Ab Initio Thermodynamics of Hydrated Calcium Carbonates and

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ab initio Thermodynamics of Hydrated Calcium Carbonates and Calcium Analogues of Mg Carbonates: Implications for Carbonate Crystallization Pathways Anne M. Chaka ACS Earth Space Chem., Just Accepted Manuscript • DOI: 10.1021/ acsearthspacechem.7b00101 • Publication Date (Web): 16 Jan 2018 Downloaded from http://pubs.acs.org on January 20, 2018

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ACS Earth and Space Chemistry

1 1

ab initio Thermodynamics of Hydrated Calcium Carbonates and Calcium

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Analogues of Mg Carbonates: Implications for Carbonate Crystallization Pathways

3 4

Anne M. Chaka

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Pacific Northwest National Laboratory, P.O. Box 999, MS K8-96, Richland, WA 99352

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[email protected]

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509-371-7104 (V)

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509-371-6354 (Fax)

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Abstract

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Formation of calcium carbonate and its hydrates are important for a wide variety of

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geological, biological, and technological concerns. Recent studies have determined that

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formation of anhydrous crystalline calcite, aragonite, and vaterite can involve a complex

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series of nonclassical pathways in which the hydrated polymorphs monohydrocalcite

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(CaCO3•H2O), ikaite (CaCO3•6H2O), and amorphous calcium carbonate (ACC) play key

16

roles and in some instances are stable or metastable endproducts. The stages of

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nucleation and crystallization along these pathways are not well understood, nor is how

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what is learned in an aqueous environment transfers to CO2-rich conditions. In this work

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ab initio thermodynamics based on density-functional theory and experimental chemical

20

potentials for H2O-rich and CO2-rich systems are used to determine the stability of

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calcium carbonate polymorphs as a function of environmental conditions. In water-

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saturated supercritical CO2, formation of ikaite and monohydrocalcite are both highly

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exothermic, yet metastable to calcite, and are therefore likely intermediates upon

ACS Paragon Plus Environment

ACS Earth and Space Chemistry 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

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carbonation of CaO and Ca(OH)2 according to the Ostwald step rule. Hence low energy

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nonclassical crystallization pathways that utilize these intermediates are available for

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calcite formation in CO2-rich environments as well as aqueous systems, particularly in

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water-saturated systems even though water is less than only 1% by mass. Formation free

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energies calculated for Ca analogues of nesquehonite (MgCO3•3H2O), lansfordite

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(MgCO3•5H2O), hydromagnesite (Mg5(CO3)4(OH)2•4H2O), and pokrovskite

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(Mg2CO3(OH)2) are exothermic in both aqueous and water-saturated scCO2 from 273-373

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K, but they are always metastable with respect to the observed Ca minerals. Hence they

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may form prenucleation clusters, transient intermediates, or localized coordination

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arrangements trapped in hydrated ACC, but will never be observed in nature. The

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arrangement of CaCO3•6H2O complexes in ikaite is proposed as the structure of

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prenucleation clusters.

36 37

KEYWORDS: Ab initio thermodynamics, calcium carbonate, DFT, prenucleation,

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crystallization, ikaite, monohydrocalcite

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3 40 41

1.0 Introduction Carbonate formation is a complex process that is important for the global carbon

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cycle, biomineralization, abiotic geochemical systems, paleoclimate indicators,1 and

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industrial processes such as carbon sequestration,2 environmental remediation3, scale

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formation in oilfields and pipelines,4 properties of concrete and cement,5 and

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development of functional materials for microlenses and other applications.6 The early

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stages of nucleation and crystallization of carbonates are not well understood, and have

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recently been the focus of an increasing number of studies on non-classical, multistep

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pathways involving amorphous and hydrated crystalline precursors.7-16 Ca carbonate -

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being widespread in nature, the most common biomineral, and an important industrial

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material - has been the subject of intensive focus as a model system to study the influence

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of factors on the nucleation and growth of carbonates at the molecular level. These

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studies have focused exclusively on aqueous systems. How Ca carbonate polymorph

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formation changes in CO2-rich environments relevant to carbon capture and geochemistry

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has not yet been investigated.

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Ca carbonate and its polymorphs are important throughout the natural

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environment. Ca carbonate can exist in an anhydrous state (calcite, vaterite, or aragonite),

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hydrated (monohydrocalcite and ikaite), or amorphous form. Calcite is the most

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thermodynamically stable polymorph except in cold waters where ikaite becomes the

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most stable. Hence ikaite is important in geochemistry as a paleoclimate and fluid

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composition indicator,1, 17 for increasing the efficiency of the sea-ice carbon pump and

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adsorption of CO2,18 and as a precursor of tufa-like mounds of calcite19-20 and glendonite

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pseudomorphs.1, 21 Monohydrocalcite, although never the most stable Ca carbonate



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polymorph, is likely a widespread metastable intermediate and may warrant

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reconsideration in pathways of geological formations as secondary origin.8, 22

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Monohydrocalcite has been shown to transform to calcite or aragonite. 23-24

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Monohydrocalcite has also been found as a biomineralization endproduct and has been

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proposed as a remediation material for anion pollutants such as arsenate and phosphate.3

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More recently it has garnered considerable interest as an intermediate in the

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transformation of amorphous Ca carbonate (ACC) to calcite8, 22, 25-26, and as a model for

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ACC.27-28 Hence delineating the conditions under which both ikaite and

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monohydrocalcite form and transform is important for interpretation of the geological

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record and understanding carbon cycling and carbonate crystallization.

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Although growth of Ca carbonate can occur by classical nucleation and growth

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under certain circumstances, in the last decade evidence has been increasing that CaCO3

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crystallization can proceed through nonclassical means via an amorphous intermediate.

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The nonclassical transformation process from soluble Ca and carbonate ions to crystalline

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CaCO3 follows a complicated sequence of steps, the majority of which are not yet

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understood. The solution phase ions are hypothesized to condense into prenucleation

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clusters, which then aggregate to form ACC.7 The prenucleation clusters have proven

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challenging to detect. The only indication of prenucleation species has come from

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isolation of ~70 micron clusters by Gebauer and coworkers via analytical ultra

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centrifugation,7 and a shift in carbonate vibration frequency in in situ time-resolved

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Raman spectroscopy observed by Montes-Hernandez and Renard that did not correspond

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to ACC or calcite.16 In either case the structure of the prenucleation species could not be

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determined. Molecular simulation has played a significant role in developing models to


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5 86

understand the early stages of this process. In the molecular dynamics (MD) simulations

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by Tribello and coworkers a barrierless aggregation of primarily neutral bidentate ( )

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Ca2+ CO3-2 ion pairs formed.29 Small amounts of water was trapped kinetically as clusters

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assembled, but was not distributed to each Ca in an approximation of the

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monohydrocalcite structure. In contrast to the results of Tribello and coworkers, MD

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simulations performed by Demichelis and coworkers resulted in prenucleation aggregates

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forming polymeric chains of Ca and carbonate ions termed Dynamically Ordered

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Liquidlike Oxyanion Polymers (DOLLOPs).30 How the DOLLOP aggregates grow,

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reorganize, and reach a critical size, however, has not yet been determined.

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In contrast to the prenucleation species, ACC has been widely observed in nature

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and in the laboratory. ACC provides a low energy pathway for CaCO3 crystallization31 to

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calcite, aragonite or vaterite, and is typically the first phase formed in biomineralization

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processes and crystallization in the laboratory at high supersaturation.14 ACC has been

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detected in more than seven animal phyla and even some plants.32-3334-353637-38 Although

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not widespread in geological settings due to its proclivity to transform to crystalline

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forms, Dupuis and coworkers hypothesized that ACC is involved in the formation of

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essentially all new geological calcareous structures.39 ACC is considered a hydrated form

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of Ca carbonate with water content typically ranging from 0.5 to 1.4 moles of water per

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mole of CaCO3.11 Biogenic ACC is observed to have nominally a stoichiometry of

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CaCO3•H2O (15.25 wgt%).40-41 Dehydration of ACC results in crystallization whether in

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air, upon heating, or even in aqueous solution. 29, 42

107 108

In 2010 Radha and coworkers determined that the order of transformations once ACC had precipitated follows the sequence of thermal stability: hydrous ACC (least



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6 109

stable) → anhydrous ACC → vaterite → aragonite → calcite (most stable),31 though

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small size for amorphous particles may invert this order. More recent results have shown

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the crystallization pathway can be even more complicated with crystalline polymorphs

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such as monohydrocalcite and ikaite playing a role.8, 14, 22, 26 Monohydrocalcite and ikaite

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can precipitate directly from solution or crystallize from ACC.8, 14 Recent work on

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CaCO3 crystallization by Rodriguez-Blanco and coworkers22, Blue et al.8, and others

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have underscored the importance of monohydrocalcite as another potentially important

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end-member phase for ACC transformation. Monohydrocalcite has also been observed as

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an intermediate between ACC and aragonite.3, 22, 24, 43 Ikaite has been reported

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infrequently as a metastable precursor in recent studies on CaCO3 crystallization, likely

119

due to its low temperature stability and highly transient nature at warmer temperatures.

