# Role of Na2CO3 as Nucleation Seeds to Accelerate the CO2 Uptake Kinetics of MgO-Based Sorbents

Annelies Landuyt, Ilia Kochetygov, Charles J. Mcmonagle, Priyank V. Kumar, Jodie A. Yuwono, Wendy L. Queen, Paula M. Abdala, Christoph R. Müller

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Annelies Landuyt, Ilia Kochetygov, Charles J. Mcmonagle, Priyank V. Kumar, Jodie A. Yuwono, et al.. Role of Na2CO3 as Nucleation Seeds to Accelerate the CO2 Uptake Kinetics of MgO-Based Sorbents. JACS Au, 2024, 4 (12), pp.4809-4820. 10.1021/jacsau.4c00782. http://dx.doi.org/10.1021/jacsau.4c00782

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# Role of Na₂CO₃ as Nucleation Seeds to Accelerate the CO₂ Uptake Kinetics of MgO-Based Sorbents

Annelies Landuyt, Ilia Kochetygov, Charles J. McMonagle, Priyank V. Kumar, Jodie A. Yuwono, Wendy L. Queen, Paula M. Abdala,* and Christoph R. Müller*

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Cite This: JACS Au 2024, 4, 4809–4820

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**ABSTRACT:** There is an urgent need for inexpensive, functional materials that can capture and release CO₂ under industrial conditions. In this context, MgO is a highly promising, earth-abundant CO₂ sorbent. However, despite its favorable carbonation thermodynamics and potential for high gravimetric CO₂ uptakes, MgO-based CO₂ sorbents feature slow carbonation kinetics, limiting their CO₂ uptake during typical industrial contact times. The addition of molten alkali metal nitrate promoters, such as NaNO₃, can partially mitigate the slow kinetics. Here, we investigate how the CO₂ uptake kinetics of NaNO₃-promoted MgO can be increased further through the addition of finely dispersed Na₂CO₃. The incorporation of Na₂CO₃ significantly increases the CO₂ uptake rate from 1.4 to 14.6 mmol MgCO₃ (mol MgO)⁻¹ s⁻¹.

Using *in situ* synchrotron X-ray powder diffraction (XRD), we track the formation of MgCO₃ and elucidate the mechanism through which Na₂CO₃ promotes the CO₂ uptake of MgO. Our findings demonstrate that Na₂CO₃ rapidly converts within seconds into Na₂Mg(CO₃)₂ during carbonation, acting subsequently as nucleation seeds for MgCO₃ formation, in turn significantly enhancing CO₂ uptake kinetics. Further, the presence of Na₂Mg(CO₃)₂ considerably enhances the mobility of ions in the sorbent, leading to sintering of MgCO₃. Importantly, Na₂Mg(CO₃)₂ promotes MgCO₃ formation even in the presence of molten RbNO₃, a salt with a limited ability to dissolve [Mg²⁺...CO₃²⁻] ion pairs, indicating that Na₂Mg(CO₃)₂ lowers the critical ion pair concentration required for MgCO₃ nucleation. Additionally, the partial dissolution of Na₂CO₃ in NaNO₃ may increase the concentration of carbonate ions in the melt, further accelerating carbonation kinetics in MgO-(Na₂CO₃/NaNO₃).

**KEYWORDS:** CO₂ capture, MgO, carbonate, X-ray diffraction, molten salts, nucleation seed, *in situ*

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## INTRODUCTION

Carbon dioxide capture is a critical technology to reach the global CO₂-emission targets and to keep global warming within a tolerable range, in particular in the hard-to-abate sectors.¹⁻⁴ Currently, the benchmark CO₂ capture technology is amine scrubbing.⁵,⁶ While amine scrubbing has been applied on the large scale for the removal of CO₂ from natural gas, it suffers from severe limitations, including the degradation of amines into products which are toxic and/or harmful greenhouse gases as well as the relatively high energy demand for their regeneration (aqueous amine solution).⁷,⁸ Solid oxide sorbents such as CaO or MgO are a promising family of materials for CO₂ capture due to their favorable thermodynamics, high theoretical CO₂ uptake capacities, environmental benignity and earth-abundance.⁹⁻¹¹ MgO-based CO₂ sorbents have attracted particular interest because of their intermediate operation temperature (typically 250–450 °C), while CaO requires higher operating temperatures (typically 600–900 °C). The lower operating temperature of MgO can translate into reduced energy requirements (for sorbent regeneration)

and may mitigate sintering effects.¹²,¹³ Further, MgO-based CO₂ sorbents have shown promise for precombustion CO₂ capture and sorption-enhanced water–gas shift reactions (SEWGS) due to their favorable, intermediate operating temperatures.¹⁴,¹⁵ Specifically, MgO reacts reversibly with CO₂ forming MgCO₃:

$$ \begin{aligned} & \mathrm{MgO}(\mathrm{s}) + \mathrm{CO}_{2}(\mathrm{g}) \\ & \rightleftharpoons \mathrm{MgCO}_{3}(\mathrm{s})\Delta H_{298 \mathrm{~K}} \\ & = -116.9 \mathrm{~kJ} \mathrm{~mol}^{-1} \end{aligned} \qquad (1) $$

