Calcium Carbonate Morphology and Structure in the Presence of

Calcium Carbonate Morphology and Structure in the Presence of Seawater Ions and Humic Acids G. Falini,*,†,‡ S. Fermani,† G. Tosi,† and E. Dinelli‡ Dipartimento di Chimica “G. Ciamician”, Alma Mater Studiorum UniVersita` di Bologna, Via Selmi 2, I-40126 Bologna, Italy, and Laboratori di Scienze Ambientali “R. Sartori”, Alma Mater Studiorum UniVersita` di Bologna, Via Sant Alberto 167, I-48100 RaVenna, Italy

CRYSTAL GROWTH & DESIGN 2009 VOL. 9, NO. 5 2065–2072

ReceiVed March 20, 2008; ReVised Manuscript ReceiVed February 16, 2009

ABSTRACT: The effectiveness of some seawater components, such as magnesium, potassium, sodium, sulfate and chloride ions, and humic acid, in the control of calcium carbonate composition, morphology, and phase distribution was studied. These components were tested singularly, in pairs, and all together. It was observed that magnesium ions phase distribution control is influenced by the presence of other ions and that in the presence of a high content of magnesium ions monohydrocalcite precipitates. Moreover, in the presence of magnesium or potassium ions the calcite crystals show modified rhomohedral morphologies, while the presence of sulfate ions favors their aggregation. Humic acids have an inhibition effect on calcium carbonate precipitation, induce the formation of empty spheres of vaterite, and modify the calcium carbonate phase distribution. The isomorphic substitution of magnesium to calcium in the calcite structure is favored by some seawater ions and enhanced by the presence of humic acids. Introduction Calcium carbonate crystallization is a ubiquitous process in numerous terrestrial and aquatic ecosystems.1 The bulk of calcium carbonate precipitation in seawater is skeletal. Organisms orchestrate crystal nucleation and growth processes (biomineralization) through a sophisticated regulation of an internal chemistry that departs significantly from the “constant ionic medium” of seawater.2 However, the interest in abiotic calcium carbonate formation, and dissolution, in the ocean is increasing, also because of the central role that these reactions will play in the ocean’s response to the increasing partial pressure of carbon dioxide (pCO2) in the atmosphere.3 The most important calcium carbonate polymorphs in seawater reactions are aragonite and calcite.1,4 Aragonite constitutes a major phase of marine sediments, occurring both as marine cement and as the principal skeletal component of many marine taxa.2a,5 Marine calcite usually contains variable amounts (10 to >30 mol %) of MgCO3 in solid solution and is generically termed magnesian calcite.4c,6 Other calcium carbonate phases are of far less abundance and importance. Vaterite is metastable with respect to calcite and aragonite and it is limited to minor biogenic occurrences;7 hydrated CaCO3 phases (mainly monohydrocalcite) are found only in the presence of high concentrations of ions and/or microrganisms.8 Surface seawater is supersaturated with respect to calcite and aragonite, and thermodynamic calculations indicate that this supersaturation should be reduced by a massive precipitation of inorganic aragonite.9 However, the inorganic precipitation is limited and almost absent. This has been attributed to the presence of soluble organic matter interacting onto surface of crystals. It may act by two mechanisms: (i) as an organic sheath, which prevents the nuclei growth by chemical-physical isolation, and (ii) by adsorption on sites where crystal growth would have occurred.10 Many studies on the precipitation of calcium carbonate in seawater have been carried out. The effects of pressure of carbon * Author for correspondence. E-mail: [email protected]. Fax: ++39 051 2099456. † Dipartimento di Chimica “G. Ciamician”. ‡ Laboratori di Scienze Ambientali “R. Sartori”.