120

Yet it has also been shown to transform to ACC14, 44-45 and subsequently to vaterite. 23, 46-

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48, 15, 31

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Even though ACC is widely observed, its metastability, complexity, and

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variability has made structural characterization challenging despite a wide range of

124

experimental techniques that have been applied.32-33, 45, 49-50 In a review by Cartwright and

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coworkers, Ca carbonate was described as exhibiting polyamorphism – amorphous

126

polymorphism – for ACC as well as polymorphism for crystalline structures calcite,

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aragonite, and vaterite. The range of Ca-O coordination reported by EXAFS is from 5 to

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9.32, 35, 41, 51-53 On average EXAFS indicates approximately seven oxygen atoms in the Ca

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coordination shell with typical Ca – O distances between 2.40 and 2.50 Å, but with

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significant variation with materials from different sources.40, 54


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7 131

Spectra of amorphous materials without sufficient reference standards can be

132

difficult to interpret. For example Michel and coworkers interpreted the 13C NMR of

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carbonate groups in ACC to be exclusively monodentate ( ) based on the similarity of

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the chemical shift with calcite, but did not examine ikaite as an NMR reference.54 Nebel

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and coworkers, however, also ran ikaite as an NMR reference and found that the ACC

136

13

C chemical shift was very close to ikaite, which is exclusively bidentate ( ).49

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Therefore the 13C chemical shift is not sufficient to distinguish between and

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carbonate coordination. Malini and coworkers utilized classical MD simulations to

139

generate models of ACC from four different starting points and obtained synchrotron X-

140

ray scattering pair distribution functions (PDF) essentially identical to the experimental

141

data from Radha et al11, despite significant structural differences. Hence a PDF is not

142

sufficiently sensitive to discriminate between ACC structural candidates.

143

Although considerable progress has been made in understanding nucleation and

144

crystallization, several questions remain regarding the process of going from fully

145

hydrated solution species to anhydrous crystalline structures. The changes in coordination

146

experienced by Ca and the accompanying thermodynamics as it goes from the fully

147

hydrated cation in solution to the monohydrated ACC and monohydrocalcite have not yet

148

been delineated. The number of coordination environments known from the three

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anhydrous and two hydrated polymorphs is limited and does not encompass the range of

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what is observed in ACC let alone partially dehydrated species that may form precursors

151

or transient intermediates. Hence there is a need to generate reasonable structural models

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for Ca-coordination beyond the known crystalline Ca carbonate polymorphs. MD

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simulations of anhydrous42, 55 and hydrated ACC40, 42, 56-58 are powerful in that they



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provide atomistic detail and dynamics, but the large system sizes required for simulation

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of ACC formation and crystallization makes sufficient time and length scales for

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configuration sampling problematic.

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To address this gap in configurations between the fully solvated cation and the

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monohydrate, we look to the type of hydrated structures formed by Mg carbonates. Ca

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and Mg exhibit the same oxide, hydroxide, and carbonate (calcite) structures, but

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completely different hydrated carbonate structures as shown in Figure 1. Not only is the

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stoichiometry complementary (mono- and hexahydrate for CaCO3, and di-, tri-, and

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pentahydrate for MgCO3 plus hydromagnesite and pokrovskite basic carbonates) but the

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local coordination chemistry is completely different, as shown in Table 1.

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Demichelis and coworkers raised the question as to why there are no Ca

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analogues of hydrated Mg carbonate polymorphs observed in nature or the laboratory.59

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The reason Mg does not crystallize in the monohydrocalcite and ikaite structures was

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determined by Chaka and Felmy using ab initio thermodynamics (AIT); the Mg-

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analogues of monohydrocalcite and ikaite were found to be thermodynamically less

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stable than the observed hydrated Mg carbonates due to Mg’s inability to accommodate

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the 8-fold coordination exhibited by the Ca carbonate hydrates. The thermodynamics of

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the Ca analogues, however, have not yet been determined. If these Ca analogues of

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hydrated MgCO3 minerals exhibit intermediate thermodynamic stability, then they may

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be reasonable models for transient metastable structures along the CaCO3

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dehydration/crystallization pathway. Formation of ACC is a nonequilibrium process and

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therefore examining other structures with different degrees of hydration may inform the

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ACC structure, even if they are highly metastable.


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9 177

In this work molecular modeling and ab initio thermodynamics are used to

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determine the structure and thermodynamics of the known hydrated Ca carbonate

179

polymorphs and Ca analogues of hydrated Mg carbonates to determine their potential role

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in forming prenucleation clusters, transient intermediates, or local coordination

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arrangements of ACC in both aqueous and CO2-rich environments. Radha and coworkers

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raised the question whether hydrated ACC formation and crystallization studies are

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relevant to carbon sequestration because the higher temperatures that exist in geological

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reservoirs (~313 – 366 K) and high CO2 activity may preclude ACC’s formation.60 To

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address this issue the ab initio thermodynamics are calculated for the observed hydrated

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Ca carbonate polymorphs and Ca-analogues of Mg carbonates in CO2-rich environments

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relevant for carbon sequestration, including dry and water-saturated supercritical CO2

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(scCO2). In addition, ab initio thermodynamics is used to expand the range of

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thermodynamic data available for species such as calcite and monohydrocalcite beyond

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300 K.



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10 (a) Lime: CaO

(d) Nesquehonite: MgCO3•3H2O

(b) Portlandite: Ca(OH)2

(e) Lansfordite: MgCO3•5H2O

(c) Calcite: CaCO3

(f) Hydromagnesite: (g) Pokrovskite: Mg5(CO3)4(OH)2•4H2O Mg2CO3(OH)2

M1 M2

(i) Ikaite: CaCO3•6H2O

(h) Monohydrocalcite: CaCO3•H2O

191 192

Figure 1. Structures of Ca and Mg oxide, hydroxide, carbonate, and hydrated carbonates.

193

Table 1. Number of Ca and Mg coordinated species in the hydroxide and carbonate polymorphs for

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observed and hypothetical mineral strutuctures.

61

Calcite Aragonite62 Vaterite63-64 Monohydrocalcite65 Nesquehonite66 Lansfordite67 Ikaite68 Hydromagnesite69 Pokrovskite70

M1 M2

Cation Ca Ca Ca Ca Mg Mg Ca Mg, Ca Mg, Ca Ca Mg Mg Ca Mg, Ca Mg, Ca

#H2O 0 0 0 2 2 2 2 6 4 6 4 1 1 0 0

#OH 0 0 0 0 0 0 0 0 0 0 0 1 1 2 4

# − CO3 6 3 4 2 4 2 1 0 2 0 0 4 2 4 2

195


# − CO3 0 3 2 2 0 1 2 0 0 1 1 0 2 0 0

Coordination 6 9 8 8 6 6 7 6 6 8 6 6 8 6 6

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11 196

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2. Methodology Demichelis and coworkers examined how well a dozen DFT functionals

198

reproduced the reaction energies for CaCO3 + nH2O → CaCO3•nH2O for

199

monohydrocalcite (n =1) and ikaite (n = 6) at 298 K for which experimental data has

200

been published, and pointed out the difficulties calculating the thermodynamics of liquid

201

water and the non-ideal behavior of water vapor via ab initio means.59 Functionals that

202

worked well for the anhydrous and low water carbonates (calcite, aragonite, and

203

monohydrocalcite) did not work well for ikaite, and vice versa, primarily due to the

204

differences in types of bonding and difficulties in treating liquids and vapors. Their work,

205

our previous study,71 and that of Costa and coworkers72 underscored the importance of

206

including dispersion into the DFT functional for these systems, as well as the necessity of

207

utilizing corrections to the heat of formation.

208

In this work the difficulties in treating liquids and vapors by ab initio means are

209

circumvented by 1) choosing well-defined reference states for water and CO2 that are

210

suitable for accurate calculations by DFT, namely isolated molecules at 0 K, and 2)

211

utilizing experimental data for the free energies of water and CO2 in the vapor, liquid, or

212

supercritical state at finite temperature and pressure. In addition, we apply corrections to

213

the 0 K heats of formation for the carbonate solids. These corrections are necessary

214

because although errors for pairwise hydrogen bonding and M•••O ligand interaction

215

energies are small – on the order of 5 kJ/mol or less consistent with the known limitations

216

of DFT – these mineral systems are large and the errors have a cumulative effect.71

217

Utilizing a 0 K reference state rather than 298 K enables the separation of error due to

218

cohesive energy versus the vibrational partition function at finite temperature.



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12 219 220 221

2.1 DFT The methodology employed in this work is the same as was developed and tested

222

in prior work on hydrated Mg carbonates in water and in CO2-rich environments.71, 73

223

Hence the overall methodology is summarized briefly here, followed by a more in depth

224

treatment of the requirements for heats of formation correction for the Ca polymorphs.