Received: August 26, 2024

Revised: October 17, 2024

Accepted: November 5, 2024

Published: November 16, 2024

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While the thermodynamics for this reaction are favorable, the slow kinetics for CO₂ uptake hamper the commercial usage of MgO at present. Nonetheless, recent research has shown that the slow kinetics of CO₂ uptake can be alleviated to some degree by promoting MgO with alkali metal nitrates (such as NaNO₃, LiNO₃ or mixtures of Li-, Na- and K-nitrates), yet only if the promoters are molten under operation conditions (300–450 °C).¹⁶⁻¹⁸ Although the exact mechanism through which molten alkali metal nitrates enhance the CO₂ uptake rate of MgO is still being investigated, it is hypothesized that the molten alkali metal nitrates facilitate MgCO₃ formation by dissolving the reactants. However, the nature and chemical environment of the dissolved species remain highly debated.¹⁷⁻²⁰ Using ¹⁸O labeling and ab initio molecular dynamics (AIMD) calculations, a recent study revealed that when a MgO surface is exposed to a CO₂ atmosphere, surface carbonates are readily formed. Subsequently, these surface carbonates dissolve in molten NaNO₃ as [Mg²⁺...CO₃²⁻] ionic pairs, which subsequently precipitate as crystalline MgCO₃.²⁰

It has been observed also that besides molten alkali metal nitrates, Na₂CO₃ and other alkali metal carbonates increase the rate of CO₂ uptake of MgO.²¹⁻²⁴ Since alkali metal carbonates exhibit high melting points (for instance, Na₂CO₃ melts at 851 °C), they are in a solid state under typical operating conditions. It has been argued that alkali metal carbonates promote the CO₂ uptake of MgO by forming a double carbonate, a reaction with reasonable kinetics also in the absence of NaNO₃.²⁵,²⁶ For example, Na₂CO₃ reacts with MgO and CO₂ forming Na₂Mg(CO₃)₂ according to

$$ \begin{aligned} & \mathrm{Na_2CO_3(s) + CO_2(g) + MgO(s)} \\ & \rightleftharpoons \mathrm{Na_2Mg(CO_3)_2(s)\Delta H_{298 K}} \\ & = -130.4 \, \mathrm{kJ \, mol^{-1}} \end{aligned} \qquad (2) $$

Generally, the promotion of MgO with alkali metal carbonates results in lower CO₂ uptakes (0.04–0.15 g gₐsorbent⁻¹) compared to its promotion with molten alkali metal nitrates (CO₂ uptakes in the range of 0.14 to 0.6 g gₐsorbent⁻¹), because their addition only leads to the formation of the respective double carbonate (e.g., Na₂Mg(CO₃)₂) and not to the formation of MgCO₃ itself.¹⁶⁻¹⁸,²⁶⁻²⁸ However, when adding a combination of alkali metal nitrates and carbonates to MgO, sorbents with both rapid CO₂ uptake kinetics and high CO₂ capacities are obtained.²⁷,²⁹⁻³³ Indeed, currently the best-performing MgO-based CO₂ sorbent is MgO promoted by a mixture of (Li, K)NO₃ and (Na, K)₂CO₃ with a reported CO₂ uptake of 0.84 g gₐsorbent⁻¹ (for a carbonation time of 4 h) and a fairly stable cyclic performance, i.e., a decrease in the CO₂ uptake by 18% over 30 cycles.²⁹ Although the results of this combined promotion effort are very encouraging, we currently have only a very limited understanding of the mechanisms through which alkali metal carbonates improve the kinetics of alkali metal nitrate-promoted MgO. However, such knowledge is critical to further advance MgO-based CO₂ sorbents and to reduce the amount of molten nitrate promoters, which may induce corrosion in the reactor vessels used.

