Enhancing Silicate Dissolution Kinetics in Hyperalkaline Environments

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C: Surfaces, Interfaces, Porous Materials, and Catalysis

Enhancing Silicate Dissolution Kinetics in Hyperalkaline Environments Erika Callagon La Plante, Tandre Oey, Yi-Hsuan Hsiao, LaKesha Perry, Jeffrey W. Bullard, and Gaurav N. Sant J. Phys. Chem. C, Just Accepted Manuscript • DOI: 10.1021/acs.jpcc.8b12076 • Publication Date (Web): 15 Jan 2019 Downloaded from http://pubs.acs.org on January 21, 2019

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The Journal of Physical Chemistry

Enhancing Silicate Dissolution Kinetics in Hyperalkaline Environments Erika Callagon La Plante1,4,5*, Tandré Oey1,4,5, Yi-Hsuan Hsiao1,4,5, LaKesha Perry2, Jeffrey W. Bullard2, Gaurav Sant1,3,4,5 1

Laboratory for the Chemistry of Construction Materials (LC2), Department of Civil and Environmental Engineering, University of California, Los Angeles, CA, USA 2 Engineering Laboratory, National Institute of Standards and Technology (NIST), Gaithersburg, MD, USA 3 Department of Materials Science and Engineering, University of California, Los Angeles, CA, USA 4 California Nanosystems Institute (CNSI), University of California, Los Angeles, CA, USA 5 Institute for Carbon Management (ICM), University of California, Los Angeles, CA, USA * Corresponding author: Erika Callagon La Plante ([email protected])

Abstract The dissolution of silicate minerals and glasses in aqueous solutions is important in many natural and engineered contexts including mineral weathering, nuclear waste stabilization, cementation, and infrastructure degradation. The influences of electrolytes on dissolution rates have been extensively studied but previous studies have used widely varying minerals and electrolytes, experimental conditions, and measurement techniques. Comparatively fewer studies have been conducted in hyperalkaline solutions that are encountered in concrete, geopolymers, and nuclear stabilization systems. This study seeks to control many of these variables to isolate the effects of electrolyte composition on altering the degree of dissolution of a soda lime silicate glass in hyperalkaline electrolyte solutions. A glass powder is dissolved in one of a series of static solutions having 10 mmol L-1 NaOH (pH 12) at (25 ± 0.2) °C and either an organic or inorganic electrolyte at a concentration of 1 mmol L-1 or 10 mmol L-1. The solution series was designed to reveal qualitatively how ion identity and concentration alter the glass’ degree of dissolution after prescribed exposure times. The results indicate that a relative dissolution enhancement of no greater than about 2.4 times can be induced at pH 12, with the greatest enhancements being observed for Na-benzoate, Na-citrate, and Na-malonate. The degree of dissolution is unaffected by most of the other salts examined. Broadly, the nucleophilic attack by OH- on Si–O bonds, and the formation of O-–Na+ surface complexes appear to be the most important factors influencing the dissolution rate at high pH. The adsorption and influence on dissolution of other electrolyte ions is comparatively weak.

