CO2 Mineralization Using Brine Discharged from a Seawater Desalination Plant

Article CO2 Mineralization Using Brine Discharged from a Seawater Desalination Plant

Jun-Hwan Bang, Yeongsuk Yoo, Seung-Woo Lee, Kyungsun Song and Soochun Chae *

Change Mitigation and Sustainability Division, Korea Institute of Geoscience and Mineral Resources (KIGAM), 124 Gwahang-no, Yuseong-gu, Daejeon 34132, Korea; [email protected] (J.-H.B.); [email protected] (Y.Y.); [email protected] (S.-W.L.); [email protected] (K.S.) * Correspondence: [email protected]

Received: 12 September 2017; Accepted: 27 October 2017; Published: 30 October 2017

Abstract: CO2 mineralization is a method of sequestering CO2 in the form of carbonated minerals. Brine discharged from seawater desalination is a potential source of Mg and Ca, which can precipitate CO2 as forms of their carbonate minerals. The concentration of Mg and Ca in brine are twice those in the seawater influent to desalination process. This study used a cycle for CO2 mineralization that involves an increase in the pH of the brine, followed by CO2 bubbling, and, finally, filtration. To the best of our knowledge, this is the first time that non-synthesized brine from a seawater desalination plant has been used for CO2 mineralization. The resulting precipitates were CaCO3 (calcite), Mg5(CO3)4(OH)2·4H2O (hydromagnesite), and NaCl (halite) with these materials being identified by X-ray Diffraction (XRD), Fourier transform infrared (FTIR) and thermo gravimetric-differentail thermal Analysis (TGA)-DTA. Despite the presence of Ca with Mg in brine being unfavorable for the precipitation of Mg carbonate, Mg reacted with CO2 to form hydromagnesite at a yield of 86%. Most of the Ca formed calcite, at 99% yield. This study empirically demonstrates that brine from seawater desalination plants can be used for CO2 mineralization.

Keywords: CO2 mineralization; desalination; carbonates; brine; seawater

1. Introduction Climate change and abnormal weather phenomena have been worsening, owing to increased CO2 emissions around the world. Typically, the Korean peninsula is not a region that is known to suffer from water shortages [1]. However, since in the last few years there have been severe annual droughts, especially in west-coast areas, with increased CO2 emissions being blamed. Several methods of preserving water resources have been examined, and seawater desalination has been found to be a successful candidate [2,3]. Seawater desalination processes remove the salts from seawater by distillation, membrane separation, and/or reverse osmosis (RO) to produce useful water that is suitable for drinking and industrial purposes [3–5]. One result of these processes is that concentrated seawater (i.e., brine) is also produced, in which the rejected salts in seawater have been concentrated. Desalination consumes large amounts of energy, meaning that large-scale desalination plants are often placed near power stations [2], and also leads to the emission of massive amounts of CO2, which will eventually exacerbate the problems that desalination is attempting to solve. Methods of limiting CO2 emissions or removing CO2 from desalination plants are, therefore, of great interest. CO2 mineralization is one such method of sequestering CO2. In this method, CO2 reacts with calcium and/or magnesium to form calcium carbonate and/or magnesium carbonate [6]. The rate of CO2 mineralization can be accelerated when the reaction takes place in the aqueous phase, as happens 2− in typical chemical reactions. Ionization of CO2 to CO3 takes place under alkaline conditions. This is 2− 2+ 2+ essential for CO2 mineralization, as it allows ionic reactions between CO3 and Ca and/or Mg .

Minerals 2017, 7, 207; doi:10.3390/min7110207 www.mdpi.com/journal/minerals Minerals 2017, 7, 207 2 of 12

For this reason, alkaline chemicals such as NaOH or NH4OH are applied to increase the pH of the aqueous phase in CO2 mineralization [7,8]. Meanwhile, using Ca and Mg from industrial waste such as waste concrete, slags from steel production, and liquid wastes is preferable when attempting to design feasible CO2 mineralization processes [9–13]. Detailed principles and applications of CO2 mineralization using alkaline waste have been summarized by Pan et al. [14]. Some pretreatment steps, such as crushing, sieving, activation, and/or digestion are required to extract Ca and Mg from solid waste [13,15,16]. When using liquid waste, such pretreatment procedures are not required because the Ca and/or Mg are already present as ions. For this reason, the use of brine for CO2 mineralization is a promising route for investigation [6,8,17]. In the implementation of this technology, both engineering and economic factors are important, as is environmentally balanced performance [18]. It is reported that the concentrations of the rejected salts in brine from desalination plants using reverse osmosis (RO) systems are 2–2.5 times higher than those in influent seawater [19], suggesting that brine could be a promising material for CO2 mineralization. If Mg alone is present in a solution, it readily reacts with CO2 to precipitate Mg carbonates; the reaction between MgCl2·6H2O solution and CO2 produced nesquehonite (MgCO3·3H2O), dypingite (Mg5(CO3)4(OH)2·5H2O), dypingite-like (Mg5(CO3)4(OH)2·8H2O), and hydromagnesite (Mg5(CO3)4(OH)2·4H2O) crystal phases, depending on the reaction conditions [20]. When considering brine for CO2 mineralization, it is worth noting that in previously reported studies, materials in a replica brine solution with coexisting Mg and Ca were not readily carbonated. Liu and Maroto-Valer [21] synthesized oil-field brine and conducted CO2 mineralization experiments under mimicked underground conditions to precipitate calcite. An artificial + 2+ 2+ + 2+ − 2− brine containing metal ions such as Na , Mg , Ca ,K , and Sr , and anions including Cl , SO4 , − 3− − − Br , BO3 ,F , and HCO3 was unfavorable for the precipitation of Mg carbonates, but was favorable for CaCO3 precipitation [22]. In the same manner, simulated brine composed of NaCl, MgCl2, CaCl2, Na2SO4, and KCl precipitated CaCO3 by the primary separation of Mg(OH)2 from the Ca2+ solution [23]. Owing to the significantly higher concentration of Mg than Ca in seawater and brine, Mg would allow much more CO2 mineralization than Ca does. However, this does not necessarily mean that more precipitating Mg carbonates than Ca carbonates will be produced, because Ca carbonates −5 are more sparingly soluble than Mg carbonates (for example, Ksp (hydromagnesite) = 1.26 × 10 [24], −9 ◦ and Ksp (calcite) = 3.36 × 10 at 25 C each [25]). This paper introduces the first use of non-synthesized brine discharged from a functional seawater desalination plant for CO2 mineralization; we have particularly focused on the precipitation of Mg carbonates. Repeated cycles involved the pH adjustment of brine, and the injection of CO2 microbubbles showed that Mg carbonation could be achieved alongside Ca. Various analytical techniques were used for the identification of the precipitates, including X-ray diffraction (XRD) analysis. We have empirically proved that CO2 mineralization using brine for CO2 sequestration.

