Direct Experimental Measurement of Water Interaction Energetics in

Article pubs.acs.org/crystal

Direct Experimental Measurement of Water Interaction Energetics in Amorphous Carbonates MCO3 (M = Ca, Mn, and Mg) and Implications for Carbonate Crystal Growth A. V. Radha and Alexandra Navrotsky* Peter A. Rock Thermochemistry Laboratory, NEAT ORU, University of California, Davis, California 95616, United States ABSTRACT: Interaction of carbonate surfaces with water plays a crucial role in carbonate nucleation and crystal growth. This study provides experimental evidence for the existence of two different types of water having distinct energetics in amorphous carbonates, MCO3 (M = Ca, Mn, and Mg). The adsorption enthalpy curves obtained using a combination of gas sorption and microcalorimetry show two different energetic regions, which correspond to weakly bound restrictedly mobile and strongly bound rigid H2O components. For weakly bound water, adsorption enthalpies of amorphous calcium carbonate (ACC) (−55.3 ± 0.9 kJ/mol), amorphous manganese carbonate (AMnC) (−54.1 ± 0.8 kJ/mol), and amorphous magnesium carbonate (AMgC) (−56.1 ± 0.4 kJ/ mol) fall in the same range, suggesting their interaction modes may be similar in all amorphous phases. Water adsorption enthalpies of crystalline nanocalcite (−96.3 ± 1 kJ/mol) and nano-MnCO3 (−65.3 ± 3 kJ/mol) measured in previous studies are more exothermic than values for ACC (−62.1 ± 0.7 kJ/mol) and AMnC (−54.1 ± 0.8 kJ/mol) and could provide a driving force for crystallization of ACC and AMnC in the presence of water. The differences in water adsorption behavior between amorphous and naocrystalline material have significant implications for crystal growth, biomineralization, and carbonate geochemistry. the presence of additives such as water, Mg2+, silica, phosphate, organic, or biomolecules and confinement in porous structures.34−41 Calcite crystal growth on solid-state transformations has been reported to occur by ion-mediated dissolution and precipitation in synthetic ACC, while in biogenic ACC, calcite may form by nucleation and growth of nanosphere.35 In the case of template-directed calcite crystal growth, presence of organic or biological macromolecules promote heterogeneous nucleation and alter the thermodynamic barrier by reducing interfacial energies at the crystalfluid, fluid-substrate, and crystal-substrate interfaces.42 Among inorganic additives, Mg2+ ions and water appear to play a crucial role in stability and transformation of both synthetic and biogenic ACC. In our previous study on crystallization, energetics of the amorphous MgCO3−CaCO3 system,37 we have identified two amorphous solid solution precursors that provide low-energy pathways for crystallization with minimum crystallization energy, one below 30 mol % and the other near 50% MgCO3. In general, biogenic ACCs contain up to 30 mol % MgCO321,22,26,34 and this thermodynamic data suggests that marine organisms may select this minimum crystallization energy pathway for biomineralization, while the solid solution region near 50% MgCO3 may be a dolomite precursor. The implications of thermodynamic studies of amorphous and nano

1. INTRODUCTION Carbonate crystallization has been a subject of interest due to the possibility of several pathways for its formation by classical and nonclassical crystal nucleation.1−5 The nonclassical pathways involve formation of prenucleation clusters,6−8 metastable liquid-like (binodal) phase separated regions,9−11 dynamically ordered liquidlike oxyanion polymers (DOLLOP),12,13 polymer-induced liquid precursors (PILP),14 mesocrystals,15 and/or amorphous and nanophase precursors16 in the early stages of crystal growth. Hamm et al. have reviewed these pathways and proposed a mechanistic interpretation to put these distinct species into context.4 Metastable amorphous calcium carbonate (ACC) has been observed in inorganic synthesis17−20 and as biominerals in marine organisms such as sea urchins, bivalves, sponges, and crustaceans.21−23 Biogenic ACC associated with magnesium and an organic or biological macromolecular matrix is found to persist longer than pure synthetic ACC, which offers organisms better control over its transformation to complex crystalline biominerals that serve as vital structural components.21,23,24 Structural analyses of synthetic, biogenic, and Mgdoped ACCs have established the existence of amorphous forms with different hydration levels as well as local structures resembling crystalline monohydrocalcite, vaterite, aragonite, dolomite, and calcite.25,18,26−33 The mechanism of ACC crystallization seems complex due to existence of multiple forms of ACC under different environments. Structural evolution and transformation of various ACC to crystalline CaCO3 appear to be controlled by © 2014 American Chemical Society

Received: June 16, 2014 Revised: November 18, 2014 Published: December 2, 2014 70

dx.doi.org/10.1021/cg500878w | Cryst. Growth Des. 2015, 15, 70−78

Crystal Growth & Design

Article

Table 1. Summary of Water Adsorption Calorimetry Dataa sample ID

water loss from TGA

sample degas condition before calorimetry

T (°C)

H2O loss (%)

T (°C)

ACC

30−125

10.6 ± 0.5 (5)

25

ACC

30−300

15.6 ± 0.7 (5)

100

AMgC AMgC AMnC nanocalciteb nano-MnCO3b

30−150 30−300 30−200

31.0 ± 1.0 (2) 35.2 ± 0.3 (2) 7.5 ± 1.5 (2)

25 100 25 150 150

water adsorption calorimety

H2O loss (%)

BET surface area (m2/g)

enthalpy of water adsorption ΔHH2Oads (kJ/mol)

surface water coverage (H2O/nm2)

10.9 ± 0.1 (3) 15.6 ± 1.7 (5) 25.0 (1) 36.91 (1) 4.97 (1)

16.77 ± 1.41 (7)

−55.31 ± 0.9 (5)

15.5 ± 4.2 (5)

14.69 ± 2.19 (4)

−62.13 ± 0.68 (8)

7.2 ± 0.5 (8)

52.46 ± 2.19 (2) 48.58 ± 0.22 (2) 50.83 ± 2.73 (3)

