Sciencemadness Discussion Board

Concentration of HCl

amateurawesomeness - 14-4-2013 at 13:42

This is my first post. Yay.

I have 200ml of 0.1M Hydrochloric acid. This is ridiculously dilute and is not very useful in experiments. I want to concentrate it to 1M without losing too much of the solution, but I am unsure about how to do this(I'm new to chemistry).
I know it will have to be done by distillation or evaporation but I am worried that I will lose most of it.
Any ideas?
Thanks

AJKOER - 14-4-2013 at 13:51

Add NaCl and some FeCl3 (very important). Distill.

The addition of these salts is reputedly reported to may it possible to concentrate, upon distillation, almost any dilute HCl.

See if works.

amateurawesomeness - 14-4-2013 at 14:09

Thanks for the quick reply.
I will try in a few days and see if it works. any idea how much of the salts I should add? I don't wanna mess it up.

Pyro - 14-4-2013 at 14:13

If he's new to chem he probably doesn't have a distilling apparatus.
I don't know if this will work if its so dilute. but heating it should drive the HCl out of solution, so heat it in a small container with a hose leading the gas into ice cold water. otherwise you could either make it from drain cleaner H2SO4 and table salt or just buy it in the hardware store, its usually very cheap and quite pure depending on the brand.

amateurawesomeness - 14-4-2013 at 14:17

I have homemade distillation stuff (old tubes and stoppers mostly). I will try these methods and if they don't work I'll give up and buy some.

Hexavalent - 14-4-2013 at 14:18

Quite frankly, I don't think a beginner to chemistry should be handling concentrated sulfuric acid...

Supposedly just distilling dilute hydrochloric acid will induce the azeotrope, which is a 20 %w/v (?) solution of HCl in water. However, as aforementioned, as your starting acid is so dilute, I'm not sure this will work.

Finally, I'd be cautious about using rubber stoppers when distilling acids...nitric acid destroys rubber stoppers, but I'm not certain about hydrochloric.

[Edited on 14-4-2013 by Hexavalent]

plante1999 - 14-4-2013 at 14:18

Don't listen to him! Better to wait until something decent come out! (refering to AJKOER)

You should consider buying sodium bisuphate, it is available as pH down for pool. Reacting this with sodium chloride will make HCl gas that can be dissolved in water.

[Edited on 14-4-2013 by plante1999]

amateurawesomeness - 14-4-2013 at 14:23

Thanks plante1999. I have some sodium bisulphate crystals and a bottle of a solution of it so that should be really useful. I'll probably just pour the whole bottle in (about 90ml) I just saw a nurdrage video explaining it and it looks quite simple. The funny thing is, I'm not sure I have any sodium chloride. :)

[Edited on 14-4-2013 by amateurawesomeness]

Hexavalent - 14-4-2013 at 14:31

No, don't use the solution. AFAIK, using the solid sodium bisulfate is more efficient, although you may need some water depending on your heat source, as NurdRage explains.

plante1999 - 14-4-2013 at 14:35

Sodium chloride is sold as Salt in the majority of grocery stores, for about 2$ per Kg.

Pyro - 14-4-2013 at 14:36

If he's new to chem he probably doesn't have a distilling apparatus.
I don't know if this will work if its so dilute. but heating it should drive the HCl out of solution, so heat it in a small container with a hose leading the gas into ice cold water. otherwise you could either make it from drain cleaner H2SO4 and table salt or just buy it in the hardware store, its usually very cheap and quite pure depending on the brand.

