Zinc Dioxide Nanoparticulates: A Hydrogen Peroxide Source at

Jul 5, 2013 - ... Chemistry, The Hebrew University of Jerusalem, Jerusalem 91904, Israel ... Health Sciences, Hadassah Academic College, Jerusalem, Is...
0 downloads 0 Views 4MB Size
Download PDF
Article pubs.acs.org/est

Zinc Dioxide Nanoparticulates: A Hydrogen Peroxide Source at Moderate pH Yitzhak Wolanov,*,† Petr V. Prikhodchenko,‡ Alexander G. Medvedev,†,‡ Rami Pedahzur,§ and Ovadia Lev† †

The Casali Center of Applied Chemistry, The Institute of Chemistry, The Hebrew University of Jerusalem, Jerusalem 91904, Israel Kurnakov Institute of General and Inorganic Chemistry, Russian Academy of Sciences, Leninskii avenue 31, Moscow 119991, Russia § Department of Environmental Health Sciences, Hadassah Academic College, Jerusalem, Israel ‡

S Supporting Information *

ABSTRACT: Solid peroxides are a convenient source of hydrogen peroxide, which once released can be readily converted to active oxygen species or to dissolved dioxygen. A zinc peroxide nanodispersion was synthesized and characterized, and its solubility was determined as a function of pH and temperature. We show that zinc peroxide is much more stable in aqueous solutions compared to calcium and magnesium peroxides and that it retains its peroxide content down to pH 6. At low pH conditions H2O2 release is thermodynamically controlled and its dissolution product, Zn2+, is highly soluble, and thus, hydrogen peroxide release can be highly predictable. The Gibbs free energy of formation of zinc peroxide was found to be −242.0 ± 0.4 kJ/mol and the enthalpy of formation was −292.1 ± 0.7 kJ/mol, substantially higher than theoretically predicted before. The biocidal activity of zinc peroxide was determined by inactivation studies with Escherichia coli cultures, and the activity trend agrees well with the thermodynamic predictions.

■

INTRODUCTION Hydrogen peroxide is an environmentally friendly oxidant that can form more potent active oxygen species1 such as hydroxyl and superoxide radicals or disproportionate to give water and dissolved oxygen, depending on the presence of appropriate catalysts. Hydrogen peroxide is more stable than other sources of active oxygen and can be stored over long periods of time at low pH and in the presence of stabilizers.2,3 The successful use of H2O2 as a source of active oxygen for environmental decontamination4,5 and disinfection6,7 as well as an oxygen source for biodegradation of organic pollutants3,8 has been amply reported. However, the use of H2O2 in aqueous environments is limited by its rapid decomposition in natural aqueous systems. Means for slow and controlled release of the hydrogen peroxide to the environment are highly desirable, but all such passive delivery means utilize solid sources of hydrogen peroxide, which are more stable and can be stored for longer duration. Indeed alkali and alkaline earth metal peroxides are well-studied.9,10 By far, the most useful solid forms of peroxides for environmental applications are calcium and magnesium peroxides,11−14 due to their stability in high pH aqueous solutions. However, at acidic and near neutral pH, rapid dissolution due to reaction 1 is preferred. MO2 (s) + 2H+(aq) → M2 +(aq) + H 2O2 (aq)

release rate is also not controllable enough when a dissolved oxygen supply is desirable. The search for an alternative solid hydrogen peroxide source that would not involve organics and will also be stable at near-neutral pH could be guided by the available thermodynamic data on metal peroxide formation. Unfortunately, the thermodynamic data on most peroxides are rather scarce15,16 or based on old thermodynamic correlations.17 The only thermodynamic data that we could find on the peroxides of d10 transition metal elements, which are the closest analogs to the alkaline earth elements, are based on a general empirical correlation with a large uncertainty level of approximately 60 kJ/mol.17 The available data for the alkaline earth and d10 transition metal peroxides and the calculated dissolution constants based on eq 1 are summarized in the Supporting Information (SI) in Table S1. The high toxicity of beryllium, cadmium, and mercury leaves zinc peroxide as the only viable d10 alternative to the widely used calcium and magnesium peroxides. ZnO2 is much less studied than MgO2 and CaO2. It is increasingly being used as a precursor for zinc oxide,18 and in a 60-year-old work, it has also been suggested for the treatment of surgical infections.19 Zinc is widely used in drinking water pipes, and it is an essential trace Received: Revised: Accepted: Published:

