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Author: Subject: Hypochlorous acid can exist even in conc. HCl?
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[*] posted on 1-1-2009 at 08:55
Hypochlorous acid can exist even in conc. HCl?


First of all, I want to wish you all a lucky, succesful and interesting new year :) I started the new year with a simple but interesting set of experiments and learned something new (at least for me).

While doing an experiment on manganese chemistry I found a remarkeble thing. Manganese(II) in conc. HCl is not converted to a dark solution with chlorine gas, nor with TCCA. But when a solution of manganese(II) is mxied with a solution of calcium hypochlorite in conc. HCl, then the liquid quickly turns dark brown/green.

Here I describe a complete set of experiments, which demonstrate this peculiarity:

<url>http://woelen.homescience.net/science/chem/exps/exppatt.cgi?compound=manganese%20chloride%20tetrahydrate</url>

Is it really the HOCl in concentrated HCl which makes this reaction possible? I always had the impression that HOCl cannot exist in HCl, certainly not in concentrated HCl.




[Edited on 2-1-09 by woelen]




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[*] posted on 1-1-2009 at 10:04


indeed interesting))
Probably, the dissolution of Ca hypochlorite(especially lime chloride) could result in some ClO3- ion formation. The solution around the solid hypochlorite particles is not too acidic, the concentration of HOCl there is sufficient and the heat is produced while dissolution, so the reaction HOCl + 2OCl- = ClO3- could take place. ClO3- is also known to oxidize HCl, but maybe, not as vigourous as hypochlorite does, so probably, your solution contained some ClO3- , and it had oxidized Mn(2) to Mn(3).
Try to do same experiment, but use sodium hypochlorite (cold) solution and add it to cold HCl with stirring, in order to deminish the ClO3- formation. Or add a wet pinch of calcium hypochlorite to HCl with stirring, maybe there could be another result..
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[*] posted on 1-1-2009 at 10:47


According to article (ACS, ja01326a002), HOCl is "stable" in strongly acidic solution.
Besides, this acid is in equilibrium with Cl2O, which can be extracted with CCl4 etc.

ja01326a002.PNG - 25kB
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[*] posted on 1-1-2009 at 12:03


I tried the experiment with a solution of NaOCl. I added some solution of NaOCl to conc. HCl and added this solution to a solution of MnCl2 in conc. HCl. This leads to formation of a brown/green complex. The mix of solution of NaOCl and HCl also has a stronger color than a plain solution of chlorine in conc. HCl.

As counter experiments I also did the same with KClO3 and NaClO2.

I added KClO3 to conc. HCl. This leads to formation of a deep yellow solution and also some gas bubbles are evolved. When this solution is added to a solution of MnCl2 in conc. HCl, then the solution at once turns almost black.

I dripped a few drops of 30% solution of NaClO2 in conc. HCl. This results in formation of a deep yellow solution and also a lot of gaseous ClO2 is formed. When this solution is added to a solution of MnCl2 in conc. HCl, then also the solution turns almost black (actually very dark brown/green).

With this renewed insight I might change my conclusion. The hypochlorites I have may contain a small amount of chlorate (due to storage) and this may explain the somewhat more intense green color of a solution of this in conc. HCl.

It seems that the common denominator is the presence of ClO2 in the conc. HCl solutions which oxidizes the Mn(2+) to Mn(3+/4+). I'll try to test this by preparing some gaseous ClO2 and dissolving this in conc. HCl, without the salts (which also may contain chlorate and hence leads to chloric acid in conc. HCl). The experiments with ClO2, however, need special care, due to its explosive properties in the gas phase.




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[*] posted on 1-1-2009 at 13:27


A happy and safe new year to all!

Interesting! If you have access to a spectrophotometer with a UV range, you can watch the HOCl vs. -OCl as absorbance at 233 and 292 nm, respectively. A nice plot of pH vs. quantity of Cl2, -OCl and HOCl is given in the paper attached.

The picture below was made using a Beckmann-Coulter DU-800 with 6Q cuvettes. The bleach sample was titrated (burette) vs. pH using standardized HCl (EM).