225

All calculations were performed using periodic all-electron density functional

226

theory as implemented in the DMol3 program. For the real-space cutoff, values ranging

227

from 5.1 to 3.5 Å were tested for a double-zeta plus polarization quality basis set with

228

respect to the ∆fH(0K) for calcite (CaCO3) and portlandite (Ca(OH)2). The difference in

229

∆fH(0K) between the more accurate 5.1 and 3.5 Å is 2.7 kJ/mol for calcite and 4.0 kJ/mol

230

for portlandite, which is negligible. An rcut of 4.3 Å was utilized for computational

231

efficiency, as computational time scales as the cutoff radius r6. Converged k-point

232

sampling using the Monkhorst-Pack74 methodology was used during lattice optimizations

233

performed with the aid of unit cell stress. The generalized gradient approximation to the

234

density functional of Perdew, Burke, and Ernzerhof (PBE)75 was utilized plus an

235

empirical dispersion term in the Hamiltonian as developed by Grimme and first published

236

in 2006.76-77 Herein this method is abbreviated PBE-G06. Table S-1 in the Supporting

237

Information (SI) shows the magnitude of the dispersion energy in these systems, as well

238

as the total PBE-G06 energies and the zero-point vibrational energy (ZPE).To improve

239

the reliability of the vibrational partition function and how the chemical potential ∆µ

240

changes with finite temperature, it is important to ensure the phonon spectrum is

241

converged from a thermodynamic perspective. This involves running increasingly large


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13 242

unit cells and longer phonon wavelengths until the free energy converges at finite

243

temperature. This is shown in Figure 2 for CaO and portlandite for which the JANAF78

244

tables provide thermodynamic data up to 1000 K. Both minerals require a 3x3x3 super

245

cell to converge the phonon spectrum. Calcite required a 2x2x1 super cell, but the

246

conventional unit cells for monohydrocalcite and ikaite were sufficiently large. 0

-20

-40 ∆µ (kJ/mol)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


CaOExpt Expt CaO CaO CaO(1x1x1) (1x1x1) CaO(2x2x2) (2x2x2) CaO

-60

CaO(3x3x3) (3x3x3) CaO Ca(OH)2Expt Expt Ca(OH)2

-80

Ca(OH)2(1x1x1) (1x1x1) Ca(OH)2 Ca(OH)2(2x2x2) (2x2x2) Ca(OH)2

-100

Ca(OH)2(3x3x3) (3x3x3) Ca(OH)2 -120

247

0

200

400

600

800

1000 K

248

Figure 2. Changes in the calculated chemical potential with temperature due to the vibrational partition function for

249

CaO and Ca(OH)2 as a function of unit cell size compared to experimental data.78

250

2.2 Thermodynamic Reference States

251

Since the enthalpy of a substance is not an absolute quantity like entropy,

252

reference states must be chosen for the enthalpy that ensures a consistent basis for

253

comparison. Thermodynamic tables typically use a standard state of 1 bar pressure and

254

298.15 K. Since DFT calculations are done at 0 K, that temperature becomes a more

255

convenient reference state for the ab initio thermodynamics framework. In this work

256

formation energies at 0 K (∆fH) are calculated relative to the oxides, i.e. H2O, CO2, and

257

CaO, consistent with much of the geochemical literature. To facilitate comparison of

258

literature thermodynamic values with the DFT results, literature values given with respect

259

to the elements at 298 K are converted to ∆fH(0K) from the oxides, as shown in Table 2.



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14 260 261

Table 2

262

Thermodynamic values in J/mol for entropy and kJ/mol for all others. HFC is the Heat of

263

Formation Correction for ∆fH°(0K) from the oxides.

CaO1

S°(298)

∆fH°(298)

∆fG°(298)

H-H°(Tr)

∆µ(298-

∆fH°(0K)

∆fH°(0K)

∆fH°(0K)

HFC

Expt

Expt

Expt

Expt

0) Calc.

Expt

Expt

Calc.

Elements

Elements

Elements

Elements

Elements

Oxides

Oxides

Oxides

38.212

-635.089

-603.501

6.749

-631.760

N/A

N/A

N/A

83.387

-986.085

-898.421

14.160

-977.358

-106.677

-92.063

14.614

91.71

-1207.590

-1129.076

14.483

-1202.262

-177.351

-142.585

34.765

Monohydrocalcite

129.7

-1497.950

-1361.218

14.652

-1489.348

-225.516

-195.801

29.715

Ikaite79

310.4

-2971.710

-2541.131

38.596

-2928.997

-470.560

-508.238

-37.679

1

Portlandite 2

Calcite

79

264

1

S°(298), ∆fH°(298), ∆fG°(298), and H-H°(Tr) from JANAF.

265

2

S°(298), ∆fH°(298), and ∆fG°(298) from Konigsberger,79 and H-H°(Tr) from Staveley.80

266

3

S°(298), ∆fH°(298), and ∆fG°(298) from Konigsberger,79 and ∆µ(298-0) from PBE-G06 calculations.

267 268 269 270 271

For a mineral such as calcite for which the ∆fH°(298K) and ∆(H - H°(Tr)) values are known for 0 and 298 K, ∆fH°(0K) from the elements is calculated as follows: ° ° 0) = ∆ ) − − ° ) )− − ° ))0)} − ∆

° ) ° 0) − − ° ) − ° 0) − 1.5 ° ) − ° 0)

272

[Eq. 1]

273

where T = 298.15 K, calcite, Ca, and C are in the solid phase, and O2 is in the gas phase.

274

For monohydrocalcite and ikaite, however, the ∆(H - H°(Tr)) values are not known

275

because the heat capacity measurements between 0 and 298 K have not been done. In

276

these cases the PBE-G06 calculated partition function is used to determine the free energy

277

difference Δ between 298 and 0 K. These values under the ∆µ(298-0) heading and the

278

resultant ∆fH°(0K) values are shown in italics in Table 2. This methodology also enables


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15 279

extrapolation of thermodynamic values beyond measured values to higher temperatures,

280

which are tabulated in Tables 3-3 and S-4 in the SI. Table 2 also shows the PBE-G06

281

calculated ∆fH°(0K) values compared with experimental values, the difference termed

282

HFC for Heat of Formation Correction. Calculated data that include these corrections in

283

this work are termed PBE-G06-HFC.

284

For Ca analogues of the Mg minerals nesquehonite, lansfordite, hydromagnesite,

285

and pokrovskite, however, there are no thermodynamic data available. In these cases the

286

HFC value for the Mg minerals was employed for the Ca-analogues as an approximation.

287

To estimate the magnitude of this assumption, we examined magnesian calcite for which

288

∆fH°(0K) is known for both calcite and magnesite. In this system applying the HFC

289

correction for Ca to Mg in calcite results in an error on the order of 8.7 kJ/mol per Mg.

290

Although this error is not large, it should be noted when considering the AIT results.

291 292 293 294 295

2.2. Chemical potentials The chemical potentials for the molecular and crystalline species used in this work are obtained from , ") = # $%&',() + # +,- + Δ , ")

[Eq 2]

296

where # $%&',() and # +,- are calculated using the PBE-G06 level of theory. For water

297

and CO2, # $%&',() and # +,- are calculated as isolated molecules at 0 K. Total energies

298

and ZPE for crystalline structures are calculated using periodic boundary conditions, as is

299

Δ , ") from the vibrational partition function. For water (. ) and CO2 ( ), the

300

effects of temperature and pressure on the chemical potentials are obtained from

301

experimental data to circumvent the difficulties cited by Demichelis and coworkers.59 For



Page 16 of 51

16 302

the water-rich and ultra-high vacuum (UHV) systems, the thermodynamic data from the

303

JANAF tables are used. The free energies at the reference pressure of 1 bar presented in

304

the JANAF tables are converted to relevant pressures using the ideal gas law:

305

[Eq. 3]

306

Chemical potentials for water and CO2 in CO2-rich environments with varying

307

amounts of water, including the supercritical region, are based on the experiments and

308

thermodynamic model described in Springer et al., and incorporated into software

309

available from OLI.81 The values of the chemical potentials used in this study are

310

tabulated in Appendix A in Chaka and Felmy.73

311 312 313

2.2 Ab initio Thermodynamics The free energies of minerals calculated with respect to the formation reaction

314

from the oxides are a function of independent variables , . , and . For Ca

315

polymorphs, the energy of the crystalline phase is defined as:

316

$%&',() 23 /0&' , ", 1 ) = #0&' + #0&' − 1 , ") − 1 . , ") − 14 , ")

317

[Eq.4]

318

where ni are the coefficients of CaO, H2O, and CO2 in the stoichiometric formula for the

319

23 mineral, #0&' is the ZPE at 0 K plus the vibrational entropy and enthalpy at finite

320

temperature and pressure for the crystalline (xtl) mineral phase, and µi(T, p) is the

321

chemical potential of species i as described above.

322 323

3.0 Results and Discussion


Page 17 of 51 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


17 324

The results are presented in two subsections. Section 3.1 describes the structures

325

of the Ca minerals and the Ca-analogues of Mg minerals presented in Figures 1 and 3 and

326

Table S-2 in the SI. Section 3.2 discusses the AIT results given in Figures 4-8 and Table

327

3.

328 329 330

3.1 Structures The issues to be discussed in this section are threefold, namely (1) how well does

331

the computational method used describe the observed mineral structures, (2) can the

332

structures provide insights into kinetic formation and transformation mechanisms in the

333

crystallization process, and (3) do the Ca-analogues of hydrated Mg carbonates provide

334

unique cation coordination environments that could be useful in interpreting experimental

335

spectra of prenucleation species and ACC.

336

The lattice constants for the observed Ca carbonate polymorphs and the

337

hypothesized Ca-analogues of Mg hydrated carbonates are shown in Table S-2. The

338

bond distances for the first Ca coordination shell are shown in Figure 3. In general, the

339

PBE-G06 level of theory describes the structures of the observed minerals with a high

340

degree of accuracy, with deviations in lattice constants being less than 1%. The

341

exceptions are the deviations for the c vector of calcite (-1.1%) and portlandite (-4.18%),

342

and the a and c vectors for ikaite at 1.34% and 1.22%, respectively, which are still quite

343

small. The portlandite structure is consistent with that calculated by Laugesen using the

344

PW91 functional.5 The structures and hydrogen bonding arrangement found for

345

monohydrocalcite and ikaite agree closely with the DFT results of Demichelis and

346

coworkers59 and Costa et al.82, as well as the neutron diffraction for monohydrocalcite by

347

Swainson65.



