Sciencemadness Discussion Board

A new remarkable riddle: what is this dark compound?

woelen - 11-5-2013 at 12:03

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]

Adas - 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.

blogfast25 - 11-5-2013 at 12:41

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

AJKOER - 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]

platedish29 - 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!

woelen - 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.

blogfast25 - 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]

Eddygp - 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]

AJKOER - 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.

woelen - 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]

papaya - 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..

bbartlog - 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?

blogfast25 - 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]

AJKOER - 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.

bbartlog - 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.

platedish29 - 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]

woelen - 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).

blogfast25 - 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]

woelen - 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.

garage chemist - 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.

blogfast25 - 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.

AJKOER - 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]

woelen - 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.

garage chemist - 18-5-2013 at 07:43

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

Eddygp - 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...

blogfast25 - 18-5-2013 at 09:13

A very nice and to the point experiment, indeed. This also appears one of the better ways to prepare ClO2.

woelen - 18-5-2013 at 10:39

I also tried the experiment with KClO3 and H2SO4. I knew this one already, but I tried it again to watch the precise color of the liquid obtained in this experiment. This liquid, however, is really different from the brown color I obtain in the other experiments. KClO3+H2SO4 gives a bright orange liquid and not a dark brown liquid. It gives ClO2 (visible yellow gas is produced and floats around the pile of KClO3 and the drop of liquid) and it gives white fumes. I think it is disproportionation of HClO3 to HClO4 and ClO2 and the orange color of is due to concentrated ClO2, dissolved in the acid.

@eddygp: I do not understand your remark. The equation, provided by blogfast25, is perfectly balanced.

blogfast25 - 18-5-2013 at 11:06

Quote: Originally posted by Eddygp  

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...


Like all redox reactions it helps a lot if you break them down into oxidation and reduction ‘half reacions’. Here:

(1) HClO2 === > ClO2 + H+ + e- (oxidation to Cl (IV) oxide)
(2) HClO2 + 3 H+ + 4 e- === > Cl- + 2 H2O (reduction to chloride)

4 x (1) + 1 x (2) and mild rearranging then yields the balanced result.

Eddygp - 18-5-2013 at 11:35

Quote: Originally posted by woelen  

@eddygp: I do not understand your remark. The equation, provided by blogfast25, is perfectly balanced.

My bad, I don't see why I said that... I thought there was a hydrogen missing or something. Sorry.

woelen - 18-5-2013 at 12:12

I also tried what happens when a big drop of Br2 is added to a solution of 30% NaClO2. When this is done, then the liquid becomes very dark, nearly black and an intensely colored yellow gas mix slowly appears above the liquid. The bromine quickly dissolves in the solution of NaClO2 and the entire liquid becomes very dark. The dark liquid slowly produces fairly big bubbles of gas.

I lighted the open end of the test tube with the dense yellow gas mix and this gives an impressive sound, the well-known WHOOSH sound when hydrogen/air mixes are ignited. The sound, however, was powerful (quite scary, I held the test tube in my hand, but wrapped in a 2 cm thick layer of towel). After the 'burning' of the ClO2, a pale brown/yellow color remained in the gas-mix in the test tube (due to some left over bromine vapor) and slowly the gas-mix above the dark liquid turns yellow again.

Most scary was that after the lighting of the ClO2 in the test tube, there was a strong crackling noise just above the dark liquid. The dark liquid was bubbling and each time when a bubble made it to the surface, it exploded with a loud CRACK noise. Probably this was due to the high temperature in the gas mix. I was afraid that the entire dark liquid would explode, but that fortunately did not happen.

So, Br2 also is capable of oxidizing ClO2(-) and then even more of the intensely dark colored compound is formed.

AJKOER - 19-5-2013 at 05:09

Woelen:

I not sure I that adding Cl2 to aqueous solution eliminates the acid question as some Cl2 can react as follows:

Cl2 + H2O <---> HCl + HOCl

-----------------------------------------------
I would also add, in my opinion, that the picture you provided appears to support my hypothesis as any Cl2 formed in the test tube (per decomposition or disproportionation reaction cited previously) would bubble to the top. Then, again per the reaction cited above, a locally higher concentration of HCl could form removing the dark brown Cl2O3 species per the reactions:

Cl2O3 + H2O = HOCl + HClO3
HOCl + HCl = Cl2 + H2O
---------------------------------
Net:
Cl2O3 + HCl = Cl2 + HClO3

Conversely, the lower half of the test tube is where the more intense darker color is expected (lower local HCl concentration), and it appears also to be so observed.
--------------------------------------------------------

[EDIT] Here is an interesting tests consistent with my reaction chain model and properties of compounds formed:

First, the simpliest test, with time/light HOCl decomposition should increase forming HCl, and the solution's color could regionally be impacted. Similarly, adding a drop of H2O2 should more rapidly form O2 and remove color.

