Next Article in Journal
Deep Structure of Nanling-Xuancheng Ore District, Eastern China: Insights from Integrated Geophysical Exploration
Next Article in Special Issue
Isolation and Characterization of Indigenous Ureolytic Bacteria from Mindanao, Philippines: Prospects for Microbially Induced Carbonate Precipitation (MICP)
Previous Article in Journal
Wastes in Underground Coal Mines and Their Behavior during Mine Water Level Rebound—A Review
 
 
Search for Articles:
Title / Keyword
Author / Affiliation / Email
Journal
Article Type
 
 
Advanced Search
 
Section
Special Issue
Volume
Issue
Number
Page
 
Journals
Minerals
Volume 13
Issue 12
10.3390/min13121497
minerals-logo
Submit to this Journal Review for this Journal Propose a Special Issue
► Article Menu

Article Menu

share Share announcement Help format_quote Cite question_answer Discuss in SciProfiles thumb_up
...
Endorse textsms
...
Comment
first_page
settings
Order Article Reprints
Font Type:
Arial Georgia Verdana
Font Size:
Aa Aa Aa
Line Spacing:
Column Width:
Background:
Article

Calcium Carbonate Precipitation Behavior in the System Ca-Me2+-CO3-H2O (Me2+ = Co, Ni, Cu, Fe): Ion Incorporation, Effect of Temperature and Aging

by
Oleg S. Vereshchagin
*,
Irina A. Chernyshova
,
Maria A. Kuz’mina
and
Olga V. Frank-Kamenetskaya
Institute of Earth Sciences, Saint Petersburg State University, Universitetskaya Emb. 7/9, 199034 St. Petersburg, Russia
*
Author to whom correspondence should be addressed.
Minerals 2023, 13(12), 1497; https://doi.org/10.3390/min13121497
Submission received: 17 October 2023 / Revised: 17 November 2023 / Accepted: 27 November 2023 / Published: 29 November 2023
(This article belongs to the Special Issue Biomineralization and MICP in Wastewater, Reclaimed Water and Seawater)
Browse Figures
Versions Notes

Abstract

:
Crystalline calcium carbonates (CCCs) are among the most widespread minerals on the Earth’s surface and play a crucial role in the global carbon cycle, heavy metal sorption and incorporation. Among the numerous factors that influence the precipitation of CCCs from solution, the most determinant are the presence of additives in the mineral-forming medium, temperature, and crystallization time (aging time). The current work fills the gaps in the study of calcium carbonate crystallization from heavy metal (Me2+ = Co, Ni, Cu Fe)-containing solutions (Me2+/Ca 0.005–1.600) at different temperatures (3 and 23 °C) and aging times (21–158 days). The resulting precipitates were studied using optical and scanning electron microscopy, powder X-ray diffraction and energy-dispersive X-ray spectroscopy. Three crystalline calcium carbonates (synthetic analogues of calcite, aragonite and monohydrocalcite), as well as amorphous carbonate (AC), were found in the resulting precipitates. Temperature and aging time have a considerable effect on the phase composition, morphology and heavy metal content in CCCs. Low temperature (3 °C) and short aging times are generally favorable for the formation of monohydrocalcite and amorphous carbonate, while calcite tends to form at a higher temperature (23 °C) and in long-term experiments. Heavy metals can be incorporated into the calcite/monohydrocalcite crystal lattice in sufficient amounts, while aragonite can host a very small amount of Me2+ (or none). Calcite can concentrate Co (up to ~0.25 atoms per formula unit (apfu)) and Ni/Cu (up to ~0.05 apfu), while its Fe content is very close to the detection limits. Calcite precipitated at a higher Me2+/Ca ratio and temperature (23 °C) contains less Me2+ compared to calcite precipitated at a lower Me2+/Ca ratio and temperature (3 °C). Monohydrocalcite can host up to ~0.1 apfu of Co/Ni/Cu with no detectable preference for Me2+. The amount of Me2+ in monohydrocalcite decreases as aging time or temperature increases. It is worth noting that AC is the main carrier of heavy metals in the system being investigated and it should be considered the main host phase in heavy metal adsorption from aqueous solutions. The results obtained can be used to solve environmental issues and in mineral resource management.

1. Introduction

Crystalline calcium carbonates (CCCs) are among the most widespread minerals on the Earth’s surface (e.g., [1]), in particular, in oceans (e.g., [2]). They are the most commonly encountered biominerals (e.g., [3]) and play a crucial role in the global carbon cycle (e.g., [4]). A substantial part of the anthropogenic CO2 has been absorbed by the oceans, which has changed carbonate ion availability and led to ocean acidification. Acidification affects the biogeochemical cycles of elements [5], including those of minor elements (e.g., [6]). These cycles are known to be CCC-type dependent, as different CCCs have different capacities to host these elements.
Five CCC minerals are known: three anhydrous polymorphs of CaCO3 (calcite, aragonite and vaterite) and two hydrated phases—CaCO3·H2O (monohydrocalcite) and CaCO3·6H2O (ikaite). Calcium is six-coordinated in the calcite crystal structure (e.g., [7]), eight-coordinated in ikaite, monohydrocalcite and nine-coordinated in aragonite crystal structures (e.g., [8,9]). This indicates the different ability of these minerals to capture “foreign” cations.
Calcite should be the only stable phase under the Earth’s surface conditions (e.g., [10,11,12]), but aragonite, vaterite, monohydrocalcite and ikaite are also present both as abiogenic phases and biominerals (e.g., [11]).
Both biogenic and abiogenic CCCs could also serve as a chemical archive of the conditions in which they were formed (e.g., [13,14]) and provide one of the most comprehensive geochemical records of the environmental and climatic conditions occurring at the time of mineral formation (e.g., [15,16]). However, most studies are focused on calcite (e.g., [17]), there is limited information available on the mechanisms of trace element incorporation into aragonite (e.g., [18,19,20]), and there are almost no data concerning hydrated calcium carbonates (monohydrocalcite and ikaite). Among the numerous factors influencing CCC precipitation, the most determinant are probably the presence of (cation) additives in the mineral-forming medium, temperature and crystallization time (aging time) (e.g., [21,22]).
Heavy metal content in rivers and lakes all over the world has dramatically increased over the past 50 years [23], making heavy metal pollution of surface water a global environmental problem (e.g., [24]). Heavy metals generally include metals with relatively high densities, which is the case with the first-row transition metals. It is important to note that most of the first-row transition metals are toxic for a number of living organisms (e.g., [25,26,27]). Cation additives have a considerable effect on CCC growth and have attracted the attention of researchers for decades (e.g., [12,21,28]). CCC crystallization in Co- (e.g., [29,30]), Ni- (e.g., [20,31]), Cu- (e.g., [21,32,33]) and Fe-rich (e.g., [34,35,36]) systems was studied previously. It was shown that the presence of some heavy metals (e.g., Co, Fe) violates the normal sequence of calcium carbonate crystallization (e.g., [20,30,37,38]). However, it was found that the effect of cations is very different due to the possibility of their incorporation into the structure of CCCs/the sorption on CCC surface, as well as different behavior in aqueous solutions (e.g., [20,21,29]). In addition, CCCs typically form small crystals (<5 µm) and have complex intergrowths (e.g., spheres) in heavy metal-rich systems, making it impossible to obtain direct data on their chemical composition using X-ray spectroscopy (EDX/WDX; typical spot size ~3 µm) or inductively coupled plasma mass spectrometry (LA-ICP MS; typical spot size ~50–70 µm). The main method that could reveal the occurrence of impurities in CCCs is the measurement of the unit cell parameters of the phases under consideration, but neither the Inorganic Crystal Structure Database (ICSD) nor the Powder X-ray Diffraction Database contains information on the unit cell parameters of CCCs enriched with heavy metals.
Temperature-dependent studies are especially important when biomineral formation and surface processes are modeled. The precipitation of hydrated CCCs (monohydrocalcite and ikaite) tends to be favored by low (<7 °C) temperatures (e.g., [11]), but anhydrous CCCs (e.g., aragonite) could also form instead of calcite under such conditions (e.g., [39]). However, the influence of low temperatures (<25 °C) has received little attention so far since the study of such systems requires a long crystallization time (e.g., [22]).
Aging time is crucial for CCC formation (e.g., [22,40]) and their chemical composition (e.g., [22,30,40]). It has been found that metastable CCC (e.g., vaterite/aragonite/monohydrocalcite) formation is favored by short aging time (<1 day, e.g., [30,40]), while long crystallization time can lead to the formation of both stable (calcite) and metastable (monohydrocalcite) CCCs (e.g., [22,30,40]). Thus, additional experiments are needed to establish the effect of crystallization time on the phase composition of sediments and the chemical composition of CCCs.
The current work fills the gaps in the study of calcium carbonate crystallization from heavy metal (Me2+ = Co, Ni, Cu, Fe)-containing solutions at different temperatures and aging times. Practically, the issues of incorporation of Me2+ into the CCC crystal lattice, temperature effect, and aging time effect were considered.

2. Materials and Methods

2.1. Synthesis

Syntheses were carried out by a precipitation method from aqueous solutions at different temperatures, Me2+/Ca ratios and aging times. Changes in the unit cell parameters and unit cell volume of the CCCs were used to identify the incorporation of heavy metals in the CCCs [22].
Synthesis was conducted in a closed volume of an aqueous solution by the rapid addition (under constant and vigorous stirring) of CaCl2 (99%, Vekton (Saint-Petersburg, Russia)) + CoCl2·6H2O (99%, Vekton)/NiCl2·6H2O (99%, Vekton)/CuCl2·2H2O (99%, Vekton)/FeSO4·7H2O (99%, Vekton) solution to a 20 mM Na2CO3 (99%, Vekton)/NaHCO3 (99%, Vekton)/NaOH (99%, Vekton) solution. The Na2CO3/NaHCO3/NaOH ratio was used to change the pH of the solution. The Ca/CO3 ratio was set as 0.5 (Table 1 and Table 2). Two series of syntheses were performed at different temperatures and Me2+/Ca ratios (Me2+ = Co, Ni, Cu, Fe): 55 syntheses at a temperature of 3 °C (Me2+/Ca 0.005–1.600, Table 1) and 26 syntheses at a temperature of 23 °C (Me2+/Ca 0.005–1.300, Table 2). Syntheses were conducted for 21–158 days (Table 1 and Table 2) and were stopped as the mother solution became clear and a layer on the bottom of the beaker was formed. The resulting precipitates were filtered, washed with deionized water several times and dried at room temperature in the air for 24 h.