Page 18 of 51

18 348

Analysis of the carbonate polymorph structures can provide insight into the

349

likelihood of rapid crystallization, namely how closely does the cation coordination in the

350

solid state resemble solution speciation. Highly hydrated structures such as ikaite,

351

lansfordite, and nesquehonite would be expected to have a higher probability of kinetic

352

formation, even if not the most thermodynamically favored. The more

353

thermodynamically favored species –except at cold temperatures- are the less hydrated

354

species monohydrocalcite, pokrovskite, and hydromagnesite, plus the anhydrous calcite,

355

aragonite, vaterite, and magnesite. These latter species, however, exhibit high carbonate

356

coordination that is not favored in solution and thus require aggregation, dehydration, and

357

extensive rearrangement to form. This is evidenced in Mg carbonation experiments where

358

nesquehonite is rapidly precipitated, followed by slow transformation to hydromagnesite,

359

and an even slower transformation to magnesite.83-85 The thermodynamic stability

360

sequence in these experiments is nesquehonite (least stable) → hydromagnesite →

361

magnesite (most stable), and their order of appearance is consistent with the Ostwald step

362

rule.71, 73, 86 The Ostwald Step Rule – also termed the Ostwald-Lussac Rule of Stages –

363

postulates that the least stable phase crystallizes first, followed by successive

364

transformations into more stable phases. 87 Discussion of the structures as a function of

365

degree of hydration is as follows: ikaite (six), lansfordite (five), nesquehonite (three),

366

monohydrocalcite (one), hydromagnesite (one), and pokrovskite (half).

367

The hexahydrate ikaite structure is extensively hydrated with each Ca coordinated

368

to six water molecules and one carbonate group. The structure is stabilized by an

369

extensive hydrogen bonding network that was described by Demichelis and coworkers.82

370

At the PBE-G06 level of theory, the Ca-Oc distance is 2.443 Å, and the Ca-Ow distance is


Page 19 of 51 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


19 371

on average 2.490 Å. This is in close agreement with synchrotron X-ray diffraction by

372

Lennie who found the mean Ca-O distance in ikaite to be 2.469(3) Å at 243 K.88 This

373

hydration structure around Ca in ikaite is nearly identical to the mean Ca-O distance of

374

2.46(2) Å for 8-fold Ca coordination in an aqueous environment determined by

375

Jalilehvand and coworkers using EXAFS, large-angle X-ray scattering, and ab initio

376

molecular dynamics.89-90 This supports the hypothesis that ikaite crystallizes out of

377

solution without changing its coordination environment or requiring water molecules to

378

be displaced in conditions that are saturated with respect to ikaite.

379

The pentahydrate lansfordite structure consists of two sixfold coordinated Mg

380

complexes hydrogen bonded together. One is a Mg(H2O)62+ complex and the other a

381

Mg(CO3)2 (H2O)42- complex. The carbonate groups are - coordinated as shown in

382

Figure 3e. The Ca analogue of lansfordite exhibited the same 6-fold coordination as the

383

Mg mineral, but with expanded M-O distances shown in Figure 3f. Ca-lansfordite may

384

form as a prenucleation aggregate if there is a sufficient population of hydrated

385

Ca(H2O)62+ and Ca(CO3)2 (H2O)42- complexes in solution. The free Ca2+ ion can

386

accommodate more than six water molecules in its solution coordination sphere, as

387

evidenced by theory and experiment, and has been observed to comprise a significant

388

fraction of the Ca2+ speciation.7 The dicarbonate Ca(CO3)2 (H2O)4+2- complex, however,

389

has to our knowledge not yet been considered as a potential Ca species in solution. As

390

shown in the next section, there are conditions under which the Ca-lansfordite analogue is

391

metastable with respect to CaO and Ca(OH)2 and thus according to the Ostwald Step

392

Rule may form as a transient intermediate. Hence an initial condensation aggregate for



Page 20 of 51

20 393

Ca carbonate may contain a significant fraction of Ca(H2O)6+2+ and Ca( - CO3)2

394

(H2O)4+2- complexes.

395

The trihydrate nesquehonite structure consists of a two-dimensional Mg carbonate

396

ribbon shown in Figure 3a surrounded by waters of hydration. Two water molecules

397

occupy the axial positions as shown in Figure 1d. The M-Ow distances from the cation to

398

the axial waters is lengthened from 2.069/2.095 for Mg to 2.323/2.340 Å for Ca. The

399

third water is not coordinated to Mg, but is hydrogen bonded to the carbonate groups and

400

axial water. This structure is discussed in more detail in Chaka and Felmy.71 The Mg is

401

six-fold coordinated by two and one carbonate groups within the plane of the

402

ribbon, plus two axial water groups. Replacement of Mg with Ca causes a significant

403

distortion of the ribbon due to a shifting of carbonate orientation to yield 7-fold

404

coordination, namely one of the carbonates becomes , shown in Figure 3b. These

405

Ca2+ carbonate ion pairs in the planar CaCO3 ribbon could readily be assembled from

406

ion pairs in solution, as shown by the circled ion pairs in Figure 3b. The axial waters

407

remain in place as the ribbon is assembled, but the planar Ca-coordinated water

408

molecules would be displaced.

409

Monohydrocalcite has a unique structure consisting of three three-fold screw axes

410

and nine formula units in the unit cell. Ca is eightfold coordinated with two water

411

molecules, two bidentate carbonate groups, and two monodentate carbonate groups

412

shown in Figure 3h. Each water molecule is coordinated to two Ca atoms, and hydrogen

413

bonded to two carbonate oxygens. One Ca-Ow bond is slightly longer than the other. The

414

low water content and coordination with four carbonate groups is not an arrangement that

415

would be found in solvated, dilute Ca ions due to an excess of negative charge.


Page 21 of 51 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


21 416

Formation of monohydrocalcite would require significant aggregation of multiple Ca

417

carbonate ion pairs and displacement of almost all waters of hydration, and thus would be

418

kinetically challenging to form despite thermodynamic drivers.

419

In addition to kinetic difficulties in formation, the monohydrocalcite structure also

420

inhibits transformation. It’s structure is unique among the Ca and Mg carbonate hydrates

421

in that the water molecules are not in contact with each other. In all the other hydrated

422

Ca and Mg carbonate polymorphs - namely ikaite, nesquehonite, lansfordite, and

423

hydromagnesite - the water molecules form columns in the structures and hence are

424

readily able to diffuse out upon heating or exposure to dehydrating conditions. These

425

structures are thus capable of undergoing solid state transformations upon facile

426

dehydration. In contrast, water molecules in monohydrocalcite are trapped by

427

surrounding carbonate groups and Ca ions, and cannot readily diffuse out. Hence

428

monohydrocalcite would have to dehydrate and transform into calcite, aragonite, or

429

vaterite via a dissolution/precipitation mechanism. These structural considerations alone

430

thus provide an explanation for why Munemoto and Fukushi observed that

431

monohydrocalcite must first dissolve in order to transform to aragonite.24 This required

432

dissolution/precipitation process to transform monohydrocalcite into anhydrous Ca

433

carbonate can also explain the release of Mg from Mg-containing monohydrocalcite just

434

prior to calcite precipitation.8 It is also consistent with the observations by Jiménez-

435

López and coworkers that monohydrocalcite dissolves prior to precipitation of calcite as

436

carbonate concentration increases.91 The stability of monohydrocalcite is underscored by

437

the observation of Neumann and Epple who found monohydrocalcite to be stable in

438

artificial seawater at room temperature in a sealed vial for 3 months.28



Page 22 of 51

22 439

Hydromagnesite also has a 1:1 ratio of Mg:H2O, but one of five water molecules

440

is dissociated to form two hydroxides. In this structure each Mg is 6-fold coordinated by

441

one water molecule, one hydroxyl group, and four carbonate groups. In the Ca

442

analogue, the coordination sphere is considerably expanded with two of the carbonate

443

groups shifting to become and yielding an 8-fold coordinated structure shown in

444

Figure 3d. The M-Ow distance increases from 2.195 Å to 2.530 Å. The kinetic issues of

445

formation for hydromagnesite are comparable to monohydrocalcite discussed above. In

446

contrast to monohydrocalcite, however, water molecules are organized in columns that

447

provide diffusion pathways for facile dehydration. Hence hydromagnesite need not

448

transform by a dissolution/precipitation mechanism, but can undergo a solid state

449

transformation.

450

Pokrovskite has a 2:1 ratio of Mg:H2O, and all water is dissociated into hydroxyl

451

groups. The pokrovskite structure shown in Figure 1g has corrugated sheets of Mg and

452

hydroxyl groups stabilized by carbonate groups that are coordinated with Mg ions within

453

the same layer and between layers. Each carbonate group has oxygens that are

454

coordinated with one, two, and three Mg ions. There are two distinct type of metal sites,

455

M1 and M2 shown in Figure 3k, with M1 exhibiting two apical hydroxyl groups and four

456

carbonate groups, and M2 having two axial carbonate groups and four hydroxyl

457

groups. The structure is described in detail in Chaka.92 In the Ca analogue of

458

pokrovskite, the rigid coordination structure and metal placement does not allow for the

459

carbonate groups to shift from to ; hence they remain sixfold coordinated.

460

Formation of this complicated structure would be expected to require a high


Page 23 of 51 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


23 461

concentration of cation and carbonate and low water, as well as a basic pH, and hence be

462

kinetically more difficult to form.

