Hydrogen-electrode Concentration Cell Research Articles

On the growth of anhydrous Mg-bearing carbonates – Implications from norsethite growth kinetics

At ambient conditions, both dolomite [CaMg(CO3)2] and magnesite [MgCO3] do not readily grow from aqueous solution. A common explanation for this is the highly hydrated character of Mg2+. The exceptionally easy growth of norsethite [BaMg(CO3)2], however, clearly shows that Mg2+ ions, in principle, can rapidly dehydrate and incorporate into anhydrous carbonate minerals even at ambient conditions. Still, the lack of reliable quantitative data prevents the necessary comparison of the reactivities of these magnesium-bearing minerals. In order to shed light on the span of possible incorporation rates of anhydrous Mg2+ ions, we present the first systematic quantitative study of norsethite growth kinetics as well as the determination of norsethite solubility product over a wide range of temperatures to replace the diffuse qualitative knowledge existing so far.Norsethite solubility was determined in 0.1 M NaCl aqueous solutions from 30 to 150 °C using a hydrogen-electrode concentration cell, which provides a continuous in-situ measurement of hydrogen ion molality. The solubility product of norsethite can be described by log10Ksp°-nrs = a + b/T + cT, where a = 31.007, b = −7321.122, and c = −0.0811. Gibbs free energy (ΔfG298.150) and enthalpy (ΔfH298.150) of norsethite formation were determined to be −2167 ± 2 kJ/mol and −2351 ± 2 kJ/mol, respectively. Growth experiments were conducted in mixed-flow reactors covering a significant span of solution compositions (pH: 7.0–8.5, [Ba]: 3 × 10−6–5 × 10−3 M, [Mg]: 1 × 10−4–9 × 10−2 M, ionic strength: 0.1 M, Ωnorsethite = 1–290) and temperatures (40, 65, and 100 °C). From the experimental data, the apparent activation energy of norsethite growth rate constant was determined to be Ea = 80 ± 7 kJ/mol. An extrapolation to 25 °C resulted in a rate constant of knrs25∘C = 1.8 × 10−2 nmol m−2 s−1 with a reaction order of 1.2 ± 0.1. These results allowed for a direct, quantitative comparison of the growth rates of different anhydrous Mg-bearing carbonate minerals. This comparison revealed that the growth rate constant of norsethite at 100 °C is approximately three orders of magnitude higher than that of magnesite and five orders of magnitude higher than that of dolomite.In the case of norsethite, obviously some effective means must exist which promotes the dehydration of the Mg ion and allows for the rapid incorporation of dehydrated Mg2+ into the growing mineral. This promotion has to take place at the norsethite surface where the hydration energy of Mg2+ can significantly differ from the well-known value in bulk solution. Consequently, not only the stability of the aqueous metal complex per se is an important factor controlling the growth rate of anhydrous magnesium bearing carbonate minerals but also the means of a given surface to weaken the stability of this complex.

Experimental determination of the solubility product of dolomite at 50–253 °C

The ‘dolomite problem’, the scarcity of present-day dolomite formation near the Earth’s surface, has attracted much attention over the past century. Solving this problem requires having reliable data on the stability and kinetics of formation of this mineral. Toward this goal, the solubility of natural dolomite (CaMg(CO3)2) has been measured from 50 to 253 °C in 0.1 mol/kg NaCl solutions using a hydrogen electrode concentration cell (HECC). The obtained apparent solubility products (Kapp-sp-dol), for the reaction: CaMg(CO3)2 = Ca2+ + Mg2+ + 2CO32−, were extrapolated to infinite dilution to generate the solubility product constants for this reaction (Ksp°-dol). The derived equilibrium constants were fit and can be accurately described by log10Ksp°-dol = a + b/T (K) + cT (K) where a = 17.502, b = −4220.119 and c = −0.0689. This equation and its first and second derivatives with respect to T were used together with corresponding aqueous species properties to calculate the revised standard state thermodynamic properties of dolomite at 25 °C and 1 bar, yielding a Gibbs energy of formation (ΔfG298.15∘) equal to −2160.9 ± 2 kJ/mol, (log10Ksp°-dol = −17.19 ± 0.3); an enthalpy of formation (ΔfH298.15∘) of −2323.1 ± 2 kJ/mol, an entropy (S298.15∘) of 156.9 ± 2 J/mol/K and heat capacity (Cp298.15∘) of 154.2 ± 2 J/mol/K (uncertainties are 3σ). The dolomite solubility product derived in this study is nearly identical to that computed using SUPCRT92 (Johnson et al., 1992) at 200 °C, but about one order of magnitude higher at 50 and 25 °C, suggesting that dolomite may be somewhat less stable than previously assumed at ambient temperatures.

