Sciencemadness Discussion Board - Making Hydrochloric Acid

Making Hydrochloric Acid

K12Chemistry - 15-2-2013 at 10:14

Hi,

I want to make hydrochloric acid by mixing sulfuric acid and sodium chloride but I only have 64% sulfuric acid(titrated by me). Do I HAVE to concentrate the acid or can I use the dilute version as long as I get the right stoichometry.

Can someone tell me the stoichometry because I have no Idea how to calculate it. I have 250ml of 64% sulfuric acid. How much salt?

elementcollector1 - 15-2-2013 at 10:25

You don't have to concentrate the acid, but yes, stoichiometry is heavily recommended.
http://theodoregray.com/PeriodicTable/MSP/BalanceReactions
Enter the equation, with HCl and K2SO4 as the products.

Diablo - 15-2-2013 at 10:34

I don't know about the concentration but I can help with the stoichiometry.
First 0.64 * 250=160 ml of acid The reaction NaCl + H2SO4-> HCl + NaHSO4 Now we multiply the volume of acid by its density of 1.84 to get 160*1.84=294.4 grams acid. Molar mass of sulfuric acid is 98 grams per mole 294.4/98= about 3 moles so 3 moles of nacl=58.44 grams per mole 58.44 * 3 =175.32 grams sodium chloride.
Edit: what he said about the concentration.

[Edited on 2-15-2013 by Diablo]

K12Chemistry - 15-2-2013 at 11:05

thank you diablo and elementcollector,

blogfast25 - 15-2-2013 at 11:41

Actually, it really works well only with at least 95 % H2SO4.

If you can't get that (Draino!), try anhydrous sodium bisulphate (NaHSO4, aka sodium hydrogen sulphate): fusing NaCl with NaHSO4 gives you dry HCl:

NaCl + NaHSO4 === > Na2SO4 + HCl

But some fumes of SO3 are likely to come over too, if you overdo it on the temperature...

[Edited on 15-2-2013 by blogfast25]

Adas - 15-2-2013 at 12:59

Quote: Originally posted by blogfast25

Actually, it really works well only with at least 95 % H2SO4.

If you can't get that (Draino!), try anhydrous sodium bisulphate (NaHSO4, aka sodium hydrogen sulphate): fusing NaCl with NaHSO4 gives you dry HCl:

NaCl + NaHSO4 === > Na2SO4 + HCl

But some fumes of SO3 are likely to come over too, if you overdo it on the temperature...

[Edited on 15-2-2013 by blogfast25]

I tried this with normal table salt, and the salt did carbonize, lol. It must have been contaminated with organics and the SO3 dehydrated them. BIG FAIL

but at the same time a good experience. Now I know that table salt is not just NaCl...

DraconicAcid - 15-2-2013 at 13:06

Quote: Originally posted by Adas

Quote: Originally posted by blogfast25

I tried this with normal table salt, and the salt did carbonize, lol. It must have been contaminated with organics and the SO3 dehydrated them. BIG FAIL

but at the same time a good experience. Now I know that table salt is not just NaCl...

I'm surprised that table salt would contain enough other stuff to carbonize with sulphuric acid. You might get iodine from the iodized salt...

Next time, recrystallize the table salt first. That should get rid of any organics (along with the iodides), and give purer HCl.

You didn't take it right out of the salt shaker, did you? Sometimes people add rice to keep it from clumping. Or maybe you accidentally grabbed the sugar container?

blogfast25 - 15-2-2013 at 13:21

Table salt is often adulterated with anti-caking agents, driers etc. Use lab grade or with kitchen grade, dissolve, filter and recrystallise as pure NaCl...

Endimion17 - 15-2-2013 at 13:36

holy shit
http://www.sciencemadness.org/talk/search.php?token=&srchtxt=hydrochloric+acid&srchfield=subject&srchuname=&f[]=all&srchfrom=0& ;filter_distinct=yes&searchsubmit=Search

[Edited on 15-2-2013 by Endimion17]

K12Chemistry - 15-2-2013 at 15:10

I am going to try it with the dilute sulfuric acid anyway. Why wouldn't it work? Also how much water should I dissolve the HCl gas in to get concentrated hydrochloric acid.

