Murexide
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Obtaining pure hydrochloric acid
Hydrochloric acid is a commonly available reagent (in Oz) but may contain various degrees of impurities, especially metal (iron) ions.
There are several main methods, as detailed below:
1. Two container equilibration. 10 M HCl is placed in one container, in the other it is filled with distilled water. After several weeks, one
container should contain pure 5 M HCl. Has anyone had success with this method, and how concentrated is the end product? I have a hard time imagining
the water to absorb HCl at an appreciable rate, and the HCl may also escape the larger container.
2. Distillation. This method certainly works and produces 20% azeotropic HCl. However, this method is quite dangerous as HCl fumes are produced in
large quantities, rendering it less suitable for the amateur lab (a boiling fuming vessel of conc. HCl is no joke!).
I propose an untested method: Removal of impurity (esp. iron) by water softening resins. This appears like quite the viable method, however the resin
may reject the metal ions in favour of the acid (H+). The affinity of the resin towards iron is much greater than to the proton, but the concentration
of protons is also much greater. What are your opinions of this?
I have not tested the degree of metal contamination for the hydrochloric acid, but I will do so before persuing any of these options.
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unionised
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If you want to "regenerate" ion exchange resins used for water softening you do it by running HCl through them.
So the reaction you want doesn't happen- it goes the other way.
However, the iron in commercial HCl is actually present as [FeCl4]-- ions and isn't in competition with the H+.
What you need to find out is whether the resins selectively bind Cl- or [FeCl4]--
Incidentally, if you dilute the conc HCl down to just below 20% before you distill it you don't have nearly as much HCl fume to deal with.
[Edited on 6-10-19 by unionised]
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Murexide
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I believe that the water softening resins are sold as primarily strongly acidic (sulphonated groups, cation exchange) and mixed bed (with strongly
basic anion exchange, quarternary amine sites). Unfortunately, I can only imagine the chloride been bound more strongly than the complex (or the
complex decomposing).
Ok, so just did some testing. Diluted the acid to 1%, added it to a test tube with a few crystals of KSCN. Then proceeded to add sodium carbonate to
sequentially raise the pH. Not detectable red colour was observed. Then, in a seperate test tube I added 1 mL 0.0002 M KMnO4 to 1 mL HCl. No
decolourisation occured, suggesting [Fe II] < 0.00005 M. Thiocyanate test has detectable limit around 0.000002 M, hence [Fe III] < 0.000002 M
For the concentrated HCl,
[Fe II] < 0.0015 M, 1.5 mM
[Fe III] < 0.00006 M, 0.06 mM
So far this HCl actually seems quite good. Might do some tests with the concentrated acid and test for combined metal ions, to give an actual number.
Might be good for immediate use 
Otherwise I'll explore the options.
Edit: Just realised that I have no way of detecting sodium contamination. Unfortunately, this is one of the rare instances where it actually matters
(at concentrations > 0.1% or 1g/L), so I might have to attempt purification
[Edited on 6-10-2019 by Murexide]
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j_sum1
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Most times I have bought hardware store Aussie HCl it has appeared remarkably clean and I have had no problens.
My current bottle does have some impurity. Not sure what. But a mauve colour can be seen after reacting with Mg or Zn. I simply distilled some for
when purity is an issue. And I can make HCl using sulfuric acid and salt if I ever need a high concentration. I see little need to devise a new
method using resins.
HCl fumes are a pain when it comes to storage. But in actual use I have found them easy to manage. They scrub easily.
I guess what I am saying is to consider your actual need. There is probably already an acceptable solution through conventional means. By all measns
experiment with ion exchange resins if that is your interest. But if your goal is clean acid then it probably is unnecessary.
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XeonTheMGPony
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If purity is such a concern the only practical solution I see is to make it in known conditions
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rockyit98
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you can use AlCl3 or FeCl3 to make HCL gas you need to heat hydrous AlCl3 or FeCl3.
Al(H2O)6Cl3 → Al(OH)3 + 3 HCl + 3 H2O
FeCl3+H2O →FeOCl+2HCl
you can use dirty HCl to dissolve Al to make AlCl3 solution and heat it to make Al2O3 and pure HCl.
since dehydrochlorination speed depend on temperature and is an equilibrium it is much safer than distillation.
if it a hassle uyou can always distill HCl from battery acid and table salt.
"A mind is a terrible thing to lose"-Meisner
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Loptr
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Dissolve calcium chloride in water, and add H2SO4 to precipitate calcium sulfate to leave behind an aqueous solution of HCl.
