The effect of pressure on the solubility of minerals in water and seawater     

Frank J Millero 《Geochimica et cosmochimica acta》1982,46(1):11-22

The effect of presure on the solubility of minerals in water and seawater can be estimated from In where the volume (ΔV) and compressibility (ΔK) changes at atmospheric pressure (P = 0) are given by Values of the partial molal volume ($\text{V}\text{?}$) and compressibilty ($\text{K}\text{?}$) in water and seawater have been tabulated for some ions from 0 to 50°C. The compressibility change is quite large (~10 × 10^?3 cm³ bar^?1 mol^?1) for the solubility of most minerals. This large compressibility change accounts for the large differences observed between values of ΔV obtained from linear plots of In K_sp versus P and molal volume data (Macdonald and North, 1974; North, 1974). Calculated values of $\text{K}{}^{\mathrm{P}}{}_{\mathrm{sp}}\text{K}{}^{\mathrm{o}}{}_{\mathrm{sp}}$ for the solubility of CaCO₃, SrSO₄ and CaF₂ in water were found to be in good agreement with direct measurements (Macdonald and North, 1974). Similar calculations for the solubility of minerals in seawater are also in good agreement with direct measurements (Ingle, 1975) providing that the surface of the solid phase is not appreciably altered.  相似文献   

Constantes de formation des complexes hydroxydés de l'aluminium en solution aqueuse de 20 a 70°C     

Yves Couturier  Gil Michard  Gérard Sarazin 《Geochimica et cosmochimica acta》1984,48(4):649-659

Stability constants of hydroxocomplexes of Al(III):Al(OH)²⁺ and A1(OH)₄^? have been measured in the 20–70°C temperature range by reactions involving only dissolved species. The stability constant ${}^1\text{K}{}_1$ of the first complex ion is studied by measuring pH of solutions of aluminium salts at several concentrations. ${}^1\text{β}{}_4$ of aluminate ion is deduced from equilibrium constants of the reaction between the trioxalato aluminium (III) complex ion and Al³⁺ in acid medium, and between the same complex ion and A1(OH)₄^? in alkaline medium. The K values and the associated ΔH are ${}^1\text{K}{}_1\text{= 10}{}^{\mathrm{?5.00}}$ and ΔH₁ = 11.8 Kcal; ${}^1\text{β}{}_4\text{= 10}{}^{\mathrm{?22.20}}$ and ΔH₄ = 42.45 Kcal. These last results are not in agreement with the values of recent tables for $\text{ΔG}{}^0\text{?}$ and $\text{ΔH}{}^0\text{?}$ of Al³⁺ and Al(OH)₄^?. We suggest a consistent set of data for dissolved and solid Al species and for some aluminosilicates.  相似文献   

Gibbsite solubility and thermodynamic properties of hydroxy-aluminum ions in aqueous solution at 25°C     

Howard M. May  Philip A. Helmke  Marion L. Jackson 《Geochimica et cosmochimica acta》1979,43(6):861-868

Solubility curves were determined for a synthetic gibbsite and a natural gibbsite (Minas Gerais, Brazil) from pH 4 to 9, in 0.2% gibbsite suspensions in 0.01 M NaNO₃ that were buffered by low concentrations of non-complexing buffer agents. Equilibrium solubility was approached from oversaturation (in suspensions spiked with Al(NO₃)₃ solution), and also from undersaturation in some synthetic gibbsite suspensions. Mononuclear Al ion concentrations and pH values were periodically determined. Within 1 month or less, data from over-and undersaturated suspensions of synthetic gibbsite converged to describe an equilibrium solubility curve. A downward shift of the solubility curve, beginning at pH 6.7, indicates that a phase more stable than gibbsite controls Al solubility in alkaline systems. Extrapolation of the initial portion of the high-pH side of the synthetic gibbsite solubility curve provides the first unified equilibrium experimental model of Al ion speciation in waters from pH 4 to 9.The significant mononuclear ion species at equilibrium with gibbsite are Al³⁺, AlOH²⁺, Al(OH)⁺₂ and Al(OH)^?₄, and their ion activity products are $\text{1K}{}_{50}\text{= 1.29 × 10}{}^8$, $\text{1K}{}_{\mathrm{s1}}\text{= 1.33 × 10}{}^3$, $\text{1K}{}_{\mathrm{s2}}\text{= 9.49 × 10}{}^{\mathrm{?3}}$ and $\text{1K}{}_{\mathrm{s4}}\text{= 8.94 × 10}{}^{\mathrm{?15}}$. The calculated standard Gibbs free energies of formation (ΔG°_f) for the synthetic gibbsite and the A1OH²⁺, Al(OH)⁺₂ and Al(OH)^?₄ ions are ?276.0, ?166.9, ?216.5 and ?313.5 kcal mol^?1, respectively. These ΔG°_f values are based on the recently revised ΔG°_f value for Al³⁺ (?117.0 ± 0.3 kcal mol_?1) and carry the same uncertainty. The ΔG°_f of the natural gibbsite is ?275.1 ± 0.4 kcal mol^?, which suggests that a range of ΔG°_f values can exist even for relatively simple natural minerals.  相似文献   

