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Article

Dissolution and Solubility of the Calcite–Otavite Solid Solutions [(Ca1−xCdx)CO3] at 25 °C

1
College of Earth Sciences, Guilin University of Technology, Guilin 541006, China
2
College of Environmental Science and Engineering, Guilin University of Technology, Guilin 541004, China
3
Guangxi Key Laboratory of Environmental Pollution Control Theory and Technology, Guilin University of Technology, Guilin 541004, China
4
Collaborative Innovation Center for Water Pollution Control and Water Safety in Karst Area, Guilin University of Technology, Guilin 541004, China
*
Authors to whom correspondence should be addressed.
Minerals 2022, 12(6), 756; https://doi.org/10.3390/min12060756
Submission received: 22 April 2022 / Revised: 6 June 2022 / Accepted: 12 June 2022 / Published: 15 June 2022

Abstract

:
A complete series of the calcite–otavite solid solutions [(Ca1−xCdx)CO3] were prepared, and their dissolution processes lasting nine months were experimentally investigated. For the dissolution in the N2-degassed water, the Ca concentrations of the aqueous phases increased up to the steady states after 5040 h of dissolution, and the Cd concentrations of the aqueous phases increased up to the highest values and then decreased gradually to the steady states of 0.017–6.476 μmol/L after 5040 h of dissolution. For the dissolution in the CO2-saturated water, the Ca and Cd concentrations of the aqueous phases increased up to the peak values and then decreased gradually to the steady states of 0.94–0.46 mmol/L and 0.046–9.643 μmol/L after 5040 h of dissolution, respectively. For the dissolution in the N2-degassed water at 25 °C, the mean solubility products (log Ksp) and the Gibbs free energies of formation (ΔGfθ) were estimated to be −8.45–−8.42 and −1129.65–−1129.48 kJ/mol for calcite [CaCO3] and −11.62–−11.79 and −671.81–−672.78 kJ/mol for otavite [CdCO3], respectively. Generally, the log Ksp values decreased non-linearly, and the ΔGfθ values increased linearly with the increasing Cd/(Ca+Cd) mole ratio (XCd) of the (Ca1−xCdx)CO3 solid solutions. In the Lippmann diagrams constructed for the sub-regular (Ca1−xCdx)CO3 solid solutions with the estimated Guggenheim coefficients a0 = −0.84 and a1 = −3.80 for the dissolution in the N2-degassed water or a0 = −1.12 and a1 = −3.83 for the dissolution in the CO2-saturated water, the (Ca1−xCdx)CO3 solid solutions dissolved incongruently, moved progressively up to the quasi-equilibrium curves for otavite and then along the quasi-equilibrium curve from right to left, approached the solutus curve and finally reached the minimum stoichiometric saturation curve for calcite. The considerably Cd-poor aqueous phases were finally in equilibrium with the CdCO3-rich solid phases.

1. Introduction

Among toxic elements, the pollution of agricultural and natural environments by cadmium (Cd) is particularly a challenge, since it easily enters the alimental cycles, causing severe effects on the public health [1]. Out of water and soil, Cd can be absorbed by plants and enriched in animal bodies [2,3], leading to bone diseases and cancers [4,5].
Naturally, Cd presents only in minor amounts at the Earth’s surface [2]. Its enrichment during weathering and soil formation of carbonate rocks, which normally contain a low Cd concentration, has caused large areas with the soil Cd concentration above the national guideline limit in many countries, including China [6,7,8]. Crystalline otavite [CdCO3] has been also found in minor amounts in rocks associated with nonferrous metal mines where natural Cd concentration is highest, and its precipitation may control Cd concentrations in polluted areas [2]. However, various anthropogenic activities, e.g., electrode manufacture, steel plating, metal smelting, printing and painting, have raised environmental Cd concentrations to an unsafe level in certain places [3,8].
Previous studies have demonstrated that the interaction of calcite with the Cd-bearing aqueous solution results in a reduction in the amount of Cd in water, which is principally of importance because calcite is almost ubiquitous in the environments and can have a strong effect on the behavior of many heavy metals, including cadmium, in shallow groundwater aquifers, marine sediments, and soils [2,3,9,10,11,12,13]. The study of the cadmium uptake by calcite suggested that the precipitation and coating of CdCO3 on calcite predominated at higher Cd additions, whereas an ideal surface solid solution was formed between CdCO3 and CaCO3 for low Cd additions (< 1 µmol/g) [9]. This dramatic Cd-removing process has been interpreted as the rapid initial chemisorption of Cd at the calcite surface and the following formation of the calcite-otavite solid solution [5,10,14]. As a result, the aqueous Cd concentration can be decreased significantly to the values restricted by the low solubility of the Cd-rich endmembers [11,12,15].
Therefore, a deep understanding of the Cd2+ distribution between the aqueous and the solid phases is essential in order to reduce the cadmium concentrations in polluted waters from an environmental point of view [3]. Although there are some arguments about the ideal or non-ideal behavior of the calcite [CaCO3] and otavite [CdCO3] solid solution, most data show that Cd2+ (ionic radius 0.97 Å) has a nearly ideal character when substituting for Ca2+ (0.99 Å) in calcite [1,15,16,17,18,19].
Unfortunately, the skill to predict and simulate equilibrium condition in the solid solution–aqueous solution (SSAS) system is constrained, since thermodynamic dealing needs information about the properties of the solid solution and its two endmembers. The published data of the solubility products (Ksp) for otavite vary from 10−13.74 [20] to 10−9.98 [21], which is a difference of more than three orders of magnitude [2,11,22]. Calcite and otavite were suggested to form a nearly ideal solid solution with a Guggenheim parameter of a0 = −0.038 ± 0.078 [16], whereas a Guggenheim parameter of a0 = −0.8 was estimated from the Cd-in-calcite distribution coefficients [23] and assuming a nearly ideal and regular solid solution of ln γCdCO3 = a0[XCd]2 [24].
Briefly, information about the thermodynamic data of the calcite–otavite solid solution [Ca1−xCdx)CO3] is still lacking due to a general lack of thermodynamic equilibrium in experiments lasting many days and the complex dissolution behavior of the solid solution, although its dissolution/precipitation can exert a considerable influence on the cadmium cycling in environments [10,11,25]. Thus, in the present work, with the aim of improving the thermodynamic data and models to describe and predict the SSAS interaction, the calcite–otavite solid solutions [(Ca1−xCdx)CO3] with the Cd/(Ca+Cd) mole ratio (XCd) of 0.00 ≈1.00 were prepared and analyzed by using different characterization techniques and then dissolved in aqueous solutions lasting nine months, in which the solution pH values and the release of different chemical species such as Cd2+, Ca2+ and CO32−/HCO3 were monitored periodically. Lastly, the Lippmann diagrams for the (Ca1−xCdx)CO3 solid solutions were calculated and constructed to analyze the interaction and reaction paths in the SSAS system to assess the solubility of Cd-carbonates and the Cd distribution in aqueous environments.