dioxide, salinity, ionic strength, main inorganic ions, and organic matter have been investigated in depth.1-8,10 Magnesium ions are the mainly responsible in controlling the kinetics and thermodynamics of the calcium carbonate precipitation. The precipitation of calcite at ambient temperature is both thermodynamically and kinetically favored from solutions containing low amounts of magnesium ions. However, in the presence of magnesium ions in high concentrations relative to calcium ions, the precipitation of aragonite is favored. Interestingly, the density of aragonite is higher than that of calcite. However, magnesium ions are not incorporated in the denser lattice of aragonite as a solid solution, in spite of Mg2+ being a smaller ion than Ca2+. Although the understanding of the mechanism of magnesium incorporation is still limited to preclude definitive statements, this occurs presumably because magnesium has a tightly bound hydration shell, and its removal during crystal growth is energetically unfavorable. Thus, hydrated Mg2+ ions may be incorporated inside the calcite lattice destabilizing its structure and favoring the precipitation of aragonite.6,11 The role of other seawater ions has also been investigated deeply. Sulfate ions are recognized as significant coprecipitating anions in marine biogenic calcites (1 mol % in magnesia calcites), and their effect on calcite precipitation has been investigated in solutions similar to seawater. It has been shown that an increase of sulfate activity reduces the precipitation rate as a function of the saturation state.12 The interactions of organic matter, such as humid acids, with calcium carbonate surfaces in marine systems have been studied primarily because of their effects on the kinetics of carbonate reactions.10d It is known that naturally occurring dissolved organic matter acts as an inhibitor in nucleation and growth of calcium carbonate crystals. However, it is not wellknown which chemical and physical processes are involved, because of the complexity and variability of chemistry and structure of the dissolved organic matter.13 In general, the research on the biotic precipitation of calcium carbonate phases in seawater has mainly addressed kinetic and thermodynamic studies of the processes, and relatively low attention has been paid to structural and morphological features of precipitates.1c,7d The goal of this study is to examine the effectiveness of seawater ions, such as magnesium, sulfate, potassium, sodium

10.1021/cg8002959 CCC: $40.75  2009 American Chemical Society Published on Web 03/24/2009

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Table 1. Molar Ratio Concentrations between the Sea Water Ions and the Calcium Ions Used in the Crystallization Experiments of Calcium Carbonatea 2+

Mg SO42Na+c Cl-c K+ H. acidsd

C1

C2

C3

C4

C5

A(H)b

1 1 12c 12 1 10

2 5 22 22 2 20

3 10 32 32 3 30

5

7

5 5 12 11 1 (40)

42 42 40

80

The concentration of calcium ions and carbonate ions was 3.3 × 10-2 M. b A indicates a mixture containing all the studied seawater ions in concentration close to that one present in seawater. A(H) indicates that humic acids (40 ppm) has been added to the A mixture. c The sodium and chloride ions are also the counterions of carbonate and calcium ions, respectively. d The concentration of humic acids is reported in ppm. a

and chloride, and humid acids in the control of calcium carbonate composition, morphology, and phase distribution. These additives were tested singularly, in pairs, and all together. Results and Discussion In this study the role of the main seawater ions, such as magnesium, potassium, sodium, chloride and sulfate, and humid acids on the control of composition, phase distribution, and morphology of calcium carbonate was investigated. These additives were studied singularly, in pairs, and all together (Table 1). Each ion was used in a range of concentrations around the value of concentration that is present in seawater (Table 1, last column), while humid acids were used up to the concentration (80 ppm) at which there was inhibition of precipitation. Since chemical constraints were not adopted during the experiments (carbonic acid system), the kinetic of precipitation processes was not followed. The precipitates were weighed and characterized by X-ray powder diffraction, Fourier transform infrared spectroscopy, and scanning electron microscopy. Almost stoichiometric amounts of rhombohedral calcite crystals precipitated from the mixture of calcium chloride and sodium carbonate solutions, the control experiment (Table 2, Figures 1, Ctrl, and 7, Ctrl). In the presence of additional ions and/or humic acids the amount of precipitate decreased (Table 2). This effect was clear using humic acids and less obvious using high concentrations of sodium ions. When all the studied sea ions were present, the amount of precipitate was reduced to about 30% with respect to the control. The addition at this mixture of humic acids (40 ppm) reduced the amount of precipitate of about 50% (Table 3, last row). A correlation is observed between the amount of precipitate and the final solution pH; in general, the higher the pH was (around 8.1) the lower the amount of precipitate became (Tables 2 and 3). This observation can be explained considering the increase of salt solubility with the solution ionic strength and the inhibition of crystal nucleation and/or growth due to a poisoning of the calcite crystal nuclei by additional ions (magnesium or potassium) and humic acids.1,6,10,11 Moreover, the precipitation of calcium carbonate phases more soluble than calcite implies an increase of the concentration of carbonate ions in solution, and as consequence of the pH.7b,d In the control experiment pure calcite precipitated. The addition of seawater ions and humic acids changed the phase distribution of the precipitates. The presence of low concentration of magnesium ions ([Mg2+]/[Ca2+] ) 1, 2) provoked the precipitation of magnesian calcite. Aragonite and magnesian calcite were detected precipitating from solutions with [Mg2+]/ [Ca2+] ) 3 and monohydrocalcite precipitated with them using