Prediction of Gibbs energies of formation of compounds from the elements—II. Monovalent and divalent metal silicates     

Yves Tardy  Robert M. Garrels 《Geochimica et cosmochimica acta》1977,41(1):87-92

A parameter ΔO^2?, defined as the difference between the Gibbs energy of formation of a given oxide and its aqueous cation, was used to obtain linear relationships among Gibbs energies of formation from the elements of hydroxides, oxides and aqueous metallic ions (Tardy and Garrels, 1976). Use of this parameter has now been extended to meta- and orthosilicates for which the Gibbs energies of formation of silicates from their oxides are shown to be linear functions of the ΔO^2? values of their constituent cations. The function obtained for metasilicates is: and that for orthosilicates is: in which Δ^o_? silicate is the Gibbs energy of formation from the elements of a silicate of a given cation and ∑ΔG^o_? oxides is the sum of the Gibbs energies of formation from the elements of the constituent oxides of the silicate considered.These functions can be used to test for consistency within and between various sources of thermodynamic data and to estimate free energy of formation values for previously unstudied species.  相似文献   

The ionization of acids in estuarine waters     

Frank J Millero 《Geochimica et cosmochimica acta》1981,45(11):2085-2089

The stoichiometric, $\text{K}{}_{\mathrm{HA}}{}^1$, and apparent, K'_HA, constants for the ionization of a number of weak acids (NH₄⁺, HSO₄^?, HF, H₂O, B(OH)₃, H₂CO₃, HCO₃^?, H₃PO₄, H₂PO₄^?, HPO₄², H₃AsO₄ H₂AsO₄^? and HAsO₄^2?) in seawater at 25°C diluted with water have been fitted to equations of the form (Millero, 1979). In $\text{K}{}_{\mathrm{HA}}{}^1\text{= In K}{}_{\mathrm{HA}}\text{+ AS}{}^{\mathrm{<mtext>1</mtext><mtext>2</mtext>}}\text{+ BS}$ where In K_HA is the thermodynamic constant in water, S is the salinity, A and B are adjustable parameters. The validity of this equation in estuarine waters has been examined by using an ion pairing model (Millero and Schreiber, 1981). The calculated values of $\text{K}{}_{\mathrm{HA}}{}^1$ and K'_HA at S = 35%. are in good agreement with the measured values for all the systems examined. The equation used to extrapolate the measured values to pure water K_HA predicted values that agreed with those determined by using the ion pairing model. The exception was the ionization of HPO₄^2? due to the strong interactions of Ca²⁺ and Mg²⁺ with PO₄^3?. The differences in the predicted values of $\text{K}{}_{\mathrm{HA}}{}^1$ in seawater diluted with pure water and average river water were very small for all the acids except HPO₄^2? (the maximum ΔpK = 0.96 in average river water). The larger difference in the $\text{K}{}_{\mathrm{HA}}{}^1$ for HPO₄^2? in river waters is due to the strong interactions of Ca²⁺ and PO₄^3?.  相似文献   

Revised values for the Gibbs free energy of formation of [Al(OH)_{4 aq}⁻], diaspore,boehmite and bayerite at 298.15 K and 1 bar,the thermodynamic properties of kaolinite to 800 K and 1 bar,and the heats of solution of several gibbsite samples     

Bruce S. Hemingway  Richard A. Robie  James A. Kittrick 《Geochimica et cosmochimica acta》1978,42(10):1533-1543