−56.13 −58.5 −54.12 −96.26 −65.3

± ± ± ± ±

0.39 (2) 0.62 (2) 0.77 (5) 0.9663 366

39.5 8.0 6.5 5.0 9

± ± ± ± ±

4.6 (2) 0.8 (3) 1.7 (5) 0.663 166

The values in parentheses are a number of measurements. The average values for ACC are obtained from calorimetric measurements on five different freshly prepared samples, for AMnC from two freshly synthesized samples, and for AMgC from a single sample. bThe data for nanophase calcite and nanophase MnCO3 are from previous studies.6366 a

samples were precipitated by adding 200 mL of 0.02 M metal chloride or sulfate (CaCl2, MgCl2, and MnSO4) solution to an equal volume of 0.02 M sodium carbonate solution with constant stirring. For ACC synthesis, a high pH solution containing 0.02 M Na2CO3 and 0.2 M NaOH was used. For AMnC synthesis, nitrogen gas was bubbled through all solutions for at least 30 min to remove any dissolved oxygen. All solutions were cooled in a refrigerator at 4 °C prior to mixing. The precipitate was immediately filtered under vacuum, washed with acetone several times, and vacuum-dried overnight at 25 °C. All sample characterization and water adsorption calorimetry on amorphous samples were carried out within 4−6 days to minimize any changes due to aging and crystallization. Amorphous samples were hydrated with composition MCO3 ·nH2O (M = Ca, Mn, and Mg), and the water content were determined from the first weight loss step in the thermogravimetric analysis (TGA) curve. Amorphous metal carbonates (M = Ca, Mn, and Mg) were freshly prepared to minimize the crystallization of initial materials. Calorimetric measurements were done on freshly prepared (a) five ACC samples as it crystallizes fast within a week, (b) two AMnC samples, and (c) a single AMgC sample. 2.2. Sample Characterization. Phase identification was done using powder X-ray diffraction (XRD) on a Bruker AXS D8 Advance diffractometer with a LynxEye solid state detector and Cu Kα radiation (Bruker AXS, Inc.; Madison, WI). The powder XRD patterns were collected using a zero-background sample holder in the 2θ range of 10−90° with a step size of 0.01° and a collection time of 0.5−2 s/step. The detection limit of XRD for a crystalline phase is typically 1−2% in volume. Thermogravimetric analysis and differential scanning calorimetry (TGA/DSC) measurements were done using a Netzsch STA 449 system (Netzsch GmbH, Selb, Germany). The sample in a platinum crucible was heated from 30 to 1200 °C at 10 °C/min with argon flow. A buoyancy correction was made using an empty platinum crucible run. A sapphire pellet in a platinum crucible was heated in argon flow from 30 to 1200 °C at 10 °C/min to create a DSC calibration file. The TGA/DSC data were analyzed using Netzsch Proteus. The water content was determined from the average of TGA weight loss curves from at least two separate sample runs. The surface area of the samples was characterized by nitrogen adsorption measurements using the Brunauer−Emmett−Teller (BET) method with a Micromeritics ASAP 2020 instrument. Approximately 0.05 g of each sample was degassed at 150 °C for 2 h. A five point N2 adsorption isotherm was collected in duplicate for all samples in the P/P0 relative pressure range of 0.05−0.3, where P0 is the saturation pressure at 196 °C. BET surface areas for all samples measured before and after water adsorption calorimety were within the uncertainties of the values reported in Table 1. 2.3. Water Adsorption Calorimetry. The instrument and methodology used for water adsorption calorimetry has been described earlier.50 A typical water adsorption calorimetric experiment involves three steps. The first step is degassing the sample placed in a silica glass forked tube under vacuum (with or without heating as appropriate) using the degas port of a Micromeritics ASAP2020

Ca−Mg−Fe−Mn carbonate system on crystallization of common rhombohedral carbonate minerals have been discussed in an earlier review article.16 Earlier studies have indicated synthetic and biogenic ACC crystallize on dehydration in air or on heating as well as in aqueous solution.6,7,20,21,32 Several combined thermal and structural evolution studies on ACC dehydration on heating have shown the existence of multiple dehydration steps involving various hydrated intermediates having different local structures.43−47 However, to gain more insight into the driving force behind different ACC crystallization pathways, it is necessary to understand the kinetics and thermodynamics of dehydration of ACC. Calorimetric studies have revealed dehydration driven crystallization of ACC is in fact thermodynamically favored as it follows an energetically downhill path via amorphous phases of decreasing water content.48 Ihli et al. have used thermal analysis and solid state NMR to determine the activation energies associated with liberation of different water fractions that exist in different structural environments during crystallization of ACC both in aqueous solution and air.44 Saharay et al.43,49 have determined the dehydration enthalpies for ACC using molecular dynamic studies and showed that dehydration enthalpies become more exothermic with decreasing water content due to formation of progressively more ordered local structures. However, there is a lack of experimental data on such interaction energetics to benchmark these calculations. In this work, we report the direct measurement of enthalpy of water adsorption on ACC using a combination of a gas sorption system and a Calvet microcalorimeter, as described in many previous studies.48,50−57 The simultaneous measurement of differential adsorption enthalpy and water pressure as a function of amount of accurately dosed water vapor gives a detailed picture of the interaction of water with carbonate surface as a function of coverage. Here we report water adsorption data for ACC and several other amorphous carbonates, followed by comparison and analysis of the water adsorption energetics of ACC with those of amorphous manganese carbonate (AMnC) and amorphous magnesium carbonate (AMgC) as well as with crystalline nanocalcite and nano-MnCO3.

2. EXPERIMENTAL SECTION 2.1. Sample Syntheses. Amorphous MCO3 (M = Ca, Mg, and Mn) samples were synthesized using earlier reported methods with reagent grade chemicals and ultrapure water.20,37,48 Briefly, amorphous 71



Article

instrument. The second step is measuring the BET surface area with nitrogen adsorption using a Micromeritics ASAP2020 instrument. The third step is measuring the enthalpy of adsorption of each incremental dose of water vapor using a coupled system consisting of the Micrormeretics ASAP2020 gas adsorption analyzer and a Setaram Sensys Calvet microcalorimeter. Each water dose (∼1−2 μmol) generates a distinct calorimetric peak, and the area under the peak provides the corresponding heat of adsorption (differential enthalpy). At least 2 runs were done on each sample under two different vacuum degas conditions (6−8 h at 25 °C and 15 min at 100 °C) between the measurements. A blank with an empty tube was run to correct the data for water adsorbed on the forked tube wall. Table 1 summarizes the BET surface area and water adsorption calorimetric data for amorphous MCO3 (M = Ca, Mn, and Mg) samples.