You will need high purity H2SO4 to keep from introducing contaminates, but it can be distilled beforehand with care. It requires high temps, so be
careful.
[Edited on 6-10-2019 by Loptr]
"Question everything generally thought to be obvious." - Dieter Rams
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woelen
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I used my 30% hardware store HCl to make appr. 25% very pure HCl.
I have a distillation setup and in the boiling flask I put the 30% HCl. In the receiver flask I put a little water. The receiver flask, I immersed in
a bucket of cold water.
Next, I started heating. Initially, you get quite a lot of fumes of HCl, which are absorbed by the water in the receiver flask. Indeed, I also had
some HCl escaping the flask. I put a long flexible 5 mm PVC tube on the receiver flask, which I lead to outside the window in my lab.
After a while, the temperature in the boiling flask rises and the mix of vapor, going over is more and more like azeotropic acid at appr. 20%
concentration. I did not use cooling in the condenser, just of the receiving flask, but the acid nicely and smoothly ran into the receiver flask. I
continued distillation, until the volume left in the boiling flask was appr. 30% of the starting volume. That liquid I kept for cleaning purposes
(nice for cleaning test tubes in which e.g. metal hydroxides are sticking to the glass). The liquid in the receiver flask is nice and colorless acid,
with a concentration well over 20%. I estimate it to be near 25%. The liquid does not fume visibly in contact with air, but it does have an acrid
smell, so some fumes do come off.
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RogueRose
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Quote: Originally posted by woelen  | I used my 30% hardware store HCl to make appr. 25% very pure HCl.
I have a distillation setup and in the boiling flask I put the 30% HCl. In the receiver flask I put a little water. The receiver flask, I immersed in
a bucket of cold water.
Next, I started heating. Initially, you get quite a lot of fumes of HCl, which are absorbed by the water in the receiver flask. Indeed, I also had
some HCl escaping the flask. I put a long flexible 5 mm PVC tube on the receiver flask, which I lead to outside the window in my lab.
After a while, the temperature in the boiling flask rises and the mix of vapor, going over is more and more like azeotropic acid at appr. 20%
concentration. I did not use cooling in the condenser, just of the receiving flask, but the acid nicely and smoothly ran into the receiver flask. I
continued distillation, until the volume left in the boiling flask was appr. 30% of the starting volume. That liquid I kept for cleaning purposes
(nice for cleaning test tubes in which e.g. metal hydroxides are sticking to the glass). The liquid in the receiver flask is nice and colorless acid,
with a concentration well over 20%. I estimate it to be near 25%. The liquid does not fume visibly in contact with air, but it does have an acrid
smell, so some fumes do come off. |
Would you be able to get a higher concentration by putting fresh HCl (30-31.45%) in the boiling flask and placing the receiver in a salt water ice
bath? Does the temp of the receiver greatly effect how much HCl gas can be held and would a 30-45 deg F drop (from regular ice to salted ice bath)
make much difference?
Also, what about using the plastic bottle the HCl came in and using warm (maybe up to 160-180?) water to increase the speed the gas comes over, would
that be enough to at least get down to 20%?
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unionised
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Yes it does.
And, when the solution warms up you have a pressurised system.
How do you plan to open that safely?
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AJKOER
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Small amounts of pure HCl (at least free from metal impurities) can be obtained in dilute amounts (which could be concentrated) by mixing chlorine gas
(produced in a variety of ways) with H2O2. The latter, for stability reasons, is free of metallic impurities. The associated chemistry:
Cl2 + H2O <--> HCl + HOCl
HOCl + H2O2 --> HCl + H2O + O2 (g)
---------------------------------------------------
Relatedly, a slower process would involve replacing H2O2 with the action of dissolved oxygen ( O2)d ) in the presence of a source of solvated
electrons, e-(aq), from say the UV photolysis of TiO2, as:
O2 (d) + e- (aq) = •O2-
and at pH < 4.88,
•O2- + H+ = •HO2
and the cited reaction:
H+ + •O2- + HOCl = H+ + Cl- + •OH + O2 (g) (Source: see Equation (1) at
https://febs.onlinelibrary.wiley.com/doi/pdf/10.1016/0014-5793(93)80394-A )
•OH + •OH = H2O2
•OH + Cl- = OH- + •Cl
•Cl + •Cl = Cl2
among other reactions.
Note, a combination method would be to add to chlorine/water a small amount of H2O2 and TiO2. Proceed to treat with UV light in a sealed system.