ClO2 is also a nasty poison, hence its use in industrial sterilization. The absorbance spectrum of NaOCl2 at equivelant concentration is in no way similar. For this pH range, ClO2 is, via my observation, not a player in a standard NaOCl system.

Cheers,

O3

[Edited on 1-1-2009 by Ozone]

Attachment: wang HOCL.pdf (287kB)
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[*] posted on 1-1-2009 at 13:29


And the plot:

vs pH_01_small.jpg - 34kB




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[*] posted on 1-1-2009 at 18:12


The most interesting literature I could find is from the Kirk‑Othmer Encyclopedia of Chemical Technology.
Dichlorine Monoxide, Hypochlorous Acid, and Hypochlorites.

I would do experiment - add some CCl4 to some concentrated HCl(aq) and add - with stirring and cooling - some hypochlorite* . Next I would add extracted Cl2O in CCl4 to Mn(II)/HCl(aq) mixture... Unfortunately I cannot do this - I used up all my Ca(ClO)2 long time ago :(

* for better "proving" existence of HOCl <--->Cl2O, hypochlorite should be added to HCl(aq) and later (after few minutes) CCl4.
Of course, question is if it work at all...... :P

[Edited on 2-1-2009 by kmno4]
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[*] posted on 2-1-2009 at 05:01


In the meantime I have done the final experiment with gaseous ClO2. I took some NaClO2, added this to 30% HCl and the intensely colored yellow gas was "poured" into a test tube with a solution of MnCl2 in conc. HCl. This experiment is quite nice on its own. On the surface of the colorless liquid dark clouds are formed, which quickly mix with the rest of the liquid. The liquid becomes dark green/brown, even with just a small amount of gaseous ClO2.

Another interesting experiment is the pouring of ClO2 over solid MnCl2.4H2O. If this is done, then nothing happens. The pale pink solid remains pale pink. When a ml of water is added, the solid dissolves and a turbid brown liquid is formed. So, the dissolved MnCl2 is oxidized also in neutral solution.

The experiment of kmno4 also sounds interesting. I can try that and will report back on that. Instead of CCl4 I'm inclined to use CH2Cl2 (much less toxic and less of a problem to get rid of), so my first attempt will be with CH2Cl2. I'll come back on this.

@Ozone: Thanks for the PDF and the plots. These confirm that HOCl indeed can exist at low pH, but the situation in conc. HCl of course is more extreme. I unfortunately have no access to any advanced equipment as a home chemist, so I cannot make plots like those you made.




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[*] posted on 2-1-2009 at 08:00


I did kmno4's experiment with extracting Cl2O and/or HOCl.

I proceeded as follows:

1) Prepare a solution of MnCl2 in conc. HCl and divide this over two test tubes.
2) Take 2 ml of conc. HCl and slowly add a spatula full of Ca(OCl)2 while swirling. Wait till the fizzling is over and the liquid is clear again. Wait one more minute. Then add 1 ml of CH2Cl2 and shake well. Let the DCM settle at the bottom.
3) Take 2 ml of conc. HCl and add a few granules of TCCA while swirling. Again, wait till most of the fizzling is over. The liquid becomes turbid (due to finely divided cyanuric acid). Add 1 ml of CH2Cl2 and shake well. Let the DCM settle at the bottom.

At this stage, there are two blobs of DCM, both clear and light green.

4) Pipette away the DCM from both test tubes and transfer the lightgreen DCM from the test tubes with the chlorine-generating solutions to the solutions of MnCl2 in conc. HCl, as prepared in step 1. The transfer of the DCM is done very carefully, assuring that no aqueous solution is transferred as well.


The result of this experiment is striking. Around the blob of DCM, resulting from step (2) a nice brown layer is formed, which on swirling makes the entire liquid look brown. The blob of DCM, resulting from step (3) does not cause any formation of brown color.

A picture is attached to show the difference. The left test tube contains the DCM, uses for extracting from the conc. HCl + calcium hypochlorite and the right test tube contains the DCM, used for extracting from the conc. HCl + TCCA. The blobs are light green/yellow. One can also see that around the blob in the left test tube, the aqueous liquid is somewhat darker than in other areas.



