Reduction of hexavalent chromium by H2O2 in acidic solutions

The rates of the reduction of Cr(VI) with H2O2 were measured in NaCl solutions as a function of pH (1.5-4.8), temperature (5-40 degrees C), and ionic strength (I = 0.01-2 M) in the presence of an excess of reductant. The rate of Cr(VI) reduction is described by the general expression -d[Cr(VI)]/dt = k2[Cr(VI)](m)[H2O2](n)[H+](z), where m = 1 and n and z are two interdependent variables. The value of n is a function of pH between 2 and 4 (n = (3 x 10(a))/(1 + 10(a)), where a = -0.25 - 0.58pH + 0.26pH2) leveling off at pH < 2 (where n approximately = 1) and pH > 4 (where n approximately = 3). The rates of Cr(VI) reduction are acid-catalyzed, and the kinetic order z varies from about 1.8-0.5 with increasing H2O2 concentration, according to the equation z = 1.85 - 350.1H2O2 (M) which is valid for [H2O2] < 0.004 M. The values of k2 (M(-(n+z)) min(-1)) are given by k2 = k/[H+](z) = k1/[H2O2](n)[H+](z), where k is the overall rate constant (M(-n) min(-1)) and k, is the pseudo-first-order rate constant (min(-1)). The values of k in the pH range 2-4 have been fitted to the equation log k = 2.14pH - 2.81 with sigma = +/- 0.18. The values of k2 are dependent on pH as well. Most of the results with H2O2 < 3 mM are described by log k2 = 2.87pH - 0.55 with sigma = +/- 0.54. Experimental results suggest that the reduction of Cr(VI) to Cr(III) is controlled by the formation of Cr(V) intermediates. Values of k2 and k calculated from the above equations can be used to evaluate the rates of the reaction in acidic solutions under a wide range of experimental conditions, because the rates are independent of ionic strength, temperature, major ions, and micromolar levels of trace metals (Cu2+, Ni2+, Pb2+). The application of this rate law to environmental conditions suggests that this reaction may have a role in acidic solutions (aerosols and fog droplets) in the presence of high micromolar concentrations of H2O2.     

Iron(VI) and iron(V) oxidation of thiocyanate

Thiocyanate (SCN-) is used in many industrial processes and is commonly found in industrial and mining waste-waters. The removal of SCN- is required because of its toxic effects. The oxidation of thiocyanate (SCN-) by environmentally friendly oxidants, Fe(VI) and Fe(V), has been studied anaerobically using stopped-flow and premix pulse radiolysis techniques. The stoichiometry with Fe(VI) was determined to be 4HFeO(4-) + SCN(-) + 5H2O-->4Fe(OH)3 + SO4(2-) + CNO(-) + O2 + 2OH-. The rate law for the oxidation of SCN- by Fe(VI) was found to be -d[Fe(VI)]/dt = k11([H+]/([H+] + Ka,HFeO4)) [Fe(VI)][SCN-] where k11 = 2.04 +/- 0.04 x 10(3) M-1 s-1 and pKa,HFeO4 = 7.33. A mechanism is proposed that agrees with the observed reaction stoichiometry and rate law. The rate of oxidation of SCN- by Fe(V) was approximately 3 orders of magnitude faster than Fe(VI). The higher reactivity of Fe(V) with SCN- indicates that oxidations by Fe(VI) may be enhanced in the presence of appropriate one-electron-reducing agents. The results suggest that the effective removal of SCN- can be achieved by Fe(VI) and Fe(V).     

Catalyzed oxidation of arsenic(III) by hydrogen peroxide on the surface of ferrihydrite: an in situ ATR FTIR study

