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Role of brucite dissolution in calcium carbonate precipitation from artificial and natural seawaters Dan Nguyen Dang, Stéphanie Gascoin, Alaric Zanibellato, Cosmelina G. Da Silva, Mélanie Lemoine, Benoît Riffault, René Sabot, Marc Jeannin, Daniel Chateigner, and Otavio Gil Cryst. Growth Des., Just Accepted Manuscript • DOI: 10.1021/acs.cgd.6b01305 • Publication Date (Web): 13 Feb 2017 Downloaded from http://pubs.acs.org on February 15, 2017
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Role of brucite dissolution in calcium carbonate
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precipitation from artificial and natural seawaters
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Dan Nguyen Danga*, Stéphanie Gascoinb, Alaric Zanibellatoc, Cosmelina G. Da Silvaa,b,
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Mélanie Lemoinea, Benoît Riffaulta, René Sabotc, Marc Jeanninc, Daniel Chateignerb, Otavio
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Gila
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a
Normandie Univ, UNICAEN, UNIROUEN, ABTE, 14000 Caen, France
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b
Normandie Univ, ENSICAEN, UNICAEN, CNRS, CRISMAT, 14000 Caen, France
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c
Univ. La Rochelle, CNRS, LaSIE, 17000 La Rochelle, France
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ABSTRACT
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Brucite is often associated to calcium carbonate in calcareous scales. In this study, we focus on
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the effects of brucite dissolution onto the carbonate precipitation. Brucite-saturated solutions are
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elaborated by adding the same amount of brucite powder in various volumes of different natural
13
and artificial seawaters (natural organic materials are also added; with or without Mg2+ ions).
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The Mg2+, Ca2+ concentrations in seawater at the beginning and end of the experiments are
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determined by ICP-AES. The solid powders are ex-situ analyzed using XRD and SEM. It is
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observed that brucite dissolution leads to pH increase and enrichment of seawaters in magnesium
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ions. The Mg2+ ions, either pre-existing in the solution or added from brucite dissolution play a
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key role in CaCO3 polymorphs selection. In order to separate the Mg2+ and OH- effects, CaCO3
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precipitation is induced by NaOH addition in artificial seawater containing or not Mg2+. Under
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both NaOH and Mg(OH)2 addition, only aragonite is precipitated from artificial and natural
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seawaters in our conditions. NaOH addition in Mg-free seawater allows to predominantly obtain
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vaterite or calcite polymorphs, while both aragonite and calcite precipitates are observed from
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Mg-free seawater when NaOH is replaced by Mg(OH)2. Natural organic materials added in
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artificial seawater have a significant effect on aragonite morphology.
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Introduction
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Calcium carbonate plays a fundamental role in marine ecosystems. It is an important and
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usually one of the dominant minerals in marine sediments.1 The mechanisms of CaCO3
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formation in both aqueous solution2-12 and seawaters1,
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authors. The rate of calcium carbonate formation in solution increases with temperature, which
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implies that its solubility decreases.5 The most important calcium carbonate polymorphs in
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seawater are aragonite and calcite. Other phases (i.e. vaterite, hydrated CaCO3) are far less
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abundant and of much lower importance, except for particular cases corresponding to
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biomineralisation processes23, high ion concentrations, and/or the presence of microrganisms.20
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Seawater is supersaturated with respect to calcite and aragonite, and this supersaturation should
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favor a massive precipitation of calcium carbonate. However, calcium carbonate precipitation in
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surface seawater is almost absent. This has been explained by the presence of soluble organic
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matter (e.g. humic compounds) and/or dissolved phosphate interacting with the surface of
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crystals.20, 24, 25 Berner et al.25 suggested that the supersaturation of surface seawater with respect
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to aragonite is maintained because suspended particles – may play a role as precipitation nuclei –
13-22
have been studied by numerous
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are coated by humic compounds and/or dissolved phosphate. Magnesium plays a key role in
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calcite–aragonite polymorph selection. The role of magnesium on the crystal growth of calcite
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and aragonite from seawater has been studied for a long time.1, 17, 20, 21, 26, 27 Berner26 concluded
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that dissolved magnesium in seawater has a strong retarding effect on calcite crystal growth rate,
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but no effect on that of aragonite. The inhibition effect of Mg2+ on calcite growth rate is due to
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its incorporation within the crystal structure, resulting in the considerably more soluble
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magnesian calcite (Mg-calcite). However, a recent study of Sun and co-workers 21 shows that the
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increased solubility of Mg–calcite has negligible impact on nucleation rates and the inhibition of
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calcite nucleation upon Mg2+ uptake is primarily due to an increase in surface energy.
