Elsevier

Geothermics

Volume 89, January 2021, 101954
Geothermics

Searching for a universal scale inhibitor: A multi-scale approach towards inhibitor efficiency

https://doi.org/10.1016/j.geothermics.2020.101954Get rights and content

Highlights

  • Four polymers were tested as inhibitors against several inorganic geothermal scales.

  • Their functional units include polyethylene glycol and phosphonic acid.

  • Inhibitor PEGPHOS-LOW (34 PEG and 14 phosphonate grafts) was the most efficient.

  • Inhibitory efficiency is proportional to the inhibitor affinity (stability constants) for the scaling metal ions.

Abstract

Scale inhibition is of crucial importance in harvesting geothermal energy, an attractive, renewable and sustainable energy source. The majority of the geothermal reservoir waters around the world contain a variety of anions and cations, which are prone to precipitation when subjected to vast transformations of both temperature and pressure, causing failures to many parts of the plants. Our goal is to prevent the formation, precipitation and deposition of such scales by stabilizing their respective components. Our working brines are designed based on the saturated conditions of geothermal well waters. Common inorganic scales tested include amorphous silica, magnesium silicate, aluminum silicate, iron (III) silicate, zinc sulfide, lead sulfide, iron (II) sulfide, and calcium carbonate. Against these scales, four interrelated methacrylate-structured polymers were tested as inhibitors, grafted either with polyethylene glycol (PEG), or phosphonic acid (PHOS) side-chains, or both. Based on several results obtained the inhibitor PEGPHOS-LOW (containing 34 PEG grafts and 14 phosphonate grafts) was selected as the most efficient, and was subsequently tested in artificial geothermal brines of variable stress, containing all the scales together.

Graphical abstract

A family of four methacrylate-structured polymers grafted with either with polyethylene glycol and/or phosphonic acid side-chains, are tested as inhibitors against geothermal scales.

  1. Download : Download high-res image (137KB)
  2. Download : Download full-size image

Introduction

Over the past few decades energy production based on geothermal sources has increased significantly (Shortall and Uihlein, 2019). The main task of geothermal power projects is to convert the energy existing below the earth’s crust into electricity, but other uses such as district heating are also important. The water inside the geothermal wells adsorbs heat from the rock and is transported to the Earth’s surface, where this heat is converted to electrical energy through turbine-generators. These fluids after being processed through energy extraction systems at the surface are injected back into the geothermal reservoir. It is worth-noting that ∼ 97 % of current geothermal reservoir production originates from magmatically-driven reservoirs (Williamson et al., 2001). The most common and modern examples of mineral-rich hydrothermal systems are the Salton Sea geothermal brines in South Eastern California, geothermal brines beneath the Cheleken Peninsula on the Caspian Sea and others. All brines are aqueous solutions containing primarily sodium, potassium, calcium, silica (either as colloidal or silicate), aluminum, zinc, sulfides and chloride (Bucher and Stober, 2010). Brines may also contain significant amounts of other alkali metals, alkaline-earth metals, transition metals, and halides (White, 1981). Not unexpectedly, scale deposits that form in geothermal brines vary considerably in composition, depending on their location and their respective surrounding environments. The precipitation of minerals and other scales from geothermal fluids commonly takes place by decreasing temperature, dilution, increasing pH, reaction with sulfides and redox reactions (Gallup, 1998a). All scale types can present challenging operational problems for geothermal plants. The extreme scaling characteristics of geothermal brines cause catastrophic failures in plant operation. Geothermal waters are very complex and problematic because they pose a risk of composite fouling, which is rare in reverse osmosis and cooling systems. The complexity of issues at hand becomes more challenging when the following issues are considered: (a) supersaturation, (b) very low solubility products of potential scaling salts, (c) high temperatures, (d) complex water chemistry, (e) potential for composite fouling, (f) underdeposit corrosion. Below, certain salient features of representative geothermal scales are presented, in a concise manner.

