Antiscalant-Driven Inhibition and Stabilization of âMagnesium

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Antiscalant-driven inhibition and stabilization of “magnesium silicate” under geothermal stresses: The role of magnesium-phosphonate coordination chemistry Argyro Spinthaki, Juergen Matheis, Wolfgang Hater, and Konstantinos D Demadis Energy Fuels, Just Accepted Manuscript • DOI: 10.1021/acs.energyfuels.8b02704 • Publication Date (Web): 25 Sep 2018 Downloaded from http://pubs.acs.org on September 30, 2018

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Antiscalant-driven inhibition and stabilization of “magnesium silicate” under geothermal stresses: The role of magnesium-phosphonate coordination chemistry Argyro Spinthaki,# Juergen Matheis,§ Wolfgang Hater§ and Konstantinos D. Demadis*# #

Crystal Engineering, Growth and Design Laboratory, Department of Chemistry, University of Crete,

Voutes Campus, Heraklion, Crete, GR-71003, Greece §

Kurita Europe GmbH, Giulinistrasse 2, 67065 Ludwigshafen, Germany

AUTHOR EMAIL ADDRESS [email protected] RECEIVED DATE TITLE RUNNING HEAD. “Phosphonate Inhibitors for Magnesium Silicate” ABSTRACT The formation, precipitation, and deposition of the so-called “magnesium silicate” in geothermal waters have been subjects of intense interest. Such scaling poses a severe threat to the smooth process of industrial systems. The aim of our approach is the systematic study of the influence of phosphonatebased chemical additives on silica polycondensation chemistry, in the presence of magnesium ions. The focus of this work is the prevention of “magnesium silicate” formation, using a variety of well-known phosphonate additives. These are PBTC (2-phosphonobutane-1,2,4-tricarboxylic acid), HEDP (hydroxyethylidene-1,1-diphosphonic acid), AMP (amino-tris(methylenephosphonic acid)), HDTMP ACS Paragon Plus Environment

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(hexamethylenediamine-tetrakis(methylenephosphonic acid)) and BHMTPAMP (bis-hexamethylenetriamine-pentakis(methylenephosphonic

acid)).

Inhibition

experiments

were

carried

out

in

supersaturated solutions of silicate (200 ppm, expressed as SiO2) and magnesium (200 ppm, as Mg) at pH 10.0. The phosphonate additives used were found to act as stabilizing agents, most likely by complexing the Mg2+ cations, and thus preventing “magnesium silicate” formation. Based on a plethora of experimental data, a number of useful functional insights have been generated, which add to building a more complete and comprehensive picture of the mechanism of “magnesium silicate” formation and stabilization.

KEYWORDS. Magnesium silicate, inhibition, additives, water treatment, geothermal, phosphonates

Introduction Metal silicates form in waters with high silica and low salinity content, in the presence of certain metal ions.1,2 Iron and magnesium silicates exhibit low solubility in brines with low enthalpy,3 while aluminum silicate usually precipitates in fluids of higher temperatures.4 Although the term “magnesium silicate” is widely recognized in the water treatment industry, its definition differs from its geological counterparts. The exact identity of the scale from the water treatment point of view remains unknown, while the geological minerals are well – defined structures. The magnesium silicate system, just like the silica system, is highly pH dependent. Precipitation practically occurs at pH range of 8.5 to 10.0, with increasing tendency as the pH increases. This is because at pH range of 8.5 to 10.0 monomeric silicic acid [Si(OH)4] becomes increasingly deprotonated,5 and as it forms silicate [Si(OH)3O-] it is reactive towards the magnesium cations due to electrostatic attraction. Below pH 7 there is practically no magnesium silicate precipitation, although silicic acid undergoes fast polycondensation.6-10 Studies have shown that Mg2+ ions actually play a ACS Paragon Plus Environment

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catalytic role in silica polycondensation chemistry.11,12 One suggestion regarding this matter is that magnesium hydroxide co – precipitates with silicic acid to form non – stoichiometric forms of the scale. Another theory suggests that magnesium hydroxide occurs first and then reacts with pre – formed colloidal silica to form magnesium silicate.13 In fact, it has been suggested that when magnesium hydroxide precipitates, any silica present will be adsorbed.14 The magnesium-to-silicate ratio is one of the most controversial topics. It was suggested that the ratio is 1:1, but shows some variation.15 Since then, many researchers proposed different ratios, such as 1:2, 1:5, 1:10 and even 1:20, but the matter still remains under consideration.16 The fact is that the scale is amorphous and it is not recognized as a member of the extensive mineral family. Amongst others, it is considered to be one of the most persistent scales. Primary methods that have been utilized to remove deposits from the industrial equipment. Mechanical scraping17 and chemical cleaning,18 or a combination of both19 are common approaches, which require a temporary suspension of operation. Regarding scale inhibition, the developed methods usually include hot brine reinjection,20 adjustment of the brine pH,21 are some approaches used. Brine acidification has been in some cases successfully applied, but one has still to consider the possibility of corrosion, the additional cost of acids and most importantly, human safety.22 The first step in solving a problem is the identification of the problem itself. The development of scale inhibitors requires an understanding of the mechanism of crystal growth, the scales true nature and its inhibition. In the prequel of the present work, we proposed the formation and characterization of magnesium silicate under geothermal stresses, coming one step closer in revealing the nature of the socalled magnesium silicate. Recently, we published detailed studies on the precipitation of “magnesium silicate” (in the absence of additives) in geothermally-relevant artificial waters.23 The variables studied were Mg2+ and silicate concentration, pH and temperature. In summary, important findings were that precipitate formation is pH-dependent (as pH increases, the precipitation driving force increases), precipitation is more pronounced with increased silica and Mg2+ concentrations, all precipitates are amorphous, Mg elemental distribution within the isolated precipitates seems to be random, pointing to a ACS Paragon Plus Environment

