Field determination of Fe2+ oxidation rates in acid mine drainage using a continuously-stirred tank reactor

A gravity-fed, battery-powered, portable continuously-stirred tank reactor has been developed to directly measure aqueous reaction rates in the field. Dye and tracer experiments indicate the reactor is well-mixed. Rates of Fe²⁺ oxidation at untreated and passively treated coal mine drainage sites in Pennsylvania were measured under ambient conditions and with the addition of either O₂ gas or NaOH solutions. Rates at 5 sites ranged from below the detection limit for this technique (approximately 10⁻⁹ mol L⁻¹ s⁻¹) to 3.27±0.01×10⁻⁶ mol L⁻¹ s⁻¹. Uncertainties in rates ranged from 70% near the lower limit of measurement to as little as 1% at higher rates of reaction. Multiple linear regressions showed no universal correlations of rates to Fe²⁺, dissolved O₂, and pH (Thiobacillus populations were not measured), although data for two more acidic sites were found to fit well for the model log rate=log K+a log [Fe²⁺]+b log [OH⁻]+c log [O₂]. Field rates of Fe oxidation from this and other studies vary by 4 orders of magnitude. A model using the ambient field rate of Fe oxidation from this study successfully reproduced independently-measured Fe²⁺ concentrations observed in a passive wetland treatment facility.     

Oxidation of pyrite in low temperature acidic solutions: Rate laws and surface textures

Rate laws have been determined for the aqueous oxidation of pyrite by ferric ion, dissolved oxygen and hydrogen peroxide at 30°C in dilute, acidic chloride solutions. Fresh, smooth pyrite grain surfaces were prepared by cleaning prior to experiments. Initial specific surface areas were measured by the multipoint BET technique. Surface textures before and after oxidation were examined by SEM. The initial rate method was used to derive rate laws.The specific initial rates of oxidation (moles pyrite cm⁻² min⁻¹) are given by the following rate laws (concentrations in molar units): r_sp,Fe³⁺ = −10^−9.74M^0.5_Fe³+M^−0.5_H+ (pH 1–2)r_sp,_o2 = −10^−6.77M^0.5_O2 (pH 2–4)r_sp,_h2o2 = −10^−1.43M_H2O2 (pH 2−4)An activation energy of 56.9 ± 7.5 kJ mole⁻¹ was determined for the oxidation of pyrite by dissolved oxygen from 20–40°C. HPLC analyses indicated that only minor amounts of polythionates are detectable as products of oxidation by oxygen below pH 4; the major sulfur product is sulfate. Ferric ion and sulfate are the only detectable products of pyrite oxidation by hydrogen peroxide. Hydrogen peroxide is consumed by catalytic decomposition nearly as fast as it is by pyrite oxidation.SEM photomicrographs of cleaned pyrite surfaces indicate that prior to oxidation, substantial intergranular variations in surface texture exist. Reactive surface area is substantially different than total surface area. Oxidation is centered on reactive sites of high excess surface energy such as grain edges and corners, defects, solid and fluid inclusion pits, cleavages and fractures. These reactive sites are both inherited from mineral growth history and applied by grain preparation techniques. The geometry and variation of reactive sites suggests that the common assumption of a first-order, reproducible dependence of oxidation rates on surface area needs to be tested.     

Solubility of Ca6[Al(OH)6]2(CrO4)3·26H2O,the chromate analog of ettringite; 5–75°C