2.2. Characterization Methods

The resulting precipitates (after washing and drying) were studied by optical microscopy, followed by powder X-ray diffraction (PXRD). After that, precipitates were studied by scanning electron microscopy (SEM) and energy-dispersive X-ray spectroscopy (EDX).
The main method that can reveal the occurrence of impurities in CCCs is the measurement of the unit cell parameters of the phases under consideration.
Optical microscopy. The morphology and optical properties (e.g., color, luster) of the synthesized products were studied using a DM 2500P (Leica, Germany) polarizing light microscope.
Powder X-ray diffraction. PXRD was used to identify the phase composition of the precipitation and determine unit cell parameters (UCPs) of newly formed phases. The regularities of change in UCPs were used in the study of the possibility of heavy metal incorporation into CCCs. PXRD patterns of samples were recorded on a Miniflex II diffractometer (Rigaku, Japan) with CoKα/CuKα radiation (2-theta range 5–110°, velocity 2°/min, step size 0.02°, Bragg-Brentano geometry). Phase identification was carried out using the PDF-2 Database. Subsequent quantitative phase content analyses of precipitates were carried out using the full-profile method using Bruker TOPAS v. 5.0 software. Unit cell parameters (for samples with a single-phase content >90%) were determined using an Ultima IV high-resolution diffractometer (Rigaku, Japan) with CuKα radiation (2-theta range 5–120°, velocity 1°/min, step size 0.02°, Bragg-Brentano geometry; PSD D-Tex Ultra detector) with internal standard (Ge).
Scanning electron microscopy and energy-dispersive X-ray spectroscopy. SEM, together with EDX spectroscopy, was used to study the morphology and chemical composition of newly formed phases by means of an S-3400N (Hitachi, Japan) SEM equipped with an AzTec Energy X-Max 20 EDX spectrometer (Oxford Instruments, UK) with the following parameters: 20 kV accelerating voltage, 1 nA beam current and 30 s data collection time (excluding dead time).

3. Results

3.1. Resulting Phase Composition and Crystal (Micro) Morphology

PXRD revealed four crystalline carbonates (synthetic analogues of calcite (space group (SpGr) R-3c, ideally CaCO3), aragonite (SpGr Pmcn, ideally CaCO3), monohydrocalcite (SpGr P31, ideally CaCO3·H2O), malachite (SpGr P21/b, ideally Cu2(CO3)(OH)2)) and two oxides (synthetic analogues of goethite (SpGr Pbnm, ideally FeO(OH)) and magnetite (SpGr Fd3m, ideally FeFe2O4)) in resulting precipitates (Table 3 and Table 4). Despite the long crystallization time, an amorphous carbonate (AC) was typically present (Table 3 and Table 4).
Optical and SEM–EDX studies showed that the studied crystalline carbonates predominantly form spherical or dumbbell-like aggregates, while AC forms crusts and shapeless particles (Figure 1, Figure 2 and Figure 3). For this reason, samples consisting of 1–2 crystalline phases were predominantly studied by optical microscopy and SEM–EDX.

3.1.1. Calcite

Calcite precipitates in Co-, Ni-, Cu- and Fe-rich systems at low (3 °C) and high (23 °C) temperatures (Table 3 and Table 4). Its amount can reach 100 vol.% (Table 3 and Table 4). This is the most common crystalline carbonate in the studied systems. However, no calcite was found in Co-rich solutions (at low temperature; Table 3) at Co2+/Ca ratio 0.10–0.12 and pHinitial 8.50–10.03 (Table 1 and Table 3). No calcite was found in Ni-rich solutions (both at low and ambient temperatures; Table 3 and Table 4) at Ni2+/Ca ratio 1.10–1.60 and pHinitial 7.55–8.18 (Table 1 and Table 3). No calcite was found in Cu-rich solutions (both at low and ambient temperatures; Table 3) at Cu2+/Ca ratio 0.90–1.50 and pHinitial 5.44–6.93 (Table 1 and Table 3).
Calcite forms rhombohedral crystals (Co-, Cu- and Fe-rich systems), spheres (Co-rich systems) and dumbbell-like aggregates (Cu- and Ni-rich systems) up to 50 µm in size (Figure 1a–d, Figure 2a and Figure 3a,b). The shape of calcite crystals clearly depends on temperature, as calcite crystals growing at 3 °C tend to form spheres (CAV 42 and 53; Figure 1a), whereas, at 23 °C, rhombohedral crystals were found (CAV 26; Figure 3a). Calcite is typically covered with AC.

3.1.2. Aragonite

Aragonite precipitates in Co-, Ni- and Cu-bearing systems at low (3 °C) and ambient (23 °C) temperatures (Table 3 and Table 4). Its amount can reach 100 vol.% (Table 3 and Table 4). No aragonite was found in Fe-rich solutions (both at low and ambient temperatures; Table 3 and Table 4). In Co-rich systems (both at low and ambient temperatures), aragonite forms at Co2+/Ca ratio 0.05–1.00 (all ranges) and pHinitial 7.50–10.45 (all ranges) (Table 1 and Table 3). In Ni-rich systems (both at low and ambient temperatures), aragonite forms at Ni2+/Ca ratio 0.60–1.20 and pHinitial 7.94–9.34. In Cu-rich systems (both at low and ambient temperatures), aragonite forms at Cu2+/Ca ratio 0.60–1.20 and pHinitial 6.27–8.72.
Aragonite forms dumbbell-like aggregates (in Co- and Ni-rich systems; Figure 2a,b) and needles (in Cu-rich systems; Figure 2d) up to 100 µm in size (Figure 2a–d and Figure 3c,d). Aragonite aggregates are always transparent and translucent under optical observation (Figure 2a–d). Aragonite is typically covered with AC (Co-/Ni-rich systems)/malachite (Cu-rich systems) and can overgrow calcite (Figure 3c,d).

3.1.3. Monohydrocalcite

Monohydrocalcite precipitates in Co-, Ni- and Cu-rich systems at low temperature (3 °C) only (Table 3). Its amount can reach 100 vol.% (Table 3). No monohydrocalcite was found at ambient temperature (23 °C) synthesis or in Fe-rich solutions (both at low and ambient temperatures; Table 3 and Table 4). In Co-rich systems, monohydrocalcite forms at Co2+/Ca ratio (solution) 0.08–0.12 and pHinitial 8.77–9.73 (Table 1 and Table 3). In Ni-rich systems, monohydrocalcite forms at Ni2+/Ca ratio 0.50–1.00 and pHinitial 8.75–8.80. In Cu-rich systems, monohydrocalcite forms at Cu2+/Ca ratio 0.50–1.00 and pHinitial 6.56–7.50.
Monohydrocalcite forms spheres up to 200 µm in size (Figure 2d–f and Figure 3e,f), which are composed of stepped twisted crystals (Ni-rich systems; Figure 3a) or thin cone-shaped fibers (50–100 nanometers in size) tightly adjacent to each other (Co- and Cu-rich systems; Figure 3c). Monohydrocalcite aggregates are pinkish (Co-rich systems), colorless (Ni-rich systems) or bluish (Cu-rich systems) (Figure 3c,d) and translucent under optical observation with a distinct opal tint. Monohydrocalcite is typically covered with AC (Figure 2d–f) and overgrown by aragonite (Figure 3c).

3.1.4. Amorphous Carbonate (AC)

Even though crystalline carbonates were the main focus of the current work, it was not possible to achieve completely crystalline samples (e.g., CAV 57, 58, Table 3; CME 22, Table 4). AC forms in all systems at low (3 °C) and ambient (23 °C) temperatures at very early stages of synthesis (aging time < 20 days) but can remain “stable” for a long time (Table 3 and Table 4). According to optical observations, the milky solution at the bottom of the beaker was free of crystalline phases for up to ~30–150 days, depending on conditions. Generally, this solution remains “stable” longer at low temperatures. However, in Fe-rich systems, this process lasted ~110 days, even at 23 °C (e.g., No. 16, Table 4). AC is almost constantly present in the resulting precipitates and forms shapeless crusts, which contain CCCs (e.g., Figure 2a and Figure 3a) or small (<<1 µm) particles on the surface of CCCs.

3.1.5. Malachite

Malachite precipitates in Cu-rich systems only (both at low (3 °C) and ambient (23 °C) temperatures; Table 3 and Table 4). Its amount can reach 100 vol.% (Table 3). Malachite appears only in the solution withhigh Cu/Ca ratio. No malachite was found at Cu2+/Ca ratio 0.20–0.85 and pHinitial 6.80–7.55 (Table 1 and Table 3). Malachite forms spheres up to 20 µm in size (Figure 2c and Figure 3d), which are composed of thin fibers (<50 nanometers in size). Malachite aggregates are greenish and translucent under optical observation (Figure 1c). Malachite forms at late stages and typically overgrows aragonite (Figure 1c and Figure 3d).

3.1.6. Iron (Hydro)oxides

Crystalline iron (hydro)oxides (goethite and magnetite) were found in Fe-rich systems, both at low (3 °C) and ambient (23 °C) temperatures (Table 3 and Table 4). Generally, goethite forms at higher Fe content (Fe/Ca 0.4–0.5) compared to magnetite (Fe/Ca 0.1–0.3). Optical observations, along with SEM studies, indicate the presence of amorphous iron (hydro)oxides in syntheses conducted at low Fe content (Fe/Ca < 0.1), as no crystalline phases were detected by PXRD, and the studied calcite crystals were typically covered by a thin orange to yellow crust (Figure 2d).

3.2. Unit Cell Parameters and Chemical Composition of Synthesized Calcium Carbonates

Calcite, aragonite and monohydrocalcite were synthesized as single phases (90–100%; Table 3 and Table 4; Figure 4a) at different Me2+/Ca ratios in solution (Table 3 and Table 4). All heavy metals (Me2+ = Co, Ni, Cu, Fe) have smaller ionic radii than calcium, both in 8- and 6-coordinated polyhedra (VIIICa2+ 1.12 > VIIIFe2+ 0.92 > VIIICo2+ 0.90 Å; VICa2+ 1.00 > VIFe2+ 0.78 > VICo2+ 0.745 > VICu2+ 0.73 > VINi2+ 0.69 Å; [41]). This means that unit cell parameters of CCCs should (1) decrease as their Me2+ content increases and (2) be smaller than that of Me2+-free CCCs.
Unfortunately, it was not possible to directly compare UCP values with the content of Me2+ in respective phases, as SEM–EDX indicated the studied phases, in most cases, form spheres/dumbbell-like aggregates and often contain impurities located near the crystals of the main phase. Because of this, only semi-quantitative data on their chemical composition were obtained, and data on the Me2+/Ca ratio in solution were used as a proxy of maximum Me2+ content in crystalline phases. We also suppose that Me2+ maximum concentrations in the studied CCCs are close to the maximum possible content of heavy metals (under given conditions). It is important to note that the metal may not be incorporated into the lattice but may be sorbed on the surface of CCCs.