2. Experimental The brine used in this study was collected from an ongoing seawater RO process for industrial water supplementation in Korea. The major cations and anions in this brine and their concentrations as measured by inductively coupled plasma atomic emission spectroscopy (ICP-OES; Optima 8300, PerkinElmer, Shelton, CT, USA) are listed in Table1. The concentrations of Sr, Si, Li, Cu, and rare earth elements such as W, Mo, and Rb were also measured, but the values obtained were not significant. Anion concentrations were measured by ion chromatography (IC; Dionex ICS-5000+, Thermo Fisher 2− Scientific, Sunnyvale, CA, USA), and the SO4 concentration was measured using IC equipped with an Ag cartridge (Dionex OnGuard II Ag, Thermo Fisher Scientific). Minerals 2017, 7, 207 3 of 12

Table 1. Concentrations of the major cations and anions in the brine.

Ions Mg Ca Na K Cl SO4 Minerals 2017, 7, 207 3 of 12 Concentration (mg/L) 2226 714 17,987 858 32,220 4766 Table 1. Concentrations of the major cations and anions in the brine. NaOH solution (1 M) was prepared for use as an alkaline chemical by dilution of NaOH Ions Mg Ca Na K Cl SO4 (96%, Junsei Chemicals,ConcentrationTokyo, (mg/L) Japan ) with deionized 2226 714 water 17,987 (Milli-Q 858 Gradient-A10, 32,220 Millipore, 4766 Billerica, MA, USA). The pH of 0.25 L of brine was increased from an initial value of 8.2 to 10–13.5 (measured by Orion StarNaOH A215, solution Thermo (1 M) Scientific, was prepared Singapore) for use as by an adding alkaline thechemical 1 M by NaOH dilution solution of NaOH to (96%, the brine in a beakerJunsei Chemicals, with a lid Tokyo, at a Japan) rate ofwith 4 deionized mL/min water using (Milli-Q a peristaltic Gradient-A10, pump Millipore, (EMP-2000W, Billerica, EMS-tech, MA, GyeongGi,USA). The Korea) pH of at 0.25 ambient L of brine temperature was increased and from pressure. an initial The value brine of 8.2 was to stirred10–13.5 (measured while the by NaOH solutionOrion was Star added, A215, Thermo and as Scientific, the pH of Singapore) the brine by increased, adding the a 1 suspension M NaOH solution of white to the precipitates brine in a was formed.beaker Once with the a lid target at a rate pH of was 4 mL/min reached, using the a peristaltic suspension pump was (EMP-2000W, filtered with EMS-tech, a 0.45 µ GyeongGi,m pore nylon Korea) at ambient temperature and pressure. The brine was stirred while the NaOH solution was membrane (Whatman, Brentford, UK) to allow the concentrations of the metal ions listed in Table1 to added, and as the pH of the brine increased, a suspension of white precipitates was formed. Once be measured by ICP-OES. However, only the concentrations of Mg, Ca, and Na changed with pH and, the target pH was reached, the suspension was filtered with a 0.45 μm pore nylon membrane as such,(Whatman, only the Brentford, concentrations UK) to of allow these the three concentrations elements are of reportedthe metal in ions the listed results in section.Table 1 to be Inmeasured the cyclic by ICP-OES. CO2 mineralization However, only experiments, the concentrations as 1 L of of Mg, brine Ca, reachedand Na changed pH 10.5, with we pH reacted and, the suspensionas such, with only COthe concentrations2 (99%, Gasco) of instantly, these three and elements monitored are reported for further in the pHresults changes section. at atmospheric ◦ pressure.In A the jacket cyclic reactor CO2 mineralization was used for experiments, CO2 mineralization, as 1 L of brine with reached a coolant pHat 10.5, 1 Cwe being reacted circulated the betweensuspension the walls with by CO a2 (99%, chiller, Gasco) and instantly, a lid being and placedmonitored on for the further upper pH part changes of the at reactor.atmospheric As CO2 microbubblespressure. A were jacket injected reactor intowas used the suspensionfor CO2 mineralization, using a microbubble with a coolant generator, at 1 °C being the pHcirculated decreased between the walls by a chiller, and a lid being placed on the upper part of the reactor. As CO2 rapidly, as has already been shown in the literature [7,26,27]. CO2 injection was stopped once the microbubbles were injected into the suspension using a microbubble generator, the pH decreased pH of the brine had decreased to 8. This was because we found that we were unable to collect the rapidly, as has already been shown in the literature [7,26,27]. CO2 injection was stopped once the pH CO2-reactedof the brine precipitates had decreased when COto 8.2 injectionThis was wasbecaus stoppede we found at pH that 7, whichwe were was unable the value to collect we had the used in previousCO2-reacted studies precipitates [7,26]. The when final CO products,2 injection which was stopped consisted at pH of 7, precipitates which was andthe value transparent we had solutions,used werein obtained previous by studies filtering [7,26]. the brineThe final through products, a 0.45 whichµm nylon consisted filter. of The precipitates experimental and transparent cycle was then begunsolutions, again, with werethe obtained pH of by the filtering solution the beingbrine through increased a 0.45 to 10.5,μm nylon CO2 filter.bubbling The experimental taking place, cycle and the precipitatewas then formed begun being again, filtered with the off. pH This of the cycle solution was repeated being increased until no to precipitates 10.5, CO2 bubbling were obtained taking on additionplace, of and NaOH, the precipitate which occurred formed duringbeing filtered the 5th of cycle.f. This Thiscycleindicated was repeated that until there no were precipitates no divalent were obtained on addition of NaOH, which occurred during the 5th cycle. This indicated that there ions such as Mg or Ca that were able to form hydroxides; thus, CO2 bubbling did not lead to the formationwere no of particles.divalent ions Therefore, such as Mg in ouror Ca experiment, that were able four to cycles form hydroxides; were sufficient thus, toCO remove2 bubbling all divalentdid not lead to the formation of particles. Therefore, in our experiment, four cycles were sufficient to ions from the brine through the formation of particles. The experimental cycle is shown in Figure1. remove all divalent ions from the brine through the formation of particles. The experimental cycle is We decided not to consider the effect of dissolved CO2 from atmosphere during the experiments, shown in Figure 1. We decided not to consider the effect of dissolved CO2 from atmosphere during 2+ 2+ becausethe weexperiments, believed because that Mg we, believed which was that 3 Mg times2+, which more was abundant 3 times more than Caabundantin the than brine, Ca2+ wouldin the not be carbonatedbrine, would by not atmospheric be carbonated CO by2 under atmospheric the experimental CO2 under the conditions. experimental conditions.