--------------------------------------------------------
My prior cited path (see https://www.sciencemadness.org/whisper/viewthread.php?tid=10... ) to hypochlorous acid, HOCl, with a sodium impurity, is by adding NaHCO3 (Baking
soda) and CaCl2 (DampRid) to NaOCl (found in 6% or higher strength Chlorine Bleach). Reactions:
NaOCl + NaHCO3 = Na2CO3 + HOCl
CaCl2 + Na2CO3 --> CaCO3(s) + 2 NaCl
In the presence of CaCl2, the reaction forms a suspension of CaCO3 which settles on cooling of the mix. Decant to collect the hypochlorous
acid/aqueous sodium chloride mix, which should be kept cool, in the dark, and used within a few hours of preparation.
Interestingly, starting with HOCl, the simple action of sunlight on HOCl will produce HCl, O2 and some HClO3 (as long has been reported, see, for
example, https://books.google.com/books?id=U502AQAAMAAJ&pg=PA329&... ).
Another source notes that in diffused light sunlight the following reaction:
5 HOCl + 5 H2O --> ClO3- + 4 Cl- + O2 + 5 H3O+
Source: https://books.google.com/books?id=Dv_F03cdKPUC&pg=PA448&...
[Edited on 17-10-2019 by AJKOER]
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Sulaiman
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Boil a small sample to dryness and if there is a residue do a flame test,
the yellow flame of sodium is very noticeable.
CAUTION : Hobby Chemist, not Professional or even Amateur
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unionised
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| Quote: Originally posted by AJKOER |
My prior cited path (see https://www.sciencemadness.org/whisper/viewthread.php?tid=10... ) to hypochlorous acid, HOCl, with a sodium impurity, is by adding NaHCO3 (Baking
soda) and CaCl2 (DampRid) to NaOCl (found in 6% or higher strength Chlorine Bleach). Reactions:
NaOCl + NaHCO3 = Na2CO3 + HOCl
CaCl2 + Na2CO3 --> CaCO3(s) + 2 NaCl
In the presence of CaCl2, the reaction forms a suspension of CaCO3 which settles on cooling of the mix. Decant to collect the hypochlorous
acid/aqueous sodium chloride mix, which should be kept cool, in the dark, and used within a few hours of preparation.
[Edited on 17-10-2019 by AJKOER] |
I presume that the power of prayer stops the ( obvious to a kid at high school) reaction of the acid with CaCO3
Or it's just another nonsense post from AJKOER.
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Deathunter88
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| Quote: Originally posted by unionised | | Quote: Originally posted by AJKOER |
My prior cited path (see https://www.sciencemadness.org/whisper/viewthread.php?tid=10... ) to hypochlorous acid, HOCl, with a sodium impurity, is by adding NaHCO3 (Baking
soda) and CaCl2 (DampRid) to NaOCl (found in 6% or higher strength Chlorine Bleach). Reactions:
NaOCl + NaHCO3 = Na2CO3 + HOCl
CaCl2 + Na2CO3 --> CaCO3(s) + 2 NaCl
In the presence of CaCl2, the reaction forms a suspension of CaCO3 which settles on cooling of the mix. Decant to collect the hypochlorous
acid/aqueous sodium chloride mix, which should be kept cool, in the dark, and used within a few hours of preparation.
[Edited on 17-10-2019 by AJKOER] |
I presume that the power of prayer stops the ( obvious to a kid at high school) reaction of the acid with CaCO3
Or it's just another nonsense post from AJKOER. |
I can't for the life of me understand where he gets his BS crap from, or why the mods haven't done anything to limit his behavior which misleads
beginners. I for one remember when I wasted a bunch of time and materials when I first joined this forum and thought he would be trustworthy.
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AJKOER
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Really???!
Try reading this:
"In the presence of CaCl2, the reaction forms a suspension of CaCO3 which settles on cooling of the mix. Decant to collect the hypochlorous
acid/aqueous sodium chloride mix, which should be kept cool, in the dark, and used within a few hours of preparation."
So to be clearer for those with low reading comprehension, one separates out the solid CaCO3 in a solution of HOCl, in the current context with a
focus on HCl. I would recommend the cited HOCl path for HCl, based on actual experimental validation and general safety!
I will also mention the action of NaHCO3 on aqueous calcium hypochlorite is the basis of an old patent on HOCl prep, so I cannot claim originality.