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A tremendous amount of work has been dedicated to the understanding of CaCO3
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biomineralization.28, 29 The mollusk shell is a product of biomineralization of CaCO3 crystals.
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Nacre is a composite consisting of aragonite pseudo-hexagonal crystals embedded in an organic
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matrix.30 This biomaterial is biocompatible and is known to have some osteogenic effect in
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sheep31, mouse32, rat33, and human34 mammals. Rousseau et al.30 have suggested that nacre
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water-soluble organic matrix (WSM) from the bivalve Pinctada margaritifera could speed up the
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differentiation and mineralization of damaged bone tissue. Although the organic matrix is
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quantitatively a minor constituent in the shell of mollusks (typically less than 5 wt%), it is
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however the major component that controls different aspects of the shell formation processes.29
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In reality the CaCO3 precipitation mechanism needs supersaturation conditions obtained by an
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increase of carbonate concentrations due to physico-chemical displacements of acid-basic
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carbonic equilibrium in seawater. This phenomenon can particularly be induced by technological
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processes involving seawater electrolysis, in which calcium carbonate precipitation is always
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associated to the one of brucite (Mg(OH)2), forming a mineral scale called calcareous deposit.
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Brucite is also a thermodynamically stable product of serpentinization in hydrothermal vent
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systems.35 The analysis of chimneys collected from the serpentinite-hosted Lost City
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hydrothermal field 36 and in the Southern Mariana Forearc37 has indicated that it consists mostly
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of brucite and aragonite. Many studies of brucite surface chemistry have focused on its
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precipitation and dissolution kinetics.38-41 Calcareous scales formed during the cathodic
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protection of metals in seawater are reported to consist of mainly CaCO3 and Mg(OH)2 (brucite).
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The Mg(OH)2/CaCO3 ratios in calcareous scales depend strongly on the applied potentials. A lot
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of studies42-44 show that brucite is predominantly observed for more cathodic potentials while
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less cathodic potentials favor CaCO3 formation. Nevertheless, few studies exist on the effect of
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brucite dissolution onto the carbonate calcium precipitation from seawater.
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The objective of this paper is to study the role of brucite dissolution in CaCO3 precipitation
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from seawaters. The effect of WSM extracted from the nacre Pinctada margaritifera onto the
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brucite dissolution and CaCO3 precipitation, especially its morphology, is also mentioned.
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Finally, the vaterite-calcite-aragonite polymorphs selection under pH monitoring is studied.
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Experimental section
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Four different seawater solutions were prepared:
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NSW: Natural SeaWater (pH~8) taken at the marine station of Luc-sur-Mer (France)
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where it was pumped directly from the English Channel and decanted to remove large-particle
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sediments.
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ASW: Artificial SeaWater prepared on the basis of the six main compounds present in the
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D1141-98 standard.45 The composition is NaCl: 0.42 mol dm−3; Na2SO4: 2.88×10−2 mol dm−3;
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CaCl2·2H2O: 1.05×10−2 mol dm−3; MgCl2·6H2O: 5.46×10−2 mol dm−3; KCl: 9.32×10−3 mol
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dm−3; NaHCO3: 2.79×10−3 mol dm−3. To avoid precipitation of CaCO3, CaSO4 while combining
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the constituents of ASW, we prepared it in two separate containers. In the first container, we
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prepared a solution containing NaCl, Na2SO4, KCl and NaHCO3 salts, while in the other, the
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desired amounts of CaCl2·2H2O and MgCl2·6H2O were dissolved. After the salts in two
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solutions were completely dissolved in pure water, they were carefully combined while stirring.
