Sciencemadness Discussion Board - Making hydrochloric acid

Making hydrochloric acid

liquidlightning - 20-5-2012 at 23:50

Is there any simple way to make hydrochloric acid at home? Unfortunately I don't have access to any commercially. How about electrolyzing a saltwater solution in a gas gen and routing the gas into distilled water?

UKnowNotWatUDo - 21-5-2012 at 01:09

What about just buying hardware store muriatic acid?

liquidlightning - 21-5-2012 at 02:01

How pure is hardware store acid?

UKnowNotWatUDo - 21-5-2012 at 02:40

There are many different brands and each has a different concentration of HCl and different additives. I would look around your local stores and see what brands are available to you. Then you can look up the MSDS sheets online to get an idea of what's in it. Most are in the range of 5-10 molar.

woelen - 21-5-2012 at 03:05

Hardware store muriatic acid is suitable for most of all home chemistry experimenting. It comes in 10% by weight HCl in many countries, sometimes one can find 30% HCl. The hardware store material frequently is pale yellow in color, this is due to some impurity, mostly iron(III), which gives the yellowish color to the acid. Sometimes the impurity is of organic nature.

I have tried distilling muriatic acid from hardware stores and my experience is that this gives really nice quality colorless acid. The 10% acid can be boiled off to appr. 60% of its original volume. Almost pure water boils off. The remaining part can be distilled such that only 10% of it remains and then you end up getting 15% or so acid, which is totally colorless. The 10% remains will have a fairly dark green/brown color and should be discarded.

if you start from 30% acid, then add appr. half the volume of water to this acid, mix well and distill until only appr. 5% of the liquid is left. You'll get very pure approximately 20% hydrochloric acid. Do not distill 30% acid. I tried once, but LOTS of choking HCl fumes escape and these do not liquefy in the receiving flask.

[Edited on 21-5-12 by woelen]

hissingnoise - 21-5-2012 at 03:21

Quote:

How about electrolyzing a saltwater solution in a gas gen and routing the gas into distilled water?

Electrolysing brine produces hypochlorite in solution, although minute quantities of HCl escape!
You could go the NaCl/H2SO4 route using the inverted funnel 'trick'!
Distilled water will give pure, strong HCl and you might find watching the visuals of the higher density solution HCl falling from the underside of the inverted funnel, er, fascinating for a while . . .

woelen - 21-5-2012 at 04:22

If you cannot get any liquid acid at all, then you can use table salt and so-called granular pH-minus for swimming pools to make HCl. Usually, pH-minus is NaHSO4.H2O (sodium bisulfate monohydrate). You really must have this sodium bisulfate monohydrate and not some liquid. Anhydrous NaHSO4 also is not what you want, because of its high meltingpoint.

Mix pH-minus and NaCl in roughly 3 to 1 weight ratio and heat in a glass flask. Soon, the NaHSO4.H2O liquefies (it seems to melt, but in reality it dissolves in its own water of hydration and it does so at relatively low temperature) and humid HCl fumes escape from the flask. These fumes can be passed through water to make hydrochloric acid. Be careful not to have suckback of water into the flask with the NaHSO4.H2O/NaCl mix.

elementcollector1 - 21-5-2012 at 05:18

I heard of a method somewhere to sort of purify hardware store grade by placing a beaker of the stuff next to a beaker of distilled water inside a sealed container, and waiting for the HCl fumes to escape from one and be absorbed by the other, thereby diluting each to half the original concentration. Is this true?

woelen - 21-5-2012 at 05:53

If you place very concentrated acid, which gives off a lot of HCl fumes next to a beaker full of water, then indeed some HCl is transferred in this way, but the process quickly comes to a near halt when the concentration in the original acid goes below 25% of HCl. More and more water will evaporate as well and the process goes very very slow. So, if you are very patient and want to wait days or even weeks, then you can transfer some of the acid to another beaker. In the long run, vapor will be everywhere in the container and there will be an equilibrium with liquid on the floor and liquid in the beaker. So, in theory you even can do with one beaker, which is made very clean on its outside and you put this in a clean and sealed container. When the floor of the container is put on relatively cold soil and the rest of the container is heated somewhat, then liquid will collect on the floor of the container. This liquid will be pure HCl. This process, however, also will be very slow.