3.2.1. Calcite

UCPs of calcite precipitated from Co-(7 syntheses, Co/Ca = 0.050–0.300), Ni- (5 syntheses, Ni/Ca = 0.500–1.050), Cu- (5 syntheses, Cu/Ca = 0.200–0.900) and Fe-rich solutions (13 syntheses, Fe/Ca = 0.100–0.005) were obtained (Table 3, Table 4 and Table 5). UCPs of synthetic Me2+-free calcite are a 4.9900(5), c 17.061(3) [42].
The values of a and c parameters change significantly (over three standard errors) within a Co-rich calcite series. UCPs of calcite precipitated from Co-rich solutions (CAV 46, 49, 52, 53, 67) vary: a 4.960(3)–4.970(3), c 16.876(3)–17.016(3) Å (Table 3). Intense changes in peak position depending on Co2+/Ca ratio (Figure 4b) and small UCPs (compared to synthetic Me2+-free calcite; Figure 5a) clearly indicate Co incorporation into the calcite crystal lattice (Figure 5b). However, it should be noted that temperature significantly affects this process, as measured Co content in calcite does not directly depend on the Co/Ca ratio in solution (see below).
Only the value of c parameter changes significantly (over three standard errors) within the Ni- and Cu-rich calcite series (Figure 5a). The UCPs of calcite precipitated from Ni-rich solutions (CAV 29, 51. 74–76) vary: a 4.985(3)–4.989(3), c 17.008(3)–17.054(4) Å (Table 3). The UCPs of calcite precipitated from Cu-rich solutions (CAV 44, 47, 50, 78–79) vary: a 4.971(3)–4.979(3), c 16.989(3)–17.022(4) Å (Table 3). The position of the calcite peaks weakly depends on the content of Ni/Cu in the solution (Figure 4c,d). However, the UCPs of synthetic calcite precipitated in a Ni-/Cu-rich system are always slightly smaller (especially c parameter) than the UCPs of synthetic Me2+-free calcite (Figure 5a,b), which indicates possible Ni/Cu incorporation into the calcite.
The values of a and c UCPs change significantly (over three standard errors) within the Fe-rich calcite series (Figure 5a,b). The UCPs of calcite precipitated from Fe-rich solutions (CAV 122–124, 137–138, 150; CME 79–81, 93–95) vary: a 4.986(3)–4.996(3), c 17.047(3)–17.060(4) Å (Table 3). Interestingly, only cell parameters of calcites synthesized at different temperatures differ, but no changes in calcite peak position depending on Fe2+/Ca ratio were found on the X-ray diagram (Figure 4e). The UCPs of synthetic calcite precipitated at 3 °C in Fe-rich solutions (Table 3) are the same as for synthetic Me2+-free calcite, while the UCPs of synthetic calcite precipitated at 23 °C in Fe-rich solutions are slightly smaller than those of synthetic Me2+-free calcite (Figure 5a). This indicates possible Fe incorporation (in very small amounts) into the calcite crystal lattice at 23 °C.
The points corresponding to calcites with cobalt and nickel impurities fit into the dependencies for the series CaCO3-CoCO3 and CaCO3-NiCO3, respectively (Figure 5a). The UCPs of calcites synthesized in the presence of copper fit into a different trend, which may indicate the possibility of the existence of a copper-enriched carbonate with a calcite structure.
SEM–EDX showed that calcite precipitated from Co-rich solutions can contain up to ~13.0 wt.% CoO and up to ~1.0 wt.% Na2O (CAV 42; Co/Ca = 0.50). It is worth noting that a cobalt content increase in the solution of 10 times leads to an increase in the content of cobalt in calcite only by a factor of two (~13.0 wt.% and ~7.0 wt.% CoO; Co/Ca = 0.50 and 0.05, CAV 42 and CAV 53, respectively). Temperature significantly affects this process. Well-shaped (rhombohedral) calcite precipitated at higher Co/Ca (CAV 26, Co/Ca 0.10) and 23 °C contains less Co (2–4 wt.% CoO) compared to sphere-like calcite precipitated at lower Co/Ca (CAV 53, Co/Ca 0.05) and 3 °C (6–7 wt.% CoO). These data are in good agreement with UCPs and support Co incorporation into the calcite crystal lattice.
SEM–EDX showed that calcite precipitated from a Ni-rich solution can contain up to ~3.0 wt.% NiO (CAV 74; Ni/Ca = 1.050) and from a Cu-rich solution—up to ~4.0 wt.% CuO (CME 77 Cu/Ca = 0.95). No iron was found in calcite precipitated from Fe-rich solutions. No relation between Ni/Cu content in calcite and Cu/Ni/Ca ratio in solution/temperature/time of exposure was found.

3.2.2. Aragonite

The UCPs of aragonite precipitated from Co-rich solution (Co/Ca = 1.00) after 21 days of exposure (CAV 43) are a 4.959(2), b 7.964(3), c 5.743(2) Å (Table 3; Figure 5c). The UCPs of aragonite precipitated from Ni-rich solutions (two syntheses, CAV 2, 7; Ni/Ca = 1.00–1.10, exposure time 30–44 days) are almost identical (a 4.964–4.966(2), b 7.970(3)–7.975(4), c 5.752(2)–5.756(2) Å (Table 3; Figure 5c)). However, these values (aragonite UCPs in Co- and Ni-rich systems) can be considered very close since the difference between them is within three standard errors. The UCPs of natural Me2+-free aragonite (Ca0.997Sr0.0021Mg0.004Pb0.0005CO3; 4.9614(3), 7.9671(4), 5.7404 (4) Å; [43]) are almost identical to our synthetic aragonite (except c parameter). We can conclude that neither Ni nor Co could be incorporated into aragonite and be reliably detected using the applied approach (Figure 5d).
SEM–EDX showed that aragonite precipitated from Co-rich solutions can contain up to ~3.0 wt.% CoO (CAV 43; Co/Ca = 1.00), from Cu-rich solution—up to ~3.0 wt.% CuO (CME 30; Cu/Ca = 1.20) and from Ni-rich solution—up to ~3.0 wt.% NiO (CME 27; Ni/Ca = 1.10). However, we cannot fully rely on these values, as individual crystals (<1 µm) are smaller than the typical EDX spot.

3.2.3. Monohydrocalcite

The UCPs of monohydrocalcite precipitated from Co-bearing solutions (three syntheses, CAV 70–72, Co/Ca = 0.10–0.11) vary: a 10.536(4)–10.553(3), c 7.534(2)–7.542(2) Å (Table 3; Figure 5e). It is worth noting that these syntheses were stopped at different times (after 37 (CAV 70) and 57 days (CAV 71 and CAV 72; Table 1)). Monohydrocalcites precipitated after 57 days (CAV 71 and CAV 72) have the same UCPs, while the one stopped at earlier stages (CAV 70) has considerably smaller values (Figure 5e). The established difference in the cell parameters within the cobalt series significantly exceeds three standard errors, which indicates significant changes in the lattice metric. Unit cell parameters of monohydrocalcite precipitated from a Cu-rich solution (Ni/Ca = 0.90) after 37 days of exposure (CAV 63) are a 10.547(4), c 7.549(2) Å. The UCPs of natural Me2+-free monohydrocalcite (a 10.5547(4), c 7.56440(29) Å, Ref. [44]) are significantly greater than those of our synthetic monohydrocalcites (Figure 5e). We can conclude that Co and Cu can be incorporated into monohydrocalcite crystal lattices, and exposure time has a great influence on this process.
SEM–EDX showed that monohydrocalcite precipitated from Co-rich solutions can contain up to ~1.0 wt.% CoO (CAV 71; Co/Cab = 0.11), from Cu-rich solutions—up to ~2.0 wt.% CuO (CAV 45; Cu/Ca = 1.00) and from Ni-rich solutions—up to ~3.0 wt.% NiO (CAV 63; Ni/Ca = 0.90). We did not observe chemical zoning or any other variation in chemical composition of monohydrocalcite within a sample.

3.2.4. Amorphous Carbonate

Amorphous carbonate can contain up to ~50 wt.% CoO/~40 wt.% NiO/~60 wt.% CuO. It is also Cl-bearing (0.5–2.0 wt.% Cl) and possibly sufficiently hydrated, which is indicated by low EDX totals (<~65 wt.%) and significant surface cracking during SEM–EDX studies.

4. Discussion

4.1. The Effect of Temperature

Temperature significantly influenced crystallization pathways and the resulting phase composition. No CoCO3 (synthetic analogue of spherocobaltite, calcite structure type), NiCO3 (synthetic analogue of gaspéite, calcite structure type), CuCO3 (no mineral known, unique structure) or FeCO3 (synthetic analogue of siderite, calcite structure type) were synthesized, despite relatively high Me2+ content in the mother solution (Table 1, Table 2 and Table 5).
Spherocobaltite occurs all over the world [46] and is widely synthesized for technical purposes (e.g., [47]). Previously, it was found that the reaction temperature (typically >> 100 °C [47,48]) is a key factor in the formation of the CoCO3 crystals, as the reaction rate decreases with the reaction temperature decreasing, thus decreasing the nucleation rate and growth rate (e.g., [49]). So, probably no spherocobaltite can be obtained at low (<30 °C) temperatures. Gaspéite quite rarely occurs in nature (e.g., [50]), probably due to its instability in air conditions (e.g., [51]). It was successfully synthesized using a hydrothermal technique (e.g., [48]). However, high temperature (>100 °C [52]) is also favored for its formation. CuCO3 was found to be formed at high pressure–temperature (PT) conditions only (20 kbar, 500 °C [53]) and has not been found in nature so far. Siderite is a quite common mineral and can be synthesized at various temperatures (15–360 °C, e.g., [36,54]) at aging time <1 day (e.g., [54]). However, a high Fe/dissolved inorganic carbon ratio is needed (>0.003 [54]), which was not the case in our experiments (Table 3 and Table 4). We can conclude that temperature could be the main factor controlling the presence of CoCO3/NiCO3/CuCO3 in resulting precipitates.
Another striking example of temperature-controlled phase composition is monohydrocalcite, which can precipitate from Co-/Ni-/Cu-rich solutions at 3 °C only (Table 3). Previously, no data on synthetic monohydrocalcite obtained in a Cu/Ni system were available, but monohydrocalcite was synthesized in a Co-rich solution at 25 °C as an intermediate phase (1–24 h) and dissoluted after 1 day of aging [30]. Additionally, natural monohydrocalcite with high Cu content ((Ca0.99Cu0.01)CO3·H2O [45]) was found in the Špania Dolina (Slovakia) deposit. This sample most likely also crystallized at lower temperatures, as it was found in underground tunnels [45]. Previous studies of a Ca–Mg–CO3–H2O system at 3 and 23 °C also showed higher abundancy of monohydrocalcite at low temperatures [22]. We can conclude that monohydrocalcite can be formed as an intermediate metastable phase in Co-/Ni-/Cu-/Mg-rich systems but remains stable (>100 days) if precipitated at low temperatures in Co-/Ni-/Cu-rich systems.
Our data indicate that aragonite content tends to increase as temperature increases in Ni-/Cu-rich systems. Comparison of syntheses provided almost the same Me/Ca ratio (e.g., CAV 63 and CME 31/CAV 66 and CME 27, Ni/Ca 1.10 and Cu/Ca 0.90, respectively; Table 3 and Table 4) and aging time (37 (CAV 63 and 66) and 44 days (CME 31 and 31 and 27); Table 1 and Table 2) and showed the content of aragonite increases by a third (from 0 to 35% and from 66 to 100%, respectively) with increasing temperature from 3 to 23 °C. Data on Co-rich systems do not contradict these findings (Table 3 and Table 4), but wide variations in pH significantly complicate the observed tendencies. This process is the result of (1) monohydrocalcite—aragonite transformation (due to monohydrocalcite instability at higher temperatures) and (2) calcite instability in Me2+-rich systems.