FigureFigure 1. Experimental 1. Experimental cycles cycles showing showing increase increase in pH, COin 2pH,injection, CO2 injection, and filtration and afterfiltration CO2 mineralization.after CO2 mineralization.

Minerals 2017, 7, 207 4 of 12

AfterMinerals the 2017 reaction, 7, 207 with CO2, the suspension was filtered, the concentrations of Mg, Ca, and4 of 12 Na in the filtered liquid samples were analyzed, and the filter cake was dried in an oven at 80 ◦C for 24 h, then characterizedAfter the reaction using XRDwith CO (X’pert2, the MPD,suspension Philips, was EAfiltered, Almelo, the concentrations The Netherlands), of Mg, aCa, field and emission Na in the filtered liquid samples were analyzed, and the filter cake was dried in an oven at 80 °C for 24 h, scanning electron microscope (FESEM; JSM-7800 FPRIME, JEOL, Tokyo, Japan) equipped with an then characterized using XRD (X’pert MPD, Philips, EA Almelo, The Netherlands), a field emission energy dispersive spectrometer (EDS; X-MaxN, Oxford), thermogravimetric analysis (TGA; DTG-60H, scanning electron microscope (FESEM; JSM-7800 FPRIME,◦ JEOL, Tokyo, Japan) equipped with an Shimadzu,energy dispersive Kyoto, Japan) spectrometer at argon (EDS; environment X-MaxN, Oxfo (10rd), C/minthermogravimetric heating rate), analysis and (TGA; Fourier DTG-60H, transform infraredShimadzu, spectroscopy Kyoto, (FTIR;Japan) NICOLETat argon environment 380, Thermo (10 Fisher °C/min Scientific heating Inc.,rate), Costaand Fourier Mesa, transform CA, USA). infrared spectroscopy (FTIR; NICOLET 380, Thermo Fisher Scientific Inc., Costa Mesa, CA, USA). 3. Results and discussion 3. Results and discussion 3.1. Concentrations of Mg, Ca, and Na in Brine with Different pH Values

Initial3.1. Concentrations experiments of Mg, aimed Ca, and to Na obtain in Brine information with Different on pH the Values changes in the concentrations of Mg and Ca asInitial the experiments pH was varied aimed from to obtain 10–13.5 information by adding on the NaOH changes solution in the concentrations to brine. This of Mg resulted and in the formationCa as the pH of awas white varied suspension. from 10–13.5 The by suspension adding NaOH thickened solution as to the brine. amount This resulted of NaOH in solutionthe addedformation increased. of aICP-OES white suspension. was used The to suspension measure the thickened concentrations as the amount of the of ions; NaOH the solution results added are shown increased. ICP-OES was used to measure the concentrations of the ions; the results are shown in in Figure2. The label “raw” on the horizontal axis represents the untreated brine (i.e., no NaOH Figure 2. The label “raw” on the horizontal axis represents the untreated brine (i.e., no NaOH solution had been added). The left vertical axis shows the concentrations of Mg and Ca, while the solution had been added). The left vertical axis shows the concentrations of Mg and Ca, while the rightright vertical vertical axis axis shows shows the the concentration concentration of of Na. Na. It Itcan can be seen be seen that thatthe concentrations the concentrations of Mg and of MgCa and Ca werewere reproducible; reproducible; however, however, the the concentration concentration of Na of showed Na showed large largestandard standard deviations deviations because becausethe the solutionssolutions werewere prepared for for ICP-OES ICP-OES measurement measurement by bydiluting diluting them them by a by factor a factor of 1000. of 1000. It was It was expectedexpected that that the the concentration concentration of of Na Na would would increaseincrease as as the the pH pH increased, increased, owing owing to the to addition the addition of of NaOHNaOH solution solution to the to the brine. brine.

20,000 2,400 2,200 Ca 19,500 2,000 Mg 1,800 Na 19,000 1,600 1,400 18,500 1,200 1,000 18,000 800 Na (mg/L) Mg, Ca (mg/L) 600 17,500 400 200 17,000 0

raw 10 10.5 11 11.5 12 12.5 13 13.5 pH of brine

FigureFigure 2. Concentrations 2. Concentrations of Mg, of Mg, Ca, Ca, and and Na Na as theas the pH pH of brineof brine was was increased increased by by the the addition addition of of NaOH NaOH solution. Error bars represent one standard deviation. solution. Error bars represent one standard deviation. It is likely that as the measured concentrations of Mg and Ca decreased as NaOH was added to Itthe is brine, likely the that Mg as and the Ca measured ions had formed concentrations their hydr ofoxides. Mg and If a Cabrine decreased suspension as with NaOH a specific was added pH to the brine,resulted the in Mg a substantial and Ca ions reduction had formed in the concentration their hydroxides. of one ion, If a brineit would suspension strongly suggest with athat speciﬁc this pH resultedwas inthe a optimal substantial pH for reduction isolating inMg the and concentration Ca from each other. of one ion, it would strongly suggest that this 2− was the optimalCO2 mineralization pH for isolating is an Mgionic and reaction Ca from between each other.CO3 and (divalent) metal ions. CO2 will change its form depending on the pH environment it is exposed2− to, as shown in the middle column CO2 mineralization is an ionic reaction between CO3 and (divalent) metal ions. CO2 will of Table 2. Therefore, increasing the pH is essential when attempting to enhance the transformation change its form depending on the pH environment it is exposed to, as shown in the middle column of of injected gaseous CO2 into its ionic form, CO32−. The pH can be increased by the addition of alkaline Table2. Therefore, increasing the pH is essential when attempting to enhance the transformation of 2− injected gaseous CO2 into its ionic form, CO3 . The pH can be increased by the addition of alkaline Minerals 2017, 7, 207 5 of 12 agents, which will dissolve and yield OH− in aqueous solutions. If a reactant contains metal ions, metal hydroxide particles will precipitate after reaction with existing OH−. Consequently, there is an increase in pH results in the formation of metal hydroxides and ionization of CO2. This is the reason that addition of NaOH is required.

Table 2. Forms of CO2 varied at pH changes and their reactions with divalent ions.

pH Forms of CO2 Ion Reactions 2+ 2+ Less than neutral (~<6.35) CO2 + H2O → H2CO3 H2CO3 + H2O + M → H2CO3 + H2O + M − − − 2+ 2+ Neutral to weak alkali H2CO3 + OH → HCO3 + H2O (pKa1 = 6.4) 2HCO3 + 2H2O + 2M → M(HCO3)2 ↓ + 2H2O + M − − 2− 2− 2+ Strong alkali (~10.33<) HCO3 + OH → CO3 + H2O (pKa2 = 10.2) 2CO3 + 2H2O + 2M → 2MCO3 ↓ + 2H2O M2+ represents divalent ions such as Mg2+ or Ca2+.