Solution calorimetric measurements compared with solubility determinations from the literature for the same samples of gibbsite have provided a direct thermochemical cycle through which the Gibbs free energy of formation of [Al(OH)_{4 aq}^?] can be determined. The Gibbs free energy of formation of [Al(OH)_{4 aq}^?] at 298.15 K is ?1305 ± 1 kJ/mol. These heat-of-solution results show no significant difference in the thermodynamic properties of gibbsite particles in the range from 50 to 0.05 μm.The Gibbs free energies of formation at 298.15 K and 1 bar pressure of diaspore, boehmite and bayerite are ?9210 ± 5.0, ?918.4 ± 2.1 and ?1153 ± 2 kJ/mol based upon the Gibbs free energy of [A1(OH)_{4 aq}^?] calculated in this paper and the acceptance of ?1582.2 ± 1.3 and ?1154.9 ± 1.2 kJ/mol for the Gibbs free energy of formation of corundum and gibbsite, respectively.Values for the Gibbs free energy formation of [Al(OH)_{2 aq}⁺] and [AlO_{2 aq}^?] were also calculated as ?914.2 ± 2.1 and ?830.9 ± 2.1 kJ/mol, respectively. The use of [AlC_{2 aq}^?] as a chemical species is discouraged.A revised Gibbs free energy of formation for [H₄SiO_4aq⁰] was recalculated from calorimetric data yielding a value of ?1307.5 ± 1.7 kJ/mol which is in good agreement with the results obtained from several solubility studies.Smoothed values for the thermodynamic functions C_P⁰, ($\text{H}{}_{\mathrm{T}}{}^0\text{- H}{}_{298}{}^0\text{)}\text{T}$, $\text{(}\text{G}{}_{\mathrm{T}}{}^0\text{- H}{}_{298}{}^0\text{)}\text{T}$, S_T⁰ - S₀⁰, $\text{ΔH}{}_{\mathrm{?,298}}{}^0$ kaolinite are listed at integral temperatures between 298.15 and 800 K. The heat capacity of kaolinite at temperatures between 250 and 800 K may be calculated from the following equation: C_P⁰ = 1430.26 ? 0.78850 T + 3.0340 × 10^?4T² ?1.85158 × 10^?4T²$\text{1}\text{2}\text{+ 8.3341 × 10}{}^6\text{T}{}^{\mathrm{?2}}$.The thermodynamic properties of most of the geologically important Al-bearing phases have been referenced to the same reference state for Al, namely gibbsite.  相似文献   

The thermodynamics of the carbonate system in seawater     

Frank J Millero 《Geochimica et cosmochimica acta》1979,43(10):1651-1661

The apparent constants (K'_i) for the ionization of carbonic acid in seawater at various salinities (S,%.) have been fit to equations of the form ln $\text{K'}{}_{\mathrm{i}}\text{= ln K}{}_{\mathrm{i}}\text{+ A}{}_{\mathrm{i}}\text{S}{}^{\mathrm{<mtext>1</mtext><mtext>2</mtext>}}\text{+ B}{}_{\mathrm{i}}\text{S}$whereK_i is the thermodynamic ionization constant in water, A_i, and B_i are adjustable parameters. The temperature dependence (TK) of K_i, A_i and B_i were of the form, a₀ + a₁/T + a₃ ln T. Equations of similar forms have been used to analyze the ionization constants for water and boric acid and the solubility product of calcite in seawater. The effect of pressure on the apparent constants ($\text{K}{}^{\mathrm{p}}{}_{\mathrm{i}}\text{K}{}^{\mathrm{o}}{}_{\mathrm{i}}$) have been fit to equations of the form ln $\text{(}\text{K}{}^{\mathrm{p}}{}_{\mathrm{i}}\text{K}{}^{\mathrm{o}}{}_{\mathrm{i}}\text{) = ? (ΔVP + 0.5 ΔK P}{}^2\text{)/RT}$ where the volume (ΔV) and compressibility (ΔK) changes are polynomial functions of temperature. The equations generated for various açids in seawater have been used to examine the carbonate system in seawater. Equations relating the NBS and Tris pH scales have been derived as well as equations of pH as a function of temperature and pressure. The equations from Hansson (1972, Ph.D. Thesis, University of Göteborg, Sweden) and Mehrbachet al. (1973, Limnol. Oceanogr.18, 897–907) have been used to examine the components of the carbonate system. At a fixed total alkalinity and total carbon dioxide, differences of ±0.01 m-equiv kg^?1 in HCO^?₃ and CO^2?₃ were found; however, the [CO₂] and P_co2 are nearly the same. The contribution of borate ion, B(OH)^?₄ determined from the equations of Hansson (1972, Ph.D. Thesis, University of Göteborg, Sweden) and Lyman (1957, Ph.D. Thesis, University of California, Los Angeles) differ by ±0.01 m-equiv kg^?1 for waters with the same salinity and temperature.  相似文献   