a solution with [Mg2+]/[Ca2+] ) 5. From a solution with [Mg2+]/ [Ca2+] ) 7 the precipitation of monohydrocalcite and aragonite was detected (Table 2, Figure 1). These phase transitions can be explained considering that the strongly hydrated magnesium ions poison the growth sites on the calcite crystals reducing the amount of precipitated carbonate and calcium ions.11 So, as a consequence of this, solution supersaturation and pH increase (Table 2), as already reported.7b,d This should be the cause of the precipitation of stable calcium carbonate phases having higher solubility than calcite, such as aragonite and monohydrocalcite.8,11 In the presence of sulfate ions, sodium (chloride) ions, potassium ions or low concentrations of humic acids (up to 30 ppm) only calcite was detected in the precipitates (Table 2, Figure 2, SO4, Na, K, H20). When 40 ppm of humic acids were used vaterite (25% w/w) precipitated together with calcite (Table 2, Figure 2, H40), while a complete inhibition of precipitation was observed using 80 ppm of humic acids (Table 2). The detection of vaterite, which is a metastable phase, implies its stabilization by the humic acids. Indeed, it is known that in solution pure vaterite transforms into calcite by a dissolutionprecipitation process.7b Thus, the humic acids may reduce the solubility (dissolution) of vaterite making a layer that envelops (or covers) the crystals. At high concentrations the humic acids may already adsorb on crystal nuclei, blocking their growth and the following precipitation process. Unfortunately, the complex chemistry of humic acids does not allow the formulation of any mechanism on their interaction with calcium carbonate crystals (or nuclei). However, because they are rich with charged functional groups, mainly carboxylates,13 nonspecific electrostatic interactions should occur. Humic acids could also interfere with the precipitation processes by chelating calcium ions in solution. However, this process, which should reduce the solution saturation, does not justify the formation of vaterite that is the most soluble calcium carbonate polymorph. Moreover, it does not agree with reported studies showing that inhibition on calcium carbonates precipitation arises from a reduction of growth sites on crystal surface.14 The FTIR spectra of these precipitates did not show any clear absorption band that could be related to the presence of organic material. In them only the characteristic absorption bands of the calcium carbonate phases were clearly observed (Figure 5).15 This is expected considering the low starting amounts of humic acids used. Thus, a low amounts of humic acids should be enough to stabilize vaterite. The effect of pairs of additional ion species was studied using concentrations similar to those ones present in seawater (Table 1, last column). In the presence of magnesium ions as unique additive and at a [Mg2+]/[Ca2+] ) 5 the precipitation of aragonite (80% w/w), magnesian calcite (10% w/w) and monohydrocalcite (10% w/w) was detected (Table 2). The presence of a second additive (ions or humic acids) with magnesium ions changed the phase distribution of the precipitates (Table 3). In the copresence of sulfates ions ([SO42-]/[Ca2+] ) 5) magnesian calcite (15% w/w) and aragonite (85% w/w) precipitated (Table 3, Figure 3, Mg/SO4); in the copresence of sodium ([Na+]/[Ca2+] ) 12) or potassium ions ([K+]/[Ca2+] ) 1) aragonite and monohydrocalcite (20% w/w and 10% w/w, respectively) precipitated (Table 3, Figure 3, Mg/K); and in the copresence of humic acids (HA ) 40 ppm) magnesian calcite (20% w/w), vaterite (10% w/w) and aragonite (70% w/w) precipitated (Table 3, Figure 3, Mg/H). The comparison of these data indicates that the presence of sulfate ions slightly favors the precipitation of magnesian calcite with respect to monohydrocalcite. This observation has been reported in a different system, and explained considering the formation of strong ionic

Calcium Carbonate Morphology and Structure

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Table 2. Crystallization Experiments of Calcium Carbonate in the Presence of Single Sea Water Ion Specie or Humic Acids at the Concentrations (C1-5) Reported in Table 1a C1

C2

pHb nmc ccpd - %

ucv

pH

nm

C3

ccp - %

ucv

pH

nm

C4

ccp - %

Mg2+

7.7 0.96 C - 100

366.2 7.7 0.93 C - 100 365.8 8.0 0.95 A - 32 C - 68

SO42Na+(Cl-) K+ H. acids

7.7 7.7 7.7 7.8

367.9 367.8 367.9 367.7

0.97 0.93 0.95 0.82

C C C C

-

100 100 100 100

7.8 7.8 7.7 7.8

0.96 0.88 0.93 0.81

C C C C

-

100 100 100 100

367.8 367.7 367.8 367.5

7.8 7.9 7.8 7.9

0.95 0.83 0.89 0.74

C C C C

-

100 100 100 100

ucv

pH

nm

ccp - %

226.2 8.0 0.87 A - 80 364.6 C - 10 M - 10 367.9 e e e 367.6 8.0 0.80 C 367.8 e e e 367.6 8.1 0.70 C - 75 V - 25

