Forsterite Dissolution Rates Research Articles

Diffusion control of quartz and forsterite dissolution rates

An established engineering model is used to identify conditions where diffusion controls the dissolution of quartz and forsterite in packed beds. The model shows that diffusion control is favored at low advection flux, large grain size, high temperature, and high pH (if the reaction consumes H+). Quartz dissolution is chemical reaction controlled for most geochemically reasonable combinations of temperature, grain size, and flow rate. On the other hand, forsterite dissolution rates can be diffusion controlled for typical advection fluxes, grain sizes, temperatures, and pH’s. The apparent activation energy for diffusion-controlled reactions in a packed bed is much higher than the <∼20kJ/mol value that is often used to identify diffusion controlled reactions. The models are quite general and can be adapted to deal with other mineral dissolution and precipitation reactions.

Forsterite dissolution rates in Mg‐sulfate‐rich Mars‐analog brines and implications of the aqueous history of Mars

AbstractHigh salinity brines, although rare on Earth's surface, may have been important in the geologic history of Mars. Increasing evidence suggests the importance of liquid brines in multiple locations on Mars. In order to interpret the effect of high ionic strength brines on olivine dissolution, which is widely present on Mars, 47 new batch reactor experiments combined with 35 results from a previous study conducted at 25°C from 1 &lt; pH &lt; 4 in magnesium sulfate, sodium sulfate, magnesium nitrate, and potassium nitrate solutions with ionic strengths as high as 12 m show that very high ionic strength brines have an inhibitory effect of forsterite dissolution rates. Multiple linear regression analysis of the data suggests that the inhibition in dissolution rates is due to decreased water activity at high ionic strengths. Regression models also show that mMg up to 4 m and mSO4 up to 3 m have no effect on forsterite dissolution rates. The effect of decreasing dissolution rates with decreasing aH2O is consistent with the idea that water acts as a ligand that participates in the dissolution process. Less available water to participate in the dissolution reaction results in a slower dissolution rate. Multiple linear regression analysis of the data produces the rate equation . Forsterite in dilute solutions with a water activity of one dissolves twice as fast as those in brines with a water activity of 0.8.

The experimental determination of hydromagnesite precipitation rates at 22.5–75ºC

Natural hydromagnesite (Mg5(CO3)4(OH)2·4H2O) dissolution and precipitation experiments were performed in closed-system reactors as a function of temperature from 22.5 to 75ºC and at 8.6 &lt; pH &lt; 10.7. The equilibrium constants for the reaction Mg5(CO3)4(OH)2·4H2O + 6H+ = 5Mg2+ + 4HCO3– + 6H2O were determined by bracketing the final fluid compositions obtained from the dissolution and precipitation experiments. The resulting constants were found to be 1033.7±0.9, 1030.5±0.5 and 1026.5±0.5 at 22.5, 50 and 75ºC, respectively. Whereas dissolution rates were too fast to be determined from the experiments, precipitation rates were slower and quantified. The resulting BET surface areanormalized hydromagnesite precipitation rates increase by a factor of ~2 with pH decreasing from 10.7 to 8.6. Measured rates are approximately two orders of magnitude faster than corresponding forsterite dissolution rates, suggesting that the overall rates of the low-temperature carbonation of olivine are controlled by the relatively sluggish dissolution of the magnesium silicate mineral.

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.