Knowledge of arsenic redox kinetics is crucial for understanding the impact and fate of As in the environment and for optimizing As removal from drinking water. Rapid oxidation of As(III) adsorbed to ferrihydrite (FH) in the presence of hydrogen peroxide (H2O2) might be expected for two reasons. First, the adsorbed As(III) is assumed to be oxidized more readily than the undissociated species in solution. Second, catalyzed decomposition of H2O2 on the FH surface might also lead to As(III) oxidation. Attenuated total reflection-Fourier transform infrared (ATR-FTIR) spectroscopy was used to monitor the oxidation of adsorbed As(III) on the FH surface in situ. No As(III) oxidation within minutes to hours was observed prior to H2O2 addition. Initial pseudo-first-order oxidation rate coefficients for adsorbed As(III), determined at H2O2 concentrations between 8.4 microM and 8.4 mM and pH values from 4 to 8, increased with the H2O2 concentration according to the equation log k(ox) (min(-1)) = 0.17 + 0.50 log [H2O] (mol/L), n = 21, r2 = 0.87. Only a weak pH dependence of log k(ox) was observed (approximately 0.04 logarithm unit increase per pH unit). ATR-FTIR experiments with As(III) adsorbed onto amorphous aluminum hydroxide showed that Fe was necessary to induce As(III) oxidation by catalytic H2O2 decomposition. Supplementary As(III) oxidation experiments in FH suspensions qualitatively confirmed the findings from the in situ ATR-FTIR experiments. Our results indicate that the catalyzed oxidation of As(III) by H2O2 on the surface of iron (hydr)oxides might be a relevant reaction pathway in environmental systems such as surface waters, as well as in engineered systems for As removal from water.     

Catalysis of elemental sulfur nanoparticles on chromium(VI) reduction by sulfide under anaerobic conditions

Chromate (CrVI) reduction by sulfide was conducted in anaerobic batch experimental systems. The molar ratio of the reduced CrVI to the oxidized S(-II) was 1:1.5 during the reaction, suggesting that the product of sulfide oxidation was elemental sulfur. Under the anaerobic condition, the reaction was pseudo first order initially with respect to CrVI, but the rate was dramatically accelerated at the later stage of the reaction. The rate acceleration was due to catalysis by elemental sulfur nanoparticles; dissolved species such as monomeric elemental sulfur and polysulfides appeared to be ineffective catalysts. Elemental sulfur nanoparticles were capable of adsorbing sulfide and such adsorbed sulfide exhibited much higher reactivity toward CrVI reduction than the aqueous-phase sulfide, resulting in the observed rate acceleration. Kinetic data under various reactant concentrations can be represented by the following empirical kinetic equation: -d[CrVI]/dt = k1 [CrVI][H2S]0.63 + k3[CrVI][triple bond S--SH]0.57. The first term on the right-hand side corresponds to the noncatalytic pathway, with k1 = 1.0 x 10(-3) (microM)(-0.63) min(-1) at pH 7.60 and 8.2 x 10(-5) (microM)-0.63 min(-1) at pH 8.10. The second term, k3[CrVI][triple bond S--SH]b, is the catalytic term with [triple bond S--SH] representing the adsorbed concentration of sulfide on the elemental sulfur nanoparticles (microM). The catalytic term is more important at the later stage of the reaction, as indicated by the observed kinetics and the enhancement of the reaction rate by externally added elemental sulfur nanoparticles. At pH 8.10, k3 = 0.0057 (microM)(-0.57) min(-1).     

Ferrate(VI) oxidation of weak-acid dissociable cyanides

Cyanide is commonly found in electroplating, mining, coal gasification, and petroleum refining effluents, which require treatment before being discharged. Cyanide in effluents exists either as free cyanide or as a metal complex. The kinetics of the oxidation of weak-acid dissociable cyanides by an environmentally friendly oxidant, ferrate(VI) (Fe(VI)O4(2-), Fe(VI)), were studied as a function of pH (9.1-10.5) and temperature (15-45 degrees C) using a stopped-flow technique. The weak-acid dissociable cyanides were Cd(CN)4(2-) and Ni(CN)4(2-), and the rate-laws for the oxidation may be -d[Fe(VI)]/dt = k[Fe(VI)][M(CN)4(2-)]n where n = 0.5 and 1 for Cd(CN)4(2-) and Ni(CN)4(2-), respectively. The rates decreased with increasing pH and were mostly related to a decrease in concentration of the reactive protonated Fe(VI) species, HFeO4(-). The stoichiometries with Fe(VI) were determined to be: 4HFeO4(-) + M(CN)4(2-) + 6H2O --> 4Fe(OH)3 + M(2+) + 4NCO(-) + O2 + 4OH(-). Mechanisms are proposed that agree with the observed reaction rate-laws and stoichiometries of the oxidation of weak-acid dissociable cyanides by Fe(VI). Results indicate that Fe(VI) is effective in removing cyanide in coke oven plant effluent, where organics are also present.     