C5 ucv 225.9 360.2 242.0 e 367.9 e 367.8 125.2

pH

nm

ccp - %

ucv

8.0 0.90 A - 70 M - 30

225.9 242.0

e e e e e e 8.2 f

e e e f

e e e f

a For each crystallization experiment the final pH of the solution (pH), the normalized mass of the precipitate (nm), the calcium carbonate phases (ccp) detected in the precipitate and their weight percentage (%), and the unit cell volume (ucv/Å3) of each phase are reported. b In the control experiment, in which only calcium chloride and sodium carbonate were mixed, the final pH was equal to 7.7. c Mass of the precipitate normalized to the mass of precipitate obtained from the control experiment (324 mg). d A, C, M, and V indicate aragonite, calcite, monohydrocalcite and vaterite, respectively. Their unit cell volumes are 226.17 Å3, 367.85 Å3, 242.57 Å3 and 125.41 Å3, respectively. e These concentrations were not studied. f No precipitation was observed after a crystallization time of 5 days.

Figure 1. X-ray powder diffraction patterns of calcium carbonate precipitated in the presence of different concentrations of magnesium ions. (Ctrl) pure calcite; (Mg1), (Mg3), (Mg5), and (Mg7) indicate [Mg2+]/[Ca2+] are equal to 1, 3, 5, and 7, respectively. (C), (A), and (M) indicate the main diffraction peaks of calcite, aragonite, and monohydrocalcite, respectively. The diffraction patterns were displaced along the vertical axis that reports intensity in arbitrary units of counts.

couples among sulfate and magnesium ions; this reduces the activity of magnesium ions in solution favoring the precipitation of magnesian calcite.16 In the presence of humic acids the deposition of vaterite could be explained taking in account the adsorption of this organic material on the crystals, as discussed above. The coupling of humic acids with sulfate or sodium ions provoked the deposition of calcite only (Table 3), while in the presence of potassium ions a small amount of vaterite (10% w/w) was also detected (Table 3, Figure 3, H/K). The absence of vaterite in the presence of sodium or sulfate ions may suggest that these ions reduce the strength of the electrostatic interactions between humic acids and crystals, which should prevent the

Figure 2. X-ray powder diffraction patterns of calcium carbonate precipitated in the in the presence of different ions: (SO4) sulfate, (Na) sodium, and (K) potassium. (H20) and (H40) are the diffraction patterns of the calcium carbonates precipitated in the presence of 20 ppm and 40 ppm of humic acids, respectively. (C) and (V) indicate the main diffraction peaks of calcite and vaterite. The diffraction patterns were displaced along the vertical axis that reports intensity in arbitrary units of counts. The species concentrations were the ones reported in the last column of Table 1.

dissolution of vaterite. Thus, during these crystallization experiments the metastable vaterite should transform in calcite.7b Also in this case the amount of adsorbed organic material was very low, because it was not revealed by the FTIR spectra of the precipitates (data not shown). The addition of two ions among sodium, potassium, and sulfate provoked only a reduction of the amount of precipitated calcite, which was always the unique phase detected (Table 3, Figure 3SO4/K). When all the studied seawater ion species were mixed together (Table 1, last column) the precipitation of magnesian

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Table 3. Crystallization Experiments of Calcium Carbonate in the Presence of Pairs of Sea Water Ions or a Sea Water Ion and Humic Acids at the Concentrations Reported in the Last Column of Table 1a Mg2+ b

c

ucv 226.2 359 226.2 242.1 226.2 242.0 226.2 356 125.2 226.9 357 241.9

pH

nm

ccp - %

SO42-

7.8

0.87

Na+ (Cl-)

7.7

0.84

K+

7.8

0.88

H. acids

8.0

0.61

Ae

8.0

0.72

A - 85 C - 15 A - 80 M - 20 A - 90 M - 10 A - 70 C - 20 V - 10 A - 30 C - 15 M - 55

Na+ (Cl-)