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The pH of the solutions was sometimes adjusted to ~8 with a few drops of 0.1M NaOH (to reach
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a similar pH to natural seawater of the English Channel).
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ASW+WSM: Water-Soluble organic Matrix (WSM) - extracted from the nacre of
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Pinctada margaritifera using a methodology described elsewhere46, 47 - was added to ASW in
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order to reach a concentration of 100mg WSM / 1dm3 ASW.
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ASW-Mg2+: ASW to which no magnesium salt was added. In this latter, ionic strength
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deficiency was compensated by addition of an excess of NaCl. The composition is NaCl: 0.53
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mol cm−3; Na2SO4: 2.88×10−2 mol dm−3; CaCl2·2H2O: 1.05×10−2 mol dm−3; KCl: 9.32×10−3 mol
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dm−3; NaHCO3: 2.79×10−3 mol dm−3.
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The experiments were carried out by adding the same amount (1g) of brucite powder (Alfa
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Aesar 95-100.5%) in different volumes (50, 200, 500 and 900 cm3) of the four above seawaters.
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This brucite amount generates in any case brucite-saturated solutions at equilibrium conditions.
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The chemical speciation of the solutions was determined using the PHREEQC computer
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program.48 More details of speciation calculations can be found in the Supporting Information.
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In order to determine Mg2+ and Ca2+ concentrations, the solutions were analyzed by ICP-AES
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(Inductively Coupled Plasma-Atomic Emission Spectrometry, Varian, Vista MPX) at the
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beginning and at the end of the experiments. The evolution of Mg2+ and Ca2+ concentrations in
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aqueous phase permitted to determine the quantities of dissolved Mg(OH)2
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and precipitated CaCO3 .
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= − " # ∗ % ∗
(1)
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= &' " − &' # ∗ % ∗
(2)
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where
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, &' are the Mg2+, Ca2+ concentrations at the end of the experiments
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" , &' " are the Mg2+, Ca2+ concentrations at the beginning of the experiments
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V is the seawater volume
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, are the molar masses of Mg(OH)2 and CaCO3.
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The total solid mass variation at the end of the experiment, ∆m, was calculated using the
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following formula:
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∆m = –
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In order to understand the roles of pH and Mg2+ concentration on CaCO3 precipitation from
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seawaters, other experiments were also carried out by gradually adding a NaOH solution in
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900mL of ASW or ASW-Mg2+ solutions to reach a pH of 9.4 or 10.3 respectively. These pHs
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have been chosen based on the values obtained at the end of the experiments of brucite
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dissolution in the seawaters.
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The pH was measured with a Cyberscan pH510 from Eutech Instruments. Low, continuous,
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solution agitation was maintained using magnetic stirring. All experiments were carried out at
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ambient temperature (20±1°C) and were performed in an open reactor during 24 hours for the
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volumes 20 cm3, 200 cm3 and 500 cm3 and 72 hours for the 900 cm3 solution.
(3)
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Precipitated powders obtained at the end of these experiments were filtered, washed with ultra-
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pure water to remove the soluble salts (in particular NaCl) and dried before X-Ray Diffraction
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(XRD) and Scanning Electron Microscopy (SEM, Carl Zeiss 5500 instrument) experiments. X-
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ray powder diffraction data were collected using a D8 Advance Vario 1 Bruker two-circle
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diffractometer (θ–2θ Bragg-Brentano mode) using the monochromatized Cu Kα radiation (λ =
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1.540598 Å) and the Lynx Eye detector. Data were collected over the angular range 5° ≤ 2θ ≤
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120° at room temperature counting for 0.8 s at each angular increment of 0.0105° (10 h/scan).
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XRD patterns analyzed using the Combined Analysis methodology49 allow the quantitative
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determination of phase fractions.