Role of organically complexed iron(II) species in the reductive transformation of RDX in anoxic environments

Organically complexed iron species can play a significant role in many subsurface redox processes, including reactions that contribute to the transformation and degradation of soil and aquatic contaminants. Experimental results demonstrate that complexation of Fe(II) by catechol- and thiol-containing organic ligands leads to formation of highly reactive species that reduce RDX (hexahydro-1,3,5-trinitro-1,3,5-triazine) and related N-heterocyclic nitramine explosive compounds to formaldehyde and inorganic nitrogen byproducts. Under comparable conditions, relative reaction rates follow HMX < RDX < MNX < DNX < TNX. Observed rates of RDX reduction are heavily dependent on the identity of the Fe(II)-complexing ligands and the prevailing solution conditions (e.g., pH, Fe(II) and ligand concentrations). In general, reaction rates increase with increasing pH and organic ligand concentration when the concentration of Fe(II) is fixed. In solutions containing Fe(II) and tiron, a model catechol, observed pseudo-first-order rate constants (k(obs)) for RDX reduction are linearly correlated with the concentration of the 1:2 Fe(II)-tiron complex (FeL2(6-)), and kinetic trends are well described by -d[RDX]/dt= k(FeL2)6-[FeL2(6-)][RDX], where k(FeL2)6- = 7.31(+/-2.52) x 10(2) M(-1) s(-1). The reaction products and net stoichiometry (1 mol of RDX reduced for every 2 mol of Fe(II) oxidized) support a mechanism where RDX ring cleavage and decomposition is initiated by sequential 1-electron transfers from two Fe(II)-organic complexes.     

Kinetics and structural constraints of chromate reduction by green rusts

Green rusts, ferrous-ferric iron oxides, occur in many anaerobic soils and sediments and are highly reactive, making them important phases impacting the fate and transport of environmental contaminants. Despite their potential importance in environmental settings, reactions involving green rusts remain rather poorly described. Chromate is a widespread contaminant having deleterious impacts on plant and animal health; its fate may in part be controlled by green rust. Here we examine chromate reduction by a series of green rust phases and resolve the reaction kinetics at pH 7. The overall kinetics of the reactions are well described by the expression d[Cr(VI)]/dt = -k[Cr(VI)][GR], and this model was successfully used to predict rates of reaction at varying chromium concentrations. The rates of reduction are controlled by the concentration of ferrous iron, surface area, and chemical structure of the green rust including layer spacing. On a mass basis, green rust (GR) chloride is the most rapid reductant of Cr(VI) followed by GRCO3 and GRSO4, with pseudo-first-order rate coefficients (k(obs)) (with respect to Cr(VI) concentration) ranging from 1.22 x 10(-3) to 3.7 x 10(-2) s(-1). Chromium(III)-substituted magnetite and lepidocrocite were identified as the major oxidation products. The nature of the oxidation products appears to be independent of the anionic class of green rust, but their respective concentrations display a dependence on the initial GR. The mole fraction of Fe(III) in the Cr(x),Fe(1-x)(OH)3 x nH2O reaction product ranged from 17% to 68%, leading to a highly stabilized (low solubility) phase.     

Kinetics of uranium(VI) reduction by hydrogen sulfide in anoxic aqueous systems

Aqueous U(VI) reduction by hydrogen sulfide was investigated by batch experiments and speciation modeling; product analysis by transmission electron microscopy (TEM) was also performed. The molar ratio of U(VI) reduced to sulfide consumed, and the TEM result suggested that the reaction stoichiometry could be best represented by UO2(2+) + HS- = UO2+ S* + H+. At pH 6.89 and total carbonate concentration ([CO32-]T) of 4.0 mM, the reaction took place according to the following kinetics: -d[U(VI)]/dt = 0.0103[U(VI)][S2-]T0.54 where [U(VI)] is the concentration of hexavalent uranium, and [S2-]T is the total concentration of sulfide. The kinetics of U(VI) reduction was found to be largely controlled by [CO32-]T (examined from 0.0 to 30.0 mM) and pH (examined from 6.37 to 9.06). The reduction was almost completely inhibited with the following [CO32-]T and pH combinations: [(> or = 15.0 mM, pH 6.89); (> or = 4.0 mM, pH 8.01); and (> or = 2.0 mM, pH 9.06)]. By comparing the experimental results with the calculated speciation of U(VI), it was found that there was a strong correlation between the measured initial reaction rates and the calculated total concentrations of uranium-hydroxyl species; we, therefore, concluded that uranium-hydroxyl species were the ones being reduced by sulfide, not the dominant U-carbonate species present in many carbonate-containing systems.     