DJF90 - 21-5-2012 at 07:05

RBT demonstrates this process on youtube as one of his videos for his amatuer chemistry book. You may want to check out the rest of the videos on his channel..

http://www.youtube.com/watch?v=jv1Ms6Subg4

Formatik - 21-5-2012 at 11:54

Quote: Originally posted by woelen

...

Mix pH-minus and NaCl in roughly 3 to 1 weight ratio and heat in a glass flask. Soon, the NaHSO4.H2O liquefies (it seems to melt, but in reality it dissolves in its own water of hydration and it does so at relatively low temperature) and humid HCl fumes escape from the flask. These fumes can be passed through water to make hydrochloric acid. Be careful not to have suckback of water into the flask with the NaHSO4.H2O/NaCl mix.

Another simple way to get hydrochloric acid is through MgSO4 and NaCl as mentioned in the primordial chemicals thread. The same precautions as by NaHSO4 and NaCl e.g. suckback would also apply here.

elementcollector1 - 21-5-2012 at 13:51

Wouldn't that just make magnesium chloride and sodium sulfate?

barley81 - 21-5-2012 at 16:33

Quote: Originally posted by elementcollector1

Wouldn't that just make magnesium chloride and sodium sulfate?

These are the posts from the primordial chemicals thread.

First one:

Quote: Originally posted by 12AX7

Quote:

Originally posted by neutrino
One I accidentally discovered: mixing calcium chloride and magnesium sulfate solutions, then heating the resulting paste to decomposition releases HCl. It goes something like this:

Ca2+ + SO42- --> CaSO4

Drying, we get solid MgCl2 . 6H2O. Finally,

2MgCl2 . 6H2O --heat --> Mg2OCl2 + H2O + HCl

Good point: some chlorides decompose a significant amount. Alkaline earth and transistion metal chlorides (and especially aluminum chloride hydrate) decompose, giving off HCl fumes with some amount of H2O, allowing muriatic acid to be distilled directly in low yield.

Sulfuric acid of course is had in higher yield from anhydrous iron, copper, or to a lesser extent zinc, sulfates.

Cinnabar is the primary mercury ore, you just have to find it -- mercury isn't very common in the Earth's crust after all, but just like silver, it's there. Mercury (elemental) and I think the oxide are also present to some extent.

Chalk, limestone and dolomite: all the calcium oxide and carbon dioxide you could hope for, most formations are several feet thick (up to a few hundred) and span thousands of miles. After calcining to yellow heat for an hour or three (big pieces may need days, cleaving into 1" slabs would help here) you're left with CaO and a lot of CO2 out the stack, which can be pumped and compressed, or bubbled into a solution for collection: read up on the Solvay process, which produces CaCl2 as a byproduct for well, given the materials CaCO3 + NaCl, I think you can guess why.

Oh, and let's not forget CaO is also the primordial alkali. Why?

The magnesium ought to be leachable, perhaps by dissolving a block of dolomite and precipitating CaCO3 with MgCO3 (from more dolomite? Dunno) to give seperate Ca and Mg (salt) products.

But this is kind of off the direction... a great conversation as well but I was thinking more what you do with it. (Or maybe I'm wasting my time explaining; this is an interesting enough angle, at least.) Like, my oxidizer example: besides oxygen in air, about the only natural thing you have is MnO2, which can oxidize Cl- to Cl2. The MnCl2 can be precipitated and reoxidized with oxygen and fire or weathering and time (as happens naturally). That Cl2 gas can go on to do just about anything, up to and including things like permanganate (which can also be made from pyrolusite and a caustic fusion, in air), perchlorate, ferrate and so on.

And heck, primordial electrochemistry is worth thinking about, too. It's a lot of effort to mine, roast, smelt, distill and cast zinc anodes, but it can be done. Copper can be mined, roasted and smelted with a bit more ease, though it needs more fire to cast it (not a problem for a firetender such as myself

). Electrolyte, well that can be made from whatever, be it acid, base or a salt. Given enough surface area and a few cells, you can do all the standard electrochemistry, well assuming you can isolate platinum,

and make chlorates, persulfates and the ever most venomous fluorine, as well as the strongest reducers, the alkali metals (which can in fact be isolated by carbothermic processes!).