Ca₆[Al(OH)₆]₂(CrO₄)₃·26H₂O, the chromate analog of the sulfate mineral ettringite, was synthesized and characterized by X-ray diffraction, Fourier transform infra-red spectroscopy, thermogravimetric analyses, energy dispersive X-ray spectrometry, and bulk chemical analyses. The solubility of the synthesized solid was measured in a series of dissolution and precipitation experiments conducted at 5–75°C and at initial pH values between 10.5 and 12.5. The ion activity product (IAP) for the reaction Ca₆[Al(OH)₆]₂(CrO₄)₃·26H₂O⇌6Ca²⁺+2Al(OH)⁻₄+3CrO²⁻₄+4OH⁻+26H₂O varies with pH unless a CaCrO_4(aq) complex is included in the speciation model. The log K for the formation of this complex by the reaction Ca²⁺+CrO²⁻₄=CaCrO_4(aq) was obtained by minimizing the variance in the IAP for Ca₆[Al(OH)₆]₂(CrO₄)₃·26H₂O. There is no significant trend in the formation constant with temperature and the average log K is 2.77±0.16 over the temperature range 5–75°C. The log solubility product (log K_SP) of Ca₆[Al(OH)₆]₂(CrO₄)₃·26H₂O at 25°C is −41.46±0.30. The temperature dependence of the log K_SP is log K_SP=A−B/T+D log(T) where A=498.94±48.99, B=27,499±2257, and D=−181.11±16.74. The values of ΔG⁰_r,298 and ΔH⁰_r,298 for the dissolution reaction are 236.6±3.9 and 77.5±2.4 kJ mol⁻¹. the values of ΔC⁰_P,r,298 and ΔS⁰_r,298 are −1506±140 and −534±83 J mol⁻¹ K⁻¹. Using these values and published standard state partial molal quantities for constituent ions, ΔG⁰_f,298=−15,131±19 kJ mol⁻¹, ΔH⁰_f,298=−17,330±8.6 kJ mol⁻¹, ΔS⁰₂₉₈=2.19±0.10 kJ mol⁻¹ K⁻¹, and ΔC⁰_Pf,298=2.12±0.53 kJ mol⁻¹ K⁻¹, were calculated.     

Solubility of ettringite (Ca6[Al(OH)6]2(SO4)3 · 26H2O) at 5–75°C

The solubility of ettringite (Ca₆[Al(OH)₆]₂(SO₄)₃ · 26H₂O) was measured in a series of dissolution and precipitation experiments at 5–75°C and at pH between 10.5 and 13.0 using synthesized material. Equilibrium was established within 4 to 6 days, with samples collected between 10 and 36 days. The log K_SP for the reaction Ca₆[Al(OH)₆]₂(SO₄)₃ · 26H₂O ⇌ 6Ca²⁺ + 2Al(OH)₄⁻ + 3SO₄²⁻ + 4OH⁻ + 26H₂O at 25°C calculated for dissolution experiments (−45.0 ± 0.2) is not significantly different from the log K_SP calculated for precipitation experiments (−44.8 ± 0.4) at the 95% confidence level. There is no apparent trend in log K_SP with pH and the mean log K_SP,298 is −44.9 ± 0.3. The solubility product decreased linearly with the inverse of temperature indicating a constant enthalpy of reaction from 5 to 75°C. The enthalpy and entropy of reaction ΔH^°_r and ΔS^°_r, were determined from the linear regression to be 204.6 ± 0.6 kJ mol⁻¹ and 170 ± 38 J mol⁻¹ K⁻¹. Using our values for log K_SP, ΔH^°_r, and ΔS^°_r and published partial molal quantities for the constituent ions, we calculated the free energy of formation ΔG^°_f,298, the enthalpy of formation ΔH^°_f,298, and the entropy of formation ΔS^°_f,298 to be −15211 ± 20, −17550 ± 16 kJ mol⁻¹, and 1867 ± 59 J mol⁻¹ K⁻¹. Assuming ΔC_P,r is zero, the heat capacity of ettringite is 590 ± 140 J mol⁻¹ K⁻¹.     

Iron control by equilibria between hydroxy-Green Rusts and solutions in hydromorphic soils

In order to verify Fe control by solution - mineral equilibria, soil solutions were sampled in hydromorphic soils on granites and shales, where the occurrence of Green Rusts had been demonstrated by Mössbauer and Raman spectroscopies. Eh and pH were measured in situ, and Fe(II) analyzed by colorimetry. Ionic Activity Products were computed from aqueous Fe(II) rather than total Fe in an attempt to avoid overestimation by including colloidal particles. Solid phases considered are Fe(II) and Fe(III) hydroxides and oxides, and the Green Rusts whose general formula is [Fe^II_1−xFe^III_x(OH)₂]^+x· [x/z A^−z]^−x, where compensating interlayer anions, A⁻, can be Cl⁻, SO₄²⁻, CO₃²⁻ or OH⁻, and where x ranges a priori from 0 to 1. In large ranges of variation of pH, pe and Fe(II) concentration, soil solutions are (i) oversaturated with respect to Fe(III) oxides; (ii) undersaturated with respect to Fe(II) oxides, chloride-, sulphate- and carbonate-Green Rusts; (iii) in equilibrium with hydroxy-Green Rusts, i.e., Fe(II)-Fe(III) mixed hydroxides. The ratios, x = Fe(III)/Fe_t, derived from the best fits for equilibrium between minerals and soil solutions are 1/3, 1/2 and 2/3, depending on the sampling site, and are in every case identical to the same ratios directly measured by Mössbauer spectroscopy. This implies reversible equilibrium between Green Rust and solution. Solubility products are proposed for the various hydroxy-Green Rusts as follows: log K_sp = 28.2 ± 0.8 for the reaction Fe₃(OH)₇ + e⁻ + 7 H⁺ = 3 Fe²⁺ + 7 H₂O; log K_sp = 25.4 ± 0.7 for the reaction Fe₂(OH)₅ + e⁻ + 5 H⁺ = 2 Fe²⁺ + 5 H₂O; log K_sp = 45.8 ± 0.9 for the reaction Fe₃(OH)₈ + 2e⁻ + 8 H⁺ = 3 Fe²⁺ + 8 H₂O at an average temperature of 9 ± 1°C, and 1 atm. pressure. Tentative values for the Gibbs free energies of formation of hydroxy-Green Rusts obtained are: Δ_fG° (Fe₃(OH)₇, cr, 282.15 K) = −1799.7 ± 6 kJ mol⁻¹, Δ_fG° (Fe₂(OH)₅, cr, 282.15 K) = −1244.1 ± 6 kJ mol⁻¹ and Δ_fG° (Fe₃(OH)₈, cr, 282.15 K) = −1944.3 ± 6 kJ mol⁻¹.     

