When exposed to UV light, hydrogen peroxide decomposes into $\ce{H2O}$ and $\ce{O}$. Why does this happen and more importantly how? Is the energy from light absorbed by the bonds which are specific to UV frequency and thus leading to decomposition?
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Chemical reactions in which, a single substance splits up into two or more simpler substances are called decomposition reactions. These reactions are carried out by energy, supplied by different sources. The required energy can be supplied by heat (thermolysis), electricity (electrolysis), or light (photolysis).
Let’s talk about photolysis reactions (not photosynthesis):
Photolysis (also called photodissociation and photodecomposition) is a chemical reaction, in which a chemical (an inorganic or an organic) is broken down by photons and is the interaction of one or more photons with one target molecule. The photolysis reaction is not limited to the effects of visible light, but any photon with sufficient energy (higher than dissociation energy of the targeted bond) can cause the chemical transformation of the said (inorganic or organic) bond(s) of a chemical. Since the energy of a photon is inversely proportional to the wavelength, electromagnetic waves with the energy of visible light or higher, such as ultraviolet light, X-rays, and $\gamma$-rays, can also initiate photolysis reactions.
Like all other peroxides, hydrogen peroxide ($\ce{H2O2}$) also consists of a relatively weaker $\ce{O-O}$ bond, which is susceptible for light or heat. In the presence of light (the UV light from the sun catalyzes the reaction), $\ce{H2O2}$ spontaneously decomposes into water and oxygen.
The net equation for the reaction is: $$\ce{ 2H2O2 -> 2H2O + O2}$$
The step-wise reaction mechanism is suggested as follows (Ref.1):
$$\ce{ H2O2 + h\nu -> 2 HO^.}$$ $$\ce{ HO^. + H2O2 -> HO-O^. + H2O}$$ $$\ce{ HO-O^. + H2O2 -> 2 HO^. + H2O + O2}$$
Using isotope studies ($\ce{^{18}O}$ labelled $\ce{H2O2}$), early work has been confirmed that the $\ce{O2}$ formed is cleanly derived from $\ce{H2O2}$ (Ref.2).
Notes: The rate increases rapidly in the presence of catalysts such as $\ce{MnO2}$ and $\ce{KI}$ (Ref.2). The rate of decomposition is slow at room temperature, but it increases with temperature. It is believed to be due to thermal decomposition of $\ce{H2O2}$, which seemingly accelerates the photolysis (Ref.3).
References:
- J. P. Hunt, H. Taube, “The Photochemical Decomposition of Hydrogen Peroxide. Quantum Yields, Tracer and Fractionation Effects,” J. Am. Chem. Soc. 1952, 74(23), 5999–6002 (https://doi.org/10.1021/ja01143a052).
- A. E. Cahill, H. Taube, “The Use of Heavy Oxygen in the Study of Reactions of Hydrogen Peroxide,” J. Am. Chem. Soc. 1952, 74(9), 2312–2318 (https://doi.org/10.1021/ja01129a042).
- F. O. Rice, M. L. Kilpatrick, “The Photochemical Decomposition of Hydrogen Peroxide Solutions,” J. Phys. Chem. 1927, 31(10), 1507–1510 (https://doi.org/10.1021/j150280a004).
answered May 27, 2019 at 17:44
Mathew Mahindaratne
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$\begingroup$ If the O-O is susceptible to decomposition wouldn’t the products be H-O and H-O (basically hydroxides)? $\endgroup$user79603– user796032019-05-27 18:18:26 +00:00Commented May 27, 2019 at 18:18
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$\begingroup$ Yes, but those are radicals ($\ce{HO^.}$). That is the initiation reaction, which then contributes to propagation as described in next two equations. $\endgroup$Mathew Mahindaratne– Mathew Mahindaratne2019-05-27 18:21:23 +00:00Commented May 27, 2019 at 18:21
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$\begingroup$ Does infra red light react with hydrogen peroxide? You say visible light and above, so I was wondering if IR exposure is acceptable or at least makes for a slow reaction 🙄 $\endgroup$Velimir Tchatchevsky– Velimir Tchatchevsky2020-12-05 02:49:02 +00:00Commented Dec 5, 2020 at 2:49
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Not quite a universally accurate reaction chain, in my opinion, as the $\ce{HO2^.}$ radical has a pKa of 4.88 and above this value, we have the superoxide radical anion ($\ce{O2^{.-}}$) as the active species whose chemistry is distinct from the $\ce{HO2^.}$ radical:
$$\ce{HO2^. = H^+ + O_2^{.-} (pKa 4.88) }$$
Interestingly, as early as 1934 Haber and Weiss (Ref. 1) proposed that $\ce{HO2^.}$ is formed in the decomposition of hydrogen peroxide.
The subsequent reaction of the transient superoxide radical anion with hydrogen peroxide has been determined to also form the hydroxyl radical (Ref. 2). The higher pH version of the last reaction is, therefore, best represented as:
$$\ce{O2^.- + H2O2 ⟶ 2HO^∙ + OH- + O2}$$
And as $\ce{H^+ + OH^- = H2O}$, the net product formation is not altered.
However, alkaline $\ce{H2O2}$ is well known to be less stable than acidic hydrogen peroxide (Ref.3) even in absence of light exposure, which accelerates its decomposition (which can involve radical pathways) liberating oxygen!
References:
- Haber Fritz and Weiss Joseph, 1934, The catalytic decomposition of hydrogen peroxide by iron salts. Proc. R. Soc. Lond. A 147: 332–351 http://doi.org/10.1098/rspa.1934.0221
- TOSHIHIKO OZAWA, AKIRA HANAKI, Reactions of Superoxide with Water and with Hydrogen Peroxide, Chemical and Pharmaceutical Bulletin, 1981, Volume 29, Issue 4, Pages 926-928. https://doi.org/10.1248/cpb.29.926
- https://www.researchgate.net/figure/Effect-of-pH-on-the-decomposition-of-hydrogen-peroxide-H-2-O-2-0-800-mg-l_fig1_234110563, from Yazıcı, Ersin & Deveci, Haci. (2010). Factors Affecting Decomposition of Hydrogen Peroxide. https://doi.org/10.13140/RG.2.1.1530.0648.