Silicon dioxide is widespread in Nature, while silicate minerals constitute over 90 % of the Earth’s crust (Gaspar, 2011). In addition, a large number of organisms, like several species of sea sponges, terrestrial plants and marine organisms, use biogenic silica to build their biological membranes and exoskeletons (Puppe, 2020). Biosilicification is the directed formation by such microorganisms of amorphous hydrated silica (biosilica) from water-soluble “Si” sources (Mann, 2002). The important process of biosilicification is considered to be a special kind of biomineralization as biosilica seems to be different from the plethora of various biogenic, metal-containing minerals (e.g. calcite, aragonite, hydroxyapatite, iron sulfides etc.) (Wysokowski et al., 2018).

Besides the biological world, silica, in its colloidal/amorphous form, is one of the most notorious scales formed in geothermal brines (Tassew, 2001). Its peculiarities include its amorphous nature, complex auto-condensation chemistry, pH-dependent formation, temperature-dependent solubility and propensity for deposition (Kristmannsdóttir, 1989). Crystalline forms of silica, such as quartz, usually are not deposited in geothermal fluids due to slow crystal formation kinetics. On the other hand, amorphous silica (or metal silicates, see below) tends to form and deposit rapidly due to fast polymerization kinetics of silicic acid and aggregation kinetics of silica nanoparticles (Iler, 1979).

Silica is formed by a complicated inorganic polymerization process and has an amorphous nature in contrast with metal carbonate and phosphate solids which are crystalline materials (Ehrlich et al., 2010). In living microorganisms, such as diatoms and sponges, the formation of unique nanopatterned silica structures is a spectacular phenomenon and a remarkable source of inspiration for biomimetic chemistry (Coradin and Lopez, 2003). However formation of silica in several engineering applications, such as industrial facilities, where the water is used as a supporting (cooling, heating) medium, is not as desirable as in biological systems (Ehrlich and Worch, 2007). In supersaturated silica-containing process waters, silicic acid polymerizes and the resulting amorphous silica precipitate is gradually transformed into an undesirable hard scale deposit (Amjad and Demadis, 2015).

Combating these issues requires understanding the process of silicon polymerization and the factors that affect it, either positively or negatively. In dilute aqueous solutions, “soluble silica” exists in the form of monosilicic acid Si(OH)4. In circumneutral pH media the monomer is found mostly in its protonated form. At higher pH values the monomer is deprotonated forming silicate ions [Si(OH)3O] (Coradin and Livage, 2001). As the silicate concentration increases, the condensation of two monomers takes place with a simultaneous loss of one water molecule. This condensation reaction involves a nucleophilic attack of a deprotonated silicic acid anion onto a fully protonated silicic acid molecule (Neofotistou and Demadis, 2004). This step is the rate-determining step and the most crucial one in the complicated silica condensation process (Tarutani, 1989). Thus, at pH values 7–8.5 the process is very fast as both protonated and deprotonated species coexist. In strongly acidic media deprotonated species cannot form, and at pH values > 9 fully protonated species do not exist, so that the nucleophilic substitution cannot take place (Perry and Keeling-Tucker, 1998).

Inhibition activity of several biopolymers and synthetic polymers on silica formation has been a part of an on-going investigation in our laboratory (Demadis et al., 2008; Mavredaki et al., 2005; Papathanasiou et al., 2019; Demadis and Neofotistou, 2007; Demadis, 2005; Demadis and Stathoulopoulou, 2006a, b). Research thus far has revealed that cationic polymers possessing protonated amine groups (Mavredaki et al., 2005; Papathanasiou et al., 2019; Demadis and Neofotistou, 2007; Demadis, 2005; Demadis and Stathoulopoulou, 2006b), quaternary ammonium groups (Demadis and Stathoulopoulou, 2006b), or phosphonium groups (Demadis et al., 2012a, 2012b) have the ability to retard silica polycondensation through electrostatic interactions with the negatively charged silicate anions (Spinthaki et al., 2016). Apart from electrostatic interactions, H-bonding has been proved to influence silica formation. For example, results on the inhibitory activity of uncharged macromolecules possessing ether moieties (PEGs) have been previously reported (Preari et al., 2014).

