Equilibrium Dissolution Rate Research Articles

Kyanite far from equilibrium dissolution rate at 0–22 °C and pH of 3.5–7.5

Kyanite is an important and slow-dissolving mineral. Earlier work has measured its dissolution rate at high temperature and acidic pH, but experimental measurements at low temperature and near neutral pH were lacking. The rate equation by Palandri and Kharaka (A compilation of rate parameters of water–mineral interaction kinetics for application to geochemical modeling. US Geological Survey, Open File Report 2004-1068, 2004) indicates that the rate of kyanite dissolution at room temperature and near neutral pH is on the order of 10−17 mol m−2 s−1, orders of magnitudes slower than most common silicate minerals such as albite and quartz. This study used an externally-stirred mixed-flow reactor, which allows high solid:solution ratios, to measure the dissolution rate of kyanite at 0–22 °C and pH of 3.5–7.5. The measured dissolution rate of kyanite is 4.6–7.6 × 10−13 mol m−2 s−1 at 22 °C, and the apparent activation energy is 73.5 kJ mol−1. This dissolution rate is close to the rate of quartz dissolution and four orders of magnitude faster than the prediction by rate equation of Palandri and Kharaka (2004). Based on our new experimental data, we recommend the following rate equation for modeling the dissolution of kyanite at ambient temperatures. $$r = ke^{{\frac{{ - E_{a} }}{R}\left( {\frac{1}{T} - \frac{1}{298.15}}\right)}}$$ where k = 5.08 × 10−13 mol m−2 s−1, and Ea = 73.5 kJ mol−1. Review of literature data (Carroll in The dissolution behavior of corundum, kaolinite, and andalusite: a surface complex reaction model for the dissolution of aluminosilicate minerals in diagenetic and weathering environs. Dissertation, Northwestern University, 1989) led to a recommended rate equation for andalusite as for T = 25 °C and pH = 2–10: $$r = k_{1} a_{{{\text{H}}^{ + } }}^{{n_{1} }} + k_{2} + k_{3} a_{{{\text{H}}^{ + } }}^{{n_{3} }}$$ where k1 = 4.04 × 10−10 mol m−2 s−1, k2 = 7.95 × 10−10 mol m−2 s−1, k3 = 1.01 × 10−17 mol m−2 s−1, n1 = 1.2 and n3 = − 0.6.

An experimental study of CO2‐oil‐brine‐rock interaction under in situ reservoir conditions

AbstractTo understand the mineralogical and chemical changes in oil‐bearing reservoirs (e.g., depleted oil reservoirs) during massive CO2 injection, we have carried out a core‐flooding experimental study of CO2‐oil‐brine‐rock interactions under a simulated reservoir condition of 100°C and 24 MPa. The experimental condition is based on field data from a CO2‐EOR project in the southern Songliao Basin. This oil‐bearing CO2‐flooding experiment used the same experimental setup, reservoir conditions, and workflow as the oil‐free experiment reported by Yu et al. (). The sandstone core samples used in the experiment have similar mineralogical compositions as that used in the previous experiment. Compared with the oil‐free experiment, the presence of oil appears to substantially reduce the reaction degree between the CO2 fluid and some sensitive minerals. The dissolution rates of the K‐feldspar and carbonate minerals for the oil‐bearing experiment are 1/5 and 1/4 of that for the oil‐free experiments, respectively. For the silicate minerals represented by the K‐feldspar, the presence of oil mainly delays the dissolution during the experiment, and reduces the equilibrium dissolution rate. For the carbonate minerals, the presence of oil appears to primarily affect the dissolution at the beginning of the experiments, and reduce the maximum dissolution rate attained. The core permeabilities for the oil‐free and oil‐bearing cases are both reduced after experiments. The reduction in permeability is probably due to the precipitation of fine siliceous mineral and clay particles released by the dissolution of the carbonate cement, which may clog some pore throats. The results provide some new insights on the fluid‐rock interaction during CO2 injection in depleted oil reservoirs or during CO2‐EOR.

Do organic ligands affect forsterite dissolution rates?