SO42-

d

pH

nm

ccp - %

ucv

7.9

0.76

C - 100

367.8

7.8

0.86

C - 100

8.1

0.56

C - 100

K+

pH

nm

ccp - %

ucv

367.9

7.9

0.79

C - 100

367.8

367.7

8.2

0.50

C - 100

367.6

A(H)e

8.2

0.46

A - 80 C - 15 V-5

226.5 354 125.2

pH

nm

ccp - %

ucv

8.2

0.55

C - 90 V - 10

367.8 125.2

a For each crystallization experiment the final pH of the solution (pH), the normalized mass of the precipitate (nm), the calcium carbonate phases (ccp) detected in the precipitate and their weight percentage (%), and the unit cell volume (ucv/Å3) of each phase are reported. In the last row the data concerning the crystallization experiments in which all the studied sea water ion species were used without (A) and with 40 ppm hunic acids (A(H)) are reported. b In the control experiment, in which only calcium chloride and sodium carbonate were mixed, the final pH was equal to 7.7. c Mass of the precipitate normalized to the mass of precipitate obtained from the control experiment (324 mg). d A, C, M and V indicate aragonite, calcite, monohydrocalcite and vaterite, respectively. Their unit cell volumes are 226.17 Å3, 367.85 Å3, 242.57 Å3 and 125.41 Å3, respectively. e The concentrations of the seawater ions are those ones reported in the last column of Table 1.

Figure 4. X-ray powder diffraction patterns of calcium carbonate precipitated in the presence of all the studied seawater ions without (A) and with humic acids (A(H)). The concentrations are those ones reported in the last column of Table 1. (C), (A), (V), and (M) indicate the main diffraction peaks of calcite, aragonite, vaterite, and monohydrocalcite, respectively. The diffraction patterns were displaced along the vertical axis that reports intensity in arbitrary units of counts.

Figure 3. X-ray powder diffraction patterns of calcium carbonate precipitated in the presence of pairs of seawater ion species or a specie of seawater ions and humic acids at the concentrations reported in the last column of Table 1. (SO4/K) sulfate and potassium ions; (H/K) humic acids and potassium ions (Mg/H); magnesium ions and humic acids; (Mg/SO4) magnesium and sulfate ions (Mg/K) magnesium and potassium ions. (C), (A), (V), and (M) indicate the main diffraction peaks of calcite, aragonite, vaterite, and monohydrocalcite, respectively. The diffraction patterns were displaced along the vertical axis that reports intensity in arbitrary units of counts.

calcite (15% w/w), aragonite (30% w/w), and monohydrocalcite (55% w/w) was detected (Table 3, Figure 4A). The high content of monohydrocalcite in the precipitate is consistent with what reported for induced calcium carbonate precipitation in natural seawater.8a The addition of humic acids (40 ppm) at the solution containing all the seawater ions changed the phase distribution

of the precipitate. A mixture of aragonite (80% w/w), vaterite (5% w/w), and magnesian calcite (15% w/w) was detected (Table 3, Figure 4, A(H)). Thus, humic acids favor the precipitation of vaterite and rule against the precipitation of monohydrocalcite, as already observed in the previous experiments. The X-ray powder diffraction patterns of the precipitates from the solutions containing all the seawater ions and the humic acids showed a marked broadening of the diffraction peaks (Figure 4) with respect to that one observed in diffraction peaks from the other precipitates (Figures 1-3). This indicates a reduction in the crystallinity of mineral phases,17 which could be related to the high number of additives in the precipitation solutions. The X-ray diffraction patterns and the FTIR spectra of the precipitates did not show any signal that could be associated with the presence of amorphous calcium carbonate.18 Thus, this phase is absent, present in low amounts, or has been transformed into crystalline phases during the precipitates’ handling.19


Figure 5. FTIR spectra of calcium carbonate precipitated in the absence of additives (Ctrl), in the presence of 40 ppm of humic acids (H40) in the presence of all the considered seawater ions without (A) and with (A(H)) humic acids. The peaks at 1485, 1420, and 1408 cm-1 correspond to the ν3 absorption band of carbonate ions. The peaks at 876 and 713 cm-1 correspond to ν2 and ν4 absorption bands of calcite, respectively. The peak at 860 cm-1 (ν2) and 745 cm-1 (ν4) are diagnostic of aragonite and vaterite, respectively. All spectra were normalized to the peak intensity of the band ν3 of carbonate and were displaced along the vertical axis that reports absorption in arbitrary units. The concentrations of the species were those ones reported in the last column of Table 1.