Is it worth boiling the sulfuric acid to concentrate and then making HCl or should I just use dilute?

thanks

Vargouille - 15-2-2013 at 15:28

If you try it with dilute sulfuric acid, you might get the 20.2% HCl/H2O azeotrope. Higher concentrations of sulfuric acid will result in higher concentrations of HCl coming over. If you calculate it just right, you can get 37% coming over at somewhere between 61 and 48C. You can check the physical properties section of the Wiki page for hydrochloric acid for more information.

AJKOER - 15-2-2013 at 19:30

Preparation of HCl

In place of H2SO4, consider using Oxalic acid, H2C2O4. React with any convenient chloride, as for the most part, there are all not very soluble. Also, can use solid chlorides (like NaCl per Watt's) to produce HCl gas.
---------------------------------------

Don't have pure H2SO4 or H2C2O4, generate Chlorine gas (at least one old thread on paths). To a vessel containing the Cl2, add a little cold distilled H2O and let seat in direct sunlight. Reactions:

Cl2 (g) + H2O <--> HOCl + HCl

2 HOCl --uv--> 2 HCl + O2 (g)

Note, in a closed vessel, the Oxygen formation will increase pressure. An idea I had, insert an open beaker in the vessel containing the chlorine water, with say aqueous FeCl2 (conveniently but slowly prepared by the action of H2CO3 on NaClO to which Fe is added forming a pale green solution). The O2 will increase pressure moving the 1st reaction to the right. Slowly over days, however, the O2 will react with the aqueous FeCl2 (changing color of the solution) removing O2 from the system. The reaction (see "Hydrometallurgy in Extraction Processes, Vol I, by C. K. Gupta and T. K. Mukherjee at http://books.google.com/books?id=F7p7W1rykpwC&pg=PA203&lpg=PA203&dq=FeCl2+%2B+O2+%3D+FeO(OH)+%2B+FeCl3&source=bl&ots=fiWLs05y8f&am p;sig=mi-pV94woVj7JABKBBzLZcqbEwM&hl=en&sa=X&ei=ZQgfUeq7BIi50AGynoDoBQ&sqi=2&ved=0CFAQ6AEwBA#v=onepage&q=FeCl2%20%2B%20O2%20%3 D%20FeO(OH)%20%2B%20FeCl3&f=false ):

6 FeCl2 + 3/2 O2 + H2O --> 2 FeO(OH) + 4 FeCl3

and, upon raising the pH of the FeCl3 (aq), say upon dilution, the same source (equations 87 and 88) gives:

4 FeCl3 + 8 H2O --> 4 FeO(OH) + 12 HCl

FeCl3 + 3 H2O --> Fe(OH)3 + 3 HCl

so dilute HCl can be recovered/prepared. Note, a potential complication of the FeCl2 synthesis occurs if the Chlorine Bleach product contains Na2CO3 in excess of the NaOCl.

[Edited on 16-2-2013 by AJKOER]

mr.crow - 15-2-2013 at 20:01

Can't you get HCl from the hardware store?

The azeotrope of HCl and water is 20.2% at 110 deg C. Water and weak HCl will distill over until this temperature is reached. Your weak acid should be fine if used in excess.

You need a proper distillation aparatus. HCl is very unpleasant for something so common. I haven't done it myself so I can only give general advice.

If you have high school chemistry education you already know enough about stoichiometry calculations. PM me if you need any more help

K12Chemistry - 16-2-2013 at 09:30

If I do get weak HCl then can I concentrate it?

I also tried to boil down my sulfuric acid but when I added it to sugar nothing happened. Do I need to boil more?

In NurdRage's method he uses 30ml of sulfuric acid. So even though my sulfuric acid will be about 98% at 100ml, I am going to boil to 30ml to GUARANTEE that I will get conc. sulfuric acid and I will have enough to make HCl

AJKOER - 16-2-2013 at 10:12

Quote: Originally posted by K12Chemistry

If I do get weak HCl then can I concentrate it?