Atmospheric chemistry of HFE-7500 [n-C3F7CF(OC2H5)CF(CF3)2]: reaction with OH radicals and Cl atoms and atmospheric fate of n-C3F7CF(OCHO*)CF(CF3)2 and n-C3F7CF(OCH2CH2O*)CF(CF3)2 radicals

Relative rate techniques were used to measure k(OH + HFE-7500) = (2.6+/-0.6) x 10(-14), k(Cl + HFE-7500) = (2.3+/-0.7) x 10(-12), k[Cl + n-C3F7CF(OC(O)H)CF(CF3)2] = (9.7+/-1.4) x 10(-15), and k[Cl + n-C3F7CF(OC(O)CH3)CF(CF3)2] < 6 x 10(-17) cm3 molecule(-1) s(-1) at 295 K [HFE-7500 = n-C3F7-CF(OC2H5)CF(CF3)2]. From the value of k(OH + HFE-7500) an estimate of 2.2 years for the atmospheric lifetime of HFE-7500 is obtained. Two competing loss mechanisms for n-C3F7-CF(OCHO.CH3)CF(CF3)2 radicals were identified in 700 Torr of N2/O2 diluent at 295 K; reaction with O2 and decomposition via C-C bond scission with kO2/k(decomp) = 0.013+/-0.006 Torr(-1). The Cl atom initiated oxidation of HFE-7500 in N2/O2 diluent gives n-C3F7CF(OC(O)CH3)CF(CF3)2 as the major product and n-C3F7CF(OC(O)H)CF(CF3)2 as a minor product. The atmospheric oxidation of HFE-7500 gives n-C3F7-CF(OC(O)CH3)CF(CF3)2 and n-C3F7CF(OC(O)H)CF(CF3)2 as oxidation products. The results are discussed with respect to the atmospheric chemistry and environmental impact of HFE-7500.     

Density functional theory study on aqueous aluminum-fluoride complexes: exploration of the intrinsic relationship between water-exchange rate constants and structural parameters for monomer aluminum complexes

Density functional theory (DFT) calculation is carried out to investigate the structures, (19)F and (27)Al NMR chemical shifts of aqueous Al-F complexes and their water-exchange reactions. The following investigations are performed in this paper: (1) the microscopic properties of typical aqueous Al-F complexes are obtained at the level of B3LYP/6-311+G**. Al-OH(2) bond lengths increase with F(-) replacing inner-sphere H(2)O progressively, indicating labilizing effect of F(-) ligand. The Al-OH(2) distance trans to fluoride is longer than other Al-OH(2) distance, accounting for trans effect of F(-) ligand. (19)F and (27)Al NMR chemical shifts are calculated using GIAO method at the HF/6-311+G** level relative to F(H(2)O)(6)(-) and Al(H(2)O)(6)(3+) references, respectively. The results are consistent with available experimental values; (2) the dissociative (D) activated mechanism is observed by modeling water-exchange reaction for [Al(H(2)O)(6-i)F(i)]((3-i)+) (i = 1-4). The activation energy barriers are found to decrease with increasing F(-) substitution, which is in line with experimental rate constants (k(ex)). The log k(ex) of AlF(3)(H(2)O)(3)(0) and AlF(4)(H(2)O)(2)(-) are predicted by three ways. The results indicate that the correlation between log k(ex) and Al-O bond length as well as the given transmission coefficient allows experimental rate constants to be predicted, whereas the correlation between log k(ex) and activation free energy is poor; (3) the environmental significance of this work is elucidated by the extension toward three fields, that is, polyaluminum system, monomer Al-organic system and other metal ions system with high charge-to-radius ratio.     