BlackDragon2712 - 14-4-2013 at 15:02

what about electrolisis of NaCl? you can react the chlorine with water to make HCl, just put a carbon anode and cathode and pass electricity to them while they are submerged into a satured solution of table salt. you put a hose in the chamber and you put the other part in ice cold water and you have HCl, is the easyest way to do it. concentrating HCl is difficult because is a gas diluted in water, considering this remember that gases are less soluble at higher temperatures, and your solution is so diluted that you must remove 99% of the water in order to reach a concentration close to 10M and 96% of the water for a solution of 10% HCl

elementcollector1 - 14-4-2013 at 15:14

Quote: Originally posted by BlackDragon2712  
what about electrolisis of NaCl? you can react the chlorine with water to make HCl, just put a carbon anode and cathode and pass electricity to them while they are submerged into a satured solution of table salt. you put a hose in the chamber and you put the other part in ice cold water and you have HCl, is the easyest way to do it. concentrating HCl is difficult because is a gas diluted in water, considering this remember that gases are less soluble at higher temperatures, and your solution is so diluted that you must remove 99% of the water in order to reach a concentration close to 10M and 96% of the water for a solution of 10% HCl

The Cl2 would disproportionate in the water to form HCl and HOCl. A better method would be to react the H2 and Cl2 in presence of UV light (in a non-glass container) to form 2HCl, and then pass that through a small amount of cold water. Unfortunately, this reaction's rather explosive.

AJKOER - 14-4-2013 at 15:38

This statement has a few issues:

"A better method would be to react the H2 and Cl2 in presence of UV light (in a non-glass container) to form 2HCl, and then pass that through a small amount of cold water. Unfortunately, this reaction's rather explosive"

First, uv light will not penetrate anything other than a clear very thin container. I know this because only then will uv rays detonate a pure H2 and Cl2 mix. By pure, I mean no O2, which apparently poisons, even in small amounts, the chain reaction leading to an explosion. Next, to say it is explosive is a big understatement. The kinetics are so high that this reaction has been suggested as possible replacement of Gasoline/O2 engines and is more powerful than H2/O2 (there are videos of H2 and Cl2 explosions on the web). No way anyone could capture the HCl, you will just lose some body parts is the more likely scenario.

Further, adding Cl2 to water to make HCl and HOCl is not that good either as chlorine is not that readily soluble in water (perhaps adding a little SO2 might help).
----------------------------------------------------

On the NaHSO4/NaCl solid reaction, why bother? The NaHSO4 is a more friendly (and weaker) substitute for H2SO4 in many reactions. If you readily need HCl in an experiment, make it in situ by adding NaCl to a hot concentrated NaHSO4 solution. Not only is this safer, but with excess NaCl, the concentrated ionic solution may actually have an 'activity level' for your HCl in excess of what you could make from the same amount of NaHSO4 by dissolving the hydrogen chloride gas in water.


[Edited on 14-4-2013 by AJKOER]

Erbium_Iodine_Carbon - 14-4-2013 at 16:00

You can put your HCl in the freezer and the water will freeze out leaving you with more concentrated acid. Keep straining out the ice chunks until the liquid volume is 1/10 of the original volume if you want 1 mol/L HCl.

I don't think you will have too many problems with freezing point depression because a 10% or 2.87 molar solution freezes at -18 degrees C. I found all this information on the Wikipedia page for hydrochloric acid.

kristofvagyok - 14-4-2013 at 16:14

Quote: Originally posted by Erbium_Iodine_Carbon  
You can put your HCl in the freezer and the water will freeze out leaving you with more concentrated acid. Keep straining out the ice chunks until the liquid volume is 1/10 of the original volume if you want 1 mol/L HCl.
Do you have any reference for this? Or have anyone done this successfully?

amateurawesomeness - 15-4-2013 at 01:49

I would like to try and do the freeing method but I would only end up with 20ml of HCl.
I know where to get sodium chloride, but my family have some super strong sea salt stuff so I need to get some table salt(I used it all up).
And I need some more sodium bisulphate. I realised I only have a couple of grams. Luckily I found it cheap on the web so I'll buy some soon.

woelen - 15-4-2013 at 02:38

Another option of concentrating 0.1 M HCl is simply boiling down. You will not drive off HCl-gas at such low concentrations, due to the azeotropic nature of the HCl/H2O mix. First you'll drive off mainly water and when the concentration comes close to the azeotropic value, there will be more and more HCl in the vapor, but the relative amount will never be more than the azeotropic percentage. A 0.1 M solution easily can be boiled down to 1 M without any noticeable loss of HCl.

amateurawesomeness - 15-4-2013 at 05:19

Thanks. I'll try that and if it doesn't work, then I'll go with the sodium bisulphate option.