The isomorphic substitution of ions in the calcium carbonate crystalline structures was evaluated by lattice parameters. Those ones of aragonite, monohydrocalcite, and vaterite were almost constant (Tables 2 and 3). Only the lattice parameters of the magnesian calcite changed, as a function of the additive(s) in solution. In general, the isomorphic substitution of magnesium to calcium in calcite increased increasing the magnesium content in solution (Figure 6). When solutions with [Mg2+]/[Ca2+] equal to 1, 2, and 3 were used the precipitated magnesian calcite had an isomorphic substitution of magnesium to calcium of 1.2, 2.6, and 4.2, atom %, respectively. When a [Mg2+]/[Ca2+] ) 5 was used an isomorphic substitution of magnesium to calcium of about 7 atom % was detected in the magnesian calcite structure. This isomorphic substitution was increased by the copresence of sulfate ions and humic acids, which provoked an isomorphic substitution to 8 and 10 atom %, respectively (Figure 6). A plausible explanation of this observation is that these additives weaken the hydration sphere of magnesium ions favoring their entrance into structure of calcite.11 In the presence of all the studied ions the magnesium substitution to calcium in the structure of calcite was about 9 atom % and increased to about 13 atom % when humic acids were added (Figure 6). This effect due to of humic acids may be explained considering that they are rich with carboxylate groups which may contribute to a reduction of the hydration of magnesium ions. Moreover, the inhibition of humic acids versus the precipitation of calcium carbonates should lead to high levels of supersatutation, in agreement with final pH of solutions. In this condition, probably through the formation of an amorphous precursor, the precipitation of magnesian calcite with high content of magnesium is favored.20 Interestingly, the observed values of isomorphic substitution of magnesium to calcium ions in the calcite structure


Figure 6. Isomorphic substitution (atom %) of magnesium to calcium in the calcite structure as function of the concentration of magnesium ions in solution. (() Only magnesium ions were present in solution: Mg1, Mg2, Mg3 and Mg5 indicate [Mg2+]/[Ca2+] molar ratio equal to 1, 2, 3 and 5, respectively. (9) Mg/SO4: Sulfate and magnesium ions were present in solution. (2) Mg/H: Magnesium ions and humic acids (40 ppm) were present in solution. (×) all the considered seawater ions (A) were present in solution; (x’s with vertical line) all the considered seawater ions and 40 ppm humic acids (A(H)) were present in solution. The concentrations of the species are those ones reported in the last column of Table 1.

are lower than those ones reported in magnesian calcites precipitated from natural seawater.1c,4,21 Calcium carbonate precipitated as single crystals and crystalline aggregates whose morphology and shape were modified by the presence of seawater ions and humic acids (Table 4, Figures 7 and 8). These forms showed a wide range of sizes that did not allow any clear correlation with the kind of additive which was present in the crystallization solution. Rhombohedra of calcite showing only {104} faces precipitated in the presence of sodium or sulfate ions, as observed in the control experiment (Figure 7, Ctrl). However, in the presence of sulfate ions the crystals were aggregated (Figure 7, SO4); the higher ion concentration was the higher aggregation occurred, as previously reported.16 The presence of potassium ions stabilized crystalline {hk0} faces on the {104} rhombohedral calcite crystals (Figure 7K). Magnesian calcite crystals showed also a modified {104} rhombohedral morphology in which {11j0} crystalline faces were observed (inset in Figure 7, Mg3).6a,b The stabilization of new crystalline faces on calcite by magnesium ions (specific) and potassium ions (nonspecific) implies a preferential interaction of these ions with specific crystal plane(s) of the calcite structure, otherwise unstable. The specificity of these ions is probably related to their ionic radius, charge density, and hydration sphere.6a,b,22 Aragonite appeared as spherical aggregates formed by small needle-like crystals, while monohydrocalcite appeared as spherulites (insets in Figure 7, Mg7). Vaterite, which precipitated only in the presence of humic acids, formed void globular spherulites (Figure 7, H40’). These spherulites were polycrystalline material (optical microscope observation, data not reported) that showed a wide distribution of sizes and a wall thickness of about 1 µm. Similar shapes for vaterite have been already observed in the presence of templating carboxylated polypeptides or in crystallization processes in which an amorphous precursor quickly converted in the crystalline phase.23 In the presented experiments a templating effect due to humic acids cannot be excluded, although they are present in small amounts (40 ppm).

sp. agg. n. m. rhomb. void sph. {104} {hk0}

A(H)e

A C(13) v

a. spher. m. rhomb. spher.

prismatic sp. agg. n. spher. A M prismatic {104} {11j0} sp. agg. n. m. rhomb. A C(8)

C C C

rhomb. rhomb. rhomb.