Far-from equilibrium, steady state forsterite dissolution rates were measured at pH ∼3 and 25°C in aqueous solutions containing 0.1m/kg NaCl and up to 0.1mol/kg of 13 distinct dissolved organic ligands in mixed-flow reactors. The organic ligands considered in this study include those common in Earth surface environments and those considered as potential catalysts for use in CO2 sequestration efforts: acetate, oxalate, citrate, EDTA4−, glutamate, gluconate, malonate, aspartate, tartrate, malate, alginate, salycilate and humate. The presence of up to 0.1mol/kg of each organic ligand altered forsterite dissolution rates less than 0.2 log units, which is the estimated uncertainty of the measured rates. Results obtained in this study, therefore, suggest that the presence of aqueous organic anions negligibly affects forsterite far-from equilibrium dissolution rates in most natural environments, and indicate that forsterite carbonation may not be appreciably accelerated by organic ligand catalysis.

Bridging the gap between laboratory measurements and field estimations of silicate weathering using simple calculations

Weathering rates of silicate minerals observed in the laboratory are in general up to five orders of magnitude higher than those inferred from field studies. Simple calculations show that even if the field conditions were fully simulated in standard laboratory experiments, it would be impossible to measure the slow rates of mineral dissolution that are observed in the field. As it is not possible to measure the dissolution rates under typical field conditions, one should extrapolate the available data to the field conditions. To do this, a rate law for the dissolution of plagioclase in the field was formulated by combining the far from equilibrium dissolution rate of weathered natural oligoclase at 25°C with the effect of deviation from equilibrium on dissolution rate of fresh albite at 80°C. In contrast to the common view that laboratory experiments predict dissolution rates that are faster than those in the field, the simulation based on this rate law indicates that laboratory dissolution experiments actually predict slower rates than those observed in the field. This discrepancy is explained by the effect of precipitation of secondary minerals on the degree of saturation of the primary minerals and therefore on their dissolution rate. Indeed, adding kaolinite precipitation to the simulation significantly enhances the dissolution rate of the plagioclase. Moreover, a strong coupling between oligoclase dissolution and kaolinite precipitation was observed in the simulation. We suggest that such a coupling must exist in the field as well. Therefore, any attempt to predict the dissolution rate in the field requires knowledge of the rate of the secondary mineral precipitation.

On the interpretation of closed system mineral dissolution experiments: comment on “Mechanism of kaolinite dissolution at room temperature and pressure part II: kinetic study” by

Huertas et al. (1999) reported the results of a series of constant-pH kaolinite dissolution experiments performed at 25°C in closed system reactors. The authors of this study chose to interpret these experiments assuming that the dissolution rates of kaolinite in their reactors, after a certain initial elapsed time, were time independent. The purposes of this comment are to (a) recall evidence available in the literature demonstrating that constant-pH kaolinite dissolution does not attain a constant rate in closed system reactors; (b) present a simple equation describing solution composition evolution during closed system multioxide mineral or glass dissolution experiments; and (c) demonstrate that Huertas et al.’s (1999) data are consistent with kaolinite rates that are proportional to cAl nAl where cAl refers to the concentration of aqueous aluminum, and nAl denotes a constant. Over the years both closed and open system reactors have been used to measure mineral and glass dissolution rates. Open system reactors such as mixed-flow reactors (e.g., Berger et al., 1994) or fluidized bed reactors (e.g., Chou and Wollast, 1985; Wogelius and Walther, 1990) provide dissolution rates at fixed solution compositions directly from the measured mass flux of the departing reactive fluids. In contrast, closed system reactors such as that used by Huertas et al. (1999) yield dissolution rates indirectly from the temporal aqueous concentration evolution of the reactive fluid. Two examples illustrating measured solution compositions taken from Huertas et al. (1999) as a function of elapsed time are given in Figure 1. The behavior illustrated in this figure is typical of the temporal aqueous concentration evolution during the closed system stoichiometric dissolution of a mineral. Early during each experiment the concentrations of metals liberated by dissolution increase at atypically fast rates; with time, the metal release rate slows, and therefore, plots of metal concentration versus time exhibit a parabolic pattern. A commonly adopted method for interpreting these observations, as well as that adopted by Huertas et al. (1999), is to assume that the observed initial rapid dissolution rates stem from temporary processes such as the relatively rapid dissolution of fine particles or high-energy surfaces sites present on the larger grains. The contributions of these relatively rapidly dissolving surfaces are assumed to be significant only during the early part of the experiment, and by the end of the experiment dissolution is presumed to proceed at a constant rate. Such an interpretation is predicated on the assumption that the overall dissolution reaction has a zeroth reaction order with respect to its constituent metals (e.g., constant-pH kaolinite dissolution rates are independent of aqueous Al and Si concentration). Although this seems to be the case for single-oxide minerals such as quartz (Berger et al., 1994), it is not generally the case for multioxide minerals and glasses; the constant-pH dissolution rates of kaolinite (Devidal et al., 1992, 1997; Devidal, 1994; Oelkers et al., 1994), alkali feldspars (Chou and Wollast, 1985; Oelkers et al., 1994; Gautier et al., 1994), kyanite (Oelkers and Schott, 1999), basaltic glass (Oelkers et al., 1999), enstatite (Oelkers and Schott, in press), forsterite (Pokrovsky and Schott, 2000), and smectite (Cama et al., 2000) were all observed to decrease in response to the increasing aqueous concentration of at least one of their constituent metals. In the absence of a true dissolution plateau, constant pHs far from equilibrium dissolution rates are consistent with the empirical expression given by (see also Lasaga et al., 1994; Oelkers, 1996)