Prediction of Gibbs free energies of calcite-type carbonates and the equilibrium distribution of trace elements between carbonates and aqueous solutions     

Dimitri A Sverjensky 《Geochimica et cosmochimica acta》1984,48(5):1127-1134

A linear correlation exists between the standard Gibbs free energies of formation of calcite-type carbonates (MCO3) and the corresponding conventional standard Gibbs free energies of formation of the aqueous divalent cations (M2+) at 25 °C and 1 bar ΔGMCO30 = m(ΔGf,M2+0) ? 141,200 cal · mole?1 where m is equal to 0.9715. This relationship enables prediction of the standard free energies of formation of numerous hypothetical carbonates with the calcite structure. Associated uncertainties typically range from about ± 250 to 600 cal · mole?1. An important consequence of the above correlation is that the thermodynamic equilibrium constant for the distribution of two trace elements M and N between carbonate mineral and aqueous solution at 25 °C and 1 bar is proportional to the free energy difference between the corresponding two aqueous ions: In Combination of predicted standard free energies, entropies and volumes of carbonate minerals at 25°C and 1 bar with standard free energies of aqueous ions and the equation of state in Helgesonet al. (1981) enables prediction of the thermodynamic equilibrium constant for trace element distribution between carbonates and aqueous solutions at elevated temperatures and pressures. Interpretation of the thermodynamic equilibrium constant in terms of concentration ratios in the aqueous phase is considerably simplified if pairs of divalent trace elements are considered that have very similar ionic radii (e.g., $\text{Sr}{}^{\mathrm{2+}}\text{Pb}{}^{\mathrm{2+}}$, $\text{Mg}{}^{\mathrm{2+}}\text{Zn}{}^{\mathrm{2+}}$). In combination with data for the stabilities of complex ions in aqueous solutions, the above calculations enable useful limits to be placed on the concentrations of trace elements in hydrothermal solutions.  相似文献   

Complex formation in the ternary system Mg(II)-CO₂-H₂O     

Walter Riesen  Heinz Gamsjäger  Paul W. Schindler 《Geochimica et cosmochimica acta》1977,41(9):1193-1200

The carbonato and hydrogencarbonato complexes of Mg²⁺ were investigated at 25 and 50° in solutions of the constant ClO₄^? molality (3 M) consisting preponderantly of NaClO₄. The experimental data could be explained assuming the following equilibria: Mg²⁺ + CO_2B + H₂O ag MgHCO⁺₃ + H⁺, $\text{log}{}^1\text{β}{}_1\text{= ?7.64}{}_4\text{± 0.01}{}_7\text{(25°), ?7.46}{}_2\text{± 0.01}{}_1\text{(50°)}$, Mg²⁺ + 2 CO_2g + 2 H₂Oag Mg(HCO₃)⁰₂ ± 2 H⁺, $\text{log}{}^1\text{β}{}_2\text{= ?15.00 ± 0.14 (25°)}$, ?15.37 ± 0.39 (50°), Mg²⁺ + CO_2g + H₂Oag MgCO⁰₃ + 2 H⁺, $\text{log}{}^1\text{k}{}_1\text{= ?15.64 ± 0.06 (25°)}$,?15.23 ± 0.02 (50°), with the assumption γ_MgCO3⁰ = γ_Mg(HCO3)⁰2, ΔG⁰(I = 0) for the reaction MgCO⁰₃ + CO_2g + H₂O = Mg(HCO₃)⁰₂ was estimated to be ?3.9₁ ± 0.8₆ and 0.6 ± 2.4 kJ/mol at 25 and 50°C, respectively. The abundance of carbonate linked Mg(II) species in fresh water systems is discussed.  相似文献   

Carbonate equilibria in hydrothermal systems: First ionization of carbonic acid in NaCl media to 300°C     

C.S Patterson  G.H Slocum  R.H Busey  R.E Mesmer 《Geochimica et cosmochimica acta》1982,46(9):1653-1663

