A surface complex reaction model for the pH-dependence of corundum and kaolinite dissolution rates

Carroll, Susan A.; Walther, John V.

doi:10.1016/0016-7037(88)90030-0

Cited by 294 publications

(111 citation statements)

References 36 publications

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“…2, the rate of kaolinite dissolution in acetic acid was dependent upon the pH value of the solution, in the range of pH 2.5-5.5, and the logarithm of the rate of kaolinite dissolution was linearly dependent upon the solution pH. The reaction order with respect to n H + of the proton-promoted dissolution reaction is 0.28, approximately close to previous reports [2,13,15,36]. Then, the dissolution of kaolinite enhanced only by protons could be described by following equation:…”

Section: Discussionsupporting

confidence: 87%

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Dissolution of kaolinite induced by citric, oxalic, and malic acids

Wang

Hu³

et al. 2005

Journal of Colloid and Interface Science

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Section: Discussionsupporting

confidence: 87%

“…Dissolution of kaolinite in solutions of oxalate, malonate, salicylate, and phthalate ligands [13][14][15] showed that oxalate enhanced dissolution of kaolinite to the greatest extent, followed by salicylate. Malonate and phthalate showed little effect on dissolution rates.…”

Section: Introductionmentioning

confidence: 99%

Dissolution of kaolinite induced by citric, oxalic, and malic acids

Wang

Hu³

et al. 2005

Journal of Colloid and Interface Science

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“…Based on the conversion rates, we estimated the reaction rate coefficient of dawsonite assuming r = kS(Ω-1), an average reactive surface area of 0.1 m 2 , and 2 000 times dawsonite supersaturation estimated by the PHREEQC speciation code (Parkhurst and Appelo, 1999) using the llnl.dat database and assuming NaHCO 3 and aluminium concentrations given by nahcolite and gibbsite saturations respectively, we get reaction rate coefficients of 1.17 × 10 -13 and 2.83 × 10 -12 moles/m 2 s respectively from kaolinite and gibbsite. These rates are indeed orders of magnitude slower than the corresponding dawsonite dissolution rate coefficients reported by Hellevang et al (2005Hellevang et al ( , 2010, but is in the same range as estimated kaolinite (e.g., Carroll-Webb and Walther, 1988;Bauer and Berger, 1998) and gibbsite (Nagy and Lasaga, 1992) dissolution rates at these conditions. At this point we can thus only assume that the precipitation rates of dawsonite at high supersaturations are faster than the aluminium source dissolution rates and we can not exclude the fast dawsonite precipitation rates as predicted from the dissolution rate experiments.…”

Section: Introductionsupporting

confidence: 45%

Why is Dawsonite Absent in CO₂Charged Reservoirs?

Hellevang

Declercq

Aagaard

2011

Oil Gas Sci. Technol. – Rev. IFP Energies nouvelles

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Résumé -Pourquoi la dawsonite est-elle absente des réservoirs chargés en CO 2 ? -La possibilité pour la dawsonite -un hydroxy-carbonate de sodium et d'aluminium (NaAl(OH) 2 CO 3 ) -de précipiter dans les aquifères salins dès lors que l'on y injecte du CO 2 a été suggérée dans maintes simulations menées sur différentes compositions (minéraux et solution aqueuse) et dans des conditions de pression et de température variées. Pourtant, sur le strict plan de la stabilité thermodynamique, la dawsonite paraît moins communément répandue que l'on pourrait s'y attendre dans les analogues naturels de stockages de CO 2 . On a cartographié la stabilité thermodynamique de la dawsonite par rapport à des phases minérales comme l'albite, la kaolinite et l'analcime, entre 37 et 200°C, puis on a simulé numériquement des évolutions minérales en système fermé à l'aide d'un nouveau formalisme cinétique pour la précipitation, formalisme qui inclut (1) un terme de nucléation, basé sur la théorie classique de la nucléation, et (2) un terme de croissance, dérivant de la théorie BCF de la croissance. Utilisant cette équation de vitesse, on a réalisé une étude de sensibilité pour examiner comment la croissance de la dawsonite varie avec la composition minérale, la température, la pression partielle de CO 2 , la vitesse de nucléation -et les variations de cette dernière avec la température et l'affinité chimique -, et enfin la constante de vitesse adoptée pour la loi de précipitation de la dawsonite. Lorsqu'on empêche la dawsonite de précipiter, le rapport de sur-saturation ne dépasse jamais 3 ou 4 pour des solutions de type eau de mer. La propension accrue à précipiter si la pression partielle de CO 2 augmente est contrebalancée par l'effet d'acidification de la solution. Diminuer jusqu'à 5 ordres de grandeur la vitesse de précipitation n'a qu'un effet limité sur le démarrage d'une croissance significative, et au bout de 1 000 ans de simulation la quantité de dawsonite formée est de 37 % ce qu'elle est dans le régime de précipitation maximale. Réduire le taux de nucléation a le même effet de retarder la croissance du minéral. Au bout du compte, sur la base de considérations thermodynamiques et au vu des simulations présentées, on suggère que le potentiel de croissance de la dawsonite est restreint à une fenêtre de températures médianes. À basse température, la stabilité relativement grande de la dawsonite vis-à-vis des minéraux en compétition est contrecarrée par la forte diminution des vitesses de nucléation et de croissance, ce qui reflète une puissante barrière énergétique. Abstract