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1. Introduction The dissolution of silicates is integral to many processes such as geological weathering and infrastructure degradation,1,2 and has therefore inspired extensive research in geochemistry3–7 and materials science.8–13 The mechanisms and rates of silicate dissolution have been studied in terms both of the solid’s properties, such as its composition and crystallinity,14 and of solution characteristics, such as pH, temperature, saturation states, and the nature and concentration of aqueous species.15–18 The goal of such studies is often to identify ways to enhance or inhibit dissolution in particular applications. For instance, reduced dissolution rates of silicates are vital both in stabilizing nuclear waste repositories and in minimizing the degradation of concrete caused by alkali-silica reaction.19–21 On the other hand, the enhanced weathering of silicate minerals has been proposed as a means to mitigate ocean acidification from increased atmospheric CO2 concentrations.22,23 In the latter situation, greater silicate dissolution rates may also stimulate the production of biomass through the dissolved cations.22 The enhanced dissolution of (alumino)silicates such as fly ash is also desired for their enhanced reaction when such materials may be used to replace ordinary portland cement (OPC) – the CO2-intensive binding agent that is commonly used in concrete24 to lower the material’s environmental impact. Most studies of silicate dissolution have been performed under acidic or circumneutral conditions relevant to surface or groundwater environments, but many engineering systems have more alkaline solutions, such as the pore solution of concrete used in the construction of nuclear power plants and nuclear waste repositories.25 Silicate dissolution in alkaline solutions is caused by nucleophilic attack on positive charge centers of the polar Si+– O- bond by OH- species.26,27 This paper will refer to this bond-breaking process as hydrolysis. Hydrolysis has been assumed to be rate-limiting,26 although the absolute rate can be modified by ion sorption or by changes in local electrochemical potential.29–32 Chemisorption involves the bonding of a solute species to a surface site, which induces local changes in electronic structure and may thereby change the susceptibility of nearby bonds to hydrolysis.33,34 Physisorption is the physical association of a solute species with a surface site, which does not alter the electronic structure but may still change (usually hinder) the hydrolysis of nearby bonds by sterically impeding the access of water or OH- to the siloxane bond.35 Cation-oxygen bonds at the mineral surface may be weakened by ion sorption if charged surface complexes further polarize the siloxane bonds.36 On the other hand, the electrochemical surface potential, which depends upon both the surface charge as well as the solvent’s composition, may affect the transport of dissolved species within the double layer. Each of these mechanisms is likely influenced by the nature and concentration of solvated ions, specifically their particular charges, sizes, and mobilities. Dissolution rates can be enhanced either kinetically, by increasing the apparent rate constant, or thermodynamically by increasing the apparent solubility. Previous studies have shown that organic and inorganic electrolytes can enhance the dissolution of silicates in either way.28,32,38–40 Inorganic salts are thought to alter silicate dissolution rates by solute adsorption, whereas organic salts are more commonly found to increase the apparent solubility by complexation with dissolved Si.38,41 For example, the dissolution rate of amorphous silica can be enhanced by up to 21 times by adding 50 mmol L-1 NaCl at circumneutral pH.40 Under hydrothermal conditions, alkaline species can increase the dissolution rate of quartz by up to 40 times.39 However, dissolution rates of more complex silicate minerals are less affected by electrolyte composition, typically by no more than a factor of five.39,42–44 Prior research39 suggests that cations increase


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the rates of quartz dissolution by increasing the cation’s concentration and water exchange rates near the surface, which increases the hydrolysis reaction attempt frequency. If so, silicate minerals containing alkali or alkaline earth cations should already have rapid surface exchange rates with water because of those cations, thereby reducing the effect of additional dissolved cations.39 Furthermore, such multi-cation silicates have an interfacial solvent layer that is highly disordered due to the presence of cations that adsorb after their detachment from the solid.32 Although there are substantial data on the effects of dissolved electrolytes on dissolution, most studies examine only a few electrolytes and are often conducted under disparate conditions or while utilizing differing means of dissolution rate assessment. In contrast, this work evaluates 28 organic or inorganic salts in hyperalkaline (pH 12) aqueous solutions in terms of their ability to enhance the hydrolysis of silicon in a soda lime silica glass network. The resulting changes in the extent of dissolution at prescribed time intervals indicate a large influence of pH, as expected, and also provide insights into strategies that may be used to enhance silicate reactivity, if so, under hyperalkaline conditions in which OH- serves as both the potential/pH determining ion. 2. Materials and methods 2.1. Materials Soda lime silicate glass microspheres (Sigmund Linder) were used as a model soda-lime silicate glass with a composition similar to that observed in commercial fly ashes.45 This solid was used to compare changes in its tendency to dissolve in aqueous solutions across a series of additives in solutions that are highly undersaturated with respect to the glass. The solid’s elemental composition in mol %, as determined by scanning electron microscopy with energy dispersive Xray spectroscopy (SEM-EDS) is: 24.35(27) % Si, 10.33(01) % Na, 0.16(08) % K, 3.61(02) % Ca, 0.34(32) % Al, 1.65(05) % Mg, and 59.55(07) % O, where the values in parentheses represent one standard deviation in the last two significant digits. The specific surface area (SSA) of the glass microspheres is (92 ± 7) m2/kg as determined by N2 sorption (Micrometrics ASAP 2020 BET analyzer)* on samples that were stored under vacuum for two hours at 300 °C prior to analysis. The particle size distribution (PSD) of the spheres was measured by laser diffraction spectroscopy using a Beckman Coulter LS 13-320 light scattering analyzer. The powder was sonicated in isopropanol for effective dispersion. The particle size analysis (Figure 1) indicates a median diameter (d50) of 53 µm. The density of the glass is 2.5 g/cm3, as specified by the manufacturer. A reference solution with pH 12 was prepared by adding 10 mmol L-1 NaOH to deionized (>18 MΩ·cm) water. Table 1 lists the salts examined in the study, all of which are ACS reagent grade or higher. The salts were chosen to span three broad categories: (1) varying alkali and alkaline earth cations and either Cl- or NO3- anions; (2) varying halide and oxy-anions with a Na+ cation; and (3) varying organic anions with a Na+ cation. Each salt was added to the reference solution by diluting a concentrated stock solution to either 1 mmol L-1 or 10 mmol L-1. Thermodynamic calculations using PHREEQC46 with the minteq.v4 database indicate that all the solutions, except for those with Mg-containing salts, are undersaturated with respect to all phases at these *