Magnesium silicates are either natural or synthetic (Ciesielczyk et al., 2007). Magnesium silicates are a family of geological entities that embraces several types of minerals consisting of both magnesium (in its Mg2+ state) and silicon (in its SiO44− state). They exist in a variety of different Mg:Si stoichiometries and hydration states. In the field of geology, these structures are crystallographically well–defined, but when it comes to inorganic precipitation/deposition in the water treatment industry, their true identity is elusive, and, occasionally, controversial. In the field, magnesium silicate is considered to be any water-formed solid that contains both Mg and Si in no specified stoichiometries. The molecular formula of such compounds may be expressed as MgSiO3·xH2O (Taspinar and Ozgul-Yucel, 2008). The magnesium/silicate system is highly pH-dependent. Below pH 7 there is essentially no precipitation, because silica exists in an unreactive (towards Mg2+), non-ionized form. At pH above 9, magnesium silicate is very likely to form because silica forms reactive silicate ions. Previous work in our group was reported on the formation and characterization of magnesium silicate precipitates/deposits under geothermal stresses (Spinthaki et al., 2018b). Co-precipitation of magnesium hydroxide, Mg(OH)2, and colloidal silica has also been observed (Dubin et al., 1985; Dubin, 1991). One scenario proposes that the formation of Mg(OH)2 takes place first, and it subsequently reacts with monomeric silicate and/or polymeric silica to form magnesium silicate (Smith, 2002). Ca2+ and Mg2+ salts were found to catalyze the silica auto-condensation process (Sheikholeslami and Tan, 1999).

There is a limited number of reports on the use of chemical additives in “magnesium silicate” stabilization. It has been reported that ethylenediaminetetraacetic acid (EDTA) can effectively compromise the catalytic role of magnesium cations that enhance colloidal silica formation, especially when the pH exceeds 9 (Demadis et al., 2012a, 2012b). This is ascribed to the well-known high affinity of EDTA for Mg2+ at high pH values (Stezowski et al., 1973). The alkaline/surfactant/polymer (ASP) flooding approach has been used to control silicate scale formation in oil recovery operations (Umar and Saaid, 2014). Lately, organic inhibitors seem to be an effective solution to mitigate silica and silica-containing scales (Gallup, 2002; Gallup and Barcelon, 2005). Amorphous precipitates obtained in the presence of Mg and silica are common (Krysztafkiewicz et al., 2004). Unfortunately, precipitate amorphicity precludes the use of X-ray powder diffraction for unequivocal material characterization. Therefore, the absence of any identifiable crystalline phase is a strong indication that none of the well-known magnesium silicate mineral phases are formed in our experiments. Magnesium silicate is a “true” compound with a Si:Mg ratio of 1:1, according to Young, 1993. Just like the (colloidal/amorphous) silica system, the magnesium silicate system is highly temperature-dependent. It has been suggested that precipitation begins at a lower pH if the temperature is sufficiently high.

Due to the high reactivity of aluminum aqueous species, Al3+ is usually found in combination with oxygen and silicon in feldspars, micas and clay minerals. Usually, in geothermal brines, aluminum is dissolved in very low concentrations (< 5 ppm) (Castet et al., 1993). This happens due to the low solubility of aluminum and aluminosilicate minerals. Aluminum’s concentration in natural waters is controlled by the pH and by very fine suspended mineral particles. It has been proposed, that during the formation of amorphous aluminum silicate, Al3+ ions are incorporated in the silica matrix, replacing Si4+ ions (Gallup, 1997, 1998b). Aluminum silicate does not form deposits at pH ≤ 4 and pH ≥ 10. Aluminum accelerates the precipitation rate of silica at pH > 9 and retards it between pH 5 and 8. Higher temperatures enhance the formation of aluminum silicate scale deposits. Also, a higher concentration of aluminum species in the brine leads to the enhanced formation of aluminum silicate. Strategies based on brine pH modification have been used to control the formation of aluminum-rich amorphous silica (Violante and Huang, 1985). Complexing/sequestering agents like EDTA, citrate and acetate can inhibit aluminum silicate scale formation (Gallup, 1997, 1998b).