Second test, cooling and warming a fresh solution should affect color.


[Edited on 19-5-2013 by AJKOER]

blogfast25 - 19-5-2013 at 06:06

Quote: Originally posted by woelen  

So, Br2 also is capable of oxidizing ClO2(-) and then even more of the intensely dark colored compound is formed.


Very interesting but can we know for sure it is the same compound generated with Br2 as with Cl2? As such, in the absence of identifying evidence like UV/VIS data, I say 'no'. It would se SOOO interesting to run these reactions at low concentration in cuvettes and take spectra, possibly as a time series?


[Edited on 19-5-2013 by blogfast25]

AJKOER - 19-5-2013 at 10:33

With respect to the Bromine experiment, here is a reference ("The Reaction of Chlorine(III) with Bromine and Hypobromous Acid: Kinetics and Mechanism", http://www.google.com/url?sa=t&rct=j&q=hclo2%20%2B%2... )

Some reactions presented include:

Br2 + ClO2− = Br2− + ClO2
Br2− + ClO2− = ClO2 + 2Br−
Br2 + Br− = Br3−
ClO2− + H+ = HClO2
H2O + Br2 = HOBr + Br− + H+
Br + Br− = Br2−
Br + Br = Br2
Br + ClO2− = Br− + ClO2
HOBr + HClO2 = BrClO2 + H2O
BrClO2 + ClO2− = 2ClO2 + Br−

where Br is a bromine radical and Br− is the bromide ion. I would add that bromine does not measurable dissolve in water until some bromide is created, then forming the tribromide. Note, several of the reactions cite the formation of ClO2 as was observed.

Another source (http://pubs.acs.org/doi/abs/10.1021/es302730h ), which includes the pesence of Chlorine states:

"HOBr, formed via oxidation of bromide by free available chlorine (FAC), is frequently assumed to be the sole species responsible for generating brominated disinfection byproducts (DBPs). Our studies reveal that BrCl, Br2, BrOCl, and Br2O can also serve as brominating agents"

As such, Br2O (and even perhaps, speculatively BrO2, formed by the action of Cl2O3/HClO2 on Br2O or BrClO2) could be responsible for the exploding bubbles. The described pale brown/yellow gas was most likely BrCl described by Wikipedia as being brownish yellow.


[Edited on 19-5-2013 by AJKOER]

woelen - 19-5-2013 at 11:30

I did another experiment, now using S2O8(2-) as oxidizer. This oxidizer has a very clean reaction:

S2O8(2-) + 2e --> 2SO4(2-)

No acids involved, no free halogens involved.

I expect that with chlorite the reaction is as follows: S2O8(2-) + 2ClO2(-) --> 2So4(2-) + 2ClO2

Now the result of my experiment:

- Prepare a 30% by weight solution of NaClO2
- Add some solid Na2S2O8 to this solution.
- As soon as the white solid is added, the solid becomes covered by a yellow/ochre layer, which soon becomes dark brown. Initially, the layer has a color like mustard and not the bright orange which appears in the KClO3/H2SO4 experiment as suggested by garage chemist. By slowly swirling the solution, nice brown 'schlieren' of dark solution are obtained. When swirled a little stronger, a homegeneous very dark solution is obtained.
- Wait some time: The reaction proceeds and goes faster. The liquid becomes nearly black and bubbles are produced. Above the liquid there is a strong and deep yellow color of (nearly) pure ClO2.
At this point I added a lot of cold water. I had the unpleasant feeling that the reaction went too fast and I did not feel comfortable with it. The concentration of ClO2 was very high and I feared a sudden explosion, hence the quenching in a lot of water.

This experiment nicely shows that just ClO2 without acid and without free halogens produces the dark brown color in a solution of NaClO2. This experiment also tells more than the experiment with Br2, because of the clean reaction of peroxodisulfate in redox reactions.