Actually and surprisingly, you can make dilute HCl act like concentrated hydrochloric acid using a chloride system (and no, I am not on drugs). See Hydrometallurgy in Extraction Processes, Vol I, by C. K. Gupta and T. K. Mukherjee, page 15 at http://books.google.com/books?id=F7p7W1rykpwC&pg=PA203&lpg=PA203&dq=FeCl2+%2B+O2+%3D+FeO(OH)+%2B+FeCl3&source=bl&ots=fiWLs05y8f&am p;sig=mi-pV94woVj7JABKBB zLZcqbEwM&hl=en&sa=X&ei=ZQgfUeq7BIi50AGynoDoBQ&sqi=2&ved=0CFAQ6AEwBA#v=snippet&q=Magnesium%20chloride%20MgCl2&f=false .

The trick is used by hydrometallurgists as dissolving metal ore with conc acid is expensive. The author says data confirms that a 2M HCl in 3M CaCl2 or MgCl2 (or FeCl3) behaves like 7M HCl.

So you may wish to try this in place of concentrating depending on the application. Note, there is an apparent more limited improvement in the presence of a monovalent chloride (like NaCl).
--------------------------------

The apparent increase in ionic strength of HCl in the presence of divalent chloride may be an explanation of why Cl2 is observed to be produced by the action of CO2 on moist Bleaching powder (Ca(OCl)2 and CaCl2). The action of CO2 in a small amount of water on Ca(OCl)2 could form unstable HOCl (in the conc range of 30% to 50% or more). Some of Hypochlorous could decompose/disproportionate, especially in light, to HCl, O2 and perhaps, a small amount of HClO3, as well. The HCl is assumed to be, in the presence of the divalent chloride, active enough to drive the reaction HCl + HOCl to Cl2 (g) and H2O, which in the presence of CO2...and the reaction chain continues, well at least in my speculation. Note, this Chlorine production (in accordance with the mechanics of the chloride system) is not similarly observed with NaClO or KClO (where more conc HCl is required to liberate Chlorine).

[Edited on 17-2-2013 by AJKOER]

blogfast25 - 16-2-2013 at 11:40

Quote: Originally posted by AJKOER

So you may wish to try this in place of concentrating depending on the application. Note, there is an apparent more limited improvement in the presence of a monovalent chloride (like NaCl).

[Edited on 16-2-2013 by AJKOER]

And contaminating everything you use it for with Ca, Mg or Fe(III)? No but thanks!

AJKOER - 16-2-2013 at 16:13

But Blogfast, if you have to choose between using Muriatic acid and a dilute, but otherwise pure, HCl with an introduced Calcium impurity (from CaCl2) not expected to be overly problematic for the synthesis in question, which would you select?

Also, if acid strength has to be increased to generate a gas (Cl2, ClO2, SO2, H2S...) does a Ca impurity in the HCl matter?

What if the reaction mixture already has a metal X salt, and there exist a not monovalent soluble chloride of metal X, which you have or can easily prepare, would you consider a chloride system to increase acid strength?

What if you have to prepare a certain quantity quickly, but do not have sufficient HCl, but could employ a chloride system, what would you do?

Perhaps yes to some of the questions. But, after all, it is just an option, not a mandate.

Poppy - 16-2-2013 at 18:25

Quote: Originally posted by blogfast25

Blogfast25, as for that retort which HCl come by, is it stainless steel pan enabled?

zed - 17-2-2013 at 15:40

Use rock salt.

repo1030 - 17-2-2013 at 17:46

Just use kosher or popcorn salt. This should be pure NaCl.

blogfast25 - 18-2-2013 at 06:08

Quote: Originally posted by Poppy

Blogfast25, as for that retort which HCl come by, is it stainless steel pan enabled?

Huh? Please rephrase, I don't know what you mean.

Poppy - 19-2-2013 at 08:26

Quote: Originally posted by blogfast25

Quote: Originally posted by Poppy

Blogfast25, as for that retort which HCl come by, is it stainless steel pan enabled?

Huh? Please rephrase, I don't know what you mean.