Here, for example, is an old source from 1889 at
https://books.google.com/books?id=JX8FAQAAIAAJ&pg=PA81&lpg=PA81&dq=2+NaHCO3+%2B+Ca(OCl)2+%3D+Na2CO3+%2B+2+HOCl+%2B+CaCO3&source=bl&
;ots=0F906Eklrg&sig=ACfU3U2c_33Tlu7AIcT4FsIyhKPJIXN6vQ&hl=en&ppis=_c&sa=X&ved=2ahUKEwjmk6i2rqTlAhWpo1kKHYVhBRYQ6AEwAHoECAgQAQ#v=on
epage&q=2%20NaHCO3%20%2B%20Ca(OCl)2%20%3D%20Na2CO3%20%2B%202%20HOCl%20%2B%20CaCO3&f=false .
My contribution is to mix CaCl2 with NaOCl, for those without access to aqueous Ca(OCl)2. See my related comments in a chlorate preparation thread at
https://www.sciencemadness.org/whisper/viewthread.php?tid=34... (note, in the latter chlorate path, it is my recommendation to keep the suspension
of CaCO3 in hypochlorous acid to address pH considerations and also, I suspect, it improves the photolysis process leading to beneficial radical
formation).
------------------------------------------------------
I should mention that for those with access to TCCA, it dissolves in water to hydrolyzes according to the following equation:
TCCA + 3 H2O → 3 HOCl + C3N3O3H3
Note, per Watts' Dictionary of Chemistry, Volume 2, page 16, the concentration of hypochlorous acid on a single distillation can be nearly doubled:
"A dilute solution of HCl0 may be distilled with partial decomposition, the distillate is richer in HCl0; Gay-Lussao found that, on distilling a
dilute solution to one-half, the distillate contained five-sixths of the total HClO (C. R. 14, 927)"
Caution: My understanding is that elevated concentrations of HOCl (over 20%, per my recollection) are increasingly NOT safe.
[Edited on 18-10-2019 by AJKOER]
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woelen
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@AJKOER: Many of the reactions and pathways you describe may occur in trace amounts, or transiently, but do you really expect them to be of practical
value for synthesis of chemicals? Certainly, some of your reactions are quite interesting from an academic point of view, but in no way I do see any
practical syntheseis.
Take your last post as an example. You state that TCCA forms HOCl when dissolved in water. This indeed is true. But did you ever try to dissolve TCCA
in water? It hardly dissolves. It is sold as slow chlorine treatment for swimming pools. A tablet of pure TCCA can be added to swimming water (they
sell special vessels for that, which allow free flow of water in/out of thewm, without allowing the tablet to go out of the vessel) and such a tablet
lasts for days. So, when we try your experiment, then maybe we can get 0.1% of free HOCl in aqueous solution, and even that is a high estimate. And do
you really think that we can increase that by repeated distillation to 20% or so? No way!
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AJKOER
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Woelen:
I have not been able to find any TCCA over the counter so far, but I do keep searching.
Recent research on prep for HOCl cites it as a source for hypochlorous acid, so I elected to mention it. Apparently, several sources are a bit remiss
in not noting that 'slow' means days!
| Quote: Originally posted by AJKOER | Small amounts of pure HCl (at least free from metal impurities) can be obtained in dilute amounts (which could be concentrated) by mixing chlorine gas
(produced in a variety of ways) with H2O2. The latter, for stability reasons, is free of metallic impurities. The associated chemistry:
.......
HOCl + H2O2 --> HCl + H2O + O2 (g)
---------------------------------------------------
..............
[Edited on 17-10-2019 by AJKOER] |
Per the above extract of my comments, an implied experiment could be to add directly TCCA to more recently available higher dose H2O2 (Chlorox
non-chlorine based laundry bleach alternative under $10). I would still argue that while the action of TCCA in water is apparently slow, this does
NOT necessarily imply that its action in say 10% H2O2, is also equally slow, as the rapid removal of the HOCl by H2O2 implies a shift of the reaction
equilibrium to the right. Interestingly, I did not mention this TCCA reaction path because I was unsure if it was even safe to perform!
So, is the described reaction too slow, or dangerous, or actually some place in between? I nevertheless personally advise employing safety measures
when performing the experiment (just in case). The product, however, is, unfortunately, dilute HCl which while actually may be free of transition
metal impurities as is the intent of this thread, is apparently not particularly cheap to produce.