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non-stoichiometric precipitation process, and temperature-dependence of precipitation is small. Our proposal was that the best term to describe the inorganic precipitates formed under certain conditions (in the presence of Mg2+ and silicate at pH 10) is “magnesium-containing amorphous silica”. In this paper, we report the influence of five phosphonate-based chemical additives on precipitate formation in the presence of Mg2+ ions and soluble silica (silicic acid). The additives are PBTC (2phosphonobutane-1,2,4-tricarboxylic acid),24 HEDP (hydroxyethylidene-1,1-diphosphonic acid),25 AMP (amino-tris(methylenephosphonic tetrakis(methylenephosphonic

acid)),27

acid)),26

HDTMP

(hexamethylenediamine-

and

BHMTPAMP

(bis-hexamethylene-triamine-

pentakis(methylenephosphonic acid))28 and their schematic structures are shown in Figure 1.

Figure 1. Schematic structures of the chemical additives used in this study with their abbreviated names.

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There are a limited number of reports on the use of chemical additives in “magnesium silicate” stabilization. We reported that EDTA can effectively compromise the catalytic role of magnesium cations that enhance colloidal silica formation, especially when pH exceeds 9.13 This is ascribed to the well-known high affinity of EDTA for Mg2+ at high pH values.29 Gallup et. al. suggested the use of a silicon–complexing amount of fluoroborate species.30 The alkaline/surfactant/polymer (ASP) flooding approach has been used to control silicate scale formation in oil recovery operations.31 Lately, organic inhibitors seem to be an effective solution to mitigating silica and silica-containing scales.32,33 Reducing agents have been tested for iron silicates inhibition,34 and chelating agents for aluminum silicates mitigation.35,36

Experimental Section

Instrumentation. ATR-IR spectra were collected on a Thermo-Electron NICOLET 6700 FTIR optical spectrometer. Measurements of soluble silicic acid (molybdate-reactive silica) were carried out using a HACH 1900 spectrophotometer from the Hach Co., Loveland, CO, USA. SEM, EDS and elemental mapping were carried out on a scanning electron microscope LEO VP-35 FEM. Hydrothermal reactions were carried out using a parallel synthesis procedure in an autoclave block made of aluminum which contains 36 reaction chambers in a 6×6 array. Teflon reactors have an inner diameter of 19 mm and a depth of 18 mm, with a total volume of about 5 mL. A thin sheet of Teflon covers the reaction vessels, which are then sealed inside a specially designed aluminum autoclave. The set-up is shown in Reference 23. Reagents and Materials. All chemicals were from commercial sources. Phosphonate additives (Figure 1) were members of the Dequest® series supplied by ThermPhos Inc. (acquired by Italmatch), with the following commercial names: PBTC Dequest 7000, HEDP Dequest 2010, AMP Dequest 2000, ACS Paragon Plus Environment

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HDTMP Dequest 2054, and BHMTPAMP Dequest 2090. 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. Magnesium chloride hexahydrate (MgCl2.6H2O) was from Riedel de Haen. Ethylenediamine tetraacetic acid (EDTA, tetrasodium salt) and Eriochrome Black T were from Alfa Aesar. All reagents were used as received from suppliers, without further purification. Acrodisc filters (0.45 µm) were from PallGelman Corporation. In-house, deionized (DI) water was used for all experiments. This water was tested for soluble silica and magnesium ions and was found to contain negligible amounts. Solution preparation. The sodium silicate stock solution (solution A) used in all experiments was prepared by dissolving 4.080 g of Na2SiO3·5H2O in 2 L DI water. The mixture is stirred for at least a day to ensure complete dissolution. This stock solution contains 500 ppm “Si” (8.33 mM), expressed as “ppm SiO2”. A 10,000 ppm in Mg2+ solution (solution B) was prepared by dissolving 8,360 g of MgCl2.6H2O in 100 mL DI water. A 0.01 M Na4EDTA solution (solution C) was prepared by dissolving 1.900 g of Na4EDTA in 500 mL DI water. The ammonium molybdate solution (solution D, kept in the refrigerator) used in the silicomolybdate test was prepared by dissolving 10 g (NH4)6Mo7O24·4H2O in 100 mL DI water, to which 24 pellets of solid NaOH were added under stirring, followed by pH adjustment to 7.7-7.8. 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 DI water. A 1 % w/v Eriochrome Black T solution (solution G) was prepared by dissolving 0.30 g in 30 mL ethanol. Stock solutions of 10,000 ppm (as actives) of each phosphonate were also prepared (solutions H). Precipitation Protocols. The protocols followed herein have been reported by us in detail elsewhere.23 All experiments were carried out either at ambient temperature, or under hydrothermal conditions (200 °C, under autogenous pressure). Precipitation reactions were carried out at pH 10.0. Precipitate characterization was carried out on precipitates that were isolated at room temperature by simple filtration through 0.45 µm membrane filters. Polyethylene (PET) containers were exclusively ACS Paragon Plus Environment

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used to avoid silicate leaching from the glass. Molybdate-reactive silicic acid was measured using the silicomolybdate spectrophotometric method, which has a ± 5 % accuracy. The formation of molybdosilicic acids from mixed solutions of molybdate and silicate.37-39 Solution pH was checked after every sampling, but no major fluctuations were measured. In general, pH adjustment was made only when solution pH was off by more than 0.2 pH units. Below, we describe in detail the protocol followed. Protocols for precipitate formation in the presence of Mg2+ and silicic acid.

