Article pubs.acs.org/IECR
Catalytic Effect of Magnesium Ions on Silicic Acid Polycondensation and Inhibition Strategies Based on Chelation Konstantinos D. Demadis,* Antonia Ketsetzi, and Eva-Maria Sarigiannidou Crystal Engineering, Growth and Design Laboratory, Department of Chemistry, University of Crete, Voutes Campus, Heraklion, Crete GR-71003, Greece ABSTRACT: The catalytic role of Mg2+ ions in the polycondensation of silicic acid to form amorphous silica has been investigated in detail. This behavior is pH-dependent. As pH increases (herein, the three pH values 8.0, 9.0, and 9.5 were tested), the catalytic effect of Mg2+ ions becomes more pronounced. Also, this behavior is directly proportional to the concentration of Mg2+. Ethylenediaminetetraacetic acid (EDTA) can inhibit this catalytic effect by strongly chelating the Mg2+ ions. This research can be further expanded to other chelating agents that have an affinity for Mg2+ ions.
■
INTRODUCTION Scaling, corrosion, and microbiological fouling present a threefold problem in industrial water systems.1 Fouled critical equipment components need to be cleaned before reinstallation for full system recovery. Mechanical or chemical cleaning are viable options, but require several person hours, total system shutdowns, high costs, and potential profit losses.2 Foulants could be organic or inorganic. Organic foulants are a result of insufficient system biocontrol or deposition of organic matter brought into the system from external sources (when river or lake water are used).3 Inorganic foulants include crystalline, sparingly soluble mineral salts, such as calcium-based carbonates (e.g., calcite,4 aragonite,5 vaterite,6 amorphous calcium carbonate,7 and the rare hexahydrate8), calcium-based sulfates (gypsum,9 hemihydrate,10 and anhydrite11), barium12 and strontium13 sulfates, as well as amorphous and colloidal deposits, such as silica14 and magnesium silicate,15 depending on chemical characteristics of local water sources16 or on the particular industry/application.17 Silica and magnesium silicate are poorly studied foulants, although recent efforts have been put forth to reveal their chemical identities.18 There are established chemical control methods for mineral salts. These commonly include use of a scale inhibitor (phosphonic acids are widely used for this purpose19), combined with use of dispersant polymers.20 Several publications have appeared on the possible mechanisms of scale inhibition, pointing toward the surface interaction of acidic groups on the inhibitor backbone with surface-exposed metal ions located at kinks and steps of the growing crystal.21 Nevertheless, satisfactory protocols for the inhibition of silica and magnesium silicate scales are much less developed. Possible reasons for the lack of universal and widely accepted methodology include (1) the amorphous nature of silica, which precludes use of mineral scale inhibitors, as there are no crystal surfaces for interaction with the inhibitor; (2) incomplete information on the true natures of silica and magnesium silicate, as these inorganic species are precisely defined in a geological context, but the precipitates/deposits that form in water systems bear little resemblance to their geological counterparts; and (3) scarcity of silica/silicate © 2012 American Chemical Society
deposits, as silicon species are not widespread in all ground or surface waters, so that their deposits create problems in certain parts of the world where such species are found in increased concentrations. Silica/silicate scales form and pose a problem in waters that satisfy at least one of the following three conditions: (1) contain high levels of silica, (2) contain high levels of magnesium, and/or (3) operate at high pH (>8.5). The objective of this work was to study the effect of Mg2+ ions on the polycondensation of silicic acid and to investigate suitable “remedies” based on chemical additives. More specifically, we studied the effects of various Mg2+ concentrations (20−100 ppm or 0.082−0.410 mM) on their ability to accelerate amorphous silica formation at the three pH values 8.00, 9.00, and 9.50. We further studied the effect of a common chelating agent (ethylenediaminetetraacetic acid, tetrasodium salt, EDTA) on the above-mentioned catalytic Mg-induced formation of silica. Based on fundamental coordination chemistry arguments, EDTA is expected to cease, or at least suppress, the effects of Mg2+.