Certain commercial equipment, instruments, or materials are identified in this paper to foster understanding. Such identification does not imply recommendation or endorsement by the University of California, Los Angeles (UCLA) or National Institute of Standards and Technology (NIST), nor does it imply that the materials or equipment identified are necessarily the best available for the purpose.


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concentrations. The measured pH of the experimental solutions was 12.0 ± 0.1 at 23 °C. The dependence on pH of the degree of dissolution was also measured in the absence of any other salts by varying the NaOH concentration. 100

25

80

20

60

15

40

10

20

5

0

0 0

25

50 75 100 125 Particle Diameter (µm)

150

Figure 1. The particle size distribution of the soda lime silicate glass beads used in this study. 2.2. Dissolution Method Glass powder was added to each solution at a liquid:solid mass ratio of 100 and stored in highdensity polyethylene (HDPE) containers for up to 168 h . The containers were either shaken after every sampling (solutions with added salts) or unstirring (solutions with varying pH) at (25  0.2) °C. A 6 mL solution aliquot was collected at times of (0, 1, 3, 6, 24, 72, 168) h, passed through a 0.2 μm filter to remove residual solids, and its chemical composition was measured by inductively coupled plasma optical emission spectrometry (ICP-OES) as described in the next section. The extent of saturation of the shaken or unstirred solution with respect to the glass increases continuously with time. The dissolution rate therefore also varies continuously and is likely controlled by the transport of dissolved ions away from the surface. Despite these complications, this study characterizes the relative dissolution tendencies in terms of a time-averaged, surfacenormalized dissolution rate, as calculated by linear regression of plots of dissolved silicon concentration versus time; by dividing the slope of the regression line by the initial BET surface area†. This way of quantifying the dissolution tendency provides a convenient way of ranking the efficacy of the various electrolytes, which is the objective of the study. In the remainder of the paper, this time-averaged rate will be called the dissolution rate for brevity.

†

The method used to determine surface area may influence comparability with other studies as reported by Fournier et al..47 This is because glass specific surface areas determined geometrically, e.g., using particle size data and assuming spherical particles are smaller than those measured by N2-sorption/BET analysis by a factor of about 2.2 to 2.5.