Errh, sorry Blogfast.
Is this set of NaHSO4 and NaCl you proposed enabled to be produced with stainless steel apparatus? Or would it corrode the pan so bad? As far as it look the pH would be buffed!

elementcollector1 - 19-2-2013 at 09:35

Glass works fine.

blogfast25 - 19-2-2013 at 12:37

Quote: Originally posted by Poppy

Errh, sorry Blogfast.
Is this set of NaHSO4 and NaCl you proposed enabled to be produced with stainless steel apparatus? Or would it corrode the pan so bad? As far as it look the pH would be buffed!

As EC1 said: glass is fine. SS will be attacked and some FeCl3 may come wafting over with the HCl. Having said that, H2SO4 + NaCl + heat in cast iron kettles is how they used to produce HCl industrially. And it was a bit yellow because of the FeCl3 in it...

[Edited on 19-2-2013 by blogfast25]

AJKOER - 22-2-2013 at 10:23

Quote: Originally posted by AJKOER

Quote: Originally posted by K12Chemistry

If I do get weak HCl then can I concentrate it?

However, if you still wish to concentrate HCl via distillation, apparently even a very dilute HCl solutions will work if you know how. Basically add MgCl2 and if very dilute, a little FeCl3. For more details see http://www.neoferric.ca/documents/Harris%20et%20al%20Iron%20... . To quote:

"Figure 4 shows the OLI simulation of the pure HCl-H2O system (it should be noted
that virtually identical values are obtained from the Chemical Engineers Handbook [35] and
the commercially available modelling software AspenPlus® [36]). The data predict that it is
not possible to distill HCl from dilute solutions (i.e. solutions below the concentration of the
azeotrope, 20% HCl). On the other hand, when magnesium chloride is added, the situation
changes quite dramatically as shown in Figure 5. In this case, it is clear that it should be
possible to distill HCl at any liquid phase concentration from magnesium chloride solutions
of concentration greater than about 2.5 m. When ferric chloride is also present, as shown in
Figure 6, then it is apparent that at high concentrations of magnesium chloride, relatively
concentrated HCl should be recovered even from brine solutions containing very little free
HCl."

Also:

"The results (Figure 7) confirmed the OLI simulation, showing
that in strong magnesium chloride solution, the free HCl was flashed off very quickly at
close to the boiling point. Interestingly, the initial concentration in the distillate was close to
36% HCl, and overall, approximately 80% of the free HCl in the solution was recovered at a
cumulative concentration in excess of the azeotrope."

[Edited on 22-2-2013 by AJKOER]

elementcollector1 - 22-2-2013 at 10:36

So, on a scientific note, what happens if you add a chloride salt to max concentration (e.g. 37%) acid? Does it behave like even more concentrated than 37%?
Also, does this process work for sulfates?

blogfast25 - 22-2-2013 at 14:15

Quote: Originally posted by elementcollector1

So, on a scientific note, what happens if you add a chloride salt to max concentration (e.g. 37%) acid? Does it behave like even more concentrated than 37%?
Also, does this process work for sulfates?

As it happens, quite a few chlorides are insoluble in conc. HCl. AlCl3 (as a hydrate) precipitates on gassing a strong solution with HCl (to saturation). Zirconyl chloride behaves also like that.

Also the formation of chloride complexes must be considered in some cases: copper, lead and silver form anions like CuCl4(2-) (tetrachlorocuprate anions).

Re. sulphates, systems like an acid + salt of the same acid + water should really be considered three phase systems. The solubilities of the two solutes can be surprising. BeSO4 is highy soluble in water (up to 44 % as BeSO4) at 20 C but it is almost insoluble in concentrated H2SO4.

[Edited on 22-2-2013 by blogfast25]

DraconicAcid - 22-2-2013 at 14:19

Quote: Originally posted by blogfast25

Quote: Originally posted by elementcollector1

So, on a scientific note, what happens if you add a chloride salt to max concentration (e.g. 37%) acid? Does it behave like even more concentrated than 37%?
Also, does this process work for sulfates?