■
FORMATION AND GROWTH OF AMORPHOUS SILICA SCALES The formation, precipitation, and deposition of amorphous silica in process industrial waters have been subjects of intense interest. In parallel, there is also a substantial focus on biosilica formation, because silica is used by nature as a structural material for several organisms, such as diatoms22 and sponges.23 Silica scale formation is a highly complex process.24 It is usually favored at pH values of less than 8.5, whereas magnesium silicate scale forms at pH values of greater than 8.5. Available data suggest that silica solubility is largely independent of pH in the range of 6−8. This is the pH range of minimum silica solubility, where silicic acid polymerization has a maximum rate.25 Silica exhibits normal solubility characteristics. Its Received: Revised: Accepted: Published: 9032
April 26, 2012 June 4, 2012 June 11, 2012 June 11, 2012 dx.doi.org/10.1021/ie3010836 | Ind. Eng. Chem. Res. 2012, 51, 9032−9040
Industrial & Engineering Chemistry Research
Article
water with the geothermal water. Magnesium silicates have low solubility in warm waters at high pH. Heating of groundwater depletes the magnesium concentration of geothermal waters to levels mostly below 0.1 mg/kg. Magnesium silicate is amorphous according to X-ray diffraction (XRD) experiments.27 However, whatever “structure” has been found resembles that of chrysotile. It was also found that the Mg/Si ratio is close to 1 with small variations. The magnesium silicate sepiolite precipitates from seawater at low temperature (down to 25 °C), as the dissolved silica concentration is increased. Increased temperature and high pH enhance the rate of precipitation. The magnesium silicate talc forms easily in hydrothermal experiments and is frequently formed outside its stability field. Several other magnesium silicates such as stevensite, saponite, and chrysotile are known to be formed hydrothermally at relatively low temperatures. Heating of fresh water also initiates precipitation, and it is wellknown that magnesium is one of the major components in “boiler stone”. The major factors controlling the degree of supersaturation are boiling temperature and pH, the latter of which is, in turn, mainly dependent on the deaeration process. Supersaturation is in all cases greater for talc than for chrysotile.28 Coprecipitation of magnesium hydroxide, Mg(OH)2, and colloidal silica has also been observed.29 One theory proposes that formation of Mg(OH)2 occurs first and then Mg(OH)2 subsequently reacts with monomeric silicate and/or polymeric silica to form magnesium silicate.30 Ca2+ and Mg2+ salts were found to catalyze the silica polymerization reaction in reverse osmosis systems.31 Higher concentrations of total hardness lead to a faster drop in dissolved silica in solution. In batch runs, Mg2+ was found to affect silica concentrations more than Ca2+. For example, runs with a given hardness level but with lower ratios of Ca to Mg caused faster declines in dissolved silica. Magnesium silicate seems to be a “true” compound according to Young et al.32 According to their results, fairly consistent amorphous precipitate was obtained. The stoichiometric ratio of silicon to magnesium was found to be 1:1. This was the same whether the mother liquor contained a 1:2 or 2:1 molar ratio of silica to magnesium and whether the precipitation took place at room temperature or 75 °C. Some comments on the possible mechanism of formation are warranted. If magnesium hydroxide precipitated out and silica simply absorbed, there should be little effect of silica on the precipitation point. By the same reasoning, the “opposite” mechanism of silica precipitation followed by magnesium absorption should be independent of magnesium concentration. In fact, increasing or decreasing the silica concentration has an effect essentially equal to that of similar increases or decreases in magnesium concentration. The precipitate was found to contain significant amounts of adventitious water, presumably in the pores of the gel. This magnesium silicate precipitate dissolved in acid. Alternatively, ethylenediaminetetraacetic acid chelated the magnesium from the precipitate, leaving a loose floc of virtually pure colloidal silica that did not redissolve in acid. It can be assumed that the magnesium silicate initially forms a loose, open gel structure with numerous hydroxide bridges. An alternative mechanism of magnesium silicate formation was proposed. According to this proposal, formation of magnesium silicate seems to be a two-step process. Under relatively high pH conditions, magnesium hydroxide is precipitated. Because magnesium hydroxide is inversely soluble with respect to temperature, the precipitation can take place near the surface of
solubility increases proportionally to temperature. In contrast, magnesium silicate exhibits inverse solubility. Other forms of silica, such as quartz (crystalline SiO2) and glass, also exhibit “normal” solubility, but they are both less soluble than amorphous silica.26 Amorphous silica formation is a polycondensation event. When silicic acid/silicate ions condense and polymerize, they form a plethora of structural motifs, including rings of various sizes, cross-linked polymeric chains of different molecular weights, and oligomeric structures. The resulting amorphous silica scale is a complicated mixture of such components. Silicic acid polymerization starts with an attack of a deprotonated, negatively charged silicate ion on a silicic acid molecule, yielding an initial silicic acid “dimer”, which then continues to undergo further attack. This results in random polymer chain growth that produces silica nanoparticles. These, in turn, can grow further (by incorporation of silicic acid onto the silica particle surface) or agglomerate with other nanoparticles to give larger particles.
■
WATER-FORMED MAGNESIUM SILICATE DEPOSITS: A SHORT LITERATURE OVERVIEW The term “magnesium silicate” is widely recognized in the water treatment industry. However, its definition differs from that in geology. In general, a deposit that contains magnesium, silicon, and oxygen is called magnesium silicate. In harsher environments, such as geothermal applications, the effect of high temperature favors the formation of geologically recognized magnesium silicates. Precipitation of magnesium silicate can cause problems in a number of water treatment applications from truck radiators to geothermal wells and plants. The magnesium silicate system is highly pH-dependent. Below pH 7, there is essentially no chance of precipitation, because the silica exists in an unreactive, un-ionized form. Above pH 9, magnesium silicate is very likely to form because silica forms reactive silicate ions. Furthermore, temperature is extremely important. Precipitation begins at lower pH if the temperature is sufficiently high. This is clearly shown in Figure 1.
Figure 1. Effect of pH on the deposition of magnesium silicate.