Alkali metal carbonates can dissolve in molten nitrate salts, which has led to the hypothesis that a dissolution effect enhances the formation of double carbonates through precipitation in the melt.²⁷,³⁴,³⁵ It was also hypothesized that the Na₂Mg(CO₃)₂ formed acts as a carrier to transfer CO₂ to MgO, thereby accelerating MgCO₃ formation.³⁶ A recent

alternative mechanism suggests that Na₂Mg(CO₃)₂ acts as a nucleation seed for MgCO₃ formation owing to the structural similarity between the two carbonates structures (with R̄c and R̄H space groups respectively).³⁷ Both carbonates are described with hexagonal (or rhombohedral) lattices whose unit cells have close dimensions (MgCO₃ has lattice parameters a = 4.637 Å and c = 15.023 Å; and Na₂Mg(CO₃)₂ has lattice parameters of a = 4.946 Å and c = 16.422 Å). This structural similarity could lower the energy barrier for MgCO₃ nucleation, resulting in faster kinetics,³⁸⁻⁴⁰ and may also explain the promoting effect of other carbonates, such as CaMg(CO₃)₂.⁴¹,⁴² Despite the observation of an improved CO₂ uptake performance of such carbonate and nitrate-promoted sorbents, the underlying mechanism of the promotional effect remains elusive. Hence, gaining insight into the structural changes of sorbents under CO₂ capture conditions is crucial for developing a detailed mechanistic understanding of the promotional effects. For example, Rekhtina et al., using in situ XRD and Raman spectroscopy on a MgO-NaNO₃ sorbent, demonstrated that Na₂Mg(CO₃)₂ forms through the partial decomposition of NaNO₃ during regeneration at temperatures ≥450 °C over repeated cycles.³⁷ The in situ formation of Na₂Mg(CO₃)₂ was linked to an accelerated rate of MgCO₃ formation, leading to the hypothesis that Na₂Mg(CO₃)₂ crystallites act as seeds for MgCO₃ nucleation and growth. However, additional mechanistic studies of the targeted addition of carbonates (such as Na₂CO₃) are required to fully elucidate the promotional effect of combining nitrate with carbonate promoters and to improve further the performance of MgO-based sorbents.

Hence, in this work we aim to shed light on the mechanism behind the cooperative promotional effect of Na₂CO₃ and NaNO₃ on the CO₂ uptake of MgO-based sorbents. We developed a model MgO-based CO₂ sorbent that is promoted with both Na₂CO₃ and NaNO₃, and tracked the evolution of the relevant phases (i.e., their amount, average crystallite sizes and lattice parameters) under CO₂ capture conditions. This was achieved using time-resolved (1 s) in situ X-ray powder diffraction (XRD) with sufficient reciprocal space resolution to monitor the temporal evolution of the average crystallite sizes. Our experimental data revealed that Na₂CO₃ fully converts into Na₂Mg(CO₃)₂ in just a few seconds, prior to the formation of crystalline MgCO₃ and provides experimental evidence that Na₂Mg(CO₃)₂ acts as a nucleation seed for MgCO₃, significantly reducing the high energy barrier associated with MgCO₃ nucleation. This, in turn, leads to a drastic increase in the nucleation and growth kinetics of MgCO₃. Finally, the discovered promotion mechanism was validated further using MgO-RbNO₃ as a benchmark for a poorly performing sorbent, demonstrating that the addition of Na₂Mg(CO₃)₂ seeds promotes MgCO₃ nucleation even in the presence of molten RbNO₃, which on its own is insufficient to promote the CO₂ uptake of MgO. These observations suggest that the role of Na₂Mg(CO₃)₂ is to act as a nucleation seed, lower the critical [Mg²⁺...CO₃²⁻] ion pair concentration in the melt required for MgCO₃ formation and to increase the availability of [Mg²⁺...CO₃²⁻] ions and their mobility in the sorbent via Na₂Mg(CO₃)₂ dissolution in the nitrate melt.
<figure><img src="image_9.png" /><figcaption>**Figure 1.** Time-resolved XRD data acquired during the carbonation of (a–d) MgO-NaNO₃ and (e–h) MgO-(Na₂CO₃/NaNO₃) at 315 °C in CO₂. (a) 2D contour plot showing the evolution of crystalline phases during the first 90 s (left) and the remaining 30 min (right) for MgO-NaNO₃ under carbonation conditions. (b) The molar fraction of MgCO₃ (which is equal to MgO conversion) and (c) the average crystallite sizes of MgCO₃ and MgO (⟨DₘgCO₃⟩ and ⟨DₘgO⟩ respectively) as a function of time, determined from the XRD data. (d) Schematic representation of the sorbent MgO-NaNO₃ during the different reaction stages. (e) 2D contour plot showing the evolution of the XRD peaks during the first 90 s (left) and the remaining 30 min (right) for MgO-(Na₂CO₃/NaNO₃) under carbonation conditions. (f) The molar fractions of MgCO₃ and Na₂Mg(CO₃)₂ and the total MgO conversion (sum of the molar fractions of MgCO₃ and Na₂Mg(CO₃)₂) and (g) the average crystallite sizes of MgCO₃, Na₂Mg(CO₃)₂ and MgO as a function of time. (h) Schematic representation of the evolution of the sorbent MgO-(Na₂CO₃/NaNO₃) during the different reaction stages. Error bars are visualized, when not visible they are smaller than the size of the symbols.</figcaption></figure>
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## RESULTS AND DISCUSSION