A. Ambient

temperature. A quantity of solution A (10 mL) was placed in a PET container. The pH of this solution was initially ~ 12.0, hence, to avoid magnesium hydroxide precipitation adjustment to pH < 9.5 was necessary. The desired volumes of solutions H and then B were placed in the container and the desired pH value of 10.0 was achieved by addition of HCl. Further pH adjustment was done by a solution of NaOH, if needed. The necessary amount of the 10,000 ppm in Mg2+ (solution B) was used (0.50 mL), such that the final Mg2+ concentration is 200 ppm (as Mg). The volume of the combined solution A+B is made up to 25 mL. The final working solution contains 200 ppm silicic acid (as SiO2) and 200 ppm of Mg2+(as Mg). Finally, the container was covered with plastic membrane and set aside without stirring. The solutions were checked for soluble silicic acid by the silicomolybdate yellow method (see below) every half hour for the first hour and every hour for the next three hours after the final pH adjustment (t = 0). In all results presented below the concentration of molybdate-reactive silica is expressed as ppm (as SiO2). B. High temperature. For the high temperature tests, the solutions are prepared according to the room temperature protocol. A standard volume of 3.5 mL of these solutions is placed in the sealed custommade bomb (see above), and heated in a temperature-controlled furnace at the desired temperature of 200 °C for 2 hours. The apparatus is then removed from the furnace and left 30 minutes to cool down. Then the device is opened and the samples are tested with the same determination procedures as those followed for the room temperature experiments (Molybdate Yellow Method and Magnesium Titration Method). ACS Paragon Plus Environment

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Quantification of “soluble (reactive) silica”. “Soluble or reactive silica” actually denotes soluble silicic acid and was measured using the “yellow” silicomolybdate spectrophotometric method. According to this method 2 mL filtered working solution through a 0.45 µm syringe filter, is diluted to 25 mL in the cell, with light path 1 cm. 1 mL of solution D and 0.5 mL of solution E are added to the sample cell, the solution is mixed well and left undisturbed for 10 min. Then 1 mL of solution F is added and mixed again. The solution is set aside for 2 min. After the second time period the photometer is set to zero absorbance with DI water. Finally, the sample absorbance is measured at 452 nm as “ppm soluble silica”. The detectable concentrations range is 0–75.0 ppm.

In order to calculate the

concentration in the original solution a dilution factor is 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 is 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 in this term includes molybdate-reactive species. “Molybdate-reactive” includes all three species, namely, monomeric (monosilicic acid, H4SiO4), dimeric (disilicic acid, H6Si2O7), and possibly trimeric (H8Si3O10) silica.40 However, since we did not detect trimeric silica in any of our present and past experiments, the term “molybdate-reactive” silica refers to a mixture of monosilicic acid (primarily) and disilicic acid, with the first being dominant under our experimental conditions.41 More specifically, it was discovered that the monomeric state will react with molybdic acid within the first 75 seconds of intact at 20 oC, disilicic acid needs about 10 minutes, while other oligomers require even longer times. It should be noted that Mg2+ does not interfere with the silicomolybdate spectrophotometric method. Quantification of magnesium. The Mg2+ concentration in the working solutions was determined by the well-known EDTA Titrimetric method.42 Hence, 1 mL sample is diluted to 10 ml with deionized water. The pH is adjusted to ~11 with 1N NaOH and 3 drops of EBT solution (solution G) are added to the solution. Slow addition of EDTA titrant (solution C), with continuous stirring to the proper end point ACS Paragon Plus Environment

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and calculation of the magnesium concentration are the last steps of the procedure. In all results presented below the concentration of Mg2+ is expressed as ppm (as Mg). Protocol for the effect of inhibitors on precipitation. In general, the above procedure is followed, except that after the pH adjustment to lower values the required amount of inhibitors solution (such as concentrations of 50, 100, 200, 500, and 1500 ppm) is added before addition of solution B. After that the same procedure as above was followed.

Results

Silica and Mg2+ stabilization by phosphonates. In a recent paper on the precipitation of “magnesium silicate” (in the absence of additives) in geothermally-relevant artificial waters we reported that the inorganic precipitates formed under certain conditions (in the presence of variable amounts of Mg2+ and silicate anions at pH 10) can best be described as “magnesium-containing amorphous silica”, with variable Mg2+ content.23 It was also found that precipitate formation was pH-dependent (enhanced precipitation as pH increases). Furthermore, precipitation was more pronounced with increased silica and Mg2+ concentrations, all precipitates were amorphous, Mg elemental distribution within the isolated precipitates was random (indicating a non-stoichiometric process), and temperature-dependence of precipitation was small. In this work, experimental conditions were fixed to [Mg2+] = [SiO2(sol)] = 200 ppm ([Mg2+] = 8.23 mM, [SiO2(sol)] = 3.33. mM) at pH = 10.0 and a range of additive concentrations was tested, 50, 100, 200, 500, and 1500 ppm. Quantification of inhibitor efficiency was based on molybdate-reactive silica and Mg2+ measurements that remained in solution for 4 hours after the onset of the precipitation reaction. Figure 2 shows representative results for HDTMP as inhibitor. Similar data can be found in the SI for inhibitors AMP, PBTC, HEDP and BHMTPAMP, Figures S-1 to S-4.