The ionization quotients of aqueous carbon dioxide (carbonic acid) have been precisely determined in NaCl media to 5 m and from 50° to 300°C using potentiometric apparatus previously developed at Oak Ridge National Laboratory. The pressure coefficient was also determined to 250°C in the same media. These results have been combined with selected information in the literature and modeled in two ways to arrive at the best fits and to derive the thermodynamic parameters for the ionization reaction, including the equilibrium constant, activity coefficient quotients, and pressure coefficients. The variation with temperature of the two fundamental quantities $\text{Δ}\text{V}\text{?}{}^{\mathrm{o}}$ and $\text{Δ}\text{C}\text{?}{}^{\mathrm{o}}{}_{\mathrm{p}}$ were examined along the saturation vapor pressure curve and at constant density. The results demonstrated again that for reactions with minimal electrostriction changes the magnitudes and variations of $\text{Δ}\text{C}\text{?}{}^{\mathrm{o}}{}_{\mathrm{p}}$ and $\text{Δ}\text{V}\text{?}{}^{\mathrm{o}}$ with temperature are small and, in addition, $\text{Δ}\text{C}\text{?}{}_{\mathrm{p}}$ and $\text{Δ}\text{V}\text{?}$ are approximately independent of salt concentration.The results have also been applied to an examination of the solubility of calcite as a function of pH (in a given NaCl medium) for the neutral to acidic region both for systems with fixed CO₂ pressure and systems where the calcium ion concentration equals the concentration of carbon. The pH of saturated solutions of calcite with P_CO2 of 12 bars increases from 5.1 to 5.5 between 100° and 300°C.  相似文献   

Geochimie et teneurs isotopiques des systèmes saisonniers de dissolution de la calcite dans un sol sur craie     

L. Dever  R. Durand  J.Ch. Fontes  P. Vachier 《Geochimica et cosmochimica acta》1982,46(10):1947-1956

In a soil developed on the Cretaceous chalk of the Eastern Paris basin, calcite dissolution begins at the surface. The soil water is rapidly saturated in calcite. Calcite dissolution follows two different pathways according to seasonal pedoclimatic conditions.During winter: the soil is only partly saturated in water and the CO₂ partial pressure is low (Ca 10^?3 atm.). As a consequence total inorganic dissolved carbon (TIDC) is a hundred times the carbon content of the gaseous phase. Equilibrium is usually observed between the two phases. It is a closed system. The measured carbon 14 activity (87,5%) and ¹³C content ($\text{δ}{}_{\mathrm{tidc}}{}^{13}\text{C = ?12,2\%0}$) of the drainage water are very close to theoretical values calculated for an ideal mixing system between gaseous and mineral phases (respectively characterized by the following isotopic values: $\text{δ}{}_{\mathrm{G}}{}^{13}\text{C = ?21,5\%0}$; $\text{A}{}_{\mathrm{G}}{}^{14}\text{C = 118\%}$; $\text{δ}{}_{\mathrm{M}}{}^{13}\text{C = +2,9\%0}$; $\text{A}{}_{\mathrm{M}}{}^{14}\text{C = 28\%}$).During spring and summer: the soil moisture decreases, the input of biogenic CO₂ induces an increase of the soil CO₂ partial pressure (Ca from 3.10^?3 atm to 7.10^?3 atm). The carbon content of the gaseous phase is higher by an order of magnitude compared to winter conditions. Therefore the aqueous phase is undersaturated in CO₂ with respect to the latter. This disequilibrium occurs as a result of unbalanced rates of CO₂ dissolution and CO₂ effusion toward atmosphère. It is an open system. The carbon isotopic ratio of the aqueous phase is regulated by that of the gaseous phase, as demonstrated by the agreement between measured and calculated isotopic compositions (respectively δ_{L mes} = from ?9,4%0 to ?11,5%0, δ_{l calc} = from ?9,8%0 to ?13,9%0 A_{L mes} = 119%, A_{L calc} = from 119% to 125%).The solutions originating from both systems (open and closed) move downwards without significant mixing together. It has also been observed that no significant variation of the TIDC isotopic composition occurs during precipitation of secondary calcite.  相似文献   

Theoretical prediction of phase relations among aqueous solutions and minerals: Salton Sea geothermal system     

Dennis K. Bird  Denis L. Norton 《Geochimica et cosmochimica acta》1981,45(9):1479-1494

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Distribution coefficients of Eu and Sr for plagioclase-liquid and clinopyroxene-liquid equilibria in oceanic ridge basalt: an experimental study     