### Phase and Average Crystallite Size Evolution during MgO Carbonation in MgO-(Na₂CO₃/NaNO₃) and MgO-NaNO₃

To provide mechanistic insight into how the addition of Na₂CO₃ affects the CO₂ uptake of NaNO₃-promoted MgO, we employed synchrotron-based, in situ, time-resolved XRD measurements (1 s time resolution) during carbonation (Figures 1 and S5). The experiments were carried out using a plug-flow capillary quartz reactor with an outer diameter of 1 mm. The high time resolution is required to resolve the rapid structural changes (phase composition, average crystallite size and lattice strain) that occur during the carbonation of the sorbents.³⁷,⁴³⁻⁴⁵ Na₂CO₃/NaNO₃-promoted MgO (referred to as MgO-(Na₂CO₃/NaNO₃)) was prepared by ball milling, resulting in Na₂CO₃ being highly dispersed within the sorbent, as confirmed by transmission electron microscopy (TEM) and XRD analysis (Figures S1 and S2). Based on a set of experiments in which the fraction of Na₂CO₃ in the sorbent was varied to optimize the CO₂ uptake and to allow for a sufficient signal intensity for XRD characterization, the optimal amount of Na₂CO₃ was determined as 5 mol % (Figure S4). NaNO₃-promoted MgO (referred to as MgO-NaNO₃) was used as a reference material and prepared via the same route, yet without the addition of Na₂CO₃ (Figures S2 and S3).

The in situ XRD data reveals that in MgO-NaNO₃ (Figure 1a), reflections due to MgCO₃ appear after an induction period of ca. 3 min. Rietveld analysis of the in situ XRD patterns (Figure S6) allows to extract the evolution of the phase composition and the average crystallite size (⟨D⟩) of each phase (Figure 1b,c) providing valuable insight into the prevailing crystallization process(es) during carbonation. For the following discussion it is important to note that when performing a line profile analysis of XRD data, we calculate a volume-weighted (apparent) average size of the crystallites (apparent average coherent domain size) without accounting for their size distribution.⁴⁶ Changes in ⟨D⟩ during reaction could be determined with reasonable accuracy for ⟨D⟩ less than approximately 100 nm. The estimated standard deviation (ESD) for each refined parameter (as a function of time) is provided in the plot (Figure 1). The ESD is <1 nm for ⟨D⟩ less than 50 nm and <6 nm for ⟨D⟩ less than 100 nm. For ⟨D⟩ larger than 150 nm, we approach the instrumental resolution leading to higher uncertainties (ESD ca. 20 nm).

Previous works have proposed that the carbonation of NaNO₃-promoted MgO involves the following reaction stages: surface carbonate formation, nucleation and growth of MgCO₃ crystallites described by an Avrami–Erofeev reaction model followed by a slow carbonation stage described by a Ginstling–Brounshtein product layer diffusion process.¹⁶,¹⁷,⁴⁷,⁴⁸ We note that the phase evolution analysis (Figure 1b) performed by XRD cannot distinguish surface carbonate formation and the nucleation period, as these steps do not involve crystalline phases. Rietveld analysis of the in situ XRD data allowed us to divide the carbonation of NaNO₃-promoted MgO into three stages. We define the induction period (stage I) as the period in which there is no presence of crystalline MgCO₃. Under the conditions applied here, the induction period for MgO-NaNO₃ lasts around 3 min and is followed by (II) MgCO₃ nucleation and growth (3 min < t < 11 min) and (III) a further, slow growth of MgCO₃, controlled by the diffusion of reactants (CO₂ and/or carbonate ions) through the product layer (t > 11 min).

