- Open Access
Quantifying hydrogen peroxide in iron-containing solutions using leuco crystal violet
© American Institute of Physics 2005
- Received: 24 February 2005
- Accepted: 26 April 2005
- Published: 14 June 2005
Hydrogen peroxide is present in many natural waters and wastewaters. In the presence of Fe(II), this species decomposes to form hydroxyl radicals, that are extremely reactive. Hence, in the presence of Fe(II), hydrogen peroxide is difficult to detect because of its short lifetime. Here, we show an expanded use of a hydrogen peroxide quantification technique using leuco crystal violet (LCV) for solutions of varying pH and iron concentration. In the presence of the biocatalyst peroxidase, LCV is oxidized by hydrogen peroxide, forming a colored crystal violet ion (CV+), which is stable for days. The LCV method uses standard equipment and allows for detection at the low microM concentration level. Results show strong pH dependence with maximum LCV oxidation at pH 4.23. By chelating dissolved Fe(II) with EDTA, hydrogen peroxide can be stabilized for analysis. Results are presented for hydrogen peroxide quantification in pyrite–water slurries. Pyrite–water slurries show surface area dependent generation of hydrogen peroxide only in the presence of EDTA, which chelates dissolved Fe(II). Given the stability of CV+, this method is particularly useful for field work that involves the detection of hydrogen peroxide.
- Electron Paramagnetic Resonance
- Ferrous Iron
- H2O2 Concentration
- Ferrous Ammonium Sulfate
Hydrogen peroxide is an important reactant in natural aquatic systems, environmental remediation technologies, and in biological systems. Reactive oxygen species (ROS), which include hydrogen peroxide (H2O2) and hydroxyl radicals (·OH) occur in rain[1, 2] and surface waters. [3, 4] ROS play an important role in natural processes in aquatic systems, including, radiolysis,[5, 6] pyrite oxidation, [7–9] and photochemical oxidation.[10, 11] Hydrogen peroxide also is intentionally added to wastewaters to promote in situ oxidation processes[12, 13] by leveraging the extreme reactivity of radicals toward organic pollutants. The generation of ·OH, via the decomposition of H2O2, degrades contaminants when methods involving microbiological degradation are ineffective. In organisms, including humans, ROS are produced during metabolic and immune system function. When ROS concentrations are above normal for prolonged periods of time, however, their presence can lead to oxidative stress. Oxidative stress is now recognized to be an important factor in the development or enhancement of many diseases.
Among ROS, H2O2 is relatively stable, in the presence of ferrous iron, H2O2 forms ·OH via the Fenton reaction
Fe(II) + H2O2 → ·OH + OH- + Fe(III) (1)
·OH is far more reactive than H2O2 and its reaction with aqueous species is diffusion limited. Therefore, H2O2, in the presence of Fe(II), represents great potential for reactivity.
In natural systems, ·OH is a transient species and its steady-state concentration may only be in the nM range. The inherent reactivity of ·OH precludes its detection via a direct measurement, and hence, methods generally rely on the detection of a stable reaction product resulting from the reaction of ·OH and a target molecule. One strategy is to add a reactant as a target species that will oxidize in the presence of ·OH and form a product that can be analyzed by UV–Vis spectroscopy or fluorescence. Verifying ·OH involvement in oxidation of the target species is difficult. Addition of competing scavengers that react with ·OH may inhibit oxidation of the target molecule; however, there is little conformity when using different scavengers and target molecules. Another strategy that is often employed is to "trap" the unpaired electron of the radical into a compound that is less reactive. By using electron paramagnetic resonance (EPR), the EPR intensity of the spin-trapped molecule can be directly related to the concentration of ·OH in the solution. EPR spin-trapping is a very sensitive technique and has been widely used in chemistry, environmental sciences, and bio-(medical) sciences; however, it does require a significant capital investment and spin-trapping cannot be conducted in the field. The spin-trapping technique has also been used for ·OH detection from aqueous mineral slurries, however reaction between Fe(III) and the spin-trap complicates interpretation of the data.
