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Oxygenated water has become popular for a variety of uses recently - for health purposes internally as well as enhanced wound healing.

Anyhow, my dad asked me recently if I knew how it is that oxygenated water does not become hydrogen peroxide. I don't know the answer to it; but it got me thinking... How does this work? Does the oxygen not bond to the water molecule? If so, how does the bottle of water not just out gas when ever it is opened until the amount of oxygen contained in the bottle is equal to that of the room air.

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    $\begingroup$ Oh, if oxygenated water did change to hydrogen peroxide, I shouldn't have been alive now! Some reactions, just don't like to happen. $\endgroup$ – M.A.R. Feb 7 '15 at 15:55
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    $\begingroup$ Sorry but anyone foolish enough to buy oxygenated water is an idiot wasting money. $\endgroup$ – MaxW Jan 7 '18 at 18:43
  • $\begingroup$ @MaxW I would have to agree! I just was more curious about the chemistry of the whole thing :) $\endgroup$ – L.B. Feb 9 '18 at 18:43
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The oxygen is dissolved in the water, just as salt can be dissolved. It does not (to any appreciable degree) combine with water molecules to form hydrogen peroxide.

The reason that oxygenated water is not fizzy like soda water (CO2) is the solubility of oxygen in water is about 2% that of CO2. So the amount of bubbles formed when the cap is removed is only about 1/50 as much as for soft drinks. Or champagne.

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  • $\begingroup$ When CO2 is dissolved in water, part of it reacts with H2O to produce H2CO3. Similarly part of the dissolved oxygen reacts with H2O to produce H2O2, doesn't it? If not why? If yes, what's the equilibrium? $\endgroup$ – Petr Pudlák Feb 7 '15 at 9:19
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    $\begingroup$ @PetrPudlák Not similar at all. The formation of peroxide is very unfavourable and requires a lot of energy. So it doesn't happen easily or spontaneously, totally unlike carbonic acid. $\endgroup$ – matt_black Feb 7 '15 at 12:16
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I think the problem is that the reaction:

$\ce{2H2O2-> 2H2O + O2}$

is spontaneous, meaning that it proceeds (slowly) in standard conditions with no other inputs required. So far from oxygen combining with water, you have hydrogen peroxide spontaneously decomposing into water and oxygen.

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First you need to know what is oxygenated water. It's actually form by adding additional oxygen to the water under pressure. It means it's a mixture, not a compound.

We know that a mixture is two or more substances mix together physically. However, hydrogen peroxide bonds chemically which is a compound.

So, we can conclude that oxygenated cannot be classified as hydrogen peroxide.

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With respect to the question, 'how it is that oxygenated water does not become hydrogen peroxide', one answer is that in natural sunlight waters it actual does create H2O2, albeit even then, in only small amounts.

More precisely, the presence of H2O2 and other reactive oxygen species (ROS) have been cited in the literature as occurring in illuminated natural waters (see, for example, ‘Reactive Oxygen Species in Natural Waters ‘, by Neil V. Blough and Richard G. Zepp). To quote from an online two-page preview :

The presence of ROS in aquatic ecosystems was first reported in 1966 by Van Baalen and Marler, who detected the presence of hydrogen peroxide (H202) in tens to hundreds of nanomolar concentrations in surface seawaters... Although their data were inadequate to specify the source of the H202, they postulated that photochemical reactions, biological processes, or atmospheric deposition might contribute to its presence in seawaters. Some years later, Swallow (1969) suggested that the hydrated electron ($\ce{e-(aq)}$) could be produced in seawater by the action of...photoionization of endogenous phenolic compounds. He further proposed that rapid reaction of $\ce{e-(aq)}$ with dioxygen (02) would yield superoxide ($\ce{.O2-}$), ultimately leading to the formation of H202 via disproportionation….

My updated understanding of these reactions is presented below which starts with a solvated electron, $\ce{e-(aq)}$ and $\ce{H+}$ per the reactions:

$\ce{O2 + e-(aq) ⇌ .O2- (aq)}$ (superoxide radical anion)

$\ce{H+ + .O2- (aq) ⇌ .HO2 (aq)}$ (hydroperoxyl radical with a pKa = 4.88)

$\ce{.HO2 (aq) + .HO2 (aq) -> H2O2 + O2 (aq)}$ Source: Equation (1)

In natural waters in sunlight, dissolved organic matter (DOM) can be a source of solvated electrons (see 'Photoproduction of hydrated electrons from natural organic solutes in aquatic environments') and electron holes:

$\ce{DOM + hv -> e- + h+}$

Also, the presence of transition metal ions and dioxygen via so-called metal auto-oxidation can be a potential source of superoxide:

$\ce{Fe(2+)/Cu(+) + O2 ⇌ Fe(3+)/Cu(2+) + .O2-}$ (Source example: Eq (Vl))

The cited work by Blough and Zepp further notes that since the 1980’s:

...environmental scientists have acquired evidence for the near-ubiquitous occurrence in surface waters of not only H202 and 02" but also singlet dioxygen..., the hydroxyl radical (OH), and organic peroxyl radicals (R02), as well as other transient intermediates that are either immediate precursors or products of the ROS. It is now recognized that the production of this diverse array of species is driven primarily by abiotic photochemical reactions involving naturally occurring organic (and sometimes inorganic) chromophores (Zepp, 1991; Zika, 1987; Zafiriou et al., 1984). Because of its high concentration in surface waters (-250 nanoM), O2 dominates the photophysics and photochemistry of these materials. The resulting reaction sequences are readily interpreted within the well established concepts of direct and sensitized photo-oxidations, potentially coupled to thermal autoxidation processes.

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