AJKOER - 19-5-2013 at 13:11

Here is a new theory based on the chemistry presented in the patent (see "Method and composition for in-situ generation of chlorous acid" by Roy W. Martin, link: http://www.faqs.org/patents/app/20120107418 ) namely we are possibly viewing an aqueous intermediary compound, to quote from the patent:

"[0049] Where X represents Bromine (Br) and Chlorine (Cl)

2ClO2- + X2(g) → 2ClO2 + 2X- (1a)

2ClO2- + HOX → 2ClO2 + X- + OH- (1b)

[0050] However, these equations give a simplistic representation of the generation process. Considering the mechanism of these reactions is important for a better understanding of the details of the generation process and the invention. The intermediate species (XClO2) forms in these reactions. This intermediate may react to give ClO2 or chlorate ion according to Equations 3-4.

Where X represents Bromine (Br) and Chlorine (Cl)

X2 + ClO2- → [XClO2] + X- (2)

2[XClO2]→2ClO2 + X2 (3a)

[XClO2] + ClO2- → 2ClO2 + X- (3b)

[XClO2] + H2O → ClO3- + X-+ 2H+ (4)

[0051] Equations 3a-b are important at high concentrations when the formation of XClO2 is rapid. On the other hand, Equation 4 is more important when the formation of XClO2 is slow, such as at low reactant concentrations or high pH values. "

That is, my new suggested theory is that one is viewing the intermediary compound Dichlorine dioxide, Cl2O2 in solution (or BrClO2) created by the action of Chlorine (or Bromine), per the reactions cited above, per the high concentration reactions with Chlorine (or Bromine), namely for NaClO2, as an example:

3 Cl2 + 3 NaClO2→ 3 [Cl2O2] + 3 NaCl (per 2)

2[Cl2O2] → 2ClO2 + Cl2 (per 3a)

[Cl2O2] + NaClO2 → 2ClO2+ NaCl (per 3b)

or net:

4 NaClO2+ 2 Cl2 → 4 ClO2 + 4 NaCl

or, rescaling:

2 NaClO2 + Cl2 → 2 ClO2 + 2 NaCl

in agreement with reaction (1a).

Note, per Wikipedia (http://en.wikipedia.org/wiki/Cl2O2 ) Dichlorine dioxide is a naturally occurring reaction intermediate formed through the combination of two chlorine monoxide radicals during atmospheric photochemical reactions.

The cited patent above describes sample solutions with varying colors of gold, darker gold and amber.

[EDIT] This model suggests that on heating the solution's color should vanish (along with the intermediary), and with varying solutions at different increasing temperatures, the color should fade more rapidly per the relative temperature.

Also, with respect to the reaction: S2O8(2-) + 2ClO2(-) --> 2So4(2-) + 2ClO2, the short answer is that this could be a net reaction also.


[Edited on 19-5-2013 by AJKOER]

papaya - 19-5-2013 at 13:32

I'm afraid to propose a stupid thing,but what is the color of solution if you pass chlorine gas into HCL?

Oscilllator - 20-5-2013 at 00:16

papaya I have done this as part of an unrelated experiment and the solution turned slightly yellow if anything.

papaya - 20-5-2013 at 01:07

That I asked to know what is the color of HCL3 because I thought in high concentrations it may become darker.

blogfast25 - 20-5-2013 at 04:09

Quote: Originally posted by papaya  
That I asked to know what is the color of HCL3 because I thought in high concentrations it may become darker.


"HCL3"? What's that supposed to be?

papaya - 20-5-2013 at 04:16

That's a chlorine analog of H[J3] , there are polyhalide anions like in KJ*J2 whichmore precisely is K[J3] complex...
https://en.wikipedia.org/wiki/Polyhalide

[Edited on 20-5-2013 by papaya]

blogfast25 - 20-5-2013 at 04:21

Quote: Originally posted by AJKOER  

That is, my new suggested theory is that one is viewing the intermediary compound Dichlorine dioxide, Cl2O2 in solution (or BrClO2) created by the action of Chlorine (or Bromine), per the reactions cited above, per the high concentration reactions with Chlorine (or Bromine), namely for NaClO2, as an example:



[Edited on 19-5-2013 by AJKOER]


The author is clearly freewheeling with regards to 'XClO2' and presents not a shred of evidence for its existence, instead piling up hypothesis upon hypothesis. Not a spectrum in sight, let alone an elemental analysis of 'XClO2', stoichiometrical proof or a MW determination. It's all conjecture. NOT impossible but UNPROVED.

