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Author: Subject: A new remarkable riddle: what is this dark compound?
woelen
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[*] posted on 11-5-2013 at 12:03
A new remarkable riddle: what is this dark compound?


I have done quite a few experiments with chlorine dioxide and some of you may have seen some web pages about this gas.

In this experiment, I studied the compound in aqueous solution. While doing so I found that the color of solutions of ClO2 (or some adduct) seems to depend strongly on pH. Usually, making ClO2 is done simply by adding excess acid to a solution of NaClO2 and then a deep yellow liquid is obtained which gives a lot of ClO2. I now did an experiment in which just a small quantity of dilute acid is added and this gives quite remarkable results! Look at the picture below. The only things used in this experiment are sodium chlorite and some acid (e.g. dilute HCl). It is remarkable that these reagents can generate such a dark compound.



The webpage describing the experiment is available at the following link:

http://woelen.homescience.net/science/chem/exps/ClO2_dark/in...

If anyone has specific knowledge about this dark compound then I would like to hear about that.

[Edited on 11-5-13 by woelen]




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[*] posted on 11-5-2013 at 12:25


Might be that ClO2 is less soluble in highly acidic solutions. Another possibility is that it interacts somehow with the unreacted NaClO2 to make this color. What can I tell for sure is that the dark color is due to the dissolved ClO2.



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[*] posted on 11-5-2013 at 12:41


It's unusual. Even Cl2O7 is reported to be colourless.



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[*] posted on 11-5-2013 at 15:55


Here is a thought, repeat the experiment by replacing the reactants with known pure products. For example, make a small amount of pure HCl in place of your current HCl which may have some Fe impurity (or, perhaps the distilled water used to wash). Make AgClO2 as your chlorite source (or, add aqueous NaOH to it and filter out Ag2O for NaClO2) by shaking a soluble aqueous Silver salt in ClO2 and quickly isolating the AgClO2:

ClO2 + H2O <---> HClO2 + HClO3

HClO2 + AgNO3 --> AgClO2 (s) + HNO3

HClO3 + AgNO3 = AgClO3 + HNO3

Any HCl formed on standing as a result of decomposition/disproportionation will form a chloride impurity, so quickly collect and wash the AgClO2,

Use a different, low level light source in the event that the ClO2 is undergoing some photolysis.

Lastly, use a new clean test tube.


[Edited on 12-5-2013 by AJKOER]
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[*] posted on 11-5-2013 at 17:13


This is totally iron. Ferratesdo not survive in acidic media, so you are simply oxydising iron so bad it turns into Fe2O3 and then decomposes. Olation is present as it is a very interactive cloud of material, as any solubilized Fe3+ becomes Fe2O3 then one of the irons are radicalized again via the following reaction:

Fe2O3 <----> Fe(OH)3
the presence of ClO2 shifts that to the left:
Fe(OH)3 + ClO2 ---> FeO4(2-) + HClO
Neverthless, the reaction:
FeO4(2-) in strongly acidic conditions tends to decompose much much more rapidily, to Fe2O3, as can be checked out on your own site woelen;
I could well post a summary on the olation of iron but I'm too lazy right now. Thats not red as indicated by the strongly dynamic nature of this reaction.

Thats just a thought, reaction rates and coefficients coudl be much different!
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[*] posted on 12-5-2013 at 02:52


Iron? This experiment has nothing to do with iron and I'm absolutely sure that none of the reagents I used contained any iron. I used reagent grade HCl, but I also used HNO3 and H2SO4, all of high purity. NaClO2 I used also is free of any iron, it may contain some NaCl though.

@AJKOER: Why do you think I should repeat the experiment? I have done this many times and I use clean glassware. I certainly do not use dirty test tubes when I find some riddle in my experiments. Especially in such cases I do everything which is possible to rule out effects due to impurities in my reagents. The only impurity I might (and almost certainly) have is chloride ion in my NaClO2, as this usually contains quite some NaCl as impurity. For the rest, I am very confident that I did not introduce any transition metal impurity, nor any colorful organics.




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[*] posted on 12-5-2013 at 05:41


Iron can be safely excluded: it wouldn't even look like that in any event.

I'm afraid that w/o identifying the VIS bands (and perhaps NIR) that are causing the colour it will be very difficult to identify what compound might be responsible for them. Something nags at me that there may be Cl-Cl bonds present: at least diatomic chlorine for instance does have a weak colour.

Can interhalogen based compounds be completely excluded? For instance may your NaClO2 contain some non-Cl halogen impurity?

It's all quite reminiscent of charge transfer bands, yet I fail to see where they would come into it.

And of course we could be looking at very small amounts of a highly coloured substance.

It's very intriguing indeed...



