Could it be possible that chlorine dioxide is able to oxidize chlorine? This is an interesting question, and there are several reasons to think both
for and against.
2 ClO2 + 3 Cl2 --> 4 Cl2O ?
It is well known in industry that chlorine dioxide can be conveniently generated by passing chlorine gas into a solution of sodium chlorite.
2 NaClO2 + Cl2 --> 2 NaCl + ClO2
Here is a reaction where both chlorine and chlorine dioxide are present, and the chlorine dioxide apparently does not oxidize the chlorine. I would
suggest three potential reasons for this. First, there is a complex equilibrium established, such that even if some hypochlorous acid does transiently
form, it can still further react, such that the net reaction is still only results in the chlorine dioxide as the sole product. The second reason is
that during the reaction process, chlorine gas is usually passed into the solution of chlorite. This ensures an ideal ratio where the chlorite is
always in excess of the chlorine, as the chlorine dioxide tends to escape as a gas as soon as it is formed. If the chlorine dioxide was not allowed to
escape the solution, it might shift the reaction equilibrium to favor the hypochlorous acid as the product. Third, it may likely be that chlorine
dioxide has a slow reaction rate. While chlorine dioxide itself is soluble in water, even while it is dissolved its hydrolysis rate may be fairly
slow. Chlorine dioxide is essentially the mixed acid anhydride of chlorous acid and chloric acid. (This may be analogous to the solubility of carbon
dioxide in water, where carbonic acid forms.) It is likely that the full oxidizing potential of chlorine dioxide is not realised until it has time to
hydrolyse into its respective acids. This may explain why chlorine dioxide does not appear to readily oxidize chlorine gas when the two are mixed. The
oxidation may only be able to proceed through a complex series of steps, with some water necessary to catalyze the reaction, and the reaction would
also likely to be very slow.
It is well known that chlorine is able to reversibly disassociate in water, into hydrochloric and hypochlorus acids.
Cl2 + 2 H2O --> H3O+(aq) + Cl- + HOCl
It is also known that hydrochloric acid can reduce both HClO2 and HClO3. In the presence of both ClO2 (which has had adequate time to hydrolyse) and
Cl2, in water, a complex equilibrium should be established, whereby the net reaction is the oxidation of the chlorine into hypochlorus acid.
ClO2 + 2 H2O --> H3O+(aq) + ClO3- + HOClO
HClO3 + 2 HCl --> 3 HOCl
Obviously the products of the above reaction would depend on the ratio of hydrochloric acid present. A lower ratio might reduce the chloric acid to
ClO2, while excess hydrochloric acid would completely reduce it to chlorine.
So now it has been established that through their respective complex equilibriums with water, ClO2 should be able to oxidize Cl2 into HOCl. If there
is only a limited ammount of moisture present, such as if ClO2 is just reacted with moist Cl2 gas, it might be possible that Cl2O would form, since it
is the acid anhydride of hypochlorous acid. However, with such a limited quantity of water the reaction might be very very slow. To determine if
chlorine dioxide can indeed oxidize chlorine gas to dichlorine monoxide, a very long reaction time would be required, which may be quite challenging
considering the instability of both these chlorine oxides. It would probably be best done in the dark, and a assessment made to determine whether
there are even any chlorine oxides left at the conclusion of the experiment. If it is found that neither Cl2O nor ClO2 are left, obviously it
does not mean that there was no oxidation, but rather that the chlorine oxides were just not able to survive to the conclusion of the experiment.
It should be quite intuitive that a higher oxide of an element should be able to oxidize an additional portion of the free element to an intermediate
oxide. However, we have not yet considered reduction potentials.
(Since dichlorine monoxide is the acid anhydride of hypochlorous acid, we can assume it has an even higher reduction potential.) These values do not
seem to bode well for the hypothesis that there could be a reaction. However, it should be remembered that the reduction potential of chlorine dioxide
may be higher in acidic pH, or its effective value in terms of reactivity may be much higher after it hydrolyses, which may be an explanation why it
is so reluctant to hydrolyse.
I also found the following information from an educational chemistry book:
HClO2 + 2 H+ + 2 e- --> HClO + H2O 1.65V
"One of the concerns in using ClO2 as a disinfectant is that the carcinogenic chlorate ion (ClO3-) might be a by-product. It can be formed by the
following reaction.
6 ClO2 + 3 H2O <==> 5 ClO3- + Cl- + 6 H+(aq)"
The equation showed the reaction as an equilibrium, although I am not entirely which direction it prefers to go under normal conditions.Traveller - 26-1-2013 at 08:05
The quote at the end of your post refers to chlorate as being carcinogenic. Are there any studies to verify this and establish a maximum exposure
level?AJKOER - 26-1-2013 at 11:11
OK, there is the following reversal reaction:
8 HOCl <==> 4 H2O + 2 ClO2 + 3 Cl2
So, then ClO2/Cl2 participate, in the presence of water, in a redox reaction.
