Peter A. Rock

Dissolution Kinetics of Calcite in 0.1 M NaCl Solution at Room Temperature: An Atomic Force Microscopic (AFM) Study

Atomic force microscopy (AFM) was used to study the rates of migration of the (10¯1 4) plane of a single-crystal of calcite dissolving in 0.1 M NaCl aqueous solutions at room temperature. The solution pH and PCO2 controlled in the ranges 4.4 < pH < 12.2 and 0 < PCO2 < 10-3.5 atm (ambient), respectively. Measured step velocities were compared with the mineral dissolution rates determined from the calcium fluxes. The step velocity is defined as the average of the velocities of the obtuse and acute steps. Rates of step motion increased gradually from 1.4(±0.2) at pH 5.3 to 2.4(±0.3) nm s-1 at pH 8.2, whereas the rates inverted and decreased to the minimum value of 0.69(±0.18) nm s-1 at pH 10.8. For pH > 10.8, only the velocity of the obtuse steps increased as pH increased, whereas that of acute steps gradually decreased.The dissolution rate of the mineral can be calculated from the measured step velocities and average slope, which is proportional to the concentration of exposed monomolecular steps on the surface. The average slope of the dissolving mineral, measured at pH 5.6 and 9.7, was 0.026 (±0.015). Using this slope, we calculate bulk dissolution rates for 5.3 < pH < 12.2 of 4.9(±3.0) × 10-11 to 1.8(±1.0) × 10-10 mol cm-2 s-1. The obtained dissolution rate can be expressed by the following empirical equation:Rdss = 10-4.66(±0.13)[H+] + 10-3.87(±0.06)[HCO3-] + 10-7.99(plusmn; 0.08)[OH-]We propose that calcite dissolution in these solutions is controlled by elementary reactions that are similar to those that control the dissolution of other amphoteric solids, such as oxides. The mechanisms include the proton-enhanced hydration and detachment of calcium-carbonate ion pairs. The detachments are enhanced by the presence of adsorbed nucleophiles, such as hydroxyl and bicarbonate ions, and by protons adsorbed to key oxygens. A molecular model is proposed that illustrates these processes.

Thermochemistry of mixing strontianite [SrCO3(s)] and aragonite [CaCO3(s)] to form CaxSr1−xCO3(s) solid solutions

Abstract Enthalpies of mixing [Δ H 298 mix ] of aragonite and strontianite to form Ca x Sr 1− x CO 3 (s) solid solutions were obtained at 298 K from drop-solution enthalpies [Δ H 975 ds ] of pellets of the respective solid solutions into molten 2PbO · B 2 O 3 at 975 K. The measured Δ H 298 mix values are positive for all measured values of x , are nearly symmetric around x = 0.50, and reach a maximum value of +3.82 ± 0.94 kJ mol −1 . Previous electrochemical studies have reported that Δ G 298 ex values are also positive over the range 0.0 x −1 at x ≈ 0.7. The general similarity between the Δ H 298 mix and the Δ G 298 ex values indicates that the excess entropy of mixing is small or zero, consistent with the regular-solution treatment. Within this regular-solution treatment, the interaction parameter is W = 13.5(±1.3) kJ mol −1 , which yields a very narrow range of stable miscible compositions at Earth surface conditions. Compositions of aragonite or strontianite with even a few percent impurity are not stable and will unmix to form a mechanical mixture of Ca-rich strontianite and a Sr-rich aragonite. It is, in general, difficult to accurately estimate the compositions of coexisting, miscible solids such as the Ca x Sr 1− x CO 3 solids at Earth-surface conditions. The predictions are quite sensitive to small uncertainties in the data and such estimates are probably beyond the capabilities of either electrochemical cell measurements or calorimetry.

A comparison of metal attenuation in mine residue and overburden material from an abandoned copper mine

Abstract The metal attenuation capacities of secondary acid mine water precipitates is dependent upon such factors as pH, ionic strength, the presence of competing ions, and tailings mineralogy. At the abandoned Spenceville Cu mine in Nevada County, California, approximately 6800 m 3 of jarosite overburden and 28,000 m 3 of hematite residue are potential sources of heavy metals loading to infiltrating surface waters. A column study was performed to assess the ability of the overburden and the residue to attenuate heavy metals from acidic mine drainage. The study information was needed as part of a remedial design for the abandoned mine, and was designed to simulate a worst-case scenario to examine the plausibility of backfilling a large open pit with the waste materials. Ten pore volumes of acidic mine drainage were allowed to pass through the materials, and the column effluents were analyzed for dissolved Fe, Al, Ca, Mg, Na, K, Mn, Cu, Zn, Pb and Ni using ICP-AES. The oxidation-reduction potential ( Eh ) was measured with a combination Pt Ag/AgCl electrode and also calculated from Fe(II) and Fe(III) measurements using the Nernst equation. Ion activities in solution and saturation index (SI) values for various solid phases were calculated using the geochemical speciation model MINTEQA2, and mineralogical compositions of fine ( 2 mm) fractions were determined by XRD. Geochemical modeling of the column effluent compositions indicate that goethite, jarosite, jurbanite and gypsum are potential solid phases that may control metal solubilities in the column effluents. Excellent agreement was observed between the measured Eh values and those calculated from the activity ratio of Fe 2+ (aq) to Fe 3+ (aq). The large attenuation capacities for Cu and Zn exhibited by the jarosite overburden also suggest that solid solution substitution plays a large role in controlling metal concentrations in the pore waters. Relatively little metal attenuation, however, was provided by the hematite residue.