3. RESULTS AND DISCUSSION 3.1. X-ray Diffraction and Thermal Analyses. The formation of X-ray amorphous metal carbonates (M = Ca, Mn, and Mg) was confirmed by the absence of Bragg reflections in the powder XRD patterns (Figure 1). Typical TGA curves of all

Figure 1. Powder XRD patterns of representative amorphous metal carbonates MCO3 (M = Ca, Mn, and Mg).

amorphous phases show two step weight loss due to dehydration and thermal decomposition of dehydrated MCO3 to MO and CO2 (M = Ca, Mn, and Mg) (see Figure 2). Simultaneous DSC curves in Figure 2 show three peaks corresponding to enthalpies associated with dehydration, crystallization, and decomposition for ACC and AMnC and two endothermic peaks for enthalpies of dehydration and decomposition for AMgC. The absence of exothermic crystallization peak in AMgC suggests that it decomposes at 400 °C without crystallization, while AMnC undergoes simultaneous crystallization and decomposition at 400 °C. ACC forms a crystalline phase (∼320 °C) before its decomposition into oxide and CO2 (∼600 °C). Therefore, the water contents of the samples were estimated from the first weight loss step up to the crystallization point for ACC and AMnC and from 30 to 200 °C for AMgC (see Table 1). The changes in slopes of the dehydration weight loss curves observed in all three amorphous samples in Figure 2 could possibly arise due to the existence of water in different environments. For ACC, the temperature range and the weight loss corresponding to dehydration steps in its TGA can be roughly identified as (a) ACC dehydration step-I: 30−125 °C

Figure 2. Typical TG_DSC profiles of amorphous metal carbonates MCO3 (M = Ca, Mn, and Mg).

with a weight loss of 10.6 ± 0.5%, (b) ACC dehydration stepII: 125−300 °C with a weight loss of 4.8 ± 0.6%, and (c) ACC and nanocalcite dehydration step-III: 300−500 °C with a 1.4 ± 0.3% weight loss (average from 5 TGA runs). The three dehydration TGA steps qualitatively indicate the existence of three types of water, including one remaining after crystallization for ACC. Recent thermal and structural studies have characterized these three steps in ACC dehydration as loss of fluidlike mobile → restrictedly mobile and rigid H2O components → final loss of hydroxyl and trapped rigid and mobile water.44,46 Further, these studies suggested that the fluidlike mobile water behaves more like physically adsorbed water and is lost in early stages of dehydration, while restrictedly mobile and rigid H2O components are lost at intermediate temperatures. The loss of the final fraction of hydroxyl/trapped water above 300 °C is found to be crucial as it triggers the crystallization.44,46,49 Similarly, Ihli et al.44 have 72



Article

identified four stages in the mechanism of dehydration of ACC leading to crystallization as stage-I: loss of surface-bound water at ∼40 °C, stage-II: loss of water from the interior of the ACC with shrinkage of ACC particles around 140 °C, stage-III: expulsion of the most deeply located water around 290 °C, and stage-IV: crystallization to calcite at 315 °C. In this work, to avoid any structural changes upon heating or crystallization, the first set of water adsorption calorimetric experiments was carried out on amorphous samples degassed 6−8 h under vacuum at 25 °C. This resulted in only partial dehydration of ACC (10.9 ± 0.1 wt %), probably due to loss of all physisorbed fluidlike water and some of loosely surface bound restrictedly mobile H2O and corresponds to the TGA weight loss observed in the ACC dehydration step-I from 30 to 125 °C. Therefore, a second set of water adsorption experiments was carried out by degassing samples by isothermal heating under vacuum at 100 °C for 15 min to remove all interior bulk water associated with the amorphous phase before crystallization (15.6 ± 1.7 wt %), and this includes TGA water loss corresponding to ACC dehydration steps I and II due to dynamic heating in argon flow from 30 to 300 °C at 10 °C/min. The samples degassed at 100 °C were X-ray amorphous. Sample degassing at 100 °C could possibly bring about subtle changes within the amorphous structure as reported in several earlier studies for ACC,33,46 but these were not considered in the present study. Interestingly, some preliminary pair distribution function (PDF) data from X-ray scattering heating experiments done by us on ACC at APS sector 11-ID-B, Argonne National Laboratory, have shown subtle changes in the PDF patterns around 100 °C, possibly due to some sort of structural changes around this temperature. These data have not been published as there were substantial variations in temperature of the furnace used for heating the sample in this temperature range. 3.2. Water Adsorption Calorimetry Data Analyses. Figure 3a shows representative profiles of enthalpies of water adsorption (differential enthalpies) as a function of surface coverage (H2O per nm2) for ACC degassed under vacuum at 25 and 100 °C. The differential enthalpies become less exothermic with successive dosing and finally reach the enthalpy of condensation of bulk water at 25 °C (−44 kJ/ mol). The total water adsorbed up to this coverage is strongly bound (chemisorbed water), and the remaining water with differential enthalpy of −44 kJ/mol represents physically adsorbed water. The sum of the differential enthalpies of adsorption divided by the total water up to this coverage (integral enthalpy) corresponds to the enthalpy of chemisorbed water. In all this work, the reference state for water is the vapor, whose enthalpy does not depend on pressure at constant temperature since H2O can be considered an ideal gas at these low pressures. The calculated integral enthalpies (ΔHads) for ACC vacuum degassed at 25 and 100 °C are −55.3 ± 0.9 and −62.1 ± 0.7 kJ/mol, and the corresponding water coverage values at this point are 15.5 ± 4.2 and 7.2 ± 0.5 H2O/nm2 respectively. The integral enthalpy (ΔHads) for the sample vacuum degassed at 25 °C is less exothermic than that of the sample degassed at 100 °C due to its partial dehydration and probably gives the energetics of loosely bound restrictedly mobile H2O. The integral enthalpy (ΔHads) becomes more exothermic on degassing at 100 °C as it removes most of the bound water and represents the dehydration energy for ACC. The differential water adsorption enthalpy curves represent the enthalpy distribution for different binding energy sites. However, both TGA and water adsorption data show the

Figure 3. Typical water adsorption enthalpy curves of amorphous metal carbonates MCO3 (M = Ca, Mn, and Mg) as a function of surface coverage.

change in slopes, which suggest that these energetics fall into two distinct energetic ranges as discussed below. Earlier experimental and computational studies on crystallization pathways of ACC have shown evidence for the presence of hydrated and anhydrous intermediates with different local structures and bonding for water.6,44,46,49 The water adsorption enthalpy profiles of ACC in Figure 3a shows the existence of two different energetic regions for both 25 and 100 °C degassed samples. The first region corresponding to the adsorption enthalpies for the initial doses of water (coverage up 73