The aqueous solubility of trichloroethene (TCE) and tetrachloroethene (PCE) as a function of temperature

Using a flexible Au bag autoclave and a precision high-pressure liquid chromatography pump to control pressure, the liquid–liquid aqueous solubilities of TCE and PCE were measured as a function of temperature from 294 to 434 K (at constant pressure). The results were used to calculate the partial molal thermodynamic quantities of the organic liquid aqueous dissolution reactions: ΔḠ_soln, ΔH̄_soln, ΔS̄_soln and ΔC̄_{p soln}. Calculated values for these quantities at 298 K for TCE are: ΔḠ_soln=11.282 (±0.003) kJ/mol, ΔH̄_soln=−3.35 (±0.07) kJ/mol, ΔS̄_soln=−49.07 (±0.24) J/mol K, and ΔC̄_{p soln}=385.2 (±3.4) J/mol K. Calculated values for these quantities at 298 K for PCE are: ΔḠ_soln=15.80 (±0.04) kJ/mol, ΔH̄_soln=−1.79 (±0.58) kJ/mol, ΔS̄_soln=−59.00 (±1.96) J/mol K and ΔC̄_{p soln}=354.6 (±8.6) J/mol K. These thermodynamic quantities may be used to calculate the solubility of TCE and PCE at any temperature of interest. In the absence of direct measurements over this temperature range, the Henry's Law constants for TCE and PCE have been estimated using the measured aqueous solubilities and calculated vapor pressures.     

Benthic fluxes and the cycling of biogenic silica and carbon in two southern California borderland basins

Benthic fluxes in two southern California borderland basins have been estimated by modeling water column property gradients, by modeling pore water gradients and by measuring changes in concentration in a benthic chamber. Results have been used to compare the different methods, to establish budgets for biogenic silica and carbon and to estimate rate constants for models of CaCO₃ dissolution. In San Pedro Basin, a low oxygen, high sedimentation rate area, fluxes of radon-222 (86 ± 8 atoms m⁻² s⁻¹), SiO₂ (0.7 ± 0.1 mmol m⁻² d⁻¹), alkalinity (1.7 ± 0.3 meq m⁻² d⁻¹), TCO₂ (1.9 ± 0.3 mmol m⁻² d⁻¹) and nitrate (−0.8 ± 0.1 mmol m⁻² d⁻¹) measured in a benthic chamber agree within the measurement uncertainty with fluxes estimated from modeling profiles of nutrients and radon obtained in the water column. The diffusive fluxes of radon, SiO₂ and TCO₂ determined from modeling the sediment and pore water also agree with the other approaches. Approximately 33 ± 13% of the organic carbon and 37 ± 47% of the CaCO₃ arriving at the sea floor are recycled. In San Nicolas Basin, which has larger oxygen concentrations and lower sedimentation rates than San Pedro, the fluxes of radon (490 ± 16 atoms m⁻² s⁻¹), SiO₂ (0.7 ± 0.1 mmol m⁻² d⁻¹), alkalinity (1.7 ± 0.3 meq m⁻² d⁻¹), TCO₂ (1.7 ± 0.2 mmol m⁻² d⁻¹), oxygen (−0.7 ± 0.1 mmol m⁻² d⁻¹) and nitrate (-0.4 ± 0.1 mmol m⁻² d⁻¹) determined from chamber measurements agree with the water column estimates given the uncertainty of the measurements and model estimates. Diffusion from the sediments matches the lander-measured SiO₂ and PO₄³⁻ (0.017 ± 0.002 mmol m⁻² d⁻¹) fluxes, but is not sufficient to supply the radon or TCO₂ fluxes observed with the lander. In San Nicolas Basin 38 ± 9% of the organic carbon and 43 ± 22% of the CaCO₃ are recycled. Approximately 90% of the biogenic silica arriving at the sea floor in each basin is recycled. The rates of CaCO₃ dissolution determined from chamber flux measurements and material balances for protons and electrons are compared to those predicted by previously published models of CaCO₃ dissolution and this comparison indicates that in situ rates are comparable to those observed in laboratory studies of bulk sediments, but orders of magnitude less than those observed in experiments done with suspended sediments.     