Where detection of ·OH is impracticable, identifying H2O2 may elucidate a reaction mechanism involving ·OH formation. In iron-containing systems, both H2O2 and ·OH are short-lived. However, chelation of iron [i.e., Fe(II)] avoids its interaction and decomposition of H2O2. Hence, we argue that hydrogen peroxide detected in iron-containing systems with an iron-chelator added is a proxy for the capacity of the system to generate ·OH in the absence of an iron chelator. This strategy may prove useful as an alternative method for radical detection in natural systems.
Several methods are available for quantifying H2O2. UV–Vis absorbance (240 nm, ε = 43.6 M-1 cm-1) is convenient for H2O2 solutions containing no other UV absorbing chromophores. Another spectrophotometric method uses copper(II) and 2,9-dimethyl-l, 10-phenanthroline (DMP) for μM H2O2 detection in wastewater. The colored complex does not change in the presence of humic acid; however, metal chelators affect the copper reactivity, complicating its use when iron chelation is necessary. For higher sensitivity detection, fluorometric techniques are available. The scopoletin/horseradish peroxidase method allows for detection of low nM concentrations, but the need for standard additions and quantification via a decrease in fluorescence prove to be time consuming. Another fluorometric technique involves the oxidation of non-fluorescent 2',7'-dichloro-fluorescin (DCFH) to fluorescent 2',7'-dichlorofluorescein (DCF) in the presence of H2O2 and peroxidase. This technique is often used to detect the formation of H2O2 in cells. The technique makes use of the non-fluorescent (DCFH-diacetate) crossing of cell membranes, which is enzymatically deacetylated to non-fluorescent DCFH. Cellular production of H2O2 produces the fluorescent DCF. The deacetylation process can be achieved chemically, but H2O2 may be produced in the process,  resulting in background levels of H2O2. Another disadvantage for using the DCFH technique is the necessity for daily preparation of the DCFH reagent.
Here, we present results on an improved method for H2O2 detection in the μM to several hundred nM range, which has several advantages over preexisting methods. The leuco crystal violet (LCV) method involves oxidation of 4,4'4"-methylidynetris (N,N-dimethylaniline) (C25H31N3) (LCV) in the presence of H2O2 and horseradish peroxidase (HRP) to form the crystal violet ion, CV+, which absorbs at 590 nm. CV+ remains stable for several days, which makes it possible to treat samples upon collection and perform the analysis at a later time. This is of value in field studies or shipboard analysis.  Daily preparation of new reagents is not required and iron chelators do not affect the analysis.
In this report, the LCV method was employed for analysis of H2O2 in pyrite/aqueous slurries. Recently, pyrite (FeS2), the most abundant metal sulfide mineral on Earth, has been shown to produce H2O2 in aqueous solutions. Several adjustments were made to the original LCV method so that it is now possible to determine H2O2 concentration in aqueous systems that contain iron-bearing minerals. Here, we determined effects of iron, pH, and addition of EDTA. To verify that our technique was specific to the presence of H2O2, and that other species were not responsible for the oxidation of LCV in the presence of HRP, we carried out experiments in the presence of catalase. This enzyme selectively decomposes H2O2, and its addition to the solution prior to the addition of LCV and HRP eliminates the production of the CV+ species.
LCV method details.a
Volume added (μl)
136.07 g + 0.5 L H2O
pH adjusted to 4.2 with H3PO4
31 mg LCV + 30 ml H2O + 19.2 ml of 0.25 N HCl
4.0 mg HRP + 50 ml H2O + 92 μl of 1 M sodium azide
0.08 mg/ ml (14.4 units/ml)
1 μg (0.18 units)
The reagent preparations and volumes used are shown in Table 1. It is useful to present our conditions in the context of prior LCV studies that were carried out by Zhang et al.  Compared to the concentrations used by Zhang et al. for pH buffer, LCV, and HRP, we used 1.25, 4, and 10 times those previously used, respectively, to increase (1) the buffer concentration, (2) the higher range of detection (more LCV), and (3) the reaction rate (more HRP). By increasing the concentration of our reagents, we were able to limit the volume of reagent addition so that a maximum volume would be available for the sample.