During carbonation of NaNO₃-promoted MgO, the average crystallite size of MgO remains constant at ~17 nm. In the first 2 min of stage II (labeled stage II.1), i.e., the period when the first MgCO₃ crystallites appear and MgO reaches 6% conversion, the average crystallite size of MgCO₃ increases rapidly to ~23 nm and it remains constant at ~23 nm in stage II.2 (5 min < t < 11 min), during which a further 39% of MgO is converted into MgCO₃. The constant average crystallite sizes of both MgO and MgCO₃ over time indicate that the individual MgO crystallites are fully and rapidly converted into MgCO₃, as opposed to the simultaneous partial conversion of many MgO crystallites, which would result in a decrease in the average size of MgO. The average crystallite size of MgCO₃ (⟨DₘgCO₃⟩ = 23 nm) is directly linked to the initial size of the MgO crystallites (⟨DₘgO⟩ = 17 nm) when considering the expansion factor of the conversion of MgO into MgCO₃ (⟨DₘgCO₃⟩ = 1.35 · ⟨DₘgO⟩, see also Figure S8). The observation that the average crystallite size of MgCO₃ does not increase with MgO conversion indicates that there is no appreciable sintering in this stage. In stage III (11 min < t < 30 min), the rate of MgCO₃ formation gradually slows down, with an additional 9% of MgO being converted into MgCO₃ (overall MgO conversion of 53%). The increase in the average crystallite size of MgCO₃ in Stage III can primarily be attributed to a continuous liquid phase sintering process, as described previously for similar systems.³³,⁴²,⁴⁹ Indeed, during stage III, the amount of MgCO₃ increases by only a factor of 1.2 (from 44 to 53%), which would theoretically result in an increase in the average crystallite size of MgCO₃ by a factor of ³√1.2 (~1.06) from 23 to 24.5 nm. However, we observe an actual increase in the average MgCO₃ crystallite size from 23 to 31 nm. A schematic representation of the structural and morphological evolution of NaNO₃-promoted MgO during the different reaction stages during carbonation is provided in Figure 1d.

Turning now to the in situ XRD data of MgO-(Na₂CO₃/NaNO₃) (Figure 1e–h and Figure S5 and S7), we observe a marked difference in the phase and average crystallite size evolution when compared to MgO-NaNO₃. In MgO-(Na₂CO₃/NaNO₃) crystalline MgCO₃ is already present after 19 s, indicating a significantly shorter induction period (as compared to ca. 3 min in MgO-NaNO₃) which is in line with thermogravimetric experiments (Figure S9). Based on the extracted phase and average crystallite size information, we can divide the carbonation of MgO-(Na₂CO₃/NaNO₃) into the following three reaction stages: (I) Na₂Mg(CO₃)₂ formation (t < 19 s), (II) MgCO₃ nucleation and growth (19 s < t < 90 s) and (III) a further, slow growth of MgCO₃, controlled by diffusion of reactants through the product layer (t > 90 s). In stage I, Na₂CO₃ and MgO rapidly react with CO₂ to form Na₂Mg(CO₃)₂, resulting in 4.5 mol % crystalline Na₂Mg(CO₃)₂ after ~19 s. This corresponds to a stoichiometric conversion, based on the Na₂CO₃:MgO ratio of 4.5% in the as-prepared sorbent, as determined by inductively coupled plasma optical emission spectroscopy (ICP-OES). Note that Na₂CO₃ was not included in the Rietveld refinement, as the peaks due to Na₂CO₃ were of very weak intensity at the onset of the carbonation reaction (Na₂CO₃ may be partially amorphous or dissolved within NaNO₃), and these peaks disappear completely by the end of stage I. Additionally, Na₂CO₃ exists in two phases, which complicates further its Rietveld analysis (see Figure S10 for more details).
<figure><img src="image_11.png" /><figcaption>**Figure 2.** Evolution of (a) *a*-parameter and (b) *c*-parameter of MgCO₃ during the early stages of the carbonation of MgO-NaNO₃ (red dots) and MgO-(Na₂CO₃/NaNO₃) (black dots) as a function of MgCO₃ conversion. The respective equilibrium values (values obtained after 60 min of carbonation) of the lattice parameters of MgO, MgCO₃ and Na₂Mg(CO₃)₂ are included in the plot (gray dashed lines). The error bars are smaller than the symbols. (c) Visualization of the unit cells of MgO (Fm$\overline{3}$m), MgCO₃ (R$\overline{3}$c) and Na₂Mg(CO₃)₂ (R$\overline{3}$H) including labeling of the respective lattice parameters.</figcaption></figure>

At t = 19 s, the stage in which MgCO₃ rapidly nucleates and grows (stage II) begins. Initially, in stage II.1 (19 s < t < 25 s), small MgCO₃ crystallites of ca. 11 nm in size form rapidly, converting 3% of MgO into MgCO₃. Subsequently, in stage II.2 (25 s < t < 90 s), the MgCO₃ crystallites rapidly grow by converting an additional 35% of MgO into MgCO₃. The size of the MgCO₃ crystallites at the end of stage II.2 is ~100 nm. When comparing the dynamics of MgO conversion and MgCO₃ crystallite size evolution in stage II.2, it is evident that the increase in the size of the MgCO₃ crystallites is not only due to an increasing MgO conversion (i.e., increasing CO₂ uptake) but also due to sintering. During stage II.2, the amount of MgCO₃ increased by a factor of 13 (from 3 to 38%), which would result in an increase in the average size of the MgCO₃ crystallites by a factor of $\sqrt[3]{13}$ (~2.35). However, the observed increase in the average size of the MgCO₃ crystallites is significantly higher (by a factor of 9), i.e., a clear sign of significant sintering (Figure S11).