Mechanism of kaolinite dissolution at room temperature and pressure Part II: kinetic study

Studies on the dissolution kinetics of kaolinite were performed using batch reactors at 25°C and in the pH range from 1 to 13. A rapid initial dissolution step was first observed, followed by a linear kinetic stage reached after approximately 600 hr of reaction during which the kaolinite dissolves congruently at pH < 4 and pH > 11. The apparent incongruency between pH 5 and 10 was due to the precipitation of an Al–hydroxide phase. The true dissolution rates were computed from the amount of Si released into solution. The rate dependence on pH can be described by: r = 10−12.19aH+0.55 + 10−14.36 + 10−10.71aOH−0.75Between pH 5 and 10, the rate is approximately constant, although a smooth minimum was observed at pH close to 9. mAn attempt was made to obtain a general rate law based on the coordination theory, which was first applied to the mineral dissolution studies by Stumm and co-workers. The kinetic data were combined with the results obtained for the surface speciation by Huertas et al. (1998). It is possible to express the linear dissolution rate as a simple power function of the concentration of the surface sites active in various pH ranges: r = 10−8.25 [>Al2OH2+] + 10−10.82 [>AlOH2+]0.5 + 10−9.1 [>Al2OH + >AlOH + >SiOH] + 103.78 [>Al2O− + >AlO−]3This equation assumes that the dissolution mechanism is mainly controlled by the two Al surface sites (external and internal structural hydroxyls, and aluminol at the crystal edges) under both acidic and alkaline conditions. The model reflects well the important contribution of the crystal basal planes to the dissolution of kaolinite.     

Coal weathering and the geochemical carbon cycle

The weathering rate of sedimentary organic matter in the continental surficial environment is poorly constrained despite its importance to the geochemical carbon cycle. During this weathering, complete oxidation to carbon dioxide is normally assumed, but there is little proof that this actually occurs. Knowledge of the rate and mechanisms of sedimentary organic matter weathering is important because it is one of the major controls on atmospheric oxygen level through geologic time.We have determined the aqueous oxidation rates of pyrite-free bituminous coal at 24° and 50°C by using a dual-cell flow-through method. Coal was used as an example of sedimentary organic matter because of the difficulty in obtaining pyrite-free kerogen for laboratory study. The aqueous oxidation rate obtained in the present study for air-saturated water (270 μM O₂) was found to be on the order of 2 × 10⁻¹² mol O₂/m²/s at 25°C, which is fast compared to other geologic processes such as tectonic uplift and exposure through erosion. The reaction order with respect to oxygen level is 0.5 on a several thousand hour time scale for both 24° and 50°C experiments. Activation energies, determined under 24° and 50°C conditions, were ≈40 kJ/mol O₂ indicating that the oxidation reaction is surface reaction controlled.The oxygen consumption rate obtained in this study is two to three orders of magnitude smaller than that for pyrite oxidation in water, but still rapid on a geologic time scale. Aqueous coal oxidation results in the formation of dissolved CO₂, dissolved organic carbon (DOC), and solid oxidation products, which are all quantitatively significant reaction products.     