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Energy & Fuels control silica control Mg silica Mg

Concentration (ppm)

150

100 ppm HDTMP 200

Concentration (ppm)

50 ppm HDTMP 200

100 50 0 0

1

2 Τime (hours)

3

150 100 50 0

4

0

1

2 Τime (hours)

3

4

500 ppm HDTMP

200 ppm HDTMP 200

200

Concentration (ppm)

Concentration (ppm)

150 100 50

150 100 50 0

0 0

1

2 Τime (hours)

3

0

4

1

2

3

4

Τime (hours)

1500 ppm HDTMP 200

Concentration (ppm)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

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150 100 50 0 0

1

2 Τime (hours)

3

4

Figure 2. Stabilization of silica and Mg2+ in the presence of various concentrations of HDTMP.

By examining the silica and Mg2+ loss in the absence of inhibitors (dotted lines), it appears that silica reduction is more pronounced than Mg2+ reduction. This is consistent with the previous observation that very little Mg2+ is incorporated in the inorganic precipitate, in Mg:Si molar ratios ranging from 0.19:1 to 0.88:1, ie. in sub-stoichiometric amounts. Even at low HDTMP concentration (50 ppm) stabilization is evident. Specifically, 45 ppm of silica and 48 ppm Mg2+ are stabilized above the control values (after the 4th hour). Silica (∆Silica) and Mg2+ (∆Mg) splits gradually increase, as inhibitor concentration increases. For example, at 1500 ppm HDTMP, 83 ppm of silica and 51 ppm Mg2+ are stabilized above the control values (after the 4th hour). The term “split” is defined as the actual measurement of either


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silica or Mg2+ (in ppm) after subtraction of the “control” value. Table 1 summarizes the silica and Mg2+ splits for all inhibitors for all concentrations.

Table 1. Silica (∆Silica) and Mg2+ (∆Mg) splits and % stabilization for all inhibitors, after 4 hours of precipitation time.a Inhibitor concentration (ppm)

50 100 200 500 1500

∆Silica, ppm (% stabilization)b PBTC

HEDP

AMP

HDTMP

BHMTPAMP

7 (4 %) 13 (10 %) 11 (7 %) 116 (73 %) 141 (86 %)

33 (22 %) 35 (24 %) 35 (21 %) 69 (45 %) 125 (76 %)

15 (12 %) 18 (13 %) 21 (15 %) 84 (54 %) 117 (71 %)

45 (27 %) 50 (33 %) 61 (37 %) 64 (41 %) 83 (51 %)

17 (10 %) 22 (16 %) 26 (16 %) 72 (46 %) 74 (45 %)

∆Mg, ppm (% stabilization) c

50 100 200 500 1500

PBTC

HEDP

AMP

HDTMP

BHMTPAMP

48 (47 %) 51 (50 %) 49 (48 %) 96 (93 %) 99 (96 %)

65 (63 %) 75 (73 %) 73 (71 %) 45 (44 %) 48 (47 %)

45 (44 %) 65 (63 %) 70 (68 %) 45 (44 %) 64 (62 %)

48 (47 %) 48 (47 %) 49 (48 %) 52 (50 %) 51 (50 %)

44 (43 %) 83 (81 %) 73 (71 %) 99 (96 %) 99 (96 %)

a

The term “split” is defined as the actual measurement of either silica or Mg (in ppm) after subtraction of the “control” value.

b

% silica stabilization is given by the equation: % stabilization =

c

% Mg2+ stabilization is given by the equation: % stabilization =

Graphical representations of silica (∆Silica) and Mg2+ (∆Mg) splits for all phosphonate additives are given in Figures S-5 and S-6 in the SI. Figure 3 presents selected % stabilization data for silica and Mg2+ for all inhibitors at the lowest (50 ppm) and highest (1500 ppm) concentration (4th hour data). All inhibitors are partially effective in controlling soluble silica at the 50 ppm dosage. The least effective inhibitor is PBTC (4 % stabilization) and the highest % stabilization is achieved by HDTMP (27 %). All inhibitors exhibit variable stabilization at the highest dosage of 1500 ppm. The least effective inhibitor ACS Paragon Plus Environment

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is BHMTPAMP (45 %) and the most effective one is PBTC (86 %). It appears that these phosphonate inhibitors are much more effective in Mg2+ stabilization. At the lowest dosage of 50 ppm, the lowest stabilization is exhibited by BHMTPAMP (43 %) and the highest by HEDP (63 %). Dosage increase to 1500 ppm has a profound effect in stabilization efficiency. The highest, and nearly quantitative stabilization is found for both PBTC and BHMTPAMP (96 %), whereas the lowest for HEDP (47 %). Dosage increase enhances stabilization, except for Mg2+ stabilization by HEDP, for which a drop from 63 % to 47 % is observed, in spite of increase in inhibitor dosage. This phenomenon may be an indication of Mg-phosphonate precipitate formation, and will be discussed extensively later. 100

% silica stabilization % Mg stabilization

90 80 70

stabilization (%)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

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60 50 40 30 20 10 0 PBTC

HEDP

AMP

HDTMP

BHMTPAMP

phosphonate inhibitor (at 50 ppm)

Figure 3. Stabilization (%) of all inhibitors at the lowest (50 ppm, left) and highest (1500 ppm, right) concentration after 4 hours of reaction time at ambient temperature. Arrows (right graph) indicate increase or decrease in silica or Mg2+ concentration upon inhibitor dosage increase from 50 ppm to 1500 ppm.