[Edited on 2-1-09 by woelen]

DCM_Mn1.jpg - 45kB




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[*] posted on 2-1-2009 at 08:06


There is another nice observation. On the blob of DCM, used for extracting from the conc. HCl + calcium hypochlorite, slowly a bubble of gas is formed. It starts as a tiny speck, slowly growing to a size of a few mm, until it looses contact and goes to the surface of the liquid.

This bubble can be tapped off the blob of DCM, but after a minute or so there is a new bubble of gas. Only after a few more minutes, no new bubble was formed anymore. The picture shows the bubble of gas on top of the blob of DCM. I can imagine that this bubble of gas is oxygen, due to slowly decomposing HOCl and/or Cl2O.


[Edited on 2-1-09 by woelen]

DCM_Mn2.JPG - 39kB




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[*] posted on 2-1-2009 at 09:59


This may be of marginal interest, but I've had this for a long time without an occasion to post it, so... Decreasing amounts of acid distilled with bleaching powder gave increasing amounts of hypochlorous acid. A warning to those who generate their chlorine from HTH, to do things in a certain way. Also the preparation of hypochlorous acid that is probably in the lab manuals.

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[*] posted on 2-1-2009 at 10:08


At pH less than ~ 3.5 HOCl should yield Cl2, which should be the major component. This is why I am intrigued with your Mn system. Perhaps a oxychloride reservoir species?

Now I remember an experiment where a 1:1 bleach solution was used to treat some wood shavings. The mixture was filtered and the liquor was acidified to pH<2 with H2SO4 which then turned yellow and emitted unholy fumes (I discovered then that the fan on that hood had crapped out and I had to run this stuff down the hall, lachrymating all the way).

Ahem.

This mixture was extracted with DCM and concentrated to 1mL of yellow liquid resembling your sample above. This mixture was injected onto the GC-MS and was un-interesting. HOWEVER...over some time, probably on exposure to light, the yellow component disappeared to yield a clear, colorless solution which upon re-analysis revealed a high yield of sym-tetrachloroethylene.

My best guess is a radical coupling of DCM (because the other route via carbene requires alkali and does not work with DCM)? Maybe you could put the extract in a window (sealed) and see if the color discharges.

I would like to try this with and without Mn, I imagine the the product distribution (or variance thereof) will be quite interesting. I have observed REDOX cycling of Mn from metal to permanganate ion, which was actually stable in the presence of NaOCl. It might also be interesting to use carefully weighed reactants and to capture and measure the gas; some stoichiometric information may be thus acquired.

The depth that can sometimes be found in the mundane never ceases to amaze me. Mn is a schizophrenic metal and bleach is underappreciated.

Cheers,

O3




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[*] posted on 2-1-2009 at 14:14


Ozone, tomorrow I will try your experiment with keeping the extracted material in the light. It is expected to be a reasonable sunny winter day, so I can give it a try.

Based on the article, posted by S.C. Wack I did another experiment, in which I used chlorine gas, made by pouring HCl over Ca(OCl)2. This chlorine gas could well contain some HOCl and/or Cl2O. I made a lot of gas in a 150 ml erlenmeyer and immediately 'poured' the gas into a test tube, containing a solution of MnCl2 in conc. HCl. The test tube was nicely filled with pale green gas-mix, but no dark color could be produced with this gas-mix.

I repeated the experiment with Ca(OCl)2 to which a few tiny specks of solid NaClO2 were added. The gas mix from this was capable of giving a brown/green coloration of the MnCl2-solution.

So, with hypochlorite, there really must be something in the solution and not in the gas-mix.

I fully agree with Ozone about the unexpected depth of the reactions of many mundane reagents. This is another example of such depths, of which I already collected quite a few.




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[*] posted on 2-1-2009 at 15:47


According to a review (ACS, cr60302a004) about Cl2O, this oxide is stable in 74% or 97% aqueous HClO4.
It also states:
"When aqueous solutions of HOCl contain chloride ions, and
are acidic, the slow oxidation of water is replaced by the rapid
oxidation of the chloride ions to chlorine"
But not so rapid (as experiment shows) and some kinetically stable "hydro-oxy-chlorides" can be trapped into inert layer . They are also stable enough to cause oxidations of Mn(II).
From the same paper - stability of Cl2O/CCl4 solutions is order of minutes in daylight.

