Review on Chemistry of Water-Containing Calcium Carbonates and Their Transformations into Amorphous and Crystalline Carbonate Modifications

Abstract

Calcium carbonate (CaCO₃) is a dominant component of sedimentary rocks and biogenic structures, and is one of the most frequently studied inorganic compounds. It also plays a key role in preparing modern engineered materials. CaCO₃ has three well-known polymorphs, calcite, aragonite, and vaterite, and four solvatomorphs with diverse crystallographic arrangements, hydration states, reactivity, and stability. Its solvatomorphs include the variable water-containing amorphous calcium carbonate (ACC—CaCO₃·xH₂O) and the crystalline monohydrocalcite (MHC—CaCO₃·H₂O), calcium carbonate hexahydrate (ikaite—CaCO₃·6H₂O), and the recently reported hemihydrate (CCHH—CaCO₃·0.5H₂O). Here, we review the preparation, crystal structure, and properties of these solvatomorphs and discuss their mutual transformations.

1. Introduction

Calcium carbonate (CaCO₃) is one of the most important inorganic compounds, playing a key role in geology [1], biology [2,3], and industry [4,5,6,7]. It is prevalent in sedimentary rocks [8,9,10], biogenic structures [11,12,13], and even in modern engineered materials [14,15]. In spite of its simple composition, CaCO₃ has a variety of polymorphs and solvatomorphs. Polymorphs are anhydrous forms, whereas solvatomorphs are hydrated calcium carbonates, and these terms are used interchangeably in the text. The diversity in crystallographic arrangements and hydration states results in unique occurrences and distinct properties, including reactivity and stability [16]. In addition to the well-known calcite, aragonite, and less-known vaterite polymorphs, the hydrated forms, including the variable crystalline water-containing amorphous calcium carbonate (ACC—CaCO₃·xH₂O) [17] and the crystalline monohydrocalcite (MHC (also called hydrocalcite)—CaCO₃·H₂O) [18], calcium carbonate hexahydrate (ikaite)—CaCO₃·6H₂O [19], and the recently reported hemihydrate (CCHH—CaCO₃·0.5H₂O) [20], are the focus of intense investigations [2,21,22,23,24,25]. Each solvatomorph is characterized by a unique structure and hydrogen-bonding systems [24,26,27,28,29,30]. They differ not only in their crystalline water content, but in their long-range ordering [31,32,33], solubility [34,35,36,37,38,39], and transformation pathways [11,21,40,41,42,43,44,45,46]. Fundamental principles such as formation energy and stability of water-containing ACC are crucial for understanding its stability and mutual interconversion pathways among ACC, crystalline solvatomorphs, and anhydrous polymorphs [21]. Here, we review the occurrence, preparation, crystallography, spectroscopy, computational analysis, and thermodynamics of various carbonate forms [38,41,42,43,44,45,46,47,48,49,50]. We summarize and compare the properties of amorphous and crystalline calcium carbonate solvatomorphs and provide insights into their transformation routes.

2. Occurrence in Geological Environment and/or Biological Systems, and Preparation, Morphology, and Properties of Hydrated Calcium Carbonates

Calcite, aragonite, and vaterite are anhydrous compounds; however, studies of their atomic force microscopy (AFM) demonstrated that they may acquire adsorbed water layers when exposed to ambient and wet atmospheric conditions or aqueous environments [51,52,53]. At ambient conditions, the stable form is calcite, and the metastable vaterite and aragonite transform to calcite on contact with water [54].

2.1. Ikaite

The preparation of ikaite preceded its natural discovery. Ikaite was synthesized in 1865 by Pelouze in the reaction of a carbonic acid stream and lime water (Ca(OH)₂ solution) at around 0 °C [55], and the “ikaite” name was given when it was identified in the Ikka fjord (Grönland) by Pauly in 1963 [56]. Its water content was first determined as five moles [57]; however, Johnston et al. (1917) demonstrated that the material, prepared from a mixture of 0.1 M of CaCl₂ and K₂CO_3, buffered with KOH solution (2 g in 600 cm³) at 2 °C, contains six moles [58]. Interestingly, needle-like “ikaite” crystals were found by Iwanoff in a Paleocene marl even in 1906, close to Puławy, Poland [59]; however, its composition was believed to be CaCO₃·5H₂O. After the discovery of natural ikaite in the Ikka fjord, many other cold-place occurrences were identified, like the Antarctic ice (Suess et al. [9]), Mono Lake (California, USA) (Bischoff et al. [60]), the Scarisoara Ice Cave (Romania) (Onac et al. [61]), alkaline artificial riverbeds (Austria) (Boch et al. [62]), and sea sediments (Rogov et al. [50]). Other occurrences include biogenic materials (Larsen [47], Mikkelsen et al. [63], Swainson et al. [64], Rickaby et al. [65], Garvie et al. [66]), washed-rind cheese [67], and steel waste leachate (Bastianini et al. [68]). Ikaite is stable <4 °C, and above this temperature it starts to decompose into calcite. However, its decomposition was confirmed even at −20 °C by solid-state NMR studies [17]. Ikaite has been used as an indicator of cold paleotemperatures (Larsen [47], Swainson and Hammond [64], and Rickaby et al. [65]). Calcite pseudomorphs after ikaite are commonly called glendonites, which have been from marine clastic rocks, caves, and lacustrine sediments (Larsen [47], Tollefsen et al. [69], Németh et al. [70], Schultz et al. [71]), and their occurrence has been correlated with cold environments and has attracted the attention of paleoclimate researchers (De Lurio and Frakes [48], Rogov et al. [50]).

Ikaite can be prepared in many ways [21,58]. For example, following the method of Johnston et al. [58], Bischoff et al. mixed dropwise a 0.1 M solution of CaCl₂ and K₂CO₃ with a 0.04 M KOH solution at 0 °C and prepared 10 g of ikaite under 6–8 h [60]. Lázár et al. also followed the procedure described by Johnston et al. [58], and successfully prepared ikaite at pH 12.7 and at temperatures between 0 and 4 °C. They demonstrated that ikaite could also be synthesized by mixing 0.02 M Na₂CO₃ and 0.02 M CaCl₂ solutions, ageing the mixture at 0–4 °C for 20–25 min, and quickly vacuum filtering the precipitated material [21]. An alternative method for ikaite synthesis is based on alkalizing a cold (2 °C), nearly pH-neutral Ca(HCO₃)₂ solution with solid KOH [72]. Following the procedure, ikaite can be synthesized at pH 12.1, 10.3, and 9.1. Tollefsen et al. showed that ikaite can be prepared at 35 °C between pH 9.3 and 10.3 by mixing natural seawater and a Na₂CO₃-rich solution [69]. Hu et al. synthesized ikaite at temperatures between 0 and 4 °C and at pH between 8 and 10 by mixing 2.5 M CaCl₂·2H₂O, 0.5 M NaHCO₃, and 0.5 M NaOH solutions and seawater—a similar mixture, prepared from distilled water containing NaCl, MgCl₂, Na₂SO₄, KCl, SrCl₂, KBr, and H₃BO₃ in various ratios [73].

Ikaite crystals are commonly prismatic and characterized by pyramidal morphology (Figure 1A). In fact, the steep and spiky canted tetragonal pyramids are uniquely preserved in glendonites (Németh et al. [70], Schultz et al. [71]) and can be associated with the precursor ikaite. In addition to prismatic grains, ikaite could also occur as platy crystals. Lázár et al. [72] demonstrated that alkalizingto a nearly pH-neutral Ca(HCO₃)₂ solution with solid KOH results in flakes with lateral dimensions of 5–10 μm and thicknesses of 0.1–0.5 μm.