Temperature effects. Silica and Mg2+ stabilization was also studied at the temperature of 200 °C in pressurized vessels, by a methodology described before.23 Results for a dosages 50 and 1500 ppm are presented in Figure 4. Regarding the 50 ppm dosage silica stabilization does not seem to be affected by temperature (compare results with those in Figure 3, left). In contrast, Mg2+ stabilization is substantially lower at 200 °C, with the most dramatic reduction observed for HEDP and AMP. At the high inhibitor dosage of 1500 ppm silica stabilization presents small differences between the two temperatures. The


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only exception is the case of PBTC, where silica stabilization drops from 86 % to 10 %. Differences in the Mg2+ stabilization are observed only for HEDP, where an increase is observed (from 47 % to 95 %). 100

100


90 80

80

70

70

60 50 40 30 20 10


90

stabilization (%)

stabilization (%)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

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60 50 40 30 20 10

0

0

PBTC

HEDP

AMP

HDTMP

BHMTPAMP

PBTC


HEDP

AMP

HDTMP

BHMTPAMP


Figure 4. Stabilization (%) of silica and Mg2+ for all inhibitors at 50 ppm (left) and 1500 ppm (right) dosages after 2.5 hours of reaction time at 200 °C. The symbol “×” denotes 0 % stabilization.

Precipitate characterization. Powder X-Ray Diffraction. Precipitates in the presence of 1500 ppm inhibitor were isolated after 4 hours of precipitation time at room temperature. They are all amorphous as indicated by powder XRD studies, see Figure 5. The broad peak around 27° corresponds to amorphous silica. No indication for the presence of crystalline magnesium silicates could be derived from the diffraction patterns, in agreement with our previous observations.23


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Figure 5. Selected powder XRD diagrams of precipitates isolated at ambient temperature after 3 hours from a pH = 10 solution containing 200 ppm SiO2, 200 ppm Mg2+ and 1500 ppm of phosphonate inhibitor. Representative diagrams are shown for AMP (red), BHMTPAMP (blue), and HDTMP (green).

Vibrational spectroscopy. ATR-IR spectroscopy was also employed to further study the precipitates that formed in the presence of inhibitors at room temperature. Comparative ATR-IR spectra are presented in Figure 6. All spectra are essentially the same, indicating the presence of Mg2+-loaded amorphous silica.

95

Reflectance (%)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

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AMP BHMTPAMP HDTMP HEDP PBTC no inhibitor

80

65

50

35 4000

3400

2800 2200 1600 wavenumbers (cm-1)

1000

400

Figure 6. ATR-IR spectra of precipitates formed without (control) and in the presence of inhibitors. Precipitates were isolated at ambient temperature after 4 hours from a pH = 10 solution containing 200 ppm silica, 200 ppm Mg2+ and 1500 ppm inhibitor.

Peaks appear at ~ 3325 cm-1 (O-H stretching vibration, very broad) and ~ 1640 cm-1 (H-O-H bending vibration, weak), due to the presence of water of hydration. The strong absorption at ~ 991 cm-1 is characteristic of Mg2+-containing amorphous silica. The bands at 887 and 610 cm-1 are assigned to the “Q2 and Q1 bending vibrations”, respectively, of the silica network. Finally, the broad band at ~ 450 cm-1 is due to the Mg-O bonds. The noise around 2200 cm-1 originates from background subtraction. ACS Paragon Plus Environment

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Scanning Electron Microscopy/Electron Dispersive Spectrometry. Precipitates isolated at room temperature were examined by SEM for morphological features, particle size and agglomeration tendency. Some representative images are shown in Figure 7. Particle morphology is similar in all samples studied displaying the common feature of particle aggregation. This phenomenon seems to be enhanced in precipitates originated from phosphonate-containing solutions, compared to precipitates in the absence of inhibitor (Figure 7, image lower right). It is likely that the phosphonate inhibitor (at least partially) is entrapped within the precipitate matrix. EDS has provided proof for inhibitor entrapment, based on the detection of P, which could only originate from the phosphonates. Precipitate samples that were isolated in the presence of any of the inhibitors showed variable, but low Mg:P atomic ratios, close to the value of 12. It is interesting to note that the presence of phosphonates was not detectable by ATRIR (Figure 6).

Figure 7. Precipitate morphology in the presence of phosphonate inhibitors (all at 1500 ppm), as indicated. The precipitates were isolated after 3 hours of precipitation time at pH = 10.0.

HDTMP was selected for studying the effect of inhibitor concentration on particle morphology. Images are shown in Figure 8 for HDTMP dosages 50, 100, 200, and 500 ppm (the image corresponding ACS Paragon Plus Environment

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to 1500 ppm is shown in Figure 7, bottom left). The same non-descript particle morphology was observed here as well. For some dosages (200 and 500 ppm) larger particles were noted, but these are not representative of the entire sample.