When the pH was increased from 10 to 10.5, the concentration of Mg decreased by approximately 2000 mg/L, and when the pH was increased to 11, the concentration decreased by a further 180 mg/L. After this, the Mg concentration remained relatively constant up to pH 13.5. The concentration of Ca did not change signiﬁcantly as the pH was varied. Consequently, we chose pH 10.5, to allow Mg from brine to form Mg(OH)2 without signiﬁcantly altering the concentration of Ca. Relatively pure Mg(OH)2 can facilitate the production of pure Mg-carbonated particles by CO2 injection. This is the reason for the separation or isolation of Mg and Ca from each other by pH adjustment before CO2 injection. We used the equations below by calculating the ion selectivity of Mg and Ca in hydroxide salts at pH 10.5 and 11.

measured ion concentration at a given pH Precipitated ratio of an ion = 1 − (1) measured ion concentration in untreated brine According to Equation (1), at pH 10.5, 90% of Mg and 23% of Ca in brine precipitated as their hydroxides. As the pH was increased to 11, 99% of Mg and 29% of Ca had formed their hydroxides. Equation (1) seems to indicate that pH 11 would be more favorable for the recovery of Mg and Ca from brine. However, the precipitated ratio does not give enough information to allow a useful pH for the separation of Mg and Ca to be selected: in this study, CO2 mineralization requires relatively pure Mg- and Ca-hydroxides to ensure that pure carbonated particles are obtained.

measured ion concentration at a give pH Ion selectivity = (2) precipitated Mg and Ca concentrations at a given pH

Equation (2) shows that at pH 10.5, the precipitated particles were composed of 92.3% Mg(OH)2 and 7.70% Ca(OH)2. Meanwhile, at pH 11, the precipitated particles were composed of 91.5% Mg(OH)2 and 8.50% Ca(OH)2. Therefore, the Mg(OH)2 obtained was purer at pH 10.5. In addition, adjusting the pH to 10.5 required 36.1 ± 0.5 mL of 1 M NaOH solution for 0.25 L brine, while adjusting to pH 11 required 42.0 ± 0.6 mL of 1 M NaOH solution. Therefore, increasing the pH from 10.5 to 11 required an additional 5.90 mL of NaOH solution: this would increase the operational cost for the process.

3.2. Bubbling CO2 through a Suspension at pH 10.5

CO2 mineralization was performed by bubbling CO2 microbubbles through a brine suspension at pH 10.5. In our previous studies on the precipitation of CaCO3 by CO2 mineralization, we found that as the pH of the suspension of Ca(OH)2 into which CO2 was injected approached neutral, all the particles had formed CaCO3, leading us to regard this as the end of the reaction. We therefore discontinued CO2 injection at this point [7,26]. On the basis of this experience, we initially intended to stop CO2 injection when the pH of the brine approached neutral. However, in contrast to our previous experiments, we found that no particles were suspended at this neutral pH. We therefore decided to MineralsMinerals2017, 20177, 207, 7, 207 6 of 126 of 12

previous experiments, we found that no particles were suspended at this neutral pH. We therefore decided to stop CO2 injection at a higher pH of 8. CO2 at 600 mL/min was injected for 4.5, 5, 7, 9 stop CO2 injection at a higher pH of 8. CO2 at 600 mL/min was injected for 4.5, 5, 7, 9 mins for four experimentalmins for four cycles. experimental cycles. The concentrations of Mg, Ca, and Na in the suspension were measured at the end of each step The concentrations of Mg, Ca, and Na in the suspension were measured at the end of each step in in the cycle shown in Figure 1. The measured results are shown in Figure 3, and the concentrations of the cycle shown in Figure1. The measured results are shown in Figure3, and the concentrations of the the ions in the raw, untreated brine are given for comparison. It can be seen that the concentrations ions inof theMg raw,and Ca untreated decreased brine during are giventhe cycles, for comparison.while the concentration It can be seen of Na that increased the concentrations owing to the of Mg and Caaddition decreased of NaOH during solution. the cycles, Filtration while was the used concentration to collect the ofparticles Na increased after the owingCO2 bubbling to the step addition in of NaOHeach solution. cycle. Filtration was used to collect the particles after the CO2 bubbling step in each cycle.

2,400 30,000

2,200 Ca 2,000 Mg Na 1,800

1,600 25,000 1,400

1,200

1,000

800 Na (mg/L) Ca, Mg (mg/L) Mg Ca, 20,000 600

400

200

0 15,000 RAW pH CO pH CO pH CO pH CO pH CO 2 2 2 2 2 1st 2nd 3rd 4th 5th Cycle Figure 3. Changes in the concentrations of Mg, Ca, and Na during the cycles shown in Figure 1. Figure 3. Changes in the concentrations of Mg, Ca, and Na during the cycles shown in Figure1. The concentration of Ca decreased to tens of mg/L after the first cycle, and then remained Theconstant. concentration This indicates of Ca that decreased all the Ca to intens the brin of mg/Le had precipitated. after the first Interest cycle,ingly, and then the concentrations remained constant. This indicatesof Mg fluctuated that all theduring Ca inthe the cycles. brine The had concentrat precipitated.ion of Interestingly, Mg in the suspension the concentrations decreased of when Mg fluctuated the duringpH the was cycles. increased The concentration to 10.5, and ofthen Mg increased in the suspension as CO2 was decreased bubbled when through thepH the was brine. increased The Mg to 10.5, concentration at pH 10.5 remained constant at around ~300 mg/L, while the concentration after CO2 and then increased as CO2 was bubbled through the brine. The Mg concentration at pH 10.5 remained bubbling decreased as the number of cycles increased. It seems that precipitation of carbonated Mg constant at around ~300 mg/L, while the concentration after CO bubbling decreased as the number of particles was less favorable than precipitation of carbonated Ca2 particles. However, it can be seen cycles increased. It seems that precipitation of carbonated Mg particles was less favorable than precipitation from the Mg and Ca concentrations measured after CO2 bubbling that Mg carbonate precipitated of carbonated Ca particles. However, it can be seen from the Mg and Ca concentrations measured after during the cycles. No particles were formed after increasing the pH and bubbling CO2 through the CO2 bubblingsolution during that Mg the carbonate fifth cycle. precipitated during the cycles. No particles were formed after increasing the pH andBased bubbling on the CO brine2 through that we theused, solution 10 million during tons the of fifthbrine cycle. contains 22,260 tons of Mg and 7140 Basedtons of onCa. the If assume brine that that weMg used, and Ca 10 precipitate million tons as MgCO of brine3 and contains CaCO3 22,260, each ion tons ideally of Mg reacts and with7140 tons of Ca.40,306 If assume and 7854 that tons Mg of andCO2. CaAccording precipitate to the as measured MgCO3 concentrationand CaCO3 ,of each remained ion ideally Mg and reacts Ca in with brine after the fourth cycle, 19,260 and 7080 tons of Mg and Ca react with 34,874 and 7788 tons of 40,306 and 7854 tons of CO2. According to the measured concentration of remained Mg and Ca in brineCO after2, respectively. the fourth The cycle, estimated 19,260 amount and 7080 of MgCO tons3 of and Mg CaCO and3 Caare 66,816 react withand 17,700 34,874 tons and by7788 CO2 tons mineralization. of CO2, respectively. The estimated amount of MgCO3 and CaCO3 are 66,816 and 17,700 tons by CO mineralization. 2 3.3. Characterization of Suspended Particles from CO2 Mineralization