2. Materials and Methods

2.1. Solid Synthesis

Well-crystallized solid samples (Cal-Ot-00–Cal-Ot-10) of the calcite–otavite solid solutions [(Ca1−xCdx)CO3] were prepared by mixing 2 mol/L Ca+Cd solution with 0.5 mol/L NH3HCO3 solution after the reaction of M2++CO32− = MCO3, where M = (Ca + Cd) (Table 1).
Pure water and chemicals of analytical reagent grade were employed to prepare the precursor solutions. The crystals were precipitated at room temperature (22 ± 1 °C) by dropping 5 mL of 2 mol/L mixture of Ca(NO3)2 and Cd(NO3)2 solutions at varied Cd/(Ca+Cd) mole ratios to 100 mL of the strongly stirred 0.5 mol/L NH3HCO3 solution under N2 (Table 1). The suspended solutions were then stirred further for ten minutes. Lastly, the precipitates were cautiously separated from the suspensions using membrane filters, rinsed with CH3CH2OH and dried at 90 °C for 1 d.

2.2. Characterization

The bulk components of the solid samples obtained were determined by using the wet chemical analytic method. First, 10 mg of each sample was decomposed in 20 mL 1 mol/L HNO3 solution and then diluted to 100 mL using pure water. The Ca and Cd concentrations were analyzed by an inductively coupled plasma-optical emission spectrometer (ICP-OES, Optima 7000DV, Perkin-Elmer Ltd., Waltham, MA, USA) with the detection limits of 0.01 mg/L or an atomic absorption spectrometer (AAS, PinAAcle 900T, Perkin-Elmer Ltd., Waltham, MA, USA) with the detection limits of 0.001 mg/L. All solid products were recognized crystallographically by comparing their X-ray diffraction spectra, which were obtained with an X-ray diffractometer (XRD, X’Pert PRO, PANalytical B.V., Almelo, The Netherlands) at the Cu Kα radiation of 40 kV and 40 mA with a scan rate of 0.09°/min, with the references 01-083-0578 for calcite and 00-042-1342 for otavite of the International Center for Diffraction Data (ICDD). The solids were embedded in resins and coated using carbon to examine their morphological characters by a field-emission scanning electron microscope (FE-SEM, Jeol JEM-7800F, Tokyo, Japan).

2.3. Dissolution Experiment

Dissolution experiments in a closed system were performed. Instrumental-grade N2 and CO2 were used to saturate the initial solutions. The solution pH values were left to change freely without adjustments and recorded regularly. In the dissolution experiment with Cal-Ot-00–Cal-Ot-10, 10 g of the solid was weighed in a 5 L polyethylene terephthalate (PET) bottle, 5 L of the N2-degassed water (initial pH 6.55) or the CO2-saturated water (initial pH 3.89) was then filled into the bottle, which was capped and placed in an air-conditioned room at 25 °C. Then, 100 mL of the aqueous solution was regularly sampled from each bottle (1 h–270 d), and the temperature and pH value were recorded at the same time of sampling. Afterwards, 20 mL of the aqueous sample was filtered using 0.22 µm membranes and instantly stabilized with 0.2% HNO3 for the determination of Ca and Cd concentrations by ICP-OES. The total dissolved H2CO30 concentrations (HCO3+CO3) were measured by an automatic potentiometric titrator (Metrohm 888 Titrando). After dissolution, the solid residues were sampled and examined using XRD and FE-SEM as well.

2.4. Thermodynamic Calculations

Firstly, the geochemical program PHREEQC (Version 3.7.3) [26] was applied to compute the free ion activities of Ca2+, Cd2+ and CO32−, and then, the ion activity product (IAP) was calculated after its definition, which equals the solubility product (Ksp) of the (Ca1−xCdx)CO3 solid solution at dissolution equilibrium. The species in the PHREEQC simulation included: Cd2+, CdOH+, CdCO30, Cd(CO3)22− and CdHCO3+ for cadmium; Ca2+, CaOH+, CaHCO3+ and CaCO30 for calcium; and CO32−, HCO3, H2CO30, CaCO30, CaHCO3+, CdCO30, Cd(CO3)22− and CdHCO3+ for carbonate. The minteq.v4.dat database covered the thermodynamic properties for all aqueous species and solid phases involved in the simulation (Supplementary Material Table S1). The ionic strengths (<0.01433 mol/L) were within the valid limit for the extended Debye–Hückel equations.

3. Results and Discussion

3.1. Solid Characterization

3.1.1. Composition

The Cd/(Ca+Cd) mole ratio (XCd) of the synthetical calcite–otavite solid solution, which was calculated for the chemical compound with the formula of (Ca1−xCdx)CO3 when it was normalized to Ca+Cd = 1.00, was close to that of the precursor solution (Table 1). After their dissolution in the N2-degassed water and the CO2-saturated water for 270 d, the Cd/(Ca+Cd) mole ratios (XCd) of the solid residues increased slightly, particularly in dissolution in the CO2-saturated water.

3.1.2. XRD

All XRD patterns of the synthetical crystals were confirmed to be the R 3 ¯ c space group and belonged to the calcite–otavite solid solutions [(Ca1−xCdx)CO3] (Figure 1). The most intense peaks (Position (°2θ)/Intensity (%)/FWHM (°2θ)/hkl) for the synthetic solid of XCd = 0.00 (Cal-Ot-00) were 23.06/9.18/0.10/012, 29.41/100.00/0.13/104, 35.96/10.92/0.13/110, 39.41/18.35/0.10/113, 43.16/13.54/0.10/202, 47.51/14.56/0.10/018, 48.51/15.62/0.08/116 and 57.40/5.91/0.12/122, agreeing well with calcite (ICSD Reference code 01-083-0578). The most intense peaks (Position (°2θ)/Intensity (%)/FWHM (°2θ)/hkl) for the synthetic solid of XCd = 1.00 (Cal-Ot-10) were 23.53/72.72/0.26/012, 30.34/100.00/0.20/104, 36.38/27.03/0.18/110, 40.09/3.56/0.51/113, 43.84/19.12/0.36/202, 48.15/11.11/0.41/018, 49.81/32.79/0.31/116 and 58.19/11.47/0.51/122, agreeing well with otavite (00-042-1342).
The reflection peaks, especially (104), moved a little to the higher angles with the increasing XCd of the solid solutions (Figure 1; Supplementary Material Figure S1), which was due to the smaller interplanar distances of Cd-calcite than pure calcite. All solid phases varied only in their peak location, intensity, and width in the XRD patterns, suggesting that they were not a simple crystal mixture of calcite [CaCO3] and otavite [CdCO3], but rather the calcite–otavite solid solutions [(Ca1−xCdx)CO3]. After dissolution in the N2-degassed water and the CO2-saturated water for 270 d, no detectable changes were observed for all solid samples (Figure 1; Supplementary Material Figure S1).