The effects of ion identity and ionic strength on the dissolution rate of a gibbsitic bauxite

The influence of cation and anion identity and concentration, on the far from equilibrium dissolution rate of gibbsite, was studied at 298°K. Input solutions, with initial pH = 3.5 and variable salt type and concentration, were flowed at different rates, through columns packed with a unconsolidated gibbsitic bauxite from Los Pijigüaos-Venezuela ore deposit. It was observed cations Na +, K +, Mg 2+ and Ca2+ have no influence on the far from equilibrium dissolution rate. Anions have two different effects: concentration increases of monovalent anions (Cl −, NO 3 − and ClO 4 −) causes a decrease in the rate, as a function of [anion] (−0.11 ± 0.01); and increases of sulfate concentration causes an increase in the rate as a function of [SO 4 =] (0.4 ± 0.1). According to our calculations, these two effects have a remarkable influence upon the lifetime of gibbsite under weathering conditions. Based on Transition State Theory, it is proposed the experimental observations are due to an electrostatic effect on the activated complex (AC ♯) of the gibbsite dissolution reaction. For this AC ♯ the product of the charge of the involved chemical entities is negative. When SO 4 = participates in the AC ♯ the product of the charges switches to positive and therefore, the electrostatic interaction increase the dissolution rate. The dissolution rates are independent of the solution saturation degree below ΔGr = − 0.74 kcal/mol. It is inferred that the critical ΔGr is a constant of the solid, not affected by the solution characteristics, e.g., pH, ionic strength, cation and anion identities.

Does organic acid adsorption affect alkali-feldspar dissolution rates?

Consideration of alkali-feldspar dissolution rates 1 The term alkali-feldspar is used in the present study to denote those feldspars that dissolve via aluminum deficient precursor complexes [Oelkers, E.H., Schott, J., Devidal, J.L., 1994. The effect of aluminum, pH, and chemical affinity on the rates of aluminosilicate dissolution reactions. Geochim. Cosmochim. Acta, 58 (1994) 2011–2024.], which includes feldspars in the system NaAlSi 3O 8–KAlSi 3O 8, as well as plagioclases of compositions ∼An 80 or less [Oelkers, E.H., Schott, J., 1995b. Experimental study of anorthite dissolution and the relative mechanism of feldspar hydrolysis. Geochim. Cosmochim. Acta, 59: (1995) 5039–5053]. 1 taken from the literature suggests that in the absence of organic ligands and at far from equilibrium conditions they can be represented by r += k( a 3 H + / a Al +3 ) n , where r + refers to the far from equilibrium dissolution rate, k designates a rate constant, a i stands for the activity of the subscripted aqueous species, and n refers to a stoichiometric coefficient. Dissolution rates of these minerals have also been observed to increase substantially with increasing aqueous organic acid concentration. To assess if this latter effect is the result of either the change in aqueous aluminum activity due to aqueous aluminum–organic acid anion complex formation or organic acid anion adsorption to the mineral surface, dissolution rates predicted using the above equation are compared to measured alkali feldspar dissolution rates. A close correspondence between these predictions and the experimental data are found, providing strong evidence that organic acid anion adsorption does not affect alkali-feldspar dissolution rates.