Photoinduced oxidation of antimony(III) in the presence of humic acid

Interactions of antimony with natural organic matter (NOM) are important for the fate of Sb in aquatic systems. The kinetics of the photosensitized oxidation of Sb(III) to Sb(V) in the presence of Suwannee River Humic Acid (SRHA) was investigated using UV-A and visible light (medium-pressure mercury lamp). At a concentration of 5 mg L(-1) dissolved organic carbon (DOC) the light-induced reaction was 9000 times faster (rate coefficient k(exp) = 7.0 +/- 0.05 x 10(-4) s(-1)) than the dark reaction and followed pseudo-first-order kinetics. Rates increased linearly with the concentration of DOC. Between pH 4 and 8 rates increased by a factor of 5. Further results and kinetic considerations indicate that singlet oxygen, hydroxyl radicals, hydrogen peroxide, and hydroperoxyl radicals/superoxide are not important photooxidants in this system, while other NOM-derived reactive species, in particular excited triplet states and/or phenoxyl radicals, seem to be relevant. The dependence of rate coefficients on Sb(III)/DOC ratio was consistent with a two binding site model including (i) a strong binding site at low concentration inducing fast oxidation, (ii) a weak binding site at high concentration inducing slower oxidation, and (iii) the even slower oxidation of Sb(OH)3. Photoirradiation of natural water samples spiked with Sb(III) showed that the oxidation rates could be well predicted based on DOC.     

Formate ion decomposition in water under UV irradiation at 253.7 nm

Formate ion (HCO2-) occurs in natural waters as a result of photooxidation of humic substances. Under UV irradiation, as applied in water purification (253.7 nm), formate ion decomposed following split-rate pseudo-zero-order kinetics (k1 and k2 are initial and final rate constants, respectively). In the presence of dissolved oxygen (DO), it was found that (a) k1 < k2, (b) k1 and k2 increased with initial formate ion concentration ([HCO2-]0 = (1.73-38.3) x 10(-5) mol L(-1)) and absorbed UV intensity (Ia = (1.38-3.99) x 10(-6) mol quanta L(-1) s(-1)), and (c) k1 and k2 were relatively insensitive to initial pH (pHo = 5.41-8.97) in buffer-free solutions. Both rate constants decreased with increasing carbonate alkalinity ((0-1.0) x 10(-3) mol L(-1)) and k1 was virtually unchanged in phosphate buffer at pH0 between 5.25 and 9.92. Carbonate buffer lowered the rate of formate ion decay, possibly due to scavenging of OH* radicals. Initial rate constant k1 slightly increased with temperature (15-35 degrees C), while k2 remained unchanged. The reaction pH increased rapidly during irradiation of buffer-free NaHCO2 solution to approach an equilibrium level as [HCO2-] reached the method detection level (MDL). The pH profile of buffer-free formate ion decay was estimated using closed-system equilibrium analysis. DO utilization during UV irradiation was 0.5 mol of O2/mol of HCO2-, while nonpurgeable organic carbon (NPOC) measurements on kinetic samples closely followed the HCO2- profile, thus strongly suggesting the transformation of HCO2- -C to CO2 in the presence of DO. In DO-free water, k1 > k2 was observed. Furthermore, k(1,DO FREE) > k(1,DO) (k(1,DO) = k1) and k(2,DO FREE) < k(2,DO) (k(2,DO) = k2). The effect of dual acid solutions on HCO2- decay was examined in a mixture of NaHCO2 and sodium oxalate (Na2C2O4). HCO2- decomposed readily until [HCO2-] approximately equal to MDL but at a lower rate than in buffer-free HCO2- solutions, while C2O4(2-) remained virtually unchanged. C2O4(2-) decay commenced following near complete conversion of HCO2-.     

Kinetics and mechanistic aspects of As(III) oxidation by aqueous chlorine, chloramines, and ozone: relevance to drinking water treatment