4.2. The Effect of Aging Time

Our data show that aging time (exposure time in solution) significantly influences the resulting phase composition. Syntheses performed at the same conditions (Me2+/Ca ratio, Ca/CO3 ratio, pH) but stopped at different stages show completely different phase compositions (Table 3 and Table 4). It seems that the studied systems do not achieve thermodynamic stability, even after ~100 days of aging, as monohydrocalcite, aragonite and amorphous carbonate were present in the resulting precipitate.
Monohydrocalcite is thought to be a metastable phase, but its content could increase as aging time increases. Monohydrocalcite becomes the main phase when the exposure time is doubled in Co-rich systems (CAV 56: Co/Ca 1.4, pHinitial 9.61, aging time 31 days, calcite 66%, aragonite 34%; CAV 72: Co/Ca 1.4, pHinitial 9.73, aging time 57 days, monohydrocalcite 100%; Table 3). However, an inverse relationship was also found: monohydrocalcite content decreases dramatically when the exposure time is doubled in Ni-rich systems (CAV 45: Ni/Ca 1.0, pHinitial 8.75, aging time 57 days, monohydrocalcite 64%, calcite 36%; CAV 98: Ni/Ca 1.0, pHinitial 8.85, aging time 114 days, calcite 85%, monohydrocalcite 15%; Table 3). Generally, the same tendency was observed in Cu-rich systems (CAV 63: Cu/Ca 0.9, pHinitial 6.93, aging time 37 days, monohydrocalcite 100%; CAV 78: Cu/Ca 0.9, pHinitial 6.75, aging time 57 days, calcite 96%, aragonite 4%; Table 3).
Amorphous carbonate with high Me2+ content (up to 60 wt.% MeO) is always present in the first stage of synthesis and can remain stable for ~100 days (Table 3 and Table 4). We think that long-term studies (aging time > 1 year) and computer simulations are needed to establish whether it is possible to completely transform amorphous carbonate into crystalline phases (under given conditions) or whether it is (thermodynamically?) limited.

4.3. Heavy Metal (Me2+ = Co, Ni, Cu, Fe) Incorporation into CCC Lattices

Some of the heavy metals under consideration may enter the CCC lattices in sufficient amounts (e.g., Co), while the entry of others is limited (e.g., Cu, Ni) or almost impossible (e.g., Fe). The type of cation and the peculiarity of the CCC crystal structure are the main limiting factors.
Calcite can host Co in sufficient amounts. Our data show that calcite unit cell volume decreases from ~366 to ~359 Å3 (our data) while Co/Ca in solution increases (Table 5; Figure 5b). This is in good agreement with Katsikopoulos et al. [29], who showed that calcite unit cell volume decreases from 369 to 362 Å3 while Co/Ca in solution increases. These data also indicate that low temperatures are favorable for Co incorporation into calcite crystal lattices. It is worth noting that Me2+-free calcite has considerably higher unit cell volume (368 Å3 [42]) than that of synthetic Co-bearing calcites, which clearly indicates Co incorporation (Table 5).
Table 5. Comparison of unit cell volumes (V) of CCCs, obtained at different Me/Ca ratios, temperatures and aging times.
Table 5. Comparison of unit cell volumes (V) of CCCs, obtained at different Me/Ca ratios, temperatures and aging times.
No.Sample
Name		T, °CAging
Time, Days		Me2+Me2+/Ca,
Solution		Main PhaseV, Å3Ref.Δ1, Å3Δ2, Å3
1-natural-0.000Calcite
(CaCO3)		367.90(1)[42]--
2CAV-4923158Co0.300362.2(3)This study6.08.3
3CAV-521580.300359.6(3)
4CAV-46570.200364.6(3)
5CAV-67370.050361.2(3)
6CAV-53310.050361.9(3)
7CME-253440.200366.5(3)
8CME-26440.100365.6(3)
9CAV-742357Ni1.050367.6(3)1.31.6
10CAV-76570.950367.2(3)
11CAV-511580.500366.3(3)
12CME-293440.600366.8(3)
13CAV-782357Cu0.900363.6(3)1.84.3
14CAV-79570.850365.4(3)
15CAV-44570.500364.8(3)
16CAV-471580.300364.8(3)
17CAV-501580.200364.3(3)
18CAV-15044Fe0.100368.6(3)1.80.9
19CAV-124660.050368.0(3)
20CAV-123660.030367.8(3)
21CAV-138900.025368.8(3)
22CAV-137900.015368.1(3)
23CAV-122660.010367.8(3)
24CAV-136900.005368.6(3)
25CME-81364Fe0.050367.8(3)
26CME-80640.030368.2(3)
27CME-95940.025367.0(3)
28CME-94940.015368.8(3)
29CME-79640.010368.0(3)
30CME-93940.005368.0(3)
31-~2222Co-Spherocobaltite (CoCO3)281.62(1)[48]-86.3
32-~2222Ni-Gaspéite
(NiCO3)		271.39(1) -96.5
33---Fe-Siderite
(FeCO3)		293.17[8]-74.7
34-natural--Aragonite
(CaCO3)		226.91(1)[43]--
35CAV-432321Co1.000226.8(3)This study1.21.1
36CME-27344Ni1.100228.0(3)
37CME-230Ni1.000227.6(3)
38-natural--Monohydrocalcite
(CaCO3·H2O)		729.79(6)[44]
39-naturalCu-728.68(18)[45] 1.11
40CAV-712357Co0.110727.4(3)This study3.15.49
41CAV-7037Co0.100724.3(3)
42CAV-7257Co0.100726.0(3)
43CAV-6337Cu0.900727.2(3)
Note: ∆1 = Vmax − Vmin, ∆2 = |me-freeV − synthVmin|.
Our data indicate limited solubility of CoCO3 in CaCO3 (both at 3 and 23 °C), as Co content in calcite does not exceed ~0.15–0.25 apfu (<13 wt.% CoO). This is consistent with previous studies by Glynn [55] and Katsikopoulos et al. [29], who showed that Co incorporation in calcite is limited to ~3–16 mol.%, while Ca incorporation in spherocobaltite is limited to ~5–7 mol.%. Katsikopoulos et al. [29] also confirmed that spherocobaltite does not precipitate directly from aqueous solution at 25 °C (and ambient pressure). Moreover, the results of a computer simulation study of the mixing of calcite and spherocobaltite suggested that experimental-based thermodynamic models significantly overestimate the solubility between the two solids and, therefore, underestimate the extension of the miscibility gap under ambient conditions [56]. Authors also conclude that carbonates of Ca1–xCoxCO3 solid solutions are metastable with respect to many compositions observed in nature [56].
Calcite can host very small amounts of Ni and Cu. The unit cell volumes of calcites precipitated from Ni-rich solutions (our data) almost do not differ (366–368 Å3) but are significantly less than that of Me2+-free calcite (368 Å3, Ref. [42]; Table 5; Figure 5b). It is important that neither Ni content in solution, temperature or aging time influence the UCPs. The unit cell volumes of calcites precipitated from Cu-rich solutions (our data) also vary insignificantly (364–365 Å3) but are significantly smaller than that of Me2+-free calcite (368 Å3 [42]). Our observations are consistent with data on natural and synthetic carbonates of (Ca,Ni)CO3/(Ni,Ca)CO3 and (Ca,Cu)CO3/(Cu,Ca)CO3 solid solutions (e.g., [55]). Previously, calcites containing 0.83 and 0.65% Ni were reported by Maksimovic [57] and Maksimovic and Stupar [58], while gaspéite can contain up to 0.45 wt.% CaO [50]. Interestingly, previous studies of NiCO3-MgCO3 systems revealed that solid solutions of intermediate compositions, such as Mg-bearing gaspéite, may form metastably at low temperatures, but they are thermodynamically unstable with respect to unmixing [52]. Copper is also not a common impurity for natural calcites (e.g., [59]), and no mineral with CuCO3 composition has been found yet. Thus, we can conclude that a miscibility gap between calcite and NiCO3/CuCO3 is even bigger than that in the case of spherocobaltite.
Iron incorporation into calcite crystal lattice is doubtful under the studied conditions. The unit cell volumes of calcites precipitated from Fe-rich solutions (our data) vary from ~366 (23 °C) to ~369 (3 °C) and are close to that of Me2+-free calcite (368 Å3 [42]; Table 5; Figure 5b). Previously, it was shown that the solid solution (Ca,Fe)CO3 is not complete (e.g., [28] and a number of works report a wide miscibility gap also at high temperatures (T > 300 °C; e.g., [60])). Glynn [55] showed that Fe incorporation in calcite is limited to ~3–18 mol.%, while Ca incorporation in siderite is limited to ~23 mol.%. So, one can suggest that the Fe content in the studied calcite should vary significantly. However, very limited data are available on (Ca,Fe)CO3/(Fe,Ca)CO3 solid solutions at low (<<25 °C) temperature. Our data show that low Fe content (Fe/Ca 0.005–0.025) at low and ambient temperatures (3–23 °C) resulted in the appearance of AC and required a very long time for CCC crystallization. This inhibitory effect induced by ferrous iron on calcite growth was previously mentioned based on growth experiments [28]. Higher Fe content in solution (Fe/Ca 0.1) resulted in the appearance of iron (hydro)oxides, but no Fe-rich calcite/aragonite crystallized (Table 3 and Table 4). Previously, it was shown that the aqueous solution becomes supersaturated with respect to siderite at lower concentrations than with respect to calcite as a result of its lower solubility product at 25 °C [61], but precipitation is hindered for kinetic reasons [62].
An aragonite crystal lattice cannot host a significant amount of Co/Ni/Cu, which is evident from UCP data on synthetic Co- (our data, Ref. [30]) and Ni-rich samples and optical/EDX data on Cu-rich samples (our data; Table 5; Figure 5d). Aging time and temperature do not affect the dependencies obtained. The unit cell volume of aragonite precipitated in a Co-rich system at 3 °C is ~226 Å3 (our data), in a Co-rich system at 25 °C is ~227 Å3 [30] and in a Ni-rich system at 23 °C is ~227 Å3 (our data; Figure 5d). These values are indistinguishable from natural Me2+-free aragonite (227 Å3 [43]). Previously, aragonite was thought to host some Co [29,30], but this synthetic “Co-aragonite” has higher UCPs than the Co-free one [43]; this fact has not yet been clearly explained. Interestingly, aragonite was successfully synthesized in an Fe-rich solution at 25 °C [28], and this aragonite was thought to be Fe-bearing. However, no UCP data were obtained during this study [28].
A monohydrocalcite crystal lattice can host Co, Ni and Cu in small amounts, which is evident from UCP data on synthetic Co- (our data, Ref. [30]), Ni- (our data) and Cu-rich (our data) and naturally Cu-rich samples [45] (Figure 5c,d). The maximum content of Me2+ is quite small and probably limited to ~0.1 apfu. It is worth noting that this value is close to that of monohydrocalcite synthesized in an Mg-rich system [22,40].
The amount of Me2+ in monohydrocalcite decreases as aging time increases or temperature increases (our data, Ref. [30]). The unit cell volume increases from 724 to 728 Å3 as aging time increases (from 37 to 57 days) in syntheses carried out at 3 °C (our data). Similarly, the unit cell volume increases from ~725 to ~729/~729 to 733 Å3 as aging time increases (from 1 to 24 h), based on data on syntheses carried out at 25 °C [30]. The unit cell volume of Me2+-free natural monohydrocalcite is ~730 Å3, which is significantly higher than the values of most synthetic monohydrocalcites. Probably, low temperature favors Me2+ incorporation into monohydrocalcite and stabilizes it.

