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Ind. Eng. Chem. Res. 2005, 44, 1660-1664
Saturation Concentration of Dissolved O2 in Highly Acidic Aqueous Solutions of H2SO4 Tomi Nath Das* Radiation Chemistry & Chemical Dynamics (RC & CD) Division, Bhabha Atomic Research Centre (BARC), Trombay, Mumbai 400 085, India
Saturation concentrations of dissolved oxygen in an aqueous-H2SO4 medium at 25 °C were measured in pulse radiolytic competition-kinetics experiments, wherein the H atom (H• radical) was generated in situ by a 50-ns electron pulse, and, subsequently, the propensity of its reactions with O2 and another solute of choice were compared on a microsecond time scale. Such direct chemical estimations reveal that an oxygen concentration of ∼1.28 mM at pH 2 steadily decreases to ∼310 µM at a H2SO4 concentration of ∼14 M, and it now provides a quantitative picture for evaluating oxidative stress in high acidity. Introduction
conditions, and the absence of experimental artifacts, as in the past, are the mainstay of this endeavor.
We have succeeded in measuring the in situ saturation concentration of dissolved oxygen in an aqueousH2SO4 medium at 25 °C for acid concentrations in the range of 100 mM to ∼14 M. Aqueous H2SO4 provides the complete range of acidity for a variety of needs in the laboratory and in industry, and such reaction media are also prevalent as atmospheric liquid sulfuric acid hydrometers. However, the dissolved O2 concentration in such matrixes, and its consequential chemical significance in the form of any oxidative stress, has not received due importance. Although the strong saltingout effect is expected to gradually reduce this chemical influence, the few available data, however, reveal a scattered and somewhat unusual trend. First such results, which are based on a volumetric comparison of the released O2 gas, suggested an initial gradual decrease in concentration with increasing acid content; however, beyond a H2SO4 concentration of ∼13 M, a reversal of this trend was reported.1,2 More recent but limited electrochemical measurements, on the other hand, show considerable disparity. For example, the ∼1.28 mM saturation solubility value in water at 25 °C and an O2 pressure of 1 atm (from ref 2) either shows negligible change with increases in the H2SO4 concentration (∼1.16 mM O2, for acid concentrations of both 1 and 2 M),3 which (i) matches the value of ∼1.15 mM O2 at 0.9 M acid reported in another study,4 but (ii) otherwise differs from it at an acid concentration of 2.6 M (∼0.85 mM O2).4 Our specific queries with the United States Geological Survey and its Office of Water Quality in Virginia5 also revealed that such seemingly simple but generally useful data are not yet available unequivocally. During our recent investigations of free-radical reactions in an aqueous-H2SO4 medium that used pulse radiolysis (PR), while measuring the dissolved O2 reactivity, we realized that the same measurements are also capable of directly revealing the variation in O2 concentration. A systematic and detailed study in this direction resulted in a set of plots of [O2] versus acidity. Their consistency under a variety of experimental * To whom correspondence should be addressed. Tel.: 91-22-25595097. Fax: 91-22-25519613. E-mail: tndas@ apsara.barc.ernet.in.
Experimental Section Method. The concerned experiments were performed based on the time-tested protocol of measuring the concentration of any dissolved solute in a solution based on its microsecond time scale (homogeneous) reaction kinetics with an in-situ-generated free radical. The PR technique in solution being a potent experimental tool for selective generation of free radical(s) kinetics, our estimation followed from its successful application. The PR experimental setup used has been previously described.6 Its salient features are as follows. The liquid sample (solution) was irradiated in a 1 cm × 1 cm square Suprasil cell by a 50-ns, 7-MeV horizontal electron pulse from a linear accelerator, thereby uniformly irradiating a sample volume of ∼1 mL. Within 0.1 µs, homogeneous radical concentrations on the order of a few micromolar were obtained in the irradiated volume.7 Although a mixture of oxidizing and reducing radicals was produced, only the reactions of H atoms (H• radicals) were selectively used in our studies, whereas all the other radicals were transformed to noninterfering types, as discussed later in this work. In the kinetic-spectrophotometric detection-analysis system that has been used, the horizontal analyzing light beam from a 450-W xenon lamp, after passing through the center of the irradiated sample, was dispersed in a Kratos monochromator that was blazed at 300 nm and the output intensity was measured by a Hamamatsu model R-955 photomultiplier tube coupled to a 200 MHz digital oscilloscope. Optical detection of transient radical species was routinely possible within the spectral range of 230-800 nm and in a real time of >500 ns. In these measurements, the spectral resolution could be maintained at e2 nm and the background signal due to scattered light was restricted to 18 MΩ/cm and an organic carbon level of 10 M) in the form PHH2•+ 335 (0.1-6 M) and 370 (>9 M) in the form CRH2•+ 430 (0.1-5.5 M) and 480 (>8 M) in the form GAH2•+ 380 in the form I2•-
a Refer to the Supporting Information for the experimental details.