Hazel et al. observed as early as 1949 the first evidence of interaction between ferric ions and silicate species in aqueous solutions (Hazel et al., 1949). They employed titrimetric procedures to qualitatively study metal–silicate interactions with metals such as Al, Fe, and Cr. No quantitative relationships were established for any of these interactions until the work of Weber and Stumm delineated the formation of a Fe(III)-silicate complex (Weber and Stumm, 1965).Fe3+ + Si(OH)4 + H2O → [FeSiO(OH)3]2+ + H3O+

The experimental pH covered was <3.5, however it is expected that similar interactions take place in natural water pH. Increased silica concentrations accelerated Fe(II) to Fe(III) oxidation in pH ranges 6.6–7.1 (Schenk and Weber, 1968; O’ Melia and Stumm, 1967). Fe(III) has a higher propensity for silicate or colloidal silica than Fe(II). This may create iron (III) silicate precipitates. If there is a significant concentration of iron in the brine fluid, deposition of iron silicates will set in at a higher temperature than the silica deposition but at lower temperatures iron tends to be deposited in the form of oxides (Gunnlaugsson et al., 2014). Chan et al. (1995) studied the effect of Fe3+ on silica fouling of a pilot heat exchanger (at 125 °C – 165 °C, 1.58 MPa and under turbulent conditions). It was shown that the addition of a very small concentration of Fe3+ into a silica solution was sufficient to induce coagulation and deposition of silica particles, which caused a significant increase in fouling rates. Gallup (1989) studied iron silicate formation in the Salton Sea geothermal field and its inhibition, suggesting that they appear as brown-black, vitreous solids resembling obsidian and have been observed depositing in brine-handling equipment and wells at rates ranging from ∼0.5 cm/yr in production wells (>200 °C) to over 50 cm/yr in injection pipes and wells (150–175 °C). These iron-rich scales deposited from the aforementioned geothermal brine at high temperatures were shown to consist of what is thought to be a compound of an empirical formula Fe(OH)3·SiO2·xH2O or Fe2O3·2SiO2·xH2O (Gallup, 1991). Several characterization techniques such as Mössbauer and IR spectroscopy and X-ray diffraction studies have been performed on samples from the same geothermal field by Gallup and Reiff (1991) in order to fully comprehend the nature of these deposits. They suggested that ferric (Fe3+) and ferrous (Fe2+) iron are present as hydrous silicates resembling the mineral hisingerite and exhibiting variable composition. They also detected other iron minerals which were microcrystalline, poorly-crystalline or glassy. Their study suggested that iron deposited in scales appears to be primarily derived from brine, but is also found to be present as poorly-crystalline steel corrosion products. Adding to these observations, Manceau et al. (1995) also used X-ray diffraction and spectroscopic techniques to confirm that a high-temperature (250 °C) scale precipitated from non-oxidized geothermal brines at Salton Sea contains polymerized ferric iron and silica and identified the mineralogical nature of this precipitate as hisingerite. Hisingerite is a poorly-crystallized, non-stoichiometric nontronite. Finally, the results of elemental analysis studies suggest that poorly crystalline species at Salton Sea that XRD cannot account for, are mainly magnesium iron silicates (Gallup and Hinrichsen, 2008). It was observed that the formation of these species is highly pH – dependent, since they mostly form when pH is high enough, so that the iron present in solution can form hydroxides (Fe(OH2) and Fe(OH)3). These species are then able to react with amorphous silica and yield proper charge balances if they include lattice cations like Mg2+. While most literature on iron silicates in geothermal brines refers to the Salton Sea geothermal field, similar scales have also been documented at other relevant sites, such as Cerro Prieto, Mexico (Mercado et al., 1989), Milos island, Greece (Karabelas et al., 1989), Tiwi, Philippines (Gallup, 1993), and Reykjanes, Iceland (Kristmannsdóttir, 1984).