Richard J. Williams 《Geochimica et cosmochimica acta》1974,38(9):1415-1433

The distribution coefficients of Eu and Sr for plagioclase-liquid and clinopyroxene-liquid pairs as a function of temperature and oxygen fugacity were experimentally investigated using an oceanic ridge basalt enriched with Eu and Sr as the starting material. Experiments were conducted between 1190° and 1140°C over a range of oxygen fugacities between 10^?8 and 10^?14 atm.The molar distribution coefficients are given by the equations: log^PL_Sr = 7320/T ? 4.62logK^CPX_Sr = 18020/T ? 13.10. Similarly, the weight fraction distribution coefficients are given by the equations: logD^PL_Sr = 6570/T ? 4.30logD^CPX_Sr = 18434/T ? 13.62.Although the mole fraction distribution coefficients have a smaller dependence on bulk composition than do the weight fraction distribution coefficients, they are not independent of bulk composition, thereby restricting the application of these experimental results to rocks similar to oceanic ridge basalts in bulk composition.Because the Sr distribution coefficients are independent of oxygen fugacity, they may be used as geothermometers. If the temperature can be determined independently — for example, with the Sr distribution coefficients, the Eu distribution coefficients may be used as oxygen geobarometers. Throughout the range of oxygen fugacities ascribed to terrestrial and lunar basalts, plagioclase concentrates Eu but clinopyroxene rejects Eu.  相似文献   

Density calculations for silicate liquids. I. Revised method for aluminosilicate compositions     

Y Bottinga  D Weill  P Richet 《Geochimica et cosmochimica acta》1982,46(6):909-919

Equations are developed for calculating the density of aluminosilicate liquids as a function of composition and temperature. The mean molar volume at reference temperature T_r, is given by $\text{V}{}_{\mathrm{r}}\text{= ∑X}{}_{\mathrm{i}}\text{V}\text{?}{}^{\mathrm{o}}{}_{\mathrm{i}}\text{+ X}{}_{\mathrm{A}}\text{V}\text{?}{}^{\mathrm{o}}{}_{\mathrm{A}}$, where the summation is taken over all oxide components except A1₂O₃, X stands for mole fraction, terms are constants derived independently from an analysis of volume-composition relations in alumina-free silicate liquids, and $\text{V}\text{?}{}^{\mathrm{o}}{}_{\mathrm{A}}$ is the composition-dependent apparent partial molar volume of Al₂O₃. The thermal expansion coefficient of aluminosilicate liquids is given by $\text{α = ∑X}{}_{\mathrm{i}}\text{}\text{?}\text{ga}{}_{\mathrm{i}}{}^{\mathrm{o}}\text{+ X}{}_{\mathrm{A}}\text{}\text{?}\text{ga}{}_{\mathrm{A}}{}^{\mathrm{o}}$, where $\text{}\text{?}\text{ga}{}_{\mathrm{i}}{}^{\mathrm{o}}$ terms are constants independent of temperature and composition, and $\text{}\text{?}\text{ga}{}^{\mathrm{o}}{}_{\mathrm{A}}$ is a composition-dependent term representing the effect of Al₂O₃ on the thermal expansion. Parameters necessary to calculate the volume of silicate liquids at any temperature T according to V(T) = V_rexp[α(T-T_r)], where T_r = 1400°C have been evaluated by least-square analysis of selected density measurements in aluminosilicate melts. Mean molar volumes of aluminosilicate liquids calculated according to the model equation conform to experimentally measured volumes with a root mean square difference of 0.28 $\text{cc}\text{mole}$ and an average absolute difference of 0.90% for 248 experimental observations. The compositional dependence of $\text{V}\text{?}{}^{\mathrm{o}}{}_{\mathrm{A}}$ is discussed in terms of several possible interpretations of the structural role of Al³⁺ in aluminosilicate melts.  相似文献   

Thermodynamic analyses of equilibria involving olivine,orthopyroxene and garnet     

Toshisuke Kawasaki  Yoshito Matsui 《Geochimica et cosmochimica acta》1983,47(10):1661-1679

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Oxygen and sulfur isotope equilibria in the BaSO4HSO4−H2O system from 110 to 350°C and applications     

Minoru Kusakabe  Brian W. Robinson 《Geochimica et cosmochimica acta》1977,41(8):1033-1040