3.3. CharacterizationAfter each pH of Suspendedincreasing/CO Particles2-bubbling from cycle, CO2 Mineralizationthe particles in the suspensions were collected via filtration and characterized. It should be noted that too few particles were formed during the first Aftercycle for each full pH characterizat increasing/COion to be2-bubbling carried out. cycle, the particles in the suspensions were collected via filtration and characterized. It should be noted that too few particles were formed during the first cycle for full characterization to be carried out. Figure4 shows the XRD results. As depicted in the figure, the particles collected from the first cycle consisted of calcite and halite. In addition, some aragonite had crystallized owing to the coexistence of Mg and Ca ions [28]. The CaCO3 successfully precipitated from the non-synthesized Minerals 2017, 7, 207 7 of 12

Figure 4 shows the XRD results. As depicted in the figure, the particles collected from the first cycle consisted of calcite and halite. In addition, some aragonite had crystallized owing to the coexistence of Mg and Ca ions [28]. The CaCO3 successfully precipitated from the non-synthesized Minerals 2017, 7, 207 7 of 12 brine during CO2 mineralization. However, particles collected from other cycles could not be identified positively by standard Joint Committee on Powder Diffraction Standards (JCPDS). Only halitebrine duringphase COwas2 mineralization. identified. As However,the number particles of cy collectedcles increased, from other the cycleshalite couldsignal not became be identiﬁed more intense.positively by standard Joint Committee on Powder Diffraction Standards (JCPDS). Only halite phase was identiﬁed.Mg carbonates As the have number been ofreported cycles increased,to form ma theny halite crystal signal phases, became depending more intense. on the conditions such Mgas pressure, carbonates temperat have beenure, reported water content, to form manyand agitation crystal phases, of the depending reactant liquid on the phase conditions [29–32]. such We as utilizedpressure, other temperature, analytical water methods content, that and would agitation provide of the reactantinformation liquid regarding phase [29– 32the]. Wepresence utilized of other Mg carbonatesanalytical methods in the particles. that would provide information regarding the presence of Mg carbonates in the particles.

C: calcite (05-0586) C A: aragonite (24-0025) H 1st cycle

A * H: halite (05-0628) *: unknown H 2nd cycle A.U.

* H * * 3rd cycle

* * H * 4th cycle

10 20 30 40 50 60 2θ (degree)

FigureFigure 4. XRDXRD results results of of the the particles collected from the first ﬁrst to fourth cycles.

Elemental mapping by SEM-EDS revealed that the particles from the second, third, and fourth Elemental mapping by SEM-EDS revealed that the particles from the second, third, and fourth cycles consisted of Ca and Mg carbonates (Figure 5). This shows that the Mg formed carbonates, cycles consisted of Ca and Mg carbonates (Figure5). This shows that the Mg formed carbonates, despite the fact that the precipitation of Mg carbonates is less favorable than the precipitation of Ca despite the fact that the precipitation of Mg carbonates is less favorable than the precipitation of Ca carbonates because Mg carbonates are more soluble than Ca carbonates and hydrated Mg ions carbonates because Mg carbonates are more soluble than Ca carbonates and hydrated Mg ions possess possess more energy [33]. Figure 5 shows that the elemental distributions of Ca, Mg, C, and O were more energy [33]. Figure5 shows that the elemental distributions of Ca, Mg, C, and O were broadly broadly similar. The second cycle produced Ca carbonate, while the third and fourth cycles similar. The second cycle produced Ca carbonate, while the third and fourth cycles produced Mg produced Mg carbonate. The samples were prepared on silicon wafers to prevent the appearance of carbonate. The samples were prepared on silicon wafers to prevent the appearance of the carbon signal the carbon signal from the commonly used carbon adhesive tape. from the commonly used carbon adhesive tape. The particles were also subjected to FTIR analysis; the results are shown in Figure 6. The The particles were also subjected to FTIR analysis; the results are shown in Figure6. The particles particles that precipitated during the second cycle showed FTIR absorption bands at 727, 873, and that precipitated during the second cycle showed FTIR absorption bands at 727, 873, and 1406 cm−1. 1406 cm−1. Several research papers that identified Mg carbonates using FTIR analysis suggest that the Several research papers that identiﬁed Mg carbonates using FTIR analysis suggest that the peak at peak at 727 cm−1 was likely caused by the presence CaCO3 in the structure. Long et al. [34] reported 727 cm−1 was likely caused by the presence CaCO in the structure. Long et al. [34] reported that that peaks at 724, 726, and 730 cm−1 were obtained3 from calcite particles including 30 mol% Mg. peaks at 724, 726, and 730 cm−1 were obtained from calcite particles including 30 mol% Mg. Similarly, Similarly, Gunasekaran et al. [34,35] and Xu and Poduska [36] reported dolomite peaks at 729 and Gunasekaran et al. [34,35] and Xu and Poduska [36] reported dolomite peaks at 729 and 730 cm−1, 730 cm−1, respectively. These peaks arose from CO32−. Consequently, it is possible to assign the IR 2− respectively. These peaks arose from CO3 . Consequently, it is possible to assign the IR peak at peak at 727 cm−1 to Mg carbonates. In addition, the peaks at 873 and 1406 cm−1 were from the CO32− −1 −1 2− 727 cm to Mg carbonates. In addition, the peaks at 873 and 1406 cm were from the CO3 ions in ions in CaCO3 [35,37], which was already known to be present due to XRD analysis. CaCO [35,37], which was already known to be present due to XRD analysis. The3 particles obtained after the third and fourth cycles both showed peaks at 852 and 877 cm−1. The particles obtained after the third and fourth cycles both showed peaks at 852 and 877 cm−1. In these cycles, both Ca and Mg carbonates formed. The peak at 852 cm−1 was due to calcite that In these cycles, both Ca and Mg carbonates formed. The peak at 852 cm−1 was due to calcite that included Mg. Long et al. [34] reported that calcite including Mg at 20 and 39 mol% absorbed IR at included Mg. Long et al. [34] reported that calcite including Mg at 20 and 39 mol% absorbed IR at 852 cm−1. CO32− ions in dolomite also absorbed IR near this position, at 853 cm−1. Meanwhile, the peak −1 2− −1 852 cm . CO3 ions in dolomite also absorbed IR near this position, at 853 cm . Meanwhile, at 877 cm−1 indicates the presence of CaCO3 [35,36,38–40]. Although we expected that the third and the peak at 877 cm−1 indicates the presence of CaCO [35,36,38–40]. Although we expected that the fourth cycles would produce Mg carbonates that excluded3 Ca, some of the Ca remaining in the third and fourth cycles would produce Mg carbonates that excluded Ca, some of the Ca remaining suspension after the second cycle managed to participate in the carbonation of Mg. The peaks at 1418 in the suspension after the second cycle managed to participate in the carbonation of Mg. The peaks at 1418 and 1479 cm−1 are typical absorption peaks of hydromagnesite. In addition, the peak at Minerals 2017, 7, 207 8 of 12 and 1479 cm−1 are typical absorption peaks of hydromagnesite. In addition, the peak at 3647 cm−1 3647provides cm− 1furtherprovides evidence further evidenceof the ofpresence the presence of hydromagnesite. of hydromagnesite. Mg Mg from from brine brine successfully precipitatedMinerals 2017as hydromagnesite, , 7, 207 indicating indicating th thatat brine brine could could be be a a useful useful resource resource for for CO CO2 2mineralization.mineralization.8 of 12