(1)

The fast disproportionation of H2O2 at high pH is a major disadvantage for advanced oxidation processes, and the oxygen © 2013 American Chemical Society

8769

May 8, 2013 June 30, 2013 July 5, 2013 July 5, 2013 dx.doi.org/10.1021/es4020629 | Environ. Sci. Technol. 2013, 47, 8769−8774

Environmental Science & Technology

Article

Determination of the Zinc Ion Concentration. The Zn2+ concentration in the solution was determined by an inductively coupled plasma (ICP) optical emission spectrometer (Varian 720-ES) with 10% inaccuracy. Instrumentation. The analytical instrumentation is described in the SI. Determination of the Biocidal Activity. The biocidal activity was quantified by the minimal inhibitory concentration (MIC) test according to a standard protocol.26 The biocidal activity was measured at pH 7. A solution of equilibrated ZnO2 at the specified pH was filtered and passed through Amberlyst 15 cation exchange column in sodium form (Fisher Chemical). Then, the pH was corrected using concentrated NaOH. The test organism was Escherichia coli MG1655. Bacteria were maintained on MacConkey agar plates from which overnight (ON) cultures were prepared in cation-adjusted Muller Hinton broth (BD) at 37 °C with shaking. Prior to MIC tests, ON cultures were refreshed by a 1:100 dilution in fresh Muller Hinton (MH) broth supplemented with 0.1 M HEPES and grown (37 °C with shaking) for 90 min to an OD590nm of 0.1. The tests were performed in transparent 96-well microtiter plates as follows: All of the wells, except the ones in row A, contained 100 μL of ×1 Muller Hinton broth supplemented with 0.1 M HEPES. The wells in row A contained 20 μL of ×10 Muller Hinton broth supplemented with HEPES and 180 μL of the different tested solutions as to achieve a ×1 strength of the MH and HEPES. Then, double dilutions of the active materials were obtained by transferring 100 μL from the first wells (A) in each row to the next well (B) and so on down the rows. The last row (H) was left without addition of active ingredient as a control. After plate construction, 20 μL of the refreshed bacteria was uniformly added to all of the wells, the plate was then incubated at 37 °C overnight, and the turbidity was assessed on the following day. The first clear well, representing growth inhibition, was determined as the MIC. HEPES and HEPES with dissolved ZnO solutions were used as control. All tests were performed in duplicates.

element for humans. The US EPA regulations, the World Health Organization guidelines, and the EU directive for drinking water do not impose health-related maximum contaminant level for zinc and zinc salts. We found only one article20 about the use of ZnO2 for environmental applications. The article reports on ZnO2 additive for antifouling paint for marine usage. This article showed inferior biofouling capabilities of zinc dioxide compared to copper oxide but did not discuss the findings in view of the underlying thermodynamics. In this study we synthesized zinc peroxide nanoparticles by a fast and simple protocol. We compared the H2O2 release of ZnO2 to that of MgO2 and CaO2 as a function of pH and showed that ZnO2 retains its peroxide content over a wide pH range and starts to decompose only in the mild acidic region, far after MgO2 and CaO2 lose their peroxide content. Furthermore, ZnO2 decomposition is shifted to lower pH at higher temperature. From the pH and temperature dependence, the heat of dissolution and the heat of formation of ZnO2 were calculated for the first time. The only available thermodynamic data on ZnO2 is an estimated value of the enthalpy of formation, −347 ± 63 kJ/mol at 25 °C.17 Finally, we demonstrate the bioactivity of the released H2O2 by testing its biocidal activity as a function of pH.