463

To summarize, the highly hydrated ikaite, Ca-lansfordite, and Ca-nesquehonite

464

structures have pathways of formation from solution that are expected to proceed with

465

low barriers due to similarities of ionic arrangements with the solvated cation. In contrast,

466

poorly hydrated species such as monohydrocalcite, Ca-hydromagnesite, and Ca-

467

pokrovskite require aggregation, extensive ionic rearrangement, and dehydration to

468

transform solution phase structures into the solid state, processes which can be expected

469

to have much higher kinetic barriers than simple ion assembly without rearrangement and

470

minimal or no dehydration. The thermodynamic attributes of these crystal structures are

471

discussed in the next section.

472



Page 24 of 51

24 MCO3•3H2 O (b). Ca (a). Mg*

M5 (CO3 )4 (OH)2 •4H2O (c). Mg* (d). Ca

3.360 2.610

2.057 2.163

2.366

2.019

2.080

2.318

2.141

2.072

2.679

2.034

2.586 2.439 2.494 2.451

2.092

2.933

2.394

2.530

2.067

2.195

2.310

2.457

MCO3 •5H2 O (f). Ca

(e). Mg*

2.384 2.028!

2.084!

2.087! 2.087! 2.090! 2.089!

2.111!

2.111!

2.089!

2.090!

2.084! 2.028!

2.300

2.434

2.341

2.434

2.374

2.341 2.300

2.339 2.374

2.339

2.384

(c)!

MCO3 •6H2O

MCO3 •H2 O

2.360

2.050 2.079

2.204

(i). Mg

(h). Ca*

(g). Mg

2.106

2.421

2.455

2.400

2.223

2.642

2.468

2.136

2.437

(j). Ca*

2.503

2.444

2.223

2.532

2.468

2.384

2.139

2.516

2.629 2.403

2.139

2.013

2.442

2.189

2.526

2.189

2.464

M2 (CO3 )(OH)2 (l). Ca

(k). Mg* 1.998

2.085 2.067

2.105 2.172

2.104

2.047 2.032

2.313

2.277

2.097

2.224 2.476 2.590

2.002

2.291

2.328

2.245 2.263

2.204

M1

2.209

2.271 2.259

M2

2.439

M1

M2

473 474

Figure 3. Ligand arrangements for Ca (blue) and Mg (green) in the hydrated carbonates calculated by

475

DFT-G06. Observed structures are indicated by an asterisk (*). Bond distances in Å. Axial waters in


Page 25 of 51 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


25 476

the nesquehonite structures (a) and (b) not shown for clarity. Circles in (b) show arrangement of

477

bidentate ion pairs.

478 479 480

3.2 ab initio Thermodynamics In this section, the thermodynamic stability of the hydrated Ca carbonates and Ca-

481

analogues of Mg carbonate polymorphs is considered as a function of environmental

482

conditions across the range of . and from pure water to dry scCO2 to ultrahigh

483

vacuum (UHV). The specific states examined are: 1) ambient pressure where pCO2 is set

484

to the atmospheric concentration of 400 ppm, and pH2O is set to 32 mbar, the pressure

485

where the vapor chemical potential is equal to that of liquid water at 298 K; 2) high pCO2

486

conditions where pCO2 equals 1, 60, and 90 bar and water is either at saturation or trace

487

amounts (mole fraction χ is 10-9); and 3) UHV where pCO2 = pH2O = 10-4 mbar. The

488

high pCO2 range considered is limited to 275-375 K due to the availability of the

489

chemical potential data. It should be noted that scCO2 calculations up to 210 bar were

490

performed, but the results did not differ significantly from 60 and 90 bar and therefore are

491

not presented. Free energy phase diagrams are presented in Figure 4-8 phase boundary

492

temperatures in Table 3.

493 494 495 496 497 498



Page 26 of 51

26 499

Table 3. Phase Boundarya and Thermal Decomposition Temperatures (K) calculated by PBE-G06-

500

HFC.

Portlandite Ca(OH)2 → CaO+H2O

" = 1 bar ". = 32 :;? :bar

640.7

640.7

393.5

>1000 NA

814.0 NA

618.2 NA

267.7 653.7 667.3

267.7 551.8 493.0

174.6 390.0 386.8

Calcite

CaCO3 → CaO + CO2 CaCO3 ↔ Ca(OH)2

Monohydrocalcite

CaCO3•H2O → CaCO3 + H2O CaCO3•H2O → CaO + CO2 + H2O CaCO3•H2O ↔ Ca(OH)2

Ikaite

501

CaCO3•6H2O → CaCO3 + 6H2O 287.7 287.7 183.2 CaCO3•6H2O → CaO + CO2 + 6H2O 392.2 373.2 245.2 CaCO3•6H2O ↔ Ca(OH)2 354.4 334.7 221.5 CaCO3•6H2O ↔ CaCO3•H2O 292.1 292.1 185.1 Ca-Nesquehonite Analogue CaCO33H2O → CaO + CO2 + 3H2O 436.5 400.1 268.1 CaCO33H2O → CaCO3 + 3H2O 238.7 238.7 150.9 CaCO33H2O ↔ Ca(OH)2 373.7 333.5 227.9 CaCO33H2O ↔ CaCO3•H2O 222.3 222.5 138.0 CaCO33H2O ↔ CaCO3•6H2O 336.1 336.0 215.0 Ca-Lansfordite Analogue CaCO35H2O → CaO + CO2 + 5H2O 415.2 360.0 238.2 CaCO3H2O → CaCO3 + 5H2O 295.2 255.0 162.4 CaCO35H2O ↔ Ca(OH)2 374.6 312.2 208.5 CaCO35H2O ↔ CaCO3•H2O 302.6 251.6 159.2 CaCO35H2O ↔ CaCO3•6H2O 248.4 (463) (289) Ca-Hydromagnesite Analogue Ca5(CO3)4(OH)24H2O ↔ CaO 492.5 422.2 293.5 Ca5(CO3)4(OH)24H2O ↔ Ca(OH)2 315.9 230.1 178.5 Ca5(CO3)4(OH)24H2O ↔ CaCO3 NA NA NA Ca5(CO3)4(OH)24H2O ↔ 359.9 355.0 228.5 CaCO3•6H2O Ca-Pokrovskite Analogue Ca2CO3(OH)2→ CaO + CO2 + H2O 466.5 395.2 280.5 a Forward arrows → indicate a reaction that becomes thermodynamically favored at the temperature indicated.

502

Bidirectional arrows ↔ indicate a phase boundary between two minerals. Note, experimental temperatures for calcite

503

have been extrapolated above 298 K using ab initio thermodynamics.

504 505 506 507


Page 27 of 51

27 508 509

3.2.1 Ambient The ambient results are presented in two stages. Initially the impact of the heat of

510

formation correction for DFT on the thermodynamics of the Ca mineral species is

511

discussed, followed by the AIT of the Ca-analogues of the Mg minerals.

512

The AIT results for Ca carbonates, hydrates, and hydroxide under ambient

513

conditions (pH2O = 32 mbar, pCO2 = 400 ppm) are shown in Figure 4 for both the heat of

514

formation corrected (PBE-G06-HFC) and the uncorrected DFT (PBE-G06) methods.

515

The uncorrected DFT results yield qualitative results consistent with experimental

516

observations – that calcite is the most thermodynamically stable across most of the

517

temperature range until decomposition at high temperature, and ikaite is the most

518

thermodynamically stable at lower temperature. Portlandite is always metastable with

519

respect to calcite, and monohydrocalcite is always metastable with respect to the most

520

stable polymorph, which is ikaite at low temperatures and calcite at higher temperatures. 10 -10 -30 µCa (kJ/mol)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


-50 -70 -90

-110 -130 -150

521 522

.

250

275

300

325

350

375

400 K

Figure 4. Comparison of PBE-G06 calculated free energies of Ca carbonate

523

minerals with (solid lines) and without (dashed lines) heat of formation

524

corrections (HFC).

525

The application of HFC brings the DFT results into quantitative agreement with



Page 28 of 51

28 526

experimental values. The HFC for calcite, monohydrocalcite, and portlandite all serve to

527

lower the energy by 34.77, 29.71, and 14.61 kJ/mol, respectively. Ikaite has the largest

528

HFC of all the polymorphs at -37.41 kJ/mol, which is not surprising as it has the largest

529

number of atoms in the stoichiometric formula. It should be noted that the HFC for

530

ikaite has the opposite sign of the other polymorphs due to the slight overbinding of

531

DFT-GGA for hydrogen bonds. Applying the HFC and thus increasing the stability of

532

calcite and decreasing the stability of ikaite results in shifting the temperature of their

533

phase boundary in liquid water under atmospheric pressure from 352.0 K to 287.7 K, a

534

much more realistic value. This lower temperature phase boundary is consistent with

535

ikaite’s utility as a paleothermometer for cold water conditions93 and the determination

536

that ikaite is the only Ca carbonate that precipitates in sea ice.44 The actual maximum

537

temperature for ikaite’s stability has been shown to depend upon conditions. In seawater,

538

the generally accepted maximum temperature for ikaite’s formation and persistence is

539

280 K.93 Ikaite was found to be stable up to 291 K in the presence of saccharose when

540

MacKenzie94 injected CO2 gas into a solution of CaO, following a method reported by

541

Pelouze in 1865.95 The maximum temperature limit for ikaite was raised to 297 K by

542

Brooks and coworkers when sodium hexametaphosphate was added during the mixing of

543

Na2CO3 and CaCl2.96 In the presence of triphosphate, Clarkson found ikaite would form

544

up to 298 K.14 The value of 287.7 K determined by AIT for the ikaite/calcite phase

545

boundary serves as a reference value that depends solely on the intrinsic free energy of

546

the solids, . and in the absence of solution kinetic effects and any added

547

compounds. For monohydrocalcite, including the HFC shifts the temperature at which it

548

becomes metastable with respect to calcite from 292.7 K to 267.7 K with HFC. Below


Page 29 of 51

29 549

this temperature, monohydrocalcite is more thermodynamically stable than calcite,

550

though not as favored as ikaite. pH2O = 32 mbar; pCO2 = 400 ppm 0

CaO

-20 -40 µCa (kJ/mol)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


-60 -80

-100 -120 -140

551

275

295

315

335

355

375 K

552

Figure 5. AIT free energy phase diagram for Ca minerals (solid lines) and Ca-analogues of Mg minerals (dashed lines)

553

calculated using PBE-G06-HFC.