2.2. Monohydrocalcite, MHC

Monohydrocalcite or hydrocalcite, identified as CaCO₃·H₂O, was first synthesized in the reaction of Na₂CO₃ and CaCl₂ by Brooks et al. (used 2 to 1 molar ratios) [75] and Van Tassel et al. (used 4 to 1 molar ratios) [76] in 1950 and 1962, respectively. The synthesis required MgCl₂ presence > 200 ppm. MHC was first found in nature by Sapozhnikov et al. in Lake Issyk-Kul (Kyrgyzstan) [77] and later by Carlström in fish otoliths [78] in 1959 and 1963, respectively. MHC was also reported in a submarine ikaite tufa column around the Arctic Ikka Fjord, Greenland [79]. The formation of MHC was found in the air scrubbers of different air-conditioning plants [80] Garvie reported its occurrence from the decay of the saguaro cactus [81], Zaytseva et al. found in the microbialites from Laguna de Cisnes [82], and Krauss et al. noted that certain evaporitic settings and lacustrine environments produced a calcium carbonate deposit, of which diffraction patterns were not similar to calcite or aragonite [18], but agreed well with that of MHC, prepared by mixing sea water and K₂CO₃ [83].

Duedall et al. [84] detected a delayed precipitation of MHC after 8 h at 20 °C, and, as Nathana mentioned earlier [85], the diffraction patterns of this product do not agree with those of the expected vaterite [84]; therefore, Duedall et al. concluded that it was a new kind of material, MHC [84]. Vereshchagin et al. found that transition metal-containing solutions can facilitate the formation of MHC [86]. Rodriguez-Blanco et al. prepared MHC at 21 °C by mixing 1 M Na₂CO₃ solution with 700 mM CaCl₂ and 300 mM MgCl₂ solutions. They found that the MgCl₂ solution promotes the formation of MHC [87]. Kitajima et al. followed the same method, but performed the synthesis at 5 °C [88]. Zhang et al. prepared MHC in the following way: they mixed Na₂Si₃O₇ (sodium silicate) solution (varying between 1 and 10 v/v%) with 0.05 M Na₂CO₃ solution and added 1M HCl solution to the mixture, that resulted in a gel. Into this gel, a 0.05 M CaCl₂ solution was injected (at pH = 10.5), and, due to the diffusion of the Ca²⁺ ion in the silicone gel, it resulted in MHC crystals after a day [89]. Kralj et al. prepared MHC under a night by mixing 0.3 g K₂CO₃ with a 250 mL seawater-similar mixture, prepared from triple-distilled water containing 29.2 g NaCl, 5.0 g MgCl₂, 1.1 g CaCl₂, and 0.024 g SrCl₂ in 1000 mL triple-distilled water [90].

Taylor reported that MHC crystals, precipitated from saline lake water, are hard pelletal aggregates with a sponge-like texture [91]. Dahl and Buchardt [79] demonstrated that MHC from the Arctic Ikka Fjord is associated with ikaite and occurs in three different forms as (1) spherulites, (2) subhedral to euhedral crystals, and (3) zoned euhedral crystals. According to Kimura et al. [74], synthetic MHC either consists of well crystalline spherical particles (Figure 1C) or poorly crystalline aggregates of twisted spindles (Figure 1D).

2.3. Calcium Carbonate Hemihydrate, CCHH

The discovery of calcium carbonate hemihydrate (CCHH, CaCO₃·0.5H₂O) [20] represents a critical missing link in the hydration–dehydration transitions of calcium carbonate containing hydrated phases. First, it was synthesized by Zou et al. in magnesium-rich aqueous systems [20], and recently, it was reported in fresh coral skeleton and nacre surfaces [22]. The high-resolution X-ray powder diffraction (HRXRPD) and automated electron diffraction tomography (ADT) measurements revealed that CCHH has a monoclinic lattice (space group found to be space group P2₁/c) and the 0.5 molar crystalline water was proved by TGA studies [20].

The stability and the electronic and mechanical properties of CCHH were determined by DFT (PBE + TS) calculations [92]. The Gibbs energy was found to be −0.052 eV/atom, for the following proposed equation:

CaCO₃ + 0.5 H₂O = CaCO₃·0.5H₂O

(1)

which is higher than the −0.055 and −0.139 eV/atom values, calculated for the MHC and ikaite, respectively [92]. The compressibility of the CCHH is large due to the monoclinic structure and is found to be 109.28 and 129. 65 GPa under the external uniaxial stress along the [001] and [010] directions, respectively, whereas, in the case of the ikaite and MHC, these values were only 48.32 and 88.97 for [001] and 91.21 and 80.81 GPa for [010], respectively [92]. The Young’s modulus was calculated to be 60.51 GPa (these values were 36.56 and 91.28 GPa for MHC and ikaite, respectively) [92].

The synthesis of CCHH is based on mixing a 5–100 mM Na₂CO₃·10H₂O solution with a solution containing MgCl₂·6H₂O and CaCl₂·2H₂O (Mg²⁺ to Ca²⁺ ratio was 5 to 1, ultra-pure water) at room temperature. The precursor ACC transforms to needle-like CCHH crystals after 2000 s if the solution is stored in an air-proof container [20].

2.4. Amorphous Calcium Carbonate (ACC)

One of the most important hydrous CaCO₃ is ACC. It lacks long-range ordering and is characterized by variable water content. In nature, ACC was first recognized during the deposition of the crystalline core material of Sponge-Spicules by Minchin in 1909 [93]. In 1912, Kendall emphasized its high instability as he could not determine its water content, but measured the composition of calcite [94]. Dorfmüller found that increasing alkalinity stabilizes the amorphous state [95]. Schmidt [96], Prenant [97], Rinne [98], and Odum [99] found an amorphous phase, which was evidenced by a diffuse scattering pattern of spicules in marine animals. Furthermore, Towe et al. found a similar material by studying the evaporation of seawater [100]. ACC was reported from cave bacteria [101], slime moulds [102], Pyura pachydermatina tunicates [103], and Biomphalaria glabrata snail embryos [104]. Scientists recognized that many organisms utilize ACC as a transient precursor, which later transforms into a crystalline phase [105,106,107,108,109,110,111,112]). The finding and study of ACC provided new insights into biomineralization [113], sediment formation [114], and atmospheric processes [115,116,117,118,119]. The transformation of the CaCO₃ polymorphs plays an important role in biomineralization, and ACC became one of the most commonly used paleoindicators, e.g., of ultra-low-temperature [8,120,121,122,123]. Studying carbonate transformation provides key insights into pressing contemporary issues such as climate change and CO₂ utilization, including carbon capture [124,125,126] and storage (CCS) technologies [127,128,129].

ACC is a precursor material which rapidly transforms to carbonate polymorphs and solvatomorphs depending on the synthesis conditions. However, its transition can be delayed, and ACC can be stabilized. Successful methods are summarized by Liu et al. [130]. Konrad et al. prepared ACC by mixing equimolar 0.25–0.25 M CaCl₂ and Na₂CO₃ solutions at pH > 10, and then rapid filtering and freeze-drying [131]. Lázár et al. [72] reported that 100 nm-sized ACC particles could be prepared by mixing K₂CO₃ and CaCl₂ solutions, and rapidly filtering and washing them with cold (2 °C) acetone (Figure 1D). Furthermore, ACC nanoparticles can be prepared in an aerosol-based way by spraying 10–100 µm-sized droplets into CO₂-enriched air from an aqueous 23 mM Ca(OH)₂ solution [132,133]. Chen et al. [134] reported an additional gas diffusion method, in which ammonia and CO₂ vapours, generated from ammonium carbonate, diffuse into abs. EtOH and a 0.0136 mM CaCl₂·2H₂O solution, leading to the formation of ACC. However, this method was found to be poorly reproducible and resulted in a limited amount of material [134]. Additives such as Mg²⁺ could stabilize ACC [41], and we will review the stabilization process in Section 4. Selected synthetic routes are summarized in

Table 1. Possible synthetic routes of hydrated CaCO₃ solvatomorphs (ikaite, monohydrocalcite, CaCO₃ hemihydrate, and amorphous CaCO₃). We highlighted the presence of other ions (like Mg²⁺) due to their importance.