3.1.3. SEM

The crystal morphologies of the synthetical calcite–otavite solid solutions [(Ca1−xCdx)CO3] showed a strong reliance on the Cd/(Ca+Cd) mole ratios of the precursor solutions, showing the influence of Cd2+ (even at low concentrations) on the {104} face growth of the calcite-type crystals [27,28,29,30]. The synthetical pure calcite (Cal-Ot-00) showed a rhombohedral morphology defined by the {104} faces with the particle sizes of ≈10 µm (Figure 2). The sizes of the (Ca1−xCdx)CO3 solid solutions decreased with the increasing XCd, although their morphologies remained rhombohedral for the (Ca1−xCdx)CO3 aggregates (Figure 2; Supplementary Material Figure S2), which indicated that cadmium can affect the growth and morphology of the calcite-type crystals. As the XCd increased, the size of the individual crystal declined, and the (Ca1−xCdx)CO3 solid solutions and the endmember CdCO3 (XCd=0.10–1.00; Cal-Ot-01–Cal-Ot-10) changed from coarse crystals to spherical aggregates of very fine small crystals (Figure 2). The spherical surfaces were characterized by the aggregation of plentiful minuscule blocks of {104} faces that seemed somewhat disoriented, since they traced an external spherical shape. In addition, the decrease in the ‘physical’ spherical diameter with the increasing XCd was also observed.
Generally, all solids of the (Ca1−xCdx)CO3 solid solutions after dissolution in the N2-degassed water (Figure 3) and the CO2-saturated water (Figure 4) presented the same morphologic character as before; i.e., the solids changed from coarse crystals to spherical aggregates of very fine small crystals with the XCd increase. After dissolution, the edges of the calcite crystals degenerated, and etch pits on the surface could be found occasionally (Supplementary Material Figure S3).

3.2. Change of the Aqueous Phase

For the dissolution in the N2-degassed pure water, the pH values of the aqueous phases increased from 6.55 up to the highest values of 9.81–8.45 in <1–1200 h and then decreased gradually to the steady states of 8.54–7.83 (Figure 5). The Ca and HCO3+CO3 concentrations of the aqueous phases increased up to the steady states after 5040 h of dissolution for all solids of the (Ca,Cd)CO3 solid solutions. The Cd concentrations of the aqueous phases increased up to the highest values of 0.049–7.588 μmol/L in <1–72 h and then decreased gradually to the steady states of 0.017–6.476 μmol/L after 5040 h of dissolution (Figure 5). The Cd/(Ca+Cd) mole ratios were significantly smaller than the Cd/(Ca+Cd) mole ratios of the corresponding solids and decreased with time to a steady state (Supplementary Material Figure S4). Generally, the Cd concentrations of the aqueous phases at the last steady state increased, while the Ca and HCO3+CO3 concentrations of the aqueous phases decreased with the increasing Cd/(Ca+Cd) mole ratios of the solids (Supplementary Material Figure S5).
For the dissolution in the CO2-saturated water, the pHs of the aqueous phases increased rapidly from 3.89 up to 5.67–4.92 in <1 h and then gradually to the steady states of 8.05–7.65 after 5040 h (Figure 6). The Ca and HCO3+CO3 concentrations of the aqueous phases increased up to the highest value of 6.14–0.58 mmol/L and 11.223–0.662 mmol/L after 48–1440 h and then decreased gradually to the last steady states of 0.94–0.46 mmol/L and 2.202~0.196 mmol/L, respectively (Figure 6). The Cd concentrations of the aqueous phases increased quickly up to the highest values of 68.943–470.239 μmol/L in <1 h and then decreased gradually to the steady states of 0.046–9.643 μmol/L after 5040 h of dissolution (Figure 6). The Cd/(Ca+Cd) mole ratios of the aqueous phases were obviously smaller than the Cd/(Ca+Cd) mole ratios of the corresponding solids and decreased with time to the steady states, indicating a non-stoichiometric release of Ca and Cd (Supplementary Material Figure S4).
Generally, the Cd concentrations of the aqueous phases at the steady state of dissolution increased with the increasing Cd/(Ca+Cd) mole ratios of the solids (Supplementary Material Figure S5). The Ca and HCO3+CO3 concentrations of the aqueous phases showed the highest values at XCd=0.31–0.40 (Cal-Ot-03–Cal-Ot-04) and then decreased as the XCd values decreased or increased (Supplementary Material Figure S5).
The dissolution of the (Ca1−xCdx)CO3 solid solutions is given as:
(Ca1−xCdx)CO3 = (1 − x)Ca2+(aq) + xCd2+(aq) + CO32−(aq)
Ca2+ + OH = Ca(OH)+
Cd2+ + OH = Cd(OH)+
Ca2+ + H+ + CO32− = CaHCO3+
Cd2+ + H+ + CO32− = CdHCO3+
nH+ + CO32−= HnCO3(2−n)− (n = 1,2)
Cd2+ + nCO32− = Cd(CO3)n2−2n (n = 1,2)
At the beginning of dissolution, all components were dissolved stoichiometrically (Equation (1)). For the dissolution in the N2-degassed water and the CO2-saturated water, the pH of the aqueous phase increases were ascribed to the formation of CaHCO3+, CdHCO3+ and HnCO3(2−n)− (Equations (4)–(6)), which resulted in the H+ depletion. The notable pH variation showed that the initial dissolution was pH controlled [31,32]. More calcium was dissolved than cadmium, indicating that the (Ca1−xCdx)CO3 solid solutions dissolved non-stoichiometrically, which suggested the existence of an interface dissolution–precipitation process, and cadmium was more favorable for nucleation on the crystal surfaces to form new precipitates with a higher XCd.