Kinetics and mechanisms of As(III) oxidation by free available chlorine (FAC-the sum of HOCl and OCl-), ozone (O3), and monochloramine (NH2Cl) were investigated in buffered reagent solutions. Each reaction was found to be first order in oxidant and in As(III), with 1:1 stoichiometry. FAC-As(III) and O3-As(III) reactions were extremely fast, with pH-dependent, apparent second-order rate constants, k'app, of 2.6 (+/- 0.1) x 10(5) M(-1) s(-1) and 1.5 (+/- 0.1) x 10(6) M(-1) s(-1) at pH 7, whereas the NH2Cl-As(III) reaction was relatively slow (k'app = 4.3 (+/- 1.7) x 10(-1) M(-1) s(-1) at pH 7). Experiments conducted in real water samples spiked with 50 microg/L As(III) (6.7 x 10(-7) M) showed that a 0.1 mg/L Cl2 (1.4 x 10-6 M) dose as FAC was sufficient to achieve depletion of As(III) to <1 microg/L As(III) within 10 s of oxidant addition to waters containing negligible NH3 concentrations and DOC concentrations <2 mg-C/L. Even in a water containing 1 mg-N/L (7.1 x 10(-5) M) as NH3, >75% As(III) oxidation could be achieved within 10 s of dosing 1-2 mg/L Cl2 (1.4-2.8 x 10(-5) M) as FAC. As(III) residuals remaining in NH3-containing waters 10 s after dosing FAC were slowly oxidized (t1/2 > or = 4 h) in the presence of NH2Cl formed by the FAC-NH3 reaction. Ozonation was sufficient to yield >99% depletion of 50 microg/L As(III) within 10 s of dosing 0.25 mg/L O3 (5.2 x 10(-6) M) to real waters containing <2 mg-C/L of DOC, while 0.8 mg/L O3 (1.7 x 10(-5) M) was sufficientfor a water containing 5.4 mg-C/L of DOC. NH3 had negligible effect on the efficiency of As(III) oxidation by O3, due to the slow kinetics of the O3-NH3 reaction at circumneutral pH. Time-resolved measurements of As(III) loss during chlorination and ozonation of real waters were accurately modeled using the rate constants determined in this investigation.     

Kinetics and modeling of the Fe(III)/H2O2 system in the presence of sulfate in acidic aqueous solutions

This work examined the effect of sulfate ions on the rate of decomposition of H2O2 by Fe(III) in homogeneous aqueous solutions. Experiments were carried out at 25 degrees C, pH < or = 3 and the concentrations of sulfate ranged from 0 to 200 mM ([Fe(III)]0 = 0.2 or 1 mM, [H2O2]0 = 10 or 50 mM). The spectrophometric study shows that addition of sulfate decreased the formation of iron(III)-peroxo complexes and that H2O2 does not form complexes with iron(III)-sulfato complexes. The rates of decomposition of H2O2 markedly decreased in the presence of sulfate. The measured rates were accurately predicted by a kinetic model based on reactions previously validated in NaClO4/HClO4 solutions and on additional reactions involving sulfate ions and sulfate radicals. At a fixed pH, the pseudo-first-order rate constants were found to decrease linearly with the molar fraction of Fe(II) complexed with sulfate. The model was also able to predict the rate of oxidation of a probe compound (atrazine) by Fe(III)/H2O2. Computer simulations indicate that the decrease of the rate of oxidation of organic solutes by Fe(III)/H2O2 can be mainly attributed to the complexation of Fe(III) by sulfate ions, while sulfate radicals play a minor role on the overall reaction rates.     

Iron-mediated oxidation of antimony(III) by oxygen and hydrogen peroxide compared to arsenic(III) oxidation

Antimony is used in large quantities in a variety of products, though it has been declared as a pollutant of priority interest by the Environmental Protection Agency of the United States (USEPA). Oxidation processes critically affect the mobility of antimony in the environment since Sb(V) has a greater solubility than Sb(lll). In this study, the cooxidation reactions of Sb(lIl) with Fe(ll) and both O2 and H2O2 were investigated and compared to those of As(III). With increasing pH, the oxidation rate coefficients of Sb(lll) in the presence of Fe(ll) and O2 increased and followed a similar pH trend as the Fe(ll) oxidation by O2. Half-lives of Sb(lll) were 35 and 1.4 h at pH 5.0 and pH 6.2, respectively. The co-oxidation with Fe(ll) and H2O2 is about 7000 and 20 times faster than with Fe(ll) and O2 at pH 3 and pH 7, respectively. For both systems, *OH radicals appear to be the predominant oxidant below approximately pH 4, while at more neutral pH values, other unknown intermediates become important. The oxidation of As(lll) follows a similar pH trend as the Sb(lll) oxidation; however, As(lll) oxidation was roughly 10 times slower and only partly oxidized in most of the experiments. This study shows that the Fe(ll)-mediated oxidation of Sb(Ill) can be an important oxidation pathway at neutral pH values.     