Systematic review of forsterite dissolution rate data

This paper demonstrates a method for systematic analysis of published mineral dissolution rate data using forsterite dissolution as an example. The steps of the method are: (1) identify the data sources, (2) select the data, (3) tabulate the data, (4) analyze the data to produce a model, and (5) report the results. This method allows for a combination of critical selection of data, based on expert knowledge of theoretical expectations and experimental pitfalls, and meta-analysis of the data using statistical methods.Application of this method to all currently available forsterite dissolution rates (0<pH<14, and 0<T<150°C) normalized to geometric surface area produced the following rate equations:For pH<5.6 and 0°<T<150°C, based on 519 datalogrgeo=6.05(0.22)-0.46(0.02)pH-3683.0(63.6)1/T(R2=0.88)For pH>5.6 and 0°<T<150°C, based on 125 datalogrgeo=4.07(0.38)-0.256(0.023)pH-3465(139)1/T(R2=0.92)The R2 values show that ∼10% of the variance in r is not explained by variation in 1/T and pH. Although the experimental error for rate measurements should be±∼30%, the observed error associated with the logr values is ∼0.5logunits (±300% relative error). The unexplained variance and the large error associated with the reported rates likely arises from the assumption that the rates are directly proportional to the mineral surface area (geometric or BET) when the rate is actually controlled by the concentration and relative reactivity of surface sites, which may be a function of duration of reaction. Related to these surface area terms are other likely sources of error that include composition and preparation of mineral starting material.Similar rate equations were produced from BET surface area normalized rates. Comparison of rate models based on geometric and BET normalized rates offers no support for choosing one normalization method over the other. However, practical considerations support the use of geometric surface area normalization. Comparison of Mg and Si release rates showed that they produced statistically indistinguishable dissolution rates because dissolution was stoichiometric in the experiments over the entire pH range even though the surface concentrations of Mg and Si are known to change with pH. Comparison of rates from experiments with added carbonate, either from CO2 partial pressures greater than atmospheric or added carbonate salts, showed that the existing data set is not sufficient to quantify any effect of dissolved carbonate species on forsterite dissolution rates.

An experimental study of magnesite precipitation rates at neutral to alkaline conditions and 100–200 °C as a function of pH, aqueous solution composition and chemical affinity

Magnesite precipitation rates were measured at temperatures from 100 to 200°C as a function of saturation state and reactive fluid composition in mixed flow reactors. Measured rates were found to increase systematically with increasing saturation state but to decrease with increasing reactive fluid aqueous CO32- activity and pH. Measured rates are interpreted through a combination of surface complexation models and transition state theory. In accord with this formalism, constant saturation state BET surface area normalized magnesite precipitation rates (rMg) are a function of the concentration of protonated Mg sites at the surface (>MgOH2+) and can be described using: rMg=kMg-KCO3KOHKCO3KOH+KOHaCO32-+KCO3aOH-n1-ΩMgn where kMg- represents a rate constant, KOH and KCO3 stand for equilibrium constants, ai designates the activity of the subscripted aqueous species, n refers to a reaction order equal to 2, and ΩMg denotes the saturation state of the reactive solution with respect to magnesite. Retrieved values of n are consistent with magnesite precipitation control by a spiral growth mechanism. The temperature variation of the rate constant can be described using kMg-=Aaexp(-Ea/RT), where Aa represents a pre-exponential factor equal to 5.9×10−5mol/cm2/s, Ea designates an activation energy equal to 80.2kJ/mol, R denotes the gas constant, and T corresponds to the absolute temperature. Comparison of measured magnesite precipitation rates with corresponding forsterite dissolution rates suggest that the relatively slow rates of magnesite precipitation may be the rate limiting step in mineral carbonation efforts in ultramafic rocks.

Oxalate-promoted forsterite dissolution at low pH

We ran a series of 124 semi-batch reactor experiments to measure the dissolution rate of forsterite in solutions of nitric and oxalic acid solutions over a pH range of 0–7 and total oxalate concentrations between 0 and 0.35 m at 25 °C. We found that the empirical rate law for the dissolution of forsterite in these solutions is r = 10 - 7.03 a H + 0.46 + 10 - 5.44 m ox 2 - 0.40 a H + 0.47 where a H + is the activity of hydrogen ions, m ox is the concentration of oxalate in molal, and the rate is given in (mol/m 2 s). The first term of the rate law expresses the dissolution rate for solutions with no oxalate and the second term expresses the additional rate produced by the presence of oxalate ions. Our experiments show that oxalate-promoted dissolution rates depend upon both oxalate concentration and pH. Based on this, we propose a reaction mechanism in which a hydrogen ion and an oxalate ion are simultaneously present in the activated complex for the reaction that releases H 4SiO 4 into solution. By analogy, we propose that water acts as a ligand in the absence of oxalate, so the first term of our rate law has a H 2 O m (=1) as a hidden term. Thus, our results suggest that both a proton and a ligand (either an anion or a water molecule) must be present in the activated complex for forsterite dissolution, suggesting that proton-promoted and ligand-promoted dissolution are simply two aspects of the same process, and that the distinction between the two is artificial. The increased dissolution rate of forsterite produced by oxalate ions at a concentration of 0.001 m, a concentration comparable to the amount of organic acids in typical soils, translates into an approximately a 6-fold increase in forsterite dissolution rates at pH > 4.2. This suggests that organic acids have a measurable, but not profound, effect on chemical weathering.