3.3. Determination of the Stoichiometric Solubility

The dissolution experiments were performed until the difference in the IAP values with respect to each solid computed from the last three solution samples (i.e., 5040 h, 5760 h and 6480 h, supposing the steady states were attained) were <0.25 log units. The results of the PHREEQC simulations revealed that all aqueous phases in this work were unsaturated with respect to any likely secondary minerals, such as Ca(OH)2 [portlandite] and Cd(OH)2.
The thermodynamic solubility product (Ksp) can be estimated from the longstanding steady state or extrapolated from the ion activity product (IAP) that corresponds to the equilibrium constant of the mineral dissolution [33]. The Ksp values for the (Ca1−xCdx)CO3 solid solutions after the dissolution reaction Equation (1) at equilibrium can be calculated according to Equation (8):
Ksp = IAP = {Ca2+}1−x{Cd2+}x{CO32−}
where {} designates the free ion activity.
The standard free energy of reaction (ΔGrθ) can be obtained from Ksp at 298.15 K and 0.101 MPa according to Equation (9):
ΔGrθ = −5.708 log Ksp
For Equation (1),
ΔGrθ = (1−x)ΔGfθ[Ca2+] + xΔGfθ[Cd2+] +ΔGfθ[CO32−] − ΔGfθ[(Ca1−xCdx)CO3]
Rearranging,
ΔGfθ[(Ca1−xCdx)CO3] = (1−x)ΔGfθ[Ca2+] + xΔGfθ[Cd2+] + ΔGfθ[CO32−] − ΔGrθ
Table 2 and Table 3 list the analytical results of the solution pH, Ca, Cd and HCO3/CO32− concentrations as well as the calculated log IAP values at the last steady state (≈log Ksp) for the dissolution of the (Ca1−xCdx)CO3 solid solutions. By using the thermodynamic data, ΔGfθ[Ca2+] = −553.54 kJ/mol, ΔGfθ[Cd2+] = −77.58 kJ/mol and ΔGfθ[CO32−] = −527.9 kJ/mol [34], the Gibbs free energies of formation for the (Ca1−xCdx)CO3 solid solutions, ΔGfθ[(Ca1−xCdx)CO3], were also estimated.
For the dissolution in the N2-degassed water at 25 °C, the mean solubility products (log Ksp) and the Gibbs free energies of formation (ΔGfθ) were estimated to be −8.45 ± 0.04–−8.42 ± 0.06 and −1129.65 ± 0.22–−1129.48 ± 0.30 kJ/mol for calcite [CaCO3] and −11.62 ± 0.05–−11.79 ± 0.10 and −671.81 ± 0.29–−672.78 ± 0.55 kJ/mol for otavite [CdCO3], respectively. For the dissolution in the CO2-saturated water at 25 °C, the mean log Ksp and ΔGfθ values were estimated to be −8.39 ± 0.08–−8.37 ± 0.09 and −1129.32 ± 0.44–−1129.22 ± 0.52 kJ/mol for calcite [CaCO3] and −11.49 ± 0.07–−11.60 ± 0.07 and −671.09 ± 0.43–−671.71 ± 0.38 kJ/mol kJ/mol for otavite [CdCO3], respectively.
The results are consistent with the various Ksp and ΔGfθ values for CaCO3 found in the literature. For instance, either the minteq.v4.dat database [35] or the phreeqc.dat database gives a log Ksp of −8.48 [26,36]. CaCO3 has a very well-defined solubility product (log Ksp) value between −8.30 [37] and −8.58 [38]. Many inconsistent values of the solubility product and ΔGfθ for otavite [CdCO3] have been reported in the literature (Supplementary Material Tables S2 and S3). Some of the log Ksp values for otavite [CdCO3] reported are: −13.74 [20]; −13.6 [39]; −12.1 ± 0.1 [2]; −12.06 [40]; −12.03 ± 0.13 [41]; −11.31 ± 0.03 [23]; −11.292 [42]; −11.284 [43] and −11.209 [44]. The various ΔGfθ values for otavite [CdCO3] found in the literature ranged from −595.07 ± 4.18 kJ/mol [45] to −779.43 kJ/mol [46], including −662.7 kJ/mol [21]; −669.44 ± 2.64 kJ/mol [47]; −670.3 ± 2.1 kJ/mol [48]; −671.1 ± 1.1 [49]; −672.79 kJ/mol [43]; −674.2 ± 0.6 kJ/mol [50]; −674.7 ± 0.6 kJ/mol [2] and −683.46 kJ/mol [39]. As a result, a great difference among the log Ksp values for otavite [CdCO3] could be also found in various databases. For instance, whereas the minteq.v4.dat database [35] compiles a log Ksp of −12.0, the wateq4f.dat database gives a log Ksp of −12.1 and the llnl.dat database gives a log Ksp of −12.2288 [26,36].
For the dissolution in the N2-degassed water, with the increasing XCd of the (Ca1−xCdx)CO3 solid solution, log Ksp values decreased non-linearly from −8.45 ± 0.04–−8.42 ± 0.06 for calcite [CaCO3] to −11.62 ± 0.05–−11.79 ± 0.10 for otavite [CdCO3] (Supplementary Material Figure S6), while the estimated ΔGfθ values increased linearly from −1129.65 ± 0.22–−1129.48 ± 0.30 kJ/mol for calcite to −671.81 ± 0.29–−672.78 ± 0.55 kJ/mol for otavite (Supplementary Material Figure S7). For the dissolution in the CO2-saturated water, with the increasing XCd, the log Ksp values decreased non-linearly from −8.39 ± 0.08–−8.37 ± 0.09 for calcite [CaCO3] to −11.49 ± 0.07–−11.60 ± 0.07 for otavite [CdCO3] (Supplementary Material Figure S6), while the estimated ΔGfθ values increased linearly from −1129.32 ± 0.44–−1129.22 ± 0.52 kJ/mol for calcite to −671.09 ± 0.43–−671.71 ± 0.38 kJ/mol for otavite (Supplementary Material Figure S7). These results were in close agreement with the experimental points and the straight line which connects the log Ksp values for CaCO3 and CdCO3, the endmembers of the solid solution series [9]. The ΔGfθ[(Ca1−xCdx)CO3] values were estimated to vary nearly linearly from −1128.79 kJ/mol for calcite to −671.06 kJ/mol for otavite with mineral composition for miscible solids [49]. Complete miscibility is consistent with the similarity in ionic radii of the two metals; the ionic radii of Ca2+ and Cd2+ in octahedral coordination to oxygens are 0.99 Å and 0.97 Å, respectively, and the endmember phases otavite and calcite are isomorphous.

3.4. Lippmann Diagram

3.4.1. Construction of the Lippmann Diagram

The solid solution–aqueous solution (SSAS) system is fundamentally important in understanding the geochemical interaction. The construction method for the Lippmann diagrams involving several SSAS systems have been depicted in detail by many previous works [11,22,24,28,51,52,53,54].
The Lippmann diagram is a chart that displays the relationship between the “solidus” and “solutus” phases in the SSAS system. The “total activity product” (ΣΠSS) is expressed by the sum of the partial activity products of the two endmembers at equilibrium. The “solidus” curve and the “solutus” curve are the plotting of ΣΠSS against the solid component and the solution component, respectively.
For the calcite–otavite solid solutions [(Ca1−xCdx)CO3], the “solidus” curve is defined as:
Σ Π SS = ( { Ca 2 + } + { Cd 2 + } ) { C O 3 2 } = K Ca X Ca γ Ca + K Cd X Cd γ Cd
where {} represents aqueous activity. K Ca and K Cd , X Ca and X Cd , γ Ca and γ Cd are the solubility products, the mole ratios (1−x, x) and the activity coefficients of CaCO3 and CdCO3 in the (Ca1−xCdx)CO3 solid solutions, respectively.
The “solutus” curve is defined as:
Σ Π SS = 1 X Ca 2 + , aq K Cd γ Ca + X Cd 2 + , aq K Cd γ Cd
where X Ca 2 + , aq and X Cd 2 + , aq are the activity ratios of the aqueous free ions of Ca2+ and Cd2+.
For the members of fixed XCd, a series of the minimum stoichiometric saturation curves are defined as:
Σ Π SS = IAP ( X Ca 2 + , aq ) X Ca ( X Cd 2 + , aq ) X Cd
The saturation curves for the two endmembers, pure CaCO3 and CdCO3, are defined as:
Σ Π CaCO 3 = Ca 2 + CO 3 2 ( X Ca 2 + , aq ) X Ca = K Ca ( X Ca 2 + , aq ) X Ca
Σ Π CdCO 3 = Cd 2 + CO 3 2 ( X Cd 2 + , aq ) X Cd = K Cd ( X Cd 2 + , aq ) X Cd

