Sciencemadness Discussion Board - Free hypochlorous acid in aqueous solution

Free hypochlorous acid in aqueous solution

woelen - 27-9-2009 at 05:14

I found something quite interesting and that is that free HOCl can be made in aqueous solution at fairly high concentration.

When you add solid Ca(OCl)2 to hydrochloric acid, then a fairly vigorous reaction starts, in which chlorine gas is evolved.

When solid Ca(OCl)2 (swimming pool chlorine, be sure to use the calciumhypochlorite stuff with 65 ... 70% active chlorine, not the TCCA-based stuff) is added to dilute HNO3 or HClO4 (15 ... 20% concentration) then some gas is evolved, mostly colorless. The amount of gas produced is much less than with HCl. I think that the main constituent is CO2 from the CaCO3 impurity of the swimming pool chlorine.
After the solid has dissolved, an almost colorless solution is obtained. When table salt is added to this colorless solution, then a violent reaction starts and chlorine is ejected from the test tube.

I did some tests with the liquid.
- brief boiling does not result in decomposition. After this treatment it still reacts violently with table salt.
- keeping it near a UV-backlight source for several minutes, or keeping it in bright sunlight does not make it less active, it still reacts violently with table salt after this treatment.

I think that this solution can be quite interesting for some syntheses. It is possible to use HOCl, even at fairly high concentration (I think I had at least several percents). It works fine with HNO3 and HClO4. With H2SO4 it also works, but you get a slurry of solid CaSO4 and that is quite annoying, the liqud gets a paste-like appearance and is hard to handle.

Is my assumption correct that the active species in the liquid is HOCl? I was surprised to see this work, I expected quick decomposition, and it also surprises me that the boiling does not destroy the HOCl.

HOCl has a very peculiar smell, certainly not like chlorine, but its smell is pungent.

chloric1 - 27-9-2009 at 06:37

What the hell?

woelen, if today was April 1st I would call your bluff. But, then again all your wonderings almost always produce valvid results. Is the smell reminisent of both chlorine and nitric acid? That is how chlorine dioxide smells. Is there any brownish fumes of chlorine monoxide? What color is your hypochlorous acid? Can we see a photo? I think the next move would be to use perchloric acid to isolate so you could do a potassium iodide/starch titration to find the concentration without the mineral acid(nitric acid) also attacking the iodide.

I was going to suggest you add ethylene glycol but that might be suicidal. Maybe adding freshly precipitated chromium (III) hydroxide or litharge to the mix and heating if it does not erupt violently of coarse.

[Edited on 9/27/2009 by chloric1]

S.C. Wack - 27-9-2009 at 09:29

It can be distilled. Boric acid is preferred.

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[Edited on 27-9-2009 by S.C. Wack]

entropy51 - 27-9-2009 at 09:34

Quote: Originally posted by S.C. Wack

It can be distilled. Boric acid is preferred.

How do you do that?? I mean come up with these obscure references just like that?

S.C. Wack - 27-9-2009 at 10:10

That's an off-topic UTFPM question, but since it's here and I've never mentioned it before

I decided to start collecting references for a book of my own, covering syntheses for chemicals that interest me, in 1984. It is as yet quite incomplete and I have decided on a web site instead, which I have started at google sites and will mention once it is more complete. [Very off-topic sample entry (acetaldehyde) showing the sort of thing I'm doing; sorry for jacking your thread, woelen. MEMEMEMEMEMEME] Might redo it at a .org domain, which I have a good name for (to me...actually an old name from the 1800's resurrected) but haven't bought yet. BTW I continue so that the many hours spent early on sitting on the floor of the library in the hope of selling a dozen copies weren't a total loss, though obviously the unforseen internet, google books, etc. has killed the ability to sell a book of out-of-copyright-syntheses.

Boric acid in particular was a brain-stored factoid, I only needed to search the inorganic folder for a file title. The pdf was downloaded some time ago, replacing handwritten notes; for JCS of that era is only on microcard at the nearest library.

It's interesting that it took 967 posts here, maybe 3000 total including a couple other places since 2004, for someone to ask WTF is up with that. People just think I'm weird.

[Edited on 27-9-2009 by S.C. Wack]

entropy51 - 27-9-2009 at 10:31

Quote: Originally posted by S.C. Wack

It's interesting that it took 967 posts here, maybe 3000 total including a couple other places since 2004, for someone to ask WTF is up with that. People just think I'm weird.

It's a thin line between genius and weird, but I recognize talent when I see it. Thanks for the off topic explanation and please let us know when we can see the website.

chloric1 - 27-9-2009 at 10:31

Thank you for the reference

One way to diffrenciate between hypochlorous acid and chloric acid is that alkaline aqueous solutions of chlorates are not oxidizing while hypochlorites are.

Ozone - 27-9-2009 at 10:40

I knew I had gotten that paper previously (here). Much more info on HOCl and associated species can be found here:

http://www.sciencemadness.org/talk/viewthread.php?tid=11651

Since then, I have ruled out chloride, chlorate, and perchlorate.