5. Conclusions

Crystalline calcium carbonates occur worldwide and play an important role in the carbon cycle and heavy metal sorption. The results of syntheses at different Me2+/Ca ratios, temperatures and aging times showed that the presence of heavy metals in the crystallization medium dramatically changes the crystallization pathways and precipitated phase composition at both low (3 °C) and ambient (23 °C) temperatures. Instead of the normal sequence of calcium carbonate (amorphous calcium carbonate (ACC) → vaterite → calcite), aragonite and monohydrocalcite were found to be the main phases in the resulting precipitate from the Co-/Ni-/Cu-rich systems. In contrast, calcium carbonate crystallization was not significantly affected in the Fe-rich systems. It is likely that the complex behavior of Fe ions in aqueous solution is of paramount importance in explaining this pattern. We can suggest that Fe ions stabilize ACC but do not kinetically/thermodynamically influence the crystallization pathways.
Our data indicate that the crystallization temperature and aging time (exposure time in solution) sufficiently affect the final phase composition of precipitates. Low temperature (3 °C) is favorable for monohydrocalcite and amorphous carbonate, while calcite tends to form more frequently at ambient temperature (23 °C). This could be an explanation for monohydrocalcite/aragonite formation in cold-water marine organisms (e.g., [14,39]) and the specific mineralogy of caves (e.g., [45]). High exposure time in solution leads to a decrease in the content of amorphous carbonate and monohydrocalcite. However, monohydrocalcite aged in solution for a long time is a stable phase and does not undergo further decomposition.
Heavy metals can be incorporated into the calcite/monohydrocalcite crystal lattice in sufficient quantities, while aragonite can host very small amounts of Me2+ (or none). Their content decreases with time if monohydrocalcite is aged in the mother solution. Calcite can concentrate sufficient amounts of Co, small amounts of Ni/Cu and almost no Fe.
It is worth noting that amorphous carbonate is the main carrier of heavy metals in the system in question. Despite the fact that heavy metals precipitate from solution (amorphous carbonate forms), we cannot state that they take a completely insoluble form. Further transformations of amorphous carbonate into crystalline phases will inevitably lead to the formation of calcite/aragonite/monohydrocalcite, the structure of which has a very low ability to absorb the cations under consideration. Simultaneously, the formation of CoCO3/NiCO3/CuCO3/FeCO3 seems to be unfavorable under low and ambient temperature conditions.
The results obtained can be used to solve environmental issues and in mineral resource management. Our studies show that crystalline calcium carbonates, in general, concentrate heavy metals quite poorly at ambient and low temperatures. Moreover, the concentrator phases (e.g., monohydrocalcite) are metastable, which can subsequently lead to the release of heavy metals during their recrystallization. It can be proposed to use the obtained results for the temporary binding into carbonate form of a significant part of transition metal cations in natural or industrial waters, with the possibility of further releasing these cations for subsequent use. One possible future direction of research could be to expand the list of cation additives to study the possibility of their extraction from the aquatic environment for the purpose of both water purification and the use of these elements for technical purposes. Further studies of amorphous calcium carbonates enriched in heavy metals are also needed.

Author Contributions

Conceptualization, O.S.V. and O.V.F.-K.; synthesis, M.A.K.; PXRD studies, I.A.C. and O.S.V.; writing—original draft preparation, O.S.V., M.A.K., I.A.C. and O.V.F.-K.; visualization, O.S.V.; funding acquisition, O.V.F.-K. All authors have read and agreed to the published version of the manuscript.

Funding

This research was funded by the Russian Science Foundation, grant number 19-17-00141.

Data Availability Statement

Data available on request.

Acknowledgments

The laboratory research was carried out in the Research Resource Centers of Saint Petersburg State University “Microscopy and microanalysis”, “Geomodel” and “X-ray Diffraction Studies”.

Conflicts of Interest

The authors declare no conflict of interest.