3.4.2. Lippmann Diagram for the Non-Ideal (Ca1−xCdx)CO3 Solid Solutions

The Lippmann diagram for the (Ca1−xCdx)CO3 solid solutions has been calculated by using the endmembers log Ksp values of −8.48 for calcite [36] and −12.1 for otavite [2] and assuming an ideal solid solution [11]. The extremely low solubility of otavite compared to calcite indicates a great preferential partitioning of CdCO3 in the solids [55] and explains the effectiveness of carbonates in eliminating cadmium from the environments [10,12,15,17,22,23]. A small variation of XCd,aq results in a big change of the XCd in the solid. Only a very limited range of the aqueous composition can exist in equilibrium with the intermediate solid solution [11]. The solid tends to be either Ca-rich (XCd < 0.1) or Cd-rich (XCd > 0.9) in a very narrow range of the fluid composition (0.000027 < XCd,aq < 0.002154) [11].
Previous experiments have revealed that a regular model was perhaps inadequate to illustrate closely thermodynamic equilibria involving carbonates [56]. The sub-regular model (two-parameter Guggenheim equation) could be more successfully employed to fit experimental solubility data for the (Ca,Mg)CO3-H2O, (Sr,Ba)CO3-H2O, (Sr,Ca)CO3-H2O systems [52]. In defect of more exact data, thus, the two Guggenheim coefficients a0 and a1 were estimated in the present work, assuming the (Ca1−xCdx)CO3 solid solutions to be sub-regular.
When the SSAS interaction achieved the stoichiometric saturation state, the two Guggenheim parameters a0 and a1 can be estimated by using:
Ln Ksp = x(1 − x)a0 + x(1 − x)(x − (1 − x))a1 + (1 − x) ln[(1 − x) KCa] + x ln[x KCd)]
The log IAPs (≈log Ksp) for the last three solution samples after 5040–6480 h dissolution in the N2-degassed water and the CO2-saturated pure water are plotted against XCd of the (Ca1−xCdx)CO3 solid solutions (Supplementary Material Figure S8), which exhibited that the experimental log Ksp values were close to and somewhat lower than the data for ideal (Ca1−xCdx)CO3 solid solutions. The Ksp values as a function of XCd could be very well fitted to Equation (17) with the two-parameter Guggenheim equations of a0 = −0.84 and a1 = −3.80, a0 = −1.12 and a1 = −3.83, respectively.

3.4.3. Solid Solution–Aqueous Solution Interaction

Figure 7 and Figure 8 illustrate the Lippmann diagrams for the (Ca1−xCdx)CO3 solid solutions at 25 °C, which were computed with the sub-regular solid solution (a0 and a1) and the endmember log Ksp values for calcite [CaCO3] and otavite [CdCO3] from the results of dissolution in the N2-degassed water and the CO2-saturated water, respectively. Along with the solutus curve and the solidus curve, the stoichiometric saturation curves for pure calcite (XCd = 0.00), the (Ca1−xCdx)CO3 solid solutions (XCd = 0.21, 0.41, 0.61 and 0.81) and pure otavite (XCd = 1.00) were also calculated and drawn in the Lippmann diagrams.
The curve for the saturation with respect to otavite and the solutus curve closely follow each other except in the part where XCd,aq is approximate to 0.00, in which the pure otavite saturation curve approaches an infinitely great ΣΠ value, while the solutus curve intercepts the ΣΠ axis when the value is equal to the pure calcite solubility product. The pure calcite saturation curve is always above the solutus curve (except at XCaCO3 = 1.00); hence, all solutions at thermodynamic equilibrium with respect to any (Ca1−xCdx)CO3 solids are always unsaturated with respect to pure calcite. These solutions are also always unsaturated with respect to pure otavite by an extremely small degree. Since the Ksp value for calcite is greatly larger than that of otavite, the solutus curve will fundamentally be identical to the pure otavite saturation curve when the solid-solution activity coefficient is unity except at the XCd,aq value very close to zero.
The quasi-equilibrium is the state when the thermodynamic equilibrium is not attained [57], but the driving force for the ion flux of one or both components between aqueous solution and solid solution can be negligible, and one of the two chemical potential differences is zero, but the other is not. At the fixed XCd of the (Ca1−xCdx)CO3 solid solutions, ΣΠCd with respect to otavite is expressed as a function of XCd,aq:
Σ Π Cd ( X Cd 2 + , aq )   =   [ K Cd γ Cd X Cd ] / [ X Cd 2 + , aq ]
The solid rose lines in the Lippmann diagrams showed the quasi-equilibrium of the solid solution of XCd = 0.10 with respect to the whole series of aqueous solutions (Figure 7 and Figure 8). The region under the solid rose lines denoted unsaturation with respect to CdCO3 and CaCO3; the region above the solid rose lines denoted oversaturation with respect to CdCO3 and unsaturation with respect to CaCO3.
In addition, the data of this work were also plotted as ({Ca2+}+{Cd2+}){CO32−} versus X Cd 2 + , aq for dissolution in the N2-degassed water (Figure 7) and the CO2-saturated pure waters (Figure 8). An increase in the released Ca2+ concentration and a highest Cd2+ concentration for the solids were generally viewed in the dissolution progress. The experimental data points indicated that the (Ca1−xCdx)CO3 solid solutions dissolved incongruently and moved progressively up to the quasi-equilibrium curve for otavite [CdCO3] and then along the quasi-equilibrium curve for otavite from right to left, approached the solutus curve and finally reached the minimum stoichiometric saturation for calcite [CaCO3]. The aqueous Cd-poor solutions were finally in equilibrium with the CdCO3-rich solids. The saturation indexes (SI) with respect to otavite [CdCO3] increased as XCd increased, and the aqueous solutions were unsaturated with respect to otavite throughout the dissolution of the solid solution in the CO2-saturated water, indicating that the dissolution of the calcite–otavite solid solutions [(Ca1−xCdx)CO3] might be controlled by a quasi-equilibrium with respect to otavite (Figure 7 and Figure 8; Supplementary Material Figure S9). The dissolution of a solid solution to stoichiometric saturation would cause either oversaturation or unsaturation with respect to the pure endmembers [25]. In the SSAS system for (Ca1−xCdx)CO3, oversaturation is expected with respect to the CdCO3 endmember if the mole fraction of CdCO3 in the solid solution is greater than ≈10−2.75 [25]. The extremely great difference between the calcite and otavite solubility products decided that the solidus curve and the solutus curve plot very far apart. Thus, the compositions of the aqueous and solid phases coexisting at equilibrium differed extremely, i.e., a considerably Cd-poor aqueous solution might be in equilibrium with a considerably Cd-rich solid phase. These results offer a deeper understanding of the geochemical cycle of cadmium in the environment.