An experimental study of forsterite dissolution rates as a function of temperature and aqueous Mg and Si concentrations

Steady state forsterite dissolution rates, at far from equilibrium conditions and pH=2±0.04, were measured as a function of aqueous magnesium and silica concentrations, and temperatures from 25°C to 65°C. All rates were measured in mixed flow reactors and exhibited stoichiometric dissolution. Measured rates are found to be independent of both aqueous magnesium and silica concentrations. The temperature dependence of the pH 2 forsterite dissolution rates obtained in this study is consistent with an Arrhenius equation of the form r=A A exp(−E A /RT) where r signifies the overall forsterite steady state dissolution rate, A A refers to a pre-exponential factor equal to 0.190 mol cm −2 s −1, E A designates an activation energy equal to 63.8 kJ/mol, R represents the gas constant, and T denotes absolute temperature. The observed variation of forsterite dissolution rates with aqueous composition is interpreted to originate from its dissolution mechanism. The forsterite structure consists of isolated silica tetrahedra that are branched together by magnesium octahedra chains. MgO bonds apparently break more rapidly than SiO bonds in this structure. The breaking of octahedra chain linking MgO bonds, which is apparently catalyzed by hydrogen ion adsorption at acidic conditions, leads directly to the destruction of the mineral. As the rate-controlling precursor complex for this mineral is formed by a hydrogen adsorption reaction, forsterite dissolution rates are unaffected by aqueous Mg and Si activities.

A high resolution study of forsterite dissolution rates

The rates of dissolution of olivine (Fo 92) were measured over the pH range of 1.8 to 3.8 and from 25 to 45°C using an externally recycled mixed flow reactor. The data were filtered to remove several sources of systematic errors including non-steady state measurements, solution concentrations that were less than 50% above the detection limit, and Mg/Si rate ratios 40% (∼3σ or 98% confidence limit) greater or smaller than 1.84, which is the stoichiometric ratio of these elements in the mineral used. The effect of random errors on calculated values of E a and n were minimized by collecting a large amount of data. Twenty-eight different runs produced 772 rate measurements of which 284 were rejected in order to remove systematic and large random errors, leaving 488 data. The errors associated with these remaining data are normally distributed. The data were regressed to the equation log r- log A− E a 2.303R 1 T −n pH to find that log A = 0.54 ± 0.14, E a = 42.6 ± 0.8 kJ/mol, and n = 0.50 ± 0.004 (R 2 = 0.97). Based on these results, we postulate that the activated complex for forsterite dissolution involves a hydronium ion adsorbed to the forsterite surface in such a way that two of the three hydrogen atoms are associated with two Si-O-Mg bridging oxygen atoms. In this configuration, the adsorbed hydronium ion dissociates to form H-O-Si bonds as water molecules move from the solution to hydrate the two Mg atoms, breaking the Mg-O-Si bonds. This dissociation of the hydronium ion to form two H-O-Si groups on the forsterite surface leads to the observed reaction order for hydrogen ions of 0.5. The rate of breakdown of this activated complex is proportional to the rate at which the water molecules move from the solution to hydrate the Mg ions. Several other silicate mineral dissolution reactions, which have reaction orders of 0.5, may involve similar activated complexes.