Sulfide scales are also encountered in geothermal operation and have been observed in high temperature as well as in low/medium temperature resources. They are commonly associated with other metal cations, forming scale deposits that are very tenacious and difficult to handle. In contrast to silicates, zinc sulfide (ZnS), as it occurs in geothermal brines, is a well – characterized crystalline mineral salt and has a solubility constant (Ksp) of ∼ 2 × 10−25 (Smith and Martel, 1976). Precipitation of zinc sulfide should also take into account Zn2+ speciation in aqueous systems (Choi et al., 2013). An increase in pH reduces the actual concentration of Zn2+, due to Zn(OH)2 precipitation, a competitive reaction to the precipitation of ZnS. In addition, S2- is the dominant species at pH > 11, while at circumneutral pH, HS is dominant. The extremely low solubility of ZnS drives the fast deprotonation of HS, which shifts the equilibrium and reduces the pH. The most common ZnS precipitates are sphalerite (or zinc blend) and wurtzite (Parvaneh et al., 2015).

Iron sulfides are often found in scale deposits formed by the flow of high-enthalpy and high-salinity geothermal brines, such as those in the widely studied Salton Sea, California, USA (Austin et al., 1977; Skinner et al., 1967), Milos Island, Greece (Karabelas et al., 1989), and Asal, Djibouti (Criaud and Fouillac, 1989). These scales are usually intermixed with amorphous silica to form iron silicates, as mentioned in the “Iron Silicate” paragraph. It has been observed that iron sulfide precipitation, in the form of a black amorphous phase, is pH-dependent. The nature of these deposits has not yet been clearly defined, in spite of several studies. So far, a phase with similar characteristics was identified as amorphous iron sulfide in the literature and has been considered as a precursor phase for mackinawite (Rickard, 1969). Iron sulfide scales are also been reported in several low temperature geothermal systems in the Dogger region (Criaud and Fouillac, 1989). Several crystalline phases have been identified, with typical examples including monosulfides such as troilite, pyrrhotite, mackinawite and pyrite. The majority of the scales that contain sulfide and have been encountered in geothermal plants consist of at least two major sulfides. Specifically for the case of iron sulfide scales, the scale composition may vary from location to location and exist as various polymorphs. The most stable phase is pyrite, a fact that does not preclude the occurrence of several other phases, some of them metastable, at low temperatures. Taylor reported that at higher temperatures, (> 200 °C), only a few sulfide phases are stable (pyrrhotite and pyrite), while at atmospheric conditions several phases are expected to occur (Taylor, 1987). It has been established that these deposits are prone to oxidation, a fact that is qualitatively proven by the change of black color to reddish-brown after air exposure.

The presence of Fe and sulfate ions in geothermal brines could be due to metallic corrosion, and/or bacterial reduction of sulfate (SO42−) to sulfide species (HS-, S2−), respectively. The rate of precipitation of iron sulfides can be influenced by several parameters, such as pH, temperature and supersaturation. Andritsos and Karabelas reported that the maximum rate for FeS formation remains virtually unaffected by temperature, for a given pH (Andritsos and Karabelas, 1995). From the point of view of morphology, the iron sulfide systems can be categorized in the group of deposits with a rather coherent deposit layer.

Several heavy metal sulfides are found in scale deposits formed by the flow of high-enthalpy and high-salinity geothermal brines. Lead sulfide is the most likely to appear of the metal sulfides family in high total dissolved solids / high enthalpy fluids. Literature provides studies of deposition rate, obtained in once-through experiments, aiming to the uncovering of the mechanism responsible for lead sulfide scaling (Andritsos and Karabelas, 1991a). PbS formation and precipitation is profoundly influenced by the solution pH, flow rate, temperature and ionic strength (salinity). The existing data show that PbS solubility is minimum at pH 7, as for Fe and Zn sulfides (Andritsos and Karabelas, 1995). Temperature and salinity have been reported to shift the “worst-case-scenario” pH to higher values (Andritsos and Karabelas, 1991b). Supersaturation and pH are factors that also affect the morphology and growth rate of PbS deposits. Specifically, formation of relatively large PbS particles is favored at low supersaturation (Andritsos and Karabelas, 1991a).