■

EXPERIMENTAL SECTION

Synthesis of Zinc Peroxide Nanoparticles. Different synthesis protocols of zinc peroxide nanoparticles are described in the literature.18,21−24 Our own proposed protocol is based on dissolution of zinc acetate at high pH using ammonia, which acts as a base and as a chelating agent. Zinc chelation by ammonia was previously reported by Singh et al.,23 but their protocol involved dissolution of the zinc precursors at low pH, whereas in our case the dissolution took place at high pH and an excess of hydrogen peroxide was then added to precipitate the zinc peroxide. The detailed synthesis protocol is provided in the SI. Magnesium Peroxide, MgO2. Since the magnesium peroxide content in the commercially available product (Sigma, Israel) was rather low, 24−28%, we decided to synthesize the magnesium peroxide in our laboratory. The detailed synthesis protocol is depicted in the SI. Solubility Tests of Solid Peroxides. Metal peroxide powder (0.4 mmol) was mixed with 0.01 M HEPES solution. The final pH was set (approximately) by addition of 1 M HCl, and distilled water was added to a final volume of 4 mL. The beakers containing the ZnO2 were left to equilibrate for at least 10 days at the specified temperatures, and the samples of calcium and magnesium peroxide were left to equilibrate for 2− 3 days only, since dissolution took place at higher pH where the hydrogen peroxide is not stable and gradually deteriorated. In all cases, the samples were then left until the concentration of hydrogen peroxide and the pH stabilized. In all cases, we verified that some solid peroxide remained after equilibration. Hydrogen peroxide, zinc, and the pH were measured after equilibration. Determination of the Peroxide Content. The hydrogen peroxide concentration of the aqueous solutions and the amount of active oxygen in the solids were determined by titration with freshly prepared 0.01N permanganate solution (permanganatometry). 25 The accuracy of the method is within 1−2% depending on the titrant volume.

■

RESULTS ZnO2 Characterization. The synthesized nanoparticles were analyzed by permanganatometry and found to contain about 12−13% active oxygen (i.e., 24−26% peroxide), which corresponds to 73−80% of the theoretical content of pure zinc peroxide. The high-resolution TEM image of the ZnO2 nanoparticles is shown in Figure 1. The image demonstrates uniform and spherically shaped nanoparticles of 30−40 nm diameter. The TEM image reveals that the observed particles are nanoaggregates of smaller crystals. Dynamic light scattering corroborated the observation that the material is nanodispersed but the average particle size was 200 nm, a result of some aggregation in the solution. Nanodispersions are favorable for our studies, since there is minimal diffusion barrier for the release of the peroxide, and nanoparticles can be more easily dispersed and applied in practical applications. However, this could be a drawback when microencapsulation is desirable to achieve slow release by mass transport limitations. X-ray diffraction of the zinc dioxide (Figure 2A) shows the diffraction pattern of cubic zinc peroxide. Although no amorphous phase is observable in the diffractogram, we believe that the 20% impurities that were calculated by permanganatometry can be attributed to small amorphous regions of zinc 8770

dx.doi.org/10.1021/es4020629 | Environ. Sci. Technol. 2013, 47, 8769−8774

Environmental Science & Technology

Article

Figure 3. STEM images of ZnO2 nanoparticles (before heat treatment) (A) and after heat treatment at 300 °C (B).

is absent in Figure 4, showing that there is no imbibed or hydrogen-bonded hydrogen peroxide. The peroxide band disappeared after heat treatment at 300 °C (black curve in Figure 4). Figure 1. High-resolution TEM image of ZnO2 nanoparticles.

hydroxide. The average grain size of the zinc peroxide was estimated by the Scherrer equation to be 4.5 nm, corroborating the TEM image that show that the nanoparticles comprise aggregates of smaller crystals. Thermal analysis (Figure 2B) revealed that the material is stable up to 160 °C, and at higher temperature the DSC curve showed exothermic decomposition (with a peak at 210 °C) accompanied by some 15% weight decrease (between T = 150 and 230 °C) corresponding to oxygen release. The product was zincite (ZnO) with hexagonal structure and approximately the same 4.5 nm size as determined by Scherrer analysis. The fact that the material maintained the same shape and size after hydrogen peroxide decomposition can be seen in parts A and B of Figure 3, which illustrate the STEM imaging of ZnO2 and the resulting ZnO nanoparticles. The Raman spectra of the ZnO2 nanoparticles exhibit the O−O stretching vibration band of the peroxide at 840 cm−1, which agrees well with previous reports.27 Free hydrogen peroxide has a strong Raman vibration band at 880 cm−1, which

Figure 4. Raman spectra of zinc peroxide before (lower curve) and after heat treatment at 300 °C).