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Table 1. List of chemical additives and the corresponding abbreviations used in the text. Purities indicated in parentheses are as reported by the manufacturer. Alkali/Alkaline Earth Salts Inorganic Sodium Salts Organic Sodium Salts Name Abbreviation Name Abbreviation Name Abbreviation Lithium LiCl Sodium NaF Sodium NaAc Chloride (>98.5%) Fluoride (>99%) Acetate (>99%) Sodium NaCl Sodium NaCl Sodium NaPr Chloride (>99%) Chloride (>99%) Propionate (>99%) Potassium KCl Sodium NaBr Sodium NaBz Chloride (>99%) Bromide (>99%) Benzoate (>99%) Magnesium MgCl2 Sodium NaI Disodium NaMal Chloride (>99%) Iodide (>99%) Malonate (>99%) Calcium CaCl2 Sodium Na2CO3 Disodium NaSuc Chloride (>99%) Carbonate (>99.5%) Succinate (>99%) Lithium LiNO3 Sodium Na2SO4 Disodium NaPht Nitrate (>99%) Sulfate (>99%) Phthalate (>99%) Sodium NaNO3 Sodium NaLac Sodium NaNO3 (>99%) Nitrate (>99%) Lactate (>99%) Nitrate Potassium KNO3 Sodium NaNO2 Sodium NaSal Nitrate (>99%) Nitrite (>97%) Salicylate (>99%) Magnesium Mg(NO3)2 Disodium NaTar Nitrate (>98%) Tartrate (>99.4%) Trisodium NaCit Calcium Ca(NO3)2 (>99%) Citrate (>99%) Nitrate 2.3. ICP-OES Elemental concentrations of Si, Al, Ca, K, and Na in solution aliquots were measured by ICPOES using a Perkin Elmer Avio 200 instrument. The filtered solution was diluted in a 5 vol% nitric acid matrix. Three spectra were collected for each sample and converted to molar concentration units by interpolation with standard solutions. The relative standard deviation (RSD) for the replicates is about 3 %. Based on this standard deviation, the relative uncertainty in dissolution rates as calculated in the previous section is estimated to be 15 %. 2.4. Zeta potential The zeta potential of the particles in several of the solutions was measured using suspensions containing 0.01 % solid by mass. The measurements were conducted using a ZetaPALS Potential Analyzer (Brookhaven Instruments Corporation) and were initiated within the first 10 min after immersion of the glass microspheres in solution. The zeta potentials were measured in solutions of sodium halides and alkali and alkaline earth chlorides and nitrates. The uncertainties in the zeta potentials reflect the RSD for 10 replicate measurements for the same sample. 3. Results and discussion 3.1. Dissolution of soda lime silicate glass Dissolution rates were characterized by the slope of a linear regression of the concentration of Si in solution as a function of time, divided by the initial glass BET surface area. In solutions with


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pH > 10, an approximately linear time dependence was observed for the 168 h duration of the experiment. However, a transient nonlinear time dependence was observed for the first 24 h in solutions with pH ≤ 10, as shown in Figure 2, suggesting that in the early stages of dissolution the diffusion of Si or water at the glass surface is rate-controlling.48 In those cases the linear regression was performed within the approximately linear period after 24 h. The reference dissolution rate at pH 12 and (25 ± 0.2) °C with occasional agitation is 2.85  10-9 mol/m2/s, and without stirring is 1.44  10-9 mol/m2/s, as inferred from the slope of the regression line. The difference between these rates indicates mass transport control; nonetheless the relative dissolution rates under the same conditions, i.e., static, or occasionally agitated, can still be compared. The congruency of dissolution was also assessed by comparing the solution composition to that of the glass. Dissolution is congruent if the molar ratios of the dissolved species in solution match those in the solid. Differences between solid and solution compositions is often observed during dissolution of mixed oxide glasses49 and can be caused either by incongruent dissolution, that is the preferential leaching of specific cations from the silicate network to form a silicate-rich surface layer,11,50 or by the formation of secondary precipitates that consume ions released from the solid. In fact, some evidence of precipitation is observed in these experiments and during similar experiments on borosilicate glass (Figure 3c), as will be discussed shortly. The solid and solution compositions in this study appear to converge only after 7 days in alkaline solutions (Figure 3), with some preferential leaching of Al and Ca being observed in most solutions at earlier times. The lone exception is at pH = 10, at which the glass at early times preferentially retains Al relative to Si. An initial preferential release of Na relative to Si was also observed at pH 10 and 12, and dissolution approached congruency with respect to Na at longer reaction times. The apparent preferential release of Al and Ca at pH > 10 (see Figure 3a,b) indicates the initial formation of a Si-rich altered surface layer.51 In any event, the molar ratios in the solid and solution approach to within a factor of two in most experiments, which suggests that any Si-rich alteration layers formed early and did not continue to grow appreciably over longer reaction times, at least not under hyperalkaline conditions. The lone exception is at pH 13, where the solution has a greater Si/Ca then the bulk solid (Figure 3b). The modified surface layer may be Ca-rich or perhaps Ca is initially precipitated from solution in greater proportions than Si to form a poorly ordered calcium silicate hydrate, for which typically 0.7 ≤ Ca/Si ≤ 2.0. The rate of Si release was constant after 24 h at high pH (Figure 2). If a surface layer does form, then either its growth stops within 24 h or, less likely, it offers no significant resistance to the diffusion of dissolved species.