Scaling of magnesium silicates has been a problem in some Icelandic district heating systems.27 This kind of scaling is not encountered in heating systems utilizing geothermal water directly but occurs by heating and deaerating fresh water. Two of the plants have heat exchangers to heat fresh water. The water in those systems is also discarded after a single use and not recirculated in the heating system. Scaling of a similar type has occurred in a few other systems as a result of mixing cold 9033
dx.doi.org/10.1021/ie3010836 | Ind. Eng. Chem. Res. 2012, 51, 9032−9040
Industrial & Engineering Chemistry Research
Article
container. The pH of this solution was initially ∼11.8 and was adjusted to the desired pH value by addition of HCl or NaOH, if needed [nine drops of concentrated HCl (10%) and then another nine drops of dilute HCl (5%)]. In this work, three pH values were tested: 7.00, 8.00, and 9.50. The volume change was taken into account in calculations. Then, the appropriate amount of solution B was added. Various amounts were tested, such that the final Mg2+ concentration was 20, 40, 60, 80, and 100 ppm. Then, the container was covered with a plastic membrane and set aside without stirring. The solutions were checked for soluble silicic acid by the silicomolybdate method (see below) every hour for the first 8 h (short-term experiments) or after 24-, 48-, and 72-h (long-term experiments) time intervals after the final pH adjustment (t = 0). Protocol for the Effect of EDTA on Mg2+-Catalyzed Silica Formation. In general, the same procedure was followed, except that, after the pH adjustment to lower values. the required amount of solution C (such that the concentration of EDTA was 20, 40, 60, 80, and 100 ppm) was added before addition of solution B. After that, the same procedure as in the preceding section was followed. Quantification of Soluble (Reactive) Silica. The term “soluble or reactive silica” actually denotes soluble silicic acid and was measured using the silicomolybdate spectrophotometric method. According to this method, 2 mL of working solution filtered through a 0.45-μm syringe filter was diluted to 25 mL in the cell with a light path of 1 cm. Then, 1 mL of solution D and 0.5 mL of solution E were added to the sample cell, and the solution was mixed well and left undisturbed for 10 min. Next, 1 mL of solution F was added, and the overall solution was mixed again. The solution was set aside for 2 min. After the second time period, the photometer was set to zero absorbance with DI water. Finally, the sample absorbance was measured at 452 nm as ppm soluble silica. The detectable concentration range was 0−75.0 ppm. To calculate the concentration in the original solution, a dilution factor was applied. The silicomolybdate method is based on the principle that ammonium molybdate reacts with reactive silica and any phosphate present at low pH (about 1.2) and yields heteropoly acids, yellow in color. Oxalic acid was added to destroy the molybdophosphoric acid, leaving silicomolybdate intact and thus eliminating any color interference from phosphates. It must be mentioned that this method measures soluble silica and that this term includes molybdate-reactive species. These include silicic (monomer) and disilicic (dimer) acid. It should be noted that Mg2+ and Na4EDTA additives do not interfere with the silicomolybdate spectrophotometric method.
the heat-transfer tubes and the maximum exchanger tube wall temperature should be ∼80 °C. Temperature has a greater influence on the deposition than any of the variables. It was reported that a hydroxylated magnesium silicate forms in seawater in which the SiO2 concentration exceeds 26 ppm at pH 8.1 and clay minerals are found (kaolinite, glauconite, and montmorillonite).33
■
EXPERIMENTAL SECTION Instrumentation. Attenuated total reflectance infrared (ATR-IR) spectra were collected on a Thermo-Electron Nicolet 6700 Fourier transform infrared (FTIR) optical spectrometer. FTIR spectra were recorded on a Perkin-Elmer FT 1760 FTIR spectrometer in KBr discs. Measurements of soluble silicic acid were carried out using a HACH 890 spectrophotometer from the Hach Co., Loveland, CO. Scanning electron microscopy (SEM) images were collected on a LEO VP-35 FEM scanning electron microscope. Reagents and Materials. All chemicals were from commercial sources. Sodium silicate (Na2SiO3·5H2O), ammonium molybdate [(NH4)6Mo7O24·4H2O], and oxalic acid (H2C2O4·2H2O) were from EM Science (Merck). Sodium hydroxide (NaOH) was from Merck, hydrochloric acid (37%) was from Riedel de Haen. Ethylenediaminetetraacetic acid (EDTA, tetrasodium salt) was from Aldrich. All reagents were used as received from suppliers. Acrodisc filters (0.45 μm) were from Pall-Gelman Corporation. In-house-deionized water was used for all experiments. This water was tested for soluble silica and was found to contain negligible amounts. Solution Preparation. The sodium silicate stock solution (solution A) used in all experiments was prepared by dissolving 4.40 g of Na2SiO3·5H2O in 2.5 L of deionized (DI) water. The mixture was stirred for at least 1 day to ensure complete dissolution. This stock solution contained 500 ppm Si (8.33 mM), expressed as ppm SiO2. A 1% w/v (10000 ppm in Mg2+) solution (solution B) was prepared by dissolving 1.014 g of MgSO4·7H2O in 10 mL of DI water. A 1% w/v (10,000 ppm in Na4EDTA) solution (solution C) was prepared by dissolving 0.100 g of Na4EDTA in 10 mL of DI water. The ammonium molybdate solution (solution D) used in the silicomolybdate test was prepared by dissolving 10 g of (NH4)6Mo7O24·4H2O in 100 mL of DI water, to which seven or eight pellets of solid NaOH had been added under stirring. A 1 + 1 HCl stock solution (solution E) was prepared by mixing equal quantities of concentrated HCl (37% w/v) and DI water. The oxalic acid solution (solution F) was prepared by dissolving 