554

According to the Ostwald Step Rule, these sequences of thermodynamic stability

555

can provide insight into species that may be kinetically formed during carbonation

556

processes, along with the structural complexity analysis of the previous section.

557

Carbonation of portlandite slurries has been shown to occur readily upon brief exposure

558

to high pCO2.16 Portlandite is much more soluble than brucite, which would facilitate the

559

carbonation reaction.97 The AIT phase diagrams enable predictions of which polymorphs

560

are more stable than portlandite and thus may be observed upon portlandite’s dissolution

561

and subsequent carbonation. In an aqueous solution under atmospheric pressure with 400

562

ppm pCO2, the sequence of thermodynamic stability changes with temperature. Below

563

287.7 K the sequence of thermodynamic stability from lowest to highest (most stable) is

564

CaO → portlandite → monohydrocalcite → calcite → ikaite. Above 287.7 K ikaite



Page 30 of 51

30 565

and calcite switch places, and above 292.1 K, ikaite switches with monohydrocalcite, to

566

result in the following sequence that holds up to 344.7 K: CaO → portlandite → ikaite

567

→ monohydrocalcite → calcite. Above 344.7 K, ikaite is less stable than portlandite,

568

and hence the thermal stability sequence is CaO → ikaite → portlandite →

569

monohydrocalcite → calcite. The thermodynamic sequence of metastability stages and

570

hence which phases may be observed during the crystallization process, change with

571

temperature under ambient conditions in water.

572

The results of the AIT calculations of the Ca-analogues of the Mg carbonate

573

minerals nesquehonite, hydromagnesite, lansfordite, and pokrovskite are shown in Figure

574

5. These results show clearly that formation free energies of these compounds from the

575

oxides are exothermic. These Ca-analogues, however, are highly metastable with respect

576

to the known minerals across the entire range of geologically relevant conditions,

577

consistent with the fact that they have never been observed in nature. In addition, they are

578

all predicted to decompose to CaO, H2O and CO2 between 360 – 422 K, much lower than

579

the observed polymorphs except for ikaite. The highly hydrated Ca-lansfordite and Ca-

580

nesquehonite are the lowest energy analogues primary due to the stabilization of

581

hydrogen bonding and nesquehonite’s ability to accommodate 7-fold coordination. The

582

least stable are the basic carbonate analogues Ca-hydromagnesite and Ca-pokrovskite,

583

which have much less structural flexibility to accommodate the larger size cation.

584

Although Ca-hydromagnesite exhibits 8-fold coordination, it has much less stabilization

585

from hydrogen bonding, and Ca-pokrovskite has none.

586 587

According to the Ostwald step rule, the sequence of thermodynamic stability provides an indication as to which phases may be observed during a carbonation of CaO


Page 31 of 51 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


31 588

or Ca(OH)2 experiment and the series of dissolution/precipitation/transformation

589

reactions. This can include the Ca-analogues of Mg polymorphs. At 278 K the stability

590

sequence from least to most stable is: CaO → Ca-pokrovskite → Ca-hydromagnesite →

591

portlandite → Ca-lansfordite → Ca-nesquehonite → monohydrocalcite → calcite →

592

ikaite. Ikaite will certainly form and persist as it is both the most thermodynamically

593

stable as well as the most kinetically accessible from associated Ca2+ and CO32- ions in

594

solution. The Ca analogues of lansfordite and nesquehonite may also be kinetically

595

accessible, and condense as components of the prenucleation clusters that contribute to

596

ACC at lower temperatures.

597

The basic carbonates Ca-hydromagnesite and Ca-pokrovskite are the most

598

metastable of the hydrates. Given their high energies, complicated structures, and high

599

degree of carbonate coordination, Ca-hydromagnesite and Ca-pokrovskite are not likely

600

to be formed in the highly hydrated initial stages of prenucleation from solution. In the

601

later stages of aggregation and dehydration, however, they may form and become trapped

602

in local regions of ACC. There is indeed evidence of hydroxyl groups in ACC. NMR has

603

indicated that in addition to water bound to Ca, a small fraction (~7±3%) of hydrogen

604

was present as rigid OH groups.54 Some of this may be due to regions in ACC with a

605

structure similar to Ca-hydromagnesite or Ca-pokrovskite. Ca-hydromagnesite has a

606

Ca:H2O ratio of 1:1, and pokrovskite has 1:0.5, both of which are in the range of the

607

observed hydration stoichiometry of ACC. It should be noted that Ca-pokrovskite’s

608

stoichiometry can also be written as CaCO3•Ca(OH)2. The atomic coordinates of these

609

structures are given in Table S-5,7 in the SI to assist with spectroscopic interpretation.

610

At a higher temperature such as 353 K in an aqueous environment the Ca-



Page 32 of 51

32 611

analogues are more destabilized relative to the observed Ca minerals with a stability

612

order of: Ca-lansfordite → CaO → Ca-pokrovskite → ikaite → Ca-hydromagnesite →

613

Ca-nesquehonite → portlandite → monohydrocalcite → calcite. Carbonation of CaO

614

might result in a very transient formation of ikaite, but at this high a temperature ikaite is

615

likely to dehydrate quickly to form monohydrocalcite or calcite via an amorphous

616

intermediate. Carbonation of portlandite would not go through an ikaite intermediate, as

617

it is higher in energy, but would likely go directly to monohydrocalcite, ACC, or

618

anhydrous CaCO3.

619

4.2.3. AIT in scCO2

620

The thermodynamics of the Ca carbonates and hydrates have not yet been

621

measured in CO2-rich conditions, which is most relevant for carbon sequestration.

622

Although scCO2 injected below ground or in above ground processing plants is dry, it is a

623

superfluid that can migrate extensively through rock and extract considerable amounts of

624

water. Even though the amount of water in saturated scCO2 is small (less than 1% by

625

weight) it exhibits the full reactivity of liquid water. Hence water in the CO2 liquid and

626

supercritical phases must be considered as well.98 The AIT results for scCO2 are

627

presented in two stages. First, the thermodynamics of observed Ca minerals are

628

discussed in both saturated and trace water conditions. In the second stage, the

629

thermodynamics of the Ca-analogues of Mg minerals are presented only for the saturated-

630

water case, as they become very unstable in the absence of water.

631

The results for the AIT of the observed Ca carbonate polymorphs in high pCO2 are

632

shown in Figure 6 for water-saturated and for trace (χ = 10-9 mole fraction) water

633

conditions. As expected, calcite is the lowest energy polymorph across the 275-375 K


Page 33 of 51

33 634

temperature range under all high pCO2 conditions. This contrasts with the ambient

635

conditions with 400 ppm pCO2, where ikaite becomes lower in free energy than calcite at

636

287.7 K. 1 bar pCO2 100

CaCO3•6H2O (tr)

µCa (kJ/mol)

-50

50

0

0

-50

-50

-100

-100

Ca(OH)2 (s)

-100

-150

-150 275

637

325

375

90 bar pCO2

100

50

CaO

0

60 bar pCO2

100

50

-150 275

325

375

275

325

375 K

638

Figure 6. Free energies of Ca carbonate minerals under high-pCO2 with water saturation (s, solid lines) and trace (tr,

639

dashed lines) water. 1 bar pCO2 10

µCa (kJ/mol)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


CaO

Water Saturated 10

-10

-10

-30

-30

-50

-50 Ca(OH)2

-70

-70

-90

-90

-110

-110

-130

-130

-150

640

90 bar pCO2

-150 275

325

375

275

325

375 K

641

Figure 7. Free energies of Ca carbonate minerals (solid lines) and Ca-analogues of Mg minerals (dashed lines) in

642

water-saturated pCO2 at 1 and 90 bar.

643



Page 34 of 51

34 644

For trace water conditions - effectively dry CO2 indicated by the dashed lines in

645

Figure 6 - calcite is by far the most stable, becoming slightly more stable (-136 to -143

646

kJ/mol) as pCO2 increases from 1 to 90 bar. The order of stability is ikaite → CaO →

647

portlandite → monohydrocalcite → calcite. Free energies of formation of calcite,

648

monohydrocalcite, and portlandite are well-separated by 50-75 kJ/mol at 300 K, yet still

649

sufficiently stable to be lower in energy (more stable) than CaO with the exception of

650

portlandite at higher temperature at 1 bar pCO2. Ikaite, however, is extremely unstable

651

under dry conditions with an endothermic free energy of formation of 138 kJ/mol,

652

making it less stable than CaO across the entire temperature range. In Figure 6 the free

653

energy for ikaite under trace water conditions is so high in energy it is off the scale and

654

not shown.