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Figure 2. The concentration of Si in solution normalized by the initial BET surface area of the solid as a function of time. The dashed black lines indicate linear regressions to either the entire data range (pH 11 to pH 13) or to the data obtained between 24 h and 168 h of dissolution (pH 5.6 to pH 10). Reported values have a relative uncertainty of about 15 % based on the relative standard deviations of repeated measurements of both concentration and specific surface area. 3.2. Effect of electrolytes on silicate dissolution rates To provide a simple basis for comparison of electrolyte effects in the sections that follow, Figure 4 summarizes the dissolution rates normalized by the steady-state glass dissolution rate of 2.85  10-9 mol/m2/s at pH 12, with occasional agitation, in the absence of other dissolve salts. The figure shows that most salts with univalent cations did not significantly affect the dissolution rates. Notable exceptions include Na-benzoate, Na-malonate, and Na-citrate. In addition, chlorides and nitrates with divalent cations significantly reduced the dissolution rates. The inhibiting effect of Mg and Ca on polymerized silicate (quartz) dissolution under alkaline conditions has been predicted by ab initio quantum chemical simulations, which suggest that the divalent cation increases the energy barrier for Si–O bond rupture.52 The formation of Si–O– Mg2+ or Si–O–Ca2+ may also inhibit the dissociation of water molecules and the subsequent hydrolysis of terminal Si–O bonds.52 At high pH, however, a more plausible explanation for the retarding effects of Mg and Ca is the precipitation of secondary phases, such as Ca(OH)2, Mg(OH)2, or Mg- and Ca-silicate hydrates, which would in time transform to a mixture of hydrous silica and calcite in the presence of dissolved CO2, as suggested by the formation of surface precipitates on borosilicate glass surfaces as shown in Figure 3c. This observation is consistent with previous studies showing the precipitation of C-S-H on an intermediate-level (nuclear) waste glass exposed to Ca-rich solutions at pH > 11, decreasing its dissolution rate.53 Similarly, a high-level waste glass dissolved by an order of magnitude slower in Ca(OH)2saturated solutions than in water, an observation which has also been attributed to the formation of C-S-H phases.54 Under similar aqueous conditions, C-S-H precipitates also formed on the surface of an international simple glass (ISG).55


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(a) (b) (c) Figure 3. The (in)congruency in glass dissolution as assessed by the evolution of: (a) Si/Al or (b) Si/Ca molar ratios in the solution as compared to the solid composition. (c) The surface of a sodium borosilicate glass (i.e., having mass-based composition: 80.6% SiO2, 12.6% B2O3, 4.2% Na2O, 2.2% Al2O3) particulate imaged using scanning electron microscopy (SEM) in secondary electron mode following exposure to 10 mmol L-1 NaOH + 100 mmol L-1 Ca(NO3)2 containing dissolved CO2 shows the presence of calcite (CaCO3) precipitates on its surface as confirmed by X-ray microanalysis and thermogravimetric analysis. Thermodynamic calculations using PHREEQC46 indicate that the 1 mmol L-1 and 10 mmol L-1 MgCl2 and Mg(NO3)2 solutions are supersaturated with respect to brucite (Mg(OH)2). The logarithm of the saturation index, SI, with respect to brucite (log10[IAP/Ksp], where IAP is the ion activity product and Ksp is the solubility product), is 3.27 in the 1 mmol L-1 MgCl2 and 3.81 when the concentration is 10 mmol L-1. Increasing the Mg concentration also decreases the pH, calculated as 11.92 and 11.62, respectively, due to the consumption of OH- to form MgOH+(aq). However, the extent of pH decrease may be even greater, the pH being about 9.7 at the higher Mg concentration when brucite reaches equilibrium with the solution. In contrast, the solutions with Ca salts all are undersaturated with respect to portlandite (Ca(OH)2). Therefore, the inhibition of rates is likely not caused by portlandite precipitation. Interestingly, dense layers of Ca-rich precipitates have been observed on alkali-silicate glass surfaces submerged in NaOH and Ca(OH)2 solutions, although those observations were made under more alkaline conditions (pH 14) and at higher temperatures of 60 ºC to 80 ºC.56 The low observed solution concentrations of Ca and Mg in that study suggest that these elements are bound in surface corrosion products.56 The current study also measured a systematic decrease in Ca concentration with time in solutions with added CaCl2 and Ca(NO3)2, which is consistent with the sequestration of Ca at the surface. Separate experiments on borosilicate glass dissolution revealed calcite crystallites at glass surfaces submerged in Ca(NO3)2 solutions with some dissolved CO2. The number density of crystallites, examples of which are shown in Figure 3c, increased with Ca concentration in solution. PHREEQC calculations indicate that those solutions are also supersaturated with respect to calcium silicate hydrates, consistent with the observations of Maraghechi et al,56 although dissolved CO2 renders the calcium silicate hydrates metastable with respect to a mixture of hydrous silica and calcite. Therefore, the observed rate reduction in the presence of soluble Ca- and Mg- salts is likely attributable to a combination of pH reduction and precipitation on the glass surfaces.