Effect of pH on the Dissolution of ZnO2. Comparative Studies of Zinc, Calcium, and Magnesium Dioxide Dissolution. As a preliminary stage, we have conducted a

Figure 2. (A) XRD patterns before (ZnO2) and after (ZnO) heat treatment at 300 °C. (B) DSC and TGA measurements of ZnO2. 8771

dx.doi.org/10.1021/es4020629 | Environ. Sci. Technol. 2013, 47, 8769−8774

Environmental Science & Technology

Article

comparative study of the dissolution of calcium, magnesium, and zinc peroxide solids as a function of pH. The hydrogen peroxide percentage in the solution as a function of pH for ZnO2 compared to MgO2 and CaO2 at 22 °C is shown in Figure 5. The figure shows that the hydrogen peroxide release

Figure 6. H2O2 concentration in solutions of ZnO2 as a function of pH at different temperatures.

log(Keq) = log{[Zn 2 +][H 2O2 ]} + 2pH

A plot of −log{[Zn ][H2O2]} vs pH should give a straight line with a slope of 2. We measured the equilibrium concentration of H+, hydrogen peroxide, and zinc in the solutions and calculated the activity coefficients of Zn2+ using the extended Debye−Huckel approximation (with α = 6 Å).33 The activity coefficients of water and hydrogen peroxide were taken as 1. Figure 7 depicts plots of −log{[Zn2+][H2O2]} vs the pH for the

Figure 5. H2O2 concentration in solution (in wt %) in the presence of ZnO2 (□), MgO2 (⊕), and CaO2 (▲) as a function of pH.

from the zinc dioxide (at T = 22 °C) occurs at a much lower pH compared to the crystalline alkaline earth peroxides and that the release takes place at a narrow pH range around pH 4− 6. Dissolution of 50% of the peroxide content of calcium and magnesium peroxide was around pH 9, whereas zinc peroxide dissolved 50% of its peroxide content at pH 5. The dissolution of the calcium and magnesium peroxides was not surprising and followed previous reports.28,29 The reported stability of calcium and magnesium peroxides in some field studies30,31 can probably be attributed to metal hydroxide deposition that creates a growing mass transport barrier,28,32 which slows down further release of the hydrogen peroxide and also helps retain the high pH at the vicinity of the peroxide salt. Although the formed hydroxide layer is thermodynamically soluble under these conditions, further generation of hydroxide due a continues dissolution of the peroxides core stabilizes to some extent the outer hydroxide protecting shell. The transport barrier formed by the hydroxide deposition in nanoparticles is too thin to decelerate hydrogen peroxide release appreciably, and thus, the technology is incompatible with the recent nanotechnology trends. The observed high stability of zinc peroxide down to moderate pH levels was never mentioned, as far as we know, in the scientific literature. It should be reemphasized that zinc peroxide dissolution was thermodynamically controlled, and the hydrogen peroxide level was determined after some 10 days of equilibration. Effect of Temperature on ZnO2 Dissolution. The effect of temperature on the solubility of ZnO2 in the pH range of 4.3−6.3 is shown in Figure 6. The graph shows lower solubility of ZnO2 at higher temperature. The dissociation of ZnO2 in a mildly acidic solution is given by eq 1a: ZnO2 (s) + 2H+(aq) → Zn 2 +(aq) + H 2O2 (aq)

Figure 7. Plots of −log{[Zn2+][H2O2]} vs pH at different temperatures and linear fits with a forced slope of 2. The results represent the average of three measurements. The error bars were always smaller than the symbols.