Fantasma4500 - 15-4-2013 at 05:51

i have been wondering about concentrating HCl some times myself tho i can get 30% HCl like if it was nothing..
if you cant find highly concentrated sulfuric acid you should go to a car store, or well a store where they sell things for cars and such
they usually have sulfuric acid standing around marked as ''battery acid''
its usually around 37% AFAIK, about 1.27 molarity or something
sulfuric acid can easily be boiled down to high concentration, the problem first starts when you get to 70% where SO3 starts forming, you dont want to breath this in also its pretty well etching
just thought about another thing...
if you can get hold of a shitload calcium chloride, its sometimes used as dessicant for drying rooms etc.
CaCl2 shouldnt decompose upon heating within 1000*C or well just past 100*C.. if you dump anhydrous calcium chloride into the HCl it should suck up the water in the solution and you can then distill the HCl
procedure can be repeated
probably possible with CuSO4 anhydrate or calcium sulfate if it doesnt make sulfuric acid and calcium chloride..

Erbium_Iodine_Carbon - 15-4-2013 at 08:39

Quote:
Do you have any reference for this? Or have anyone done this successfully?


Solid-Liquid phase diagram for HCl:


http://en.wikipedia.org/wiki/File:Phase_diagram_HCl_H2O_s_l.PNG

Sources, as quoted on page:
"erstellt aus Daten vom Gmelins Handbuch der anorganischen Chemie
plotted with data from the German books Gmelins Handbuch der Anorganischen Chemie, Systemnummer 6 Chlor, Verlag Chemie Berlin 1927 and Gmelins Handbuch der Anorganischen Chemie, Systemnummer 6 Chlor, Ergänzungsband Teil B - Lieferung 1, Verlag Chemie Weinheim 1968. Since 1990 the Gmelin was published as Gmelin Handbook of Inorganic and Organometallic Chemistry"

Morgan - 15-4-2013 at 09:42

Quote: Originally posted by AJKOER  
This statement has a few issues:

"A better method would be to react the H2 and Cl2 in presence of UV light (in a non-glass container) to form 2HCl, and then pass that through a small amount of cold water. Unfortunately, this reaction's rather explosive"

First, uv light will not penetrate anything other than a clear very thin container. I know this because only then will uv rays detonate a pure H2 and Cl2 mix. By pure, I mean no O2, which apparently poisons, even in small amounts, the chain reaction leading to an explosion. Next, to say it is explosive is a big understatement. The kinetics are so high that this reaction has been suggested as possible replacement of Gasoline/O2 engines and is more powerful than H2/O2 (there are videos of H2 and Cl2 explosions on the web). No way anyone could capture the HCl, you will just lose some body parts is the more likely scenario.

Further, adding Cl2 to water to make HCl and HOCl is not that good either as chlorine is not that readily soluble in water (perhaps adding a little SO2 might help).
----------------------------------------------------

On the NaHSO4/NaCl solid reaction, why bother? The NaHSO4 is a more friendly (and weaker) substitute for H2SO4 in many reactions. If you readily need HCl in an experiment, make it in situ by adding NaCl to a hot concentrated NaHSO4 solution. Not only is this safer, but with excess NaCl, the concentrated ionic solution may actually have an 'activity level' for your HCl in excess of what you could make from the same amount of NaHSO4 by dissolving the hydrogen chloride gas in water.