In 1952, J.P. Miller (1952) calculated the solubility of CaCO3 at temperatures ranging from 0 °C up to 105 °C and pressures of 1–100 bars. It was concluded that higher temperatures result in a lower CaCO3 solubility. Hence, it is not surprising that CaCO3 scales appear in geothermal systems (Ueckert et al., 2020). Calcium carbonate deposition is influenced by degassing of CO2 shifting the pH to more alkaline values. Calcite scale is very common in production wells for high-medium temperature reservoirs. When geothermal fluids reach the flash point in the well, water and steam phases separate from each other. As a result of the CO2 degassing, pressure drops and pH changes occur (Haklidir and Haklidir, 2017). Unlike silica, calcite is a less soluble mineral at higher temperatures. Efforts to control CaCO3 scaling are available in the literature in harsh conditions (geothermal and oilfield) (Fan et al., 2012; Kelland, 2011; Matty and Tomson, 1988; Khormali and Petrakov, 2016; Spinthaki and Demadis, 2020; Zotzmann et al., 2018). Acrylate-based, hydrophilic polymers have been used to control CaCO3 in the presence of other scales (Wang et al., 2017).

Section snippets

Rationale of the study

The main objective of the present study was to discover and evaluate scale inhibitor(s) that will possess the ability to inhibit multiple scaling salts, ie. a “universal” scale inhibitor. An alternative approach would be the utilization of inhibitor blends (Zhang et al., 2020; Gwak and Hong, 2017; Amjad, 2019). This approach has been shown to be effective, but for much simpler systems (Mavredaki et al., 2005). Such antiscalant mixtures must contain a multitude of scale inhibitors, each one

Instrumentation

ATR-IR spectra were collected on a Thermo-Electron NICOLET 6700 FTIR optical spectrometer. Determinations of soluble silicic acid (molybdate-reactive silica) and soluble sulfide were carried out with a HACH 1900 spectrophotometer from the Hach Co., Loveland, CO, USA. SEM (Scanning Electron Microscopy) images and EDS (Electron Dispersive Spectrometry) data were collected on a scanning electron microscope LEO VP-35 FEM.

Reagents and materials

All chemicals were from commercial sources, and were used as received. Sodium

Polymeric inhibitors

The polymers HOMOPEG, HOMOPHOS possess the same polymethacrylate backbone, but with different grafted esters (see Fig. 2). HOMOPEG contains a short-chain PEG graft with 20 ethylene glycol units, while HOMOPHOS contains a methylphosphonic acid graft. HOMOPEG is pH-insensitive and remains uncharged regardless of solution pH, in contrast to HOMOPHOS, which contains acidic phosphonic acid moieties. Based on the pKa values of methylphosphonic acid (pKa1 = 2.21 ± 0.02 and pKa2 = 7.48 ± 0.06) (Popov

Discussion

In the previous section inhibition results were presented and categorized according to each individual scale. It would be very informative to view inhibition efficiency based on the individual inhibitors. This is attempted in Fig. 12.

HOMOPEG is an uncharged methacrylate-based polymer with 20-unit PEG grafts. As seen in Fig. 12 upper left, it is virtually inactive towards all scales at pH = 8.5, even at a high dosage of 400 ppm. This is not unexpected, at least for the crystalline (metal

Conclusions

Herein, we implemented an initial systematic attempt to study the inhibitory effects of four graft co-polymers based on a methacrylate backbone, and possessing phosphonate and PEG grafts. The individual scales against which these four inhibitors were tested were amorphous silica, magnesium silicate, aluminum silicate, iron (III) silicate, zinc sulfide, lead sulfide, iron(II) sulfide, and calcium carbonate. Based on the results obtained, the most efficient inhibitor is PEGPHOS-LOW. This

Author contributions

The manuscript was written through contributions of all authors. All authors have given approval to the final version of the manuscript.

Declaration of Competing Interest

The authors report no declarations of interest.

References (104)

  • D.L. Gallup

    Investigations of organic inhibitors for silica scale control in geothermal brines

    Geothermics

    (2002)
  • D.L. Gallup et al.