Here is an extract by you on my suggested reaction above, but with 30% H2O2 and Na-DCCA, relatedly mentioning TCCA, but not as a path to iron-free
HCl, but Singlet Oxygen (a highly reactive gas) with associated chemiluminescence on creation in minutes:
| Quote: Originally posted by woelen | Another nice property of hydrogen peroxide is that it shows red chemiluminescence in some of its reactions.
One particularly good example is simply pouring 30% H2O2 over crushed swimming pool chlorine tablets. Best results are obtained with granules of
Na-DCCA, but TCCA and Ca(ClO)2 also work. The effect must be watched in a dimly lit or even dark room. Use 30% H2O2, the 3% stuff is not sufficiently
strong and then the effect hardly is noticeable. Also, bleach is not sufficiently strong. You really need the strong hydrogen peroxide and the solid
"chlorine" granules |
at https://www.sciencemadness.org/whisper/viewthread.php?tid=99... .
---------------------------------------------
As demonstrated in my chlorate prep thread, the action of NaOCl mixed with CaCl2 in the presence of CO2 upon shaking rapidly, and safely, forms HOCl
in as large amounts as desired, nothing impractical or theoretical. The use of NaHCO3 is more convenient but adds more sodium ions.
I believe in Europe, the 6% aqueous NaOCl bleaching product is not available, so you are not likely to be able to verify my claims, although it is
noted in the literature for over a century (in an ebook that may not be available in Europe). However, my pictures of successful chlorate production
(and subsequent explosive testing of the product) implies strongly a starting mix with hypochlorous acid that is mixed with excess NaOCl among other
agents.
-----------------------------------------
A lower cost and easy path I still contend to iron-free HCl is from HOCl (from the action of CO2 on aqueous Ca(OCl)2, from CaCl2 + 2 NaOCl, minus the
CaCO3, where the hypochlorous acid was distilled to half to nearly double its concentration, as noted above, owing, I suspect, to the early release of
some Cl2O) in the presence of a small amount O2 (generated from adding some H2O2 to the HOCl) and a photocatalyst, treated with diffused sunlight (or
UV) in a sealed chamber. Possible contamination of the HCl with some HClO3. Per one of my cited sources above, the action of light takes but a few
hours!
But why is starting with hypochlorite a good transition metal-free source? Answer: Because the presence of Fe, Cu, and Ni are known to decompose
bleach leading to oxygen buildup in storage containers (see, for example, https://www.powellfab.com/technical_information/files/3670.p... ). Chemically, I explain it by the fact that HOCl engages in a Fenton-type
reaction with transition metals (see comments and references at https://www.sciencemadness.org/whisper/viewthread.php?tid=15... ).
Caution: The half distillation of dilute HOCl (doubling concentration) should not be repeated producing concentrated hypochlorous acid over 20% due to
a possible increased risk of an explosion! My rationale for this hazard follows from the noted disproportionation of HOCl to HClO3 (see https://patents.google.com/patent/EP0490978A1/en ), and the possible further disproportionation of concentrated HClO3 to HClO4 (see https://en.wikipedia.org/wiki/Chloric_acid ). I suspect that the latter perchloric acid is potentially the major problem here, as it apparently
readily forms highly explosive compounds with metals, organics and especially plastic. Employing a distillation setup containing any such material,
other than glass, could be especially problematic.
[Edited on 18-10-2019 by AJKOER]
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unionised
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| Quote: Originally posted by AJKOER |
So to be clearer for those with low reading comprehension, one separates out the solid CaCO3 in a solution of HOCl,
[Edited on 18-10-2019 by AJKOER] |
And, once again for those with poor reading or poor chemistry skills.
HOCl is an acid and, in accordance with high school chemistry, it reacts with carbonates like calcium carbonate to give calcium hypochlorite and
carbon dioxide and water.
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AJKOER
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| Quote: Originally posted by unionised |
And, once again for those with poor reading or poor chemistry skills.
HOCl is an acid and, in accordance with high school chemistry, it reacts with carbonates like calcium carbonate to give calcium hypochlorite and
carbon dioxide and water. |
Finally, the cause of the confusion!
The answer is that hypochlorous acid is effectively best written as HOCl and NOT H+ + ClO- suggesting acidic ability.
In fact, carbonate acid is a stronger acid than hypochlorous acid and CaCO3 is an insoluble precipitate. So, the correct statement is that calcium
hypochlorite and carbon dioxide and water forms HOCl and a precipitate of CaCO3 as my cited source also notes.