References

  1. Summerfield, M.A. Global Geomorphology: An Introduction to the Study of Landforms; Longman Sc.: New York, NY, USA, 1991. [Google Scholar]
  2. Sulpis, O.; Jeansson, E.; Dinauer, A.; Lauvset, S.K.; Middelburg, J.J. Calcium Carbonate Dissolution Patterns in the Ocean. Nat. Geosci. 2021, 14, 423–428. [Google Scholar] [CrossRef]
  3. Lowenstam, H.A.; Weiner, S. On Biomineralization; Oxford University Press: New York, NY, USA, 1989. [Google Scholar]
  4. Bindschedler, S.; Cailleau, G.; Verrecchia, E. Role of Fungi in the Biomineralization of Calcite. Minerals 2016, 6, 41. [Google Scholar] [CrossRef]
  5. Millero, F.J.; Woosey, R.; Ditrolio, B.; Water, J. Effect of Ocean Acidification on the Speciation of Metals in Seawater. Oceanography 2009, 22, 72–85. [Google Scholar] [CrossRef]
  6. Stumm, W.; Morgan, J.J. Aquatic Chemistry: Chemical Equilibria and Rates in Natural Waters; John Wiley & Sons: Hoboken, NJ, USA, 2012. [Google Scholar]
  7. Falini, G.; Fermani, S.; Gazzano, M.; Ripamonti, A. Structure and Morphology of Synthetic Magnesium Calcite. J. Mater. Chem. 1998, 8, 1061–1065. [Google Scholar] [CrossRef]
  8. Effenberger, H. Crystal Structure and Infrared Absorption Spectrum of Synthetic Monohydrocalcite, CaCO3 H2O. Monatshefte Für Chem. 1981, 112, 899–909. [Google Scholar] [CrossRef]
  9. Caspi, E.N.; Pokroy, B.; Lee, P.L.; Quintana, J.P.; Zolotoyabko, E. On the Structure of Aragonite. Acta Crystallogr. Sect. B Struct. Sci. 2005, 61, 129–132. [Google Scholar] [CrossRef] [PubMed]
  10. Boettcher, A.L.; Wyllie, P.J. The Calcite-Aragonite Transition Measured in the System CaO-CO2-H2O. J. Geol. 1968, 3, 314–330. [Google Scholar] [CrossRef]
  11. Garvie, L.A.J. Seasonal Formation of Ikaite in Slime Flux Jelly on an Infected Tree (Populus fremontii) Wound from the Sonoran Desert. Sci. Nat. 2022, 109, 48. [Google Scholar] [CrossRef]
  12. Rodriguez-Navarro, C.; Kudłacz, K.; Cizer, Ö.; Ruiz-Agudo, E. Formation of Amorphous Calcium Carbonate and Its Transformation into Mesostructured Calcite. CrystEngComm 2015, 17, 58–72. [Google Scholar] [CrossRef]
  13. Ontl, T.A.; Schulte, L.A. Soil Carbon Storage. Nat. Educ. Knowl. 2012, 3, 35. [Google Scholar]
  14. Little, S.H.; Wilson, D.J.; Rehkämper, M.; Adkins, J.F.; Robinson, L.F.; van de Flierdt, T. Cold-Water Corals as Archives of Seawater Zn and Cu Isotopes. Chem. Geol. 2021, 578, 120304. [Google Scholar] [CrossRef]
  15. Défarge, C. Organomineralization. In Encyclopedia of Geobiology; Reitner, J., Thiel, V., Eds.; Springer: Dodrecht, The Netherlands, 2011; pp. 697–701. [Google Scholar]
  16. Sondi, I.; Mikac, N.; Vdović, N.; Ivanić, M.; Furdek, M.; Škapin, S.D. Geochemistry of Recent Aragonite-Rich Sediments in Mediterranean Karstic Marine Lakes: Trace Elements as Pollution and Palaeoredox Proxies and Indicators of Authigenic Mineral Formation. Chemosphere 2017, 168, 786–797. [Google Scholar] [CrossRef] [PubMed]
  17. de Nooijer, L.J.; Reichart, G.J.; Duenas-Bohorquez, A.; Wolthers, M.; Ernst, S.R.; Mason, P.R.D.; van der Zwaan, G.J. Copper Incorporation in Foraminiferal Calcite: Results from Culturing Experiments. Biogeosciences 2007, 4, 493–504. [Google Scholar] [CrossRef]
  18. Raiswell, R.; Brimblecombe, P. The Partition of Manganese into Aragonite between 30 and 60 °C. Chem. Geol. 1977, 19, 145–151. [Google Scholar] [CrossRef]
  19. Mavromatis, V.; Goetschl, K.E.; Grengg, C.; Konrad, F.; Purgstaller, B.; Dietzel, M. Barium Partitioning in Calcite and Aragonite as a Function of Growth Rate. Geochim. Cosmochim. Acta 2018, 237, 65–78. [Google Scholar] [CrossRef]
  20. Brazier, J.M.; Mavromatis, V. Effect of Growth Rate on Nickel and Cobalt Incorporation in Aragonite. Chem. Geol. 2022, 600, 120863. [Google Scholar] [CrossRef]
  21. Gutjahr, A.; Dabringhaus, H.; Lacmann, R. Studies of the Growth and Dissolution Kinetics of the CaCO3 Polymorphs Calcite and Aragonite II. The Influence of Divalent Cation Additives on the Growth and Dissolution Rates. J. Cryst. Growth 1996, 158, 310–315. [Google Scholar] [CrossRef]
  22. Vereshchagin, O.S.; Frank-Kamenetskaya, O.V.; Kuz’mina, M.A.; Chernyshova, I.A.; Shilovskikh, V.V. Effect of Magnesium on Monohydrocalcite Formation and Unit-Cell Parameters. Am. Mineral. 2021, 106, 1294–1305. [Google Scholar] [CrossRef]
  23. Li, Y.; Zhou, Q.; Ren, B.; Luo, J.; Yuan, J.; Ding, X.; Bian, H.; Yao, X. Trends and Health Risks of Dissolved Heavy Metal Pollution in Global River and Lake Water from 1970 to 2017. In Reviews of Environmental Contamination and Toxicology; Springer: Cham, Switzerland, 2020; Volume 251, pp. 1–24. [Google Scholar]
  24. Liao, Z.; Wu, S.; Zhang, H.; Chen, F. Removal of Aqueous Cu2+ by Amorphous Calcium Carbonate: Efficiency and Mechanism. Minerals 2022, 12, 362. [Google Scholar] [CrossRef]
  25. Macomber, L.; Hausinger, R.P. Mechanisms of Nickel Toxicity in Microorganisms. Metallomics 2011, 3, 1153–1162. [Google Scholar] [CrossRef]
  26. Leyssens, L.; Vinck, B.; Van Der Straeten, C.; Wuyts, F.; Maes, L. Cobalt Toxicity in Humans—A Review of the Potential Sources and Systemic Health Effects. Toxicology 2017, 387, 43–56. [Google Scholar] [CrossRef]
  27. Sun, Z.; Gong, C.; Ren, J.; Zhang, X.; Wang, G.; Liu, Y.; Ren, Y.; Zhao, Y.; Yu, Q.; Wang, Y.; et al. Toxicity of Nickel and Cobalt in Japanese Flounder. Environ. Pollut. 2020, 263, 114516. [Google Scholar] [CrossRef] [PubMed]
  28. Di Lorenzo, F.; Burgos-Cara, A.; Ruiz-Agudo, E.; Putnis, C.V.; Prieto, M. Effect of Ferrous Iron on the Nucleation and Growth of CaCO3 in Slightly Basic Aqueous Solutions. CrystEngComm 2017, 19, 447–460. [Google Scholar] [CrossRef]
  29. Katsikopoulos, D.; Fernández-gonzález, Á.; Carmelo, A.; Prieto, M. Co-Crystallization of Co (II) with Calcite: Implications for the Mobility of Cobalt in Aqueous Environments. Chem. Geol. 2008, 254, 87–100. [Google Scholar] [CrossRef]
  30. González-López, J.; Fernández-González, Á.; Jiménez, A. Precipitation Behaviour in the System Ca2+-Co2+-CO32−-H2O at Ambient Conditions—Amorphous Phases and CaCO3 Polymorphs. Chem. Geol. 2018, 482, 91–100. [Google Scholar] [CrossRef]
  31. Zhang, X.; Guo, J.; Wu, S.; Chen, F.; Yang, Y. Divalent Heavy Metals and Uranyl Cations Incorporated in Calcite Change Its Dissolution Process. Sci. Rep. 2020, 10, 16864. [Google Scholar] [CrossRef]
  32. Liu, Y.; Zou, Y. Experimental Study on the Relationship between Metal Ions and Formation of CaCO3 Crystalline Fouling under Boiling Scaling System. Heat Mass Transf. Stoffuebertragung 2019, 55, 3077–3085. [Google Scholar] [CrossRef]
  33. Lin, P.Y.; Wu, H.M.; Hsieh, S.L.; Li, J.S.; Dong, C.; Chen, C.W.; Hsieh, S. Preparation of Vaterite Calcium Carbonate Granules from Discarded Oyster Shells as an Adsorbent for Heavy Metal Ions Removal. Chemosphere 2020, 254, 126903. [Google Scholar] [CrossRef]
  34. Rosenberg, P.E. Subsolidus Relations in the System CaCO3-FeCO3. J. Sci. 1963, 261, 683–689. [Google Scholar] [CrossRef]
  35. Goldsmith, J.R.; Graf, D.L.; Witters, J.; Northrop, D.A. Studies in the System CaCO3-MgCO3-FeCO3: 1. Phase Relations; 2. A Method for Major-Element Spectrochemical Analysis; 3. Compositions of Some Ferroan Dolomites. J. Geol. 1962, 70, 659–688. [Google Scholar] [CrossRef]
  36. Neerup, R.; Løge, I.A.; Fosbøl, P.L. FeCO3 Synthesis Pathways: The Influence of Temperature, Duration, and Pressure. ACS Omega 2023, 8, 3404–3414. [Google Scholar] [CrossRef] [PubMed]
  37. Mejri, W.; Ben Salah, I.; Tlili, M.M. Speciation of Fe(II) and Fe(III) Effect on CaCO3 Crystallization. Cryst. Res. Technol. 2015, 50, 236–243. [Google Scholar] [CrossRef]
  38. Zeppenfeld, K. Prevention of CaCO3 Scale Formation by Trace Amounts of Copper (II) in Comparison to Zinc (II). Desalination 2010, 252, 60–65. [Google Scholar] [CrossRef]
  39. Iglikowska, A.; Bełdowski, J.; Chełchowski, M.; Chierici, M.; Kędra, M.; Przytarska, J.; Sowa, A.; Kukliński, P. Chemical Composition of Two Mineralogically Contrasting Arctic Bivalves’ Shells and Their Relationships to Environmental Variables. Mar. Pollut. Bull. 2017, 114, 903–916. [Google Scholar] [CrossRef] [PubMed]
  40. Rodriguez-Blanco, J.D.; Shaw, S.; Bots, P.; Roncal-Herrero, T.; Benning, L.G. The Role of Mg in the Crystallization of Monohydrocalcite. Geochim. Cosmochim. Acta 2014, 127, 204–220. [Google Scholar] [CrossRef]
  41. Shannon, R.D. Revised Effective Ionic Radii and Systematic Studies of Interatomic Distances in Halides and Chalcogenides. Acta Crystallogr. Sect. A 1976, 32, 751–767. [Google Scholar] [CrossRef]
  42. Graf, D.L. Crystallographic Tables for the Rhombohedral Carbonates. Am. Mineral. 1961, 46, 1283–1316. [Google Scholar]
  43. De Villiers, J.P.R. Crystal Structures of Aragonite, Strontianite, and Witherite. Am. Mineral. 1971, 56, 758–767. [Google Scholar]
  44. Swainson, I.P. The Structure of Monohydrocalcite and the Phase Composition of the Beachrock Deposits of Lake Butler and Lake Fellmongery, South Australia. Am. Mineral. 2008, 93, 1014–1018. [Google Scholar] [CrossRef]
  45. Mikuš, T.; Patúš, M.; Luptáková, J.; Bancík, T.; Biroň, A. Mineralogical Characteristics of the Secondary Calcium Carbonates Association from the Špania Dolina-The First Occurrence of Monohydrocalcite in Ore Deposits in Slovakia. Bull. Miner. Petrol. 2017, 25, 318–326. [Google Scholar]
  46. Hazen, R.M.; Hystad, G.; Golden, J.J.; Hummer, D.R.; Liu, C.; Downs, R.T.; Morrison, S.M.; Ralph, J.; Grew, E.S. Cobalt Mineral Ecology. Am. Mineral. 2017, 102, 108–116. [Google Scholar] [CrossRef]
  47. Zhang, Q.; Yu, W.; Zhang, D.; Liu, M.; Wang, J.; Meng, K.; Yang, C.; Jin, X.; Zhang, G. Recent Advances on Synthesis of CoCO3 with Controlled Morphologies. Chem. Rec. 2022, 22, e202200021. [Google Scholar] [CrossRef] [PubMed]
  48. Pertlik, F. Structures of Hydrothermally Synthesized Cobalt (II) Carbonate and Nickel (II) Carbonate. Acta Crystallogr. 1986, C42, 4–5. [Google Scholar] [CrossRef]
  49. Li, H.; Duan, X.; Ma, J.; Zheng, W. A Controllable Ionic Liquid-Assisted Hydrothermal Route to Prepare CoCO3 Crystals and Their Conversion to Porous Co3O4. Cryst. Res. Technol. 2012, 47, 25–30. [Google Scholar] [CrossRef]
  50. Bermanec, V.; Sijarić, G.; Kniewald, G.; Mandarino, J.A. Gaspéite and Associated Ni-Rich Minerals from Veins in Altered Ultrabasic Rocks from Duboštica, Bosnia and Herzegovina. Can. Mineral. 2000, 38, 1371–1376. [Google Scholar] [CrossRef]
  51. Seguin, M.K. The Stability of Gaspeite in Inert Atmospheres and in Air. Can. Mineral. 1973, 12, 26–32. [Google Scholar]
  52. Gaines, A.M.; Goldsmith, J.Κ. Crystal Chcmistry and Stability Relations in the System MgCO3—NiCO3. Z. Für Krist.-Cryst. Mater. 1971, 133, 432–444. [Google Scholar] [CrossRef]
  53. Seidel, H.; Ehrhardt, H.; Viswanathan, K.; Johannes, W. Darstellung, Struktur Und Eigenschaften von Kupfer(II)-Carbonat. ZAAC 1974, 410, 138–148. [Google Scholar] [CrossRef]
  54. Holme, E.A.; Henkes, G.A.; Tosca, N.J.; Rasbury, E.T.; Young, J.M.; Schaub, D.R.; Nekvasil, H.; Hurowitz, J.A. Experimental Constraints on Siderite Clumped Isotope Thermometry. Geochim. Cosmochim. Acta 2023, 343, 323–340. [Google Scholar] [CrossRef]
  55. Glynn, P. Solid-Solution Solubilities and Thermodynamics: Sulfates, Carbonates and Halides. Rev. Mineral. Geochem. 2000, 40, 481–511. [Google Scholar] [CrossRef]
  56. Gonzalez-Lopez, J.; Ruiz-Hernandez, S.E.; Fernandez-Gonzalez, A.; Jimenez, A.; Leeuw, N.H.D.; Grau-Crespo, R. Cobalt Incorporation in Calcite: Thermochemistry of (Ca,Co)CO3 Solid Solutions from Density Functional Theory Simulations. Geochim. Cosmochim. Acta 2014, 142, 205–216. [Google Scholar] [CrossRef]
  57. Maksimovic, Z.J. Resultats Preliminaires de l’examen Des Affleurements de Nickel Du Village de Ba, Pres de Ljig, Dans La Serbie Occidentale. Recl. Trav. Acad. Derbe Sci. 1952, 23, 21–48, 49–52. [Google Scholar]
  58. Maksimovic, Z.J.; Stupar, J. La Calcite Nickelomagnesienne et l’aragonite de Rujevac (Village de Ba, Serbie Occidentale). Recl. Trav. Acad. Derbe Sci. 1953, 33, 220–221. [Google Scholar]
  59. Ciurej, A.; Struska, M.; Wolska, A.; Szczerba, M.; Olszak, J. Copper-Bearing Mineralisation in the Upper Devonian Limestones: A Case Study from the Historical Teresa Adit in the Świętokrzyskie Mountains, Poland. Minerals 2023, 13, 54. [Google Scholar] [CrossRef]
  60. Rosenberg, P.E. Subsolidus Relations in the System CaCO3-MgCO3-FeCO3 between 350° and 550 °C. Am. Mineral. 1967, 52, 787–796. [Google Scholar]
  61. Radha, A.V.; Navrotsky, A. Thermodynamics of Carbonates. Rev. Mineral. Geochem. 2013, 77, 73–121. [Google Scholar] [CrossRef]
  62. Fosbøl, P.L.; Thomsen, K.; Stenby, E.H. Review and Recommended Thermodynamic Properties of FeCO3. Corros. Eng. Sci. Technol. 2010, 45, 115–135. [Google Scholar] [CrossRef]
Minerals 13 01497 g001