    Investigations of organic inhibitors for silica scale control from geothermal brines–II

    Geothermics

    (2005)
  • D.L. Gallup et al.

    Characterization of geothermal scale deposits by Fe-57 Mössbauer spectroscopy and complementary X-ray diffraction and infra-red studies

    Geothermics

    (1991)
  • G. Gwak et al.

    New approach for scaling control in forward osmosis (FO) by using an antiscalant-blended draw solution

    J. Memb. Sci.

    (2017)
  • G.A. Icopini et al.

    Kinetics of silica oligomerization and nanocolloid formation as a function of pH and ionic strength at 25 °C

    Geochim. Cosmochim. Acta

    (2005)
  • A.J. Karabelas et al.

    Characteristics of scales from the Milos geothermal plant

    Geothermics

    (1989)
  • A. Ketsetzi et al.

    Being “Green” in chemical water treatment technologies: issues, challenges and developments

    Desalination

    (2008)
  • H. Kristmannsdóttir

    Types of scaling occurring by geothermal utilization in Iceland

    Geothermics

    (1989)
  • A. Krysztafkiewicz et al.

    Amorphous magnesium silicate, synthesis, physicochemical properties and surface morphology

    Adv. Powder Technol.

    (2004)
  • S. Lacour et al.

    Complexation of trivalent cations (Al(III), Cr(III), Fe(III)) with two phosphonic acids in the pH range of fresh waters

    Talanta

    (1998)
  • J.M. Matty et al.

    Effect of multiple precipitation inhibitors on calcium carbonate nucleation

    Appl. Geochem.

    (1988)
  • S. Mercado et al.

    Scale incidence on production pipes of Cerro Prieto geothermal wells

    Geothermics

    (1989)
  • E. Neofotistou et al.

    Silica Scale Growth Inhibition By Polyaminoamide starburst® Dendrimers

    Colloids Surf. A Physicochem. Eng. Asp.

    (2004)
  • A.H. Nielsen et al.

    Effects of pH and Iron concentrations on sulfide precipitation in wastewater collection systems

    Water Environ. Res.

    (2008)
  • C.R. O’ Melia et al.

    Aggregation of silica dispersions by Fe(III)

    J. Colloid Interface Sci.

    (1967)
  • C.C. Perry et al.

    Aspects of the bioinorganic chemistry of silicon in conjunction with the biometals calcium, iron and aluminium

    J. Inorg. Biochem.

    (1998)
  • D. Puppe

    Review on protozoic silica and its role in silicon cycling

    Geoderma

    (2020)
  • R. Sheikholeslami et al.

    Effects of water quality on silica fouling of desalination plants

    Desalination

    (1999)
  • R.M. Smith et al.

    Critical Stability Constants, Volume 4: Inorganic Complexes

    (1976)
  • M.F.B. Sousa et al.

    New methodology based on static light scattering measurements for evaluation of inhibitors for in bulk CaCO3 crystallization

    J. Colloid Interface Sci.

    (2014)
  • A. Spinthaki et al.

    The precipitation of magnesium silicate under geothermal stresses: formation and characterization

    Geothermics

    (2018)
  • A. Tramaux et al.

    Synthesis of phosphonated comb-like copolymers and evaluation of their dispersion efficiency on CaCO3 suspensions. Part I: effect of an increasing phosphonic acid content

    Powder Technol.

    (2018)
  • W.J. Weber et al.

    Formation of a silicato-iron (III) complex in dilute aqueous solution

    J. Inorg. Nucl. Chem.

    (1965)
  • Z. Amjad

    Scale inhibitor blends for industrial water systems

    Mater. Perform.

    (2019)
  • N. Andritsos et al.

    An experimental study of sulfide scale formation in pipes

  • A.L. Austin et al.

    The Geothermal Status Report, Report UCRL-50046-76

    (1977)
  • J.F. Blais et al.

    Metals precipitation from effluents: Review. Practical periodical of hazardous

    Toxic Radioactive Waste Manage.

    (2008)
  • K. Bucher et al.

    Fluids in the upper continental crust

    Geofluids

    (2010)
  • S.H. Chan et al.