In stage III ($t > 90$ s) of the carbonation of MgO-(Na₂CO₃/NaNO₃), the rate of MgCO₃ formation is significantly slower, with 10% of MgO being converted into MgCO₃ in 28.5 min (as compared to 43% MgO conversion in the first 90 s of the reaction), resulting in a total MgO conversion of 53% after 30

min of carbonation. In stage III, the average MgCO₃ crystallite size increases from 100 to 160 nm. Again, the increase in the size of the MgCO₃ crystallites cannot be explained solely by the increasing amount of MgCO₃ formed but must also be due to liquid phase sintering. Compared to MgCO₃, the sintering of Na₂Mg(CO₃)₂ is even more pronounced with its average crystallite size increasing from 36 to 130 nm during stages II and III (sintering in the absence of any additional formation of Na₂Mg(CO₃)₂). The pronounced sintering of Na₂Mg(CO₃)₂ is likely explained by the high mobility of ions, possibly linked to Na₂Mg(CO₃)₂ dissolution in molten NaNO₃.³³,⁴⁹ Nevertheless, scanning transmission electron microscopy with energy-dispersive x-ray (STEM-EDX) maps of carbonated MgO-(Na₂CO₃/NaNO₃) (Figure S13) reveal that while a fraction of Na₂Mg(CO₃)₂ is in a highly agglomerated/sintered state, there is still well-distributed Na₂Mg(CO₃)₂ within the sorbent. Unfortunately, it was not feasible to probe the reaction mechanism using STEM-EDX analysis due to several limitations. First, the overlapping signals of different phases (e.g., NaNO₃ and Na₂Mg(CO₃)₂ as well as MgO and MgCO₃) made it difficult to distinguish between them. Additionally, STEM-EDX does not provide detailed structural information and certain components of the sorbent exist in different phases
at room temperature compared to reaction conditions. For example, while $NaNO_3$ is liquid under reaction conditions, it crystallizes into large agglomerates at room temperature.

Notably, after 30 min of carbonation the average crystallite size of $MgCO_3$ is larger in $MgO-(Na_2CO_3/NaNO_3)$ compared to $MgO-NaNO_3$ (160 nm versus 31 nm), although the conversion of MgO is similar in both materials. Most of the sintering in $MgO-(Na_2CO_3/NaNO_3)$ takes place during the formation of $MgCO_3$ in reaction stage II, resulting in $MgCO_3$ crystals with an average crystallite size of 100 nm. We propose that the dissolution of $Na_2CO_3/Na_2Mg(CO_3)_2$ in $NaNO_3$ facilitates the diffusion/mobility of cationic and carbonate ions ($Na^+$, $Mg^{2+}$, $CO_3^{2-}$) in the melt leading to an accelerated $MgCO_3$ growth and liquid phase sintering.<sup>35,50</sup>

To assess the impact of sintering of $MgCO_3$ and $Na_2Mg(CO_3)_2$ in $MgO-(Na_2CO_3/NaNO_3)$ on its cyclic performance, we recorded the cyclic $CO_2$ uptake for $MgO-(Na_2CO_3/NaNO_3)$ over 20 cycles of $CO_2$ uptake and regeneration using a TGA (Figure S15). The data show that the $CO_2$ uptake of $MgO-(Na_2CO_3/NaNO_3)$ decreases by 46% after 20 cycles. Nevertheless, the positive effect of the addition of $Na_2CO_3$ on the $CO_2$ uptake kinetics of $NaNO_3$-promoted MgO remains evident even after 20 cycles, as reflected by the faster $CO_2$ uptake kinetics and overall higher $CO_2$ uptake of $MgO-(Na_2CO_3/NaNO_3)$ compared to $MgO-NaNO_3$.