655

In a water-saturated high pCO2 environment, the relative stability of the Ca

656

polymorphs changes dramatically from the trace-water series to CaO → portlandite →

657

ikaite → monohydrocalcite → calcite. In equilibrium with water, the free energy of

658

ikaite is lowered (made more stable) to a point where ikaite and monohydrocalcite are

659

essentially isoenergetic at -131 and -129 kJ/mol at 1 bar respectively, and slightly lower

660

in energy at 90 bar at -138 and -137 kJ/mol. Hence ikaite and monohydrocalcite are both

661

likely precursors to calcite at low temperatures. Ikaite would likely be formed first as

662

structural simplicity favors fast formation. Upon standing, however, thermodynamics will

663

drive ikaite to dehydrate and convert to monohydrocalcite or ACC, followed by

664

transformation to calcite. Monohydrocalcite’s transformation will likely occur via a

665

dissolution/precipitation process due to the lack of diffusion channels for water as

666

previously discussed. As the temperature increases, ikaite destabilizes more rapidly than


Page 35 of 51 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


35 667

monohydrocalcite and calcite, as would be expected due to entropy favoring free over

668

bound water.

669

Experimentally, carbonation of portlandite occurs rapidly.16 Montes-Hernandez

670

and Renault used in situ Raman spectroscopy to observe the transformations that

671

occurred upon exposure of a portlandite slurry to a pCO2 of 50 bar at 298 K. Within the

672

first five minutes Ca carbonate clusters and/or ACC was observed as well as an

673

unidentified species referred to as a complex carbonate, prior to growth of vibrational

674

modes consistent with calcite. The complex carbonate vibrational mode peaked at 1073

675

cm-1. The ab initio thermodynamics predicts that ikaite may be a transient intermediate

676

under these conditions. The ikaite Raman spectrum obtained by Coleyshaw and

677

coworkers is consistent with the complex carbonate intermediate, as the carbonate

678

vibrational mode was observed at 1072 cm-1.23

679 680 681

3.2.4. AIT of Ca analogues of Mg carbonate polymorphs As shown by the dashed lines in Figure 7, under high pCO2 water-saturated

682

conditions between 275 and 375 K the Ca analogues of nesquehonite and lansfordite are

683

more stable than portlandite but less than ikaite, monohydrocalcite, and ikaite. There are

684

no significant differences as pCO2 increases from 1 to 90 bar. The exothermic formation

685

energy of lansfordite indicates that the fully hydrated Ca2+ ion as well as the Ca( -

686

CO3)2 (H2O)4+2- complex may form in water-saturated scCO2 as portlandite or CaO

687

dissolves. In carbonation experiments of the Mg minerals brucite and forsterite, water in

688

scCO2 has been shown by Loring and coworkers to form films at surfaces and surround

689

ions as they are solvated rather than be uniformly dispersed in the supercritical fluid as



Page 36 of 51

36 690

isolated molecules.99 Hence it is conceivable that hydrated complexes such as ikaite’s

691

Ca( - CO3) (H2O)6+, and Ca-lansfordite’s Ca( - CO3)2 (H2O)4+2- and Ca(H2O)6+2+

692

species will form and contribute to prenucleation clusters, followed by dehydration and

693

formation of monohydrocalcite, ACC, and calcite.

694 695 696

3.3. UHV The thermal stabilities of the Ca carbonate polymorphs UHV conditions where

697

pH2O and pCO2 are equal to 10-4 mbar is shown in Figure 8 and Table 3. This state also

698

corresponds to an inert atmosphere at higher pressures such as may be employed in

699

thermal decomposition experiments. Under UHV conditions, the relative stabilities of

700

the Ca polymorphs and Ca-analogues of Mg polymorphs follow the same general trends

701

as under ambient conditions, albeit at much lower temperatures. Ikaite is the most stable

702

below 183.2 K, and calcite above that until it decomposes at 618.2 K. Monohydrocalcite

703

will decompose to calcite and water at 174.4 K. Portlandite is remarkably stable up to

704

393.5 K upon which it decomposes to CaO and water. The Ca-analogues of Mg

705

polymorphs are highly metastable, decomposing at very low temperatures under UHV

706

conditions.

707


Page 37 of 51

37

0

pH2O = pCO2 = 10-4 mbar CaO

-50 µCa (kJ/mol)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


-100

-150

-200 100

200

300

400K

708 709

Figure 8. Free energies of Ca carbonate minerals (solid lines) and Ca-analogues of Mg minerals (dashed lines) under

710

UHV conditions.

711 712 713

3.4. Implications for Prenucleation Although it is metastable relative to calcite, the structure and thermodynamics of

714

ikaite have implications for the prenucleation clusters observed by Gebauer and

715

coworkers7 as well as Montes-Hernandez and Renaud.16 Ca carbonate in solution has

716

been found to resemble the ikaite structure with each Ca coordinated to six water

717

molecules and one bidentate carbonate group as shown in Figure 3j and discussed Section

718

3.1. Ikaite is 52% water by weight, with the bidentate Ca2+ CO32- ion pair dipoles

719

arranged linearly as shown in Figure 9. Adjacent rows of dipoles are oriented in opposing

720

directions to ensure a nonpolar structure. The dipoles are separated by 4.118 Å due to the

721

coordinated waters in between, and provide an electrostatic driving force for alignment of

722

the complexes in solution. We postulate that this driving force – and significant

723

exothermic heat of formation for ikaite in water-saturated conditions with high carbonate

724

concentration – is sufficient for the ion pairs in solution to self-organize into the ikaite



Page 38 of 51

38 725

arrangement and form clusters. At lower temperatures these clusters can grow and

726

precipitate out as ikaite. At higher temperatures they will dehydrate and lead to

727

formation of ACC, monohydrocalcite, and ultimately anhydrous Ca carbonate. This

728

mechanism can provide a low energy pathway for formation of ACC and CaCO3

729

polymorph crystallization in aqueous solutions with added carbonate and in water-

730

saturated liquid and scCO2.

731

a c b

4.118 Å

b c

732 733

a

Figure 9. Arrangement of Ca2+ and CO32- ion pairs in ikaite. Water molecules not shown for clarity.

734 735

4.0 Conclusions

736

Understanding how Ca carbonate coordination and energies change successively

737

from the fully hydrated structures to the minimally and fully dehydrated stages provides

738

insight into the underlying mechanisms of carbonate crystallization. AIT calculations

739

show that in water-saturated scCO2 nonclassical low energy pathways are available for

740

calcite formation. The formation energies of the hydrated Ca carbonate polymorphs

741

monohydrocalcite and ikaite are exothermic even at temperatures up to 373 K. Even

742

though they are metastable with respect to calcite, they can form as transient (ikaite) or


Page 39 of 51 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


39 743

persistant (monohydrocalcite) intermediates upon carbonation of Ca(OH)2 or CaO. Hence

744

formation of CaCO3•6H2O complexes are kinetically as well as thermodynamically

745

feasible in water-saturated scCO2. We postulate that these CaCO3•6H2O complexes

746

aggregated in the ikaite arrangement are one of the major constituents of the

747

prenucleation clusters observed by Gebauer et al.15 and Montes-Hernandez and

748

coworkers16 in aqueous systems, and can occur in scCO2 as well. Ikaite is likely to form

749

rapidly as its coordination structure is exhibited in solution and its crystallization is

750

essentially an alignment and ordering of solution complexes. This same structure enables

751

facile water diffusion and dehydration in all directions. Upon dehydration ikaite has been

752

shown to result in formation of ACC, monohydrocalcite, vaterite, and calcite under other

753

conditions21, 45-46, 48, 100, and is likely to do so in CO2-rich environments as well.

754

In aqueous systems with 400 ppm pCO2, ikaite is a likely intermediate for

755

carbonation of Ca(OH)2 at low temperatures. As carbonate concentration is increased in

756

aqueous solution towards a chemical potential equivalent to pCO2 of 1 bar as in

757

concentrated carbonate solutions, the ikaite stability is dramatically increased and

758

becomes a more likely intermediate across the temperature range from 273 to 373 K.

759

Ikaite or prenucleation clusters with the ikaite structure become more susceptible to rapid

760

dehydration the higher the temperature. This dehydration can lead directly to formation

761

of ACC, or crystalline monohydrocalcite, vaterite, or calcite.

762

The monohydrocalcite structure is complicated and more difficult to form than

763

ikaite despite being more thermodynamically favored. There is no direct relationship

764

between a solution Ca carbonate ion complex and monohydrocalcite coordination, thus it

765

is unlikely to comprise a significant fraction of the prenucleation clusters in solution.



Page 40 of 51

40 766

Once a critical mass of aggregated carbonate and Ca ions has developed with most of the

767

waters of hydration removed, monohydrocalcite can form and trap water in its structure

768

since there is no clear diffusion pathway out of the crystal. Monohydrocalcite is therefore

769

likely to persist much longer until transforming via a dissolution/precipitation mechanism

770

in both aqueous and CO2-rich systems. The free energy of monohydrocalcite is likely to

771

be a lower bound for ACC with a comparable water content, as bond distances and angles

772

are less than ideal in amorphous systems. As ACC dehydrates, calcite becomes a lower

773

bound for its free energy.