Patents, by their very nature, can be very deceitful.

[Edited on 20-5-2013 by blogfast25]

AJKOER - 20-5-2013 at 06:15

Blogfast:

Thanks for asking for further research on Cl2O2. I found the precise reactions (and more) discussed in detail in a good book "Inorganic Chemistry", edited by Arnold F. Holleman, Egon Wiber, Nils Wiberg at link: http://books.google.com/books?id=Mtth5g59dEIC&pg=PA459&a... at the bottom of page 460.

The author clearly states that Cl2O2 (or ClClO2) is actually reputedly formed in several reactions of concern here. It also is stated to readily decomposes. There is, however, a slow rate determining step that may keeps the dichlorine dioxide visible for awhile depending on reaction conditions. For example:

HClO2 + HClO2 --H+--> HClO + HClO3 (slow rate determing step)

HClO + HClO2 --> ClClO2 + H2O

ClClO2 ---> 1/2 Cl2 + ClO2 (rapid)

Net to this point:

3 HClO2 --H+--> HClO3 + H2O + 1/2 Cl2 + ClO2

There is also a reaction forming more ClClO2 from the Cl2:

1/2 Cl2 + 1/2 HClO2 --> HCl + ClClO2

The author also directly states that the reaction cited in the patent (Equation 10 on page 460):

2 NaClO2 + Cl2 --> 2 NaCl + 2 ClO2

has ClClO2 also as an intermediary.
------------------------------------------------------

OK, so the chemistry look good, but as to whether it is the actual cause of the color, I am not certain, but it appears to qualify as a possible candidate.

[Edited on 20-5-2013 by AJKOER]

bbartlog - 20-5-2013 at 06:50

Mind you, ClClO2 is not the same compound as the symmetrical Cl2O2 in the wikipedia link that you posted earlier. It looks like it would involve one chlorine atom with nominal oxidation state +1 and another at +3, which at least looks interesting as a possible charge transfer complex.
Another relevant paper I found, with a nice set of cited work that I haven't time to look through right now:
http://hopf.chem.brandeis.edu/pubs/pub293%20rep.pdf

AJKOER - 20-5-2013 at 08:40

Bbartlog:

I have looked at the excellent research paper you supplied.

I also found interesting that Cl2O3 can be created via the pathway:

ClO2- + Cl2O2 --> Cl- + 2 *ClO2
HClO2 --> *ClO + *OH

Leading to:

*ClO2 + *ClO --> Cl2O3

so I would add Cl2O3, as one of the "longer-lived" (per the authors words on the first page) intermediary species back on the candidate list as well. Also, around equation E3 on page 6970, the authors cite three reason for re-introducing Cl2O3.

Bottom line, for this well-studied decomposition reaction the species observed are HOCl, Cl2O2, Cl2O3, *ClO and *OH. And, as the radical *ClO interact with itself to form the symmetric form of Cl2O2 in a termination step, I think we have at least mentioned all the possible candidates responsible for the color.
-------------------------------------------------

Here is the last candidate and may even be the answer. The conditions are some ClO2 is formed in solution containing a high level of chlorite. Some reactions:

HClO2 <--> H+ + ClO2-
ClO2 + H2O <--> HClO2 + HClO3

[EDIT] My research (http://www.google.com/url?sa=t&rct=j&q=hydrolysis%20... ) notes an interesting property of ClO2, to quote:

"One of the most important physical properties of chlorine dioxide is its high solubility in water,
particularly in chilled water. In contrast to the hydrolysis of chlorine gas in water, chlorine dioxide
in water does not hydrolyze to any appreciable extent but remains in solution as a dissolved gas
(Aieta and Berg, 1986). It is approximately 10 times more soluble than chlorine (above 11°C), while
it is extremely volatile and can be easily removed from dilute aqueous solutions with minimal
aeration or recarbonation with carbon dioxide (e.g. softening plants). Above 11 to 12°C, the free
radical is found in gaseous form."

As such, the last hydrolysis reaction above is pretty much to the left and with a high chlorite concentration, basically just unreactive ClO2 gas in water.