Relative contributions of abiotic and biological factors in Fe(II) oxidation in mine drainage

The oxidation of Fe(II) is apparently the rate-limiting step in passive treatment of coal mine drainage. Little work has been done to determine the kinetics of oxidation in such field systems, and no models of passive treatment systems explicitly consider iron oxidation kinetics. A Stella II^™ model using Fe(II)_init concentration, pH, temperature, Thiobacillus ferrooxidans and O₂ concentration, flow rate, and pond volume is used to predict Fe(II) oxidation rates and concentrations in seventeen ponds under a wide range of conditions (pH 2.8 to 6.8 with Fe(II) concentrations of less than 240 mg L⁻¹) from 6 passive treatment facilities. The oxidation rate is modeled based on the combination of published abiotic and biological laboratory rate laws. Although many other variables have been observed to influence Fe(II) oxidation rates, the 7 variables above allow field systems to be modeled reasonably accurately for conditions in this study.Measured T. ferrooxidans concentrations were approximately 10⁷ times lower than concentrations required in the model to accurately predict field Fe(II) concentrations. This result suggests that either 1) the most probable number enumeration method underestimated the bacterial concentrations, or 2) the biological rate law employed underestimated the influence of bacteria, or both. Due to this discrepancy, bacterial concentrations used in the model for pH values of less than 5 are treated as fit parameters rather than empirically measured values.Predicted Fe(II) concentrations in ponds agree well with measured Fe(II) concentrations, and predicted oxidation rates also agree well with field-measured rates. From pH 2.8 to approximately pH 5, Fe(II) oxidation rates are negatively correlated with pH and catalyzed by T. ferrooxidans. From pH 5 to 6.4, Fe(II) oxidation appears to be primarily abiotic and is positively correlated with pH. Above pH 6.4, oxidation appears to be independent of pH. Above pH 5, treatment efficiency is affected most by changing design parameters in the following order: pH>temperature≈influent Fe(II)>pond volume≈O₂. Little to no increase in Fe(II) oxidation rate occurs due to pH increases above pH 6.4. Failure to consider Fe(II) oxidation rates in treatment system design may result in insufficient Fe removal.     

Biogeochemical cycling in an organic-rich coastal marine basin. 6. Temporal and spatial variations in sulfate reduction rates

Rates of sulfate reduction were measured over a 3 year period in the anoxic nearshore sediments of Cape Lookout Bight, North Carolina, using both a tube incubation method and a ³⁵S-sulfate direct injection technique. The methods yielded similar depth-integrated rates over the upper 30 cm ranging from less than 10 mol SO⁼₄ · m⁻² · y⁻¹ in winter to greater than 50 mol SO⁼₄ · m⁻² · y⁻¹ in summer. There were also seasonal changes in the Arrhenius activation energies for the sulfate reduction rates indicating that the assumption that E_a is constant with temperature is not always valid. The time averaged annual turnover rate for all three years was 20.4 (±11.4) mol SO⁼₄ · m⁻² · y⁻¹. Surface rates ranged seasonally from less than 0.01 to over 3 mM SO⁼₄ · d⁻¹ between winter and summer, respectively. A subsurface rate maximum was observed to develop during the summer months which accounted for 28 percent of the annual depth integrated sulfate reduction rate. Subsurface rate maxima are the result of changes in the chemistry (substrate type and/or concentration) and the microbiology in the sediments. The possibility of the subsurface maximum being an artifact of the ³⁵S method is also discussed. However, the sulfate reduction rates compare well with previous measurements of the carbon sediment-water plus burial fluxes and with a depth integrated CO₂ production rate modelled from a ΣCO₂ concentration profile from the same site.     

Quantification of oxygen consumption and sulphate release rates for waste rock piles using kinetic cells: Cluff lake uranium mine,northern Saskatchewan,Canada

Oxidation rates of low sulphide (<0.5 wt.%) gneissic waste rock from the Cluff lake U mine, northern Saskatchewan, Canada were determined using 3 independent methods: O₂ consumption rates in kinetic cells, SO₄ measurements of kinetic cell effluent and humidity cell SO₄ release rates. The O₂ consumption measurements demonstrated that the oxidation of pyrite was strongly dependent on grain size and moderately dependent on water content, temperature and microbiology. Oxygen consumption rates were highest at water contents of 5–10 wt.% (12–25% saturation). Measured SO₄ release rates (3.1–91 mg SO₄ kg⁻¹ wk⁻¹) for the kinetic cells were comparable to rates calculated from the O₂ consumption values (6.9–70 mg SO₄ kg⁻¹ wk⁻¹). Sulphate release rates determined from humidity cells were generally higher than those obtained from the kinetic cells, ranging from 6 to 64 mg SO₄ kg⁻¹ wk⁻¹ for the coarsest and finest fraction, respectively. These differences were attributed to sample heterogeneity.     