Name	Method 1	Method 2	Method 3
Ikaite	0.1 M CaCl2 + 0.1 M K2CO3 + 0.04 M KOH; T = 0 °C [60]	0.02 M Na2CO3 + 0.02 M CaCl2; pH = 12.8; T = 0–4 °C; t = 20–25 min [21]	2.5 M CaCl2·2H2O + 0.5 M NaHCO3 + 0.5 M NaOH + seawater-similar mixture (presence of, e.g., PO43−), T = 0 and 4 °C [73]
Monohydrocalcite (MHC)	700 mM Na2CO3 + 700 mM CaCl2 + MgCl2 (or CoCl2 2H2O [86]); T = 5 or 21 °C [87,88]	0.05 M Na2CO3 + sodium silicate gel with 0.05 M CaCl2; pH = 10.5 [89]	0.3 g K2CO3 + 250 mL seawater-similar mixture [90]
Calcium carbonate hemihydrate (CCHH)	5–100 mM Na2CO3 ·10H2O + MgCl2·6H2O + CaCl2·2H2O (Mg2+ to Ca2+ ratio was 5 to 1); T = 25 °C; t = 4200 s [20]	------	------
Amorphous Calcium carbonate (ACC)	0.25 M CaCl2·2H2O + Na2CO3, pH~10; T = 20 °C, filtration and freeze-drying [131]	23 mM Ca(OH)2 solution, microfluidic spray-drying [132,133]	Gas diffusion of NH3/CO2 through 0.0136 mM CaCl2·2H2O in abs. EtOH, pH~8 [134]

.

ACC commonly occurs as nm-sized spheres (Figure 1B). However, Lázár et al. [72] demonstrated that rapid dehydration of ikaite by organic solvents and vacuum pumping at room temperature, and by increasing the temperature from 5 to 30 °C within 1 min, results in micron-sized, porous ACC grains that preserve the prismatic ikaite grain morphology.

3. Spectroscopic and Structural Studies on Hydrated Calcium Carbonates

3.1. IR and Raman Spectroscopy

Raman and IR spectroscopies are excellent techniques to study the different polymorphs and solvatomorphs of CaCO₃ (Figure 2). These techniques can detect the vibrational modes of carbonate and water (

Table 2. The bands in the IR and Raman spectra of the different hydrated CaCO₃ solvatomorphs (ikaite, monohydrocalcite, CaCO₃ hemihydrate, and amorphous CaCO₃). Reproduced from Coleyshaw et al. [136] and Zou et al. [20].

Modes	Ikaite [136]		MHC [136]		ACC (Hydrated)		CCHH [20]
	IR	Raman	IR	Raman	IR [138]	Raman [139]	IR	Raman
υas,OH	3543, 3502, 3468, 3404, 3361, 3119	3423, 3182	3400, 3327	3425, 3326	3401, 3130	3430	3379	3382
υs,OH	3216	3240	3236	3224	3226	---	3290	3285
δs,HOH	1673, 1644, 1616	---	1700	---		1650	1659	1661
υas,CO (υ3)	1425, 1411	1483	1492, 1401	---	1485	1540, 1460, 1390	1524, 1490, 1423, 1392,	1579, 1483, 1423
υs,CO (υ1)	1085	1072	1063	1069	1064	1077	1096	1102
π,CO (υ2)	876	873	872	876	876, 856	868	866, 860	862
δas,CO (υ4)	743, 720	722	762, 698	723, 699	698	723, 698	723, 692	731, 700

), confirming that while local arrangements resemble those of crystalline calcium carbonate, the overall structure is disordered [135]. For example, ikaite and MHC can easily be distinguished by IR and Raman spectroscopy. The υ_s (O-H) stretching modes of ikaite appear as multiple bands between 3550 and 3200 both in IR and Raman spectra [136], whereas MHC gives rise to three well-separated υ_OH modes at 3400, 3327, and 3236 cm⁻¹ in the IR spectrum (Table 2) [136]. Coleyshaw et al. found well-resolved Raman bands for ikaite and MHC at 3423, 3240, and 3182 cm⁻¹ and 3425, 3326, and 3224 cm⁻¹, respectively [136]. CO₃²⁻ modes are positioned between 1500–1400 cm⁻¹ (υ_as,_CO) and 1100–1050 cm⁻¹ (υ_s,_CO), at ~870 cm⁻¹ (π,_CO), 760, and 725–690 cm⁻¹ (δ_as,_CO) [136] for both compounds. The IR and Raman data of hydrated CaCO₃ compounds are listed in Table 2 [136].

Effenberger and Coleyshaw et al. showed that the IR and Raman studies can be used to identify and distinguish MHC from ikaite and CCHH (Table 2) [136,140,141]. Solid-state NMR studies of MHC confirmed the presence of H₂O and CO₃²⁻ ions in different chemical environments [17].

3.2. Structure of Ikaite

Ikaite has monoclinic (C2/c) flat, platy, or tabular crystals with a prismatic habit, both at −120 °C (Dickens et al. [142]) and −25 °C (Hesse et al. [143]). Its main crystallographic parameters are summarized in

Table 3. Crystalline parameters of the different hydrated CaCO₃ solvatomorphs (ikaite, monohydrocalcite, and CaCO₃ hemihydrate).

Name	Calcium Carbonate Hexahydrate [143]	Monohydrocalcite [140]	Calcium Carbonate Hemihydrate [20]
Acronym	Ikaite	MHC	CCHH
Formula	CaCO3·6H2O	CaCO3·H2O	CaCO3·0.5H2O
Shape of the crystalline	flat, platy, or tabular crystals with a prismatic habit	thin, sometimes rounded or plate-like crystals	needle-like morphology
Crystal System	Monoclinic	Trigonal	Monoclinic
Space group	C2/c	P31	P21/c
Lattice parameters	a = 8.78 Å b = 8.28 Å c = 10.88 Å β = 109.59°	a = 10.554 Å b = 10.554 Å c = 7.454 Å	a = 9.33 Å b = 10.44 Å c = 6.16 Å β = 90.5°
Unit Cell Volume (Å3)	~743.7	~751.6	~600.0

in comparison to MHC and CCHH.