Gibbs energies of formation for hydrocerussite [Pb(OH) 2 . (PbCO 3 ) 2 (S)] and hydrozincite [Zn(OH) 2 ] 3 . (ZnCO 3 ) 2 (S)] at 298 K and 1 bar from electrochemical cell measurements

Abstract New values are reported for the Gibbs energies of formation from the elements for hydrocerussite Pb(OH)2 · (PbCO3)2 and hydrozincite [Zn(OH)2]3 · (ZnCO3)2. These ΔG0f values were obtained from electrochemical cells without liquid junction. We determined ΔG0f [Pb(OH)2 · (PbCO3)2(S)] = -1699.8 ± 1.6 kJ/mol for hydrocerussite and ΔG0f ([Zn(OH)2]3 · (ZnCO3)2) = -3163.3 ±4 kJ/mol for hydrozincite. These results allow future electrochemical cell experiments to be performed to determine the ΔG0f values of other hydroxycarbonate minerals using either the Pb amalgam-hydrocerussite or the Zn amalgam-hydrozincite as reference electrodes. These reference electrodes provide a strategy for establishing Gibbs energies for phases with two different anions, which are geochemically interesting but difficult to study experimentally.

A new method for determining Gibbs energies of formation of metal-carbonate solid solutions: 1. The CaxCd1−xCO3(s) system at 298 K and 1 bar

New thermodynamic data are presented for mixing of calcite (CaCO[sub 3](s)) and otavite (CdCO[sub 3](s)) to form Ca[sub x]Cd[sub 1-x]CO[sub 3](s) solutions. A reversible equilibrium was achieved between the aqueous solution and the metastable solid solution using electrochemical double cells and a new electrode of the third kind. Evidence that the cell reactions are reversible includes demonstration of Nernstian response to changes in electrolyte composition and an absence of significant hysteresis in cell current-voltage and current-temperature (10-35[degrees]C) plots. The authors observe negative excess Gibbs energies for mixing on the order of 0 to -6.6 kJ[center dot]mol[sup [minus]1]; these magnitudes are sensitive to uncertainties in the estimates of [Delta]G[sub f][degrees][CdCO[sub 3](s)]. The double-cell method is potentially useful in other systems where, through incorporation of impurities, the activity of the CaCO[sub 3](s) is not unity relative to the standard state of the pure, well-crystallized, endmember mineral. It remains to be seen whether the double cell method is equally useful for systems, such as the Ca[sub x]Mg[sub 1-x]CO[sub 3](s) or Ca[sub x]Sr[sub 1-x]CO[sub 3](s) solid solutions, that have variable and positive excess Gibbs energies, but some of the same advantages would apply.

Thermodynamics of Lithium Isotope Exchange Reactions. II. Electrochemical Investigations in Diglyme and Propylene Carbonate

This paper reports our final results for the direct experimental determination of the equilibrium constants for the lithium isotope exchange reactions 7Li(s)+6LiBr(digl)=6Li(s)+7LiBr(digl) 7Li(s)+6Li+(PC)=6Li(s)+7Li+(PC) involving the solvents diglyme (digl, e=7.2) and propylene carbonate (PC, e=65.5). These reactions were studied in electrochemical double cells without liquid junction. The experimentally determined values of the equilibrium constants at 296.6°K are K1=1.035± 0.007 and K2=1.030± 0.005. The experimental values of K1 and K2 are compared with values calculated from statistical thermodynamic theory for various model reactions. A Born—von Karman lattice calculation has been carried out for the isotopic lithium metals.

Lattice energies of calcite-structure metal carbonates II. Results for CaCO3, CdCO3, FeCO3, MgCO3, and MnCO3

Abstract Lattice energies ( ΔU ) of five calcite-structure metal carbonates (i.e., CaCO3 3(s) , CdCO 3(s) , FeCO 3(s) , MgCO 3(s) , and MnCO 3(s) ) have been calculated theoretically by separating ΔU into the following three energy components: electrostatic, polarization, and repulsive. Experimental crystal structure data and previously determined carbonate ion charge distributions (Mandell, G. K. and Rock, P. A., Lattice energies of calcitestructure metal carbonates I. Calculation of carbonate ion charge distributions from oxygen polarizabilities. J. Phys. Chem. Solids , 1998, 59 , 695) serve as the primary input data to the model. The electrostatic portion of the lattice energy ( ΔU electrostatic ) is evaluated by a point charge approach, while the polarization energy contribution ( ΔU polarization) is computed via a self-consistent method, which considers both dipole/dipole and ion/dipole interactions. Lastly, the repulsive energy segment ( ΔU repulsive ) is evaluated by means of the Slater potential (Slater, J. C., Introduction to Chemical Physics , Ch. 13. McGraw-Hill, New York, 1939). The ΔU values obtained from this theoretical approach compare favorably with the corresponding experimentally based values.