Article

to 3 H2O/nm2) is more exothermic and shows a sharp decrease in magnitude from −100 to −60 kJ/mol. The second energetic region (−60 to −44 kJ/mol) shows different behavior depending on the sample degas conditions. The sample degassed at room temperature shows an extended tail with coverage of about 15.5 ± 4.2 H2O/nm2 with less exothermic adsorption enthalpy (−55.31 ± 0.9 kJ/mol), indicating the existence of a large number of low-energy sites with weak water interactions. The sample degassed at 100 °C has more exothermic adsorption enthalpy (−62.1 ± 0.7 kJ/mol) with a coverage of 7.2 ± 0.5 H2O/nm2, which probably indicates that heating not only removes all adsorbed water but also eliminates low-energy adsorption sites for water. This measured trend in the adsorption energetics of ACC (ranging from −44 to −94 kJ/mol) supports the recently reported computed hydration enthalpies for ACC by Saharay et al.,49 which showed progressively more exothermic enthalpies (−65 to −85 kJ/ mol for H2O/CaCO3 = 1.049 to 0.25). They attributed this energetic trend to evolution of more ordered local structure as a result of favorable H-bond interactions and increasing coordination of carbonate group oxygens with Ca 2+ . Furthermore, the different energetic regions in water adsorption enthalpy curves of ACC can be correlated to the TGA dehydration steps and consequently to different types of water reported in the earlier structural studies.44,46,49 The energetics of water adsorption in the first step probably correspond to the strongly bound restrictedly mobile and rigid H2O components (dehydration step-II in TGA of ACC). ACC degassed at 100 °C appear to have a higher number of these sites and hence has slightly more exothermic adsorption enthalpy than the 25 °C degassed ACC sample. The energetics of the second step in the 25 °C vacuum degassed sample corresponds to the existence of a large number of weakly bound restrictedly mobile H2O components (dehydration step-I in TGA of ACC). Therefore, the trends in ACC water adsorption curves provide direct experimental evidence for the existence of different types of water with distinct energetics in the ACC structure. 3.3. Comparison of Water Adsorption Energetics of ACC with AMnC and AMgC. Figure 3 shows the representative enthalpies of water adsorption as a function of surface coverage (H2O per nm2) for AMnC and AMgC. The integral enthalpies (ΔHads) of AMnC and AMgC for vacuum degassed at 25 °C are −54.12 ± 0.77 kJ/mol and −56.13 ± 0.39 kJ/mol and the corresponding water coverages are 6.5 ± 1.7 and 39.5 ± 4.6 H2O/nm2, respectively. Interestingly, the water adsorption enthalpies of AMnC and AMgC fall in the same range as that of ACC (ΔHads = −55.3 ± 0.9 kJ/mol and water coverage = 15.5 ± 4.2 6 H2O/nm2) for samples degassed in vacuum at 25 °C. This suggests that the nature of bonding interactions of loosely bound water is similar in Ca, Mn, and Mg amorphous phases. However, these samples show significant differences in their water coverage values with amorphous MgCO3 having the maximum and amorphous MnCO3 with minimum water coverage. Similar to earlier observation for ACC, the integral enthalpy (ΔHads) of AMgC became slightly more exothermic (−58.5 ± 0.6 kJ/mol) upon degassing at 100 °C with corresponding decrease in the water coverage (8.0 ± 0.8 H2O/nm2). Unlike values for 25 °C vacuum degassed samples, the adsorption enthalpies of ACC and AMgC are different for samples degassed at 100 °C, which indicates that the interaction energetics for strongly bound water are different

in Ca and Mg amorphous phases. The adsorption enthalpy of ACC (−62.1 ± 0.7 kJ/mol) is more exothermic than that of AMgC (−58.5 ± 0.6 kJ/mol) for samples degassed at 100 °C, even though the ionic radius of Mg2+ is smaller than that of Ca2+. In general, one would expect hydration enthalpies to be ionic size-dependent and follow the trends observed in absolute hydration enthalpies of gaseous ions (ΔHh°) with smaller cations having more exothermic ΔHh° values than larger cations (Mg2+ = −1926 < Mn2+ = −1851 < Ca2+ = −1579 kJ/ mol).58 However, several theoretical and experimental studies have shown the reversal of hydration energy trends, especially for Mg2+ and Ca2+ cations.59,60 Rodriguez-Cruz et al.59 attributed reversal in binding energies of hexahydrated divalent Mg2+ and Ca2+ cations at high temperature (Ba2+ < Sr2+ < Mg2+ < Ca2+ at T > 80 °C) to the presence of two magnesium hexahydrate isomers with six water molecules in the first solvation shell at low temperature and a different isomer with four water molecules in the inner shell and two water molecules in the second shell at high temperature. Peschke and coworkers60 used gas phase ion−water molecule equilibria to show that the first six water molecules are in the inner shell for M(H2O)62+ with M = Mg2+, Ca2+, and Sr2+, and the sequential hydration enthalpy for Ca(H2O)62+ (−105.86 kJ/mol) was found to be more exothermic than that for Mg(H2O)62+ (−102.9 kJ/mol) due to crowding in the inner shell in the Mg2+ complex. Among various hydration enthalpy studies on crystalline rhombohedral carbonate surfaces, atomistic computer modeling by De Leeuw61 reported some deviations in hydration energetics for magnesian calcites and dolomite. The average hydration energy for magnesite, MgCO3 (−134 kJ/ mol) was more exothermic than for calcite (−79 kJ/mol), reflecting the strong affinity of the smaller Mg2+ cations for water. Interestingly, the average hydration enthalpies for dolomite, Ca0.5Mg0.5CO3 (−101 kJ/mol) and magnesian calcite (−94 kJ/mol) surfaces were much lower compared to that for magnesite because of inaccessibility of the surface magnesium ions due to surface relaxations and rotation of the carbonate groups. These observations imply that interaction of Mg2+ species with water in different structures is complex and hence their hydration energies vary depending on coordination number, temperature, and presence of water in first or outer solvation shells and accessibility of Mg2+ cations. Unfortunately, no such information about water structure and co-ordination in AMgC is available to explain the anomalous hydration enthalpy values of ACC and AMgC at 100 °C observed in this work. A new study on the cation hydration effect on crystallization of Ca−Mg−CO3 systems by Xu et al.62 has revealed the existence of an additional thermodynamic barrier (along with cation hydration effect) to form long-range ordering of Mg and CO3 ions due to entropy decrease resulting from the reduced freedom of carbonate caused by the lattice limitation of the compact Mg−O octahedra in crystalline magnesite. This implies that the Mg−O octahedra in AMgC are probably more relaxed, which could explain the lower hydration enthalpy observed in AMgC. 3.4. Comparison of Water Adsorption Energetics of ACC and Nano-Calcite. In order to understand the role of water in carbonate crystal growth from the amorphous phase, we compared the water adsorption energetics of ACC and nanophase calcite. Our previously measured water adsorption enthalpy curves for nanophase calcite showed the existence of three energetically distinct water coverage regions, which agreed with the molecular dynamic simulation for the existence 74