five different temperatures that were denoted in Figure 5. The hydrogen peroxide levels were determined for each data point, and the dissolved zinc was determined in six measurements for each temperature. Note that the amount remaining in the solid form and the possible formation of zinc hydroxide (which anyway should not take place at the relevant pH range) do not affect the validity of eq 3, and in fact, the same applies for the possible dismutation of hydrogen peroxide during the prolonged tests. By fitting linear regressions with a slope of 2, we could extract the equilibrium constants (Keq) of this reaction for the five different temperatures. The Keq at 22 °C is 8.9 × 106. The Gibbs free energy of the reaction was calculated from Keq and found to be −39.3 kJ/mol. The standard enthalpy of the

(1a)

The equilibrium constant of eq 1a is given by Keq = [Zn 2 +][H 2O2 ]/[H+]2

(3)

2+

(2)

Expression of eq 2 in logarithmic form gives a linear equation: 8772

dx.doi.org/10.1021/es4020629 | Environ. Sci. Technol. 2013, 47, 8769−8774

Environmental Science & Technology

Article

reaction was calculated by plotting the natural logarithm of the five equilibrium constants as a function of 1/T as shown in Figure 8.

Figure 8. Weighted linear regression of natural logarithm of the five equilibrium constants as a function of 1/T. Figure 9. Plot of the maximal number of dilutions that still show antibacterial activity (○) and the original concentrations of the samples (■) as a function of pH (A) and a photograph of the microtiter plates after incubation (B).

From the weighted linear fit using the inverse of the variance as weights,34 we could extract the enthalpy according to the Van’t Hoff equation: d(ln K ) 1 dT

=−

ΔH ° R

■

ASSOCIATED CONTENT

S Supporting Information *

(4)

Additional information as indicated in the text. This material is available free of charge via the Internet at http://pubs.acs.org.

We calculated ΔH° = −51.9 ± 0.5 kJ/mol from eq 4. The standard deviations of the different Keq(T) are denoted by bars in Figure 8. According to eq 1 and the Gibbs free energy and enthalpy of formation of Zn2+(aq) and H2O2(aq),35 we could calculate the enthalpy of formation of ZnO2 to be −292.1 ± 0.7 kJ/mol and the Gibbs free energy of −242.0 ± 0.4 kJ/mol. As noted above, the Gibbs free energy of formation of zinc peroxide was never reported in the literature. The value of the enthalpy of formation of zinc dioxide is very different from the value calculated in ref 17, but it falls well within the uncertainty range of the calculations indicated in this reference. Biocidal Activity: MIC Experiment by Dilution Method. The dilution method is the most commonly used technique to determine the minimal inhibitory concentration (MIC) of antimicrobial agents. MIC represents the concentration of antimicrobial at which there is complete inhibition of organism growth. E. coli solution was added to the 2-fold serial dilutions of hydrogen peroxide samples as described in the Experimental Section and the visual results are shown in Figure 9. The yellowish turbidity indicates bacterial growth. The pictures agree well with the hydrogen peroxide release curve at 22 °C (Figure 9A) and so does the good fit of the maximal number of dilutions that still show antibacterial activity in the MIC experiment as a function of pH (Figure 9B). The nanoparticulate solid zinc peroxide that was studied is stable under natural pH conditions and becomes activated in a mildly acidic environment. Thus ZnO2 can be a used for pHcontrolled release of hydrogen peroxide for various environmental applications, including advanced oxidation processes or for remediation in cases where high levels of dissolved zinc are tolerated or can be precipitated after the treatment at higher pH.

■

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]; phone: +972-2-6584191; fax: +972-2-6586155. Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS This research is supported by the Singapore National Research Foundation under CREATE programme: Nanomaterials for Energy and Water Management and by the Russia−Israel Binational Program in Nanotechnology. We thank The Harwey M. Kreuger Family Center for Nanoscience and Nanotechnology of the Hebrew University of Jerusalem. We thank the Russian Foundation for Basic Research (grant 11-0300551), the Council on Grants of the President of the Russian Federation (NSh-1670.2012.3), the Ministry of Education and Science of the Russian Federation (SC-8437), the Target Programs for Basic Research of the Presidium of the Russian Academy of Sciences, and the Division of Chemistry and Materials Science of the Russian Academy of Sciences.