774

Nature exhibits a gap with respect to intermediate degrees of hydration for Ca

775

carbonate minerals, but not for Mg. Yet understanding intermediate stages of hydration

776

and sampling a greater variety of coordination environments can provide insight into the

777

nucleation and crystallization process. Formation energies of the Ca analogues of Mg

778

carbonate minerals are exothermic in both aqueous and water-saturated CO2-rich

779

systems, though in general metastable with respect to the observed Ca carbonate

780

polymorphs. In the carbonation process it is reasonable to postulate that the polymorphs

781

that most closely resemble structures in solution, namely Ca-lansfordite and Ca-

782

nesquehonite, may also play a role in the initial formation of prenucleation clusters or as

783

rapidly transforming intermediates. This progression would be consistent with Ostwald’s

784

Step Rule. Ca-hydromagnesite and Ca-pokrovskite have complicated structures with no

785

direct relationship to solution complexes, and hence are not expected to constitute an

786

early phase or highly hydrated prenucleation cluster. Given that the cation:water ratio is

787

less than one for both Ca-hydromagnesite and Ca-pokrovskite, their local Ca coordination

788

structure may appear in low-water ACC and account for the rigid OH groups observed by


Page 41 of 51 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


41 789

NMR. Atomic coordinates for all the Ca-analogues of hydrated Mg carbonates are

790

provided in the SI so that these structures can be factored into a fit for a PDF or EXAFS

791

spectral interpretation.

792

Although solution thermodynamics, kinetics, surface reactivity and enthalpy, and

793

deposition/dissolution mechanisms have a great influence on the precipitation of

794

carbonates, understanding the bulk thermodynamics and structure-property relationships

795

provides a framework for understanding processes that are too fast, too slow, or two

796

difficult to measure in the laboratory or on geological timescales.

797 798

Acknowledgments:

799

This work was supported by the U.S. Department of Energy, Office of Basic Energy

800

Sciences, Division of Chemical Sciences, Geosciences & Biosciences, the Geosciences

801

Research Program. This research was performed using the computing facilities at

802

EMSL, a national scientific user facility sponsored by the U.S. Department of

803

Energy's Office of Biological and Environmental Research and located at Pacific

804

Northwest National Laboratory (PNNL), and PNNL Institutional Computing. PNNL

805

is a multiprogram national laboratory operated for DOE by Battelle.

806 807

Supporting Information contains tables of the 1) Total, ZPE, and dispersion energies

808

for all species in this work (Table S-1); 2) Optimized lattice constants of Ca carbonate

809

minerals and Ca-analogues of Mg carbonate minerals (Table S-2)

810

3) Free energies of portlandite, calcite, monohydrocalcite, and ikaite up to 1000 K at 1

811

bar and UHV conditions (Table S-3), as well as water-saturated conditions for pCO2 = 1,



Page 42 of 51

42 812

60, and 90 bar (Table S-4); and 3) Atomic coordinates for the Ca-analogues of hydrated

813

Mg carbonate structures (Tables S-5, 6, and 7).

814 815 816 817

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(2) Benson, S. M.; Cole, D. R., CO2 Sequestration in Deep Sedimentary Formations. Elements 2008, 4, 325-331.

820 821 822

(3) Fukushi, K.; Munemoto, T.; Sakai, M.; Yagi, S., Monohydrocalcite: A Promising Remediation Material for Hazardous Anions. Sci. Technol. Adv. Mat. 2011, 12, 53185327.

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(4) Chen, T.; Neville, A.; Yuan, M. D., Influence of Mg2+ on CaCO3 Formation-Bulk Precipitation and Surface Deposition. Chem. Eng. Sci. 2006, 61, 5318-5327.

825 826

(5) Laugesen, J. L., Density Functional Calculations of Elastic Properties of Portlandite, Ca(OH)2. Cement Concr. Res. 2005, 35, 199-202.

827 828 829

(6) Lee, K.; Wagermaier, W.; Masic, A.; Kommareddy, K. P.; Bennet, M.; Manjubala, I.; Lee, S. W.; Park, S. B.; Colfen, H.; Fratzl, P., Self-Assembly of Amorphous Calcium Carbonate Microlens Arrays. Nat. Commun. 2012, 3, 1-7.

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(7) Gebauer, D.; Volkel, A.; Colfen, H., Stable Prenucleation Calcium Carbonate Clusters. Science 2008, 322, 1819-1822.

832 833 834 835

(8) Blue, C. R.; Giuffre, A.; Mergelsberg, S.; Han, N.; De Yoreo, J. J.; Dove, P. M., Chemical and Physical Controls on the Transformation of Amorphous Calcium Carbonate into Crystalline CaCO3 Polymorphs. Geochim. Cosmochim. Ac. 2017, 196, 179-196.

836 837

(9) Blue, C. R.; Dove, P. M., Chemical Controls on the Magnesium Content of Amorphous Calcium Carbonate. Geochim. Cosmochim. Ac. 2015, 148, 23-33.

838 839 840

(10) Rodriguez-Blanco, J. D.; Shaw, S.; Benning, L. G., The Kinetics and Mechanisms of Amorphous Calcium Carbonate (ACC) Crystallization to Calcite, Via Vaterite. Nanoscale 2011, 3, 265-71.

841 842 843

(11) Radha, A. V.; Fernandez-Martinez, A.; Hu, Y.; Jun, Y.-S.; Waychunas, G. A.; Navrotsky, A., Energetic and Structural Studies of Amorphous Ca1−xMgxCO3·nH2O (0⩽x⩽1). Geochim. Cosmochim. Ac. 2012, 90, 83-95.


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(12) Gale, J. D.; Demichelis, R.; Raiteri, P.; Quigley, D.; Gebauer, D., Nucleation of Amorphous Calcium Carbonate: A Combined Theoretical and Experimental Perspective. Abstr. Pap. Am. Chem. S. 2011, 242.

847 848 849 850

(13) Cartwright, J. H. E.; Checa, A. G.; Gale, J. D.; Gebauer, D.; Sainz-Diaz, C. I., Calcium Carbonate Polyamorphism and Its Role in Biomineralization: How Many Amorphous Calcium Carbonates are There? Angew. Chem. Int. Edit, 2012, 51, 1196011970.

851 852 853

(14) Clarkson, J. R.; Price, T. J.; Adams, C. J., Role of Metastable Phases in the Spontaneous Precipitation of Calcium Carbonate. J. Chem. Soc. Faraday T. 1992, 88, 243-249.

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(15) Gebauer, D.; Colfen, H., Prenucleation Clusters and Non-Classical Nucleation. Nano Today 2011, 6, 564-584.

856 857 858

(16) Montes-Hernandez, G.; Renard, F., Time-Resolved in Situ Raman Spectroscopy of the Nucleation and Growth of Siderite, Magnesite, and Calcite and Their Precursors. Cryst. Growth Desi. 2016, 16, 7218-7230.

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(17) Ito, T., Factors Controlling the Transformation of Natural Ikaite from Shiowakka, Japan. Geochem. J. 1998, 32, 267-273.

861 862 863 864

(18) Rysgaard, S.; Sogaard, D. H.; Cooper, M.; Pucko, M.; Lennert, K.; Papakyriakou, T. N.; Wang, F.; Geilfus, N. X.; Glud, R. N.; Ehn, J.; McGinnis, D. F.; Attard, K.; Sievers, J.; Deming, J. W.; Barber, D., Ikaite Crystal Distribution in Winter Sea Ice and Implications for CO2 System Dynamics. Cryosphere 2013, 7, 707-718.

865 866 867

(19) Bischoff, J. L.; Fitzpatrick, J. A.; Rosenbauer, R. J., The Solubility and Stabilization of Ikaite (CaCO3•6H2O) from 0˚C to 25˚C - Environmental and Paleoclimatic Implications for Thinolite Tufa. J. Geol. 1993, 101, 21-33.

868 869 870

(20) Buchardt, B.; Seaman, P.; Stockmann, G.; Vous, M.; Wilken, U.; Duwel, L.; Kristiansen, A.; Jenner, C.; Whiticar, M. J., Submarine Columns of ikaite Tufa. Nature 1997, 390, 129-130.

871 872 873

(21) De Lurio, J. L.; Frakes, L. A., Glendonites as a Paleoenvironmental Tool: Implications for Early Cretaceous High Latitude Climates in Australia. Geochim. Cosmochim. Ac. 1999, 63, 1039-1048.

874 875 876

(22) Rodriguez-Blanco, J. D.; Shaw, S.; Bots, P.; Roncal-Herrero, T.; Benning, L. G., The Role of Mg in the Crystallization of Monohydrocalcite. Geochim. Cosmochim. Ac. 2014, 127, 204-220.

877 878 879

(23) Coleyshaw, E. E.; Crump, G.; Griffith, W. P., Vibrational Spectra of the Hydrated Carbonate Minerals Ikaite, Monohydrocalcite, Lansfordite and Nesquehonite. Spectrochim. Ac. A: Molec. Biomolec. Spectr. 2003, 59, 2231-2239.



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Page 51 of 51

51 1094

For TOC only

1095 1096 90 bar pCO 2 Water Saturated

-50

Free Energy (kJ/mol)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


Ca(OH)2 (xtl)

-70

CaCO 3•6H 2O Prenucleation clusters

-90

CaCO3 •6H2 O (xtl)

-110

ACC•H 2O

CaCO3 •H2 O (xtl) CaCO3 (xtl)

-130 -150

1097

275

325

375K


Ab Initio Thermodynamics of Hydrated Calcium Carbonates and

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