Preparation and solubility of northupite from brine and its adsorption properties for Cu(II) and Cd(II) in seawater

The mineral northupite Na₃Mg(CO₃)₂Cl was synthesized from a solar Adriatic seawater brine pond to which Na₂CO₃ was added at 373°K. The precipitated northupite had a surface area (P) of 6.0 ± 0.4 m²g⁻¹, and the thermodynamic solubility product was estimated to be log K Na₃Mg(CO₃)₂Cl = −4.8 ± 0.3 at 25°C. This value was used to calculate the interfacial energy (σ = 50 erg cm⁻²) for the homogeneous nucleation of northupite. The solubility constant determined in this study has been used to examine the saturation state of Mahega Lake and Lake Katwe (Uganda). The waters from Lake Katwe were found to be supersaturated with respect to northupite.The adsorption of Cu and Cd onto northupite particles was studied in seawater. Both metals are strongly adsorbed. Adsorption constants and the specific area of northupite occupied by Cd and Cu using Langmuir adsorption isotherms and equilibrium constants for surface complex formation have been determined.     

The heat capacity of a natural monticellite and phase equilibria in the system CaO-MgO-SiO2-CO2

The heat capacity of a natural monticellite (Ca_1.00Mg_.09Fe_.91Mn_.01Si_0.99O_3.99) measured between 9.6 and 343 K using intermittent-heating, adiabatic calorimetry yields C_p⁰(298) and S₂₉₈⁰ of 123.64 ± 0.18 and 109.44 ± 0.16 J · mol⁻¹K⁻¹ respectively. Extrapolation of this entropy value to end-member monticellite results in an S⁰₂₉₈ = 108.1 ± 0.2 J · mol⁻¹K⁻¹. High-temperature heat-capacity data were measured between 340–1000 K with a differential scanning calorimeter. The high-temperature data were combined with the 290–350 K adiabatic values, extrapolated to 1700 K, and integrated to yield the following entropy equation for end-member monticellite (298–1700 K): S_T⁰(J · mol⁻¹K⁻¹) = S₂₉₈⁰ + 164.79 In T + 15.337 · 10⁻³T + 22.791 · 10⁵T⁻² − 968.94. Phase equilibria in the CaO-MgO-SiO₂ system were calculated from 973 to 1673 K and 0 to 12 kbar with these new data combined with existing data for akermanite (Ak), diopside (Di), forsterite (Fo), merwinite (Me) and wollastonite (Wo). The location of the calculated reactions involving the phases Mo and Fo is affected by their mutual solid solution. A best fit of the thermodynamically generated curves to all experiments is made when the S⁰₂₉₈ of Me is 250.2 J · mol⁻¹ K⁻¹ less than the measured value of 253.2 J · mol⁻¹ K⁻¹.A best fit to the reversals for the solid-solid and decarbonation reactions in the CaO-MgO-SiO₂-CO₂ system was obtained with the ΔG⁰₂₉₈ (kJ · mole⁻¹) for the phases Ak(−3667), Di(−3025), Fo(−2051), Me(−4317) and Mo(−2133). The two invariant points − Wo and −Fo for the solid-solid reactions are located at 1008 ± 5 K and 6.3 ± 0.1 kbar, and 1361 ± 10 K and 10.2 ± 0.2 kbar respectively. The location of the thermodynamically generated curves is in excellent agreement with most experimental data on decarbonation equilibria involving these phases.     

In situ push-pull tests for the determination of TCE degradation and permanganate consumption rates

A simple, single-well push-pull test was conducted at a TCE-contaminated site to estimate the site-specific TCE degradation and permanganate (MnO₄⁻) consumption rate. Known quantities of a conservative tracer (Br⁻) and permanganate were rapidly injected into a saturated aquifer then periodically sampled during extraction from the same well. Concentrations of Br⁻, TCE, and MnO₄⁻ were measured; breakthrough curves (BTCs) for all species of solute were determined. Data analysis of BTCs for bromide and TCE showed that the first-order rate constant of TCE degradation by MnO₄⁻ is 1.67 ± 0.152 h⁻¹. Further, the in situ MnO₄⁻ demand rate by TCE and aquifer materials is estimated to be 0.54 ± 0.371 h⁻¹. This study demonstrates that in situ push-pull tests are useful and economical tools for field investigations to determine contaminant reaction and oxidant consumption rates, which may then be used to optimize groundwater remediation designs.     