## Evolution of $MgCO_3$ Lattice Parameters and Its Implications for the Nucleation and Growth Mechanism

The observation that the nucleation and growth of $MgCO_3$ proceeds much faster in $MgO-(Na_2CO_3/NaNO_3)$ than in $MgO-NaNO_3$ could be explained by two factors: In $MgO-(Na_2CO_3/NaNO_3)$, there is (i) a higher solubility of $[Mg^{2+} \cdots CO_3^{2-}]$ ion pairs in molten $NaNO_3$ (it has been reported that $Na_2CO_3$ has a relatively high solubility in $NaNO_3$ (3.5 mol % at 315 °C and 9 mol % at 450 °C) which could lead to a higher concentration of dissolved carbonate ions in the nitrate melt, yielding in turn higher $MgCO_3$ formation rates), (ii) a higher density of nucleation points for $MgCO_3$ formation, or (iii) a combination of both. Our XRD measurements show that $Na_2Mg(CO_3)_2$ formation precedes the crystallization of $MgCO_3$, making it likely to act as a nucleation seed. Considering that $MgCO_3$ formation is a heterogeneous nucleation process, the availability and density of nucleation points can play a key role in accelerating the $MgCO_3$ formation process. Owing to the structural similarity between $Na_2Mg(CO_3)_2$ and $MgCO_3$, both crystallize in closely related rhombohedral structures (Figure 2c), $Na_2Mg(CO_3)_2$ crystallites may act as nucleation seeds by lowering the energy barrier for $MgCO_3$ nucleation, resulting in faster nucleation kinetics.<sup>37–40</sup>

To gain insight into the lattice strain of $MgCO_3$ formed in the two sorbents, we compare the evolution of the $MgCO_3$ lattice parameters obtained from Rietveld analysis of the *in situ* XRD data during the carbonation of $MgO-NaNO_3$ and $MgO-(Na_2CO_3/NaNO_3)$. At the early stages of crystal growth, $MgCO_3$ formed on a substrate can be strained due to a lattice mismatch with the substrate. Figure 2 plots the lattice parameters of $MgCO_3$ (*a*-axis and *c*-axis) for $MgO-(Na_2CO_3/NaNO_3)$ and $MgO-NaNO_3$ as a function of $MgCO_3$ conversion. Note that the XRD analysis yields the average strain over all crystallites formed and not an individual crystallite lattice strain.

In the initial stages of $MgCO_3$ formation, both parameters a and c differ from the values reached at high degrees of $MgCO_3$ formation, confirming that $MgCO_3$ is initially in a strained state that relaxes as the amount of $MgCO_3$ increases. For $MgO-(Na_2CO_3/NaNO_3)$, the a-parameter of $MgCO_3$ ($a$ = 4.68063(6) Å) is initially larger than its final value, while for $MgO-NaNO_3$, the a-parameter of $MgCO_3$ ($a$ = 4.59852(6) Å) is initially lower than its final value (Figure 2a). In both materials, the a-parameter relaxes with increasing conversion and the final value is ~4.65 Å (4.65207(3) Å for $MgO-(Na_2CO_3/NaNO_3)$ and 4.65391(3) Å for $MgO-NaNO_3$, see Figure S16).

Turning to the c-parameter (Figure 2b), we observe that the initial value of the c-parameter for $MgO-(Na_2CO_3/NaNO_3)$ (15.631(7) Å) is larger than for $MgO-NaNO_3$ (15.433(1) Å), yet in both materials the c-parameter shows a similar trend with $MgCO_3$ formation reaching ultimately an equilibrium value of $c$ ~ 15.17 Å (15.1700(3) Å for MgO $MgO-(Na_2CO_3/NaNO_3)$ and 15.1725(3) Å for $MgO-NaNO_3$). As cell parameters may also depend on the size of the crystallites, we also plot the cell parameters as a function of the average crystallite size (Figure S17). Yet we cannot observe a clear correlation between the cell parameters and the average size of the $MgCO_3$ crystallites.<sup>51,52</sup> Instead, we hypothesize that the difference in the cell parameters of $MgCO_3$ at low MgO conversions between the two sorbents arises from the different substrates onto which $MgCO_3$ crystallites grow. As the initial a-parameter of $MgCO_3$ in $MgO-NaNO_3$ and $MgO-(Na_2CO_3/NaNO_3)$ is similar to the a-parameter of MgO and $Na_2Mg(CO_3)_2$, respectively, we argue that in $MgO-NaNO_3$, $MgCO_3$ nucleates onto MgO, while in $MgO-(Na_2CO_3/NaNO_3)$, $MgCO_3$ additionally nucleates onto $Na_2Mg(CO_3)_2$. Importantly, the lattice parameter of MgO does not change during carbonation and shows similar values for both materials ($a_{(MgO)}$ is 4.23101(3) Å and 4.23331(3) Å for $MgO-NaNO_3$ and $MgO-(Na_2CO_3/NaNO_3)$ respectively), indicating that the observed differences in the lattice parameters of $MgCO_3$ in the two sorbents studied cannot be explained by temperature or instrumental effects (Figure S18).