4. Conclusions

For the dissolution of the calcite–otavite solid solutions [(Ca1−xCdx)CO3] in the N2-degassed water, the Ca concentrations of the aqueous phases increased up to the steady states after 5040 h of dissolution and the Cd concentrations of the aqueous phases increased up to the highest values and then decreased gradually to the steady states of 0.017–6.476 μmol/L after 5040 h of dissolution. For the dissolution in the CO2-saturated water, the Ca and Cd concentrations of the aqueous phases increased up to the highest value and then decreased gradually to the steady states of 0.94–0.46 mmol/L and 0.046–9.643 μmol/L after 5040 h of dissolution, respectively. The Cd/(Ca+Cd) mole ratios of the aqueous phases were obviously smaller than the Cd/(Ca+Cd) mole ratios of the corresponding solid phases and decreased with time to the steady states, indicating a strong non-stoichiometric release of Ca and Cd.
For the dissolution in the N2-degassed and CO2-saturated pure waters at 25 °C, the mean log IAP values at the longstanding steady state (≈log Ksp) were estimated for calcite [CaCO3] to be −8.45 ± 0.04–−8.42 ± 0.06 and −8.39 ± 0.08–−8.37 ± 0.09 with the ΔGfθ values of −1129.65 ± 0.22–−1129.48 ± 0.30 kJ/mol and −1129.32 ± 0.44–−1129.22 ± 0.52 kJ/mol, respectively. For the dissolution in the N2-degassed and CO2-saturated pure waters at 25 °C, the mean log IAP values at the longstanding steady state (≈log Ksp) were estimated for otavite [CdCO3] to be −11.62 ± 0.05–−11.79 ± 0.10 and −11.49 ± 0.07–−11.60 ± 0.07 with the ΔGfθ values of −671.81 ± 0.29–−672.78 ± 0.55 kJ/mol and −671.09 ± 0.43–−671.71 ± 0.38 kJ/mol, respectively. Generally, the log IAP values at the final longstanding steady states decreased non-linearly, and the estimated ΔGfθ values increased linearly with the increasing XCd of the (Ca1−xCdx)CO3 solid solution.
In the Lippmann diagrams constructed for the sub-regular (Ca1−xCdx)CO3 solid solutions with the estimated Guggenheim coefficients a0 = −0.84 and a1 = −3.80 or a0 = −1.12 and a1 = −3.83, the (Ca1−xCdx)CO3 solid solutions dissolved incongruently, moved progressively up to the quasi-equilibrium curve for otavite [CdCO3], and then along the quasi-equilibrium curves for otavite from right to left, approached the solutus curves and finally reached the minimum stoichiometric saturation curves for calcite [CaCO3]. The dissolution of the calcite–otavite solid solutions [(Ca1−xCdx)CO3] might be controlled by a quasi-equilibrium with respect to otavite. The considerably Cd-poor aqueous phases were finally in equilibrium with the CdCO3-rich solid phases.

Supplementary Materials

The following supporting information can be downloaded at: https://www.mdpi.com/article/10.3390/min12060756/s1, Table S1: Speciation reaction in the PHREEQC simulation; Table S2: Solubility constants (log K) of otavite at standard state; Table S3: Gibbs free energies of formation (ΔGfθ) of otavite; Figure S1: Position variation of the strongest peak (104) with the solid composition XCd: (a) before dissolution, and (b) after dissolution in the N2-degassed water for 270 d and (c) after dissolution in the CO2-saturated water for 270 d; Figure S2: SEM image of the (Ca1−xCdx)CO3 solid solutions before dissolution; Figure S3: SEM image of calcite (Cal-Ot-00) after dissolution for 270 d; Figure S4: Change of the Cd/(Ca+Cd) mole ratios of the aqueous phases during the dissolution of the (Ca1−xCdx)CO3 solid solution; Figure S5: Variation of the final aqueous phases with the solid XCd mole ratios during the dissolution of the (Ca1−xCdx)CO3 solid solutions in water; Figure S6: Change of log IAP (≈log Ksp) for the (Ca1−xCdx)CO3 solid solution with the CdCO3 mole fraction (XCdCO3); Figure S7: Change of the Gibbs free energy of formation (ΔGfθ) for the (Ca1−xCdx)CO3 solid solution with the CdCO3 mole fraction (XCdCO3); Figure S8: Estimation of the Guggenheim coefficients (a0 and a1) for the non-ideal (Ca1−xCdx)CO3 solid solution; Figure S9: Saturation index of the aqueous solution with respect to calcite and otavite during the dissolution of the (Ca1−xCdx)CO3 solid solution.

Author Contributions

Conceptualization, Y.Z.; Data curation, H.Y., C.M. and P.N.; Formal analysis, H.Y., P.N. and F.X.; Funding acquisition, Y.Z.; Investigation, C.M. and S.T.; Methodology, C.M.; Project administration, Y.Z.; Visualization, Z.Z. and Y.Z.; Writing—Original draft, Y.Z.; Writing—Review and Editing, Z.Z., L.Z. and Z.K. All authors have read and agreed to the published version of the manuscript.

Funding

The present research was funded by the National Natural Science Foundation of China (42063003) and the Science & Technology Planning Projects of Guangxi (2018GXNSFAA050044).

Data Availability Statement

Not applicable.

Acknowledgments

The manuscript has greatly benefited from insightful comments by editors and three anonymous reviewers. The authors would like to acknowledge the Science and Education Combined with Science and Technology Innovation Base of Guangxi Key Laboratory of Environmental Pollution Control Theory and Technology for the help in solid and solution analysis.

Conflicts of Interest

The authors declare no conflict of interest.