A kinetic study of the oxidation of arsenopyrite in acidic solutions: implications for the environment

Arsenopyrite is an important component of many ore deposits and dissolves in the O₂-rich, acidic surface waters that are commonly found in the vicinity of active mines, releasing As, Fe and S to the environment. However, despite the potentially serious effect of this pollution on the human and animal population, the rate at which such oxidation occurs is poorly known. Kinetic experiments were therefore conducted in a mixed flow reactor to investigate the oxidation of arsenopyrite in Fe₂(SO₄)₃ solutions (pH=l.8) having a concentration of l×l0⁻² to 1 ×l0⁻⁵ mol kg⁻¹ at temperatures of 45, 35, 25 and 15 °C. The results of these experiments show that the rate of oxidation of arsenopyrite increases with increasing concentration of dissolved Fe₂(SO₄)₃ and temperature. They also show that As released during the oxidation of arsenopyrite has the form As(III), and that the rate of conversion of As(III) to As(V) is relatively low, although it tends to increase with increasing concentration of dissolved Fe₂(SO₄)₃ and temperature. In the presence of Cl⁻, oxidation of arsenopyrite is accelerated, as is the conversion of As(III) to As(V). These findings indicate that exploitation of arsenopyrite-bearing ores will cause contamination of groundwaters by As at levels sufficient to have a major negative effect on the health of humans and animals.     

Kaolinite and smectite dissolution rate in high molar KOH solutions at 35° and 80°C

Experiments measuring kaolinite and smectite dissolution rates were carried out using batch reactors at 35° and 80°C. No potential catalysts or inhibitors were present in solution. Each reactor was charged with 1 g of clay of the ≤2 μm fraction and 80, 160 or 240 ml of 0.1–4 M KOH solution. An untreated but sized kaolinite from St. Austell and two treated industrial smectites were used in the experiments. One smectite is a nearly pure montmorillonite, while the second has a significant component of beidellitic charge (35%). The change in solution composition and mineralogy was monitored as a function of time. Initially, the 3 clays dissolved congruently. No new formed phases were observed by XRD and SEM during the pure dissolution stage. The kaolinite dissolution is characterized by a linear release of silica and Al as a function of the log of time. This relationship can be explained by a reaction affinity effect which is controlled by the octahedral layer dissolution. Far from equilibrium, dissolution rates are proportional to a^0.56±0.12_OH⁻ at 35°C and to a^0.81±0.12_OH⁻ at 80°C. The activation energy of kaolinite dissolution increases from 33±8 kJ/mol in 0.1 M KOH solutions to 51±8 kJ/mol in 3 M KOH solutions. In contrast to kaolinite, the smectites dissolve at much lower rates and independently of the aqueous silica or Al concentrations. The proportionality of the smectite dissolution rate constant at 35 and 80°C was a^0.15±0.06_OH⁻. The activation energy of dissolution appears to be independent of pH for smectite and is found to be 52±4 kJ/mol. The differences in behavior between the two kinds of minerals can be explained by structural differences. The hydrolysis of the tetrahedral and the octahedral layer appears as parallel reactions for kaolinite dissolution and as serial reactions for smectite dissolution. The rate limiting step is the dissolution of the octahedral layer in the case of kaolinite, and the tetrahedral layer in the case of smectite.     

Sulfide mineral oxidation in freshly processed tailings: batch experiments

This work focuses on sulfide mineral oxidation rates under oxic conditions in freshly processed pyrite-rich tailings from the ore concentrator in Boliden, northern Sweden. Freshly processed tailings are chemically treated in the plant to kill bacteria and to obtain increased metal yields, resulting in a high pH level of 10–12 in the process water. Different oxidation experiments (abiotic oxidation in untreated tailings, acid abiotic oxidation and acid microbial oxidation), containing the Boliden tailings, were performed at room temperature with dissolved oxygen (0.21 atm O₂) for 3 months. The different pyrite oxidation rates given from the study were 2.4×10⁻¹⁰ mol m⁻² s⁻¹ for the microbial, 5.9×10⁻¹¹ mol m⁻² s⁻¹ for the acidic abiotic and 3.6×10⁻¹¹ mol m⁻² s⁻¹ for the untreated experiments. Because of the potential precipitation of gypsum in the batch solutions, these oxidation rates are considered minimum values. The release rates for copper and zinc from chalcopyrite and sphalerite in the acid experiments were also investigated. These rates were normalized to the metal concentration in the tailings, and then compared to the release rate for iron from pyrite. These normalized results indicated that metal release decreased in the order Cu>Zn>Fe, demonstrating that pyrite is more resistant to oxidation than sphalerite and chalcopyrite. Pyrite was also more resistant to acidic dissolution than to microbial dissolution, while a significant fraction of sphalerite and chalcopyrite dissolved in the acid abiotic solutions.     