This site notes that at high concentrations of ClO2, a significant brown appearance (link: http://www.linkedin.com/groups/Can-anyone-explain-difference... ). To quote:

"H. Peter H. • You are walking a dangerous line when you see brown fumes - chlorine dioxide is not just toxic, it is also explosive. In dilute solution or as a dilute gas, it is yellow, and with increasing concentration, it is getting more and more orange-coloured. When you see brown, it is reason to be alarmed."

So the answer could just be free ClO2. A test for this would be to check the color sensitivity to HCl, HOCl and/or Cl2 as these would promote the ClO2 hydrolysis by attacking the products. Another test, add a few drops of seltzer water (CO2/H2O). If the color rapidly vanishes, it is ClO2 and not an intermediary.


[Edited on 21-5-2013 by AJKOER]

blogfast25 - 20-5-2013 at 10:29

AJ: I didn't really ask you for anything but anyway.

Although all this makes 'Cl2O2' (but which of the three?) a potential candidate to explain the dark substance, without more direct 'fingerprinting' evidence re. the exact nature of that substance everything here remains speculation. Unless woelen can come up with indirectly corroborating experiments that an 'XClO2' type species is at play here, we've basically come to the end of the line as fat as I'm concerned.

kmno4 - 20-5-2013 at 11:37

Reaction ClO2 + chlorite -> "deeply brown something" is known (in chemical literature).
But I do not want to break up the party here.
Have a nice searching (and finding).
:P

woelen - 22-5-2013 at 12:52

Indeed, the combination of ClO2(-) and ClO2 causes the brown color. If you prepare a neutral solution of ClO2 in water then the solution becomes deep yellow. If you want to repeat this experiment, be careful! let concentrated ClO2 diffuse into a small pool of water, do not suck ClO2 in syringes or other plastic/rubber things!

If, however, NaClO2 is present in the water, then the color is different, it is more brown/mustard. I did not do experiments with very concentrated solutions of ClO2 with NaClO2, because that requires handling larger amounts of high concentrations of ClO2 and I fear its explosive power too much. The experiments I did with lower concentrations, however, are quite convincing and the other experiments I did before also give very good evidence (especially the one with Na2S2O8 which affords production of ClO2 only without adding any acid in the mix).

I also read about this subject on internet. Remarkably, the brown compound is mentioned nowhere in scientific documents, which I can access. I'm quite sure other people must have noticed it as well, but I did not find any scientific paper describing it. I found it mentioned once, not in a reputable paper, but on some site which describes making MMS (a miracle mineral solution which is supposed to be good for almost everything and more of that kind of crap). This site warns for the brown solutions and the yellow gas, telling that it can explode and that it is dangerous and if this color is observed that the concentration used in the preparation is way too high. More serious documents, however, describe single electron transfer reactions in redox systems:

ClO2(-) <---> ClO2 + e(-)

This reaction occurs very easily, according to literature and can go in both directions. Many oxidizers easily make ClO2 from ClO2(-) and many reductors easily make ClO2(-) from ClO2.

I have the impression that the brown color is due to some single electron transfer reaction between ClO2 and ClO2(-). I can imagine that ClO2 and ClO2(-) become bonded to each other in some way and that the electron 'flips' from one side to the other at extremely high frequency (i.e. a resonance structure is formed which at least partially overlaps both ClO2-entities) and that this process leads to strong absorption in the visible spectrum (due to the relatively long distance, covered by this resonance system). I, however, did not find a true reference which confirms my theory, nor do I have the equipment to validate this. My measurement devices are my eyes and other more advanced equipment I have not at home for doing this kind of observations.
I speculate that this species will have a structure like O2Cl...ClO2 with ... being the special bond which is formed by the lone electron of ClO2 on its Cl-atom and a free electron pair of ClO2(-), also on its Cl-atom. The molecular orbital may extend even further, into the oxygen parts of the molecules.

As I said, there is quite some speculation from my side, but given the limited resources I have, I think it is the best I can offer.

AJKOER - 22-5-2013 at 14:51

I think there may also be an interesting safety question here as well. Obviously, dissolving a large amount of ClO2 in water can form a dangerous/potentially explosive solution. However, in the presence of a high level of chlorite, much less ClO2 may be able to produce a similar colored solution (or, as I would describe it, per a visible small equilibrium amount of free ClO2 effected by a limited hydrolysis in the presence of a soluble chlorite). The safety question is does this latter solution present the same level of danger, or is this the case of a false positive test?