8.750 g of solid hydrated oxalic acid [(COOH)2·2H2O] in 100 mL of DI water. All of these stock solutions were kept in a refrigerator. Silicification Protocols. The basic protocols for all experiments and measurements described herein have been reported in detail elsewhere.34 All experiments were carried out at room temperature. Molybdate-reactive silicic acid was measured using the silicomolybdate spectrophotometric method, which has a ±5% accuracy.35 Reproducibility was excellent. Solution pH was checked after addition of the additives (Mg, EDTA, or both), but no major changes were noted. Also, after every sampling, pH was again checked, but no major fluctuations were measured. In general, pH adjustments were made only when solution pH was off by more than 0.2 pH units. Because the present silicification protocols involve the addition of Mg2+, we briefly outline the procedures followed. Protocol for the Effect of Mg2+ on Silica Formation. Solution A (100 mL) was placed in a polyethylene (PET)
■
RESULTS Catalytic Role of Mg2+ in Colloidal Silica Formation. In these studies, we varied the concentration of Mg2+ (20, 60, and 100 ppm; these are Mg2+ concentrations commonly found in industrial waters), and we followed its effect on silicic acid condensation at three pH values (8.0, 9.0, and 9.5), where magnesium silicate has been reported to form.36 Figure 2 clearly shows that, at pH 8.0, Mg2+ up to 100 ppm had virtually no effect on the silicic acid condensation reaction. The values measured spectrophotometrically for silicic acid are indistinguishable from those corresponding to the control (no Mg2+ present). When the pH was increased to 9.0 (Figure 3), the catalytic effects of Mg2+ started to appear. This is evidenced by the consistently lower measured values for silicic acid than for the 9034
dx.doi.org/10.1021/ie3010836 | Ind. Eng. Chem. Res. 2012, 51, 9032−9040
Industrial & Engineering Chemistry Research
Article
dramatic effect. Another significant conclusion derived from Figure 4 is that, at pH 9.5, the effect of Mg2+ was measurable and significant. There seemed to be a rapid decrease in soluble silicic acid levels as the Mg2+ concentration increased. For example, at 20 ppm, Mg2+ caused an additional 40 ppm silicic acid to be polymerized. At 20 ppm Mg2+, this value increased to 87 ppm, and finally, at the level of 100 ppm Mg2+, soluble silicic acid levels dropped ∼140 ppm lower than the control. This is convincing evidence that Mg2+ is an effective catalyst for silicic acid polymerization at pH values greater than 9.0. Effect of EDTA on Mg-Catalyzed Silicic Acid Condensation. Magnesium silicate scale control was pursued in our laboratories by use of EDTA as a Mg2+ sequestering agent. First, it was ensured that EDTA alone had no effect (catalytic or retarding) on the condensation of silicic acid. Hence, by following the procedure outlined in the Experimental Section, EDTA was tested at concentrations up to 100 ppm, and no effect was noted at pH values 8.0, 9.0, and 9.5 (results not shown). Figure 5 shows that addition of EDTA at a
Figure 2. Effect of Mg2+ [20 (0.082), 60 (0.246), and 100 (0.410) ppm (mM)] on silicic acid condensation at pH 8.0. No catalytic effect is observed.
Figure 3. Effect of Mg2+ [20 (0.082), 60 (0.246), and 100 (0.410) ppm (mM)] on silicic acid condensation at pH 9.0. Note that, in the presence of Mg2+, soluble (unpolymerized) silicic acid was consistently lower than that measured for the control (no Mg2+ present).
Figure 5. Effect of EDTA [20 (0.0048), 40 (0.0144), 60 (0.0240) ppm (mM)] on silicic acid condensation at pH 8.0 in the presence of Mg2+ [20 (0.082), 60 (0.246), and 100 (0.410) ppm (mM)]. The Mg/ EDTA molar ratio was ∼17 in all experiments.
control. The reduction in silicic acid levels was ∼50 ppm. However, an increase of the Mg2+ concentration seemed to have no measurable effect. Increasing the operating pH to 9.5 had a dramatic change on the catalytic effects of Mg2+. Figure 4 demonstrates this
concentration level equal to that of Mg had no effect on soluble silicic acid. These experiments were performed by monitoring soluble silicic acid levels (with a starting concentration of silicic acid of 500 ppm as SiO2). EDTA was demonstrated to be ineffective at the dosages tested. Soluble silicic acid levels were the same as those in the absence of EDTA. When the operating pH was increased to 9.0, the same situation was observed. As illustrated in Figure 6, no increase in soluble silicic acid levels was observed, and these silicic acid values were the same as those without EDTA present. When the pH was increased to 9.5, a profound, dosagedependent effect of EDTA was observed (Figure 7). All three EDTA dosages (20, 40, and 60 ppm) caused soluble silicic acid to rise above the control level. An interesting observation warrants further discussion. The dosage dependence seemed to have an inverse relationship. The higher the Mg/EDTA combination dosage, the lower the soluble silica observed. Therefore, the most effective Mg/EDTA combination for maximum soluble silica was 20/20 ppm. A possible explanation for this inverse effect might be that, at increased Mg/EDTA levels (40/40 and 60/60 ppm), possible precipitation of a Mg− EDTA complex might occur. EDTA is well-known to be an
Figure 4. Effect of Mg2+ [20 (0.082), 60 (0.246), and 100 (0.410) ppm (mM)] on silicic acid condensation at pH 9.5. Note the dramatic concentration-dependent effect of Mg2+ on the reduction of soluble (unpolymerized) silicic acid. 9035
dx.doi.org/10.1021/ie3010836 | Ind. Eng. Chem. Res. 2012, 51, 9032−9040
Industrial & Engineering Chemistry Research
Article
Figure 6. Effect of EDTA [20 (0.0048), 40 (0.0144), 60 (0.0240) ppm (mM)] on silicic acid condensation at pH 9.0 in the presence of Mg2+ [20 (0.082), 60 (0.246), and 100 (0.410) ppm (mM)]. The Mg/ EDTA molar ratio was ∼17 in all experiments.
