Some background facts:

$\ce{ClO_2}$ is a gas that is highly soluble in water (some 8 g/L at 20 °C). However, per Wikipedia , to quote:

It does not hydrolyze when it enters water, and is usually handled as a dissolved gas in solution in water.

Although technically, there are many (including important) produced products with water, which I account for by the dissociation of chlorine dioxide, a stable free radical and subsequent hydrolysis:

$\ce{.ClO2 + .ClO2 <=> Cl2O4}$

$\ce{Cl2O4 + H2O <=> HClO2 + HClO3}$

A government source makes the statement:

In air, chlorine dioxide readily dissociates both thermally and photochemically and may form chlorine, oxygen, hydrogen chloride, HClO3, HClO4.ClO, chlorine peroxide, and/or chlorine trioxide, dependent on temperature and humidity. Chlorine dioxide dissociates in water into chlorite and chloride, and to a lesser extent into chlorate (Budavari et al. 1996).

Per a Dutch study: Chlorine Dioxide as a Post-Disinfectant for Dutch Drinking Water, to quote:

In this paper, results are presented of experiments into the consumption and reaction kinetics of chlorine dioxide in a number of (drinking) waters in The Netherlands. It was found that chlorine dioxide consumption is related to the dissolved oxygen content (DOC) of the water and the reaction time.

Lastly, a cited reaction between hydrogen peroxide and chlorine dioxide:

$\ce{H2O2 + 2 .ClO2 -> O2 + 2 HClO2 }$

Now, based on the Dutch study suggested connection to dissolved oxygen, I start with the last reaction above and use it as a thought tool to arrive at my postulate reaction as to how ClO2 decomposes in natural waters. The first step removes H+, next removes an electron from HO2-, next dissociates the radical $\ce{.HO2}$, then cancel H+ from both sides, merge an electron into ClO2 to produce ClO-, and in the final step cancel part of the chlorite ion present. The manipulation steps are:

$\ce{HO2- + 2 .ClO2 -> O2 + ClO2- + HClO2}$

$\ce{.HO2 + e- + 2 .ClO2 -> O2 + 2 ClO2- + H+}$

$\ce{H+ + .O2- + e- + 2 .ClO2 -> O2 + 2 ClO2- + H+}$

$\ce{ .O2- + e- + 2 .ClO2 -> O2 + 2 ClO2-}$

$\ce{.O2- + .ClO2 + ClO2- -> O2 + 2 ClO2-}$

Producing the final postulated reaction of this question:

$\ce{.O2- + .ClO2 -> O2 + ClO2-}$

Namely, I postulate that the presence of generated superoxide from various sources in natural oxygen-rich waters (from sunlight and transition metals,..., interacting with O2), which has been treated with chlorine dioxide, where superoxide concentration is connected to dissolved oxygen presence, is one of the more likely pathways to account for the reported removal of $\ce{ClO_2}$.

Comments welcome or evidence to assert otherwise.

[EDIT] Given interest per comments on the presence of superoxide in natural waters, I refer those interested to Measurement of Antioxidant Activity toward Superoxide in Natural Waters, which mentions a photochemical pathway for superoxide, which is likely the largest source for sunlit waters.

Superoxide can also be created by a so-called metal auto-oxidation reaction with say ferrous (as in $\ce{Fe(HCO3)_2}$):

$\ce{Fe(II) + O2 <=> Fe(III) + .O2-}$ (See also, Figure 1 here)

The presence of ferrous in natural waters is discussed in this paper, where, it should be noted, that absence light, transition metals and dissolved oxygen could be sources of superoxide. If this is but minimal, then chlorine dioxide should not be significantly consumed (as it is reported).

I would like to also propose a modified system, where oxygen and associated superoxide are but catalysts (note, the question as presented remains valid). In particular, the interaction of ferrous with chlorine dioxide:

$\ce{Fe(II) + .ClO2 -> Fe(III) + ClO2-}$

And, the cycling of a soluble ferric via the reversible metal auto-oxidation reaction:

$\ce{Fe(III) + .O2- <=> Fe(II) + O2}$

assuming such a soluble ferric species exists here (perhaps ferric bicarbonate), else my original path proposal is unchallenged. Note, the net of the last two reactions closely resembles the prior proposal reaction, but cites a transition metal presence:

$\ce{.O2- + .ClO2 -(Fe(II),Fe(III))-> O2 + ClO2-}$

  • 1
    $\begingroup$ I think I´m missing sth here. Where is this supposed to take place? Your municipal tap water should not contain larger amounts of transition metals, and sees no sunlight. $\endgroup$ – Karl Apr 5 '20 at 22:07
  • $\begingroup$ Thanks for the review. And yes, likely limited superoxide presence which is why ClO2 is likely employed (so called post-treatment) and remains largely effective. The science and particulars of superoxide generation are not included (but could be if someone wishes to base a master thesis on the suggested mechanics). $\endgroup$ – AJKOER Apr 5 '20 at 22:24
  • $\begingroup$ Karl, I have noticed some curious criticism from you, and just noticed, we are neck and neck in the reputation points for the quarter! Any connection? But congratulations on your success!! $\endgroup$ – AJKOER Apr 5 '20 at 22:28
  • $\begingroup$ I have edited my question with an article, 'Measurement of Antioxidant Activity toward Superoxide in Natural Waters', targetting the measurement of superoxide which, per the source, can also be photochemically produced . $\endgroup$ – AJKOER Apr 6 '20 at 0:43

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