[Edited on 3-1-2009 by kmno4]
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[*] posted on 2-1-2009 at 16:49


I eagerly await your findings, Woelen!

KMnO4, that agrees with my findings. Although, from pH 3.5-12, I ruled out Cl2O. Of course, in concentrated HCl, it's anyone's guess. I'll have to get a scan and see. Unfortunately, that will have to wait until monday.

I might try this with reduced (bisulfite) KMnO4 (the only Mn salt I have handy) tomorrow. Confirmation is good;).

Cheers,
\
O3



[Edited on 2-1-2009 by Ozone]




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[*] posted on 3-1-2009 at 07:30


Ozone, unfortunately we had a dark and cloudy day, despite the nice weather forecast of yesterday and over here in the Netherlands that means that it is rather dark at this time of year, even in the daytime (on such dark days we need to have artificial light on all day). Tomorrow's weather forecast is even worse, a dark grey and rainy day.
I'll try working with a UV-lamp instead (I have a blacklight lamp) and see if the DCM-extract still shows activity after several minutes of irradiation with the blacklight.

It would be very nice if you could repeat my experiments. The more people try this, the more insight we will gain. You may notice other interesting things, which I overlooked.
If you try to repeat this with reduced KMnO4, then don't use bisulfite. Bisulfite is oxidized to a mix of sulfate, SO4(2-), and dithionate, S2O6(2-). The latter may spoil the outcome of the experiment, because it can be oxidized further. Instead, I would use H2O2 added to a solution of KMnO4 in dilute sulphuric acid. Add the H2O2 slowly and after each drop stir well. The endpoint can be detected very well and the reaction is clean. No incompletely oxidized byproducts. When the reaction is at its end, boil down the solution to make it somewhat more concentrated and at the same time destroying the tiny remains of H2O2 of the last drop.

[Edited on 3-1-09 by woelen]




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[*] posted on 3-1-2009 at 08:00


The weather is also vile here. I'll see if I can get some more DCM and I'll repeat what I can. IIRC, there was a lengthy list of experiments that you did with a link from the first post. Today, I can only find experiments 1 and 2? It would be a good idea to review the previous experiments in detail before attempting to repeat them.

Good call with the bisulfite;).

Cheers,

[Edited on 3-1-2009 by Ozone]




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[*] posted on 3-1-2009 at 08:12


Sorry for the inconvenience of the link. You should click experiment 2 and then you see what you used to see. I need to make a technical change to my website. Right now, all experiments have an index (the particular experiment you mean has index 517), but this index only is valid as long as the server is running. Each time when the server is restarted another index may be assigned to experiments, making all existing links to such experiments incorrect.

The technical solution to this is to have an index not pointing directly to an experiment description, but to another index, where the second index points to the experiments. This technique allows for constant URL's for experiments. It works a little bit like a Java Virtual Machine where each object has a reference to a reference, allowing moving around of objects through memory (e.g. for garbage collecting), while the most outer reference remains the same. I need some time though to make such a technical change, it is not trivial and there are more interesting things to do ;). If you have ever programmed in C and you have done some more advanced memory management tricks then you most likely exactly understand what I mean, and otherwise simply forget my blurp on this subject :P




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[*] posted on 3-1-2009 at 08:59


Thanks, that did the trick. I saved the page this time. OK so the experimental sequence should look like this (let me know if I missed anything):

To HCl (at various strengths ranging from 36-0 %) is added Ca(OCl)2, TCCA or Cl2 (KMnO4/HCl). The solutions are split into two tubes. To one is added Mn2+. The other is available for reference. Observations are recorded (photos, etc.) and if possible the solutions should be checked for absorbance at 233 (HOCl) and 292 (-OCl).

This can be calibrated vs Ca(OCl)2 (or NaOCl) at various pH with a quantitative correction applied via reduction of the -OCl with H2O2 standardized via thiosulfate/I2.