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Results and discussion
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pH evolution
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In a preliminary test, brucite powder is added to stirred ultra-pure water. The test is operated in a
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closed reactor to avoid the CO2(g) dissolution and MgCO3 or relevant phases precipitation. The
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pH increases abruptly to 10.3 ± 0.1 in just a few minutes and remains stable for the total duration
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of experiments. This pH is similar to the value numerically determined using the PHREEQC
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computer program (Table S1 of the Supporting Information). The pH rising is due to the brucite
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solubility (solubility product pK°sp = 10.5±0.2 at 25°C) and relatively fast dissolution kinetics:40
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() ⇌ + 2() -
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Dissolution of the brucite powder in NSW, ASW and ASW+WSM solutions (initial pH ~8)
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results in a rapid increase in pH to 9.4 ± 0.1 after a few minutes. These pHs remain almost stable
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until the end of the tests (24 hours for V = 50, 200, 500 cm3 and 72 hours for V = 900 cm3).
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Figure 1 shows the final pHs obtained at the end of the experiments. These measured pHs are
(reaction 1)
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consistent with the values simulated using PHREEQC (table S1 and figure S1 of the Supporting
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Information). A slight pH difference is observed among different volumes for the same seawater
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or among different seawaters but smaller than the measured pH standard deviation (± 0.1). The
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difference between the pH obtained in the ultra-pure water (~10.3) and seawaters (~9.4) can be
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partially explained by the pre-existence of Mg2+ ions in seawaters (~1300 mg.cm-3), which
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reduced the OH- concentration according to the displacement of brucite dissolution equilibrium
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(reaction 1).
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This remark is consistent with the pH~10.3 experimentally reached in the ASW-Mg2+ solution
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during the first 6 hours. This value is similar with that obtained for the case of brucite dissolution
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in ultra-pure water. Indeed, the absence of Mg2+ ions favors the equilibrium displacement in the
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direction of the dissolution and thus to a higher release of OH-. However, after 6h of experiment,
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the pH begins to decrease slightly and reaches ~9.8 at the end of the tests. This pH decrease (that
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is not previously observed in ultra-pure water) suggests that the buffer effect of seawater also
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played an important role. The reaction rate between )&(.- and &(. - in the ASW-Mg2+
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solutions (reaction 2), which is slower than that of brucite dissolution, leads to a slight and
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progressive decrease of pH and reaches ~9.8 at the end of the tests. Dreybrodt et al.50 have found
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that the precipitation kinetics of calcite is controlled by the slow kinetics of the overall reaction
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&( + ) ( ⇌ )&(.- + ) . PHREEQC modeling of brucite dissolution equilibrium indicates
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that the pH obtained in ASW-Mg2+ solution (10.13) is slightly lower than the one reached in pure
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water (10.37). This difference confirms the role of the )&(.- / &(. - equilibrium in the ASW-
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Mg2+ solution.
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)&(.- + () - ⇌ &(. - + ) (
(reaction 2)
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&' + &(. - ⇌ &'&(.
(reaction 3)
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Thus, the higher pH in ASW-Mg2+ solutions allows to suggest that the amount of brucite
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dissolved in this solution is more important than that in other seawaters. Mass evolution
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(calculated form ICP-AES) and XRD results hereafter will allow validating this hypothesis.
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Mass evolution
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The quantities of dissolved brucite, precipitated calcium carbonate and the total mass variation
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∆m are determined based on the ICP-AES analysis (Figure 2).