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Figure 1. Diffraction patterns of the (Ca1−xCdx)CO3 solid solutions: (a) before dissolution and (b) after dissolution in the N2-degassed water for 270 d and (c) after dissolution in the CO2-saturated water for 270 d.
Figure 1. Diffraction patterns of the (Ca1−xCdx)CO3 solid solutions: (a) before dissolution and (b) after dissolution in the N2-degassed water for 270 d and (c) after dissolution in the CO2-saturated water for 270 d.
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Figure 2. SEM images of the (Ca1−xCdx)CO3 solid solutions before dissolution.
Figure 2. SEM images of the (Ca1−xCdx)CO3 solid solutions before dissolution.
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Figure 3. SEM images of the (Ca1−xCdx)CO3 solid solutions after dissolution in the N2-degassed water for 270 d.
Figure 3. SEM images of the (Ca1−xCdx)CO3 solid solutions after dissolution in the N2-degassed water for 270 d.
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Figure 4. SEM images of the (Ca1−xCdx)CO3 solid solutions after dissolution in the CO2-saturated water for 270 d.
Figure 4. SEM images of the (Ca1−xCdx)CO3 solid solutions after dissolution in the CO2-saturated water for 270 d.
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Figure 5. Change of the aqueous pH, the total dissolved calcium, the total dissolved cadmium and the total dissolved H2CO30 during the dissolution of the (Ca1−xCdx)CO3 solid solutions in the N2-degassed water.
Figure 5. Change of the aqueous pH, the total dissolved calcium, the total dissolved cadmium and the total dissolved H2CO30 during the dissolution of the (Ca1−xCdx)CO3 solid solutions in the N2-degassed water.
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Figure 6. Change of the aqueous pH, the total dissolved calcium, the total dissolved cadmium and the total dissolved H2CO30 during the dissolution of the (Ca1−xCdx)CO3 solid solutions in the CO2-saturated water.
Figure 6. Change of the aqueous pH, the total dissolved calcium, the total dissolved cadmium and the total dissolved H2CO30 during the dissolution of the (Ca1−xCdx)CO3 solid solutions in the CO2-saturated water.
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Figure 7. Lippmann diagram for the non-ideal (Ca1−xCdx)CO3 solid solutions together with the plot of some stoichiometric saturation curves and the experimental data for dissolution in the N2-degassed water. The graph on the left is a zoom of the one on the right.
Figure 7. Lippmann diagram for the non-ideal (Ca1−xCdx)CO3 solid solutions together with the plot of some stoichiometric saturation curves and the experimental data for dissolution in the N2-degassed water. The graph on the left is a zoom of the one on the right.
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Figure 8. Lippmann diagram for the non-ideal (Ca1−xCdx)CO3 solid solutions together with the plot of some stoichiometric saturation curves and the experimental data for dissolution in the CO2-saturated water. The graph on the left is a zoom of the one on the right.
Figure 8. Lippmann diagram for the non-ideal (Ca1−xCdx)CO3 solid solutions together with the plot of some stoichiometric saturation curves and the experimental data for dissolution in the CO2-saturated water. The graph on the left is a zoom of the one on the right.
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Table 1. Summary of synthesis and composition of the calcite–otavite solid solutions [(Ca1−xCdx)CO3].
Table 1. Summary of synthesis and composition of the calcite–otavite solid solutions [(Ca1−xCdx)CO3].
Sample No.Volume of Precursor (mL)Amount of Precipitate (g)Cd/(Ca+Cd) Mole Ratio
2 mol/L
Ca(NO3)2
2 mol/L
Cd(NO3)2
0.5 mol/L
NH4HCO3
Precursor
Solution
Synthetic
Solid
Solid after Dissolution in Water
N2-DegassedCO2-Saturated
Cal-Ot-005.00.01000.960.000.000.000.00
Cal-Ot-014.50.51001.120.100.100.110.12
Cal-Ot-024.01.01001.150.200.210.210.24
Cal-Ot-033.51.51001.230.300.310.310.34
Cal-Ot-043.02.01001.290.400.400.410.45
Cal-Ot-052.52.51001.360.500.500.510.55
Cal-Ot-062.03.01001.440.600.600.610.66
Cal-Ot-071.53.51001.490.700.700.710.75
Cal-Ot-081.04.01001.560.800.790.810.84
Cal-Ot-090.54.51001.630.900.890.900.92
Cal-Ot-100.05.01001.681.001.001.001.00
Table 2. Analytical data and solubility for the calcite–otavite solid solutions [(Ca1−xCdx)CO3] in the N2-degassed water at 25 °C.
Table 2. Analytical data and solubility for the calcite–otavite solid solutions [(Ca1−xCdx)CO3] in the N2-degassed water at 25 °C.
SampleDissolution Time (h)pHConcentration# log IAPMean log IAPΔGfθ (kJ/mol)Mean ΔGfθ (kJ/mol)
Ca
(mmol/L)
Cd
(μmol/L)
HCO3+CO3
(mmol/L)
a (Ca1.00Cd0.00)CO350408.250.527 1.043−8.48−8.45−1129.84−1129.65
57608.230.536 1.153−8.45±0.04−1129.68±0.22
64808.240.562 1.195−8.41 −1129.43
b (Ca1.00Cd0.00)CO350408.290.511 1.035−8.45−8.42−1129.70−1129.48
57608.270.512 1.145−8.43±0.06−1129.57±0.30
64808.320.534 1.154−8.36 −1129.18
* (Ca0.90Cd0.10)CO350408.370.4450.0170.987−8.95−8.91−1080.14−1079.93
** (Ca0.89Cd0.11)CO357608.380.4540.0231.080−8.88±0.04−1079.76±0.21
64808.350.4720.0181.095−8.90 −1079.88
* (Ca0.79Cd0.21)CO350408.500.3780.0250.836−9.35−9.29−1034.88−1034.50
** (Ca0.79Cd0.21)CO357608.510.4220.0350.942−9.23±0.06−1034.19±0.38
64808.440.4380.0310.996−9.27 −1034.43