Can accurate kinetic laws be created to describe chemical weathering?

Knowledge of the mechanisms and rates of mineral dissolution and growth, especially close to equilibrium, is essential for describing the temporal and spatial evolution of natural processes like weathering and its impact on CO2 budget and climate. The Surface Complexation approach (SC) combined with Transition State Theory (TST) provides an efficient framework for describing mineral dissolution over wide ranges of solution composition, chemical affinity, and temperature. There has been a large debate for several years, however, about the comparative merits of SC/TS versus classical growth theories for describing mineral dissolution and growth at near-to-equilibrium conditions. This study considers recent results obtained in our laboratory on oxides, hydroxides, silicates, and carbonates on near-equilibrium dissolution and growth via the combination of complementary microscopic and macroscopic techniques including hydrothermal atomic force microscopy, hydrogen-electrode concentration cell, mixed flow and batch reactors. Results show that the dissolution and precipitation of hydroxides, kaolinite, and hydromagnesite powders of relatively high BET surface area closely follow SC/TST rate laws with a linear dependence of both dissolution and growth rates on fluid saturation state (Ω) even at very close to equilibrium conditions (|ΔG|<500J/mol). This occurs because sufficient reactive sites (e.g. at kink, steps, and edges) are available at the exposed faces for dissolution and/or growth, allowing reactions to proceed via the direct and reversible detachment/attachment of reactants at the surface. In contrast, for magnesite and quartz, which have low surface areas, fewer active sites are available for growth and dissolution. Such minerals exhibit rates dependencies on Ω at near equilibrium conditions ranging from linear to highly non-linear functions of Ω, depending on the treatment of the crystals before the reaction. It follows that the form of the f(ΔG) function describing the growth and dissolution of minerals with low surface areas depends on the availability of reactive sites at the exposed faces and thus on the history of the mineral-fluid interaction and the hydrodynamic conditions under which the crystals are reacted. It is advocated that the crystal surface roughness could serve as a proxy of the density of reactive sites. The consequences of the different rate laws on the quantification of loess weathering along the Mississippi valley for the next one hundred years are examined.

Experimental determination of the solubility product of magnesite at 50 to 200 °C

Accurate knowledge of magnesite thermodynamic properties over a wide range of temperatures is crucial for characterizing mineral stabilities in the system MgO–CO 2–H 2O and for modeling carbon dioxide fate in important natural and industrial processes. However, available databases, especially for its solubility product, are sparse and contradictory, leading to considerable uncertainties in the calculation of chemical equilibria and phase transformations among carbonates. In this study, the solubility of synthetic magnesite was investigated from 50 to 200 °C in 0.1 mol kg − 1 NaCl solutions and in some cases under constant CO 2 partial pressure (4–30 bars) both by means of a hydrogen electrode concentration cell (HECC) and a traditional batch Ti-reactor. The obtained apparent solubility products ( Q sp-mgs) were extrapolated to infinite dilution to generate the solubility products ( K sp°-mgs), allowing calculation of the thermodynamic properties of magnesite. Of all the temperature functions tested, the equation giving a reliable fit of our data in the investigated temperature range (50–200 °C) has the following form: log 10 K sp°-mgs = a + b / T (K) + cT (K) with: a = 7.267, b = − 1476.604 and c = − 0.033918. Based on this equation and its first and second derivatives with respect to T, we were able to derive values at 25 °C, 1 bar for magnesite thermodynamic functions: Δ f G 298.15 o = (− 1026.48 ± 2) kJ mol − 1 (log 10 K sp°-mgs = − 7.80 ± 0.3), Δ f H 298.15 o = (− 1111.75 ± 2) kJ mol − 1 , S 298.15 o = (60.00 ± 2) J mol − 1 K − 1 , and C p 298.15 o = (75.91 ± 2) J mol − 1 K − 1 (uncertainties are 3 σ).