[Edited on 12-5-2013 by blogfast25]




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[*] posted on 12-5-2013 at 06:06


Some sort of chloryl complex, maybe? Sodium, chlorine, oxygen, chlorine and hydrogen forming in some way a dark substance. Just thought about the red-coloured ClO2+ cation. However, I don't think that chloryl chloride would be that easy to make especially without chlorates.

EDIT: Whoops, forgot that it reacts with water. So not a chloryl compound then.

[Edited on 12-5-2013 by Eddygp]




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[*] posted on 12-5-2013 at 09:02


Quote: Originally posted by woelen  
....
@AJKOER: Why do you think I should repeat the experiment? I have done this many times and I use clean glassware. I certainly do not use dirty test tubes when I find some riddle in my experiments. Especially in such cases I do everything which is possible to rule out effects due to impurities in my reagents. The only impurity I might (and almost certainly) have is chloride ion in my NaClO2, as this usually contains quite some NaCl as impurity. For the rest, I am very confident that I did not introduce any transition metal impurity, nor any colorful organics.


Does a change in lighting conditions have any impact?

Next, I did not mean to imply that you have introduced any such impurities. What I am suggesting as a next step is that by using alternate preparations of the starting reactants, under your control, are the experimental results reproducible?

If yes, then the understanding of what is produced is most likely due to the reactants solely and experimental conditions, and one can discard questions of impurities, no matter how unlikely (especially in your case) to begin with.

I am sure you known what to do next, like change the NaClO2 for another chlorite and note results...

This systematic process is, unfortunately, only suggestive as to the source of the mystery.
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[*] posted on 12-5-2013 at 09:59


@AJKOER: I tried the experiment with another source of NaClO2 (I have samples from two completely different sources). The result is exactly the same.

@blogfast25: I have severe doubts that there are other halogens involved in my experiments. One of my NaClO2 samples is intended for human consumption (as MMS, containing according to spec at least 80% NaClO2, the balance being plain NaCl with traces of Na2CO3) and having bromine or iodine in this does not seem something which is acceptable for a consumer product. The other sample is reagent grade NaClO2. I also expect this to be free of bromine or iodine.

Even if there were some other halogen, could this lead to such very dark compounds in aqueous solution? This dark material is a total riddle to me, I was very surprised to see such a dark compound from chlorine-based inorganic chemicals only.

[Edited on 12-5-13 by woelen]




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[*] posted on 12-5-2013 at 10:15


Woelen, I don't know much about how compound H[Cl3] looks like, but conditions may be favourable for it's formation, check out it's color..
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[*] posted on 12-5-2013 at 10:32


Cl2O3 is supposed to be dark brown. Maybe it's possible that HCl can reduce ClO2 somehow to yield this?



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[*] posted on 12-5-2013 at 12:26


Quote: Originally posted by bbartlog  
Cl2O3 is supposed to be dark brown. Maybe it's possible that HCl can reduce ClO2 somehow to yield this?


Aha. At least one coloured Cl, O based compound. So bbart is proposing some reaction, in which in acid conditions Cl (+IV) is reduced to Cl (+III) and Cl (-I) oxidised to Cl (0) (or Cl (+I)). So far the best, yet unproven, hypothesis, methinks... Especially as interhalogens can be ruled out.

This hypothesis could in part be tested by increasing the amount of Cl (-I) present (as NaCl), because that should increase the reaction speed (formation of the coloured species). Alternatively, a purer sample of NaClO2 (with less free chloride) should then produce the opposite effect.

To observe any influence of free chloride on discolouration speed it may be necessary to run the experiments at lower concentration.



[Edited on 12-5-2013 by blogfast25]




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[*] posted on 12-5-2013 at 15:04


As I previously posted (http://www.sciencemadness.org/talk/viewthread.php?tid=18911#... ) on the color of HOCl:

"Dilute HOCl solutions are colorless; at higher concentrations the color ranges from yellow to yellow-orange due to small equilibrium amounts of Cl2O."

This could be part of the answer.
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[*] posted on 12-5-2013 at 15:57


In Brauer's second prep for ClO2 [p302 of his book in the forum library], which involves passing chlorine gas through a 10% solution of sodium chlorite, he writes '...when the NaClO2 solution in the first wash bottle changes suddenly from brown to a weak yellowish-green, it is exhausted...'. So it seems that gassing the NaClO2 with chlorine also produces this brown color. Also it seems to me on reconsideration that no reduction needs to be occuring in order for Cl2O3 to be present since it would simply be the anhydride of HClO2. On the other hand, Cl2O3 hardly sounds stable and I have no idea whether it could exist in significant quantities in aqueous solution at room temperature.



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[*] posted on 12-5-2013 at 17:47


Quote: Originally posted by woelen  
@AJKOER: I tried the experiment with another source of NaClO2 (I have samples from two completely different sources). The result is exactly the same.