Figure 8. Precipitate morphology in the presence of variable concentrations of HDTMP. The precipitates were isolated after 3 hours of precipitation time at pH = 10.0.

An additional issue to be addressed was the dependence of particle size/morphology evolution on time. Hence, precipitates from HDTMP-containing (1500 ppm) working solutions were isolated after 0.5, 1.0 and 2.0 hours, and were studied by SEM. Images are shown in Figure 9 (the image corresponding to the 3-hour sample is shown in Figure 7, bottom left). Interestingly, no differences were evident pointing to the conclusion that particles appearing after 0.5 hours do not grow in size, neither their morphology is altered.

Figure 9. Time-dependence of precipitate morphology in the presence of HDTMP (1500 ppm).


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Study of the Mg:Si atomic ratio in the isolated precipitates can also reveal important characteristics. Below, Mg:Si atomic ratio in precipitates isolated in the presence of HDTMP will be discussed. Figure 10 shows that the Mg:Si atomic ratio remains constant for the duration of the experiment (3 hours). Hence, it appears that no transformation or composition change takes place. The observed Mg:Si atomic ratio (~ 0.15) is much lower than that observed in well-characterized magnesium silicates (see upper part of Figure 10). It should also be noted that the “theoretical” Mg:Si ratio (based on the salts added at the onset of precipitation) is ~ 2.5 (excess Mg2+). The low Mg2+ content in the isolated precipitates is another indication of Mg2+ stabilization by phosphonates in solution. 1.00

enstantite (Mg:Si = 1)

0.90 0.80 Mg:Si molar ratio

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

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talc (Mg:Si = 0.75)

0.70

sepiolite (Mg:Si = 0.67)

0.60 0.50 0.40 0.30 0.20 0.10 0.00 0

0.5

1

1.5 2 time (hours)

2.5

3

3.5

Figure 10. Time-dependence of Mg:Si atomic ratio in precipitates isolated in the presence of HDTMP (1500 ppm) after 0.5, 1, 2, and 3 hours.

Interestingly, HDTMP dosage increase induces a systematic reduction in Mg2+ content in the isolated precipitates (Figure 11). The Mg:Si atomic ratio drops gradually from 0.72 to 0.20 upon HDTMP dosage increase. The Mg:Si ratio is close to that observed in sepiolite (0.67) for the low phosphonate dosages 50 and 100 ppm, but there is no definitive proof that sepiolite actually forms. The drop in Mg:Si ratio in the precipitates upon HDTMP dosage increase supports the hypothesis that Mg2+ is stabilized by the phosphonate by generation of water soluble Mg-HDTMP “complexes”. There is an apparent contradiction between the results presented in Figure 11 and those depicted in Figure S-5 of the SI. In Figure S-5 Mg2+ and silica splits are plotted vs. HDTMP dosage, and while the Mg2+ split seems quite ACS Paragon Plus Environment

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constant, the silica split increases. We assign this discrepancy to the fact that the silicomolybdate photometric test measures only monomeric and dimeric silicic acids. There may be other silica species which are unreactive to molybdate, but still part of the solution and not the precipitate. 1.00

enstantite (Mg:Si = 1) Mg:Si range in the control (no inhibitor)

0.90 0.80 Mg:Si molar ratio

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

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0.70 0.60 0.50

talc (Mg:Si = 0.75) sepiolite (Mg:Si = 0.67)

0.40 0.30 0.20 0.10 0.00 0

200

400 600 800 1000 1200 HDTMP concentration (ppm)

1400

1600

Figure 11. Influence of HDTMP dosage on the measured Mg:Si atomic ratio in precipitates isolated after 3 hours. The dotted line has been added to aid the reader.

Discussion Previously, we reported that Mg2+ ions actually catalyze silicic acid polycondensation generating amorphous silica precipitates that contain variable amounts of Mg2+ adsorbed/embedded in the silica matrix.23 This catalytic effect of Mg2+ ions can be turned-off by addition of EDTA.13 The latter acts as a chelating agent for Mg2+ due to its high affinity for alkaline-earth metal ions.43 We propose that a similar effect is exerted by the phosphonate additives on the stabilization of Mg2+ in the system studied herein. The initial strategy to approach this issue was to study solutions containing Mg2+ (200 ppm) and phosphonate (1500 ppm) at pH = 10.0, in the absence of silica. All phosphonate additives are fully deprotonated at such high pH.44 Small amounts of Mg-phosphonate precipitates formed and were


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isolated after 3 hours. In some cases, e.g. BHMTPAMP, the precipitation reaction had to be scaled up because of the extremely small amount of precipitate formed. It is expected that the various phosphonate additives create Mg-phosphonate precipitates of varying solubility. Hence, the “leftover” Mg2+ content of each solution was measured and the results are presented in Figure 12 (upper). Evidently, most of the Mg2+ remains in solution, most likely complexed by the phosphonate (see comments on NMR results below), and very little ends-up in the precipitate. This is in agreement with the EDS results (Mg:P ratio ~ 12) on precipitates formed in the presence of Mg2+, silica and phosphonate. For comparison the % Mg2+ stabilization data are presented (taken from Table 1 for the 1500 ppm inhibitor dosage). The similarities in trends are obvious, pointing to the conclusion that the higher the ability of an inhibitor is to maintain high Mg2+ in solution, the higher its effectiveness is to maintain high levels of Mg2+ in the presence of silica.

soluble Mg2+ (%)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

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100 90 80 70 60 50 40 30 20 10 0

phosphonate additive


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Mg2+ stabilization (%)

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100 90 80 70 60 50 40 30 20 10 0

phosphonate additive

Figure 12. Upper: Percent Mg2+ in solution after the “Mg2+ + phosphonate” precipitation reaction. Conditions: Mg2+ 200 ppm, phosphonate 1500 ppm, pH = 10.0. The “control” solution contains no additives. Lower: Percent Mg2+ stabilization in the presence of phosphonate. These data were taken from Table 1. Conditions: Silica 200 ppm, Mg2+ 200 ppm, phosphonate 1500 ppm, pH = 10.0.