As it happens, quite a few chlorides are insoluble in conc. HCl. AlCl3 (as a hydrate) precipitates on gassing a strong solution with HCl (to saturation). Zirconyl chloride behaves also like that.

So does sodium chloride. If you add a few drops of conc. HCl(aq) to a solution of saturated sodium chloride, you get salt precipitating out.

AJKOER - 22-2-2013 at 15:03

Quote: Originally posted by elementcollector1

So, on a scientific note, what happens if you add a chloride salt to max concentration (e.g. 37%) acid? Does it behave like even more concentrated than 37%?
Also, does this process work for sulfates?

Not sure to be honest, notice that a lot of this work is based on simulations and relatively recent. I would, however, not be surprised of an acid could have an 'activity' level that is sustained in a reaction that would otherwise not be observed with an acid starting at its max concentration. This, I would guess, could occur as the 'activity' level of H2O, normally around 1, could be at a much lower value (apparently somewhere between 0.3 and 0.4, for example, for very conc brine solutions which is far from unity).
-----------------------------------------------

To answer your second question, maybe, check out thiosulphate leaching. To quote " Ammoniacal thiosulphate solution allows the solubilization of gold as stable anionic complex. Leaching of gold occurs at appreciable dissolution rates." See http://www.sciencedirect.com/science/article/pii/0304386X950... .

[Edited on 22-2-2013 by AJKOER]

AJKOER - 23-2-2013 at 07:44

For those interested in MgCl2 in cases where CaCl2 presents solubility issues, can prepare MgCl2 as follows:

MgSO4 + CaCl2 --> CaSO4 (s) + MgCl2

Don't have CaCl2, make Magnesium hypochlorite (from NaOCl + NaCl + MgSO4 actually forms the dibasic form, Mg(OCl)2.2Mg(OH)2 ) and then add H2O2, or heat to disproportionate into MgCl2 and MgClO3.

Why is the hydrochloric acid green ?

experimenter - 18-8-2014 at 08:51

I have tried to make hydrochloric acid by distilling table salt and concentrated sulfuric acid (battery acid). I had few mililiters of deionised water in the receiver to catch possible HCl fumes.

The experiment was a success, but the acid produced has a green color(?). Please, do you know if this is normal? I think it should be transparent.

Sometimes commercial hydrochloric acid is green due to iron impurities, but
in my case, distillation should have left back possible iron salts.

Another possibility is that the hydrochloric acid is contaminated with Cl2 gas. How can I find out if this is the case? How can I remove it? Thanks.

Zyklon-A - 18-8-2014 at 10:51

Perhaps the chloride ion was oxidized by an impurity in your battery acid (or salt).
If your sulfuric acid was pure, chlorine would not form.
I think one way to clear the color, is to put it in sunlight for a few hours.
AJKOER, I'm sure could explain how this works better than I could.

[Edited on 19-8-2014 by Zyklon-A]

experimenter - 18-8-2014 at 11:51

Thanks for your response. In an attempt to find the "culprit", I tried two experiments.

- I distilled commercial, low grade hydrochloric acid mixed with table salt. The green color appeared again in the distillate, especially during the end of the distillation.

- Distilled the commercial, low grade hydrochloric acid by itself in order to check for impurities in it. No green color appeared in the distillate. It left behind a very small amount of whitish salt only.

It seems that the impurity came from the salt. Some chemical that can volatilize with HCl?

Texium - 18-8-2014 at 12:31

Quote: Originally posted by experimenter

Thanks for your response. In an attempt to find the "culprit", I tried two experiments.

- I distilled commercial, low grade hydrochloric acid mixed with table salt. The green color appeared again in the distillate, especially during the end of the distillation.

- Distilled the commercial, low grade hydrochloric acid by itself in order to check for impurities in it. No green color appeared in the distillate. It left behind a very small amount of whitish salt only.

It seems that the impurity came from the salt. Some chemical that can volatilize with HCl?

If you think it might be the salt, try to find some picking salt. It doesn't have additives like iodine compounds and free-flowing agents in it. Probably still wouldn't be as good as real reagent grade NaCl, but it's a lot better than normal salt.