liquidlightning
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Making hydrochloric acid
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?
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UKnowNotWatUDo
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What about just buying hardware store muriatic acid?
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liquidlightning
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How pure is hardware store acid?
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UKnowNotWatUDo
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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.
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woelen
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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]
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hissingnoise
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| 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/H<sub>2</sub>SO<sub>4</sub> 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 . . .
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woelen
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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.
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elementcollector1
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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?
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
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woelen
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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.
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DJF90
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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
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Formatik
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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.
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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.
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elementcollector1
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Wouldn't that just make magnesium chloride and sodium sulfate?
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Next in Line: Nd
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barley81
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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:
Ca<sup>2+</sup><sub></sub> + SO<sub>4</sub><sup>2-</sup> --> CaSO<sub>4</sub>
Drying, we get solid MgCl<sub>2</sub> . 6H<sub>2</sub>O. Finally,
2MgCl<sub>2</sub> . 6H<sub>2</sub>O --heat --> Mg<sub>2</sub>OCl<sub>2</sub> +
H<sub>2</sub>O + 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: | 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:
Ca<sup>2+</sup><sub></sub> + SO<sub>4</sub><sup>2-</sup> --> CaSO<sub>4</sub>
Drying, we get solid MgCl<sub>2</sub> . 6H<sub>2</sub>O. Finally,
2MgCl<sub>2</sub> . 6H<sub>2</sub>O --heat --> Mg<sub>2</sub>OCl<sub>2</sub> +
H<sub>2</sub>O + HCl |
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]
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AJKOER
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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.
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woelen
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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]
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AJKOER
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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]
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tetrahedron
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hydrothermal regeneration of spent pickle liquor
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]
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IanCaio
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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?
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Bot0nist
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Due to the azeotrope, water vapor, and HCl will be released past a certian point.
U.T.F.S.E. and learn the joys of autodidacticism!
Don't judge each day only by the harvest you reap, but also by the seeds you sow.
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IanCaio
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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)
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Bot0nist
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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.
U.T.F.S.E. and learn the joys of autodidacticism!
Don't judge each day only by the harvest you reap, but also by the seeds you sow.
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tetrahedron
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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.
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'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]
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tetrahedron
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| 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.
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tetrahedron
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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]
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