Article

nanosecond timescales and hence could spontaneously become amorphous. However, not much information is available on mechanism of subsequent crystal growth of ordered crystalline phase from the disordered ACC. In general, crystallization of ACC in the presence of water is attributed to dissolution− reprecipitation reactions.35 A recent investigation by Ihli et al.44 concluded that both in solution and in air, the dehydration and associated structural changes are similar, and the loss of final fraction of water (15 wt % of total) triggers the crystallization via “solid-state” transformation. The final dehydration step was found to have very high activation energy (∼245 kJ/mol), and hence at room temperature the initial crystals form by dissolution−reprecipitation in solution or on surfaces.44 However, calorimetric studies conducted in our group on energetics of amorphous and nanocrystalline carbonates have revealed the existence of dissimilar thermodynamic behavior for crystal growth of ACC under different circumstances. The hydration enthalpies of ACC (−62.1 ± 0.7 kJ/mol) and calcite (−96.3 ± 1.0 kJ/mol)63 clearly illustrate that the more exothermic hydration enthalpy of calcite provides a thermodynamic driving force for ACC crystallization in the presence of water or in aqueous solution. The dehydration of ACC in air or on heating triggers the crystallization of ACC since it is an energetically downhill process as reported in our previous study and hence thermodynamically favored.48 However, these thermodynamic arguments do not provide direct information on the heights of kinetic barriers to crystallization. The role of amorphous Ca−Mg−Mn carbonates in the crystallization of rhombohedral carbonates is not fully understood. In recent years, our group has identified several amorphous and nanophase precursors in the Ca−Mg−Fe− Mn carbonate system and conducted a series of thermodynamic studies to show that these precursors could provide low-energy pathways for crystallization of common rhombohedral carbonate minerals such as calcite (CaCO3), magnesian calcites (Ca1−xMgxCO3), dolomite (Ca0.5Mg0.5CO3), siderite (FeCO3), and rhodochrosite (MnCO3).16,37,48,63,66,69 The crystallization enthalpies appear to be controlled by cation size and become less exothermic with increase in ionic radius.66,69 These thermodynamic studies have also made significant contribution toward understanding the role of ACC in biomineralization. We have performed an experimental thermodynamic study of both synthetic and biologically produced (biomineralization by California purple sea urchin spicules) amorphous calcium carbonate and identified that hydrated and dehydrated amorphous calcium carbonate precursors provide the energetically favorable pathway for calcite crystallization in both chemical and biologically mediated processes.48 Biogenic ACCs are commonly found associated with magnesium ions, phosphates, and macromolecules (proteins) and need careful extraction since they crystallize on exposure to air or water.21,23,34 The goal of the present work is to understand the role of water in stability and crystallization of amorphous carbonates. The water adsorption energetics of amorphous and nanocrystalline carbonates reported in this work provides experimental evidence for existence of an additional thermodynamic driving force for amorphous carbonate crystallization in the presence of water. The trends in water interaction energetics have greater significance as it could probably explain the driving force behind the rapid ACC crystallization observed during biomineralization and in laboratory settings in the presence of water.

of three different conformational modes for water adsorption on different active sites of calcite surface.63−65 The differential enthalpies of nanocalcite range from −160 to −44 kJ/mol,63 while those of ACC fall in the range of −99 to −44 kJ/mol. This enthalpy difference at low coverage suggests the presence of high-energy strong binding sites in calcite which either do not exist or are less prevalent in ACC. Unlike the water adsorption enthalpy curve of ACC with two different energetic regions, nanophase calcite shows an additional, more exothermic adsorption region that favors water adsorption. The differential enthalpies of nanophase calcite up to 3 H2O per nm2 coverage are more exothermic, whereas values between 3 and 5.1 H2O per nm2 coverage increase rapidly and reach a plateau of −44 kJ/mol.63 This suggests that about 60% of the chemisorbed water (up to 3 H2O per nm2) on the calcite surface is bound strongly on high-energy active sites compared to water molecules between 3 and 5.1 H2O per nm2, which are bound less strongly. The nonlinear trend in enthalpies between 3 and 5.1 H2O per nm2 suggests the existence of a range of lower-energy sites. Therefore, the water adsorption enthalpy of nanocalcite (−96.3 ± 1.0 kJ/mol)63 is more exothermic than that of ACC (−62.13 ± 0.68 kJ/mol), and hence in the presence of water nanocalcite formation is energetically additionally favored over ACC. Further evidence for this can also be seen in the method used for synthesis of ACC. Freshly precipitated ACC persists only with immediate filtration and subsequent drying under vacuum at room temperature after washing with acetone; otherwise, it crystallizes to calcite on aging in water within a few minutes. These results clearly demonstrate that the exothermic hydration enthalpy provides an additional thermodynamic driving force for ACC crystallization in the presence of water. A similar trend was also observed in the case of amorphous and nanophase MnCO3, where the water adsorption enthalpy of nano-MnCO3 (−65.3 ± 3 kJ/mol)66 is more exothermic than that of AMnC (−54.1 ± 0.8 kJ/mol). Interestingly, ACC crystallizes in 3−4 days as opposed to 40 days for AMnC, probably due to larger driving force for ACC crystallization as a consequence of more exothermic water adsorption enthalpy of calcite compared to rhodochrosite. 3.5. Implications of Thermodynamic Approach for Carbonate Crystal Growth, Biomineralization, and Geochemistry. Interaction of surfaces with an aqueous phase has been reported to play a crucial role in nucleation and crystal growth of carbonate as well as in other systems.67 The water adsorbed on the surface can act as a reaction medium and could influence the mobility of ions on the solid surface or hinder the surface reaction by blocking the surface active sites.65 Therefore, the stability and residence time of the adsorbed water, which depends on the hydration energy, can strongly influence crystal growth and dissolution on the surface.64,65,68 For calcium carbonate system, earlier simulation studies by Raitri and Gale,6 suggested that addition of both ion pairs as well as water into ACC nanoparticle clusters during carbonate nucleation is thermodynamically driven. Interestingly, the incorporation of water into ACC clusters is shown to be sizedependent and ACC clusters are found to be more stable than the corresponding clusters for bulk calcite only below a critical diameter of 4 nm.6 This is supported by several experimental evidence that show the formation of smaller clusters in the early stages of nucleation.3,8,9 A recent molecular simulation study by Bano et al.68 showed that at small sizes in the range of 1.8−4.1 nm, nanocalcite may not retain an ordered structure on 75