Solubility Measurements of Crystalline NiO in Aqueous Solution as a Function of Temperature and pH

Results of solubility experiments involving crystalline nickel oxide (bunsenite) in aqueous solutions are reported as functions of temperature (0 to 350 °C) and pH at pressures slightly exceeding (with one exception) saturation vapor pressure. These experiments were carried out in either flow-through reactors or a hydrogen-electrode concentration cell for mildly acidic to near neutral pH solutions. The results were treated successfully with a thermodynamic model incorporating only the unhydrolyzed aqueous nickel species (viz., Ni2+) and the neutrally charged hydrolyzed species (viz., $\mathrm{Ni(OH)}_{2}^{0})$ . The thermodynamic quantities obtained at 25 °C and infinite dilution are, with 2σ uncertainties: $\log_{10}K_{\mathrm{s0}}^{\mathrm{o}} = (12.40 \pm 0.29),\varDelta_{\mathrm{r}}G_{m}^{\mathrm{o}} = -(70. 8 \pm 1.7)$ kJ⋅mol−1; $\varDelta_{\mathrm{r}}H_{m}^{\mathrm{o}} = -(105.6 \pm 1.3)$ kJ⋅mol−1; $\varDelta_{\mathrm{r}}S_{m}^{\mathrm{o}} =-(116.6 \pm 3.2)$ J⋅K−1⋅mol−1; $\varDelta_{\mathrm{r}}C_{p,m}^{\mathrm{o}} = (0 \pm 13)$ J⋅K−1⋅mol−1; and $\log_{10}K_{\mathrm{s2}}^{\mathrm{o}} = -(8.76 \pm 0.15)$ ; $\varDelta_{\mathrm{r}}G_{m}^{\mathrm{o}} = (50.0 \pm 1.7)$ kJ⋅mol−1; $\varDelta_{\mathrm{r}}H_{m}^{\mathrm{o}} = (17.7 \pm 1.7)$ kJ⋅mol−1; $\varDelta_{\mathrm{r}}S_{m}^{\mathrm{o}} = -(108\pm 7)$ J⋅K−1⋅mol−1; $\varDelta_{\mathrm{r}}C_{p,m}^{\mathrm{o}} = -(108 \pm 3)$ J⋅K−1⋅mol−1. These results are internally consistent, but the latter set differs from those gleaned from previous studies recorded in the literature. The corresponding thermodynamic quantities for the formation of Ni2+ and $\mathrm{Ni(OH)}_{2}^{0}$ are also estimated. Moreover, the $\mathrm{Ni(OH)}_{3}^{ -}$ anion was never observed, even in relatively strong basic solutions ( $m_{\mathrm{OH}^{ -}} = 0.1$ mol⋅kg−1), contrary to the conclusions drawn from all but one previous study.

Solubility measurements of crystalline β-Ni(OH)2 in aqueous solution as a function of temperature and pH

Experimental results on the solubility of crystalline nickel hydroxide (theophrastite), according to the dissolution reaction, β-Ni(OH)2(cr) + 2H+ ⇄ Ni2+ + 2H2O(l), are reported as functions of temperature (0–200°C) and pH at pressures slightly exceeding saturation vapor pressure. The experiments were carried out in either various flow-through cells or a hydrogen-electrode concentration cell (HECC). The results were treated with a thermodynamic model incorporating only the unhydrolyzed aqueous nickel species (namely, Ni2+) whose solubility constants were fitted as a function of temperature. The thermodynamic quantities obtained at infinite dilution are: (25°C) = (11.67 ± 0.20), (25°C) = −(66.6 ± 1.1) kJ mol−1, (mean value 5–200°C) = −(82.1 ± 0.8) kJ mol−1, = (mean value 5–200°C) = −(52 ± 5) J K−1 mol−1, (mean value 5–200°C) = 0 J K−1 mol−1, where the uncertainties quoted are 2σ. These results are internally consistent, but differ from most of those reported in the literature. A significantly lower temperature for the thermodynamic stability of β-Ni(OH)2(cr) versus NiO(cr), namely, 77°C, is proposed based on the current measurements. Additional batch experiments at 50°C and 80°C within aqueous NaClO4 media established and quantified a particle-size dependence of the solubility constant of β-Ni(OH)2(cr) that extended to larger particle sizes than normally observed for metal oxide/hydroxides.

Effect of Amines on the Surface Charge Properties of Iron Oxides

Specific studies detailing the effects of amines, used as pH control agents for corrosion inhibition in power plants, on the surface charge of iron oxides provide data to assess the mechanism of how these amines impact deposition rate. The current study was undertaken in order to determine accurately the dissociation constants of the relevant amines at Pressurized Water Reactor (PWR) operating conditions and to investigate the effect of sorption of two of these amines (morpholine and dimethylamine) by magnetite. The acid-dissociation equilibria of morpholine (MOR), dimethylamine (DMA) and ethanolamine (ETA) were measured potentiometrically with a hydrogen-electrode concentration cell (HECC) from 0 to 290 °C in sodium trifluoromethanesulfonate (NaTr) solutions at ionic strengths up to 1 mol⋅kg−1. Magnetite surface titrations were performed at an ionic strength of 0.03 mol⋅kg−1 (NaTr medium) in the presence or absence of morpholine and dimethylamine buffers over a wide range of pH and total amine concentrations at 150–250 °C.