Cheers,

O3

woelen - 27-9-2009 at 11:58

Funny to read that old thread again. Yes, this is interesting chemistry.

I also tried the Mn-reaction of that old thread with my newly prepared HOCl in dilute HClO4. Surprisingly, this does NOT lead to the formation of the dark brown species! I need to do much more investigations now on this subject, it is interesting and intriguing, but I need more time to test a large number of possible combinations.

densest - 27-9-2009 at 14:28

Quote: Originally posted by S.C. Wack

I decided to start collecting references for a book of my own, covering syntheses for chemicals that interest me, in 1984. It is as yet quite incomplete and I have decided on a web site instead, which I have started at google sites and will mention once it is more complete. Might redo it at a .org domain, which I have a good name for (to me...actually an old name from the 1800's resurrected) but haven't bought yet. BTW I continue so that the many hours spent early on sitting on the floor of the library in the hope of selling a dozen copies weren't a total loss, though obviously the unforseen internet, google books, etc. has killed the ability to sell a book of out-of-copyright-syntheses.

Don't be quite so discouraged immediately; the effort of consolidating the data could be worth a lot. The latest thing in web based business (so I read today on an investing website) is selling organized data/databases on obscure or hard-to-access topics. You could give a discount to us impoverished amateurs

The Merck Index contains zero (AFAIK) original content, but is very valuable to many people!

Quote: Originally posted by S.C. Wack

It's interesting that it took 967 posts here, maybe 3000 total including a couple other places since 2004, for someone to ask WTF is up with that. People just think I'm weird.

[Edited on 27-9-2009 by S.C. Wack]

Or have a really, really good memory!

[Edited on 27-9-2009 by densest]

kmno4 - 27-9-2009 at 14:30

If it is HOCl sol. then you should be able to extract some part of Cl2O from it with CCl4 and make experiment with Mn(II).
BTW:
Solutions of HOCl are "stable" up to 25%, have smell different than Cl2O and Cl2, concentrated ones are yellow,
diluted are colorless.
You can read more about it in Brauer, IS, Kariakyn & Angielov.... etc.

Jor - 27-9-2009 at 14:53

That would be very nice. In that case you just need CCl4 to make Cl2O solution, instead of oxidising chlorine over a bed of yellow HgO (you need ALOT HgO, very nasty stuff).
But I wonder if it converts to the anhydride so easily, especially in aqeous solution.

entropy51 - 27-9-2009 at 15:39

Quote: Originally posted by Jor

HgO, very nasty stuff

Compared to what? Formaldehyde? Potassium dichromate? Chlorine? Lead acetate? Almost anything else used in half the threads here?

Apparently lots of people here distill HNO3, and heaven forbid, H2SO4. That's nasty for sure.

Sure, HgO is toxic as all hell if you eat it, breathe the dust, or rub it on your skin. Otherwise, no. Methylmercury, now that's very nasty! But not HgO. It won't penetrate gloves as methyl mercury will. HgO is not volatile, as is formaldehyde and many other things people use routinely.

A standard experiment in high school was to heat a gram or so in a test tube to show that oxygen was released and tiny droplets of Hg coated the test tube. Now that was very nasty. But we did it, and forty years later I'm no worse off than other people my age.

The problem with Hg compounds is how the heck can you dispose of the waste?

I don't understand why rather ordinary chemicals are stigmatized (Mercury! run for your lives) and others that just as toxic are handled without a second thought.

Basically all chemicals are very nasty and that's the attitude that we should all have.

[Edited on 27-9-2009 by entropy51]

Jor - 27-9-2009 at 17:16

I was referring to the disposal ofcourse. I have no problem with the toxicity. I have worked with mercury compounds at home, doing chemistry test-tube scale (starting from my HgO), and also one larger scale synthesis:
http://amateurchemie.nl/viewtopic.php?f=20&t=155

The scary thing is, when you spill some mercury compounds. Hard to clean up, a mess. I once had beaker containing 10mL conc. CuSO4 fall over. I would not want to face a mercuric nitrate solution falling over, especially if it also contains precipate (harder to clean).
When you spill hydrazine hydrate, bromine, formaldehyde, benzene, carbon tetrachloride, it's no problem. Just let it evaporate (unless you spill it outside a hood ofcourse!).

I remember you many grams of HgO for 1g of Cl2O. That's a PAIN to dispose or recycle so much mercury.

Sorry for going offtopic.

entropy51 - 27-9-2009 at 17:34

No, my fault for going off topic. Just wanted to make the point that pretty much all chemicals are nasty. Hg is not the safest element to work with, but no worse than many other chemicals discussed on the forum. I guess I'd rather work with Hg salts than with Cl2O.

As far as spills go, one can work in a shallow metal or plastic tray lined with absorbent paper. That's standard for radioisotope work and makes spills very manageable.