Atmospheric chemistry of 2-ethoxy-3,3,4,4,5-pentafluorotetrahydro-2,5-bis[1,2,2,2-tetrafluoro-1-(trifluoromethyl)ethyl]-furan: kinetics, mechanisms, and products of Cl atom and OH radical initiated oxidation

Smog chamber/FTIR techniques were used to study the atmospheric chemistry of the title compound which we refer to as RfOC2H5. Rate constants of k(Cl + RfOC2H5) = (2.70 +/- 0.36) x 10(-12), k(OH + RfOC2H5) = (5.93 +/- 0.85) x 10(-14), and k(Cl + RfOCHO) = (1.34 +/- 0.20) x 10(-14) cm3 molecule(-1') s(-1) were measured in 700 Torr of N2, or air, diluent at 294 +/- 1 K. From the value of k(OH + RfOC2H5) the atmospheric lifetime of RfOC2H5 was estimated to be 1 year. Two competing loss mechanisms for RfOCH(O*)CH3 radicals were identified in 700 Torr of N2/O2 diluent at 294 +/- 1 K; decomposition via C-C bond scission giving a formate (RfOCHO), or reaction with 02 giving an acetate (RfOC(O)CH3). In 700 Torr of N2/O2 diluent at 294 +/- 1 K the rate constant ratio k(O2)/k(diss) = (1.26 +/- 0.74) x 10(-19) cm3 molecule(-1). The OH radical initiated atmospheric oxidation of RfOC2H5 gives Rf0CHO and RfOC(O)CH3 as major products. RfOC2H5 has a global warming potential of approximately 55 for a 100 year horizon. The results are discussed with respect to the atmospheric chemistry and environmental impact of RfOC2H5.     

Arsenic(III) oxidation by birnessite and precipitation of manganese(II) arsenate

Solution chemical techniques were used to investigate the oxidation of As(III) to As(V) in 0.011 M arsenite suspension of well-crystallized hexagonal birnessite (H-birnessite, 2.7 g L(-1)) at pH 5. Products of the reaction were studied by scanning electron microscopy coupled with energy dispersive spectroscopy (SEM-EDS), atomic force microscopy (AFM), and X-ray absorption near-edge structure spectroscopy (XANES). In the initial stage (first 74 h), chemical results have been interpreted quantitatively, and the reaction is shown to proceed in two steps as suggested by previous authors: 2>Mn(IV)O2 + H3AsO3 + H2O --> 2>Mn(III)OOH + H2AsO4- + H+ and 2>Mn(III)OOH + H3AsO3 + 3H+ --> 2Mn2+ + H2AsO4- + 2H2O. The As(III) depletion rate was lower (0.02 h(-1)) than measured in previous studies because of the high crystallinity of the H-birnessite sample used in this study. The surface reaction sites are likely located on the edges of H-birnessite layers rather than on the basal planes. The ion activity product of Mn(II) and As(V) reached after 74 h reaction time was the solubility product of a protonated manganese arsenate, having a chemical composition close to that of krautite as identified by XANES and EDS. Krautite precipitation reaction can be written as follows: Mn2+ + H2AsO4- + H2O = MnHAsO4 x H2O + H+ log Ks approximately -0.2. Equilibrium was reached after 400 h. The manganese arsenate precipitate formed long fibers that aggregated at the surface of H-birnessite. The oxidation reaction transforms a toxic species, As(III), to a less toxic aqueous species, which further precipitates with Mn2+ as a mixed As-Mn solid characterized by a low solubility product.     