* (Ca0.69Cd0.31)CO350408.340.4260.0520.955−9.75−9.68−989.52−989.13
** (Ca0.69Cd0.31)CO357608.440.4230.0491.046−9.63±0.07−988.85±0.39
64808.450.4190.0361.049−9.66 −989.04
* (Ca0.60Cd0.40)CO350408.450.3720.0560.844−10.11−10.05−943.99−943.65
** (Ca0.59Cd0.41)CO357608.550.3650.0580.899−9.99±0.06−943.34±0.34
64808.510.3720.0500.916−10.04 −943.62
* (Ca0.50Cd0.50)CO350408.480.3370.0570.741−10.54−10.46−898.84−898.39
** (Ca0.49Cd0.51)CO357608.600.3380.0620.821−10.37±0.09−897.91±0.48
64808.540.3390.0510.833−10.46 −898.42
* (Ca0.40Cd0.60)CO350408.570.2610.0500.629−10.98−10.86−853.75−853.09
** (Ca0.39Cd0.61)CO357608.720.2890.0540.696−10.78±0.12−852.62±0.66
64808.510.3080.0710.791−10.82 −852.89
* (Ca0.30Cd0.70)CO350408.430.3230.0940.733−11.20−11.17−807.45−807.26
** (Ca0.29Cd0.71)CO357608.560.3300.0810.800−11.10±0.07−806.88±0.38
64808.530.3220.0640.804−11.20 −807.45
* (Ca0.21Cd0.79)CO350408.460.2210.0900.537−11.69−11.59−762.65−762.05
** (Ca0.19Cd0.81)CO357608.490.2310.1170.612−11.52±0.10−761.68±0.60
64808.540.2410.0900.645−11.55 −761.84
* (Ca0.11Cd0.89)CO350408.360.1400.2310.424−11.82−11.81−720.57−720.48
** (Ca0.10Cd0.90)CO357608.380.1400.2280.455−11.78±0.03−720.34±0.14
64808.480.1420.1700.448−11.82 −720.52
(Ca0.00Cd1.00)CO350407.83 5.3200.149−11.66−11.62−672.06−671.81
57607.84 5.4350.155−11.63±0.05−671.86±0.29
64807.83 6.4760.152−11.57 −671.52
(Ca0.00Cd1.00)CO350407.79 3.4610.151−11.89−11.79−673.33−672.78
57607.89 3.6210.147−11.78±0.10−672.70±0.55
64807.86 4.3950.152−11.71 −672.32
Note: a,b Duplication; * Bulk component before dissolution; ** Bulk component after dissolution. # Calculated with respect to the bulk component after dissolution.
Table 3. Analytical data and solubility for the calcite–otavite solid solutions [(Ca1−xCdx)CO3] in the CO2-saturated water at 25 °C.
Table 3. Analytical data and solubility for the calcite–otavite solid solutions [(Ca1−xCdx)CO3] in the CO2-saturated water at 25 °C.
SampleDissolution Time (h)pHConcentration# log IAPMean log IAPΔGfθ (kJ/mol)Mean ΔGfθ (kJ/mol)
Ca
(mmol/L)
Cd
(μmol/L)
HCO3+CO3
(mmol/L)
a (Ca1.00Cd0.00)CO350407.910.803 1.645−8.47−8.39−1129.76−1129.32
57607.970.817 1.716−8.38±0.08−1129.28±0.44
64808.050.801 1.668−8.32 −1128.93
b (Ca1.00Cd0.00)CO350407.900.825 1.741−8.44−8.37−1129.62−1129.22
57607.930.853 1.776−8.39±0.09−1129.33±0.52
64808.060.837 1.721−8.28 −1128.70
* (Ca0.90Cd0.10)CO350407.900.7740.0741.896−8.92−8.88−1075.24−1075.02
** (Ca0.88Cd0.12)CO357607.890.8220.0462.023−8.91±0.06−1075.17±0.37
64808.010.7940.0471.931−8.82 −1074.65
* (Ca0.79Cd0.21)CO350407.870.8770.0512.097−9.40−9.38−1020.85−1020.74
** (Ca0.76Cd0.24)CO357607.870.8700.0682.092−9.37±0.02−1020.70±0.11
64807.990.8160.0411.899−9.36 −1020.65
* (Ca0.69Cd0.31)CO350407.820.8890.1112.202−9.74−9.70−975.19−975.01
** (Ca0.64Cd0.34)CO357607.850.9220.1012.196−9.71±0.04−975.06±0.23
64808.020.8850.0632.005−9.66 −974.78
* (Ca0.60Cd0.40)CO350407.830.9380.1322.001−10.15−10.13−925.21−925.07
** (Ca0.55Cd0.45)CO357607.830.9160.1232.053−10.16±0.06−925.26±0.32
64807.980.9020.1031.964−10.07 −924.75
* (Ca0.50Cd0.50)CO350407.830.9320.1921.884−10.48−10.53−879.46−879.74
** (Ca0.45Cd0.55)CO357607.820.8800.1481.933−10.55±0.05−879.86±0.28
64807.920.8530.1031.878−10.55 −879.90
* (Ca0.40Cd0.60)CO350407.830.8370.1371.856−11.00−10.98−830.09−830.00
** (Ca0.34Cd0.66)CO357607.870.8420.1291.870−10.97±0.02−829.95±0.09
64808.030.8030.0821.780−10.98 −829.97
* (Ca0.30Cd0.70)CO350407.840.8750.1611.936−11.26−11.26−788.75−788.72
** (Ca0.25Cd0.75)CO357607.880.9020.1311.950−11.29±0.04−788.89±0.19
64807.990.8650.1291.796−11.22 −788.53
* (Ca0.21Cd0.79)CO350407.860.7630.3741.691−11.33−11.47−746.30−747.10
** (Ca0.16Cd0.84)CO357607.850.7890.2101.702−11.55±0.14−747.54±0.80
64807.960.7710.1741.606−11.53 −747.47
* (Ca0.11Cd0.89)CO350407.940.4810.4631.051−11.62−11.64−709.89−710.00
** (Ca0.08Cd0.92)CO357607.940.4620.4161.054−11.66±0.02−710.13±0.13
64808.040.4580.3631.014−11.64 −709.99
(Ca0.00Cd1.00)CO350407.66 9.6430.219−11.42−11.49−670.66−671.09
57607.65 8.5130.196−11.53±0.07−671.29±0.43
64807.73 6.9570.197−11.53 −671.31
(Ca0.00Cd1.00)CO350407.65 6.8940.214−11.58−11.60−671.60−671.71
57607.65 6.0230.201−11.67±0.07−672.09±0.38
64807.75 6.6360.186−11.56 −671.45
Note: a,b Duplication; * Bulk component before dissolution; ** Bulk component after dissolution. # Calculated with respect to the bulk component after dissolution.
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Ma, C.; Xu, F.; Zhu, Z.; Yang, H.; Nong, P.; Kang, Z.; Tang, S.; Zhang, L.; Zhu, Y. Dissolution and Solubility of the Calcite–Otavite Solid Solutions [(Ca1−xCdx)CO3] at 25 °C. Minerals 2022, 12, 756. https://doi.org/10.3390/min12060756

AMA Style

Ma C, Xu F, Zhu Z, Yang H, Nong P, Kang Z, Tang S, Zhang L, Zhu Y. Dissolution and Solubility of the Calcite–Otavite Solid Solutions [(Ca1−xCdx)CO3] at 25 °C. Minerals. 2022; 12(6):756. https://doi.org/10.3390/min12060756

Chicago/Turabian Style

Ma, Chengyou, Fan Xu, Zongqiang Zhu, Hongqu Yang, Peijie Nong, Zhiqiang Kang, Shen Tang, Lihao Zhang, and Yinian Zhu. 2022. "Dissolution and Solubility of the Calcite–Otavite Solid Solutions [(Ca1−xCdx)CO3] at 25 °C" Minerals 12, no. 6: 756. https://doi.org/10.3390/min12060756

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