Formatik - 30-9-2009 at 14:18

In Gmelin it's stated a conc. HClO solution even at 100 deg. its decomposition is incomplete so that by distillation, a concentrated solution is obtainable. Dilute solutions can be distilled without noticeable decomposition. Pretty conc. solns. can be stored in the dark with nearly no decomposition for several months.

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Taoiseach - 1-10-2009 at 03:31

I once made a pretty conc. solution of HOCl in an attempt to make chlorates from just sodium carbonate and chlorine. Basically Cl2 was bubbled into a suspension of Na2CO3 at low temperature and then brought to a boil to decompose the sodium hypochlorite, forming chloride and chlorate. The chlorate would then be precipated as the K salt. Well that was the theory

I felt pretty smart having discovered a way to make chlorates from chlorine without the need of using expensive KOH

But in practice the synthesis turned into a complete mess. The reason is: there was no sodium hypochlorite to decompose, but rather a deeply yellow solution of badly corrosive HClO (a drop on my hand made me run for the sink instantaneously) that smelled decidedly different from chlorine. It was incredibly stable and I had to boil this shit for several hours until finally the color and smell disappeared. A shitty yield of potassium chlorate was recovered in the form of hygroscopic crystals. Only by strongly heating the moisture went off slowly, forming dense fumes of HCl. So there was still free hypochlorous acid in this solution, after at least 2 hours of boiling! The experiment was a complete failure, except it gave some interesting lessons.

Now here comes my explanation:
When chlorine is reacted with water, an equilibrium is formed that lies pretty far on the left side:

Cl2 + H2O ---> HClO + HCl

However, by removing HCl it is shifted to the right and HClO is formed. Neutralizing the HCl can be accomplished by anything that does *not* react with HClO. Hydroxides dont work as HClO acts upon them forming hypochlorites which then auto-oxidize to form chloride and chlorate. Thats the famous lecture experiment of chlorate being formed from chlorine in KOH solution. Metal oxides which don't hydrolize can be used. HgO has already been mentioned in this thread. However sodium carbonate will also work nicely, and this procedure is described in literature as well. I found it in an old chemistry book quite some time after the failed experiment. It seems as if HClO is an extremely weak acid, much weaker than H2CO3 actually.

Here's the supposed reaction scheme:

2Cl2 + 2H2O + Na2CO3 ---> 2HClO + 2NaCl + H2O + CO2

It'd be interesting to destill such a solution and form exotic hypochlorites from it. There's not much info on hypochlorites other than Ca, Na and K (the latter apparently only being known in solution). Tetrammine copper hypochlorite pops to mind... wonder how stable that beast'd be

Also I wonder if chlorites could somehow be made from it...

[Edited on 1-10-2009 by Taoiseach]

woelen - 1-10-2009 at 05:29

I expect that the supposed reaction equation only approximately can be correct. There also is HCO3(-) and that is an even weaker acid than H2CO3 (and most likely also than HOCl). So, I expect that this method can work, but only when not too much Na2CO3 is used. Initially I think that NaOCl is formed, but when more chlorine is added, then it might revert to HOCl.

In one of my experiments with the HOCl solutions (in dilute HClO4) I noticed an interesting effect when N2H4.2HCl is added. When this is done, then of course you get chlorine, due to the chloride in this chemical, but also ClO2 is formed. The liquid turns intensely yellow and the color of the gas mix also becomes much stronger than the color of chlorine gas. I know that ClO2 has an intense color, and I think that the deep color is due to the presence of ClO2. I'm almost 100% sure that this is not from the perchlorate, the latter is very inert in aqueous solution (I tested that by boiling 60% HClO4 with KI and it does not oxidize the iodide ions).

Tetrammine copper complex cannot exist in the presence of hypochlorite. Just add some to bleach. Ammonia reacts with hypochlorite, giving mostly nitrogen, but also some nasty chloramine is formed (very pungent and unpleasant smell, probably carcinogenic as well).

You can make organic hypochlorites, IIRC there is some thread around here about that subject, but I did not try that myself. Organic hypochlorites are explosive and I do not want an accident with that stuff. Inorganic hypochlorites might be another matter, I can imagine that it is possible to make Sr- or Ba-hypochlorites, maybe also Cs-hypochlorite. These hypochlorites could be very energetic compounds. Ca-hypochlorite already is. When crushed Ca-hypochlorite is mixed with powdered red P, then the mix slowly heats up and suddenly it sets off with an orange flame. There is no "might set off", there is a "will set off".

chloric1 - 1-10-2009 at 06:12

The Merck Index says that zinc sulfate is used with calcium hypochlorite to bleach paper. I would presume an unstable zinc hypochlorite was formed. I was cleaning my trash can the other day and with a couple gallons of water and some very old calcium hypochlorite. I added some zinc sulfate to get a flocculate slurry and all organic odors where INSTANTLY dispensed. More so than with Ca hypochlorite alone. As a plus, the can had a very faint sweet odor

Although the zinc sulfate lowered the pH of the solution, no chlorine gas wad evolved nor was there any pungent fumes. Might need to reevaluate this with fresh Ca hypochlorite though.