The structure of ikaite is composed of hydrated calcium carbonate chains, aligned along the axis a and linked together by a hydrogen bond system [144]. The thermal expansion of the unit cell and its relationship with the hydrogen-bonding system of ikaite were studied by several groups. Swainson et al. used neutron powder diffraction between 4 and 270 K [64], Lennie et al. applied synchrotron X-ray powder diffraction (S-PXRD) [144] between 114 and 293 K, and Tateno et al. performed Single-Crystal X-ray Diffraction (SC-XRD) [145] studies between 223 and 263 K. Temperature increase results in anisotropic unit cell expansion of ikaite along the a and c axes (Table 3). The largest structural modification occurred near the dehydration temperature of the ikaite, which suggests that the water, and especially the H-bonding system, is a strong cohesive force in the structure [145]. Chahi et al. refined the structure parameters of ikaite by DFT calculations. They have found that the a, b, and c cell parameters increased by +1.7, +0.4, and +0.03% applying DFT-PBE calculation and decreased by −1.1, −0.5, and −1.4% applying DFT-D2 calculation as a result of the differences in the van der Waals interactions [146]. Rietveld structural refinement of PXRD data, applying the structure model determined by Hesse [143], resulted in Ca-O_carbonate distances of 2.43–2.44 Å [66,75]. The strong anisotropic thermal expansion suggested the H-bonds play an important role in the structure; therefore, neutron diffraction measurements were performed to determine the positions of the H(D) atoms in the structure. The D-O distances varied between 1.761 and 1.962 Å in the ikaite structure (Figure 3A), and the hydrogen bond distance between the D3 and O2 atoms was found to be 2.508(9) Å [64,144].

Further neutron diffraction studies on ikaite showed the precise locations of hydrogen atoms within the lattice, clarifying the unique arrangement of water molecules [17,147]. The presence of six water molecules results in a looser, more open structure compared to anhydrous carbonates, and they are responsible for the high reactivity and sensitivity to temperature and anisotropic thermal expansion. The structure of ikaite was refined with DFT calculations as well, in a detailed manner [146,148]. Costa et al. performed DFT-GGA-PBE calculations and found that the band gap of ikaite did not change compared to calcite, but the band type changed to direct from indirect (Figure 4) [148]. The complex electronic structure of calcite was simplified upon hydration, and only two flat bands were formed very close to each other in the case of MHC and a single band in the case of ikaite (Figure 4B,C) due to the O_2p states of crystalline water. The simplification of the electronic structure results from a structural rearrangement due to the entrance of crystalline water [148]. Furthermore, this simplification implies that water molecules more strongly and directly influence the electronic structure of ikaite compared to MHC [148]. The calculated elastic data are in good agreement with the experimental elastic parameters obtained by Sekkal et al. [28]. In the case of ikaite, a smaller (5.069 eV instead of 5.14 eV) value was obtained by Zhou et al. by DFT-PBE + TS calculations [92].

3.3. Structure of MHC

MHC adopts a trigonal lattice (space group P3₁) similar to calcite (Table 3) [140,149,150]. Its structure is similar to that of ikaite, as the Ca²⁺ ion is surrounded by eight oxygen atoms, i.e., can be considered as a distorted tetragonal antiprism, and shares two O-O edges between the Ca coordination polyhedron and a carbonate group (Figure 3B). A primitive trigonal supercell was also detected but could not be determined properly; the proposed space group was P3₁ with a = 10.5536 Å and c = 7.5446 Å. The Ca-O_carbonate distances are between 2.41 Å and 2.47 Å, and the two Ca-O(w) distances are 2.49 Å (Figure 3B) [149].

Swainson et al. performed X-ray and neutron powder diffraction studies and Rietveld structure refinement on natural MHC samples from South Australia [149]. The P3₁ structure model consisted of three independent networks [149]. There was one water molecule and one carbonate group (Figure 3B) in each network, and every hydrogen atom of water molecules belonged to two different hydrogen bond systems, including a linear one (H3···O4) with one carbonate ion and a second one showing a bifurcated hydrogen (H3···O5 and H3···O6) bond to another carbonate ion (Figure 3B) [149]. The same interaction was detected by Makovicky, who described the structure of MHC as consisting of rods with three different Ca²⁺ ions, accompanied by spiral arrangements of CO₃²⁻ ions and crystalline water molecules [151]. These alternating carbonate environments in MHC were also detected by ESR (electron spin resonance) and 1H ENDOR (electron-nuclear double resonance) (accompanied by TGA) studies performed by Callens et al. [151]. The water molecule in the structure of MHC is coordinated in a manner that facilitates extensive hydrogen-bonding interactions with surrounding carbonate oxygens, thereby modifying the local geometry and overall lattice dynamics. The lattice and structural relations were studied by Demichelis et al. with the DFT computation method at the PBE0 level of theory, using all-electron Gaussian-type basis sets [152,153]. The calculated structural parameters were in good agreement with the experimental ones (Table 3 [149]), e.g., the 8-fold Ca coordination sphere and the shortest distance between two water molecules was at least 4 Å; each water molecule has two H-bonds with two different carbonate ions [152,153]. The calculation, similar to the measurement, also showed that there was a three-fold rotational axis parallel to the c lattice parameter, which created a triangular motif projected on the plane perpendicular to the [001] direction [149]. However, the calculation indicated that the symmetry of the “helical” chains, similar to what occurs in vaterite, could rather be described with the space group of P3₁ or P3₂ [153]. The occurrence of these helical chains might explain why ACC transforms into calcite or aragonite via MHC and vaterite intermediate phases [153]. The nature of van der Waals interactions was also studied by DFT-GGA-PBE calculation, and found that the presence of water increased the main band gap of MHC by 0.4 eV relative to the anhydrous structure (Figure 4) [148]. The structural parameters of the MHC were refined by DFT-PBE (standard) and -D2 (dispersion-corrected) methods [146]. Calculation demonstrated that the a, b, and c cell parameters increased by +0.7, +0.7, and +0.9% applying the DFT-PBE method and decreased by −0.4, −0.4, and −0.09% applying DFT-D2 calculation as a result of the differences in the van der Waals interactions [146]. The values of elastic moduli of MHC were reported to be lower than those of calcite, which is explained by the difference in the strength of the ionic (calcite) and hydrogen (hydrated phase) bonds [146]. Furthermore, MHC has greater elastic moduli, acoustic wave velocities, and Debye temperature than ikaite due to the distinct complex hydrogen-bonding network [146].

The electronic properties of MHC were determined by GGA-PBE (Costa et al., Figure 4B) [148] and PBE + TS calculations (Zhou et al.) [92], and the slightly different band gaps of 5.53 and 5.366 eV, respectively, were explained by the different O_2p states [92,148].

3.4. Structure of CCHH

The CCHH has needle-like monoclinic (P2₁/c) crystals (Table 3), and in its structure, each Ca²⁺ ion is surrounded by eight oxygen atoms, and the Ca-O_carbonate distances were similar to those of other solvatomorphs and varied between 2.09 and 2.68 Å (Figure 3C) [20]. This structural arrangement contributes to the unique hydration state of CCHH, positioning it as an intermediate phase among the amorphous or fully hydrated states (such as ACC and ikaite) and the anhydrous crystalline polymorphs (like calcite). Under ambient conditions, CCHH is metastable; however, the presence of Mg²⁺ ions can stabilize it. The orientation of the water molecules in the structure remained unresolved [20], which prompted Aufort et al. to study the structure of CCHH with Force Field-Based (FF) and PBE0-DC (DFT) geometry optimization methods. Aufort et al. determined CCHH has an orthorhombic Pbcn space group instead of the monoclinic P2₁/c [137], and the crystalline water forms hydrogen-bonding interactions. This water lies within channels parallel to the b-axis (Figure 3C), unlike ikaite, where water molecules form part of the rigid coordination environment, or monohydrocalcite, where water is more loosely associated [137]. Aufort et al. suggested that some Mg substitution resulted in β angle distortion and lowered the symmetry compared to pure CCHH [137]. The electronic properties of CCHH were determined by DFT (PBE + TS) calculation, and found that the band gap of the O_2p orbital was around 4.571 eV, which was much smaller than those in ikaite and MHC [92].