[Edited on 14-4-2013 by AJKOER]


Slight tangent on the main topic ...
I was wondering if you had any more info on the oxygen posioning of the H2/Cl2 reaction.
"Vacuum UV is so-named because it is absorbed strongly by air, and is therefore used in a vacuum. In the long-wave limit of this region, roughly 150 – 200 nm, the principal absorber is the oxygen in air. Work in this region can be performed in an oxygen-free atmosphere (commonly pure nitrogen), avoiding the need for a vacuum chamber."
https://en.wikipedia.org/wiki/Ultraviolet
I remember that H2/Cl2 demo on youtube where the fellow fires off a test tube using little LED lights. Blue sometimes works, the ultraviolet producing the bang. Also a quartz test tube was employed to allow the UV.
http://www.youtube.com/watch?v=NN82GoBG98s
"Ordinary glass is partially transparent to UVA but is opaque to shorter wavelengths, whereas silica or quartz glass, depending on quality, can be transparent even to vacuum UV wavelengths. Ordinary window glass passes about 90% of the light above 350 nm, but blocks over 90% of the light below 300 nm."
https://en.wikipedia.org/wiki/Ultraviolet

So I came across this title but you can't read the entire article.
"The Photosensitized Explosion of Hydrogen and Oxygen by Chlorine." I don't know if it's pertinent or not but it seemed curious.
http://www.jstor.org/discover/10.2307/95863?uid=3739600&...

I wonder though if the UV might work to initiate H2/Cl2 by creating a tiny electrostatic charge? Would water vapor/humidity have any inhibiting effect on the hydrogen/chlorine sensitivity to detonation using UV initiation?
"On the Moon, there is no rubbing. The dust is electrostatically charged by the Sun in two different ways: by sunlight itself and by charged particles flowing out from the Sun (the solar wind)."
http://science.nasa.gov/science-news/science-at-nasa/2005/30...

amateurawesomeness - 15-4-2013 at 11:05

So many options!
I had not heard about calcium chloride being used like that before. It looks like it would work but I don't have enough calcium chloride. I might just go with the heating option and if that doesn't work then I'll make some from sodium bisulphate and salt.

AJKOER - 16-4-2013 at 16:41

Quote: Originally posted by woelen  
Another option of concentrating 0.1 M HCl is simply boiling down. You will not drive off HCl-gas at such low concentrations, due to the azeotropic nature of the HCl/H2O mix. First you'll drive off mainly water and when the concentration comes close to the azeotropic value, there will be more and more HCl in the vapor, but the relative amount will never be more than the azeotropic percentage. A 0.1 M solution easily can be boiled down to 1 M without any noticeable loss of HCl.


Actually, believe it or not, adding the NaCl/FeCl3 salts as I recommended, and distilling is claimed not to be limited to azeotropic concentration. Look it up.

woelen - 16-4-2013 at 22:39

I think that adding NaCl + FeCl3 to such a dilute solution is just spoiling things. If this is done, then ALL of the HCl must be distilled over into another liquid. If you just boil down the liquid then no HCl needs to be distilled over into another liquid.

Maybe the NaCl/FeCl3 method works with moderately concentrated HCl, but here we are talking about 0.1 M HCl.

Personally I would go for NaCl + NaHSO4 and lead the HCl-fumes from the heated mix through distilled water. The 0.1 M HCl is so dilute that it hardly is worth the effort and energy to gain back the HCl.

amateurawesomeness - 17-4-2013 at 10:57

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]

subsecret - 14-12-2013 at 09:32

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!

confused - 14-12-2013 at 09:48

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]

blogfast25 - 14-12-2013 at 14:26

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.

AJKOER - 22-12-2013 at 17:34

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.


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]

subsecret - 22-12-2013 at 20:47

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?

AJKOER - 23-12-2013 at 06:51

Quote: Originally posted by Awesomeness  
....

Could the concentration of HCl in the azeotrope be raised by distilling over concentrated sulfuric acid?


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]

bfesser - 23-12-2013 at 07:24

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]

blogfast25 - 23-12-2013 at 10:31

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]

AJKOER - 24-12-2013 at 18:11

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]

blogfast25 - 25-12-2013 at 05:21

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).

Random - 25-12-2013 at 05:27

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.

blogfast25 - 25-12-2013 at 05:29

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.


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]

bfesser - 25-12-2013 at 06:59

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.

AJKOER - 25-12-2013 at 07:37

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]

blogfast25 - 25-12-2013 at 08:15

Quote: Originally posted by bfesser  
If you have something to add, write it out.


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.