To measure [O2] vs [H2SO4], the competitiveness of the H-atom reactions to form respective adducts with O2 (reaction 3) and various RS (reaction 4) was compared.
H• + O2 (+ H+) f H2O2•+
(3)
H• + RS f RSH•
(4)
Selection of the RS compounds were based on (i) their chemical stability in high acid solutions, (ii) their nonreactivity toward the O2 that is present, and (iii) their individual fast reaction kinetics, k4 g 109 M-1 s-1, which is comparable to the O2 reaction kinetics (k3 ) 1.2 × 1010 M-1 s-1 at pH 2).13 In addition to the direct measurement of the in situ H-atom concentration values in appropriate acid solutions,10,11b the current experimental estimation of [O2] was possible because of the direct measurement of relevant acid-dependent bimolecular rate constant (k4) and later estimation of the acid-dependent k3 values from these sets of data. In the literature, only a few such relevant k4 values are currently available (only at pH >0 for NP, PH, I2);13 all other k4 values were measured in this study. However, before such kinetic measurements, mandatory unequivocal physical characterization of each of the RSH• species (along with that of the respective RSH2•+ or RS•species, because of possible radical protonation/deprotonation reactions) constituted the first set of measurements. The relevant radical λmax values are shown in Table 1. In such argon-saturated solutions, the •OH radical could be selectively transformed to the chemically inert and noninterfering 2-hydroxy-2-methylpropyl radical, •CH2C(OH)(CH3)2, in the presence of the added 0.15 M TB.13 The SO4•- radical contributions in the concerned time scales (windows) and appropriate wavelengths were separately estimated (then conveniently minimizing the H-atom reactions by replacing argon with O2, but maintaining the same RS and TB concentrations) and later subtracted verbatim to reveal only the H-atom reactions.17 The latter was necessary because of the moderate SO4•- radical absorption at 150 µM and the solutions containing TB were purged with argon to minimize the loss of H atoms, following reaction 3. Subtracting the SO4•- radical contribution,17 the second-order k4 in eq 6 was obtained by following the increase in the pseudofirst-order formation kinetics (kobs) with increasing [RS]. The measured rate constants (in units of M-1 s-1) are plotted in Figure 1, using the Hammett acidity scale (H0):19
d[RSH•] ) k4 × [RS] × [H•] ≈ kobs × [H•] dt
(6)
In every case, the plots reveal that the actual decrease of the k4 values, even at an acid concentration of ∼14 M, was smaller (30%-70%) than the expected change (∼90%). Such an observation probably is in accordance with the quantum nature of the H atom, as compared to the solvent. However, an interesting but significant trend emerged, showing an inverse [H2SO4] dependence of the ∆k4 value, with respect to the k4 value at pH 2, plotted in Figure 2. As a corollary to such an observation, it is proposed that the magnitude of k3 would also follow this trend. It may be noted here that a direct experimental measurement of the k3 values, otherwise following a similar experimental protocol as that used for RS, was not possible, because of weak H2O2•+ radical absorbance at λ < 250 nm. The vertical dashed line in Figure 2 (≡ k3 value at pH 2) helps to reveal the other k3 values, where it meets the different H0 curves. For
3 2.5 2 1.5 1 0.75 0.7 0.3 0 -0.5 -0.56 -0.6 -0.8 -1.0 -1.4 -1.6 -1.9 -2.0 -2.5 -2.6 -2.9 -3.0 -3.7 -4.0 -4.2 -4.5 -4.7 -4.8 -5.0 -5.3 -5.4 -5.7 -6.0 -6.4 -6.5 -6.7 -7.0 96 wt % a
[H2SO4] (M)
[O2] vs NP
[O2] vs I2
[O2] vs CR
[O2] vs PH
0.001
1.28
0.0093
1.27
[O2] vs GA 1.31 1.29
0.09 0.096 0.1 0.22 0.43 1.06 1.15 1.23 1.61 2.07 3.11 3.57 4.19 4.39 5.4 5.6 6.2 6.41 7.81 8.4 8.79 9.38 9.76 9.95 10.33 10.89 11.07 11.61 12.14 12.83 12.99 13.32 13.8 ∼17
1.27
1.26
1.28
1.26 1.25 1.21
1.27 1.25 1.23
1.28 1.24 1.23
1.26 1.29 1.25 1.02
1.24
1.06 1.02 0.95 0.94 0.88 0.80
0.86
0.73 0.61
0.70 0.61
0.76
0.59
0.59 0.52
0.65
0.56 0.48 0.45
0.50
0.52
0.46
0.45
0.43
0.44
0.43
0.4
0.46 0.37 0.40
0.40
0.35
0.38
0.33
0.34
0.35
0.32
0.35 0.36 0.31 0.34
0.31
0.26
Data taken from Figure 3.