Could it be possible that chlorine dioxide is able to oxidize chlorine? This is an interesting question, and there are several reasons to think both
for and against.
2 ClO2 + 3 Cl2 --> 4 Cl2O ?
As:
8 HClO = 4 Cl2O + 4 H2O
substituting the above expression for HClO and cancelling out the common 4 H2O in the equation:
8 HOCl <==> 4 H2O + 2 ClO2 + 3 Cl2
appears to answer your question in the affirmative as for very concentrated HOCl solution, free Cl2O exists, but some water may still need to be
present, or perhaps UV light. Note, the use of polarized light irradiation (See Provisional Patent 60/747,705, filed May 19, 2006, link: http://www.freshpatents.com/-dt20100204ptan20100025226.php), may favor a single direction for the reaction.
For those who would appreciate an added reference, here is a reference from Mellor, page 288, (link: http://books.google.com/books?pg=PA288&lpg=PA288&dq=... )that mentions when the presence of sunlight, heat, chlorine, chlorides, ClO2 or
platinum can act as catalyst:
"The aq. soln. [referring to ClO2 in water] is fairly stable in darkness; in sunlight, it decomposes rapidly in a few hours; and slowly in diffused
daylight into chloric acid, HCl03, chlorine, oxygen: 6Cl02 + 2H20 = Cl2 + 02 +4HCl03. Some perchloric acid is formed at the cost of the chloric acid:
2HCl03 +02=2HCl04. The presence of chlorides accelerate the rate of decomposition such that a soln. with 0.15 mol. of chlorine dioxide suffered a 2
per cent, decomposition in five weeks in darkness at 0°, while with a normal soln. of chloride, there was a 70 per cent, decomposition. In the
presence of chlorides the reaction is represented: 6Cl02 + 3H20 = 5HCl03 + HCl; the velocity constants follow the relation d[C102]/dt =
—K[Cl02]2[HCl], and accordingly it is inferred that there is a slow reaction : 2Cl02+H20+HCl=2HCl02+H0Cl, followed by a rapid change:
6HCl02+3H0Cl=5HCl03+4HCl. Platinized asbestos also accelerates the reaction like chlorides. In the presence of chlorine, the reaction progresses: Cl02
+ 1/2Cl2 + H20=HCl03 + HCl, with the side reactions: 6Cl02 + 3H20 = 5HCl03 + HCl, and 3Cl2+3H20=HCl03+5HC1. At 60° another reaction : Cl02=1/2Cl2+02,
sets in. Consequently, the decomposition of aq. soln. of chlorine dioxide is very complex, for there are (i) 2Cl02=Cl2+202, which is accelerated by
raising the temp, or exposure to sunlight; (ii) 6Cl02 + 3H20=5HCl03+HCl, which is accelerated by the presence of chlorides or by platinum; (iii)
2Cl02+1/2Cl2+2H20=2HCl03 +2HCl, which is accelerated by chlorine; (iv) 3Cl2 + 3H20 = HCl03 + 5HCl, which is accelerated by platinum or chlorine
dioxide; and (v) 2Cl2+2H20=4HCl+02, which is accelerated by light."
[Edited on 26-1-2013 by AJKOER]AndersHoveland - 26-1-2013 at 21:53
While chlorine dioxide may have some complex equilibriums with chlorine, I think under most conditions ClO2 is very reluctant to react, and for the
most part it just exists as relatively inert ClO2 molecules solubilized in water. Similarly, while solutions of chlorous acid are very reactive, and
so too chloric acid, I suspect that if they are mixed together in equal ratios we would observe a sudden drop in reactivity (or at least reaction
rate), as the oxidizing acids would be tied up in the less reactive form of ClO2.
The quote at the end of your post refers to chlorate as being carcinogenic.
It is not really that dangerous. In fact there used to be chewable chlorate tablets to treat mouth sores, which were supposed to be rinsed out, not
swallowed. Chlorate is poisonous if ingested in more than very small ammounts. I do not think it is really carcinogenic unless it is chronically
ingested over a long period of time.
[Edited on 27-1-2013 by AndersHoveland]AJKOER - 28-1-2013 at 07:32
Yes, I agree under most conditions. But does 'most' include sunlight and or diffused light?
The quote at the end of your post refers to chlorate as being carcinogenic. Are there any studies to verify this and establish a maximum exposure
level?
Chlorate is not carcinogenic. It is somewhat toxic, but not so much that it is not safe to handle it.
Just use your common sense when working with it. Avoid exposure to larger amounts and if you work very frequently with it, then work clean and try to
avoid long term exposure. Avoiding that is easy, just regularly mop up your lab with a humid towel to collect any dust from the floor and from
workbenches. AJKOER - 29-1-2013 at 08:46