Figure 1. Calcite crystals formed in (a) Co-rich solution (sample CAV 42), (b) Ni-rich solution (sample CME 29), (c) Cu-rich solution (sample CAV 50) and (d) Fe-rich solution (sample CME 99). Note: CAL—calcite, AC—amorphous carbonate. Sample numbers (in parentheses) correspond to Table 1 and Table 2.
Figure 1. Calcite crystals formed in (a) Co-rich solution (sample CAV 42), (b) Ni-rich solution (sample CME 29), (c) Cu-rich solution (sample CAV 50) and (d) Fe-rich solution (sample CME 99). Note: CAL—calcite, AC—amorphous carbonate. Sample numbers (in parentheses) correspond to Table 1 and Table 2.
Minerals 13 01497 g001
Minerals 13 01497 g002
Figure 2. Morphology of synthetic carbonates: (a) aragonite, Co-rich solution (CAV 43), (b) aragonite, Ni-rich solution (CME 27), (c) aragonite, Cu-rich solution (CME 30), (d) monohydrocalcite, Co-rich solution (CAV 71), (e) monohydrocalcite, Ni-rich solution (CAV 45) and (f) monohydrocalcite, Cu-rich solution (CAV 63). Note: MHC—monohydrocalcite, AR—aragonite, CAL—calcite, MLC—malachite. Sample numbers (in parentheses) correspond to Table 1 and Table 2.
Figure 2. Morphology of synthetic carbonates: (a) aragonite, Co-rich solution (CAV 43), (b) aragonite, Ni-rich solution (CME 27), (c) aragonite, Cu-rich solution (CME 30), (d) monohydrocalcite, Co-rich solution (CAV 71), (e) monohydrocalcite, Ni-rich solution (CAV 45) and (f) monohydrocalcite, Cu-rich solution (CAV 63). Note: MHC—monohydrocalcite, AR—aragonite, CAL—calcite, MLC—malachite. Sample numbers (in parentheses) correspond to Table 1 and Table 2.
Minerals 13 01497 g002
Minerals 13 01497 g003
Figure 3. Morphology of synthetic carbonates: (a) calcite, Fe-rich solution (CME 93), (b) calcite, Co-rich solution (CAV 26), (c) aragonite on monohydrocalcite, Co-rich solution (CAV 71), (d) malachite on aragonite, Cu-rich solution (CME 30), (e) monohydrocalcite and amorphous carbonate, Ni-rich solution (CAV 45), (f) amorphous carbonate and malachite on monohydrocalcite, Cu-rich solution (CAV 63). Note: 1—calcite, 2—aragonite, 3—monohydrocalcite, 4—malachite, 5—amorphous carbonate, Cu-C?—Cu-rich carbonate phase (malachite?). Sample numbers (in parentheses) correspond to Table 1 and Table 2.
Figure 3. Morphology of synthetic carbonates: (a) calcite, Fe-rich solution (CME 93), (b) calcite, Co-rich solution (CAV 26), (c) aragonite on monohydrocalcite, Co-rich solution (CAV 71), (d) malachite on aragonite, Cu-rich solution (CME 30), (e) monohydrocalcite and amorphous carbonate, Ni-rich solution (CAV 45), (f) amorphous carbonate and malachite on monohydrocalcite, Cu-rich solution (CAV 63). Note: 1—calcite, 2—aragonite, 3—monohydrocalcite, 4—malachite, 5—amorphous carbonate, Cu-C?—Cu-rich carbonate phase (malachite?). Sample numbers (in parentheses) correspond to Table 1 and Table 2.
Minerals 13 01497 g003
Minerals 13 01497 g004
Figure 4. Representative room-temperature PXRD patterns: (a) calcite (CME-27), aragonite (CAV 20) and monohydrocalcite (CAV 71), (b) calcite precipitated from Co-bearing solutions (CME-68 and CAV-49, respectively), (c) calcite precipitated from Ni-bearing solutions (CME-29 and CAV 74, respectively), (d) calcite precipitated from Cu-bearing solutions (CME-50 and CAV 78, respectively) and (e) calcite precipitated from Fe-bearing solutions (CME-136 and CAV 124, respectively). Note: Cal—calcite, Ge—germanium.
Figure 4. Representative room-temperature PXRD patterns: (a) calcite (CME-27), aragonite (CAV 20) and monohydrocalcite (CAV 71), (b) calcite precipitated from Co-bearing solutions (CME-68 and CAV-49, respectively), (c) calcite precipitated from Ni-bearing solutions (CME-29 and CAV 74, respectively), (d) calcite precipitated from Cu-bearing solutions (CME-50 and CAV 78, respectively) and (e) calcite precipitated from Fe-bearing solutions (CME-136 and CAV 124, respectively). Note: Cal—calcite, Ge—germanium.
Minerals 13 01497 g004
Minerals 13 01497 g005
Figure 5. Unit cell parameter variation in calcite (a,b), aragonite (c,d) and monohydrocalcite (e,f) precipitated from Me2+-bearing solutions (Me2+= Co, Ni, Cu, Fe) in comparison with published data [29,30,42,43,44,45]: (a) a vs. c in calcite, (b) Me2+/Ca vs. unit cell volume of calcite, (c) a vs. b in aragonite, (d) Me2+/Ca vs. unit cell volume of aragonite, (e) a vs. c in monohydrocalcite and (f) Me2+/Ca vs. unit cell volume of monohydrocalcite. Note: Numbers are aging time (in days).
Figure 5. Unit cell parameter variation in calcite (a,b), aragonite (c,d) and monohydrocalcite (e,f) precipitated from Me2+-bearing solutions (Me2+= Co, Ni, Cu, Fe) in comparison with published data [29,30,42,43,44,45]: (a) a vs. c in calcite, (b) Me2+/Ca vs. unit cell volume of calcite, (c) a vs. b in aragonite, (d) Me2+/Ca vs. unit cell volume of aragonite, (e) a vs. c in monohydrocalcite and (f) Me2+/Ca vs. unit cell volume of monohydrocalcite. Note: Numbers are aging time (in days).
Minerals 13 01497 g005
Table 1. Experimental conditions for the carbonate syntheses carried out at 3 °C.
Table 1. Experimental conditions for the carbonate syntheses carried out at 3 °C.
No.Sample
Name (CAV)		Exposure
Time, Day		Me2+,
mmol/L		Ca,
mmol/L		CO3,
mmol/L		Me2+/Ca,
Solution		pH
Initial		pH
Final
Me2+ = Co
1432110.00010.00020.0001.0007.507.95
242215.00010.00020.0000.5008.558.68
3491583.32010.00020.0000.3008.758.73
4521584.64014.00028.0000.30010.059.90
546572.00010.00020.0000.2008.608.20
654312.80014.00028.0000.2009.449.33
755312.10014.00028.0000.1509.548.80
841A342.80014.00028.0000.1208.508.32
9971031.61214.00028.0000.1159.389.65
1071571.54014.00028.0000.1109.809.73
11961031.47214.00028.0000.10510.039.75
1269371.40014.00028.0000.1009.238.41
1370371.40014.00028.0000.1008.778.54
1472571.40014.00028.0000.1009.739.53
1556311.40014.00028.0000.1009.619.58
16951031.33214.00028.0000.0959.969.84
1773571.26014.00028.0000.0909.509.53
1840A341.20014.00028.0000.0809.158.93
1968370.70014.00028.0000.0508.988.52
2067370.70014.00028.0000.0509.378.62
2153310.70014.00028.0000.0509.639.76
Me2+ = Ni
22573116.00010.00020.0001.6007.557.82
23583114.00014.00028.0001.4007.837.75
24593112.00010.00020.0001.2007.948.06
25663711.00010.00020.0001.1008.088.20
26745710.50010.00020.0001.0508.909.33
27455710.00010.00020.0001.0008.759.10
289811412.00012.00024.0001.0008.859.31
29755710.00010.00020.0001.0008.459.02
3076579.50010.00020.0000.9508.509.09
3165379.00010.00020.0000.9008.417.75
3260318.00010.00020.0000.8008.638.33
33511587.00014.00028.0000.5008.808.46
Me2+ = Cu
34623115.00010.00020.0001.5005.447.49
35643711.00010.00020.0001.1006.357.87
36613110.00010.00020.0001.0006.568.10
3777579.50010.00020.0000.9506.609.06
3863379.00010.00020.0000.9006.937.74
399911410.80012.00024.0000.9007.509.34
4078579.00010.00020.0000.9006.759.25
4179578.50010.00020.0000.8506.809.26
4244575.00010.00020.0000.5006.958.50
43471583.00010.00020.0000.3007.488.87
44501582.80014.00028.0000.2007.559.03
Me2+ = Fe
45154445.00010.00020.0000.5009.3010.49
46153444.00010.00020.0000.4009.5010.40
47152443.00010.00020.0000.30010.159.84
48151442.00010.00020.0000.20010.3510.81
49150441.00010.00020.0000.10010.5010.69
50124660.50010.00020.0000.05011.1510.40
51123660.30010.00020.0000.03011.4011.02
52138900.25210.00020.0000.02511.059.70
53137900.15210.00020.0000.01511.119.99
54122660.10010.00020.0000.01011.2010.09
55136900.05210.00020.0000.00511.059.78
Table 2. Experimental conditions for the carbonate syntheses carried out at 23 °C.
Table 2. Experimental conditions for the carbonate syntheses carried out at 23 °C.
No.Sample
Name (CME)		Exposure
Time, Day		Me2+,
mmol/L		Ca,
mmol/L		CO3,
mmol/L		Me2+/Ca,
Solution		pH
Initial		pH
Final
Me2+ = Co
120425.00010.00020.0000.5009.748.21
221424.00010.00020.0000.4009.929.53
31305.00015.00030.0000.3009.508.40
425442.00010.00020.0000.20010.369.74
526441.00010.00020.0000.10010.458.98
Me2+ = Ni
6224213.00010.00020.0001.3007.937.82
7274411.00010.00020.0001.1008.188.24
823015.00015.00030.0001.0008.208.00
928449.00010.00020.0000.9008.407.95
1023428.00010.00020.0000.8008.317.89
1129446.00010.00020.0000.6009.348.52
Me2+ = Cu
12304412.00010.00020.0001.2006.278.13
1331449.00010.00020.0000.9006.918.06
1432445.40010.00020.0000.7008.807.99
1524426.00010.00020.0000.6008.727.74
Me2+ = Fe
16111415.00010.00020.0000.5009.209.66
17110404.00010.00020.0000.4009.5010.00
18109403.00010.00020.0000.30010.2010.37
19108402.00010.00020.0000.20010.4010.60
20107401.00010.00020.0000.10010.6010.99
2181640.50010.00020.0000.05010.7511.15
2280640.30010.00020.0000.03010.5011.07
2395940.25210.00020.0000.02511.2011.14
2494940.15210.00020.0000.01511.0310.77
2579640.10010.00020.0000.01010.9511.36
2693940.05210.00020.0000.00511.1511.39
Table 3. Phase composition (PXRD data) of precipitates synthesized at 3 °C and UCPs of the main CCCs.
Table 3. Phase composition (PXRD data) of precipitates synthesized at 3 °C and UCPs of the main CCCs.
No. *Sample
Name (CAV)		Me2+/Ca,
Solution		Crystalline
Phase Composition, %		Amorphous
Phase 		Unit Cell Parameters, Å
abc
Me2+ = Co
1431.000Arg 91 Cal 9□4.959(2)7.964(3)5.743(2)
2420.500Cal 81 Arg 19+
3490.300Cal 90 Arg 10-4.970(3)=a16.930(2)
4520.300Cal 97 Arg 3-4.960(3)=a16.876(2)
5460.200Cal 100-4.979(3)=a16.981(2)
6540.200Cal 78 Arg 22+
7550.150Cal 63 Arg 37+
841A0.120Mhcal 86 Arg 14+
9970.115Cal 72 Arg 28±
10710.110Mhcal 100-10.553(3)=a7.542(2)
11960.105Mhcal 70 Cal 27 Arg 3 ±
12690.100Arg 65 Mhcal 23 Cal 12+
13700.100Mhcal 98 Arg 2±10.536(4)=a7.534(2)
14720.100Mhcal 100±10.545(3)=a7.539(2)
15560.100Cal 66 Arg 34+
16950.095Cal 77 Arg 23-
17730.090Mhcal 52 Cal 28 Arg 20□
1840A0.080Cal 60 Arg 35 Mhcal 5+
19680.050Cal 81 Arg 19-
20670.050Cal 95 Arg 5-4.966(2)=a16.911(3)
21530.050Cal 94 Arg 6-4.969(2)=a16.926(3)
Me2+ = Ni
22571.600-■
23581.400-■
24591.200Arg 10 Cal 90■
25661.100Arg 66 Cal 34■
26741.050Cal 90 Arg 10+4.989(3)=a17.054(4)
27451.000Mhcal 64 Cal36■
28981.000Cal 85 Mhcal 15±
29751.000Cal 100■
30760.950Cal 99 Arg 1±4.987(3)=a17.040(4)
31650.900Cal 84 Arg 16+
32600.800Cal 69 Arg 31±
33510.500Cal 96 Mhcal 4±4.985(3)=a17.020(3)
Me2+ = Cu
34621.500Mlc 100±9.536(2)11.870(4)3.270(1)
35641.100Mlc 78 Arg 22+
36611.000Amorphous/Mhcal ■
37770.950Cal 87 Mhcal 7 Arg 6+
38630.900Mhcal 100□10.547(4)=a7.549(2)
39990.900Cal 54 Mhcal 46 □
40780.900Cal 96 Arg 4-4.971(3)=a16.991(3)
41790.850Cal 86 Arg 14-4.979(3)=a17.022(4)
42440.500Cal 100-4.976(2)=a17.012(3)
43470.300Cal 100-4.977(2)=a17.004(3)
44500.200Cal 100-4.976(2)=a16.989(3)
Me2+ = Fe
451540.500Cal 72 Gth 28±
461530.400Cal 78 Gth 22±
471520.300Cal 82 Mgt 18±
481510.200Cal 84 Mgt 16-
491500.100Cal 100-4.993(2)=a17.051(3)
501240.050Cal 100-4.992(2)=a17.054(3)
511230.030Cal 100-4.991(2)=a17.050(3)
521380.025Cal 100-4.996(3)=a17.060(4)
531370.015Cal 100-4.992(2)=a17.051(3)
541220.010Cal 100-4.991(2)=a17.051(3)
551360.005Cal 100-4.995(2)=a17.060(3)
Note: *—numbers are the same as in Table 1; Mhcal—monohydrocalcite, Cal—calcite, Arg—aragonite, Gth—goethite, Mgt—magnetite; -—not detectable (<3 vol.%); ±—traces (3–10 vol.%); +—small amount (10–30 vol.%); □—significant amount (30–50 vol.%); ■—main phase (>50 vol.%).
Table 4. Phase composition (PXRD data) of precipitates synthesized at 23 °C and UCPs of the main CCCs.
Table 4. Phase composition (PXRD data) of precipitates synthesized at 23 °C and UCPs of the main CCCs.
No *Sample
Name (CME)		Me2+/Ca,
Solution		Crystalline
Phase Composition, %		Amorphous
Phase 		Unit Cell Parameters, Å
abc
Me2+ = Co
1200.500Cal 88 Arg 12±
2210.400Cal 89 Arg 11±
310.300Cal 88 Arg 12±
4250.200Cal 97 Arg 3±4.987(3)=a17.016(2)
5260.100Cal 95 Arg 5-4.982(3)=a17.010(3)
Me2+ = Ni
6221.300-■
7271.100Arg 100□4.966(2)7.975(4)5.756(2)
821.000Arg 90 Cal 10□4.964(2)7.970(3)5.752(2)
9280.900Arg 78 Cal 22+
10230.800Cal 61 Arg 39+
11290.600Cal 91 Arg 9-4.986 (3)=a17.039(4)
Me2+ = Cu
12301.200Mlc 88 Arg 12-
13310.900Mlc 65 Arg 35±
14320.700Mlc 47 Cal 38 Arg 15-
15240.600Cal 60 Mlc 34 Arg 6±
Me2+ = Fe
161110.500Cal 63 Gth 37±
171100.400Cal 83 Gth 17-
181090.300Cal 84 Mgt 16-
191080.200Cal 79 Mgt 21-
201070.100Cal 82 Mgt 18-
21810.050Cal 100+4.989(3)=a17.058(4)
22800.030Cal 100+4.991(3)=a17.047(4)
23950.025Cal 100+4.986(4)=a17.048(3)
24940.015Cal 100-4.996(4)=a17.058(4)
25790.010Cal 100-4.991(4)=a17.060(4)
26930.005Cal 100-4.992(3)=a17.052(4)
Note: *—numbers are the same as in Table 2; Mhcal—monohydrocalcite, Cal—calcite, Arg—aragonite, Mlc—malachite, Gth—goethite, Mgt—magnetite. -—not detectable (<3 vol.%); ±—traces (3–10 vol.%); +—small amount (10–30 vol.%); □—significant amount (30–50 vol.%); ■—main phase (>50 vol.%).
Disclaimer/Publisher’s Note: The statements, opinions and data contained in all publications are solely those of the individual author(s) and contributor(s) and not of MDPI and/or the editor(s). MDPI and/or the editor(s) disclaim responsibility for any injury to people or property resulting from any ideas, methods, instructions or products referred to in the content.