Article

(10) Faatz, M.; Gröhn, F.; Wegner, G. Amorphous Calcium Carbonate: Synthesis and Potential Intermediate in Biomineralization. Adv. Mater. 2004, 16 (12), 996−1000. (11) Wolf, S. E.; Muller, L.; Barrea, R.; Kampf, C. J.; Leiterer, J.; Panne, U.; Hoffmann, T.; Emmerling, F.; Tremel, W. Carbonatecoordinated metal complexes precede the formation of liquid amorphous mineral emulsions of divalent metal carbonates. Nanoscale 2011, 3 (3), 1158−1165. (12) Demichelis, R.; Raiteri, P.; Gale, J. D.; Quigley, D.; Gebauer, D. Stable prenucleation mineral clusters are liquid-like ionic polymers. Nat. Commun. 2011, 2, 590. (13) Wallace, A. F.; Hedges, L. O.; Fernandez-Martinez, A.; Raiteri, P.; Gale, J. D.; Waychunas, G. A.; Whitelam, S.; Banfield, J. F.; De Yoreo, J. J. Microscopic Evidence for Liquid-Liquid Separation in Supersaturated CaCO3 Solutions. Science 2013, 341 (6148), 885−889. (14) Gower, L. B. Biomimetic model systems for investigating the amorphous precursor pathway and its role in biomineralization. Chem. Rev. 2008, 108 (11), 4551−627. (15) Meldrum Fiona, C.; Colfen, H. Controlling mineral morphologies and structures in biological and synthetic systems. Chem. Rev. 2008, 108 (11), 4332−432. (16) Radha, A. V.; Navrotsky, A. Thermodynamics of carbonates. Rev. Mineral. Geochem. 2013, 77, (Geochemistry of Geologic CO2 Sequestration), 73−121. (17) Donners, J. J. J. M.; Meijer, E. W.; Nolte, R. J. M.; Roman, C.; Schenning, A. P. H. J.; Sommerdijk, N. A. J. M.; Heywood, B. R. Amorphous calcium carbonate stabilized by poly(propylene imine) dendrimers. Chem. Commun. (Cambridge) 2000, No. 19, 1937−1938. (18) Lam, R. S. K.; Charnock, J. M.; Lennie, A.; Meldrum, F. C. Synthesis-dependant structural variations in amorphous calcium carbonate. CrystEngComm 2007, 9 (12), 1226−1236. (19) Michel, F. M.; MacDonald, J.; Feng, J.; Phillips, B. L.; Ehm, L.; Tarabrella, C.; Parise, J. B.; Reeder, R. J. Structural Characteristics of Synthetic Amorphous Calcium Carbonate. Chem. Mater. 2008, 20 (14), 4720−4728. (20) Koga, N.; Nakagoe, Y.; Tanaka, H. Crystallization of amorphous calcium carbonate. Thermochim. Acta 1998, 318, 239−244. (21) Addadi, L.; Raz, S.; Weiner, S. Taking advantage of disorder: Amorphous calcium carbonate and its roles in biomineralization. Adv. Mater.(Weinheim, Germany) 2003, 15 (12), 959−970. (22) Beniash, E.; Aizenberg, J.; Addadi, L.; Weiner, S. Amorphous calcium carbonate transforms into calcite during sea urchin larval spicule growth. Proc. R. Soc. London, Ser. B 1997, 264 (1380), 461− 465. (23) Weiner, S.; Dove, P. M. An overview of biomineralization processes and the problem of the vital effect. Rev. Mineral. Geochem. 2003, 54, (Biomineralization), 1−29. (24) Gong, Y. U. T.; Killian, C. E.; Olson, L. C.; Appathurai, N. R.; Amasino, A. L.; Martin, M. C.; Holt, L. J.; Wilt, F. H.; Gilbert, P. U. P. A. Phase transitions in biogenic amorphous calcium carbonate. Proc. Natl. Acad. Sci. U.S.A. 2012, 109 (16), 6088−6093. (25) Fernandez-Martinez, A.; Kalkan, B.; Clark, S. M.; Waychunas, G. A. Pressure-Induced Polyamorphism and Formation of ‘Aragonitic’ Amorphous Calcium Carbonate. Angew. Chem., Int. Ed. 2013, 52 (32), 8354−8357. (26) Politi, Y.; Metzler, R. A.; Abrecht, M.; Gilbert, B.; Wilt, F. H.; Sagi, I.; Addadi, L.; Weiner, S.; Gilbert, P. Transformation mechanism of amorphous calcium carbonate into calcite in the sea urchin larval spicule. Proc. Natl. Acad. Sci. U.S.A. 2008, 105 (45), 17362−17366. (27) Hasse, B.; Ehrenberg, H.; Marxen, J. C.; Becker, W.; Epple, M. Calcium carbonate modifications in the mineralized shell of the freshwater snail Biomphalaria glabrata. Chemistry (Weinheim, Germany) 2000, 6 (20), 3679−85. (28) Weiss, I. M.; Tuross, N.; Addadi, L.; Weiner, S. Mollusk larval shell formation: amorphous calcium carbonate is a precursor phase for aragonite. J. Exp. Zool. 2002, 293 (5), 478−491. (29) Rodriguez-Blanco, J. D.; Shaw, S.; Bots, P.; Roncal-Herrero, T.; Benning, L. G. The role of Mg in the crystallization of monohydrocalcite. Geochim. Cosmochim. Acta 2014, 127 (0), 204−220.

4. CONCLUSIONS The energetics of water adsorption (hydration energy) on material surfaces is important as it determines the stability and reactivity of the adsorbed water layer. Direct calorimetric measurement of enthalpies of water adsorption generates distinct energetic trends due to different water bonding modes on different sites. This technique made it possible to characterize the existence of distinct water interaction energetics for loosely and strongly bound water in amorphous Ca, Mg, and Mn carbonates.

■

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS The funding for this work is provided by Center of Nanoscale Control of Geologic CO2, an Energy Frontier Research Center funded by the U.S. Department of Energy, Office of Science, Office of Basic Energy Sciences under Award DE-AC0205CH11231. Preliminary X-ray scattering heating experiments on ACC and PDF data analyses were done in collaboration with Glenn A. Waychunas, Alejandro Fernandez-Martinez, and Alexis Loulier. Authors thank all collaborators for preliminary PDF data and Karena Chapman, Peter Chupas, and Kevin Beyer for their help during data collection at synchrotron X-ray scattering experiments at sector 11-ID-B, Advanced Photon Source, an Office of Science User Facility operated for the U.S. Department of Energy (DOE) Office of Science by Argonne National Laboratory.