Oxygen isotope exchange between BaSO₄ and H₂O from 110 to 350°C was studied using 1 m H₂SO₄-1 m NaCl and 1 m NaCl solutions to recrystallize the barite. The slow exchange rate (only 7% exchange after 1 yr at 110°C and 91% exchange after 22 days at 350°C in 1 m NaCl solution) prompted the use of the partial equilibrium technique. However, runs at 300 and 350°C were checked by complete exchange experiments. The temperature calibration curve for the isotope exchange is calculated giving most weight to the high temperature runs where the partial equilibrium technique can be tested. Oxygen isotope fractionation factors (α) in 1 m NaCl solution (110–350°C), assuming a value of 1.0407 for α_CO2H2O at 25°C, are: .These data, when corrected for ion hydration effects in solution (Truesdell, 1974), give the fractionation factors in pure water: .In the 1 m H₂SO₄-1 m NaCl runs, sulfur isotope fractionation between HSO^?₄ and BaSO₄ is less than the detection limit of 0.4%. A barite-sulfide geothermometer is obtained by combining HSO^?₄H₂S and sulfide-H₂S calibration data.Barite in the Derbyshire ore field, U.K., appears to have precipitated in isotopic equilibrium with water and sulfur in the ore fluid at temperatures less than 150°C. At the Tui Mine, New Zealand, the barite-water geothermometer indicates temperatures of late stage mineralization in the range 100–200°C. A temperature of 350 ± 20°C is obtained from the barite-pyrite geothermometer at the Yauricocha copper deposit, Peru, and oxygen isotope analyses of the barite are consistent with a magmatic origin for the ore fluids.  相似文献   

Sedimentary iron monosulfides: Kinetics and mechanism of formation     

Albert J Pyzik  Sheldon E Sommer 《Geochimica et cosmochimica acta》1981,45(5):687-698

The reaction between hydrous iron oxides and aqueous sulfide species was studied at estuarine conditions of pH, total sulfide, and ionic strength to determine the kinetics and formation mechanism of the initial iron sulfide. Total, dissolved and acid extractable sulfide, thiosulfate, sulfate, and elemental sulfur were determined by spectrophotometric methods. Polysulfides, S₄^2? and S₅^2?, were determined from ultraviolet absorbance measurements and equilibrium calculations, while product hydroxyl ion was determined from pH measurements and solution buffer capacity.Elemental sulfur, as free and polysulfide sulfur, was 86% of the sulfide oxidation products; the remainder was thiosulfate. Rate expressions for the reduction and precipitation reactions were determined from analysis of electron balance and acid extractable iron monosulfide vs time, respectively, by the initial rate method. The rate of iron reduction in moles/liter/minute was given by $\text{d(}\text{reduction}\text{Fe)}\text{dt}\text{= kS}{}_{\mathrm{t}}{}^{0.5}\text{(J}{}^{\mathrm{+}}\text{)}{}^{0.5}\text{A}{}_{\mathrm{FeOOH}}{}^1$ where S_t was the total dissolved sulfide concentration, (H⁺) the hydrogen ion activity, both in moles/ liter; and A_FeOOH the goethite specific surface area in square meters/liter. The rate constant, k, was 0.017 ± 0.002m^?2 min^?1. The rate of reduction was apparently determined by the rate of dissolution of the surface layer of ferrous hydroxide. The rate expression for the precipitation reaction was $\text{d(FeS)}\text{dt}\text{= kS}{}_{\mathrm{t}}{}^1\text{(H}{}^{\mathrm{+}}\text{)}{}^1\text{A}{}_{\mathrm{FeOOH}}{}^1$ where $\text{d(FeS)}\text{dt}$ was the rate of precipitation of acid extractable iron monosulfide in moles/liter/minute, and k = 82 ± 18 mol^?1l²m^?2 min^?1.A model is proposed with the following steps: protonation of goethite surface layer; exchange of bisulfide for hydroxide in the mobile layer; reduction of surface ferric ions of goethite by dissolved bisulfide species which produces ferrous hydroxide surface layer elemental sulfur and thiosulfate; dissolution of surface layer of ferrous hydroxide; and precipitation of dissolved ferrous specie and aqueous bisulfide ion.  相似文献   

Stability of ethylxanthate ion in neutral and weakly acidic media. Part 1: Influence of pH     