    Effect of ferric chloride on silica fouling

    Int. J. Heat Mass Transf.

    (1995)
  • C.H. Choi et al.

    Effects of fluid flow on the growth and assembly of ZnO nanocrystals in a continuous flow microreactor

    CrystEngComm

    (2013)
  • F. Ciesielczyk et al.

    Physicochemical studies on precipitated magnesium silicates

    J. Mater. Sci.

    (2007)
  • T. Coradin et al.

    Biogenic Silica Patterning: Simple Chemistry or Subtle Biology?

    ChemBioChem

    (2003)
  • K.D. Demadis

    A structure/function study of polyaminoamide (PAMAM) dendrimers as silica scale growth inhibitors

    J. Chem. Technol. Biotechnol.

    (2005)
  • K.D. Demadis et al.

    Synergistic effects of combinations of cationic polyaminoamide dendrimers/anionic polyelectrolytes on amorphous silica formation: a bioinspired approach

    Chem. Mater.

    (2007)
  • K.D. Demadis et al.

    Novel, multifunctional, environmentally friendly additives for effective control of inorganic foulants in industrial water and process applications

    Mater. Perform.

    (2006)
  • K.D. Demadis et al.

    Solubility enhancement of silicate with polyamine/polyammonium cationic macromolecules: relevance to silica-laden process waters

    Ind. Eng. Chem. Res.

    (2006)
  • K.D. Demadis et al.

    Inhibitory effects of multicomponent, phosphonate-grafted, zwitter-ionic chitosan biomacromolecules on silicic acid condensation

    Biomacromolecules

    (2008)
  • K.D. Demadis et al.

    Catalytic effect of magnesium ions on silicic acid polycondensation and inhibition strategies based on chelation

    Ind. Eng. Chem. Res.

    (2012)
  • K.D. Demadis et al.

    Promiscuous stabilisation behavior of silicic acid by cationic macromolecules: the case of phosphonium-grafted dicationic ethylene oxide bolaamphiphiles

    RSC Adv.

    (2012)
  • Cited by (19)

    • An antiscalant with chelating residues of amino acid glycine

      2022, Desalination
      Citation Excerpt :

      A review article summarizes the nuances of scale deposition in oil field reservoirs and their mitigation using synthetic and natural-based environmentally benign green scale inhibitors [28]. In search for a universal inhibitor for scale formation in geothermal energy plants, a polymethacrylate grafted with polyethylene glycol or phosphonic acid has been shown to inhibit scale formation [29]. Although very efficient in mitigating the formation of scale in the desalination plants, nonbiodegradable phosphorous containing antiscalants are not environmentally friendly since their discharge in the marine ecosystems causes eutrophication associated with the increased growth of algae [20,30,31].

    • Testing the stability of chemical inhibitors at geothermal conditions and their efficiency to prevent galena formation

      2022, Geothermics
      Citation Excerpt :

      The amount of galena precipitated in the Balmatt geothermal brine is, however, much higher compared to galena precipitation in the Soultz-sous-Fôrets and Rittershoffen geothermal plant. Galena scaling can be often related to geothermal systems with high enthalpy/high amount of totally dissolved solids (Spinthaki et al., 2021). The Balmatt brine contains approximately 165 g L-1 TDS (Bos and Laenen 2017), which is higher compared to the Rittershoffen brine and the Soultz-sous-Fôrets brine (both 100 g L-1 TDS) (Genter et al., 2000; Ledésert and Hébert 2012; Mouchot et al., 2021).

    • Synthesis, characterization and scale inhibition performance evaluation of novel dendrimers with the initiator core of pentaerythritol derivative

      2022, Desalination
      Citation Excerpt :

      In summary, it's critically important to inhibit the scale deposition/precipitation for the heat exchange process, RO membrane separation process, and the ASP process of the petroleum industry. Scale inhibitors, a series of fine chemicals, have been widely applied to achieve scaling prevention [22]. At present, the most commonly used scale inhibitors are still phosphates, phosphonates, and polycarboxylates [23].

    View all citing articles on Scopus
    View full text