Figure 7. Effect of EDTA [20 (0.0048), 40 (0.0144), 60 (0.0240) ppm (mM)] on silicic acid condensation at pH 9.5 in the presence of Mg2+ [20 (0.082), 60 (0.246), and 100 (0.410) ppm (mM)]. The Mg/ EDTA molar ratio was ∼17 in all experiments. Figure 8. Silica precipitates in the presence of Mg2+ at pH 9.5. Upper images: 100 ppm Mg2+. Middle images: 100 ppm Mg2+ + 60 ppm EDTA. Lower image: XRD powder pattern of the amorphous precipitate obtained at pH 9.5 in the presence of 100 ppm Mg2+, in either the presence or absence of EDTA.
effective chelator of Mg at high pH values. A Mg−EDTA complex has been structurally characterized.37 Deposit Characterization. Silica precipitates in the presence of Mg2+ appeared as white, loose precipitates, as shown in a representative example in Figure 8 (left). In the presence of EDTA, the amount of precipitate decreased, as corroborated by the results in Figure 7. The precipitates obtained at pH 9.5 when Mg2+ was present were amorphous as determined by XRD, in either the presence or absence of EDTA (see Figure 8, lower image). However, their size and shape changed, depending on whether EDTA was present. In the absence of EDTA. the particles appeared as polydisperse blocks with sizes ranging from 10 to 30 μm. In the presence of EDTA, the particles appeared to have a round shape with a diameter ∼40 μm (see Figure 8, middle image). Upon closer examination, however, the particles appeared as aggregates of much smaller particles, Li+ > Na+ > K+.43 The principal findings of this research are summarized as follows: (1) Mg2+ affects the condensation of silicic acid in a pH- and concentration-dependent manner. Its effect is minimal at pH 8.0 but profound at pH 9.5. (2) Magnesium is not incorporated in the final silica precipitate; therefore, its role is truly catalytic. (3) The tetracarboxylic acid EDTA can inhibit the catalytic effect of Mg2+ on silica polymerization. It acts as a strong chelant for Mg2+ and, hence, renders the Mg2+ ions “inactive”. (4) In addition to stopping the catalytic effect of Mg2+, EDTA seems to offer additional stabilization of silicic acid. This rather surprising result certainly needs further exploration. Taking into account that all EDTA under these experimental conditions is complexed by the Mg2+, it is probable that the formed Mg−EDTA complex could exert some silicic acid stabilization. Magnesium ions can cause precipitation problems in highsilica waters at high pH values. Based on the work described herein, possible solutions to this issue could include lowering
Figure 10. Inhibition of the catalytic effect of Mg2+ on silica polymerization with the addition of EDTA at pH 9.5.
In fact, the addition of EDTA seems not only to cease the catalytic effect, but also to cause further inhibition of silicic acid polymerization. For example, in the presence of 20 ppm Mg2+ and 20 ppm EDTA, there is an additional stabilization of >100 ppm of silicic acid. However, this level drops as the level of Mg2+ increases. EDTA is a well-known chelator of metal ions.39 Its affinity for alkaline-earth metal cations is exceptional. Therefore, under the experimental conditions of our experiments, EDTA chelates and “immobilizes” Mg2+. Hence, Mg2+ can no longer interact with silicic acid, so its catalytic activity for the polymerization reaction is diminished. Although a precise mechanism of the Mg2+ catalytic effect and its EDTA-induced suppression is not available at the moment, certain arguments can be put forth to propose possible pathways for these processes. These are shown in Figure 11 and outlined as follows: As stated above, only at pH 9.0 and above does the catalytic effect of Mg2+ start to become 9037
dx.doi.org/10.1021/ie3010836 | Ind. Eng. Chem. Res. 2012, 51, 9032−9040
Industrial & Engineering Chemistry Research
Article
Figure 11. Possible pathways that are consistent with the catalytic effects of Mg2+ ions on the polycondensation of silicic acid.
■
ACKNOWLEDGMENTS We thank the Department of Chemistry, University of Crete, and the GSRT (Contract GSRT-2006-170c) for partial financial support.
the operating pH below 8. However, in this case use of mineral acids is necessary and often unacceptable by certain end users. Addition of chelating agents could mitigate the problem, as shown here with EDTA. It is certain that this “chelating therapy” can be extended to several other metal chelators. Such experiments are currently underway in our laboratory, with an emphasis on “green” and environmentally acceptable chelating agents.
■
■
REFERENCES