■

REFERENCES

(1) De Yoreo, J. Crystal nucleation: More than one pathway. Nat. Mater. 2013, 12 (4), 284−285. (2) Hu, Q.; Nielsen, M. H.; Freeman, C. L.; Hamm, L. M.; Tao, J.; Lee, J. R. I.; Han, T. Y. J.; Becker, U.; Harding, J. H.; Dove, P. M.; De Yoreo, J. J. The thermodynamics of calcite nucleation at organic interfaces: Classical vs. non-classical pathways. Faraday Discuss. 2012, 159, 509−523. (3) Gebauer, D.; Kellermeier, M.; Gale, J. D.; Bergstrom, L.; Colfen, H. Pre-nucleation clusters as solute precursors in crystallisation. Chem. Soc. Rev. 2014, 43, 2348−2371. (4) Hamm, L. M.; Bourg, I. C.; Wallace, A. F.; Rotenbert, B., Molecular simulation of CO2- and CO3-brine-mineral systems. Rev. Mineral. Geochem. 2013, 77, (Geochemistry of Geologic CO2 Sequestration), 189−228. (5) Navrotsky, A. Energetic clues to pathways to biomineralization: Precursors, clusters, and nanoparticles. Proc. Natl. Acad. Sci. U.S.A. 2004, 101 (33), 12096−12101. (6) Raiteri, P.; Gale, J. D. Water is the key to nonclassical nucleation of amorphous calcium carbonate. J. Am. Chem. Soc. 2010, 132 (49), 17623−17634. (7) Tribello, G. A.; Bruneval, F.; Liew, C. C.; Parrinello, M. A Molecular Dynamics Study of the Early Stages of Calcium Carbonate Growth. J. Phys. Chem. B 2009, 113 (34), 11680−11687. (8) Pouget, E. M.; Bomans, P. H. H.; Goos, J. A. C. M.; Frederik, P. M.; de With, G.; Sommerdijk, N. A. J. M. The Initial Stages of Template-Controlled CaCO3 Formation Revealed by Cryo-TEM. Science (Washington, DC, U.S.) 2009, 323 (5920), 1455−1458. (9) Bewernitz, M. A.; Gebauer, D.; Long, J.; Colfen, H.; Gower, L. B. A metastable liquid precursor phase of calcium carbonate and its interactions with polyaspartate. Faraday Discuss. 2012, 159, 291−312. 76



Article

(30) Seto, J.; Azaïs, T.; Cölfen, H. Formation of Aragonitic Layered Structures from Kaolinite and Amorphous Calcium Carbonate Precursors. Langmuir 2013, 29 (24), 7521−7528. (31) Nebel, H.; Neumann, M.; Mayer, C.; Epple, M. On the Structure of Amorphous Calcium Carbonate: A Detailed Study by Solid-State NMR Spectroscopy. Inorg. Chem. 2008, 47 (17), 7874− 7879. (32) Guenther, C.; Becker, A.; Wolf, G.; Epple, M. In vitro synthesis and structural characterization of amorphous calcium carbonate. Z. Anorg. Allg. Chem. 2005, 631 (13−14), 2830−2835. (33) Gebauer, D.; Gunawidjaja, P. N.; Ko, J. Y. P.; Bacsik, Z.; Aziz, B.; Liu, L.; Hu, Y.; Bergström, L.; Tai, C.-W.; Sham, T.-K.; Edén, M.; Hedin, N. Proto-Calcite and Proto-Vaterite in Amorphous Calcium Carbonates. Angew. Chem., Int. Ed. 2010, 49 (47), 8889−8891. (34) Aizenberg, J.; Lambert, G.; Addadi, L.; Weiner, S. Stabilization of amorphous calcium carbonate by specialized macromolecules in biological and synthetic precipitates. Adv. Mater. (Weinheim, Germany) 1996, 8 (3), 222−6. (35) Gal, A.; Habraken, W.; Gur, D.; Fratzl, P.; Weiner, S.; Addadi, L. Calcite Crystal Growth by a Solid-State Transformation of Stabilized Amorphous Calcium Carbonate Nanospheres in a Hydrogel. Angew. Chem. Int. Ed. 2013, 52 (18), 4867−4870. (36) Wang, D.; Wallace, A. F.; De Yoreo, J. J.; Dove, P. M. Carboxylated molecules regulate magnesium content of amorphous calcium carbonates during calcification. Proc. Natl. Acad. Sci. U.S.A. 2009, 106 (51), 21511−21516. (37) Radha, A. V.; Fernandez-Martinez, A.; Hu, Y. D.; Jun, Y. S.; Waychunas, G. A.; Navrotsky, A. Energetic and structural studies of amorphous Ca1‑xMgxCO3· nH2O (0 ⩽ x ⩽ 1). Geochim. Cosmochim. Acta 2012, 90, 83−95. (38) Rieger, J.; Frechen, T.; Cox, G.; Heckmann, W.; Schmidt, C.; Thieme, J. Precursor structures in the crystallization/precipitation processes of CaCO 3 and control of particle formation by polyelectrolytes. Faraday Discuss. 2007, 136 (0), 265−277. (39) Tomono, H.; Nada, H.; Zhu, F.; Sakamoto, T.; Nishimura, T.; Kato, T. Effects of Magnesium Ions and Water Molecules on the Structure of Amorphous Calcium Carbonate: A Molecular Dynamics Study. J. Phys. Chem. B 2013, 117 (47), 14849−14856. (40) Han, N.; Blue, C. R.; De Yoreo, J. J.; Dove, P. M. The Effect of Carboxylates on the Mg Content of Calcites that Transform from ACC. Procedia Earth Planet. Sci. 2013, 7, 223−227. (41) Popescu, D. C.; van Leeuwen, E. N. M.; Rossi, N. A. A.; Holder, S. J.; Jansen, J. A.; Sommerdijk, N. A. J. M. Shaping amorphous calcium carbonate films into 2D model substrates for bone cell culture. Angew. Chem., Int. Ed. 2006, 45 (11), 1762−1767. (42) De Yoreo, J. J.; Waychunas, G. A.; Jun, Y.-S.; FernandezMartinez, A. In situ investigations of carbonate nucleation on mineral and organic surfaces. Rev. Mineral. Geochem. 2013, 77, (Geochemistry of Geologic CO2 Sequestration), 229−257. (43) Saharay, M.; James Kirkpatrick, R. Onset of Orientational Order in Amorphous Calcium Carbonate (ACC) upon Dehydration. Chem. Phys. Lett. 2014, 591 (0), 287−291. (44) Ihli, J.; Wong, W. C.; Noel, E. H.; Kim, Y.-Y.; Kulak, A. N.; Christenson, H. K.; Duer, M. J.; Meldrum, F. C. Dehydration and crystallization of amorphous calcium carbonate in solution and in air. Nat. Commun. 2014, 5. (45) Reeder, R. J.; Tang, Y.; Schmidt, M. P.; Kubista, L. M.; Cowan, D. F.; Phillips, B. L. Characterization of Structure in Biogenic Amorphous Calcium Carbonate: Pair Distribution Function and Nuclear Magnetic Resonance Studies of Lobster Gastrolith. Crys. Growth Des. 2013, 13 (5), 1905−1914. (46) Schmidt, M. P.; Ilott, A. J.; Phillips, B. L.; Reeder, R. J. Structural Changes upon Dehydration of Amorphous Calcium Carbonate. Crys. Growth Des. 2014, 14 (3), 938−951. (47) Singer, J. W.; Yazaydin, A. O.; Kirkpatrick, R. J.; Bowers, G. M. Structure and transformation of amorphous calcium carbonate: A solid-state 43Ca NMR and computational molecular dynamics investigation. Chem. Mater. 2012, 24 (10), 1828−1836.