The Mg-phosphonates were studied by powder XRD (see Figure S-7 in the SI), ATR-IR (see Figure S-8 in the SI), and elemental analysis (EDS). All isolated precipitates are amorphous. Their vibrational spectra showed the expected bands in the region 900-1100 cm-1 attributed to the phosphonate moiety.4547

Mg:P molar ratios were: 1:1 (AMP, theoretical 3:3), 3:1 (PBTC, theoretical 1:1), 1:1 (BHMTPAMP,

theoretical 5:5), 1:1 (HDTMP, theoretical 4:4) and 1:1 (HEDP, theoretical 2:2). These values are expected if it is assumed that each fully deprotonated phosphonate moiety (-PO32-) can be chargedbalanced by one Mg2+ ion. Elemental mapping (Si, Mg, and P) was performed on precipitates isolated in the absence (control) and in the presence of HDTMP (1500 ppm). These images are provided as Figures S-9 (A, B, C, D) and confirm the uneven distribution of Mg and P in the isolated precipitates. 31

P solution NMR provides direct evidence of Mg2+ coordination by phosphonates. Again, HDTMP

was used as the example. The 31P solution NMR of a solution of HDTMP (D2O, 1500 ppm, pH adjusted to 10.0 with ethylenediamine) showed a single peak at 6.86 ppm (see Figure S-10-upper in the SI). In ACS Paragon Plus Environment

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the presence of Mg2+ (same conditions as before, 200 ppm Mg2+) this peak substantially broadens and shifts to 7.46 ppm (see Figure S-10-lower in the SI). The same experiments were repeated with NaOH as the pH-adjusting reagent. The single peak in the absence of Mg2+ appeared at 6.40 ppm (see Figure S11-upper in the SI), and at 8.06 ppm in the presence of Mg2+ (see Figure S-11-lower in the SI). Similar shifts of +0.60 ppm and +0.66 ppm, respectively, and similar line-broadening were noted. Our findings are in concert with previously published work on the interaction of DTPMPA [diethylenetriamineN,N,N′,N′′,N′′-pentakis(methylenephosphonic) acid] with selected metal ions (namely Sn2+, Zn2+, Cu2+, Fe2+,

Fe3+

and

Al3+),48

and

on

the

interaction

of

α,ω-alkylenediamine-N,N,N′,N′-

tetramethylenetetraphosphonates with cobalt complexes.49 Based on the measurements of stabilized Mg2+ in solution (Figure 2, Table 1, and Figures S-1 to S-4 in the SI) the number of Mg2+ ions stabilized by one molecule of inhibitor were calculated and compared to the “theoretical” number of Mg2+ ions. For this calculation the following hypotheses were taken into account: (a) each fully deprotonated phosphonate moiety (-PO32-) can be charged-balanced by one Mg2+ ion. Hence, one molecule of AMP6- stabilizes three Mg2+ ions, one molecule of BHMTPAMP10- stabilizes five Mg2+ ions, etc. (b) All phosphonate molecules are complexed by Mg2+, and no “free” inhibitor exists in solution. (c) Inhibitor loss due to precipitation (as Mg-phosphonate precipitates) is not taken into account. (d) the Mg:P ratio is the same in solution as it is in the precipitates. The outcome of this comparison is shown in Figure 13.


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Figure 13. Number of Mg2+ ions stabilized by one molecule of phosphonate inhibitor (blue bars) vs. the “theoretical” number of Mg2+ ions (red bars).

It is apparent that in some cases (HEDP, AMP, PBTC) the experimentally determined number of Mg2+ ions is higher than the “theoretical” one. This phenomenon could be rationalized based on the wellknown ability of the phosphonate group to bind to more than one Mg2+ ions (and several polyvalent metal ions as well), or the presence of other binding groups (such as –OH in HEDP). It is obvious that such binding influences phosphonate precipitation as well. The bridging capability of the phosphonate group has been well established in metal phosphonate chemistry.46,50-53 In the case of HDTMP the experimental number of Mg2+ ions is lower than the “theoretical” value. We attribute this observation to precipitation of a Mg-HDTMP “complex” out of solution. This is in agreement with the fact that ~ 20 % of Mg2+ is lost to precipitation in the presence of HDTMP, Figure 12. An additional important issue is the correlation between the ability of each phosphonate to stabilize Mg2+ ions and its affinity for Mg2+ ions in solution. Metal-ligand affinity is quantified by the stability constant. The aforementioned correlation is displayed in Figure 14. Mg-phosphonate stability constants were taken from the literature for HEDP,54 HDTMP,55 AMP,56 and diethylenetriamineN,N,N′,N″,N″-pentakis(methylenephosphonic acid) (a similar pentaphosphonate to BHMTPAMP).57 No data are available for PBTC. It is clearly seen that the higher the Mg-phosphonate stability constant is, ACS Paragon Plus Environment