My suspicion per Patent US 4892148 A at http://www.google.com/patents/US4892148 from the example cited below to quote:

"EXAMPLE 3
A granulated Sodium Chlorite was dissolved in water to form a 48% sodium cholite 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. This solution was tested and the presence of ClO.sub.2 was detected. No attempt was made to ascertain the ClO.sub.2 ppm of this solution."

is that the solution is not dangerous (in other word, a false positive) given the particular nature of the patent involving bulk mixing for oil recovery operations.

[EDIT] Further, given that the solution above has high concentrations of chlorite (from aqueous NaClO2) and HClO2 (via the lactic acid and NaClO2 reaction), it is also my opinion that aqueous Cl2O3 (the acid anhydride of HClO2 isolated as a dark brown explosive solid) is as plausible an explanation for the immediately observed deep brown color as is ClO2, being a yellowish-green gas which crystallizes as bright orange crystals at −59 °C, which "in aqueous solution, depending on concentration, the colour varies fiom green-yellow to orange-red" per http://www.google.com/url?sa=t&rct=j&q=%22chlorous%2... page 11. I would also note that I would suspect ClO2 to be present in much smaller amounts via decomposition/disproportionation reactions, in this particular example.


[Edited on 23-5-2013 by AJKOER]

ScienceSquirrel - 22-5-2013 at 16:03

You are never going to get an explosion from a solution of sub 10% of anything.
Water is a fantastic thermal buffer.
Mixing 10ml of 35% hydrogen peroxide with 30ml of 10% sodium hypochlorite, roughly stochiometric proportions, produces a vigorous effervescence of oxygen and a very hot solution but it does not explode.
A mixture of 35% hydrogen peroxide with 98% sulphuric acid, acid Piranha, can explode in contact with acetone, etc but it not an aqueous solution.
The fact is that; solid or solution, and if a solution, concentration play an important role in a materials properties.

phlogiston - 23-5-2013 at 04:15

Quote: Originally posted by woelen  
I can imagine that ClO2 and ClO2(-) become bonded to each other in some way and that the electron 'flips' from one side to the other at extremely high frequency (i.e. a resonance structure is formed which at least partially overlaps both ClO2-entities)


It is incorrect to think of molecules rapidly switching between resonance structures. This concept originates from quantum chemical descriptions of molecules. The molecule is instead a static object (in the sense that no bonds are broken or formed, i.e. all electrons remain in their respective orbitals). Intuitively, you can view a molecule as being a weighted average of all the resonance structures that you can draw. The structures with the lowest energy are more likely and contribute more, (and very unlikely structures contribute very little).

As an example, consider the nitrate ion:


In reality, the negative charge on the oxygen atoms is static in time and distributed equally among the oxygen atoms. The N-O bonds are not a single bond or a double bond but more like a 1.333 bond (in terms of electron density).

I am not qualified to comment on the likelyhood of your Cl2O4- structure, but interesting observations, as always.

[Edited on 23-5-2013 by phlogiston]

[Edited on 23-5-2013 by phlogiston]

kmno4 - 23-5-2013 at 05:13


"Resonance structures" are only mathematical trick, not existing in reality. It is commonly known, master.

But woelen's " electron 'flips' " are something different and correspond to charge-transfer: really existing phenomenon.
See here, for some informations:
http://en.wikipedia.org/wiki/Intervalence_charge_transfer
or
http://en.wikipedia.org/wiki/Mixed-valence



AJKOER - 23-5-2013 at 05:49

Quote: Originally posted by ScienceSquirrel  
You are never going to get an explosion from a solution of sub 10% of anything....


OK, ScienceSquirrel yes and no. More ClO2 in solution and given its ability to readily exit solution (per a decrease in pressure or increase in temperature, as it does not hydrolyze like Cl2) is where the danger lies. To quote from someone in the business of using ClO2:

Bhupen P.: "When in gas phase about 10 kPa partial pressure, roughly 30% in air ClO2 may decompose from a shock or hotspots into chlorine and oxygen with explosive force."

Source: http://www.linkedin.com/groups/Is-Clo2-in-Aquous-form-178984...