(1) (a) Demadis, K. D. Water Treatment Processes; Nova Science Publishers, Inc.: New York, 2012. (b) Demadis, K. D. Water Treatment’s “Gordian Knot”. Chem. Process. 2003, 66 (5), 29−32. (c) Frenier, W. W.; Ziauddin, M. Formation, Removal, and Inhibition of Inorganic Scale in the Oilfield Environment; Society of Petroleum Engineers: Richardson, TX, 2008. (d) Bott, T. R. Fouling Notebook; Institution of Chemical Engineers: Rugby, U.K., 1990. (e) Bott, T. R. Fouling of Heat Exchangers; Elsevier Science: New York, 1995. (2) (a) Mavredaki, E.; Neofotistou, E.; Demadis, K. D. Inhibition and Dissolution as Dual Mitigation Approaches for Colloidal Silica (SiO2) Fouling and Deposition in Process Water Systems: Functional Synergies. Ind. Eng. Chem. Res. 2005, 44, 7019−7026. (b) Demadis,
AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected]. Notes
The authors declare no competing financial interest. 9038
dx.doi.org/10.1021/ie3010836 | Ind. Eng. Chem. Res. 2012, 51, 9032−9040
Industrial & Engineering Chemistry Research
Article
scaling at high-barium formation water. J. Pet. Sci. Eng. 2012, 70, 74− 83. (c) Shen, D.; Fu, G. M.; Al-Saiari, H.; Kan, A. T.; Tomson, M. B. Barite Dissolution/Precipitation Kinetics in Porous Media and in the Presence and Absence of a Common Scale Inhibitor. SPE J. 2009, 14, 462−471. (13) (a) BinMerdhah, A. B. Inhibition of Calcium Sulfate and Strontium Sulfate Scale in Waterflood. SPE Prod. Oper. 2012, 25, 545− 552. (b) Savenko, A. V. Solubilities of strontium carbonate and sulfate in seawater. Geochem. Int. 2004, 42, 78−187. (14) (a) Fleming, B. A. Kinetics of Reaction between Silicic Acid and Amorphous Silica Surfaces in NaCI Solutions. J. Colloid Interface Sci. 1986, 110, 40−64. (b) Meyers, P. Behavior of Silica Technologies Available and How They Rate. Water Cond. Purif. 2004, March, 22− 24. (c) Darton, E. G. RO plant experiences with high silica waters in the Canary Islands. Desalination 1999, 124, 33−41. (15) Demadis, K. D. Recent Developments in Controlling Silica and Magnesium Silicate in Industrial Water Systems. In Science and Technology of Industrial Water Treatment; CRC Press: London, 2010; Chapter 10, pp 179−203. (16) Abu Ghunmi, L.; Zeeman, G.; Fayyad, M.; van Lier, J. B. Grey Water Treatment Systems: A Review. Crit. Rev. Environ. Sci. Technol. 2011, 41, 657−698. (17) (a) Demadis, K. D. Scale Formation and Removal. Power 2004, 148 (6), 19−23. (b) Demadis, K. D. Investing in Chemical Cooling Water Treatment. Water Wastewater Int. 2005/2006, December/ January, 21−22. (18) (a) Demadis, K. D.; Neofotistou, E. Inhibition and Growth Control of Colloidal Silica: Designed Chemical Approaches. Mater. Performance 2004, 43 (4), 38−42. (b) Stathoulopoulou, A.; Demadis, K. D. Enhancement of Silicate Solubility by Use of “Green” Additives: Linking Green Chemistry and Chemical Water Treatment. Desalination 2008, 224, 223−230. (19) (a) Demadis, K. D.; Katarachia, S. D. Metal-Phosphonate Chemistry: Preparation, Crystal Structure of Calcium−Amino−tris− Methylene Phosphonate and CaCO3 Inhibition. Phosphorus Sulfur Silicon 2004, 179, 627−648. (b) Demadis, K. D.; Baran, P. Chemistry of Organophosphonate Scale Growth Inhibitors: Two Dimensional, Layered Polymeric Networks in the Structure of Tetrasodium 2Hydroxyethyl-amino-bis(methylenephosphonate). J. Solid State Chem. 2004, 177, 4768−4776. (c) Demadis, K. D.; Raptis, R. G.; Baran, P. Chemistry of Organophosphonate Scale Growth Inhibitors: 2. Structural Aspects of 2-Phosphonobutane-1,2,4-Tricarboxylic Acid Monohydrate (PBTC·H2O). Bioinorg. Chem. Appl. 2005, 3, 119−134. (20) (a) Kuila, D.; Blay, G. A.; Borjas, R. E.; Hughes, S.; Maddox, P.; Rice, K.; Stansbury, W.; Laurel, N. Polyacrylic Acid (Poly-A) as a Chelant and Dispersant. J. Appl. Polym. Sci. 1999, 73, 1097−1115. (b) Amjad, Z. Factors influencing the performance of natural and synthetic additives as iron dispersants. Tenside, Surfactants, Deterg. 2007, 44, 88−93. (c) Amjad, Z.; Guyton, D. Biopolymers and Synthetic Polymers as Iron Oxide Dispersants for Industrial Water Applications. Mater. Perform. 2012, 51 (3), 48−53. (d) Zhou, X.; Sun, Y.; Wang, Y. Inhibition and dispersion of polyepoxysuccinate as a scale inhibitor. J. Environ. Sci. 2011, 23 (Suppl.), S159−S161. (21) (a) Martinod, A.; Euvrard, M.; Foissy, A.; Neville, A. Progressing the understanding of chemical inhibition of mineral scale by green inhibitors. Desalination 2008, 220, 345−352. (b) Yang, Q.; Liu, Y.; Gu, A.; Ding, J.; Shenz, Z. Investigation of Calcium Carbonate Scaling Inhibition and Scale Morphology by AFM. J. Colloid Interface Sci. 2001, 240, 608−621. (22) (a) Mann, S.; Perry, C. C. Structural aspects of biogenic silica. In Silicon Biochemistry; Evered, D., O’Connor, M., Eds.; Ciba Foundation Symposium 121; John Wiley and Sons Ltd.: Chichester, U.K., 1986; pp 40−58. (b) Hildebrand, M. Diatoms, Biomineralization Processes, and Genomics. Chem. Rev. 2008, 108, 4855−4874. (23) Ehrlich, H. Biological Materials of Marine Origin. Invertebrates. Springer: Heidelberg, 2010. (24) Bergna, H. E. Colloid Chemistry of Silica: An Overview. In The Colloid Chemistry of Silica; Bergna, H. E., Ed.; Advances in Chemistry