Solubility product of siderite (FeCO 3) as a function of temperature (25–250 °C)

The solubility of natural siderite was investigated from 25 to 250 °C in 0.1 mol kg − 1 NaCl aqueous solutions using a hydrogen-electrode concentration cell, which provided continuous in situ measurement of hydrogen ion molality. Iron(II) was analyzed by a revised Ferrozine-spectrophotometric method. The obtained apparent solubility products ( Q sp-siderite) were extrapolated to infinite dilution to generate the solubility products ( K sp°-siderite), allowing calculation of the thermodynamic properties of siderite. Of all the temperature functions tested, the equation giving a reliable fit of our data in the investigated temperature range (25–250 °C) has the following form: log 10 K sp°-siderite = a + b ∙ ( T / K ) + c ∙ ( T / K) − 1 + d ∙ log 10 (T / K) with: a = 175.568, b = 0.0139, c = − 6738.483 and d = − 67.898. Based on this equation and its first and second derivatives with respect to T, we were able to derive values for the Gibbs energy of formation: Δ fG 298.15 o = (− 680.71 ± 2) kJ mol − 1, enthalpy of formation: Δ fH 298.15 o = (− 749.59 ± 2) kJ mol − 1, entropy: S 298.15 o = (109.54 ± 2) J mol − 1 K − 1 and heat capacity: C p 298.15 o = (83.26 ± 2) J mol − 1 K − 1 of siderite (uncertainties are 3 σ). The values of Δ fG 298.15 o and Δ fH 298.15 o are in very good agreement with the values reported by Robie et al. [Robie, R.A., Haselton, H.T. Jr., Hemingway, B.S., 1984. Heat capacities and entropies of rhodochrosite (MnCO 3) and siderite (FeCO 3) between 5 and 600 K. Am. Mineral. 69, 349–357] and Chai and Navrotsky [Chai, L., Navrotsky, A., 1994. Enthalpy of formation of siderite and its application in phase equilibrium calculation. Am. Mineral. 79, 921–929], respectively. The density model [Anderson, G.M., Castet, S., Schott, J., Mesmer, R.E., 1991. The density model for estimation of thermodynamic parameters of reactions at high temperatures and pressures. Geochim. Cosmochim. Acta 55, 1769–1779] reproduced correctly our experimental data and allowed the extrapolation of the siderite solubility product up to 350 °C by using our values of the Gibbs energy and enthalpy of formation of siderite combined with its entropy and the heat capacity equation given by Robie et al. [Robie, R.A., Haselton, H.T. Jr., Hemingway, B.S., 1984. Heat capacities and entropies of rhodochrosite (MnCO3) and siderite (FeCO3) between 5 and 600 K. Am. Mineral. 69, 349–357].

Dissolution/precipitation kinetics of boehmite and gibbsite: Application of a pH-relaxation technique to study near-equilibrium rates

The dissolution and precipitation rates of boehmite, AlOOH, at 100.3°C and limited precipitation kinetics of gibbsite, Al(OH)3, at 50.0°C were measured in neutral to basic solutions at 0.1 molal ionic strength (NaCl+NaOH+NaAl(OH)4) near-equilibrium using a pH-jump technique with a hydrogen-electrode concentration cell. This approach allowed relatively rapid reactions to be studied from under- and over-saturation by continuous in situ pH monitoring after addition of basic or acidic titrant, respectively, to a pre-equilibrated, well-stirred suspension of the solid powder. The magnitude of each perturbation was kept small to maintain near-equilibrium conditions. For the case of boehmite, multiple pH-jumps at different starting pHs from over- and under-saturated solutions gave the same observed, first order rate constant consistent with the simple or elementary reaction: Al(OOH)(cr)+H2O(l)+OH-⇄Al(OH)4-.This relaxation technique allowed us to apply a steady-state approximation to the change in aluminum concentration within the overall principle of detailed balancing and gave a resulting mean rate constant, (2.2±0.3)×10−5kgm−2s−1, corresponding to a 1σ uncertainty of 15%, in good agreement with those obtained from the traditional approach of considering the rate of reaction as a function of saturation index. Using the more traditional treatment, all dissolution and precipitation data for boehmite at 100.3°C were found to follow closely the simple rate expression:Rnet,boehmite=10-5.485{mOH-}{1-exp(ΔGr/RT)}, with Rnet in units of molm−2s−1. This is consistent with Transition State Theory for a reversible elementary reaction that is first order in OH− concentration involving a single critical activated complex. The relationship applies over the experimental ΔGr range of 0.4–5.5kJmol−1 for precipitation and −0.1 to −1.9kJmol−1 for dissolution, and the pHm≡−log(mH+) range of 6–9.6. The gibbsite precipitation data at 50°C could also be treated adequately with the same model:Rnet,gibbsite=10-5.86{mOH-}{1-exp(ΔGr/RT)}, over a more limited experimental range of ΔGr (0.7–3.7kJmol−1) and pHm (8.2–9.7).