Addition of catalase is important for verifying the presence of H2O2, since other ROS or other reactions may also oxidize LCV. When catalase is added, no H2O2 is detected, which limits the possibility of false positives.
For H2O2 detection in natural waters, the LCV technique may be one of the more suitable methods. It has already been shown that the method can be used for seawater analyses  and here we show that it can be used for a pH range of around 3.5 – 6.0 and in the presence of EDTA. Compared to techniques involving fluorescence, where reagents are prepared on the day of analyses, our solutions of LCV and HRP remained stable for months at 4°C in opaque centrifuge vials. The stability of CV+ upon reaction of H2O2 and the LCV reagents makes it possible to quantify H2O2 several days after sampling. This can be exploited in field studies. For example, the LCV technique could be used to study the spatial and temporal distribution of H2O2 in hot spring waters in Yellowstone National Park. It has been shown that steady-state levels of photochemically produced H2O2 in the surface geothermal waters at Yellowstone National Park reach 200 – 600 nM by late afternoon and decrease to less than 50 nM during the night. With the LCV technique a large number of water samples can be collected and prepared for later analysis. Temperature is expected to affect the rate of the HRP-mediated reaction and it could possibly also affect the stability of the LCV or CV+. Probably the best strategy working with hydrothermal waters is to rapidly cool the sample down to a temperature between 20 and 30°C before adding the reagents. Cooling the samples to much lower temperature may impede the enzyme reaction. Further experimental work would be needed to resolve this temperature dependence. The LCV technique may also prove to be useful to evaluate the performance of environmental remediation projects involving the injection of H2O2 into contaminated waters. Many groundwaters contain dissolved iron, which could make it difficult to determine the residual H2O2 concentration. With the LCV techniques samples can be treated with ETDA and preserved for latter analysis.
This study demonstrates a reliable and efficient method for quantifying H2O2 from iron-containing mineral slurries and waste by use of separate calibration curves to account for pH and iron concentrations. The stability of the colored CV+ makes this method suitable for the field or when immediate access to a spectrophotometer is not possible. Relative to the concentration of H2O2 consumed, the reported high molar absorptivity of 75 000 M-1 cm-1 for CV+ makes it possible to determine H2O2 at sub-μM concentration levels. In iron-containing systems at low pH, hydrogen peroxide reacts to form ·OH. Under those conditions, the presence of H2O2 as an intermediate to ·OH formation would be extremely difficult to detect. The LCV technique as outlined provides a relatively simple method to demonstrate the involvement of H2O2 under those conditions.
This work was funded by the Department of Energy through grants to D.R.S. and M.A.S., Basic Energy Sciences Grant Nos. DEFG029ER14644 and DEFG0296ER14633, respectively. The Center for Environmental Molecular Science (NSF CHE 0221934) facilitated the contribution of A.P. to this project. C.C. would like to acknowledge support from a National Defence Science and Engineering Fellowship.