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Figure 4. The dissolution rates of soda lime silicate glass at pH 12 in the presence of different salts at 1 mmol L-1 and 10 mmol L-1 concentrations, relative to a control case with no electrolytes present. Measured zeta potentials are shown in Figure 5 and can be used to roughly infer the extent of ion sorption on the glass surfaces and the consequences for surface chemistry. The zeta potential is the electric potential, relative to that of the bulk solution, measured at the shear plane in the electric double layer, and is proportional to the electrostatic forces between particles in a suspension. Although distinct from the electric surface potential, the zeta potential nevertheless describes the particles’ effective relative electrostatic potential .57 The zeta potential of the soda lime silicate glass particles in a pH 12 solution without other dissolved salts is (-103.5  3.5) mV. For comparison, that of amorphous silica in a similar solution is (-37.0  3.0) mV. Addition of inorganic electrolytes generally decreases the magnitude of the zeta potential (Figure 5) and therefore indicates the preferential association of cations and compression of the electrical double layer upon increasing the ionic strength. These effects are enhanced in solutions with higher additive concentrations. Moreover, potential reversals were observed in solutions of Ca or Mg chlorides and nitrates. Cations with greater charge are known to more strongly affect the zeta potential58–60 because they have smaller ionic radii and can therefore congregate more densely at particle surfaces.61 The Debye length decreases from 3.0 nm to 1.8 nm or 1.2 nm upon adding 20 mmol L-1 of a monovalent or divalent electrolyte, respectively, in 10 mmol L-1 NaOH with pH 12.62 Thus, at the same electrolyte concentration, the observed differences in the zeta potentials are caused primarily by the interaction of the specific ions with the glass surface. A correlation between zeta potential and dissolution rate might be expected if the zeta potential is controlled by the solid’s intrinsic surface charge, as has been observed for some silicates as a function of pH.31 No such correlation was observed in this study, but some insights into the controls on average dissolution rate can still be obtained. The most telling aspect of Figure 5 is that the greatest changes in zeta potential are caused by the same salts that cause the greatest reductions in dissolution rate—MgCl2, Mg(NO3)2, CaCl2, and Ca(NO3)2. Counterion migration toward the particle surfaces could be responsible for the trends observed in Figure 3, but the trends could also be explained by precipitation of surface layers of a different composition, which would also generally hinder dissolution of the underlying glass. For example, zeta potentials measured in the Ca- and Mg- salt solutions are similar, and are also similar to those of calcite63,64 and of brucite65,66 in Ca- and Mg-rich solutions, respectively. Zeta potential alone is


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not a means of compositional identification, but the values measured in this study are at least consistent with the presence of these secondary phases.