and 1479 cm−1 are typical absorption peaks of hydromagnesite. In addition, the peak at 3647 cm−1 provides further evidence of the presence of hydromagnesite. Mg from brine successfully precipitated as hydromagnesite, indicating that brine could be a useful resource for CO2 mineralization.

Figure 5. Elemental distributions of Mg, Ca, C, and O in particles produced during the experimental cycles. FigureFigure 5. 5.Elemental Elemental distributions distributions of of Mg, Mg, Ca, Ca,C, C,and and OO inin particlesparticles produced during the the experimental experimental cycles. cycles.

1418 877 852 798 708 1418 877 852 798 708 36473647 1479 590590

2nd 2nd cycle cycle 3rd 3rd cycle cycle 4th cycle 4th cycle

Transmittance Transmittance

727 1406 873 727 4000 3500 3000 2500 2000 15001406 1000873 500

-1 4000 3500 3000Wavenumber 2500 2000 (cm ) 1500 1000 500 -1 FigureFigure 6. FTIR 6. FTIR spectra spectra of particlesof particles fromWavenumber from the the second, seco nd,(cm third, ) and and fourth fourth carbonation carbonation cycles. cycles.

Figure 6. FTIR spectra of particles from the second, third, and fourth carbonation cycles.

Minerals 2017, 7, 207 9 of 12

InMinerals addition, 2017, 7, 207 characterization by TGA was used to conﬁrm the presence of hydromagnesite.9 of 12 The weight losses of the particles we obtained, shown in Figure7a, were compared with those of commercialIn addition, Mg(OH) characterization2, CaCO3, MgCO by3 TGA, and was NaCl, used which to confirm are shown the presence in Figure of hydromagnesite.7b. Weight loss The due to ◦ ◦ ◦ Mg(OH)weight2 occurred losses of between the particles 328 and we 422 obtained,C. Losses show atn 589–778in Figure C7a, and were 180–524 comparedC were with due those to CaCOof 3 and hydromagnesite,commercial Mg(OH) respectively.2, CaCO3, MgCO By3 comparing, and NaCl, which Figure are7a,b, shown it can in beFigure seen 7b. that Weight the particlesloss due to from Mg(OH)2 occurred between 328 and 422 °C. Losses at 589–778 °C and 180–524 °C were due to CaCO3 the second cycle were composed of Mg(OH)2 and CaCO3. A small amount of NaCl was included and hydromagnesite, respectively. By comparing Figure 7a,b, it can be seen that the particles from in all the particles, as shown by the weight losses that began at 810 ◦C. The TGA curves of the the second cycle were composed of Mg(OH)2 and CaCO3. A small amount of NaCl was included in particlesall the from particles, the third as shown and by fourth the weight cycles losses were that quite began different at 810 °C. from The that TGA obtained curves of forthe theparticles particles ◦ fromfrom the secondthe third cycle. and fourth The weightcycles were losses quite at differ ~450entC from in the that curves obtained of thefor the particles particles from from the thethird and fourthsecond cyclescycle. The did weight not quite losses match at ~450 with °C what in the would curves beof expectedthe particles for from the lossthe third of hydromagnesite. and fourth Montes-Hermandezcycles did not quite et al. [match41] synthesized with what hydromagnesitewould be expected and for characterized the loss of particleshydromagnesite. of it using TGA.Montes-Hermandez They found that hydromagnesite et al. [41] synthesized exhibited hydromagnesite weight loss atand 450 characterized◦C with some particles ﬂuctuation of it using at around 375 ◦CTGA., which They precisely found that matched hydromagnesite the results exhibited shown for weight the particles loss at 450 from °C thewith third some and fluctuation fourth cycles at in Figurearound7a. There 375 °C, were which also precisely some differences matched the between results shown the TGA for the curves particles obtained from the for third the particlesand fourth from the thirdcycles and in fourthFigure cycles.7a. There The were curve also for some cycle diffe threerences indicated between an additionalthe TGA curves weight obtained loss of for 4.5% the in the particles from the third and fourth cycles. The curve for cycle three indicated an additional weight range of 635–665 ◦C, which indicates the coexistence of calcite and hydromagnesite. As the number of loss of 4.5% in the range of 635–665 °C, which indicates the coexistence of calcite and hydromagnesite. cyclesAs increased, the number the of puritycycles increased, of the hydromagnesite the purity of the that hydromagnesite precipitated that increased. precipitated increased.

110 (a)

100

60 Weight (%) 2nd cycle 50 3rd cycle 4th cycle 40

200 400 600 800 1000

o Temperature ( C) 110 (b) 100

60 weight (%) weight 50

40 Mg(OH) CaCO 2 3 30 MgCO NaCl 3

200 400 600 800 1000

ο Temperature ( C)

FigureFigure 7. Thermogravimetric 7. Thermogravimetric analysis analysis (TGA) (TGA) curvescurves of ( (aa)) the the precipitates precipitates obtained obtained from from the second, the second, third, and fourth cycles; and (b) the commercial reagents of Mg(OH)2, MgCO3, CaCO3, and NaCl. third, and fourth cycles; and (b) the commercial reagents of Mg(OH)2, MgCO3, CaCO3, and NaCl.

