10 molar HCl at a hardware store. It's sold as "Muriatic Acid"
(just an old name for HCl solutions), and it will definitely last you a while. It's technical grade, however, and it often comes with a slightly
green/yellow tint (mostly due to iron salts). It would be much easier to distill this to obtain a purer product.
Good luck!
Fear is what you get when caution wasn't enough.
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confused
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what would be the typical concentration of HCl after the distillation?
i ask because its very nearly impossible to get HCl where i live, would it be comparable to concentrated lab grade HCl?
Awesomeness, won't distilling the HCl drive it off in the form of HCL gas?
[Edited on 14-12-2013 by confused]
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blogfast25
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Quote: Originally posted by confused  | what would be the typical concentration of HCl after the distillation?
i ask because its very nearly impossible to get HCl where i live, would it be comparable to concentrated lab grade HCl?
Awesomeness, won't distilling the HCl drive it off in the form of HCL gas?
[Edited on 14-12-2013 by confused] |
HCl/water form an azeotrope (look it up) at around 20 % HCl. Sooner or later during distillation this fixed composition starts distilling over.
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AJKOER
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| Quote: Originally posted by blogfast25 |
HCl/water form an azeotrope (look it up) at around 20 % HCl. Sooner or later during distillation this fixed composition starts distilling over.
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While a true statement, there are possible paths to defeating the azeotrope. For example, Wikipedia (see http://en.wikipedia.org/wiki/Azeotrope ) describes several methods including pressure swing distillation, azeotropic distillation involving the
introduction of an an entrainer (an additional agent) impacting the volatility of one of the azeotrope constituents, chemical action separation where
an entrainer is added also having a strong chemical affinity for one of the constituents, distillation using a dissolved salt where a salt is
dissolved in a solvent to alter its relative boiling point, extractive distillation (similar to azeotropic distillation except that the entrainer is
less volatile than any of the azeotrope's constituents) to mention a few.
In the case of dilute HCl, the addition of anhydrous calcium chloride is one possibility.
[Edited on 23-12-2013 by AJKOER]
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subsecret
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It depends on the concentration you need. For most purposes, concentrated HCl is not needed, and ~20% HCl would work fine. If you need concentrated
acid, you could bubble HCl gas through cold water, after passing it through concentrated sulfuric acid to remove extra water. To obtain HCl gas, just
heat HCl (aq) to reduce the solubility of hydrogen chloride in the water. You can also drip concentrated sulfuric acid over NaCl, and I've heard that
you can drip concentrated sulfuric acid into an HCl solution (though I assume this only works with a relatively high concentration of HCl).
Could the concentration of HCl in the azeotrope be raised by distilling over concentrated sulfuric acid?
Fear is what you get when caution wasn't enough.
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AJKOER
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If this technique worked (I would not recommend it as my 1st choice), it would fall in the chemical action separation category where the H2SO4
entrainer has a strong chemical affinity for water.
For your information, I have previously given references on the topic of 'activity level'. The addition of salts like dry MgCl2, for example, have
been demonstrated to significantly raise the activity level making the weak HCl seemingly behave as a much stronger acid. This is all without even
the need for distillation. The result is cool, but incurs the expense of adding the appropriate salt. The required math to account for the result can
be both tedious and advanced (even with a powerful computer).
[Edited on 23-12-2013 by AJKOER]
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bfesser
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I strongly recommend <strong>Chemical Equilibrium</strong> by Allen J. Bard, but was unable to locate a full <a
href="http://books.google.com/books?id=x4f0UcZDVVQC" target="_blank">digital copy</a> <img src="../scipics/_ext.png" />.
<a href="http://books.google.com/books?id=d6eCW8mhP8QC&pg=PA257" target="_blank">A Deeper Look at Chemical Equilibrium</a> <img
src="../scipics/_ext.png" /> (<strong>Exploring Chemical Analysis</strong> by Daniel C. Harris; Google Books)
<a href="http://en.wikipedia.org/wiki/Ionic_strength" target="_blank">Ionic strength</a> <img src="../scipics/_wiki.png" />
<a href="http://en.wikipedia.org/wiki/Thermodynamic_activity" target="_blank">Thermodynamic activity</a> <img
src="../scipics/_wiki.png" />
[Edited on 23.12.13 by bfesser]
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blogfast25
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| Quote: Originally posted by AJKOER | | The addition of salts like dry MgCl2, for example, have been demonstrated to significantly raise the activity level making the weak HCl seemingly
behave as a much stronger acid. This is all without even the need for distillation. The result is cool, but incurs the expense of adding the
appropriate salt. [Edited on 23-12-2013 by AJKOER] |
"the weak HCl"... the mind boggles, truly. Hydrochloric acid is so strong its pKa can only be estimated from theory. One estimate is pKa = - 4, or Ka
= 10<sup>4</sup>, with Ka for relatively dilute solutions:
Ka = [H<sub>3</sub>O<sup>+</sup>] x [Cl<sup>-</sup>] / [HCl]
for the equilibrium:
HCl(aq) + H<sub>2</sub>O(l) < === > H<sub>3</sub>O<sup>+</sup>(aq) + Cl<sup>-</sup>(aq)
BTW, 10<sup>4</sup> = 10000.