Share and Cite

MDPI and ACS Style

Vereshchagin, O.S.; Chernyshova, I.A.; Kuz’mina, M.A.; Frank-Kamenetskaya, O.V. Calcium Carbonate Precipitation Behavior in the System Ca-Me2+-CO3-H2O (Me2+ = Co, Ni, Cu, Fe): Ion Incorporation, Effect of Temperature and Aging. Minerals 2023, 13, 1497. https://doi.org/10.3390/min13121497

AMA Style

Vereshchagin OS, Chernyshova IA, Kuz’mina MA, Frank-Kamenetskaya OV. Calcium Carbonate Precipitation Behavior in the System Ca-Me2+-CO3-H2O (Me2+ = Co, Ni, Cu, Fe): Ion Incorporation, Effect of Temperature and Aging. Minerals. 2023; 13(12):1497. https://doi.org/10.3390/min13121497

Chicago/Turabian Style

Vereshchagin, Oleg S., Irina A. Chernyshova, Maria A. Kuz’mina, and Olga V. Frank-Kamenetskaya. 2023. "Calcium Carbonate Precipitation Behavior in the System Ca-Me2+-CO3-H2O (Me2+ = Co, Ni, Cu, Fe): Ion Incorporation, Effect of Temperature and Aging" Minerals 13, no. 12: 1497. https://doi.org/10.3390/min13121497

Note that from the first issue of 2016, this journal uses article numbers instead of page numbers. See further details here.

Article Metrics

Back to TopTop