Organic chemistry of course all starts with organic chemicals, since it's a waste to start with CO2 and there's so much plant and animal life available to the desperate Mad Scientist. Finding the source of reagents (oh my, and glassware!

) for these synthesis might prove an interesting angle however.

Tim

Second one:

Quote: Originally posted by Formatik

Quote:

I've recently done this using NaCl. The decomposition equation given in the Handbook of Inorganic Chemicals by Pradyot Patnaik forming also the basic salt is: MgCl2.6 H2O -> Mg(OH)Cl + HCl + 5 H2O.

I mixed powdered MgSO4.7 H2O and NaCl into a paste using water. And then heated them on a hotplate, after the water evaporated, and heating continued significant amounts of HCl evolved, also recognized by red litmus and the irritating odor.

But anyone have ideas on how to further work up the basic salt and use up the other Cl?

[Edited on 22-5-2012 by barley81]

AJKOER - 21-5-2012 at 21:38

First, make Chlorine. There is at least one prior Sciencemadness thread on ways to prepare Cl2. I would suggest NaHSO4 (use to control pH for a Spa) plus Bleach (NaOCl and NaCl). Reaction sequence:

NaHSO4 + NaCl + H2O --> Na2SO4 + HCl + H2O

NaHSO4 + NaOCl ---> Na2SO4 + HOCl

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

More exotic, FeSO4 and Bleach, has been reported in that reference thread (Link: https://www.sciencemadness.org/whisper/viewthread.php?tid=13... ).

Second, allow the generated Cl2 to dissolve in cold water.

Cl2 + H2O --> HOCl + HCl

Third, let sunlight convert the HOCl to HCl:

HOCl --uv--> HCl + O

You may also get some HClO3 as an added kicker to your HCl solution as:

3 HOCl ---Diffused Sunlight---> 2 HCl + HClO3

So, depending on the strength of your solution, beware of any ClO2 (explosive) gas formation as:

6 ClO2 + 3 H2O <---> 5 HClO3 + HCl

Best to use an open container preferably outdoors as ClO2 is also more toxic than Cl2.

woelen - 21-5-2012 at 22:39

AJKOER, your method does not sound like a very useful method of making HCl. If you use NaHSO4, then why not simply add NaCl and collect the HCl from that. The reaction you mention is a very slow one and Cl2 does not dissolve that good in water (one liter of cold water dissolves appr. 3 liters of Cl2 gas, which is only a few grams of gas). Making concentrated acid is very hard in this way, because the solubility of Cl2 even becomes lower when HCl is dissolved in the water.

Then I also want to make a remark about toxicity of ClO2. This is much less toxic than Cl2. But you hardly need to fear formation of ClO2. If any is formed, then the concentration will be so low that it is hardly visible.

-----------------------------------------------------------------------------

If you want to watch a funny effect then do the following:
- Prepare a liter or so of Cl2 gas.
- Bubble the Cl2 gas under water in a 500 ml bottle of clear colorless glass with a decent screw cap, such that it is completely filled with Cl2 gas and the only water must be the droplets still sticking to the glass.
- Take a magnifying glass and focus sunlight, such that the point of focus is inside the chlorine gas.

At the focus point you see formation of a dense white fume. Most likely this is due to reaction of Cl2 with water vapor with formation of HCl and O2. The HCl, formed at that spot is responsible for the fumes.
This reaction is not useful for preparative purposes, but it is fun to see the formation of fumes apparently out of nothing.

[Edited on 22-5-12 by woelen]

AJKOER - 22-5-2012 at 15:02

Woelen:

Thanks for reviewing my suggested HCl synthesis. Yes, if one has NaHSO4 then reacting it with NaCl and condensing the vapors (assuming one has the apparatus to perform this operation), is indeed a much more straightforward path to even more concentrated HCl. However, I only used this as an example of one of many possible paths to forming Chlorine, and not necessarily assuming the available of this resource or of the equipment so needed.