(48) Radha, A. V.; Forbes, T. Z.; Killian, C. E.; Gilbert, P. U. P. A.; Navrotsky, A. Transformation and crystallization energetics of synthetic and biogenic amorphous calcium carbonate. Proc. Natl. Acad. Sci. U.S.A. 2010, 107 (38), 16438−16443. (49) Saharay, M.; Yazaydin, A. O.; Kirkpatrick, R. J. DehydrationInduced Amorphous Phases of Calcium Carbonate. J. Phys. Chem. B 2013, 117 (12), 3328−3336. (50) Ushakov, S. V.; Navrotsky, A. Direct measurements of water adsorption enthalpy on hafnia and zirconia. Appl. Phys. Lett. 2005, 87 (16), 164103-1−164103-3. (51) Radha, A. V.; Bomati-Miguel, O.; Ushakov, S. V.; Navrotsky, A.; Tartaj, P. Surface enthalpy, enthalpy of water adsorption, and phase stability in nanocrystalline monoclinic zirconia. J. Am. Ceram. Soc. 2008, 92 (1), 133−140. (52) Castro, R. H. R.; Quach, D. V. Analysis of Anhydrous and Hydrated Surface Energies of γ-Al2O3 by Water Adsorption Microcalorimetry. J. Phys. Chem. C 2012, 116, 24726−24733. (53) Castro, R. H. R.; Ushakov, S. V.; Gengembre, L.; Gouvea, D.; Navrotsky, A. Surface Energy and Thermodynamic Stability of γAlumina: Effect of Dopants and Water. Chem. Mater. 2006, 18, 1867− 1872. (54) Hayun, S.; Shvareva, T. Y.; Navrotsky, A. Nanoceria: energetics of surfaces, interfaces and water adsorption. J. Am. Ceram. Soc. 2011, 94, 3992−3999. (55) Levchenko, A. A.; Li, G.; Boerio-Goates, J.; Woodfield, B. F.; Navrotsky, A. TiO2 Stability Landscape: Polymorphism, Surface Energy, and Bound Water Energetics. Chem. Mater. 2006, 18, 6324−6332. (56) Sahu, S. K.; Maram, P. S.; Navrotsky, A. Thermodynamics of Nanoscale Calcium and Strontium Titanate Perovskites. J. Am. Ceram. Soc. 2013, 96, 3670−3676. (57) Tavakoli, A. H.; Maram, P. S.; Widgeon, S. J.; Rufner, J.; van Benthem, K.; Ushakov, S.; Sen, S.; Navrotsky, A. Amorphous Alumina Nanoparticles: Structure, Surface Energy, and Thermodynamic Phase Stability. J. Phys. Chem. C 2013, 117, 17123−17130. (58) Smith, D. W. Ionic hydration enthalpies. J. Chem. Educ. 1977, 54 (9), 540. (59) Rodriguez-Cruz, S. E.; Jockusch, R. A.; Williams, E. R. Hydration Energies and Structures of Alkaline Earth Metal Ions, M2+(H2O)n, n = 5−7, M = Mg, Ca, Sr, and Ba. J. Am. Chem. Soc. 1999, 121 (38), 8898−8906. (60) Peschke, M.; Blades, A. T.; Kebarle, P. Hydration Energies and Entropies for Mg2+, Ca2+, Sr2+, and Ba2+ from Gas-Phase Ion−Water Molecule Equilibria Determinations. J. Phys. Chem. A 1998, 102 (48), 9978−9985. (61) De Leeuw, N. H. Surface structures, stabilities, and growth of magnesian calcites: A computational investigation from the perspective of dolomite formation. Am. Mineral. 2002, 87, 679−689. (62) Xu, J.; Yan, C.; Zhang, F.; Konishi, H.; Xu, H.; Teng, H. H. Testing the cation-hydration effect on the crystallization of Ca−Mg− CO3 systems. Proc. Natl. Acad. Sci. U.S.A. 2013, 110 (44), 17750− 17755. (63) Forbes, T. Z.; Radha, A. V.; Navrotsky, A. The energetics of nanophase calcite. Geochim. Cosmochim. Acta 2011, 75 (24), 7893− 7905. (64) Kerisit, S.; Parker, S. C. Free Energy of Adsorption of Water and Metal Ions on the {1014} Calcite Surface. J. Am. Chem. Soc. 2004, 126 (32), 10152−10161. (65) Rahaman, A.; Grassian, V. H.; Margulis, C. J. Dynamics of Water Adsorption onto a Calcite Surface as a Function of Relative Humidity. J. Phys. Chem. C 2008, 112 (6), 2109−2115. (66) Radha, A. V.; Navrotsky, A. Manganese carbonate formation from amorphous and nanocrystalline precursors: Thermodynamics and geochemical relevance. Am. Mineral. 2014, 99 (5−6), 1063−1070. (67) Schaef, H. T.; McGrail, B. P.; Owen, A. T. Basalt-CO2-H2O interactions and variability in carbonate mineralization rates. Energy Procedia 2009, 1 (1), 4899−4906. 77



Article

(68) Bano, A. M.; Rodger, P. M.; Quigley, D. New Insight into the Stability of CaCO3 Surfaces and Nanoparticles via Molecular Simulation. Langmuir 2014, 30 (25), 7513−7521. (69) Sel, O.; Radha, A. V.; Dideriksen, K.; Navrotsky, A. Amorphous iron (II) carbonate: Crystallization energetics and comparison to other carbonate minerals related to CO2 sequestration. Geochim. Cosmochim. Acta 2012, 87, 61−68.

78


Direct Experimental Measurement of Water Interaction Energetics in

Recommend Documents