Potentiometric Study of the Dissociation Quotient of the Aqueous Ethanolammonium Ion as a Function of Temperature and Ionic Strength

The acid dissociation equilibrium involving aqueous ethanolammonium ion in the reaction, HO(CH2)2NH3+ ⇄ HO(CH2)2NH2 + H+, was measured potentiometrically with a hydrogen-electrode concentration cell from (0 to 290) °C in sodium trifluoromethanesulfonate (NaTr) solutions at ionic strengths of (0.03, 0.1, and 0.3) molal. The molal dissociation quotients and selected literature data at infinite dilution were fitted by an empirical equation with five adjustable parameters involving functions of temperature, solvent density, and ionic strength. This treatment yielded the following thermodynamic quantities at 25 °C and infinite dilution: log Ka = −9.48 ± 0.03, ΔHa = (51.3 ± 1.5) kJ·mol-1, ΔSa = (−9 ± 5) J·K-1·mol-1, ΔCpa = (8 ± 17) J·K-1·mol-1.

Potentiometric studies of the rutile–water interface: hydrogen-electrode concentration-cell versus glass-electrode titrations

Abstract This paper represents a comparison of surface protonation studies of rutile in NaCl media obtained using a conventional glass-electrode autotitrator system from 10 to 50 °C, and hydrogen-electrode concentration cells from 25 to 250 °C [J. Colloid Interface Sci., 200 (1998) 298]. Experimental conditions were matched as closely as possible between the two techniques, permitting a direct comparison of the results. Values for the pH of zero net proton charge (pH znpc ) of the rutile surface obtained at 10 and 35 °C were consistent with the temperature trends observed previously using hydrogen-electrode concentration cells. The pH znpc of rutile decreases systematically from 5.7 to 4.2 as temperature increases from 10 to 250 °C. Moreover, the experimentally determined pH znpc values agree with independent estimates of the pH at the pristine point of zero charge (pH znpc ) calculated from an extension of the revised Multi-Site Complexation (MUSIC) Model of Hiemstra et al. [J. Colloid Interface Sci., 184 (1996) 680]. Surface protonation curves obtained from the glass-electrode titration results were rationalized using surface protonation constants derived from the MUSIC Model, in conjunction with a Basic Stern representation of the electrical double layer (EDL) structure. Best-fit parameters (Stern layer capacitance values, and electrolyte cation and anion binding constants) are consistent with those obtained from fits to titration curves obtained using hydrogen-electrode concentration cells at 25 and 50 °C. Consequently, this comparison demonstrates that independent conventional glass-electrode and hydrogen-electrode concentration cell titrations provide completely compatible results despite the intrinsic differences in the two techniques (pH calibration, equilibration times, stirring rates, gas phase composition, etc).

New Measurements of the Solubility of Zinc Oxide from 150 to 350°C

The solubility of zincite (ZnO) has been re-evaluated in noncomplexing solutions over a wide range of pH and from 150 to 350°C at pressures ranging from slightly above saturation vapor pressure to significantly higher pressures in NaOH, NH3, F3CSO3H/F3CSO3 Na (HTr/NaTr), CH3COOH/NaOH, and HTr/NH3 solutions using a hydrogen-electrode concentration cell (HECC) and a flow-through cell with downstream acid injection. The new results, coupled with our earlier results at 75–100°C, indicate that the solubility of zinc oxide is higher than we recently reported,(1) but substantially lower than previous estimates available in the scientific literature. It is believed that deposition of ZnO or a zinc hydroxide phase must have occurred near the solubility minimum in the HECC sample lines in our earlier study; a problem that is overcome in the flow cell by injecting acid into the sample line before the solution is depressurized and cooled to room temperature.