convenience, these acid-dependent k3 values are compared to different k4 values in Figure 1 (representing O2). Conclusion The final results from all these sets of [O2] measurements are consolidated in Figure 3 and Table 2, wherein the H0 scale has been correlated with the solution molarity,19 for easy reference and universal appeal. These plots show a steady decrease in [O2]saturation values with increases in acidity, suggesting a steady change
Ind. Eng. Chem. Res., Vol. 44, No. 6, 2005 1663
in the oxygen partial pressure. At lower acidity (i.e., up to ∼2 M), our results are similar to the results obtained in past studies,1-4 wherein, mainly because of the salting-out effect, oxygen solubility decreases linearly with respect to the value in pure water (with an empirical constant value of 0.0845, following the Setchenow model). However, at higher acidity, considerable differences are revealed, both in terms of the concentration values as well as the solubility trend discussed previously. Contrary to these reports,1,2 the current decreasing trend in the [O2] value with increasing solution acidity continued, and at an acid concentration of ∼17 M, the measured value was ∼260 µM (value not included in Figure 3, because only I2 could be used, as the other organic reactive solutes (RS) then were found to be thermally unstable). Keeping in mind the reported pKa values for sulfuric acid (pKa1 ) -10.0, pKa2 ) 1.9),20 as the solution acidity increases, it is expected that the composition of the solvent medium would undergo a gradual change from mainly the H2O/H2n+1On+/HSO4-/ SO42- type to the H2O/H2n+1On+/HSO4-/H2SO4 type (including different H2O-H2SO4 clusters21) with an increasing fraction of the undissociated acid (and, consequently, less ions), thus changing the solvent character and effecting a continuous change (decrease) in O2 solubility. Although our free-radical reaction kinetics methodology has been observed to compete favorably and also score better than the other, more conventional methods used in the past, the use of these O2 solubility data (with additional inputs at other temperatures) is envisaged in widely different areas such as hydrometallurgy and other related industries, studies in atmospheric22 and other sulfuric-acid-laden mists, and liquid hydrometers, and also in the laboratory, for a quantitative understanding and control of the influence of dissolved O2. Acknowledgment I thank Dr. T. Mukherjee, Associate Director, Chemistry Group, and my RC & CD Division colleagues for their support. I also thank the reviewer of this manuscript for valuable comments and suggestions. Supporting Information Available: Figures representing radiation dosimetry (radical concentrations) in an aqueous-H2SO4 medium and RSH• radical characteristics (PDF). This material is available free of charge via the Internet at http://pubs.acs.org. Literature Cited (1) Bohr, C. U ¨ ber die Lo¨slichkeit von Gasen in Konzentrierter Schwefelsa¨ure und in Mischungen von Schwefelsa¨ure und Wasser. Z. Phys. Chem. 1910, 71, 47. (2) Linde, W. F. Solubilities: Inorganic and Metal-organic Compounds. A Compilation of Solubility Data from the Periodical Literature, Vol. II, 4th Edition; American Chemical Society: Washington, DC, 1965; p 1229. (3) Kaskiala, T. Determination of Oxygen Solubility in Aqueous Sulphuric Acid Media. Miner. Eng. 2002, 15, 853. (4) Geffcken, G. Beitra¨ge zur Kenntnis der Lo¨slichkeitsbeeinflussung. Z. Phys. Chem. 1904, 49, 257. (5) Wilde, F. United States Geological Survey, private communication. (E-mail addresses:
[email protected] and h2osoft@ usgs.gov.) (6) (a) Guha, S. N.; Moorthy, P. N.; Kishore, K.; Naik, D. B.; Rao, K. N. One-Electron Reduction of Thionine Studied by Pulse Radiolysis. Proc. Indian Acad. Sci. (Chem. Sci.) 1987, 99, 261. (b) Das, T. N. Reactivity and Role of SO5•- Radical in Aqueous
Medium Chain Oxidation of Sulfite to Sulfate and Atmospheric Sulfuric Acid Generation. J. Phys. Chem. A 2001, 105, 9142. (7) (a) Keene, J. P. Radiation Chemistry: Kinetics of Radiation Induced Chemical Reactions. Nature 1960, 188, 843. (b) Matheson, M. S.; Dorfman, L. M. Pulse Radiolysis; MIT Press: Cambridge, MA, 1969. (c) Spinks, J. W. T.; Woods, R. J. An Introduction to Radiation Chemistry, 3rd Edition; Wiley: New York, 1990; pp 178-209. (8) Braun, W.; Herron, J. T.; Kahaner, D. Documentation and Sample Run File for ACUCHEM/ACUPLOT. A Computer Program for Modeling Complex Reaction Systems. Int. J. Chem. Kinet. 1988, 20, 51. (9) Buxton, G. V.; Stuart, C. R. Re-evaluation of the Thiocyanate Dosimeter for Pulse Radiolysis. J. Chem. Soc., Faraday Trans. 1995, 91, 279. (10) Das, T. N.; Ghanty, T. K.; Pal, H. Reactions of Methyl Viologen Dication (MV2+) with H Atoms in Aqueous Solution: Mechanism Derived from Pulse Radiolysis Measurements and Ab Initio MO Calculations. J. Phys. Chem. A 2003, 107, 5998. (11) (a) See data regarding radiation dosimetry (radical concentrations) in an aqueous-H2SO4 medium in the Supporting Information. (b) See data regarding RSH• radical characteristics in the Supporting Information. (12) (a) Allen, A. O. The Radiation Chemistry of Water and Aqueous Solutions; Van Nostrand: New York, 1961; p 58. (b) Kevan, L.; Moorthy, P. N.; Weiss, J. J. Formation and Reactions of Radiation-Produced Electrons and Atomic Hydrogen in γ-Irradiated Ice. J. Am. Chem. Soc. 1964, 86, 771. (c) Spinks, J. W. T.; Woods, R. J. An Introduction to Radiation Chemistry, 3rd Edition; Wiley: New York, 1990; pp 258-263. (d) Jiang, P.-Y.; Katsumura, Y.; Nagaishi, R.; Domae, M.; Ishikawa, K.; Ishigure, K.; Yoshida, Y. Pulse Radiolysis Study of Concentrated Sulfuric Acid Solutions. Formation Mechanism, Yield and Reactivity of Sulfate Radicals. J. Chem. Soc., Faraday Trans. 1992, 88, 1653. (e) Katsumura. Y. Radiation Chemistry of Concentrated Inorganic Aqueous Solutions in Radiation Chemistry: Present Status and Future Trends; Elsevier Science B.V.: Amsterdam, The Netherlands, 2001; p 163. (13) Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B. Critical Review of Rate Constants for Reactions of Hydrated Electrons, Hydrogen Atoms and Hydroxyl Radicals (•OH/O•-) in Aqueous Solution. J. Phys. Chem. Ref. Data 1988, 17, 513. (14) (a) Evers, E. L.; Jayson, G. G.; Robb, I. D.; Swallow, A. J. Determination by Pulse Radiolysis of the Distribution of Solubilizates between Micellar and Nonmicellar Phases. Naphthalene and Its Reduced Free Radical in Aqueous Sodium Dodecyl Sulphate Solutions. J. Chem. Soc., Faraday Trans. 