Silica dissolution reaction rates

During caustic waterflooding the alkali can be consumed by the dissolution of clays and is lost in this way. The amount lost depends on the kinetics of the particular reaction. Several studies have been performed with kaolinite, using quartz as a yardstick, because the kinetic data are documented in the literature. The initial reaction rate has been found pH independent in the pH range of 11 to 13 [517]. The kinetics of silica dissolution could be quantitatively described in terms of pH, salinity, ion-exchange properties, temperature, and contact time [1549]. 

Ionic Medium. Silica dispersions were freshly prepared for each experiment in solutions buffered with 10"3M HC03"/C02. The amount of species dissolved from the amorphous silica surface during the experiment was negligible because of the small rate of dissolution reactions. The ionic medium in which coagulation and adsorption studies were carried out was kept constant I = 1.0 to 2.0 X 10 3M. The conditions in all agglomeration and adsorption experiments were such that no Al(OH)3 precipitated within the period of observation. 

Nonetheless, the general understanding of magnesium and calcium carbonation reactions has improved significantly (see also the studies by Hanchen et al. [107-110] on the relative importance of process parameters such as temperature, C02 pressure and particle size distribution). Studies involving a three-step process of olivine carbonation, involving (i) dissolution of olivine (ii) precipitation of magnesite and (iii) precipitation of silica in an aqueous solution, were recently reported from Norway [69], where the process proceeds without chemical additives at 10-15 MPa and 403-523 K. No reaction rates were reported, however. 

At present, many studies are ongoing to identify a means of enhancing the carbonation chemistry of magnesium silicates in aqueous systems, using weak acids and additives that will improve silica dissolution, such as citrates, oxalates, and EDTA [105, 111]. In this case, a near-complete recovery and reuse, thereby minimizing the losses of such chemicals, will be essential for viable process economics. Likewise, there is much to improve with regards to the reaction rates and/ or times. 

Tiemann (T8) studied the dissolution of silica from a siliceous iron ore by sintering the ore with sodium carbonate followed by leaching the sodium silicate with water. The reaction rates were found to be low after sintering for 4 hr at 1450°F. The residual concentrate was analyzed to be 56% iron, corresponding to 88% dissolution of the silica. Partial to complete fusion resulted when the temperature was increased. 

Stone et al. (S29) developed by a mathematical analysis the functional relationship between the rate of extraction of silica from pure quartz in sodium hydroxide solution and time, temperature, sodium hydroxide concentration, and particle size. With the use of response surface methodology, a comprehensive picture of this dissolution process was obtained from a few well-chosen experiments. The fractional extraction of silica can be expressed by a second-order equation. The effect of quartz particle size and temperature are predicted to be about equal and greater than the influence of sodium hydroxide concentration and reaction time. The reaction rate is controlled by the surface area of the quartz. An increase in sodium hydroxide concentration increases the activation energy for the reactions and is found to be independent of quartz size. 

That nucleation and growth rate are the limiting steps in quartz cementation has no particular imphcations with respect to the ultimate source of the sihca or the mechanism of transport. Potential sources of silica for quartz cementation are numerous (McBride, 1989) and include all documented silicate dissolution reactions in sandstones and shales. 

Figure 2,5 (a) An Arrhenius plot of log k versus I/TXK) for the dissolution rates of various silicate rocks and minerals. The data points and curves for rhyolite, basalt glass, and diabase are from Apps (1983), as is the curve labeled silicates, which Apps computed from the results of Wood and Walther (1983). Curves for the S1O2 polymorphs are based on Rimstidt and Barnes (1980). Modified from Langmuir and Mahoney (1985). Reprinted from the National Well Water Assoc. Used by permission, (b) An Arrhenius plot of log k versus 1 /T(K) for the precipitation of quartz and amorphous silica based on Rimstidt and Barnes (1980). Reprinted from Geochim. Cosmochim. Acta, 44, J.D. Rimstidt and H.L. Barnes, The kinetics of silica water reactions, 1683-99, 1980, with permission from Elsevier Science Ltd, The Boulevard. Langford Lane. Kidlington OXS 1GB, U.K. 

The rate-controlling step for dissolution of an oxide or primary silicate mineral generally involves a surface reaction. For surface-controlled dissolution, the rate-controlling step is either the detachment of silica or a metal ion from the surface or the attack of the surface to form precursor sites for detachment. Surface detachment controlled kinetics can be modelled using the surface complexation rate model (Wieland et al., 1988) that models rates as a function of the surface concentration of surface complexation sites that are precursors for dissolution. In this model, the formation of precursor sites is rapid compared to the rate of detachment and the concentration of sites can be described by surface complexation theory (Sposito, 1983). 

This phenomenon was confirmed by the introduction of symmetric tetraalkylammonium hydroxides in the dissolution of silica gel. In TMAOH the observed rate of dissolution was slow compared to the rate observed for cesium hydroxide dispersions, and cesium hydroxide has the lowest rate for the different alkali metal hydroxides. Results in Figure 3 clearly reveal an inhibition time between mixing of the silica gel with the aqueous TMAOH and the onset of dissolution. This observation is attributed to the strong interaction of the rather apolar TMA cation with the negatively charged silica gel surface. Because in this case no hydration shell is present, dissolution only occurs very slowly. The observed inhibition period of the dissolution reaction can be related to specific interactions of TMA cations with relatively large oligomeric species of the monomeric... 