Formation Quotients of Aluminum Sulfate Complexes in NaCF3SO3 Media at 10, 25, and 50 °C from Potentiometric Titrations Using a Mercury/Mercurous Sulfate Electrode Concentration Cell

Mercury/mercurous sulfate electrode concentration cells (with liquid junction) are employed in this study to determine the formation constants of aluminum sulfate complexes, with the principal advantage that the change in the free sulfate concentration is measured directly without the need to know the standard potential of the electrode. Potentiometric titrations were conducted at temperatures of 10, 25, and 50 degrees C and ionic strengths of approximately 0.3, 0.5, and 1.0 molal in aqueous solutions of the inert 1:1 electrolyte sodium trifluoromethanesulfonate (NaTr). Stoichiometric molal formation quotients Q1 and Q2, respectively, for the reactions Al3+(aq) + SO4(2-)(aq) AlSO4+(aq) and Al3+(aq) + 2SO4(2-)(aq) <= >Al(SO4)2-(aq) were determined. The values of log Q1 obtained from this work in NaTr media at ionic strengths of 0.3 and 1.0 mol x kg(-1) and 50 degrees C (1.72 +/- 0.08 and 1.35 +/- 0.06, respectively) are in excellent agreement with the values (1.71 +/- 0.2 and 1.32 +/- 0.1) determined in NaCl media from the recent potentiometric study conducted in the same laboratory using a hydrogen electrode concentration cell by Ridley et al. (Ridley, M. K.; Wesolowski, D. J.; Palmer, D. A.; Kettler, R. M. Geochim. Cosmochim. Acta 1999, 62, 459-472). The value of log Q2 (2.05 +/- 0.05) in 1.0 mol x kg(-1) from this work is smaller than the value reported by Ridley et al. (2.6 +/- 0.5) but within the combined experimental error. Empirical isothermal equations are presented to permit calculation of the equilibrium quotients as a function of ionic strength (0-1 mol x kg(-1)), giving log K1 and log K2 values at 25 degrees C and infinite dilution of 3.84 +/- 0.12 and 5.58 + 0.09, respectively. The value for log K1 obtained in this study at 25 degrees C is bracketed within experimental uncertainty by values reported by Kryzhanovskii et al. (Kryzhanovskii, M. M.; Volokhov, Y. A.; Pavlov, L. N.; Eremin, N. I.; Mironov, V. E Zh. Prikl. Khim. 1971, 44,476-479) and Nishide and Tsuchiya (Nishide, T.; Tsuchiya, R. Bull. Chem. Soc. Jpn. 1965, 38, 1398-1400), namely, 3.89 and 3.73, respectively. All other literature values for the first aluminum sulfate association constant are considerably lower than these results, which is also true for the second association constant, although there are few experimental data available for the latter. Empirical equations are also presented for calculating values of log Q1 and log Q2 from 0 to 1 molal ionic strength and from 10 to 125 degrees C, spanning the range of most environmental conditions at which these reactions are important.

Measurement of bacterial surface protonation constants for two species at elevated temperatures

This study reports the first potentiometric titrations of bacterial surfaces at elevated temperatures. We used a hydrogen electrode concentration cell to examine the protonation behavior of Bacillus subtilis (an aerobic, gram-positive species present in a wide variety of low-temperature, near-surface environments) at 30, 50, and 75°C. We also investigate the protonation of TOR-39 (an anaerobic, Fe-reducing, thermophylic bacteria isolated from a deep sedimentary basin) at 50°C and compare its surface properties to those of B.subtilis. Using a surface complexation approach that has been previously applied only to room temperature experiments involving bacteria, we determine the number of functional group types present on each cell wall, their acidity constants, and their site concentrations. We also characterize the temperature dependence of the acidity constants for B.subtilis. Our results indicate that the surface acidity can be successfully described using site-specific reactions involving discrete surface functional groups.Our results indicate that the acidity constants for the surface functional groups on B. subtilis do not change significantly over the temperature interval studied, and that the surface protonation at each temperature can be estimated using a single set of acidity constants and site densities. The best fitting model for B. subtilis is a three- site model with pKa values of 4.1, 5.8, and 7.7. Average site abundances for the three dominant functional groups on B. subtilis were determined to be 9.5 · 10−5, 1.0 · 10−4, and 8.8 · 10−5 moles of sites/gm. For TOR-39, the best fitting model was also a 3-site model with pKa values of 4.5, 5.8 and 8.2, with site abundances of 6.0 · 10−5, 5.0 · 10−5 and 3.0 · 10−5. This study suggests that bacterial surface site protonation reactions are not strongly temperature dependent, at least to 75°C. The lack of significant temperature dependence could greatly simplify the task of modeling bacteria-water-rock interactions at elevated temperature.