M. Maillot  J.-L. Cecile  R. Bloise 《International Journal of Mineral Processing》1984,13(3):193-210

Xanthates are used in the flotation of sulfide ores although their aqueous solutions are not stable under certain conditions. Their stability in acidic and weakly acidic aqueous solutions was therefore investigated, as these media are required for some processes.The peak absorbances of ethylxanthate ion and carbon disulfide were first determined in aqueous solution. The decomposition of ethylxanthate ion was analyzed by measuring variations in absorbance (at 301 nm) and pH with respect to time. A pH regulation system was then used while measuring variations in absorbance and productions of protons caused by xanthate decomposition.The results concerning xanthate half-lives show good agreement with the literature, but the kinetic results deviate substantially. The following relation was obtained for half-life: We established that ethylxanthate decomposition at pH 4 is a first order reaction with respect to ethylxanthate concentration, and postulating this order to the other pH values, the following kinetic relation was found: where v is the rate of decomposition (mol l^?1 min^?1), and [EtX^?]_∞ is the ethylxanthate concentration when the decomposition equilibria are reached (mol l^?1). The better concentration was found to obey the law:   相似文献   

Experimental hydrogen isotope studies—I. Systematics of hydrogen isotope fractionation in the systems epidote-H₂O,zoisite-H₂O and AlO(OH)-H₂O     

Colin M. Graham  Simon M.F. Sheppard  Timothy H.E. Heaton 《Geochimica et cosmochimica acta》1980,44(2):353-364

$\text{H}\text{D}$ Fractionation factors between epidote minerals and water, and between the AlO(OH) dimorphs boehmite and diaspore and water, have been determined between 150 and 650°C. Small water mineral ratios were used to minimise the effect of incongruent dissolution of epidote minerals. Waters were extracted and analysed directly by puncturing capsules under vacuum. Hydrogen diffusion effects were eliminated by using thick-walled capsules.$\text{H}\text{D}$ Exchange rates are very fast between epidote and water (and between boehmite and water), complete exchange taking only minutes above 450°C but several months at 250°C. Exchange between zoisite and water (and between diaspore and water) is very much slower, and an interpolation method was necessary to determine fractionation factors at 450 and below.For the temperature range 300–650°C, the $\text{H}\text{D}$ equilibrium fractionation factor (α^e) between epidote and water is independent of temperature and Fe content of the epidote, and is given by 1000 In α_epidote-H2O^e = ?35.9 ± 2.5, while below 300°C 1000 In , with a ‘cross-over’ estimated to occur at around 185°C. By contrast, zoisite-water fractionations fit the relationship 1000 In .All studied minerals have hydrogen bonding. Fractionations are consistent with the general relationship: the shorter the O-H -- O bridge, the more depleted is the mineral in D.On account of rapid exchange rates, natural epidotes probably acquired their H-isotope compositions at or below 200°C, where fractionations are near or above 0%.; this is in accord with the observation that natural epidotes tend to concentrate D relative to other coexisting hydrous minerals.  相似文献   

Rates of reaction of pyrite and marcasite with ferric iron at pH 2     

C.L Wiersma  J.D Rimstidt 《Geochimica et cosmochimica acta》1984,48(1):85-92

The relative reactivities of pulverized samples (100–200 mesh) of 3 marcasite and 7 pyrite specimens from various sources were determined at 25°C and pH 2.0 in ferric chloride solutions with initial ferric iron concentrations of 10^?3 molal. The rate of the reaction: was determined by calculating the rate of reduction of aqueous ferric ion from measured oxidation-reduction potentials. The reaction follows the rate law: where m_Fe³⁺ is the molal concentration of uncomplexed ferric iron, k is the rate constant and $\text{A}\text{M}$ is the surface area of reacting solid to mass of solution ratio. The measured rate constants, k, range from 1.0 × 10^?4 to 2.7 × 10^?4 sec^?1 ± 5%, with lower-temperature/early diagenetic pyrite having the smallest rate constants, marcasite intermediate, and pyrite of higher-temperature hydrothermal and metamorphic origin having the greatest rate constants. Geologically, these small relative differences between the rate constants are not significant, so the fundamental reactivities of marcasite and pyrite are not appreciably different.The activation energy of the reaction for a hydrothermal pyrite in the temperature interval of 25 to 50°C is 92 kJ mol^?1. This relatively high activation energy indicates that a surface reaction controls the rate over this temperature range. The BET-measured specific surface area for lower-temperature/early diagenetic pyrite is an order of magnitude greater than that for pyrite of higher-temperature origin. Consequently, since the lower-temperature types have a much greater $\text{A}\text{M}$ ratio, they appear to be more reactive per unit mass than the higher temperature types.  相似文献