{104} {104} {104}

C C

rhomb. rhomb.

{104} {104}

C V A M

m. rhomb void sph. sp. agg. n. spher.

prismatic

A C(10) V

shape phase morph. shape phase morph. shape phase morph. shape phase

prismatic

morph.

{104} {104} {104} {104} {hk0} {104} prismatic {104} {11j0}

shape

rhomb.b rhomb. ag.c rhomb. m. rhomb.e rhomb. void sph.d sp. agg. n.f m. rhomb. spher. a. spher. m. rhomb. spher. Ae

Mg2+

phase

C C C C C V A C(7)g M A C(9) M

K+ Na+ (Cl-)

a These additives were tested singularly, in pairs and all together. Their concentrations are reported in the last column of Table 1. b rhomb. indicates a rhombohedral shape. c rhomb. ag. indicates a shape generated by the aggregation of rhomhedra. d void sph. indicates spherical shapes which are void. e m. rhomb. indicates a modified rhombohedral shape. f sp. agg. n. indicates spherical shapes generated by the aggregation of needle-like crystals. g The value in parentheses reports the isomorphic substitution of magnesium to calcium in the magnesian calcite structure. It is expressed as atom %.

morph. H Ctrl SO42Na+ (Cl-) K+ H. acids

In the copresence of sulfate and magnesium ions the crystals of aragonite and magnesian calcite aggregated forming compact spherulites (Figure 8, Mg/SO4), similarly to what observed in the presence of only magnesium ions, at the same concentration (Figure 7, Mg5). Also the crystals formed in the presence of potassium ions aggregated when sulfate ions were present (Figure 8, SO4/K). These crystals conserved the modified rhombohedral morphology, due to the presence of potassium ions. However, the effect of potassium ions as crystal morphology modifiers was lost when a high concentration of sulfate ions was used ([SO42-]/[K+] ) 10) (data not shown). In the copresence of potassium and sodium ions the calcite crystals conserved the morphology observed in the presence of potassium ions only (Figures 7, K and 8, Na/K). In the copresence of magnesium and potassium ions the presence of distinct aragonite needles and monohydrocaltes sperulites was observed (Figure 8, Mg/K), while in the copresence of humic acids and

SO42-

Figure 7. Scanning electron micrographs calcium carbonates precipitated in the presence of single seawater ions or humic acids. (Ctrl) pure calcite, the control experiment. (Mg3), (Mg5), and (Mg7) calcium carbonate precipitated from a solution with [Mg2+]/[Ca2+] equal to 3, 5, and 7, respectively. (K) and (SO4) calcium carbonate precipitated in the presence of potassium ions and sulfate ions, respectively. (H40) and (H40′) calcium carbonate precipitated in the presence of 40 ppm humic acids; two different magnifications are reported. In the inset in (Mg3) a magnesian calcite crystal exposing {11j0} faces is shown. The insets in (Mg7) show needle-like crystals of aragonite (A) and spherulites of monohydrocalcite (M). In the inset in (K) a calcite crystal showing {hk0} faces is shown and therein the arrow indicates the c-axis of calcite. In the inset in (H40′) empty spherulites of vaterite are shown. The concentrations of the species were those ones reported in the last column of Table 1.

Falini et al. prismatic {104} {11j0}


Table 4. Effects of Seawater Ions, Magnesium, Potassium, Sodium and Sulphate, and Humic Acids on Composition, Shape, Morphology and Phase Distribution of Calcium Carbonatea

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Figure 8. Scanning electron micrographs of calcium carbonate precipitate in the presence of pairs of seawater ions: (Mg/SO4) sulfate and magnesium ions; (SO4/K) potassium and sulfate ions; (Na/K) potassium and sodium ions; (H/K) potassium ions and humic acids; (Mg/K) potassium and magnesium ions. (A) calcium carbonate precipitated from all the seawater ions. (A(H)) calcium carbonate precipitated in the presence of all seawater ions plus humic acids. The inset in (K/H) shows a empty spherulite of vaterite. The insets in (A) and (A(H) show a low magnification view of the corresponding precipitated aggregates. The concentrations of the species were those ones reported in the last column of Table 1.