3.5. Structure of ACC

X-ray, electron, or neutron diffraction studies demonstrated that only a short-range order can be found in the ACC structure [96,97,98,99]. In spite of the long-range periodicity, X-ray absorption near-edge structure (XANES) and extended X-ray absorption fine structure spectroscopy (EXAFS) demonstrated that ACC has an MHC and anhydrous CaCO₃ similar coordination environments, and water molecules occur in the close proximity to Ca²⁺ and CO₃²⁻ clusters [31,32,33]. The oxygen K-edge and calcium L-edge EXAFS data showed a close structural relationship among the CaCO₃ polymorphs, ACC, and MHC (Figure 5) [33].

The CaCO₃·xH₂O formula of ACC could be described as Ca(OH)(HCO₃), but the presence of CO₃²⁻ ion instead of HCO₃⁻ was proved by solid-state ¹³C NMR studies. Only one peak appeared at the NMR spectra between 166 and 174 ppm (carbonate ion range) [17]. The same conclusion was given by the ¹H and the ¹³C CP MAS studies [17]. The chemical shift in ¹³C NMR of ACC is dissimilar to any other known CaCO₃ polymorphs and solvatomorphs, and proton NMR shows a peak between 5.0 and 5.6 ppm, whereas the proton of HCO₃⁻ is expected above 13 ppm [154]. Molecular modelling (MM) and molecular dynamics (MD) studies were also used to determine the basic structural features of ACC [29,155,156,157,158,159,160,161,162,163]. Bushuev et al. used the MD method and found that the Ca-Ca distances are around 5–7 Å [29], which agreed well with previous findings [164]. Jensen et al. determined the two Ca-O and Ca-Ca distances to be 2.4, 4.0, and 6.2 Å, respectively, by neutron and X-ray total scattering techniques combined with MD studies. The Ca-O_carbonate peaks showed two coordination spheres centred at 2.37 and 4.11 Å [165], similar to what was found by Goodwin et al. [157]. Michel et al. determined that the short- and intermediate-range orders were not extended beyond 15 Å, and the NMR studies showed that most of the hydrogen in ACC occurs as structural water, and the amount of rigid O-H groups was about 7 (±3)% [154]. All these values are similar to the other carbonate phases, except ikaite, which lacks the 4.0 Å Ca-Ca distance [143,154].

The variable water content within ACC is one of its hallmark features and is believed to be a critical factor in its role as a precursor in biomineralization processes. The investigation of its transformations demonstrated the strong influence of pH, temperature, pressure, presence of additives, water content, and mechanical conditions [112,166,167,168,169,170].

Bushuev et al. [29] and Jensen et al. studied the hydrogen bond system of ACC by neutron and X-ray total scattering techniques and molecular modelling (empirical potential structural refinement (EPSR)) [165]. The Ca-O_water distances showed one coordination sphere at 2.43 Å, which was a little bit longer than that of ikaite and MHC, whereas the Ca-H_w distance gave rise to a single broad coordination sphere at 2.9 Å [165]. The n = 0.5 and 1.1 values for ACC·xH₂O, 2.7 and 10% of the water molecules were calculated, respectively, and they were not bound directly to the Ca²⁺ ion, but all the water molecules were coordinated to the carbonate oxygens. Accordingly, two types of hydrogen bonds were identified in ACC: one between the O_carbonate and water, and another one between the water molecules (Figure 6).

All water molecules have hydrogen bonds with at least one carbonate ion [29,165]. It was concluded that the hydrogen bonds between the O_carbonate and H_water are more preferable, calculating x = 1 for CaCO₃·xH₂O [16,142,143], than between water molecules. Additional calculations showed that 63% and 20% of the hydrogen takes part in the carbonate–water and water–water hydrogen bonds, respectively, and these values decreased to 78% and to 10% upon dehydration, calculating x = 1 to x = 0.5 for CaCO₃·xH₂O [165]. These results agreed well with the NMR studies [17,154]. However, there were controversies between Bushuev et al. [29] and Jensen et al. [165] regarding the long-range ordering of the H-bond network in the CaCO₃·xH₂O systems. Bushuev et al. proposed the presence of water clusters with large numbers in the channels in both ACCs (Figure 6C) [29] (the same was proposed by Gale et al. [156], Goodwin et al. [157], and Innocenti et al. [161]), whereas, Jensen et al. found ca. 220 water molecules (about 80–90% of all water H₂O) in clusters for x = 1.1, and not more than 40 water molecules were found in the clusters for x = 0.5 in the CaCO₃·xH₂O system [165].

Katsikopoulos et al. studied the incorporation of foreign ions into calcium carbonate structures, found that Co^II ions enter calcite and aragonite, and reported ACC formation at Co^II/Ca^II = 5/1 [171]. Shuseki et al. reported that the Sr²⁺ ions have an influence on the ACC structure [172], and X-ray and neutron diffraction studies aided by MD simulation revealed that the Sr-containing ACC has a partially similar structure to MHC, and, upon the removal of the foreign ion, the structure transformed into calcite [172]. Molnár et al. [173] studied the incorporation of Sr²⁺ and Ba²⁺ in ACC, reported the partition coefficients (D_Ca, D_Sr, and D_Ba) decreased with increasing ionic radius between the solution and the solid, and demonstrated that Sr²⁺ and Ba²⁺ cations significantly modified the structural properties of ACC, and the changes in short-range structure resulted in exceptional kinetic stability of the ternary mixed amorphous Ca–Sr–Ba carbonates.

Lopez-Berganza et al. studied the influence of the amounts (n) of water molecules in the clusters on the stability of CaCO₃·xH₂O compounds between n = 1 and 18 with a combination of Monte Carlo simulation and DFT calculation (DFTB+ and B3LYP) [160]. They found that incremental binding energy (ΔE) became less negative and reached a plateau when n > 11; therefore, water up to n = 12 could fit into the first hydration shell. The simulation proved that for n = 18, 12 water molecules were coordinated to the CaCO₃ molecules, and the rest were found in the second hydration shell [160]. Furthermore, incremental binding free energy (ΔG) was close to zero (and even larger) when n ≥ 15, thus larger clusters were thermodynamically unfavourable in size [160]. The Ca-O_c distance increased with increasing n, and for n > 9, it varied between 2.305 and 2.585 Å, which was in good agreement with the Ca-O_c distances mentioned previously [160].

4. Transformations of Hydrated Calcium Carbonates

The stability of each calcium carbonate polymorph is governed by a combination of intrinsic lattice energy and extrinsic environmental factors such as temperature, pressure, pH, and chemical composition of the surrounding solution. Calcite is widely recognized as the most thermodynamically stable polymorph under ambient conditions [16]. Its low solubility and high resistance to dissolution explain why it occurs in long-term geological processes [174]. The stability of aragonite is profoundly influenced by environmental conditions (T, pH) and the presence of magnesium or strontium (e.g., ionic strength) [175,176]. Furthermore, in alkaline (above pH = 8.3) conditions, aragonite is more likely to form and persist, whereas a low pH (most typically below pH = 7.5) can promote its dissolution and transformation to calcite [177]. Vaterite transformation to calcite was observed to be faster than aragonite crystallization to calcite. Once the local environment, such as pH or temperature, changes, the disordered structure of vaterite can quickly reorganize into calcite [178].