@blogfast25: I have severe doubts that there are other halogens involved in my experiments. One of my NaClO2 samples is intended for human consumption (as MMS, containing according to spec at least 80% NaClO2, the balance being plain NaCl with traces of Na2CO3) and having bromine or iodine in this does not seem something which is acceptable for a consumer product. The other sample is reagent grade NaClO2. I also expect this to be free of bromine or iodine.

Even if there were some other halogen, could this lead to such very dark compounds in aqueous solution? This dark material is a total riddle to me, I was very surprised to see such a dark compound from chlorine-based inorganic chemicals only.

[Edited on 12-5-13 by woelen]


Hm... take note even ridiculous trace ammounts of certain elements can fade colors to a lattice.
I still take the position that this strong iron transition may even provide a very useful dye laser media/ tunner.
You cannot fake out the presence of impurities stating sources that do not stablish well suitable standards in the description.

[Edited on 13-5-2013 by platedish29]
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[*] posted on 12-5-2013 at 23:04


Cl2O3 is an interesting suggestion. I can imagine formation of that, it simply is the anhydride of HClO2. I did not know about Cl2O3, but indeed, it is mentioned on wikipedia and there is some (albeit very sparse) information about this compound on other sites. It is a dark brown compound according to what is written.

The only thing which makes me doubt somewhat is the low stability of this compound. I certainly would not expect such a compound to be stable at all in aqueous solution. But on the other hand, I have seen many other surprises with halogens (e.g. easy formation of choclate brown ONBr from aqueous NaNO2 and aqueous HBr). I'll try to find more info on Cl2O3 and I'll also try bubbling Cl2 through a solution of NaClO2 as bbartlog mentions.

@platedish29: I rule out the presence of iron for this very dark compound. Why are you so sure that it is the presence of iron causing this strong color? I know that small amounts of impurities can cause strong colors, but I see no reason why my chemicals would contain iron. I did experiments with different acids from different sources, NaClO2 from different sources and all experiments have the same result. The strength of the effect also is the same and if the effect were due to trace amounts of an impurity, then with different sources of chemicals I would expect strong differences of strength of effect. No, this strange effect really is due to the chlorite (and possibly in connection with the presence of chloride).




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[*] posted on 13-5-2013 at 04:54


Woelen:

Did you know Wiki mentions weakly acidified NaClO2 solutions (with citric acid) as commercial solutions? No mention of the dark colour though.

I'm even a bit miffed at the generation of ClO2:

ok, so we have HClO2 < === > H+ + ClO2 + e-, that's an oxidation, Cl (+III) to Cl (+IV). So what gets reduced here? Holleman states: chlorine + chlorite === > chloride + ClO2 and that makes sense, as it's the chorine (Cl2) that gets the electron. But in the absence of an oxidiser, how can ClO2 form? Only straight decomposition can then explain it: like 2 HClO2 < === > ClO2 + 1/2 Cl2 + H2O + 1/2 O2, or some disproportionation involving HClO or Cl2O as reaction product.

It seems a plausible explanation that Cl2O3 forms at modest acid pH via: HClO2 < === > 1/2 H2O + 1/2 Cl2O3 but why does it disappear when more acid is added?



[Edited on 13-5-2013 by blogfast25]




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[*] posted on 13-5-2013 at 06:37


HClO2 is unstable and tends to disproportionate to chloride and ClO2. One of the chlorine atoms at oxidiation state goes from oxidation state +3 to -1 and acts as oxidizer and the others go from oxidation state +3 to +4 and act as reductor. This reaction is not immediate, but it easily occurs. This causes formation of ClO2.

In the presence of chloride and a large excess amount of acid, there is fast formation of nearly pure ClO2. E.g. adding 30% HCl to a solution of NaClO2 yields nearly pure ClO2.




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[*] posted on 13-5-2013 at 06:53


Try adding a few drops of conc. sulfuric acid to powdered KClO3 (caution, may explode). This also turns a dark color. I can't try this out myself right now, but I think I remember this gives a similar color as your sodium chlorite + HCl riddle.
I also think that the dark color is due to an oxide of chlorine, perhaps a very unstable one that cannot be isolated.




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[*] posted on 13-5-2013 at 09:35


Quote: Originally posted by woelen  
HClO2 is unstable and tends to disproportionate to chloride and ClO2.


Basically:

5 HClO2 === > 4 ClO2 + HCl + 2 H2O

That makes sense.




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[*] posted on 13-5-2013 at 10:53


From an old text some related comments on HClO2 and its properties ("Records of General Science", Volume 2, edited by Robert Dundas Thomson, pages 341 to 342 at http://books.google.com/books?pg=PA342&lpg=PA342&dq=... )

"3. Properties of the Aqueous Solution of Chlorous Acid.— Chlorous acid, when diluted with water, is a transparent liquid, slightly coloured yellow when in a concentrated state."