Once one has generated Cl2, bubbling into a solution is not the only (albeit also, I agree, inefficient) way to proceed. I was thinking again along more simple apparatus route like having the Cl2 generator container in a larger covered vessel containing a small amount of cold distilled water. With time, the Cl2 is absorbed forming:

Cl2 + H2O <----> HOCl + HCl

Or better, possibly adding a tiny amount of Sulfur, which would form SO2 upon reaction with the HOCl moving the reaction above to the right. Reaction sequence:

2 HOCl + S --> 2 HCl + SO2

SO2 + H2O + HOCl ---> H2SO4 + HCl

Normally, in the next step with sunlight:

HOCl --uv--> HCl + O

and more rapidly, in the presence of a catalyst, like Tartaric or Citric acid, see "A comprehensive treatise on inorganic and theoretical chemistry", Volume 2, by Joseph William Mellor, top of page 82, to quote: "According to C. Lowig,{27} bromine water in light behaves in a similar way to that of chlorine water, but as J. M. Eder showed, bromine water is much less sensitive to light in that it decomposes with but one-sixth or one-twelfth the speed of chlorine water. The presence of tartaric or citric acid accelerates the decomposition of chlorine or bromine water in light." Link: http://books.google.com/books?pg=PA82&lpg=PA82&id=An...

Note, water is consumed both by the reaction of Chlorine and water, and also by the reaction of Hypochlorous acid and SO2 & water.

Source: Per Watts' Dictionary of Chemistry Volume 2 Page 16, even dilute solution of HClO can oxidize Sulfur all the way to H2SO4. To quote the relevant section from Watts':

"Reactions.--1. HClOAq acts generally as an oxidiser; it easily parts with 0 while HClAq remains. Thus, As is rapidly oxidised with evolution of light; P, S, Se, Br, I are converted to H3P04Aq, H2S04Aq, &c., even by dilute HClOAq; lower oxides or salts are converted into higher, e.g. SO2Aq to H2SO4Aq, FeO to Fe203, As203Aq to As2O5Aq, FeS04Aq to Fe2(S04)3Aq, Fe2Cl6Aq, and Fe2O3, MnSO4Aq to MnO2; sulphides yield sulphates, c.g. H2SAq gives" H2SO4,Aq and S; "
Link:
http://books.google.com/books/reader?id=ijnPAAAAMAAJ&dq=...
-------------------------------------------------------

On the comparative toxicity of Cl2 vs ClO2 with respect to inhalation, I believe my comment on the greater toxicity of Chlorine dioxide is valid based on a comparison of several MSDS listings showing the recommended thresholds for ClO2 at a fraction for those of Cl2. For example:
EXPOSURE LIMITS:
CHLORINE:
1 ppm (3 mg/m3) OSHA ceiling
0.5 ppm (1.5 mg/m3) OSHA TWA (vacated by 58 FR 35338, June 30, 1993)
1 ppm (3 mg/m3) OSHA STEL (vacated by 58 FR 35338, June 30, 1993)
0.5 ppm ACGIH TWA
1 ppm ACGIH STEL
0.5 ppm (1.45 mg/m3) NIOSH recommended ceiling 15 minute(s)
Link: http://www.clean.cise.columbia.edu/msds/chlorine.pdf

CHLORINE DIOXIDE (10049-04-4)
OSHA Pel 0.100 ppm – TWA
ACGIH TLV 0.100 ppm – TWA
ACGIH TLV 0.300 ppm – STEL
Link: http://www.thesabrecompanies.com/literature/clo2_msds.pdf

[Edited on 23-5-2012 by AJKOER]

hydrothermal regeneration of spent pickle liquor

tetrahedron - 28-9-2012 at 17:29

i'm not talking about an alcoholic drink made from salty cucumbers. in metallurgy, a pickle is a HCl solution used to remove rust from iron. during this process, the solution accumulates iron chloride and loses its strength. the fancy sounding regeneration process involves nothing but (1) boiling the solution in the presence of air (to ensure that all iron turns to the ferric state) and (2) distilling off the hydrochloric acid, leaving insoluble Fe2O3 behind (it's well known that anhydrous FeCl3 hydrolyzes, that's why it cannot be obtained by crystallization). The hydrochloric acid can thus be reused for pickling. The Fe2O3 also finds uses.