Aqueous high-temperature solubility studies. II. The solubility of boehmite at 0.03 m ionic strength as a function of temperature and pH as determined by in situ measurements

The solubility of pure synthetic boehmite (γ-AlOOH) has been measured over a wide range of pH (2–10, depending on temperature), and temperature (100–290°C) at 0.03 mol · kg −1 ionic strength (NaCl) in a hydrogen-electrode concentration cell (HECC), which provided continuous in situ measurement of hydrogen ion molality (Palmer et al. 2001). A least-squares regression of the results obtained was used to determine the molal solubility quotients (Q s0 to Q s4) of boehmite. The solubility products (Q sn) were extrapolated to infinite dilution (K sn), permitting calculation of the thermodynamic properties of aqueous species of aluminum for comparison with previous works, such as Bourcier et al. (1993) and Castet et al. (1993). These results are generally consistent with the latter study conducted near infinite dilution, although some significant differences are apparent, particularly near the solubility minimum.

Aqueous high-temperature solubility studies. I. The solubility of boehmite as functions of ionic strength (to 5 molal, NaCl), temperature (100–290°C), and pH as determined by in situ measurements

The solubility of boehmite, γ-AlOOH, was measured in a modified hydrogen-electrode concentration cell, which provided continuous in situ measurements of hydrogen ion molality over the range of ionic strengths from 0.1 to 5.0 mol · kg −1 (NaCl) at temperatures from 100 to 290°C. A series of conventional solubility measurements was also carried out in acidic solutions over the same temperature range (i.e., pH was not monitored, but rather calculated from mass balance). The combined results yielded the molal solubility quotients, Q s0 and Q s4 for the equilibria: AlOOH (cr) + 3H + ⇄ Al 3+ + 2H 2O Q s0 = [Al 3+]/[H +] 3 AlOOH (cr) + 2H 2O ⇄ Al(OH) 4 − + H + Q s4 = [Al(OH) 4 −][H +]In the regression of each isothermal data set, the values for the first hydrolysis quotient, Al 3+ + H 2O ⇄ Al(OH) 2+ + H + Q 1,1 = [Al(OH) 2+][H +]/[Al 3+]were fixed according to a previous potentiometric study (Palmer and Wesolowski, 1993). Moreover, for one series of titrations at 0.1 mol · kg −1 ionic strength at 150°C, the remaining two solubility quotients, Q s2 and Q s3, were determined simultaneously from the regression. However, at all other conditions, the values of Q s2 and Q s3 were also fixed at values consistent with the corresponding 0.03 mol · kg −1 ionic strength results (Bénézeth et al., 2001) by invoking the isocoulombic assumption (i.e., an assumption of minimal ionic strength dependence for reactions with no net change in charge). The stability field with respect to pH of the Al(OH) 2 + and Al(OH) 3 0 aqueous species were found to be very narrow, and hence assumptions concerning their stabilities had little effect on the predicted shape of the solubility profiles at high ionic strength. Global fits were made of the log Q s0 and log Q s4 values as functions of temperature and ionic strength after combining with corresponding values from an analogous study at 0.03 mol · kg −1 ionic strength (Bénézeth et al., 2001), as well as appropriate constants taken from the literature. The fits were further constrained by inclusion of gibbsite solubility data (<100°C) in NaCl media (Wesolowski, 1992; Palmer and Wesolowski, 1992) after adjustment for the relative free energies of formation of gibbsite and boehmite. This treatment ensured that a continuous empirical model exists for Al 3+ and Al(OH) 4 − from ambient conditions to 300°C and infinite dilution to five mol · kg −1 ionic strength.

Potentiometric Study of the Dissociation Quotients of Aqueous Dimethylammonium Ion As a Function of Temperature and Ionic Strength

The acid dissociation equilibria involving aqueous dimethylammonium ion in the reaction (CH3)2NH2+ ⇄ (CH3)2NH + H+ were measured potentiometrically with a hydrogen-electrode concentration cell from (0 to 290) °C in sodium trifluoromethanesulfonate (NaTr) solutions at ionic strengths of (0.1, 0.3, and 1) molal. The molal dissociation quotients and selected literature data at infinite dilution were fitted by an empirical equation involving six adjustable parameters involving functions of temperature, solvent density, and ionic strength. This treatment yielded the following thermodynamic quantities at 25 °C and infinite dilution: log Kd = −10.77 ± 0.02, ΔHd = (50.8 ± 0.7) kJ·mol-1, ΔSd = (−35.8 ± 2.5) J·K-1·mol-1, ΔCp,d = (116 ± 11) J·K-1·mol-1, and ΔVd = (−4.3 ± 2.5) cm3·mol-1.