Raman spectra in the libration region of concentrated aqueous solutions of lithium-6 and lithium-7 halides. Evidence for tetrahedral Li(OH2)4+

The Raman spectra of saturated solutions of6LiCl and7LiCl have been decomposed into Gaussian components, one of which is a polarized band that occurs at 360 cm−1 when the ion is6Li+ and shifts to 335 cm−1 when the ion is7Li+. Equivalent bands occur in the spectra of saturated solutions of6LiBr and7LiBr at 343 and 320 cm−1, respectively. These bands are assigned to solvent-separated ion aggregates. The Raman spectra of 8.0 and 3.5 m solutions of the isotopic lithium chlorides have been decomposed into five Gaussian components, three of which are assigned to water librations. In addition, there is a polarized band at 440 cm−1 independent of the lithium isotope used, and a depolarized band which occurs at 385 cm−1 in the6LiCl solutions and 360 cm−1 in the7LiCl solutions. We interpret these two additional bands as theA1 andF2 stretching modes of Li+ tetrahedrally solvated by water molecules.

THERMODYNAMICS OF LITHIUM-ISOTOPE-EXCHANGE REACTIONS. III. ELECTROCHEMICAL STUDIES OF EXCHANGE BETWEEN ISOTOPIC METALS AND AQUEOUS IONS.

This paper reports our results for the direct experimental determination of the equilibrium constant for the lithium‐isotope‐exchange reaction 7Li(s)+6LiCl(aq)=6Li(s)+7LiCl(aq). This reaction was studied in an electrochemical quadruple cell without liquid junction of the type 7Li(s)|7LiBr(PC)| 7Li(Hg)|7LiCl(aq)| Hg2Cl2 (s)| Hg(l)| Hg2Cl2 (s)|6LiCl(aq)|6Li(Hg)|6LiBr(PC)|6Li(s), where PC denotes the solvent propylene carbonate. The experimental value of the equilibrium constant at 296.6°K for this reaction is K = 1.046±0.013. The experimental value of K is compared with values for K calculated from statistical thermodynamic theory for various model reactions. The results of these calculations are consistent with a tetrahedrally coordinated structure for the aqueous lithium ions.

Standard Gibbs energies of formation of ZnC2O4 · 2H2O(s), CdC2O4 · 3H2O(s), Hg2C2O4(s), and PbC2O4(s) at 298 K and 1 bar

Abstract We test the hypothesis that cycling of some heavy metals in soils and aquifers is affected by equilibrium with oxalate solids. The hypothesis was tested using electrochemical cells that allow us to determine Δ G f ∘[PbC 2 O 4 (s)], Δ G f ∘[CdC 2 O 4 · 3H 2 O(s)], Δ G f ∘[ZnC 2 O 4 · 2H 2 O(s)], and Δ G f ∘[Hg 2 C 2 O 4 (s)] and hence, the solubilities in soil solutions. The approach was stepwise. First, reversible equilibria was achieved using an electrochemical cell: Pb(Hg),2-phase ¦PbC 2 O 4 (s), CaC 2 O 4 · H 2 O(s)¦CaCl 2 (aq,m)¦Hg 2 Cl 2 (s)¦ Hg(1), in order to calculate Δ G f ∘[PbC 2 O 4 (s)] from the relatively well-known value of G f ∘[CaC 2 O 4 · H 2 O(s)]. This value of Δ G f ∘[PbC 2 O 4 (s)] then allowed us to obtain values of Δ G f ∘[CdC 2 O 4 · 3H 2 O(s)], Δ G f ∘[ZnC 2 O 4 · 2H 2 O(s)], and Δ G f ∘[Hg 2 C 2 O 4 (s)] from suitable electrochemical cells. When these values of Δ G f ∘ are incorporated into multicomponent speciation calculations, we find that CdC 2 O 4 · 3H 2 O(s), Hg 2 C 2 O 4 (s), and ZnC 2 O 4 · 2H 2 O(s) are unlikely to form except in highly contaminated soils, or in intercellular environments where the concentration of dissolved oxalate is very high. Of these heavy-metal-oxalate minerals, it is conceivable that the PbC 2 O 4 (s) may reach equilibrium in soil solutions. Although the metal-oxalate solids may precipitate locally in the rhizosphere, these solids would not be in equilibrium with the adjacent soil solution.