The differential reactivity of Cl2 bubbled mixture and the -OCl/TCCA mixtures rules out oxidation, etc. resulting from chlorine (which is the expected outcome of -OCl at very low pH). I would like to see that HOCl peak at 233nm; this should help to close the case.

Whew,

O3




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[*] posted on 3-1-2009 at 09:20


Yes, I think that what you describe should be conclusive. The only strange thing is that HCL, treated with TCCA does not lead to the oxidation of Mn(2+), while HCl, treated with Ca(OCl)2 does oxidize Mn(2+).

There may be a practical problem with the TCCA. The liquid with TCCA becomes turbid, the cyanuric acid, formed in the reaction, hardly is soluble, so I'm afraid that absorbance measurements are not possible with TCCA.

Also keep in mind, that with TCCA no oxidation occurs, not even when solid pieces are thrown in a solution of MnCl2 in conc. HCl.




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[*] posted on 3-1-2009 at 09:51


I suppose I could filter everything through a 0.5 um TFE membrane directly into the cuvette (or I could centrifuge them). I am using 1.5 mL cells, so I only need to pass a small amount.

The inability of the TCCA to oxidize the Mn(II) backs up the Cl2 experiment, it should not lead to HOCl, where, hypothetically, the -OCl might. My guess is that the TCCA goes straight to Cl2 and cyanuric acid in conc. HCl which results in a test outcome similar to that of the Cl2 bubbled mixtures.

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O3




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[*] posted on 3-1-2009 at 11:19


Ah yes, if you can do absorption measurements, then surely you also have the suitable filtering equipment :P

I have done the experiment with DCM again. I added a pinch of solid Ca(OCl)2 to conc. HCl swirled a little and waited till all bubbling was over. Then I added the DCM and swirled again for quite some time to extract as much as possible.

Next, I transferred the DCM to a separate test tube. Half of the amount I added to a solution of MnCl2 in conc. HCl. The other half I treated with UV-A light from a blacklight tube for 2 minutes and then I added that to a similar solution of MnCl2 in conc. HCl. The DCM, treated with UV-A light lost its ability to oxidize the Mn(2+) completely, while the other half gave the same results as yesterday.

I also took Ca(OCl)2 from a totally different source, just to be sure that not a specific batch of this substance gives such a peculiar reaction, but that is not the case. I now have tried with three sources of hypochlorite (HTH 65% active chlorine, Choc POOL, 70% active chlorine, and ordinary househoold bleach). All show the same peculiar reaction.

Btw, the Choc bottle of calcium hypochlorite tells that the main ingredient is Ca(OCl)2.2H2O instead of Ca(OCl)2.

[Edited on 3-1-09 by woelen]




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[*] posted on 3-1-2009 at 13:56


Fantastic! This agrees with my observation. I was reluctant to go forward as my Ca(OCl)2 contains ~45% of "other" stuff. I was going to wait until I can get some lab-grade material monday. I'll do the replicates and get MS on the DCM extracts.

I am sure now that the reaction is peculiar to hypochlorite/HOCl at low pH and that the oxidative "intermediate" is photochemically labile. From my experiments, H2SO4 will also do (although I will test this again with NaOCl this evening when my friend brings over some DCM) which rules out a specificity for HCl. I won't be able to rule for or against Cl2O (as KMnO4 suggests) until I see some UV-Vis.

I'll try to dig up my earlier GC-MS for reference.

Yea!

O3




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[*] posted on 6-1-2009 at 12:05


The funny thing is that the aqueous solution also looses its activity when it is exposed to UV-A radiation. The following very simple experiment demonstrates this.

Take 3 ml of 30% HCl
Add a small amount of solid Ca(OCl)2 and swirl until the bubbling stops.
Divide the solution over 2 test tubes.
Set 1 test tube aside in a dark place and expose the other to the light of a blacklight for 2 minutes or so.
Add a pinch of solid MnCl2 or MnSO4 to both test tubes and dissolve the solid.

In the test tube, exposed to the blacklight, no dark brown color appears. The solution remains pale green. In the other test tube, the solution becomes dark when the MnCl2 dissolves.

[Edited on 7-1-09 by woelen]




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[*] posted on 6-1-2009 at 13:58


I assume you mean Ca(ClO)2 ?
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