Potentiometric determination of cadmium–acetate complexation in aqueous solutions to 250°C

The molal formation quotients of cadmium–acetate complexes were measured potentiometrically (hydrogen-electrode concentration cell) from 50 to 250°C at ionic strengths of 0.1, 0.3, and 1.0 molal in aqueous sodium trifluoromethanesulfonate (NaTr) media. Two cadmium–acetate species, namely Cd(Ac) +and Cd(Ac) 2°, were identified by regressional analysis of the data. Their formation quotients were modeled by empirical equations to describe their temperature and ionic strength dependencies. The thermodynamic quantities, obtained by differentiating these equations with respect to temperature at 25°C and at infinite dilution are for Cd(Ac) +: log K 1=1.96±0.24, Δ H°=−17±17 kJ·mol −1, Δ S°=−21±53 J·K −1·mol −1 and ΔCp°=660±370 J·K −1·mol −1; and for Cd(Ac) 2°: log K 2=3.15±0.10, Δ H°=4±14 kJ·mol −1, Δ S°=75±47 J·K −1·mol −1 and ΔCp°=−400±260 J·K −1·mol −1.

Magnetite surface charge studies to 290°C from in situ pH titrations

The proton-induced surface charge of magnetite was investigated in 0.03 and 0.30 molal sodium trifluoromethanesulfonate solutions from 25°C to 290°C by potentiometric titrations using a stirred hydrogen electrode concentration cell. Pure magnetite with excellent crystallinity was produced by reaction with the Ni/NiO/H 2O hydrogen fugacity buffer at 500°C. Inflection points in the 0.03 molal proton sorption isotherms (pH infl) at 6.50, 6.24, 5.65, 5.47, 5.31 and 5.55 at temperatures of 50°C, 100°C, 150°C, 200°C, 250°C and 290°C, respectively, were used as estimates of the pristine point of zero charge (pH ppzc) for modeling purposes. These pH infl values parallel 1/2 p K w and agree within the assigned uncertainty (±0.3 pH units) at all temperatures with independent estimates of the pH ppzc calculated from an extension of 88the revised MUSIC model. The surface charging can be adequately described by a one-p K model with a surface protonation constant fitted to the pH infl values, and giving the standard state thermodynamic properties log K H,298=7.00, Δ H 298°=−32.4±0.8 kJ/mol and constant Δ C p=128±16 J K −1 mol −1, with Δ S 298° assumed to be equal to that of rutile protonation (25.5±3.4 J K −1 mol −1. The 0.03 and 0.30 molal proton sorption isotherms also exhibit pHs of common intersection (pH cip) at 6.33, 5.78, 5.37, 4.82, 4.62 and 4.90 at 50°C, 100°C, 150°C, 200°C, 250°C and 290°C, respectively. The difference between the pH cip and pH ppzc≅pH infl values can be related to specific binding of Na + on the negatively charged surface, which increases with increasing temperature, although the pH cip values may also be affected by dissolution of the solid. The electrical double layer model includes a basic Stern layer capacitance, with specific cation and anion binding at the Stern layer, and a fixed diffuse layer capacitance computed from Guoy–Chapman theory. To fit the steepness and asymmetry of the charging curves above the pH ppzc, an additional cation binding constant was invoked, which allows the cation to experience the surface potential. Significant kinetically controlled dissolution of magnetite was observed below the pH ppzc, which may be a result of leaching of Fe 2+ from the surface, to produce a magnetite+hematite assemblage, despite the high hydrogen partial pressures (ca. 10 bars) used in these experiments.

The solubility of zinc oxide in 0.03 m NaTr as a function of temperature, with in situ pH measurement

The solubility of zincite (ZnO) has been measured in noncomplexing solutions over a wide range of pH m (4–11), and temperature (75–200°C) at 0.03 mol · kg −1 ionic strength in NaTr media (sodium trifluoromethanesulfonate, a noncomplexing 1:1 electrolyte), in a hydrogen electrode concentration cell (HECC), which provided continuous in situ measurement of hydrogen ion molality. Total zinc content was analyzed by atomic absorption using graphite furnace, flame, and inductively coupled plasma (ICP) spectrometers. The direction of approach to the equilibrium saturation state was varied to demonstrate that the system was reversible thermodynamically. Separate experiments were performed in alkaline solutions (0.03 mol · kg −1 NaOH) at 25 and 50°C in polypropylene syringes, and between 50 and 290°C in a Teflon-lined pressure vessel. The aim of these experiments was to reach higher pH m (>8 depending on the temperature) to determine the thermodynamic properties of the negatively charged species, Zn(OH) 3 −. A least-squares regression of the results obtained at this ionic strength was used to determine the molal solubility products (Q sn) of zincite. The solubility products (Q sn) were extrapolated to infinite dilution (K sn), permitting calculation of the thermodynamic properties of aqueous species of zinc for comparison with previous work.