1980, 76, 528. (b) Ross, A. B.; Bielski, B. H. J.; Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Huie, R. E.; Grodkowski, J.; Neta, P.; Mallard, W. G. NDRL-NIST Solution Kinetics Database, Version 3; NIST Standard Reference Database 40; National Institute of Standards and Technology: Gaithersburg, MD, 1994. (15) Wavelength shift that is due to the protonation reaction RSH• + H+ S RSH2•+ in higher acidity. (16) Hug, G. L. Optical Spectra of Nonmetallic Inorganic Transient Species in Aqueous Solution. NSRDS-NBS 69, U.S. Department of Commerce, National Bureau of Standards: Gaithersburg, MD, 1981; p 159. (17) In the presence of a few hundred micromolar RS and a few micromolar of H atoms (low dosage), the kinetic traces were analyzed using the ACUCHEM multicomponent kinetic data analysis software.8 The RS solution was divided into two parts: one part was saturated with O2, and the other was saturated with argon. Different volume ratios of these two were then mixed as required, prior to irradiation. Analysis of such kinetic traces obtained at the respective λmax values (see Table 1) under O2 saturation provided the contribution of the SO4•- radical toward its reaction with GA, PH, CR, and NP, including contributions from any unreacted SO4•- radicals and also products from the SO4•- + RS reaction (and the •CH2C(OH)(CH3)2 radical) to give the ∆Azero value of the measurements. Measurements under Ar + O2 saturation helped to calculate the loss in [RSH•] and appropriate kinetics of its various reactions. Incorporating such loss processes into the respective radical formation kinetics under argon saturation, the [RSH•]corrected (i.e., the ∆ARSH• value) and relevant kinetic values were finally obtained. (18) Sulco Chemicals, Ltd. H2SO4 Technology Bulletin, 2002, p 7. (Available via the Internet at www.sulcochemicals.com.)
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(19) (a) Jorgenson, M. J.; Hartter, D. R. A Critical Re-evaluation of the Hammett Acidity Function at Moderate and High Acid Concentrations of Sulfuric Acid. New H0 Values Based Solely on a Set of Primary Aniline Indicators. J. Am. Chem. Soc. 1963, 85, 878. (b) Ryabova, R. S.; Medvetskaya, I. M.; Vinnik, M. I. Acidity Functions of Aqueous H2SO4 Solutions. Russ. J. Phys. Chem. 1966, 40, 182. (c) Johnson, C. D.; Katritzky, A. R.; Shapiro, S. A. The Temperature Variation of the H0 Acidity Function in Aqueous Sulfuric Solution. J. Am. Chem. Soc. 1969, 91, 6654. (d) Rochester, C. H. Organic Chemistry. Acidity Functions, Vol. 17; Academic Press: London, 1970. (e) Gillespie, R. J.; Peel, T. E.; Robinson, E. A. The Hammett Acidity Function for Some Superacid Systems. I. The Systems H2SO4-SO3, H2SO4-HSO3F, H2SO4-HSO3Cl, and H2SO4-HB(HSO4)4. J. Am. Chem. Soc. 1971, 93, 5083. (20) Serjeant, E. P., Dempsey, B., Eds. Ionization Constants of Organic Acids in Solution; IUPAC Chemical Data Series No. 23; Pergamon Press: Oxford, U.K., 1979.
(21) Kathmann, S. M.; Hale, B. N. Monte Carlo Simulations of Small Sulfuric Acid-Water Clusters. J. Phys. Chem. B 2001, 105, 11719. (22) (a) Arnold, F.; Fabian, R. First Measurement of Gas-Phase Sulfuric Acid in the Stratosphere. Nature 1980, 283, 55. (b) Arnold, F.; Fabian, R.; Joos, W. Measurement of the Height Variation of Sulfuric Acid Vapor concentrations in the stratosphere. Geophys. Res. Lett. 1981, 8, 293. (c) Park, J. K.; Cho, S. Y. A Long-Range Transport of SO2 and Sulfate between Korea and China. Atmos. Environ. 1998, 32, 2745. (d) Lovejoy, E. R.; Curtius, J.; Froyd, K. D. Atmospheric Ion-Induced Nucleation of Sulfuric Acid and Water. J. Geo. Phys. Res. D 2004, 109, D08204.
Received for review May 28, 2004 Revised manuscript received December 14, 2004 Accepted December 24, 2004 IE049539M