■

REFERENCES

(1) Lin, S. S.; Gurol, M. D. Catalytic decomposition of hydrogen peroxide on Iron Oxide: Kinetics, mechanism, and implications. Environ. Sci. Technol. 1998, 32, 1417−1423. (2) Schumb, W. C.; Satterfield, C. H.; Wentworth, R. L. Hydrogen Peroxide; Reinhold Publishing Corp.: New York, 1955. (3) Pardieck, D. L.; Bouwer, E. J.; Stone, A. T. Hydrogen peroxide use to increase oxidant capacity for in situ bioremediation of

8773

dx.doi.org/10.1021/es4020629 | Environ. Sci. Technol. 2013, 47, 8769−8774

Environmental Science & Technology

Article

contaminated soils and aquifers: A review. J. Contamin. Hydrol. 1992, 9, 221−242. (4) Pignatello, J. J. Dark and photoassisted iron (3+)-catalyzed degradation of chlorophenoxy herbicides by hydrogen peroxide. Environ. Sci. Technol. 1992, 26, 944−951. (5) Zepp, R. G.; Faust, B. C.; Hoigne, J. Hydroxyl radical formation in aqueous reactions (pH 3−8) of iron(II) with hydrogen peroxide: The photo-Fenton reaction. Environ. Sci. Technol. 1992, 26, 313−319. (6) Armon, R.; Laot, N.; Lev, O.; Shuval, H.; Fattal, B. Controlling biofilm formation by hydrogen peroxide and silver combined disinfectant. Water Sci. Tech. 2000, 42, 187−192. (7) Pedahzur, R.; Katzenelson, D.; Bamea, N.; Lev, O.; Shuval, H.; Fattal, B.; Ulitzur, S. The efficacy of long-lasting residual drinking water disinfectants based on hydrogen peroxide and silver. Water Sci. Technol. 2000, 42, 293−298. (8) Piotrowski, M. R. Bioremediation: Testing the waters. Civ. Eng. 1989, 59, 51−53. (9) Woad, R. H.; D’Orazio, L. A. Thermodynamics of the higher oxides. I. The heats of formation and lattice energies of the superoxides of potassium, rubidium, and cesium. J . Phys. Chem. 1965, 69, 2550−2557. (10) Woad, R. H.; D’Orazio, L. A. Thermodynamics of the higher oxides. II. Lattice energies of the alkali and alkaline earth peroxides and the double electron affinity of the oxygen molecule. J. Phys. Chem. 1965, 69, 2558−2561. (11) Kao, C. M.; Chen, S. C.; Su, M. C. Laboratory column studies for evaluating barrier system for providing oxygen and substrate for TCE biodegradation. Chemosphere. 2001, 44, 925−934. (12) Kao, C. M.; Chen, S. C.; Wang, J. Y.; Chen, Y. L.; Lee, S. Z. Remediation of PCE-contaminated aquifer by an in situ two-layer biobarrier: Laboratory batch and column studies. Water Res. 2003, 37, 27−38. (13) Liu, S. J.; Jiang, B.; Huang, G. Q.; Li, X. G. Laboratory column study for remediation of MTBE-contaminated groundwater using a biological two-layer permeable barrier. Water Res. 2006, 40, 3401− 3408. (14) Cassidy, D. P.; Irvine, R. L. Use of calcium peroxide to provide oxygen for contaminant biodegradation in a saturated soil. J. Hazard. Mater. 1999, 69, 25−39. (15) Brewer, L. The thermodynamic properties of the oxides and their vaporization process. Chem. Rev. 1953, 52, 1−75. (16) Rossini, F. D.; Wagman, D. D.; Evans, W. H.; Levine, S.; Jaffe, I. Selected Values of Chemical Thermodynamic Properties; National Bureau of Standards Circular 500; U.S. Government Printing Office: Washington, DC, 1952. (17) Wilcox, D. E.; Bromley, L. A. Computer estimation of heat and free energy of formation for simple inorganic compounds. Ind. Eng. Chem. 1963, 55, 32−39. (18) Uekawa, N.; Kajiwara, J.; Mochizuki, N.; Kakegawa, K.; Sasaki, Y. Synthesis of ZnO nanoparticles by decomposition of zinc peroxide. Chem. Lett. 2001, 30, 606−607. (19) Meleney, F. L. Present role of zinc peroxide in treatment of surgical infections. J. Am. Med. Assoc. 1952, 149, 1450−1453. (20) Olsen, S. M.; Pedersen, L. T.; Hermann, M. H.; Kiil, S.; DamJohansen, K. Inorganic precursor peroxides for antifouling coatings. JCT Res.. 2009, 6, 187−199. (21) Sun, M.; Hao, W.; Wang, C.; Wang, T. A simple and green approach for preparation of ZnO2 and ZnO under sunlight irradiation. Chem. Phys. Lett. 2007, 443, 342−346. (22) Rosenthal-Toib, L.; Zohar, K.; Alagem, M.; Tsur, Y. Synthesis of stabilized nanoparticles of zinc peroxide. Chem. Eng. J. 2008, 136, 425−429. (23) Singh, N.; Mittal, S.; Sood, K. N.; Rashmi; Gupta, P. K. Controlling the flow of nascent oxygen using hydrogen peroxide results in controlling the synthesis of ZnO/ZnO2. Chalcogenide Lett. 2010, 7, 275−281. (24) Chen, W.; Lu, Y. H.; Wang, M.; Kroner, L.; Paul, H.; Fecht, H. J.; Bednarcik, J.; Stahl, K.; Zhang, Z. L.; Wiedwald, U.; Kaiser, U.; Ziemann, P.; Kikegawa, T.; Wu, C. D.; Jiang, J. Z. Synthesis, thermal