Figure 5. The measured zeta potentials for soda lime silicate glass particulates suspended in 10 mmol L-1 NaOH solutions containing 2 mmol L-1 (light) and 20 mmol L-1 (dark) inorganic electrolytes. In general, the addition of electrolytes decreased the magnitude of the zeta potential from the reference value of (-103.5  3.5) mV (dashed red line). 3.3. Enhancement of silicate dissolution rates and the effect of pH Separate experiments on the soda lime silicate glass were undertaken to clarify why mere ion adsorption is relatively ineffective at modifying dissolution rates in hyperalkaline conditions and why certain organic salts of Na enhance the rates. These experiments are designed to assess the effect of pH on dissolution rates either in the absence of any additional salts (Figure 6a) or with NaBr at twice the concentration used before, 20 mmol L-1 (Figure 6b). As reported for many silicate minerals and glasses,67–69 the dissolution rate increases substantially for pH > 10 (see Figure 6a). In contrast, the rate is low and nearly constant for 5.6 ≤ pH ≤ 10, a behavior similar to that observed for other silicate glasses.48,70 This behavior supports the idea that a critical hydroxyl activity must be achieved before hydrolysis can happen at appreciable rates. Glass dissolution rates in the presence of 20 mmol L-1 NaBr are about 1.3 times greater than in a salt-free solution for pH between 6 and 10. However, the rates are about the same as in the saltfree solution for pH > 10 (see Figure 6b). The speciation of NaBr is not sensitive to pH, which implies that the effects of NaBr on dissolution rate are not related to bulk solution composition effects. However, the ability of a particular electrolyte like NaBr to enhance dissolution may be related to the abundance of surface sites with which the salt ions may interact, as illustrated in Figure 7, and the surface site speciation changes systematically with increasing pH. Specifically, the relative fractions of surface species on amorphous or crystalline silica shift systematically from protonated (>Si–OH2+) to neutral (>Si–OH) and then to deprotonated (>Si–O-) as pH increases,71 as also suggested by zeta potential measurements in Figure 6(c). For comparison, pure amorphous silica has an isoelectric point (IEP, defined by zero zeta potential) at pH = 1.9.72 The magnitude of surface charge increases at pH above or below the IEP. The partitioning of these surface sites has been shown to correlate with pH-dependence of the dissolution rate of quartz71 and, more recently, almandine.31 At near-neutral and slightly alkaline pH’s, silicate surface sites are primarily neutral >Si–OH with a small fraction of >Si–O- sites.71 Under these conditions, addition of sodium electrolytes causes hydrated Na+ ions to adsorb. The consequent


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increases in the concentrations of water and hydroxide ions near the surface promote dissolution by increasing rates of surface hydrolysis.39,73 The complexation of Na+ with available >Si–Osites further increases rates. Therefore, in the near-neutral pH region organic and inorganic ions are most effective at modifying silicate dissolution rates.38,40,74

sest = 1.93E-10

(a) (b) (c) Figure 6. (a) The dissolution rate of the soda lime silicate glass as a function of pH, and (b) the dissolution rate of the soda lime silicate glass for solutions containing 20 mmol L-1 NaBr relative to equivalent solutions of the same pH but without NaBr. Little if any change in rate is induced by NaBr for pH > 10. (c) The zeta potential of soda lime silicate glass as a function of pH, showing significant increase in magnitude at pH > 10, consistent with the prevalence of negative Si–O- sites. At high pH the surface, with mostly >Si–O- sites, has a large negative charge density that controls bulk dissolution rates.73 In addition, in high pH-solutions with high concentrations of Na+ , >Si–O-–Na+ species are also abundant on the surface.26,71 Other studies have proposed that outer-sphere sorption has an effect similar to that of ionic strength on dissolution rates, both of which increase the density of >Si–O- sites.69,75 However, in NaOH above pH 10 the surface is already composed primarily of >Si–O- sites that have already formed stable complexes with Na+, so added salts alter neither the nature of surface sites nor the rate at which they dissolve, as shown illustratively in Figure 7. Sodium benzoate has the greatest effect on dissolution of any of the electrolytes studied here (Figure 4). Its influence, like that of disodium malonate to a lesser extent, could be due to either ion-specific interactions with the glass network or complexation with ions in solution. If the organic anion complexes with any of the free ionic products of glass dissolution, the concentration of those free ions is decreased. This would, in turn, both increase the driving force for continued dissolution and increase the apparent solubility of the glass in terms of total elemental concentrations in solution.38,76 However, the benzoate ligand, like other carboxylic acids, likely forms only very weak complexes with Si in solution,77 so Na-benzoate probably does not influence dissolution of silicate glass in this way. Therefore, it likely enhances dissolution rate by some ion-specific interaction with the silicate surface and the organic ligand, despite the fact that previous studies in acidic or neutral conditions have shown that organic ligands do not interact appreciably with silicate surfaces. For example, Poulson et al. showed that sodium oxalate has an enhancing effect on quartz dissolution at pH < 7 similar to that of NaCl, although the oxalate anion does not adsorb at the surface.78 Similarly, Kubicki et al. observed no significant adsorption of carboxylic acid at silicate surfaces and argued that