Also:

"By the action of a strong light it is converted into chlorine and chloric acid, and sometimes, also, deutoxide of chlorine is formed."

Per Wikipedia on Chlorous acid (see http://en.wikipedia.org/wiki/HClO2 ):
"The pure substance [HClO2] is unstable, disproportionating to hypochlorous acid (Cl oxidation state +1) and chloric acid (Cl oxidation state +5):

2 HClO2 → HClO + HClO3 "

Also, "Chlorous acid is a powerful oxidizing agent, although its tendency to disproportionation counteracts its oxidizing potential" confirming the disproportionation comment by Thomson.

The alluded to formation of ClO2 follows from:

HClO2 + HOCl <--> HClO3 + HCl
HClO3 + HClO2 → 2 ClO2 + Cl2 + 2 H2O

Now, here is an educational reference (link: http://www.google.com/url?sa=t&rct=j&q=the%20decompo... ) on the hydrolysis of Cl2O3, to quote:

"5. The decomposition of chlorous acid, HClO2 in water has been suggested to proceed by the following mechanism

2 HClO2 ⇒ Cl2O3 + H2O rate coefficient k1
Cl2O3 + H2O ⇒ 2 HClO2 rate coefficient k–1
Cl2O3 + H2O ⇒ HOCl + HClO3 rate coefficient k2"

so there is both an equilibrium reaction between Cl2O3 and water along with a possible disproportionation reaction.

Now, as I previously noted, to quote again:

"Dilute HOCl solutions are colorless; at higher concentrations the color ranges from yellow to yellow-orange due to small equilibrium amounts of Cl2O."

so perhaps similarly increasing equilibrium amounts of Cl2O3 (dark brown) can add more intense color to explain the observed results.
---------------------------------

A way to possibly test this explanation, performing the experiment with strong (or no) light, or with hot solutions (or cold) also should effect the color intensity observed. Also, increasing HCl should move the reaction:

Cl2O3 + H2O ⇒ HOCl + HClO3

to the right (as Hypochlorous acid is removed via HCl + HOCl --> Cl2 + H2O), and reduces the amount of Cl2O3, impacting the solution's color. Per the same reaction, adding a little hypochlorite to the solution may be able to preserve the color intensity over a non-treated solution.
---------------------------------------------------

I found another reference describing a brown solution, to quote:

"A granulated Sodium Chlorite was dissolved in water to form a 48% sodium chlorite solution according to standard published data on solubility of Sodium Chlorite. This solution was then combined with a solution of 88% lactic acid. An immediate reaction occurred forming a deep brown solution."

Source: http://www.google.com/patents/US4892148


[Edited on 13-5-2013 by AJKOER]
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[*] posted on 18-5-2013 at 07:31


I did the following experiment:

Prepare 100 ml of pure chlorine gas (made by adding 10% HCl to solid Ca(ClO)2). Suck this gas in a syringe and wash with a little water, just to be sure that any HCl in the gas dissolves in the water.
Transfer the cleaned Cl2-gas into another syringe, assuring that the water with the HCl dissolved in it does not go into the other syringe. The gas in the second syringe is fairly pure Cl2, with possible contaminants being air and water vapor. None of these is of any concern for this experiment.

Prepare a solution of 30% NaClO2 and put this in a test tube. Appr. 3 ml of such a solution was prepared. Slowly press the gas from the second syringe into the 30% solution of NaClO2. Bubbles of Cl2 go to the surface. These bubbles quickly become smaller and are absorbed by the solution. Brown 'schlieren' are formed at the bubble and sink to the bottom.

I did the experiment until all of the chlorine from the syringe (which is 60 ml) was bubbled into the solution of NaClO2. After this, the liquid is dark brown and there is a strong yellow color of ClO2 above the liquid.

The reaction of Cl2 with chlorite is simple: 2ClO2(-) + Cl2 --> 2ClO2 + 2Cl(-).
No acid is produced in this reaction, just ClO2 and chloride. The ClO2 dissolves in the liquid and gives rise to the very dark brown color.

This experiment indicates against the hypothesis of Cl2O3 being the brown species. No acid is formed in this reaction and hence no HClO2 is formed. ClO2 itself seems to be the cause of the brown color. It might be that it is the combination of ClO2 and ClO2(-) which gives the dark brown color. The riddle is not solved yet.





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[*] posted on 18-5-2013 at 07:43


Did you try the H2SO4 + KClO3 experiment that I suggested?



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[*] posted on 18-5-2013 at 07:51


Quote: Originally posted by blogfast25  
Quote: Originally posted by woelen  
HClO2 is unstable and tends to disproportionate to chloride and ClO2.


Basically:

5 HClO2 === > 4 ClO2 + HCl + 2 H2O

That makes sense.


Uhh that reaction cannot be balanced. I have tried to balance similar equations with HClO2 yielding HCl and ClO2 but I can't balance any...




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