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“…Uptake is reduced by order-of-magnitude increases in the concentration of the 80 supporting electrolyte below pH ~ 7-8, but uptake is insensitive to change in electrolyte concentration above this pH range (Figure 2a) Cowan " 60 et al, 1992;O'Day, 1992;Zachara et al, 1994). Based o. on solution uptake experiments and potentiometric ti-_~ trations, uptake of cations, anions, and/or protons from o solution onto kaolinite as a function ofpH is generally O 40 modeled assuming two classes of sites: 1) ion exchange *~ or nonspecific adsorption sites that exchange background electrolyte cations with weakly bound, hydrat-20 ed adions ("outer-sphere" complexes) and 2) specific adsorption at amphoteric surface hydroxyl sites (e.g., AI-OH, St-OH) in which surface sites hydrolyze and adions bond directly to surface oxygens such that they are not easily displaced by electrolyte ions ("innersphere" complexes) (Farrah et al, 1980;Riese, 1982;Sposito, 1984;Zachara et al, i 1988;Carroll-Webb and Walther, 1988;Cowan et al, 1992;Singh and Mattigod, 1992;Wieland and Stumm, 1992;Xie and Walther, 1992;Zachara et al, 1994). In these previous macroscopic studies, however, different models of ion uptake have reproduced experimental ~" -5 data equally well and do not supply a unique solution.…”

Section: Sorption and Sorption Sites On Kaolinitementioning

confidence: 99%

Molecular Structure and Binding Sites of Cobalt(II) Surface Complexes on Kaolinite from X-Ray Absorption Spectroscopy

O’Day

Parks

Brown

1994

Clays and clay miner.

127

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Abstract--X-ray absorption spectroscopy (XAS) was used to determine the local molecular environment of Co(II) surface complexes sorbed on three different kaolinites at ambient temperature and pressure in contact with an aqueous solution. Interatomic distances and types and numbers of backscattering atoms have been derived from analysis of the extended X-ray absorption fine structure (EXAFS). These data show that, at the lowest amounts of Co uptake on kaolinite (0.20--0.32 ~mol m-2), Co is surrounded by ~60 atoms at 2.04-2.08 A and a small number orAl or Si atoms (N = 0.6-1.5) at two distinct distances, 2.67-2.72 A and 3.38-3.43/~. These results indicate that Co bonds to the kaolinite surface as octahedrally coordinated, bidentate inner-sphere mononuclear complexes at low surface coverages, confirming indirect evidence from solution studies that a fraction of sorbed Co forms strongly bound complexes on kaolinite. In addition to inner-sphere complexes identified by EXAFS spectroscopy, solution studies provide evidence for the presence of weakly bound, outer-sphere Co complexes that cannot be detected directly by EXAFS. One orientation for inner-sphere complexes indicated by XAS is bidentate bonding of Co to oxygen atoms at two A1-O-Si edge sites or an A1-O-Si and AI-OH (inner hydroxyl) edge site, i.e., comersharing between Co octahedra and AI and Si potyhedra. At slightly higher surface sorption densities (0.51-0.57/~mol m-2), the presence of a small number of second-neighbor Co atoms (average Nco < 1) at 3.10-3.13/~ indicates the formation of oxy-or hydroxy-bridged, multinuclear surface complexes in addition to mononuclear complexes. At these surface coverages, Co-Co and Co-AI/Si distances derived from EXAFS are consistent with edge-sharing between Co and AI octahedra on either edges or (001) faces of the aluminol sheet in kaolinite. Multinuclear complexes form on kaolinite at low surface sorption densities equivalent to < 5% coverage by a monolayer of oxygen-ligated Co octahedra over the N2-BET surface area. These spectroscopic results have several implications for macroscopic modeling of metal ion uptake on kaolinite: 1) Primary binding sites on the kaolinite surface at low uptake are edge, non-bridging AI-OH inner hydroxyl sites and edge AI-O-Si bridging oxygen sites, not Si-OH sites typically assumed in sorption models; 2) specific adsorption of Co is via bidentate, inner-sphere complexation; and 3) at slightly higher uptake but still a small fraction of monolayer coverage, formation of Co multinuclear complexes, primarily edge-sharing with A1-OH octahedra, begins to dominate sorption.

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A surface complex reaction model for the pH-dependence of corundum and kaolinite dissolution rates

Cited by 294 publications

References 36 publications

Dissolution of kaolinite induced by citric, oxalic, and malic acids

Dissolution of kaolinite induced by citric, oxalic, and malic acids

Why is Dawsonite Absent in CO₂Charged Reservoirs?

Molecular Structure and Binding Sites of Cobalt(II) Surface Complexes on Kaolinite from X-Ray Absorption Spectroscopy

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A surface complex reaction model for the pH-dependence of corundum and kaolinite dissolution rates

Cited by 294 publications

References 36 publications

Dissolution of kaolinite induced by citric, oxalic, and malic acids

Dissolution of kaolinite induced by citric, oxalic, and malic acids

Why is Dawsonite Absent in CO2Charged Reservoirs?

Molecular Structure and Binding Sites of Cobalt(II) Surface Complexes on Kaolinite from X-Ray Absorption Spectroscopy

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Why is Dawsonite Absent in CO₂Charged Reservoirs?