Formation of ferriolivine and magnesioferrite from Mg — Fe-olivine: Reactions and kinetics of oxidation

Olivine samples (Fa 11) have been oxidized in air (f _O2 = 0.2 atm) at temperatures ranging from 350–700 °C and examined by Mössbauer spectroscopy, transmission electron microscopy, X-ray powder diffraction and thermomagnetic analysis. Oxidation of olivine was found to result in ferriolivine, magnesioferrite (major oxide phase) and magnetite (minor oxide phase) formation. Ferriolivine forms planar (001) precipitates, 0.6 nm in thickness, in the olivine host; the composition is likely to be Mg_0.5 v _0.5(Fe³⁺)_1.0SiO₄. Magnesioferrite MgFe₂O₄ exsolves as fine-grained precipitates (5–6 nm in size) filling interstices between the ferriolivine planar precipitates. Oxidation kinetic data at 700 °C show two stages of oxidation corresponding to formation of ferriolivine in the first stage and magnesioferrite in the second stage. The linear rate law with a rate constant k _Fol = 1.23 · 10^-3 s^-1 was found for the first stage whereas a parabolic rate-law with a constant of k _oxi = 3.28 · 10^-3 s^-1 was determined for the second stage of oxidation. It was found that ferriolivine is not an intermediate metastable phase in the oxidation process, terminated by magnesioferrite formation. The ferriolivine and magnesioferrite are considered to have formed by independent reactions which do not necessarily proceed simultaneously.     

The solubility of BaCO3(cr) (witherite) in CO2-H2O solutions between 0 and 90°C,evaluation of the association constants of BaHCO3+(aq) and BaCO30(aq) between 5 and 80°C,and a preliminary evaluation of the thermodynamic properties of Ba2+(aq)

One hundred and fifty new measurements of the solubility of witherite were used to evaluate the equilibrium constant of the reaction BaCO₃(cr) = Ba²⁺(aq) + CO₃²⁻(aq) between 0 and 90°C and 1 atm total pressure. The temperature dependence of the equilibrium constant is given by logK = 607.642 + 0.121098T − 20011.25/T − 236.4948 logT where T is in degrees Kelvin. The logK of BaCO₃(cr), the Gibbs energy, the enthalpy and entropy of the reaction at 298.15 K are −8.562, 48.87 kJ · mol⁻¹, 2.94 kJ · mol⁻¹ and −154.0 J · mol⁻¹ · K⁻¹, respectively. The equilibrium constants are consistent with an aqueous model that includes the ion pairs BaHCO₃⁺(aq) and BaCO₃⁰(aq) Three different methods were used to evaluate the association constant of BaHCO₃⁺(aq), and all yielded similar results. The temperature dependence of the association constant for the reaction Ba²⁺(aq) + HCO₃⁻(aq) = BaHCO₃⁺(aq) is given by logK_BaHCO3⁺ = −3.0938 + 0.013669T.The log of the association constant, the Gibbs energy, the enthalpy and entropy of the reaction at 298.15°K are 0.982, −5.606 kJ · mol⁻¹, 23.26 kJ · mol⁻¹ and 96.8 J · mol⁻¹ · K⁻¹, respectively. The temperature dependence of the equilibrium constant for the reaction Ba²⁺(aq) + CO²⁻₃(aq) = BaCO₀³(aq) is given by logK_BaCO3⁰ = 0.113 + 0.008721T.The log of the association constant, the Gibbs energy, the enthalpy and entropy of the reaction at 298.15° K are 2.71, −15.49 kJ · mol⁻¹, 14.84 kJ · mol⁻¹ and 101.7 J· mol⁻¹ · K⁻¹.The above model leads to reliable calculations of the aqueous speciation and solubility of witherite in the system BaCO₃-CO₂-H₂O from 0 to more than 90°C. Literature data on witherite solubility were re-evaluated and compared with the results of this study.Problems in the thennodynamic selections of Ba compounds are considered. Newer data require the revision of Δ_fH° and Δ_fG° of Ba²⁺(aq) to −532.5 and −555.36 kJ · mol⁻¹, respectively, for agreement with solubility data.