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enhancement by organic salts is not caused by surface Si complexation with the organic ligand at low pH.79 Instead, they proposed that a more likely role of the organic ligand could be complexation with the charge compensating alkali and alkaline earth cations in the solid.79 Together, these studies implied that complexation involving the Na+ cation, or possibly complexation of the organic ion with charge compensating ions in the mineral, was likely responsible for the rate enhancement under those experimental conditions. However, in the high-pH NaOH solutions used here, the surface is already populated with abundant >Si–O-–Na+ sites. Also, as shown in Figure 4, rate increases are not observed in the presence of NaBr or other inorganic sodium salts at high pH. Therefore, the current results are consistent with a mechanism involving organic anion complexation with either surface sites or with the charge compensating cations in the glass. Either way, the effect is specifically induced by the organic anion at high pH. 4. Summary and conclusions This study has demonstrated that dissolved organic and inorganic salts often have a minor influence on soda lime silicate glass dissolution rates under hyperalkaline conditions, up to a pH of 13. The solution’s pH has a much larger effect in altering the structure, availability and distribution of silicate surface sites than any other solution parameter. Particularly, under basic conditions, the surface is primarily composed of >Si–O- groups. The rate of dissolution is controlled mostly by the rate of attack by OH- ions, which are abundant in hyperalkaline solutions, on the polar siloxane bonds. In contrast, under near-neutral and mildly alkaline (pH < 10) conditions, added salts more effectively enhance silicate dissolution rates by adsorption (both inner-sphere and outer-sphere) of ions and the attending changes in local concentration of OHions and water at the silicate surface. The low rate enhancements at pH > 10 show that dissolution is not accelerated by outer-sphere adsorption of ions because the surface already has a highly negative charge and strong complexation with the alkaline solution’s cation. Dissolution rates are directly correlated with the >Si–O- site density, which is not significantly affected by outer-sphere adsorption, so further increases in dissolution rates are not observed by adding an electrolyte. Therefore, electrolytes that merely adsorb on the surface are relatively ineffective at modifying glass dissolution rates. In contrast, substantial rate enhancement can be induced by sodium benzoate additions due either to the formation of strong inner-sphere surface complexes with Si or charge compensating cations within the glass. These findings should be generally applicable to silicate surfaces if their Si surface site densities are known. This has implications on developing new routes for accelerating silicate dissolution rates in applications such as cementation and chemical-mechanical polishing (CMP).


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Near-neutral solutions

Highly alkaline solutions

Soda lime silicate glass

Soda lime silicate glass

Figure 7. An illustration of silicate dissolution at (left) near-neutral and (right) highly alkaline pH levels as promoted by metal adsorption and hydroxyl attack. The illustration also clarifies why in a hyperalkaline environment that is overabundant in OH- and Na+ the addition of other ions (unless they form surface precipitates or strong complexes with Si or charge compensating cations in the solid) is typically ineffective in altering dissolution rates. Shown in the schematic are H (white), O (red), Si (gray), Na or other alkali metals (purple), Ca or other alkaline earth metals (green), and Br (dark red) atoms. 5. Acknowledgements The authors acknowledge financial support for this research provisioned by the COMAX Consortium: A joint UCLA-NIST Initiative that is supported by its industry and government agency partners, the Department of Energy’s Nuclear Energy University Program (DOE-NEUP: DE-NE0008398), National Science Foundation (CAREER Award: 1253269), and the U.S. Department of Transportation (U.S. DOT) through the Federal Highway Administration (DTFH61-13-H-00011). The contents of this paper reflect the views and opinions of the authors who are responsible for the accuracy of data presented. This research was carried out in the Laboratory for the Chemistry of Construction Materials (LC2) and Molecular Instrumentation Center at UCLA. As such, the authors gratefully acknowledge the support that has made these laboratories and their operations possible. References (1) White, A. F.; Brantley, S. L. The Effect of Time on the Weathering of Silicate Minerals: Why Do Weathering Rates Differ in the Laboratory and Field? Chem. Geol. 2003, 202, 479–506. (2) Glasser, F. P.; Marchand, J.; Samson, E. Durability of Concrete — Degradation Phenomena Involving Detrimental Chemical Reactions. Cem. Concr. Res. 2008, 38, 226– 246. (3) Busenberg, E.; Clemency, C. V. The Dissolution Kinetics of Feldspars at 25°C and 1 atm CO2 Partial Pressure. Geochim. Cosmochim. Acta 1976, 40, 41–49.


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Enhancing Silicate Dissolution Kinetics in Hyperalkaline Environments

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