So even a 10% solution of dissolved ClO2 in solution could potentially rapidly lead to a explosive level of ClO2 and air above the solution (per, for example, a sharp drop in pressure or increase in temperature)which could further trigger an explosion from a pressure wave or local hot spot. For safety, keep aqueous solutions of ClO2 cold.

Per another source (see http://www.scribd.com/doc/30120967/Chlorine-Oxygen-Acids-and... ) to quote:

"An equation for calculating the partial pressure of chlorine dioxide above specified chlorine dioxide solutions at various temperatures based on the data from reference [24] has been developed (25):

pClO2 =(g/L ClO2) e^[10.717−(3102/T)] (1)

where pClO2 is the partial pressure of chlorine dioxide gas in kPa, g/L ClO2 is the chlorine dioxide solution concentration in grams per liter, and T is the absolute temperature in Kelvin."

This is important as another source (see http://www.google.com/url?sa=t&rct=j&q=%22chlorous%2... ) to quote from page 11, Chlorine dioxide "decomposes explosively at >300 mm Hg partial pressure" so there is a valid safety concern relating a ClO2's solution concentration (and/or temperature), and proneness to explosion.

[EDIT] Hopefully, the foregoing discussion will foster more knowledge of ClO2 and some of its unique properties (particularly its ability to rapidly exit solution) to facilate safer experimentation with Chlorine dioxide.


[Edited on 23-5-2013 by AJKOER]

phlogiston - 23-5-2013 at 06:06

Quote: Originally posted by kmno4  

But woelen's " electron 'flips' " are something different and correspond to charge-transfer: really existing phenomenon.


Indeed. I just mentioned it because Woelen called it resonance structures, which they aren't.

woelen - 23-5-2013 at 06:26

@kmno4: If I read your links, then you are suggesting a structure with bridging oxygen atoms between the Cl-atoms instead of an O2Cl...ClO2 structure? This is what the last link you provided describes.

I cannot draw here, but I then would either suggest a structure like O-Cl-O...O-Cl-O, or even two bridges in parallel to each other with the Cl-atoms at the end sites, something like

Cl(-O...O-)2Cl ,

where the total charge of this object is -1.

The structure (-O...O-) is not a peroxo structure, but part of an intervalence charge transfer structure as described in the wiki page.

If you have a literature reference for me, then I would like to have that. Then I'll modify my webpage and add the reference to my webpage (of course not a copy of the article itself, due to copyright ownership issues).

I have the feeling that I am coming closer to the solution of this riddle.

[Edited on 23-5-13 by woelen]

woelen - 28-5-2013 at 01:45

I finally found a reference which mentions the charge transfer between ClO2 and ClO2(-). They did analysis with a special radioactive isotope of chlorine:

http://pubs.acs.org/doi/abs/10.1021/ja01175a077?journalCode=...

Unfortunately I cannot read the entire article, but the first page gives sufficient information. I still have no 100% certainty about the nature of the compound, most likely it is something like [Cl(-O...O-)2Cl](-), which is in equilibrium with ClO2 + ClO2(-).

Bedlasky - 8-3-2021 at 02:33

I found document which talks about formation of [Cl2O4]- complex.

https://www.oieau.org/eaudoc/system/files/documents/39/19857...

Similar complex between NO and NO+ was disscused here:

https://www.sciencemadness.org/whisper/viewthread.php?tid=15...

[Edited on 8-3-2021 by Bedlasky]

woelen - 8-3-2021 at 12:58

Interesting read about the Cl2O4(-) complex. In that article it simply is taken for granted, but it certainly is not common knowledge. This complex allows one to reach very high concentrations of ClO2, bound to a solution of chlorite. I have done further experimenting with this a few months ago (also with Br2 and NO2). Quite a few oxidizers are capable of oxidizing ClO2(-) to ClO2 in neutral solution and if that is done, you get a nearly black liquid. I once had a nasty near-explosion from a saturated solution of NaClO2 through which I passed chlorine gas. Suddenly the black liquid made a lot of loud crackling noises and then the entire contents was ejected from the test tube. Droplets of the black liquid expelled clearly visible yellow gas. Scary stuff when at such high concentrations.

Cl2 + 2ClO2(-) --> 2Cl(-) + 2ClO2

ClO2 + ClO2(-) <---> Cl2O4(-) (the latter is nearly black when this is done with appr. 30% NaClO2 in water).