K. D.; Mavredaki, E. Green Additives to Enhance Silica Dissolution During Water Treatment. Environ. Chem. Lett. 2005, 3, 127−131. (c) Demadis, K. D. Combating Heat Exchanger Fouling and Corrosion Phenomena in Process Waters. In Compact Heat Exchangers and Enhancement Technology for the Process Industries; Shah, R. K., Ed.; Begell House Inc.: New York, 2003; pp 483−491. (d) Demadis, K. D. Silica Scale Inhibition Relevant to Desalination Technologies: Progress and Recent Developments. In Desalination Research Progress; Delgado, D. J., Moreno, P., Eds.; Nova Science Publishers, Inc.: New York, 2008; Chapter 6, pp 249−259. (3) Park, N.; Lee, S.; Yoon, S. R.; Kim, Y. H.; Cho, J. Foulants analyses for NF membranes with different feed waters: Coagulation/ sedimentation and sand filtration treated waters. Desalination 2007, 202, 231−238. (4) Yang, Q. F.; Liu, Y. Q.; Gu, A. Z.; Ding, J.; Shen, Z. Q. Investigation of induction period and morphology of CaCO3 fouling on heated surface. Chem. Eng. Sci. 2002, 57, 921−931. (5) Cho, Y. I.; Choi, B. G. Validation of an electronic anti-fouling technology in a single-tube heat exchanger. Int. J. Heat Mass Transfer 1999, 42, 1491−1499. (6) Li, J.; Liu, J.; Yang, T.; Xiao, C. Quantitative study of the effect of electromagnetic field on scale deposition on nanofiltration membranes via UTDR. Water Res. 2007, 41, 4595−4610. (7) (a) Wang, Y. W.; Kim, Y. Y.; Stephens, C. J.; Meldrum, F. C.; Christenson, H. K. In Situ Study of the Precipitation and Crystallization of Amorphous Calcium Carbonate (ACC). Cryst. Growth Des. 2012, 12, 1212−1217. (b) Rodriguez-Blanco, J. D.; Shaw, S.; Benning, L. G. The kinetics and mechanisms of amorphous calcium carbonate (ACC) crystallization to calcite, via vaterite. Nanoscale 2011, 3, 265−271. (c) Nebel, H.; Neumann, M.; Mayer, C.; Epple, M. On the structure of amorphous calcium carbonateA detailed study by solid-state NMR spectroscopy. Inorg. Chem. 2008, 47, 7874−7879. (8) (a) Kojima, Y.; Endo, N.; Yasue, T.; Arai, Y. Morphological controls of calcium carbonate hexahydrate and its dehydrating substance. J. Ceram. Soc. Jpn. 1995, 103, 1282−1288. (b) Shahar, A.; Bassett, W. A.; Mao, H. K.; Chou, I. M.; Mao, W. The stability and Raman spectra of ikaite, CaCO3·6H2O, at high pressure and temperature. Am. Mineral. 2005, 90, 1835−1839. (9) (a) Mi, B.; Elimelech, M. Gypsum Scaling and Cleaning in Forward Osmosis: Measurements and Mechanisms. Environ. Sci. Technol. 2010, 44, 2022−2028. (b) Le Gouellec, Y. A.; Elimelech, M. Calcium sulfate (gypsum) scaling in nanofiltration of agricultural drainage water. J. Membr. Sci. 2002, 205, 279−291. (c) Brusilovsky, M.; Borden, J.; Hasson, D. Flux decline due to gypsum precipitation on RO membranes. Desalination 1992, 86, 187−222. (d) Akyol, E.; Ö ner, M.; Barouda, E.; Demadis, K. D. Systematic Structural Determinants of the Effects of Tetraphosphonates on Gypsum Crystallization. Cryst. Growth Des. 2009, 9, 5145−5154. (10) (a) Wang, Y. W.; Kim, Y. Y.; Christenson, H. K.; Meldrum, F. C. A new precipitation pathway for calcium sulfate dihydrate (gypsum) via amorphous and hemihydrate intermediates. Chem. Commun. 2012, 48, 504−506. (b) Al-Hadhrami, L. M.; Quddus, A. Role of solution hydrodynamics on the deposition of CaSO4 scale on copper substrate. Desalin. Water Treat. 2012, 21, 238−246. (11) (a) Morales, J.; Manuel Astilleros, J.; Fernandez-Diaz, L. Nanoscopic Characteristics of Anhydrite (100) Growth. Cryst. Growth Des. 2012, 12, 414−421. (b) Fisher, R. D.; Mbogoro, M. M.; Snowden, M. E.; Joseph, M. B.; Covington, J. A.; Unwin, P. R.; Walton, R. I. Dissolution Kinetics of Polycrystalline Calcium Sulfate-Based Materials: Influence of Chemical Modification. ACS Appl. Mater. Interf. 2011, 3, 3528−3537. (c) Shindo, H.; Igarashi, T.; Karino, W.; Seo, A.; Yamanobe-Hada, M.; Haga, M. Stabilities of crystal faces of anhydrite (CaSO4) compared by AFM observation of facet formation processes in aqueous solutions. J. Cryst. Growth 2012, 312, 573−579. (12) (a) Barouda, E.; Demadis, K. D.; Freeman, S.; Jones, F.; Ogden, M. I. Barium Sulfate Crystallization in the Presence of Variable Chain Length Aminomethylenetraphosphonates and Cations (Na+ or Zn2+). Cryst. Growth Des. 2007, 7, 321−327. (b) BinMerdhah, A. B.; Yassin, A. A. M.; Muherei, M. A. Laboratory and prediction of barium sulfate 9039
dx.doi.org/10.1021/ie3010836 | Ind. Eng. Chem. Res. 2012, 51, 9032−9040
Industrial & Engineering Chemistry Research
Article