In contrast, the calcium carbonate solvatomorphs, including ikaite and ACC, were mostly observed only transiently during rapid precipitation or biomineralization, and they eventually transform into the more stable calcite or aragonite forms [110]. Monohydrocalcite exhibits intermediate stability, and its formation is favoured under conditions of moderate saturation and in environments with elevated magnesium-to-calcium ratios, where the structural water in the lattice stabilizes it relative to the anhydrous calcite [179]. The transformation from ikaite to calcite involves a substantial loss of crystalline water and occurs via ACC (Lázár et al. [72]). This dehydration process is accompanied by significant changes in the vibrational spectra, as captured by IR and Raman measurements, and has been modelled extensively using both experimental kinetics and theoretical approaches [180]. Synthetic conditions strongly influence the morphology of calcium carbonate polymorphs and solvatomorphs. For instance, Sand et al. found that not only the purity of ethanol used, but agitation of the solution (gentle or vigorous) influenced the shape of the crystalline products. In the following subsection, we review the possible transformation routes of each hydrated CaCO₃ and their mutual transformation (Figure 7).

4.1. Transformation of Ikaite

Ikaite is stable only under low-temperature or high-pressure conditions (~300 MPa) [21,73,181,182,183]. Johnston et al. found that it slowly decomposed at and above 0°C in different organic liquids (diethyl ether, benzene) and transformed into calcite within two days.

Lázár et al. [72] reported that the preparation conditions, such as pH, temperature, Mg-content, and the media, e.g., air or solution, strongly influenced the transformation time and the structure of the final carbonate. For example, ikaite prepared at pH 10.3 and placed on a glass slide was stable for several weeks at 20 °C, whereas it quickly (within 4 h) transformed calcite or vaterite in solution. In clove oil, it was stable for months if the crystals were kept separately [58]. Slight warming led to rapid dehydration and transformation into calcite, often preserving the original crystal morphology in a pseudomorph calcite phase known as glendonite, which could be used to infer cold paleoclimate conditions [73,183].

Ikaite typically forms below 7 °C in sediments [114] and is generally prepared between 0 and 10 °C by direct reaction of carbon dioxide and calcium oxide or calcium carbonate solutions/suspensions [21]. Tollefsen et al. found that ikaite was formed even at 35 °C when pH was between 9.3 and 10.3, and the starting materials were a mixture of natural seawater and sodium carbonate-rich solution [69]. The kinetics of the ikaite transformation have been the subject of many studies, employing both laboratory experiments and field observations in polar and deep-sea environments. Purgstaller et al. studied the formation and transformation of ikaite in various pH, T, and Ca²⁺ to Mg²⁺ ratios [184]. Strohm et al. studied how phosphate could influence the growth of ikaite crystals [185]. Brečević et al. found the solubility of ikaite (−logK_s°) at 25 °C was 7.1199 [16], which is higher than the previously determined value (6.62) by Clarkson et al. [186]. Lázár et al. [72] reported that ikaite transformation in solution and in air occurs via ACC. Magnesium has a stabilizing effect on ikaite, which is similar to what was found for other hydrated calcium carbonates (MHC and CCHH) [187]. The transformation routes between the calcium carbonate polymorphs and solvatomorphs are complex and influenced by both kinetic and thermodynamic factors. In many natural systems, ACC is an initial transient phase deposited quickly under supersaturated conditions. Over time, ACC transforms into crystalline phases such as vaterite or directly into calcite, depending on the presence of inhibitors (such as organic macromolecules or specific ions) and the prevailing environmental conditions. Tang et al. [188] and Shaikh et al. [189] showed that ikaite can transform into vaterite, and Lázár et al. [21] reported the occurrence of two ACC types that differ in their mode of formation, morphology, particle size, water content, and stability, and found that ikaite formation starts with ACC (I) and ikaite transforms to calcite via ACC (II).

Ikaite transforms into vaterite and calcite at room temperature [189] during heating [136,186,188], or upon dropping it into boiling water [75,105]. Lázár et al. [72] demonstrated that ikaite can also transform into aragonite if the ikaite parent solution contains some Mg²⁺ (26 mg/L) and found that ikaite transforms into calcite via the nanosized ACC (I). However, micron–sized ACC (II) forms if ikaite is treated with polar organic solvents (MeOH, abs. EtOH, n-propanol, and DMSO) [21] and during vacuum pumping at room temperature. ACC (I) is spherical and contains 1.12 mol adsorbed and 0.26 mol chemically bonded water, whereas ACC (II) preserves the morphology of ikaite and is characterized by 0.42 mol absorbed and 0.35 mol chemically bonded water [21].

Sekkal studied different models (Raiteri’s and Xiao’s) of the surface stabilities of ikaite and found that (001) was the most stable surface with 0.2 J/m² surface energy and −2.57 eV attachment energy, which is lower than the (10–14) surface energy of calcite (0.51 J/m² and −8.69 eV) [28]. In contrast, (100) and (010) surfaces of ikaite were found to be 0.52 J/m² and −8.19 eV, and 0.37 J/m² and −5.79 eV values, respectively [28].

4.2. Transformation of MHC

The thermodynamic data of MHC were determined by Hull et al. Its free energy of formation (ΔG_f°) and the standard enthalpy of formation (ΔH_f°) were found to be −325.430 ± 270 cal mol⁻¹ and −358.100 ± 280 cal mol⁻¹ [190], respectively. Its solubility product (−log(K_s°) was reported as 7.60 ± 0.03 [190]. However, according to Brečević et al. [16], it is 7.050 ± 0.029 at 15 °C. Krajel et al. [90] found that the different accompanying ions influenced its solubility; the lgK values changed from 7.1772 to 7.4890 on increasing the temperature from 25 °C to 50 °C [90]. Liu et al. showed that the MHC (prepared from artificial seawater), in the absence of any additives, transforms into aragonite over 72 h in solution, even at 4 °C, and then to calcite at the end [191].

The structure of MHC can be stabilized by the presence of Mg²⁺ ions [87,192,193], polymeric substrates like sulfonated polystyrene and polystyrene-divinylbenzene [194], oxalate ions [195], or silica-rich environments [89,196]. Rodriguez-Blanco et al. studied the effect of the Ca²⁺ to Mg²⁺ ratio in aqueous solutions between 65 and 0.17 [87], and found that the magnesium ions inhibited calcite formation [197,198].

Based on PXRD measurements, Rodriguez-Blanco et al. [87] reported that magnesium-containing ACC (ACMC) was stable for 8 h and transformed into MHC in 10 h (Figure 8). Von Euw et al. followed the transformation of ACMC into MHC with solid-state NMR and found that 40% of the carbonate ions of ACMC converted into MHC after 20 h [199]. Ghilardi et al. found that marine species, such as coral reef fishes, mostly produced MHC and ACMC; however, their ratio was species dependent [13].

Korneva et al. [141] and Kimura et al. [200] followed the thermal transformations of synthetic MHC by TG/DTA-MS, PXRD, and FT-IR [200], and the kinetic parameters of the thermal dehydration were determined [200]. Heating of MHC at 443 K for 60 min resulted in the growth of calcite peaks in PXRD [200]. The thermal dehydration-initiated calcite nucleation on the MHC particle surfaces was characterized by a distinguished induction period and self-induced gelation with E_ip = 228.1 ± 8.6 kJ mol⁻¹ and E_a = 210.3 ± 5.1 kJ mol⁻¹, respectively. As it was mentioned in Section 2, Kimura et al. showed two different MHC morphologies, spherical and spherical polygons (Figure 1C), and twisted spindles (Figure 1D) were formed at 288 and 303 K, respectively. They characterized the two kinds of grains with TG-DTA-MS, and two distinct curves were found (Figure 9A and B, respectively). The sample, prepared at 288 K, started to lose water at 432 K and decomposed between 460 and 485 K with calcite formation [74]. In contrast, the sample, prepared at 303 K, decomposed in a wide temperature range (400–560 K), and its PXRD data showed no diffraction peak up to 523 K, consistent with the decomposition product of ACC [74].