Minerals 2017, 7, 207 10 of 12

4. Conclusions

Non-synthesized brine from a functional RO seawater desalination plant was used for CO2 mineralization to fix CO2 under ambient temperature and pressure. To the best of our knowledge, this is the first time that non-synthesized brine has been used for CO2 mineralization. Cycles consisting of an increase in pH followed by CO2 injection and finally filtration precipitated 99% of Ca and 86% of Mg from the brine, using them to fix CO2 in the form of carbonated minerals. CO2 was successfully mineralized with Ca to yield CaCO3 and with Mg to produce hydromagnesite. This was shown by the characterization of the precipitates obtained from the cycles by XRD, SEM-EDS, FTIR, and TGA-DTA. It was thought that co-presence of Ca with Mg was not favorable for the precipitation of Mg carbonates. In this study, we precipitated hydromagnesite followed by precipitation of calcite from brine. This means that the concentrations of Mg and Ca in brine are twice those in influent seawater, making it a good source of divalent ions for CO2 mineralization. According to our experiments, 42,662 tons of CO2 could be sequestered from 10 million tons of the brine used for a year. Moreover, 17,700 tons of CaCO3 and 66,816 tons of MgCO3 would be obtained.

Acknowledgments: This research was supported by the Basic Research Project (17-3422) of the Korea Institute of Geoscience and Mineral Resources (KIGAM), funded by the Ministry of Science and Information and Communications Technologies (ICT). Author Contributions: Jun-Hwan Bang and Soochun Chae conceived and designed the experiments; Yeongsuk Yoo, Seung-Woo Lee and Kyungsun Song performed experiments and analyzed the data; Jun-Hwan Bang wrote the paper. Conﬂicts of Interest: The authors declare no conﬂict of interest.

References

1. Ko, D.-H. Drought Causes Water Shortage Crisis in Southern Korea; The Korea Times: Seoul, Korea, 2017. 2. Papapetrou, M.; Cipollina, A.; la Commare, U.; Micale, G.; Zaragoza, G.; Kosmadakis, G. Assessment of methodologies and data used to calculate desalination costs. Desalination 2017, 419, 8–19. [CrossRef] 3. Charcosset, C. A review of membrane processes and renewable energies for desalination. Desalination 2009, 245, 214–231. [CrossRef] 4. Alkhudhiri, A.; Darwish, N.; Hilal, N. Membrane distillation: A comprehensive review. Desalination 2012, 287, 2–18. [CrossRef] 5. Piyadasa, C.; Ridgway, H.F.; Yeager, T.R.; Stewart, M.B.; Pelekani, C.; Gray, S.R.; Orbell, J.D. The application of electromagnetic ﬁelds to the control of the scaling and biofouling of reverse osmosis membranes—A review. Desalination 2017, 418, 19–34. [CrossRef] 6. Sanna, A.; Uibu, M.; Caramanna, G.; Kuusik, R.; Maroto-Valer, M.M. A review of mineral carbonation

technologies to sequester CO2. Chem. Soc. Rev. 2014, 43, 8049–8080. [CrossRef][PubMed] 7. Bang, J.-H.; Kim, W.; Song, K.S.; Jeon, C.W.; Chae, S.C.; Cho, H.-J.; Jang, Y.N.; Park, S.-J. Effect of experimental

parameters on the carbonate mineralization with CaSO4·2H2O using CO2 microbubbles. Chem. Eng. J. 2014, 244, 282–287. [CrossRef] 8. Druckenmiller, M.L.; Maroto-Valer, M.M. Carbon sequestration using brine of adjusted pH to form mineral carbonates. Fuel Process. Technol. 2005, 86, 1599–1614. [CrossRef] 9. Dri, M.; Sanna, A.; Maroto-Valer, M.M. Mineral carbonation from metal wastes: Effect of solid to liquid ratio on the efﬁciency and characterization of carbonated products. Appl. Energy 2014, 113, 515–523. [CrossRef] 10. Zevenhoven, R.; Wiklund, A.; Fagerlund, J.; Eloneva, S.; Veen, B.I.; Geerlings, H.; van Mossel, G.; Boerrigter, H. Carbonation of calcium-containing mineral and industrial by-products. Front. Chem. Eng. China 2010, 4, 110–119. [CrossRef] 11. Baciocchi, R.; Costa, G.; di Bartolomeo, E.; Polettini, A.; Pomi, R. The effects of accelerated carbonation on

CO2 uptake and metal release from incineration APC residues. Waste Manag. 2009, 29, 2994–3003. [CrossRef] [PubMed] 12. Baciocchi, R.; Costa, G.; Lategano, E.; Marini, C.; Polettini, A.; Pomi, R.; Postorino, P.; Rocca, S. Accelerated carbonation of different size fractions of bottom ash from RDF incineration. Waste Manag. 2010, 30, 1310–1317. [CrossRef][PubMed] Minerals 2017, 7, 207 11 of 12

13. Bang, J.-H.; Lee, S.-W.; Jeon, C.; Park, S.; Song, K.; Jo, W.; Chae, S. Leaching of Metal Ions from Blast Furnace

Slag by Using Aqua Regia for CO2 Mineralization. Energies 2016, 9, 996. [CrossRef] 14. Pan, S.-Y.; Chang, E.; Chiang, P.-C. CO2 capture by accelerated carbonation of alkaline wastes: A review on its principles and applications. Aerosol Air Qual. Res. 2012, 12, 770–791. [CrossRef] 15. Bałdyga, J.; Henczka, M.; Sokolnicka, K. Utilization of carbon dioxide by chemically accelerated mineral carbonation. Mater. Lett. 2010, 64, 702–704. [CrossRef]

16. Lee, S.; Kim, J.-W.; Chae, S.; Bang, J.-H.; Lee, S.-W. CO2 sequestration technology through mineral carbonation: An extraction and carbonation of blast slag. J. CO2 Util. 2016, 16, 336–345. [CrossRef] 17. Soong, Y.; Fauth, D.L.; Howard, B.H.; Jones, J.R.; Harrison, D.K.; Goodman, A.L.; Gray, M.L.; Frommell, E.A.