potassium ions void spherulites of vaterite and modified rhombohedral crystals of calcite were observed (Figure 8, H/K). The copresence of humic acids and magnesium ions provoked the precipitation of aggregates similar to those ones observed in the presence of all the seawater ions and humic acids (Figure 8, A(H)). The presence of all the studied seawater ions provoked the precipitation of aggregates, in which it was possible to observe the presence aragonite needles and monohydrocaltes sperulites localized in different regions (Figure 8, A). When humic acids were also used the aggregation of the calcium carbonate particles increased and it was not possible anymore to discriminate among phases (Figure 8, A(H)). Interestingly, the calcium carbonate precipitated from seawater appears as cement in which the different crystalline phases are closely packed.21 The results of the calcium carbonate crystallization experiments in the presence of seawater ions and humic acids are summarized in the Table 4. It reports shape, aggregation status, and phase distribution of the precipitates obtained in the presence of a single additive, a couple of additives, and all the additives. From its analysis it can be seen that the effectiveness of individual ions in the control of calcium carbonate precipitation is modulated by the presence of other ions. Moreover, some ions have a key role in controlling the characteristics of the precipitate. Magnesium ions are the controllers of crystal morphology and phase distribution, potassium ions control crystal morphology, and sulfate ions act as controllers of crystal

aggregation. Also the humic acids have an important role inhibiting the precipitation process and controlling the phase distribution of calcium carbonate. Interestingly, the isomorphic substitution of magnesium to calcium in the magnesian calcite structure is influenced by the presence of the seawater ions and humic acids, and reaches the highest value when all the additives are present. Conclusion This study shows how seawater ions and humic acids can control calcium carbonate composition, morphology, and phase distribution. The control of magnesium ions over the phase distribution is modulated by the presence of humic acids, sodium or sulfate ions. The morphology and aggregation of calcite crystals are controlled by magnesium ions, potassium ions, and sulfate ions. Humic acids provoke inhibition of the calcium carbonate precipitation, favor the formation of void spherulite of vaterite, and modify the calcium carbonate phase distribution. The isomorphic substitution of magnesium to calcium in the calcite structure is favored by the presence of sulfate ions and humic acids. Experimental Procedures The chemicals used, CaCl2, MgCl2, Na2CO3, NaCl, KCl, Na2SO4, HCl, and NaOH, were analytically pure (Sigma-Aldrich) and the solutions were made using Milli-Q water. Humic acids, as sodium salt,

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were purchased from Sigma. 3.3 × 10-2 M CaCl2 and 3.3 × 10-2 M Na2CO3 solutions were used as stock solutions. The solutions containing additional solutes were prepared adding the salts to the CaCl2 or Na2CO3 stock solution. The additions were made avoiding the formation of precipitates. Humic acids were added to the carbonate solution. The final pH of the solutions was adjusted to 8.2 by addition of 0.1 M NaOH or 0.1 M HCl solution. The experiments were carried out pouring 100 cm3 of the CaCl2 solution into the same volume of the Na2CO3 solution. After mixing the final solution was kept for five days in a closed bottle at room temperature. The formed precipitate was filtered and washed with Milli-Q water, then it was dried at 105 °C for 12 h and finally weighed. The surnatant pH was measured. The X-ray powder diffraction (XRD) patterns were recorded using a X’Celerator diffractometer (PANalytical) with Cu KR radiation and a Ni filter. In the FTIR analysis each powdered sample (approximately 0.1 mg) was mixed with about 10 mg of anhydrous KBr. The mixtures were pressed into 7 mm diameter discs. Pure KBr discs were used as background. The analysis was performed at 4 cm-1 resolution using a Nicolet 380 FT-IR spectrometer. Scanning electron micrographs (SEM) of the samples were recorded by a Philips 515 scanning electron microscope. The samples were glued on aluminum stubs and gold sputtered prior the observations. A Rietveld program (Quanto) for quantitative phase analysis of polycrystalline mixtures from powder diffraction data was used to quantify calcium carbonate phases and to evaluate the unit cell parameters.24a The isomorphic substitution of magnesium to calcium in the calcite structure was estimated from the unit cell parameters according to Goldsmith et al.24b

Acknowledgment. We thank the Alma Mater Studiorum Universita` di Bologna (Funds for Selected Topics), the Ministero dell’Istruzione, dell’Universita` e della Ricerca (MIUR), and Consorzio Interuniversitario di Ricerca Chimica dei Metalli nei Sistemi Biologici (CIRCMSB) for financial support.

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Calcium Carbonate Morphology and Structure in the Presence of

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