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the higher the number of Mg2+ ions stabilized by one molecule of phosphonate inhibitor. This is a reflection of the affinity of each phosphonate additive for Mg2+, and hence the stability of the produced complex in solution. 6.0

4.0

HDTMP

3.0

BHMTPAMP

AMP

5.0 HEDP

number of Mg ions stabilized

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

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2.0 1.0 0.0 5

6

7

8

9 10 11 stability constant

12

13

14

Figure 14. Correlation between the number of Mg2+ ions stabilized by one molecule of phosphonate inhibitor and stability constants. Red circles: experimental, Blue squares: theoretical. The lines have been added to aid the reader.

It has been established that the phosphonate group58 (in additives HEDP, AMP, HDTMP, PBTC, and BHMTPAMP) and the carboxylate group59 (in the PBTC additive) can bind metal ions in a plethora of modes, Figure 15. Depending on the protonation state, the phosphonate moiety can bind one to nine metal ions, whereas the carboxylate moiety can coordinate one to four metal centers in a variety of structural motifs, from monodentate to polydentate, chelating and bridging. Although it is a difficult challenge to discern which mode(s) the phosphonate additives bind Mg2+ ions in solution, it reasonable to assume that all aforementioned binding modes are possible. Furthermore, it appears that the higher number of binding moieties present on the inhibitor backbone, the more effective Mg2+ stabilizer that inhibitor is. For example, in Figure 13 HEDP (with two phosphonate and a hydroxyl groups) stabilizes three Mg2+ ions, whereas BHMTPAMP (with five phosphonate groups) stabilizes five Mg2+ ions.


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Page 24 of 32 M O

M O

O

M

M

O

O

C

C

C

R

R

R

M

M

M O

M O

O C

O M

R

O C

O

R

M M

M

O

O

O

M

R

M O

M

M

O

O

C

C

R

R

M M

O C

M

M

M

O

O

C

C

R

R

M

Figure 15. Metal binding modes of the phosphonate (left) and the carboxylate (right) group. In this case M represents a Mg2+ ion.

Conclusions

Herein, we implemented an initial systematic attempt to study the effect of a variety of antiscalant phosphonate additives on the inhibition of silica formation in the presence Mg2+ ions. The conclusions derived are: (1) All phosphonate additives demonstrate variable stabilization effectiveness for Mg2+ and molybdate-reactive silica (silicic acid). (2) The inhibition efficiency of the phosphonates, quantified by ∆Silica (% silica stabilization), and ∆Mg (% Mg2+ stabilization) (see Table 1) generally increases with inhibitor concentration. However, in the case of % Mg2+ stabilization no such enhancement is observed beyond the 500 ppm dosage (Figure S-6 in the SI). (3) The phosphonate additives act as chelators/ligands for Mg2+ ions, and thus they turn off its catalytic effect in silicic acid polycondensation. However, this may not be the only effect present. As observed before,13 EDTA has a similar “turn-off” effects for Mg2+ catalysis in silica polycodensation. Interestingly, a silica inhibitory effect was observed for Mg-EDTA complexes. ACS Paragon Plus Environment

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This invokes the possibility that in the present work Mg-phosphonate complexes may act as silica inhibitors as well, with variable efficiency. For example, PBTC shows a 96 % Mg2+ stabilization and a 86 % silica inhibition at 1500 ppm (see Table 1). BHMTPAMP, on the other hand shows a 96 % Mg2+ stabilization but only a 45 % silica inhibition at 1500 ppm (Table 1). Possibly, soluble Mg-silicate and/or Mg-phosphonate-silicate complexes may be involved. PBTC (with 4 binding sites, 3 carboxylates, one phosphonate) may allow silicate binding to Mg2+, whereas the more bulky BHMTPAMP (5 phosphonate binding sites), may hinder silicate binding to Mg2+. Silicate binding to Mg2+ has been invoked before.13 (4) Mg2+ ion stabilization is dependent on the Mg-phosphonate stability constants, based on available data. Additives with high stability constants seem to be the most efficient Mg2+ stabilizers. (5) The silica precipitates that form in the presence of Mg2+ ions are Mg2+-absorbed amorphous silica, with variable Mg2+ content. (6) Limited phosphonate entrapment into the precipitates occurs due to formation of sparingly-soluble Mg-phosphonate “complexes” that are found in the precipitated solid (based on the P analyses).

Acknowledgment. K.D.D. and A.S. thank Kurita Europe GmbH for financial support of the research project GEOSCALE (KA 4521). Author Contributions. The manuscript was written through contributions of all authors. All authors have given approval to the final version of the manuscript. Supporting Information. Figures with stabilization of silica and Mg2+ in the presence of various concentrations of phosphonate additives, silica (∆Silica) and Mg (∆Mg) splits for all inhibitors, powder XRD diagrams and ATR-IR spectra of Mg-phosphonate precipitates, and selected 31P NMR spectra (for the Mg2+ + HDTMP system). ACS Paragon Plus Environment

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TOC image and text Phosphonate additives can control the precipitation of “magnesium silicate” by complexing Mg2+, thus “turning-off” its catalytic effect in silica polycondensation.


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Antiscalant-Driven Inhibition and Stabilization of âMagnesium

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