In a study [6] of the dissolution of amorphous silica gels in aqueous alkali metal hydroxides, the rate of dissolution was found to depend on the cation used in the dissolution reaction. A maximum in dissolution rate was found for potassium hydroxide solutions, whereas both intrinsically smaller and larger cations (lithium-sodium and rubidium-cesium) showed slower dissolution rates, as can be concluded from the concentration of dissolved silicate species (normalized peak areas) as a function of alkali metal cation (Figure 45.2). This result is contradictory to the expectation that a monotonic increase or decrease in dissolution rate is to be observed for the different cations used. One major effect that occurs at the high pH values of this study is that the majority of silanol... 

More elaborate and ambitious studies on the dissolution reactions of silica were conducted by Xiao and Lasaga (1994, 1996). Their objective was to provide full descriptions of the reaction pathway of quartz dissolution in acidic and basic solutions, from the adsorption of H2O or OH on a site, the formation of possible reaction intermediates and transition states, to the hydrolysis of the Si-O-Si bonds. Also, their aim was to extract kinetic properties such as changes in activation energy, kinetic isotope effects, catalytic and temperature effects, and the overall rate law form. The reaction mechanisms investigated were... 

Her examined the pH-tltration behavior of silicic acid in the presence of 2-hydroxypyridine 1-oxide by titrating 16 mA/ (1000 ppm SiO ) silicic acid silica from pH 10.5 to 3.0 in the presence and absence of a 43 mA/ concentration of the N-oxide. At no point did the titration curves differ, indicating that no complex had formed. In another experiment, a solution of Si(OH) containing 100 ppm as SiOj was mixed with a 200-fold excess of the above -oxidc at pH 6.15 and aged for a few hours. Tests with molybdic acid showed that the reaction rate with silica monomer was the same as a control, indicating either that no complex was formed at this pH or that it dissociated very rapidly. However, the rate of dissolution of monomer from colloidal silica particles at pH 1.4 was apparently doubled in the presence of a 20 mA/ concentration of the yV-oxide, indicating some type of interaction at low pH. -... 

Instead of an alkaline solution, a dilute solution of NaF and HCl to generate HF was employed by Goto (216a) to measure the specific surface area of colloidal silica from the rate of dissolution in this acid medium. The samples from the reaction mi.x-ture were removed, the reaction was stopped by adding aluminum salt to combine with the fluoride, and the dissolved silica was then determined by the molybdic acid method. 

The high surface area and rate of dissolution of amorphous silica permits reactions at much lower temperatures than required with pulverized crystalline silica. The high chemical reactivity of colloidal silicas thas been reviewed in Chapter 4. Transparent fused silica can be formed at 2000 psi and 1200 C from silica powder of 15 nm ultimate particles, whereas 2000 C is required for blown or cast material (655). 

Fukui et al. [33] have investigated the effects of NaOH concentration on the crystal stracture and the rate of reaction of the synthesized zeolite from fly ash with a hydrothermal treatment method. They have reported that fly ash or the mixture of fly ash and silica powder results in an increase in the reaction rates with the increase of NaOH concentration due to increase of the dissolution rate of silicate ion and aluminate ion. It has been clarified that the NaOH concentration also affects the crystal structure of synthesized zeolites. It has been concluded that the proportion of Phillipsite continues to be lower than the increasing proportion of Hydroxy-sodalite in the product with the increase in the concentration of NaOH. 

In Chapter 16, we wrote rate laws for simple dissolution and precipitation reactions, such as those for the silica minerals forming from SiC>2(aq). Rewriting Equation 16.22 in terms of volumetric concentration C , assuming the activity coefficient Yi does not vary over the reaction, gives the rate law,... 

In the calculation results (Fig. 26.6), the initial segment of the path is marked by the disappearance of the amorphous silica as it reacts to form cristobalite. The amorphous silica is almost completely consumed after about 10000 years of reaction. The mineral s mass approaches zero asymptotically because (as can be seen in Equation 26.1) as its surface area As decreases, the dissolution rate slows proportionately. During the initial period, only a small amount of quartz forms. 

Fig. 2. Plot of normalized rate vs. the activity of silicic acid for the LAWABP1 (see Table 1) glass composition at two temperatures (26 and 40 °C). Rates are all computed at steady-state conditions. Boron and Na release rates are identical at low silica activities, then decrease, and become constant at or near saturation with respect to amorphous silica (vertical dot-dashed line). Note that the B rate decreases more than the Na rate. This behaviour can be rationalized as competition between two concurrent reactions alkali-hydrogen exchange and matrix dissolution (see text). Error bars represent 2-

Finally, our observations regarding the longterm impact of alkali ion exchange on glass dissolution now provide a mechanistic basis for the empirical residual rate of reaction appended to the TST rate law articulated by Grambow (1985). The residual rate was appended to prevent calculated glass dissolution rates from dropping to zero under silica-saturated conditions, which is not in accord with experimental observations. 

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