CO2 sequestration with brine solution and ﬂy ashes. Energy Convers. Manag. 2006, 47, 1676–1685. [CrossRef] 18. Li, P.; Pan, S.-Y.; Pei, S.; Lin, Y.J.; Chiang, P.-C. Challenges and perspectives on carbon ﬁxation and utilization technologies: An overview. Aerosol Air Qual. Res. 2016, 16, 1327–1344. [CrossRef] 19. Shenvi, S.S.; Isloor, A.M.; Ismail, A.F. A review on RO membrane technology: Developments and challenges. Desalination 2015, 368, 10–26. [CrossRef]

20. Morrison, J.; Jauffret, G.; Galvez-Martos, J.L.; Glasser, F.P. Magnesium-based cements for CO2 capture and utilisation. Cem. Concr. Res. 2016, 85, 183–191. [CrossRef]

21. Liu, Q.; Maroto-Valer, M.M. Studies of pH buffer systems to promote carbonate formation for CO2 sequestration in brines. Fuel Process. Technol. 2012, 9, 6–13. [CrossRef] 2+ 2+ 22. Wang, W.; Liu, X.; Wang, P.; Zheng, Y.; Wang, M. Enhancement of CO2 Mineralization in Ca -/Mg -Rich Aqueous Solutions Using Insoluble Amine. Ind. Eng. Chem. Res. 2013, 52, 8028–8033. [CrossRef] 23. Xie, H.; Liu, T.; Hou, Z.; Wang, Y.; Wang, J.; Tang, L.; Jiang, W.; He, Y. Using electrochemical process to 2+ 2+ mineralize CO2 and separate Ca /Mg ions from hard water to produce high value-added carbonates. Environ. Earth Sci. 2015, 73, 6881–6890. [CrossRef]

24. Langmuir, D. Stability of Carbonates in the System MgO-CO2-H2O. J. Geol. 1965, 73, 730–754. [CrossRef] 25. Lide, D.R. CRC Handbook of Chemistry and Physics; CRC Press: Boca Raton, FL, USA, 1997. 26. Bang, J.-H.; Jang, Y.N.; Kim, W.; Song, K.S.; Jeon, C.W.; Chae, S.C.; Lee, S.-W.; Park, S.-J.; Lee, M.G. Precipitation of calcium carbonate by carbon dioxide microbubbles. Chem. Eng. J. 2011, 174, 413–420. [CrossRef]

27. Bang, J.-H.; Song, K.; Park, S.; Jeon, C.; Lee, S.-W.; Kim, W. Effects of CO2 Bubble Size, CO2 Flow Rate and Calcium Source on the Size and Speciﬁc Surface Area of CaCO3 Particles. Energies 2015, 8, 12304. [CrossRef] 28. Santos, R.M.; Bodor, M.; Dragomir, P.N.; Vraciu, A.G.; Vlad, M.; van Gerven, T. Magnesium chloride as a leaching and aragonite-promoting self-regenerative additive for the mineral carbonation of calcium-rich materials. Miner. Eng. 2014, 59, 71–81. [CrossRef]

29. Cheng, W.; Li, Z. Nucleation kinetics of nesquehonite (MgCO3·3H2O) in the MgCl2-Na2CO3 system. J. Cryst. Growth 2010, 312, 1563–1571. [CrossRef] 30. Hales, M.C.; Frost, R.L.; Martens, W.N. Thermo-Raman spectroscopy of synthetic nesquehonite —Implication for the geosequestration of greenhouse gases. J. Raman Spectrosc. 2008, 39, 1141–1149. [CrossRef] 31. Coleyshaw, E.E.; Crump, G.; Grifﬁth, W.P. Vibrational spectra of the hydrated carbonate minerals ikaite, monohydrocalcite, lansfordite and nesquehonite. Spectrochim. Acta Part A Mol. Biomolecul. Spectroscop. 2003, 59, 2231–2239. [CrossRef]

32. Hopkinson, L.; Kristova, P.; Rutt, K.; Cressey, G. Phase transitions in the system MgO–CO2–H2O during CO2 degassing of Mg-bearing solutions. Geochim. Cosmochim. Acta 2012, 76, 1–13. [CrossRef] 33. Xu, J.; Yan, C.; Zhang, F.; Konishi, H.; Xu, H.; Teng, H.H. Testing the cation-hydration effect on the

crystallization of Ca–Mg–CO3 systems. Proc. Natl. Acad. Sci. USA 2013, 110, 17750–17755. [CrossRef] [PubMed] 34. Long, X.; Nasse, M.J.; Ma, Y.; Qi, L. From synthetic to biogenic Mg-containing calcites: A comparative study using FTIR microspectroscopy. Phys. Chem. Chem. Phys. 2012, 14, 2255–2263. [CrossRef][PubMed] 35. Gunasekaran, S.; Anbalagan, G.; Pandi, S. Raman and infrared spectra of carbonates of calcite structure. J. Raman Spectrosc. 2006, 37, 892–899. [CrossRef] 36. Xu, B.; Poduska, K.M. Linking crystal structure with temperature-sensitive vibrational modes in calcium carbonate minerals. Phys. Chem. Chem. Phys. 2014, 16, 17634–17639. [CrossRef][PubMed] 37. Pandey, S.; Sengupta, J. Fourier transform infrared spectroscopy: A tool for detection of lime content in hot mix asphalt. In Proceedings of the 26th ARRB Conference, Sydney, Australia, 19–22 October 2014; p. 11. Minerals 2017, 7, 207 12 of 12

38. Bourrata, X.; Qiao, L.; Feng, Q.; Angellier, M.; Dissaux, A.; Beny, J.-M.; Barbin, V.; Stempﬂé, P.; Rousseau, M.; Lopez, E. Origin of growth defects in pearl. Mater. Charact. 2012, 72, 94–103. [CrossRef] 39. Junfeng, J.; Yun, G.; Balsam, W.; Damuth, J.E.; Jun, C. Rapid identiﬁcation of dolomite using a Fourier transform infrared spectrophotometer (FTIR): A fast method for identifying Heinrich events in IODP Site U1308. Mar. Geol. 2009, 258, 60. 40. Bang, J.-H.; Song, K.S.; Lee, M.G.; Jeon, C.W.; Jang, Y.N. Effect of Critical Micelle Concentration of Sodium Dodecyl Sulfate Dissolved in Calcium and Carbonate Source Solutions on Characteristics of Calcium Carbonate Crystals. Mater. Trans. 2010, 51, 1486–1489. [CrossRef] 41. Montes-Hernandez, G.; Renard, F.; Chiriac, R.; Findling, N.; Toche, F. Rapid Precipitation of

Magnesite Microcrystals from Mg(OH)2-H2O-CO2 Slurry Enhanced by NaOH and a Heat-Aging Step (from ~20 to 90 ◦C). Cryst. Growth Des. 2012, 12, 5233–5240. [CrossRef]

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