Zhang et al. studied how silica-gel media could influence the morphology of MHC [89,196] between 25 and 70 °C (Figure 10) [201]. The MHC nanoparticles, formed upon mixing the starting solutions, were self-aggregated into nanorods and spontaneously self-organized into particles with distinct morphology depending on temperature, which can be explained by, e.g., the dipole–dipole interactions [201].

Chaka et al. performed DFT calculations [187] to determine the effect of Mg²⁺ and H₂O contents during the formation of MHC. They found that the Mg atom had a higher water binding energy than Ca²⁺, which promoted the formation of MHC. However, the thermodynamic stability of MHC decreased due to the presence of a small amount of Mg (11.1 mol%), which resulted in a structural distortion. This destabilizing effect was enhanced if the CO₂ pressure was increased from 1 to 90 bar [187]. Sekkal et al. determined that the most stable surface of the MCH by the DFT method was (001) with 0.99 J/m² surface energy and −5.66 eV attachment energy, which was higher than the most stable (10–14) surface of calcite [28]. The (100) and (010) surfaces of MHC ((100) and (010)) were determined to be 1.54 J/m² with −7.82 eV and 1.21 J/m² with −6.62 eV, respectively [28].

4.3. Transformation of Various ACCs

Synthesis of ACC was recently reviewed by Jiang et al. [24]. Pure ACC is metastable, typically lasting from a few minutes to several hours. However, the presence of magnesium or phosphate ions, or organic macromolecules, can enhance its stability, significantly extending its lifetime. Additionally, factors such as pH play a crucial role by influencing the structure and properties of ACC through the carbonate/bicarbonate equilibrium and the availability of hydroxyl groups [25,29,202,203,204,205]. Furthermore, water has a significant influence on the stability and transformation routes of the ACC [206]. Recently, Maslyk et al. presented a study that describes a way to prepare an intermediate phase, which was proved to be an amorphous calcium carbonate monohydrate (aMHC) that was suggested by the combination of PXRD, FT-IR, and solid-state NMR [206].

In addition to the stabilization effect of the phosphate ion [202,204], some biological entities produce different solvatomorphs of CaCO₃, including ACC, which is the so-called biogenic ACC [109]. Furthermore, these biogenic products were found to be more stable than those without any organic stabilizer. The same stabilization effects of different polymer and organic materials were reported [44,207,208,209,210,211], For example, Yasue et al. used PEG (in ethanol) during the production of ACC [212]. Laboratory experiments unambiguously showed that the transformation of ACC occurs by dissolution–reprecipitation mechanisms, in which local restructuring and water reorganization lead to the nucleation of a more stable phase. Similar pathways were observed in biomineralization, and organisms control the phase transformation by modulating the local chemical environment and by secreting organic matrices that stabilize ACC until crystal nucleation is triggered.

Various synthetic reaction routes yielded ACCs with distinct water contents. Radha et al. determined that the H₂O content of the synthetic and the biogenic ACC were between 1.20 and 1.58 moles and 0.16 to 0.35 moles per formula unit, respectively [12]. They found that the dehydration of ACC had an exothermic nature, which means that irreversible structural reorganization occurred during the dehydration reaction. The measured enthalpy of the crystallization at ambient conditions for the synthetic and biogenic ACC samples was found to be from −17 ± 1 to −24 ± 1 kJ/mol and −14.3 ± 0.97 kJ/mol, respectively. Synthetic ACCs were less ordered and more metastable than the biogenic ACCs [12]. Brečević et al. [16] and Purgstaller et al. [205] determined the solubility of ACC with about 0.5 mol of crystalline water as logK = −6.1987 and −6.20 ± 0.02 [205], respectively. Purgstaller et al. found that the logK value was changed to −4.42 ± 0.01 as a result of amorphous magnesium carbonate (AMC) formation. They investigated the effect of different magnesium ion content and determined the stability and υ₁ mode position in Raman spectra (

Table 4. The summary of the effect of the Mg²⁺ ion content on the position of υ₁ (Raman shift), stability, and solubility of the ACC [205].

Sample Composition	υ1 (Raman Shift)	Stability in Solution (Min)	logK
ACC	1080 ± 1	15	−6.20 ± 0.02
Ca0.91Mg0.09	1082 ± 1	24	−6.13 ± 0.02
Ca0.85Mg0.15	1082 ± 1	24	−6.13 ± 0.02
Ca0.78Mg0.22	1083 ± 1	54	−6.04 ± 0.03
Ca0.69Mg0.31	1084 ± 1	85	−6.00 ± 0.03
Ca0.61Mg0.39	1085 ± 1	64	−5.81 ± 0.07
Ca0.47Mg0.53	1086 ± 1	51	−5.71 ± 0.03
Ca0.20Mg0.80	1088 ± 1	20	−5.43 ± 0.02
AMC	1095 ± 1	9	−4.96 ± 0.01

) [205].

Radha et al. showed that although the difference in the amount of crystalline water in ACC did not influence the XRD positions of the amorphous humps centred at 20 and 45° 2Θ, the intensity ratio of the IR active υ₂ to υ₄ modes changed from around 9.5 to around 6.2 if less crystalline water occurred in the structure [12]. Ihli et al. studied the dehydration and crystallization of ACC in solution and in air with TGA, DSC, SEM, NMR, and IR spectroscopy and developed a model for the ACC to calcite crystallization Figure 11a–e [213].

The overall ACC dehydration process is accompanied by a structural rearrangement, which resulted in a sharpening of the υ₃ IR band, reducing the intensity υ₁ IR peak, and a slight shifting in the υ₂ IR band positions toward higher frequencies (Figure 11f). However, the DSC results did not show any crystallization process up to 290 °C [213]. The activation energies (E_a) of the different dehydration steps are summarized in Figure 11g, which shows a general increase in the E_a values with increasing dehydration. Three well-separated dehydration steps (Figure 11g) are attributed to the presence of three different amounts of H₂O (noted with x) in ACC (Figure 11, where (a): x = 1.4 − 0.98 and E_a = 80 kJ/mol, (b): x = 0.98 − 0.25 and E_a = 145 kJ/mol, and (c): x = 0.25 − 0.08 and E_a = 245 kJ/mol [213]). According to Ihili et al., ACC fully transforms into calcite above 290 °C, a bit lower temperature than that found by Radha et al. [12] or Reeder et al. [213] (330–340 and 290–400 °C, respectively), with E_a 125 J/mol.

ACC, prepared without any additives, is commonly reported to be stable only for minutes to 4 days at ambient temperature and atmosphere, and after this period, it transforms into crystalline CaCO₃ [214,215]. The alternative crystallization pathways of ACC, which involve stable clusters, were reviewed by Gebauer et al. [216]. Rodriguez-Blao et al. studied the ACC transformation into calcite via vaterite at 7.5 °C in air with the Energy Dispersive X-ray diffraction (EDXRD) technique and identified two transformation stages: Stage I, representing the ACC → ACC + vaterite → vaterite crystallization, which took 14–45 min, and Stage II, representing the vaterite → vaterite +calcite → calcite recrystallization, which took about 18 h. The increase in temperature up to 25 °C increased the velocity of transformation to 12 min and 3 h, respectively [180]. It is of interest that Besselink et al. identified a “short-lived” ikaite formation before Stage I, which was facilitated by the low temperature conditions and the presence of chelating citrate ligands [167]. Bots et al. followed the transformation of ACC with small- and wide-angle X-ray scattering (SAXS/WAXS) and identified three stages, including the following: (I) the dehydration of hydrated and disordered ACC, which did not show Bragg peaks and occurred 60s to 90s after the mixing; (II) the transformation of ordered ACC to vaterite, evidenced by X-ray background decrease, and the occurrence of weak vaterite peaks about 4–6 min after mixing and their increase 35 min after the mixing; and (III) the formation of calcite (Figure 12A) [217]. The SEM pictures showed that the 30–40 nm-sized ACC spheres aggregated and transformed into micron-sized vaterite (Figure 12B–D) [217]. An identical morphological transformation was found by Galan et al. [41] and Cai et al. [218].