stability and properties of ZnO2 nanoparticles. J. Phys. Chem. C. 2009, 113, 1320−324. (25) Schumb, W. C.; Satterfield, C. H.; Wentworth, R. L. Hydrogen Peroxide; Reinhold Publishing Corp.: New York, 1955; 553 pp. (26) Andrews, J. M. Determination of minimum inhibitory concentrations. J. Antimicrob. Chemother. 2001, 48, 5−16. (27) Bai, H.; Liu, X. Green hydrothermal synthesis and photoluminescence property of ZnO2 nanoparticles. Mater. Lett. 2010, 64, 341−343. (28) Waite, A. J.; Bonner, J. S.; Austenite, R. Kinetics and stoichiometry of oxygen release from solid peroxides. Environ. Eng. Sci. 1999, 16, 187−199. (29) Northup, A.; Cassidy, D. Calcium peroxide (CaO2) for use in modified Fenton chemistry. J. Hazard. Mater. 2008, 152, 1164−1170. (30) Nykänen, A.; Nykänena, Anne; Kontio, H.; Klutas, O.; Penttinen, O. P.; Kostia, O. P.; Mikola, J.; Romantschuk, M. Increasing lake water and sediment oxygen levels using slow release peroxide. Sci. Total Environ. 2012, 429, 317−324. (31) Schmidtke, T.; White, D.; Woolard, C. Oxygen release kinetics from solid phase oxygen in Arctic Alaska. J. Hazard. Mater. 1999, 64, 157−165. (32) Elprince, A. M.; Mohamed, W. H. Catalytic decomposition kinetics of aqueous hydrogen peroxide and solid magnesium peroxide by birnessite. Soil. Sci. Soc. Am. J. 1992, 56, 1784−1788. (33) Benjamin, M. M. Water Chemistry; Waveland Press, Inc.: Long Grove, IL, 2010. (34) Miller, J. C., Miller, J. N. Statistics for Analytical Chemistry; Ellis Horwood PTR Prentice Hall: New York, 1993. (35) Bard, A. J.; Parsons. R.; Jordan, J. (Eds.) Standard Potentials in Aqueous Solution; Marcel Dekker: New York, 1985.

8774

dx.doi.org/10.1021/es4020629 | Environ. Sci. Technol. 2013, 47, 8769−8774