Reduction of aqueous chromate by Fe(II)/Fe(III) carbonate green rust: kinetic and mechanistic studies

This work describes the heterogeneous reaction between FeII in carbonate green rust and aqueous chromate, in NaHCO3 solutions at 25 degrees C, and at pH values of 9.3-9.6. Evidence for reduction of CrVI to CrIII and concomitant solid-state oxidation of lattice FeII to FeIII was found from FeII titration and from structural analysis of the solids using FTIR, XRD, SEM, and XPS methods. Results indicate the formation of ferric oxyhydroxycarbonate and the concomitant precipitation of CrIII monolayers at the surface of the iron compound that induce passivation effects and progressive rate limitations. The number of CrIII monolayers formed at the completion of the reaction depends on [FeII]t=0, the molar concentration of FeII(solid) at t=0; on [n(o)]t=0, the molar concentration of reaction sites present at the surface of the solid phase at t=0; and on [CrVI]t=0, the molar concentration of CrVI at t=0. Kinetic data were modeled using a model based on the formation of successive CrIII monolayers, -(d[CrVI]/dt) = sigma(1)j k(i)[S] [CrVI]([n(i - 1)] - [n(i)]) with k(i)[S] (in s(-1) L mol(-1)), the rate coefficient of formation of CrIII monolayer i, and [n(i)] and [n(i - 1)], the molar concentration of CrIII precipitated in monolayer i and monolayer i - 1, respectively. Good matching curves were obtained with kinetic coefficients: k(1)[S] = 5-8 x 10(-4), k(2)[S] = 0.5-3 x 10(-5), and k(3)[S] about 1.7 x 10(-6) s(-1) m(-2) L. The CrVI removal efficiency progressively decreases along with the accumulation of CrIII monolayers at the surface of carbonate green rust particles. In the case of thick green rust particles resulting from the corrosion of iron in permeable reactive barriers, the quantity of FeII readily accessible for efficient CrVI removal should be rather low.     

Rates of arsenopyrite oxidation by oxygen and Fe(III) at pH 1.8-12.6 and 15-45 degrees C

The oxidation rate of arsenopyrite by dissolved oxygen was measured using a mixed flow reactor at dissolved O2 concentrations of 0.007-0.77 mM, pH 1.8-12.6, and temperatures of 15-45 degrees C. As(III) was the dominant redox species (>75%) in the experimental system, and the As(III)/As(V) ratio of effluent waters did not change with pH. The results were used to derive the following rate law expression (valid between pH 1.8 and 6.4): r = 10((-2211 +/- 57)T) (mO2)(0.45 +/- 0.05), where r is the rate of release of dissolved As in mol m(-2) s(-1) and T is in Kelvin. Activation energies (Ea) for oxidation of arsenopyrite by 02 at pH 1.8 and 5.9 are 43 and 57 kJ/mol, respectively, and they compare to an Ea value of 16 kJ/mol for oxidation by Fe(III) at pH 1.8. Apparent As release rates passed through a minimum in the pH range 7-8, which may have been due to oxidation of Fe2+ to hydrous ferric oxide (HFO) with attenuation of dissolved As onto the freshly precipitated HFO.     

Mechanism and kinetics of cyanogen chloride formation from the chlorination of glycine

Glycine is an important precursor of cyanogen chloride (CNCl)--a disinfection byproduct (DBP) found in chlorinated drinking water. To model CNCl formation from glycine during chlorination, the mechanism and kinetics of the reaction between glycine and free chlorine were investigated. Kinetic experiments indicated that CNCI formation was limited by either the decay rates of N,N-dichloroglycine or a proposed intermediate, N-chloroiminocarboxylate, CIN=CHCO2-. Only the anionic form of N,N-dichloroglycine, NCl2CH2CO2-, however, decays to form CNCl, while the protonated neutral species forms N-chloromethylimine. At pH > 6, glycine-nitrogen is stoichiometrically converted to CNCI, while conversion decreases at lower pH due to the formation of N-chloromethylimine. Under conditions relevant to drinking water treatment, i.e., at pH 6 to 8 and with free chlorine in excess, a simplified rate expression for the concentration of glycine-nitrogen converted to CNCl, [CNCl]f, applies: dt/d[CNCl]f = k2*[Cl2-Gly](T,o)exp(-k2*t) where [Cl2-Gly]T,o is the initial concentration of total N,N-dichloroglycine, k2* is the first-order decay constant for CIN=CHCO2-, k2*(s(-1)) = 10(12)(+/-4) exp(-1.0(+/-0.3) x 10(4)/T), and T is the absolute temperature in K. Kinetic expressions for d[CNCl]/dt when free chlorine is in excess, however, must also account for the significant decay of CNCl by hypochlorite-catalyzed hydrolysis, which has been characterized in previous studies. Although CNCl formation is independent of the free chlorine concentration, higher chlorine concentrations promote its hydrolysis.