Studying ikaite formation and its recrystallization, Lázár et al. recognized two morphologically and chemically distinct ACCs [21]. They identified ACC (I), which occurred below 5 °C as 50–100 nm-sized spheres, contained 1.12 mol absorbed and 0.26 mol chemically bonded water, and transformed to ikaite within 10–25 min, and ACC(II), which was micron-sized and preserved the morphology of ikaite, contained 0.42 mol absorbed and 0.35 mol chemically bonded water, and slowly (within hours) transformed to calcite [21]. ACC (I) could be dehydrated to 0.03 mol water upon heating and before its transformation to calcite at 300 °C in air, whereas for ACC(II), calcite transformation was temperature dependent and took more than 24 h at 75 °C and 10–30 min at 150 °C [21]. Interestingly, Lázár et al. [72] reported that ACC (I) occurs during the ikaite-to-calcite transformation, also, and not only as a precursor phase before ikaite formation. In the presence of strontium ions, vaterite did not occur as a transient phase, but the anhydrous ACC immediately transformed into calcite [172]. However, based on Raman studies, Molnár et al. [173] reported the formation of an aragonite-type structure from increased Sr-containing ACC.

Konrad et al. studied ACC stability and the effect of freeze-drying to remove physiosorbed surface water by PXRD, in situ Raman, SEM, and TEM methods [131]. They reported that gastight sealing could stabilize ACC even up to 150 days, and ACC started to transform into anhydrous polymorphs of CaCO₃ after 51 days in a semi-tight container (Figure 13B), but there was some ACC leftover even after 87 days [131]. In an open container, only 5 days of stability could be reached (Figure 13A). Freeze-drying by itself did not have a significant effect on the stability of the ACC [131].

Galan et al. studied the transformation of ACC in solution in the presence of vaterite with and without MgCl₂ additive up to 120 h [41]. The presence of vaterite accelerated the ACC transformation reactions and the calcite formation, whereas the presence of Mg²⁺ inhibited vaterite precipitation at the early stages and promoted the direct transformation of ACC into calcite [41] (Figure 14).

4.4. Transformation of CCHH

The last known member of the CaCO₃ solvatomorph series is the CCHH, which is stabilized by Mg²⁺ ions and prepared from a premixed solution of CaCl₂ and MgCl₂ by adding Na₂CO₃ after the immediately precipitated, 6.5% Mg-containing ACC [20]. The chemical composition of the amorphous phase was found to be Ca₀.₉₈Mg₀.₀₂CO₃·xH₂O, where x = 0.48. It transformed into a new needle-like, ~ 200 nm diameter and 1–5 mm long crystalline material in ~1 h. The transformation was marked by a sharp decrease in Ca²⁺ activity in the solution and was accompanied by an increase in pH, suggesting a mechanism in which Mg²⁺ ions were being excluded from the growing crystals. The initial solution contained Mg²⁺ at high concentration, whereas the final CCHH phase contained only about 1.5 n% Mg [20]. When the concentrations of CaCl₂, MgCl_2, and Na₂CO₃ were selected to be 0.05, 0.06, and 0.05 mol/kg, respectively, the CCHH formed from ACC after 60–110 min and was found to be stable over 5 h (Figure 15E–H) [23]. The Ca²⁺ concentration decreased while the Mg²⁺ concentration was set back to almost the initial concentration, and pH increased when CCHH crystallization started. The Mg²⁺ ions were released during the transformation, and only about 3–4% Mg²⁺ was left in the CCHH (Figure 15H) [23]. When the aragonite peaks appeared on the diffractograms (and, therefore, in the system) after 400 min, the pH, alkalinity, and Ca²⁺ concentration suddenly decreased in the solution (Figure 15E) [23], which agrees well with the findings of Zou et al. [20] and a kinetic control by magnesium. Mg²⁺ ions hindered the complete dehydration of ACC, thus maintaining a partial level of hydration in the final crystalline product [20]. Suyama et al. determined its solubility (logK = −6.80 ± 0.05) and its thermodynamic properties: ΔG_f° = −1238.2 ± 0.3 kJ/mol [23]. Comparing the standard free enthalpies of formation for hydrated CaCO₃ solvatomorphs, the thermodynamic stability order was found to be as follows: ACC < ikaite < CCH < MHC < vaterite < aragonite < calcite, which agrees well with the transformation line ACC → CCHH → Aragonite → Calcite [23]. The thermodynamic modelling and experimental decomposition results of CCHH showed that this solvatomorph is stable only within a narrow window of humidity and temperature, generally between 10 and 30 °C, and at intermediate relative humidities. Outside these conditions, CCHH tends to dehydrate into calcite, or, under high humidity, it may hydrate further into monohydrocalcite. CCHH is sensitive to both drying and heating, decomposing by releasing its bound water and transforming into anhydrous crystalline form. As a result, the CCHH only appeared in specific micro-environments in nature, such as magnesium-rich sediments [20], but its formation was predicted during evaporation in highly alkaline natural lakes [23].

The biomineralization of MHC and CCHH was followed by Schmidt et al. in different natural materials (e.g., coral skeleton) with Myriad Mapping (MM) and FT, and Synchrotron Infrared Nano Spectroscopy (SINS) [22]. A multi-step route was found; the main difference was that aragonite could be formed in many different ways, but calcite formed in two ways only (Figure 15A–D). The same transformation route of CCHH was detected by PXRD [23]. After 370 min of the CCHH synthesis, aragonite appeared and became the dominant phase with calcite after about 430–480 min [22]. The pH, Ca, or Mg concentration was reported to influence these transformations (Figure 15E,H).

5. Conclusions and Future Perspectives

Ikaite, MHC, CCHH, and ACC are hydrated calcium carbonates and represent a complex and dynamic group of solvatomorphs, characterized by distinct morphologies and structural, chemical, and thermodynamic properties. Despite their relatively simple chemical compositions, these phases exhibit unique hydration states, coordination environments, transformation pathways, and environmental stabilities. Recent advances in characterization techniques—such as IR, Raman, and NMR spectroscopies, PXRD, and EXAFS—DFT and MD computational modelling, and in situ monitoring have significantly enhanced our understanding of the formation mechanisms, stability, and mutual interconversions of CaCO₃ solvatomorphs. The discovery of CCHH and the identification of distinct ACC forms, such as ACC I and ACC II, have highlighted the complex influence of water content, hydrogen bonding, and environmental conditions on the structure–property relationships within the hydrated CaCO₃ system. The role of these hydrated phases as precursors in biomineralization, along with their transient occurrence in cold or chemically specific environments, highlights their significance in both biological and geological contexts. Moreover, the transformation of ACC and related precursors into more stable crystalline polymorphs, such as calcite and aragonite, proceeds through kinetically sensitive pathways, governed by factors including temperature, pH, Mg²⁺/Sr²⁺ content, and the presence of organic additives. The metastability of ACC, its intermediate role, and its capacity to incorporate foreign ions make it a particularly compelling target for ongoing research in the fields of biomineralization and climate proxy development.

Advancing our understanding of hydrated calcium carbonate phases will require sustained interdisciplinary research that integrates crystallography, spectroscopy, geochemistry, and computational chemistry—bridging natural and engineered systems alike.