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It can be seen that ∆m evolves (Figure 2a) roughly linearly with the initial volume of seawaters
8
and is coherent with the final pH values (Figure 1). The increase of experiment time from 24 to
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72 hours seems to have minor effect on the solid mass evolution. These ∆m variations are
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resulting from the competition between brucite dissolution (Figure 2b) and calcium carbonate
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precipitation induced by the pH change (Figure 2c). It is important to note that only Mg(OH)2
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and CaCO3 (calcite and aragonite polymorphs) are found by SEM and XRD analysis (detailed in
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next section), other numerically supersaturated phases (ex. magnesite, dolomite and/or relevant
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phases – table S1 and figure S1) are not experimentally detected. Figure 2b indicates that
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are similar in natural and artificial seawaters with or without WSM from
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Pinctada margaritifera. This nacre’s WSM contains mainly two amino acids, Glycine (42.5
17
wt%) and Alanine (21.4 wt%), and )&(.- (20 wt%).47 The amino groups of amino acids
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(Glycine pKa2 9.87; Alanine pKa2 9.7851) and bicarbonate ions might participate to buffer effects
19
in seawater and seem to modify the amount of dissolved brucite (Figure 2b) and the final pH
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(Figure 1). However, this present study is not detailed enough to conclude on the effects of this
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specific WSM on brucite dissolution. The amount of brucite dissolved via reaction 1 depends
22
strongly on the pre-existing Mg2+ ions. The lower the pre-existing Mg2+ ions are present, the
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higher the brucite is dissolved (Figure 2b). Thus, the highest and the highest
2
pH are observed in ASW-Mg2+ solutions (Figure 1). The initial Mg2+ concentration, before
3
adding brucite, in NSW (ICP-AES, 1287 mg.cm-3) is slightly lower than that in ASW (ICP-
4
AES, 1335 mg.cm-3). Therefore, the quantity of dissolved brucite is slightly larger in NSW
5
compared to ASW (with or without WSM). On the other hand, the WSM added in ASW
6
significantly affects the CaCO3 precipitation (Figure 2c):
7
#
8
be partially explained by the presence of )&(.- in WSM (about 20 wt% 47). The WSM addition
9
leds to a weak increase of [)&(.- ] in the ASW+WSM solution. Therefore, ∆m exhibits its largest
10
values in the ASW+WSM solutions (Figure 2a). The lowest ∆m obtained from NSW is the result
11
of the smallest amount of CaCO3 precipitated because of the presence of soluble organic matter
12
(i.e. humic compounds) and phosphate. 20, 24, 25 The lack of these organic matters in the artificial
13
seawaters might be the origin of the larger amount of precipitated CaCO3 (Figure 2c) and ∆m
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(Figure 2a) in these solutions. We also see in Figure 2c that CaCO3 precipitation is the largest in
15
the ASW-Mg2+ (and in ASW+WSM) solution. The amount of CaCO3 precipitated in ASW-Mg2+
16
(without Mg2+) is larger than that obtained in ASW (with Mg2+). This result suggests that the pH
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of the solution (controlled by brucite dissolution) has a significant effect on the rate of CaCO3
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precipitation (reaction 3). So, the ∆m obtained from ASW-Mg2+ (without magnesium ions) is
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larger than that from the ASW solution (containing Mg2+). Many studies have focused on the
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role of magnesium ions on CaCO3 crystal growth in seawaters.1, 17, 26, 52 Berner26 have found that
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magnesium ions at seawater concentration levels appear to have no effect on the rate of crystal
22
growth of aragonite, but a strong retarding effect on that of calcite. In agreement with many
23
previous results from literature, it can be observed from XRD and SEM analysis hereafter that
/01
#
/01 10
>
. The larger amount of precipitated CaCO3 in the ASW+WSM solution might
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both calcite and aragonite polymorphs precipitate after brucite dissolution in ASW-Mg2+ solution
2
in which no Mg ions pre-existed.
3
In this present study, the excessive amount of brucite powder has been chosen hoping that Ca2+
4
in the seawaters could totally react to form solid CaCO3. Then, we assume that
5
6
present in the seawaters. However, the number of moles of )&(.- ion, 6 7 in seawaters is
7
lower than that of Ca2+ and the amount of precipitated CaCO3 should be controlled by the initial
8
[)&(.- ] concentration (with the assumption that )&(.- is the main carbonate form present in the
9
seawaters).
Consequently
10
comparison
between
11
12
13
It can be seen that the amounts of precipitated CaCO3 are partially controlled by [)&(.- ] in the
14
seawaters:
15
849 precipitated in ASW+WSM (Figure 3c) and ASW-Mg2+ (Figure 3d) are slightly
16
larger than those calculated from )&(.- content of seawaters (and of WSM). This phenomenon
17
is also found in ASW and NSW solutions (Figure 3a,b) when the lowest seawater volumes is
18
used (i.e. 50, 200 cm-3). These results suggest that the increase of seawaters pH due to brucite
19
dissolution facilitates the atmospheric CO2(g) absorption into the seawaters via the following
20
equilibria:
21
&( ⇌ &(