- Willey JD, Kieber RJ, Lancaster RD, Atmos J: Chem. 1996, 25: 149.Google Scholar
- Faust BC, Allen JM: Environ Sci Technol. 1993, 27: 1221-10.1021/es00043a024.View ArticleGoogle Scholar
- Hakkinen PJ, Anesio AM, Graneli W: Can J Fish Aquat Sci. 2004, 61: 1520-10.1139/f04-098.View ArticleGoogle Scholar
- Gerringa LJA, Rijkenberg MJA, Timmermans KR, Buma AGJ: Netherlands J Sea Res. 2004, 51: 3.View ArticleGoogle Scholar
- Amme M, et al: Environ Sci Technol. 2005, 39: 221.View ArticleGoogle Scholar
- Clarens F, et al: Environ Sci Technol. 2004, 38: 6656-10.1021/es0492891.View ArticleGoogle Scholar
- Ahlberg E, Broo AE: Int J Min Process. 1996, 47: 49-10.1016/0301-7516(95)00100-X.View ArticleGoogle Scholar
- Allen JM, Lucas S, Allen SK: Envir Toxicol Chem. 1996, 15: 107-10.1897/1551-5028(1996)015<0107:FOHROI>2.3.CO;2.Google Scholar
- Borda MJ, Elsetinow AR, Strongin DR, Schoonen MA: Croat Chem Acta. 2003, 67: 935.Google Scholar
- Stumm W, Morgan JJ: Aquatic Chemistry: Chemical Equilibria and Rates in Natural Waters. 1995, Wiley-Interscience, New York, 1022-3Google Scholar
- Vaughan P, Blough N: Environ Sci Technol. 1998, 32: 2947-10.1021/es9710417.View ArticleGoogle Scholar
- Gogate PR, Pandit AB: Adv Environ Res. 2004, 8: 501-10.1016/S1093-0191(03)00032-7.View ArticleGoogle Scholar
- Arienzo M: Chemosphere. 1999, 39: 1629-10.1016/S0045-6535(99)00061-2.View ArticleGoogle Scholar
- Dröge W: Physiol Rev. 2002, 82: 47.View ArticleGoogle Scholar
- Nordberg J, Arnér ESJ: Free Radic Biol Med. 2001, 31: 1287-10.1016/S0891-5849(01)00724-9.View ArticleGoogle Scholar
- Finkel T, Holbrook NJ: 2000, Nature (London), 408: 239.Google Scholar
- Pryor WA: Annu Rev Physiol. 1986, 48: 657-10.1146/annurev.ph.48.030186.003301.View ArticleGoogle Scholar
- Lebel CP, Ischiropoulos H, Bondy SC: Chem Res Toxicol. 1992, 5: 227-10.1021/tx00026a012.View ArticleGoogle Scholar
- Winterbourn CC: Free Radic Biol Med. 1987, 3: 33-10.1016/0891-5849(87)90037-2.View ArticleGoogle Scholar
- Makino K, Hagiwara T, Hagi A, Nishi M, Murakami A: Biochem Biophys Res Commun. 1990, 172: 1073-10.1016/0006-291X(90)91556-8.View ArticleGoogle Scholar
- Hildebrandt AG, Roots I: Arch Biochem Biophys. 1975, 171: 385-10.1016/0003-9861(75)90047-8.View ArticleGoogle Scholar
- Kosaka K, Yamada H, Matsui S, Echigo S, Shishida K: Environ Sci Technol. 1998, 32: 3821-10.1021/es9800784.View ArticleGoogle Scholar
- Holm T, George G, Barcelona M: Anal Chem. 1987, 59: 582-10.1021/ac00131a010.View ArticleGoogle Scholar
- Rota C, Chignell CF, Mason RP: Free Radic Biol Med. 1999, 27: 873-10.1016/S0891-5849(99)00137-9.View ArticleGoogle Scholar
- Mottola HA, Simpson BE, Gorin G: Anal Chem. 1970, 42: 410-10.1021/ac60285a017.View ArticleGoogle Scholar
- Zhang LS, Wong GTF: Talanta. 1994, 41: 2137-10.1016/0039-9140(94)00199-5.View ArticleGoogle Scholar
- Gregg SJ, Sing KSW: Adsorption Surface Area and Porosity. 1982, Academic, LondonGoogle Scholar
- Borda M, Elsetinow A, M Schoonen, Strongin D: Astrobiology. 2001, 1: 283-10.1089/15311070152757474.View ArticleGoogle Scholar
- Cohn CA, Borda MJ, Schoonen MA: Earth Planet Sci Lett. 2004, 225: 271-10.1016/j.epsl.2004.07.007.View ArticleGoogle Scholar
- Wilson CL, et al: Environ Sci Technol. 2000, 34: 2655-10.1021/es9906397.View ArticleGoogle Scholar