Here is yet another source noting that Ca(OCl)2 decomposes with moisture in the presence of CO2 forming HOCl at https://books.google.com/books?id=LA0NdnTHmLMC&pg=PA48&a... .
[Edited on 18-10-2019 by AJKOER]
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woelen
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With the latter reaction, AJKOER is right. HOCl is a VERY weak acid, and even carbonic acid (from CO2, dissolved in water) is capable of releasing
HOCl from water and ClO(-) ions. For the same reason, bleach has its particular smell. What you actually smell is HOCl, whose smell differs quite a
lot from the smell of Cl2.
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unionised
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| Quote: Originally posted by AJKOER | | Quote: Originally posted by unionised |
And, once again for those with poor reading or poor chemistry skills.
HOCl is an acid and, in accordance with high school chemistry, it reacts with carbonates like calcium carbonate to give calcium hypochlorite and
carbon dioxide and water. |
Finally, the cause of the confusion!
...
In fact, carbonate acid is a stronger acid than hypochlorous acid and CaCO3 is an insoluble precipitate. So, the correct statement is that calcium
hypochlorite and carbon dioxide and water forms HOCl and a precipitate of CaCO3 as my cited source also notes.
...
[Edited on 18-10-2019 by AJKOER] |
OK.
Here's a table of acid strengths
https://depts.washington.edu/eooptic/links/acidstrength.html
And from that table, an acid in the third column can protonate the any conjugate base (in the 4th column) below it.
Now here's the interesting bit.
If you look at the table, you see that HCl is a stronger acid than H2SO4.
So, the equilibrium
H2SO4 + NaCl <--> NaHSO4 + HCl
actually lies to the left. The reaction "shouldn't" work.
Yet it's common knowledge that it does.
And the reason for that is that HCl is volatile- the HCl leaves the reaction as a gas.
That reduces its concentration, which enhances the reaction (in the direction as written).
Now consider the case of HOCl and CaCO3
There's a "choice".
Both acids have volatile anhydrides.
So the system could lose Cl2O or it could lose CO2
And the next question is which one is more soluble?
Well, the solubility of Cl2O is about 1430 grams per litre
And the solubility of CO2 is about 1.4 grams per litre.
Which one is more likely to leave the solution?
More importantly, has anyone actually done the experiment?
I have done the analogous experiment with (even weaker) boric acid and sodium carbonate.
I got CO2 and borax
[Edited on 19-10-19 by unionised]
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draculic acid69
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posted on 6-10-2019 at 05:52
Dissolve calcium chloride in water, and add H2SO4 to precipitate calcium sulfate to leave behind an aqueous solution of HCl.
You will need high purity H2SO4 to keep from introducing contaminates, but it can be distilled beforehand with care. It requires high temps, so be
careful.
[Edited on 6-10-2019 by Loptr]
rockyit98
posted on 6-10-2019 at 05:24
you can use AlCl3 or FeCl3 to make HCL gas you need to heat hydrous AlCl3 or FeCl3.
Al(H2O)6Cl3 → Al(OH)3 + 3 HCl + 3 H2O
FeCl3+H2O →FeOCl+2HCl
you can use dirty HCl to dissolve Al to make AlCl3 solution and heat it to make Al2O3 and pure HCl.
since dehydrochlorination speed depend on temperature and is an equilibrium it is much safer than distillation.
if it a hassle uyou can always distill HCl from battery acid and table salt
Both of these methods generate hydrogen chloride gas long before anything else occurs which if bubbled thru distilled water gives u what you want
without going the long way around to reach the same conclusion not to mention boiling a hydrochloric acid solution will cause a heap of acidic steam
to come out the opening of your apparatus.just bubble HCL thru distilled water.simple.no heat.can be done outside with two bottles and a hose.simple
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AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
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Unionised:
Your work is insightful.
However, the main reason why HCO3- (from in particular basic NaHCO3) in the presence of Ca(OCl)2 and H2O causes a precipitate of CaCO3 (and also
results more rapidly in HOCl formation), in my opinion, relates to the lack of stability of in situ formed Ca(HCO3)2, especially with reference to pH
(and an associated absence of CO2 to stabilize the bicarbonate).
See the collected articles at https://www.sciencedirect.com/topics/engineering/calcium-bic... relating to the properties of Ca(HCO3)2 (interestingly, there is a reference to
similar properties for bicarbonates of Zn and Mg, where I have also prepared Mg(HCO3)2 and have observed its decomposition on mild heating or just on
standing for a day or more).
[Edited on 20-10-2019 by AJKOER]
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