http://en.wikipedia.org/wiki/Hydrochloric_acid_regeneration#...

has anyone tried this route to HCl(aq)?

question. for our purpose of HCl preparation we'll need to regenerate the FeCl3 solution. would it suffice to bubble chlorine gas (from a simple electrolytic generator) into a cold aqueous suspension of Fe2O3 (which can be recycled)? my guess is that the iron will catalyze the dissociation of any chlorate species formed.

[Edited on 29-9-2012 by tetrahedron]

IanCaio - 30-9-2012 at 11:17

Sorry if it might sound stupid,
but I was wondering why not getting a low concentration
muriatic acid, and heat it so the HCl would go out and make
this gas go through a beaker with less water then the original
acid. Wouldnt it work to get a higher concentration and purity
HCl solution? What would be the problem of such process?

Bot0nist - 30-9-2012 at 12:51

Due to the azeotrope, water vapor, and HCl will be released past a certian point.

IanCaio - 30-9-2012 at 16:23

Got it man, thanks

So theres no viable way to concentrate a low concentrated
Hydrochloric acid? Maybe using vacuum to remove the HCl?
But this would bring the problem of transfering the gas to the
right place I guess..
In the worst case, heating it would get me a 20% acid? (The
azeotropic concentration)

Bot0nist - 30-9-2012 at 16:36

Standard Distillation should give the highest concentration possible (azeotrope), or lead evolved HCl gas into ice cold distilled water via inverted funnel trick or other suckback prevention methods. If you cant buy muriatic at around 30% then maybe sulfuric acid + NaCl into water is the best bet.

tetrahedron - 30-9-2012 at 16:36

Quote: Originally posted by IanCaio

So theres no viable way to concentrate a low concentrated Hydrochloric acid?

there are many ways, just not by simple distillation as you describe here:

Quote: Originally posted by IanCaio

why not getting a low concentration muriatic acid, and heat it so the HCl would go out and make this gas go through a beaker with less water then the original
acid

you need to scavenge that water somehow

Quote: Originally posted by Bot0nist

Standard Distillation should give the highest concentration possible (azeotrope), or lead evolved HCl gas into ice cold distilled water via inverted funnel trick or other suckback prevention methods. If you cant buy muriatic at around 30% then maybe sulfuric acid + NaCl into water is the best bet.

'azeotropic' (20%) is not the same as 'concentrated' (37%). in this case consider azeotropic to be low concentration.

[Edited on 1-10-2012 by tetrahedron]

tetrahedron - 1-10-2012 at 15:41

Quote: Originally posted by AJKOER

possibly adding a tiny amount of Sulfur, which would form SO2 upon reaction with the HOCl moving the reaction above to the right. Reaction sequence:

2 HOCl + S --> 2 HCl + SO2

SO2 + H2O + HOCl ---> H2SO4 + HCl

a tiny amount? assuming your second reaction really takes place under the given conditions, for which i haven't found any evidence, the overall equation is

3 Cl2 + 4H2O + S ---> 6 HCl + H2SO4

now, if all the sulfur is present in suspension from the beginning (hence in excess), all you'll get is SO2 evolution.

you need to maintain a stoichiometric ratio for HOCl:S of 3:1 or higher if you wanna give your second equation a fighting chance.

tetrahedron - 23-10-2012 at 01:29

Quote: Originally posted by tetrahedron

http://en.wikipedia.org/wiki/Hydrochloric_acid_regeneration#...

more info on the process, also in relation to other cations (Ni, Cu, Co):

http://www.neoferric.ca/documents/Harris%20et%20al%20Metal%2...

Quote:

2FeCl3 + 3H2O → Fe2O3 + 6HCl

the solution was heated up to 175-180°C at atmospheric pressure, and the
HCl stripped off at a concentration of 30% with >99% recovery

another one:

http://digitool.library.mcgill.ca/webclient/StreamGate?folde...

[Edited on 23-10-2012 by tetrahedron]