Series; American Chemical Society: Washington, DC, 1994; Vol. 234, Chapter 1, pp 1−47. (25) Ketsetzi, A.; Stathoulopoulou, A.; Demadis, K. D. Being “green” in chemical water treatment technologies: Issues, challenges and developments. Desalination 2008, 223, 487−493. (26) (a) Mroczek, E.; Christenson, B. Solubility of quartz in the hypersaline brine Implications for fracture permeability. Presented at the World Geothermal Congress, Kyushu, Tohoku, Japan, May 28−Jun 10, 2000. (b) Drever, J. I. The Geochemistry of Natural Waters: Surface and Groundwater Environments, 3rd ed.; Prentice-Hall: Upper Saddle River, NJ, 1997. (c) Dove, P. M. The dissolution kinetics of quartz in aqueous mixed cation solutions. Geochim. Cosmochim. Acta 1999, 63, 3715−3727. (27) Kristmanndóttir, H.; Ó lafsson, M.; Thórhallsson, S. Magnesium silicate scaling in district heating systems in Iceland. Geothermics 1989, 18, 191−198. (28) Kent, D. B.; Kastner, M. Mg2+ removal in the system Mg2+− amorphous SiO 2−H2O by adsorption and Mg-hydroxysilicate precipitation. Geochim. Cosmochim. Acta 1985, 49, 1123−1136. (29) (a) Dubin, L. Silica Stabilization in Cooling Water Systems. In Surface Reactive Peptides and Polymers: Discovery and Commercialization; Sikes, C. S., Wheeler, A. P., Eds.; ACS Symposium Series; American Chemical Society: Washington, DC, 1991; Vol. 444, pp 355−379. (b) Meier, D. A.; Dubin, L. A novel approach to silica scale inhibition. Presented at Corrosion/87, San Francisco, CA, Mar 9−13, 1987; Paper 344. (c) Dubin, L.; Dammeier, R. L.; Hart, R. A. Deposit control in high silica water. Mater. Perform. 1985, 24 (10), 27−33. (30) Smith, C. W. Usage of a polymeric dispersant for control of silica. Ind. Water Treat. 1993, No. July/August, 20−26. (31) Sheikholeslami, R.; Tan, S. Effects of water quality on silica fouling of desalination plants. Desalination 1999, 126, 267−280. (32) (a) Young, P. R. Magnesium Silicate Precipitation. Presented at Corrosion/93, New Orleans, LA, Mar 8−12, 1993; Paper 466. (b) Brooke, M. Magnesium Silicate Scale in Circulating Cooling Systems. Presented at Corrosion/84, New Orleans, LA, Apr 2−6, 1984; Paper 327. (33) MacKenzie, F. T.; Garrels, R. M.; Bricker, O. P.; Bickley, F. Silica in Sea Water: Control by Silica Minerals. Science 1967, 155, 1404−1405. (34) Demadis, K. D.; Neofotistou, E. Use of Antiscalants for Mitigation of Silica (SiO2) Fouling and Deposition: Fundamentals and Applications in Desalination Systems. Desalination 2004, 167, 257− 272. (35) Eaton, A. D.; Clesceri, L. S.; Rice, E. W.; Greenberg, A. E.; Franson, M. H. Standard Methods for Examination of Water & Wastewater; American Public Health Association: Washington, DC, 2005. (36) (a) Cowan, J. C.; Weintritt, D. J. Water Formed Scale Deposits; Gulf Publishing Co.: Houston, TX, 1976; p 245. (b) Hann, W. M.; Robertson, S. T.; Bardsley, J. H. Recent Experiences in Controlling Silica and Magnesium Silicate Deposits with Polymeric Dispersants. In Official Proceedings: The International Water Conference, 54th Annual Meeting; Engineers' Society of Western Pennsylvania: Pittsburgh, PA, 1993; Paper 59, pp 358−370. (37) Stezowski, J. J.; Countryman, R.; Hoard, J. L. Structure of the ethylenediaminetetraacetato-aquomagnesate(II) ion in a crystalline sodium salt. Comparative stereochemistry of the seven-coordinate chelates of magnesium(II), manganese(II), and iron(III). Inorg. Chem. 1973, 12, 1749−1754. (38) Sheikholeslami, R.; Al-Mutaz, I. S.; Koo, T.; Young, A. Pretreatment and the Effect of Cations and Anions on Prevention of Silica Fouling. Desalination 2001, 139, 83−95. (39) (a) Nowack, B. Environmental Chemistry of Aminopolycarboxylate Chelating Agents. Environ. Sci. Technol. 2002, 19, 4009−4016. (b) Knepper, T. P. Synthetic chelating agents and compounds exhibiting complexing properties in the aquatic environment. Trends Anal. Chem. 2003, 22, 708−724. (c) Perry, T. D.; Duckworth, W.; Kendall, T. A.; Martin, S. T.; Mitchell, R. Chelating Ligand Alters the
Microscopic Mechanism of Mineral Dissolution. J. Am. Chem. Soc. 2005, 127, 5744−5745. (40) Koo, T.; Lee, Y. J.; Sheikholeslami, R. Silica fouling and cleaning of reverse osmosis membranes. Desalination 2001, 139, 43−56. (41) Allen, L. H.; Matijevic, E. Stability of colloidal silica. II. Ion exchange. J. Colloid Interface Sci. 1970, 33, 420−429. (42) Chan, S. H.; Chen, Z. J.; He, P. Effects of sodium and potassium chlorides on silica fouling. Presented at the Winter Annual Meeting of the American Society of Mechanical Engineers, Dallas, TX, Nov 25−30, 1990; Paper 90-WA/HT-1. (43) (a) Marshall, W. L.; Warakomski, J. M. Amorphous silica solubilitiesII: Effect of aqueous salt solutions at 25 °C. Geochim. Cosmochim. Acta 1980, 44, 915−917. (b) Chan, S. H.; Neusen, K. F.; Chang, C. T. The solubility and polymerization of amorphous silica in geothermal energy applications. In Proceedings of the ASME−JSME Thermal Engineering Joint Conference; ASME Press: New York, 1987; Vol. 3, p 103. (c) Chan, S. H. A review on solubility and polymerization of silica. Geothermics 1989, 18, 49−56.
9040
dx.doi.org/10.1021/ie3010836 | Ind. Eng. Chem. Res. 2012, 51, 9032−9040