Now go and calculate which percentage of HCl of a 0.1 M HCl solution isn't dissociated (deprotonated). Please???
Oh, and can you post that reference the MgCl2 again, because I'm fairly sure you are misinterpreting that...
[Edited on 23-12-2013 by blogfast25]
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AJKOER
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A reference on my 'activity level' remark is found in the hydrometallury text provided below.
Not exactly a lot of detail provided in this source, unfortunately, around the claim that one can significantly raise the 'activity level' of dilute
HCl by adding MgCl2 (alternatively, to a lesser extent using NaCl). It references most likely relates to its importance in this field where leaching
out minerals from ores efficiently and cheaply is desirable.
Source: See Hydrometallurgy in Extraction Processes, Vol I, by C. K. Gupta and T. K. Mukherjee, page 15 at
<del>http://books.google.com/books?id=F7p7W1rykpwC&pg=PA203&lpg=PA203&dq=FeCl2+%2B+O2+%3D+FeO(OH)+%2B+FeCl3&source=bl&ots=fi
WLs05y8f&sig=mi-pV94woVj7JABKBB
zLZcqbEwM&hl=en&sa=X&ei=ZQgfUeq7BIi50AGynoDoBQ&sqi=2&ved=0CFAQ6AEwBA#v=snippet&q=Magnesium%20chloride%20MgCl2&f=false</
del> http://books.google.com/books?id=F7p7W1rykpwC&pg=PA15 .
Note, the author claims there is data confirming that a 2M HCl in 3M CaCl2 or MgCl2 (or FeCl3) behaves like 7M HCl.
Also, hopefully the corrected link will work properly (something about issues with Google books across jurisdictions).
[Edit] Here is a quote on the matter of discussion in one of the reference sources kindly provided above by Bfesser on Thermodynamic Activity (see http://en.wikipedia.org/wiki/Thermodynamic_activity ):
"When a 0.1 M hydrochloric acid solution containing methyl green indicator is added to a 5 M solution of magnesium chloride, the color of the
indicator changes from green to yellow—indicating increasing acidity—when in fact the acid has been diluted. Although at low ionic strength
(<0.1 M) the activity coefficient approaches unity, this coefficient can actually increase with ionic strength in a high ionic strength regime. For
hydrochloric acid solutions, the minimum is around 0.4 M.[1]"
[Edited on 25-12-2013 by AJKOER]
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blogfast25
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I'd really like to see how this activity translates into the rate a chloride/HCl mixture dissolves a given metal, compared to simple HCl (all other
things being equal).
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Random
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| Quote: Originally posted by blogfast25 | | I'd really like to see how this activity translates into the rate a chloride/HCl mixture dissolves a given metal, compared to simple HCl (all other
things being equal). |
Someone should take for example 15% HCl and record the time when the small part of iron nail dissolves alone or with added NaCl in one case and MgCl2
in other. Cacl2 would be also interesting to check if it's the property of the compound alone or the amount of chloride ions in the mole of salt.
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blogfast25
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| Quote: Originally posted by Random |
Someone should take for example 15% HCl and record the time when the small part of iron nail dissolves alone or with added NaCl in one case and MgCl2
in other. Cacl2 would be also interesting to check if it's the property of the compound alone or the amount of chloride ions in the mole of salt.
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Yes, but who will be the 'someone'? 
Adding a chloride to HCl to increase its activity does of course have one massive drawback that doesn't have to be spelled out here, I think.
[Edited on 25-12-2013 by blogfast25]
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bfesser
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| Quote: Originally posted by blogfast25 | | Yes, but who will be the 'someone'? Adding a chloride to HCl to increase its activity does of course have one massive drawback that doesn't have to
be spelled out here, I think. |
blogfast25, please don't make useless cryptic replies. What
is immediately obvious to you may not be so to others. If you have something to add, write it out.
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AJKOER
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Blogfast25 expresses a valid concern except in the case when one is attempting to expel a gas (H2, Cl2, CO2, SO2, ...) from the solution.
My experience is that having prepared one reactant from a Calcium salt (and apparently didn't wash sufficiently), I was surprised when it decided to
show up later in a different reaction using the impure reactant. As such, an inert salt is not always subsequently 'inert'.
In addition to Hydrometallurgy, there may be other instances where such 'impure' acids have niche applications (probably relating to cost and
accessible issues).
[Edited on 25-12-2013 by AJKOER]
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blogfast25
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If, for instance, you increase the activity of HCl solution with e.g. CaCl2 to dissolve a substance, you end up with CaCl2 in the solution. For most
chemical purposes that isn't really a desirable outcome.
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