An analysis of the $MgCO_3$ lattice parameters suggests that $MgCO_3$ can nucleate on MgO in $MgO-NaNO_3$ and on $Na_2Mg(CO_3)_2$ in $MgO-(Na_2CO_3/NaNO_3)$. The lattice mismatch between the substrate and the nucleus directly influences the energy barrier for nucleation.<sup>53,54</sup> For example, for the $MgCO_3-MgO$ interface, the lattice mismatch $\epsilon_a$ is 9.1% ($\epsilon_a$ = $(a_{MgCO_3} - a_{MgO}) \times a_{MgCO_3}^{-1} \times 100$). For $MgCO_3-Na_2Mg(CO_3)_2$, $\epsilon_a$ is 6.6% ($\epsilon_a$ = $(a_{MgCO_3} - a_{Na_2Mg(CO_3)_2}) \times a_{MgCO_3}^{-1} \times 100$) and $\epsilon_c$ = 9.1 $(c_{MgCO_3} - a_{Na_2Mg(CO_3)_2}) \times c_{MgCO_3}^{-1} \times 100$. The lower lattice mismatch for the $MgCO_3-Na_2Mg(CO_3)_2$ interfaces, in combination with the structural similarity between the two carbonates, could make $Na_2Mg(CO_3)_2$ a more effective nucleation seed compared to MgO. However, further computational studies are required to rationalize in more detail the experimental findings obtained here regarding energy barriers for $MgCO_3$ nucleation. Furthermore, the lattice parameter analysis does not rule out the potential contribution of small amounts of dissolved $Na_2CO_3$ in molten $NaNO_3$ in increasing the kinetics of $MgCO_3$ formation in $MgO-(Na_2CO_3/NaNO_3)$, e.g., by facilitating the dissolution (and potentially also increasing the concentration) of $[Mg^{2+} \cdots CO_3^{2-}]$ ion pairs.
<figure><img src="image_12.png" /><figcaption>**Figure 3.** *In situ* synchrotron XRD data collected during the first 12 min of carbonation for MgO-(Na₂CO₃/RbNO₃): (a) 2D contour plot of the acquired XRD patterns and (b) the molar fractions of MgCO₃ (xNa₂CO₃) and Na₂Mg(CO₃)₂ (xNa₂Mg(CO₃)₂) as well as the total MgO conversion (xMgCO₃ + xNa₂Mg(CO₃)₂). (c) MgCO₃ formation rate (first derivative of the MgCO₃ molar fraction as obtained from the *in situ* XRD data) during the first 12 min of the carbonation reaction for MgO-(Na₂CO₃/NaNO₃) (green), MgO-NaNO₃ (red) and MgO-(Na₂CO₃/RbNO₃) (black). (d) 2D contour plot of the *in situ* XRD patterns (synchrotron) collected during the first 12 min of carbonation of MgO-Na₂CO₃.</figcaption></figure>

## Probing the Mechanism through Which Na₂CO₃ Promotes the Carbonation of MgO in the Presence of RbNO₃

To understand better how Na₂CO₃/Na₂Mg(CO₃)₂ promotes the formation of MgCO₃, we examined the effect of the nature of the molten nitrate on the CO₂ uptake and the dynamics of the phase evolution. We chose RbNO₃ as it was found not to enhance the CO₂ absorption of MgO, despite having a melting point that is very similar to that of NaNO₃, (i.e., Tₘelting(RbNO₃) = 310 °C and Tₘelting(NaNO₃) = 308 °C).⁵,⁵ Interestingly, while MgO-RbNO₃ shows no appreciable CO₂ uptake, MgO-(Na₂CO₃/RbNO₃) exhibits a high CO₂ uptake at a relatively fast rate, as evidenced by thermogravimetric analysis (TGA) under carbonation conditions at 315 °C (see Figure S19).

To probe the structural evolution of MgO-(Na₂CO₃/RbNO₃) under CO₂ uptake conditions, synchrotron-based *in situ* XRD experiments were conducted (Figures 3a and S20). As the presence of Rb led to a fluorescence-induced, complex background in the XRD data and reduced the signal-to-noise ratio, the level of detail that could be extracted from the *in situ* XRD data was lower when compared to the Rb-free samples. Nonetheless, despite these limitations, we were able to extract key information on the phase evolution during the initial stages of MgCO₃ formation. Figure 3a plots the *in situ* XRD data during the first 12 min of carbonation and Figure S21 shows the phase quantification at selected points during carbonation. Similar to MgO-(Na₂CO₃/NaNO₃), we observe also for MgO-

(Na₂CO₃/RbNO₃) a rapid conversion of Na₂CO₃ into Na₂Mg(CO₃)₂ during the first 20 s of carbonation (stage I), followed (after ~1.3 min.) by the appearance of Bragg reflections due to MgCO₃ (stage II). The amount of MgCO₃ formed increases gradually, resulting in an overall MgO conversion of 40% after 12 min of carbonation. When compared to MgO-(Na₂CO₃/NaNO₃) (Figure 1e), which reaches 40% conversion in approximately 1 min, the formation of MgCO₃ is significantly slower in MgO-(Na₂CO₃/RbNO₃). Figure 3c compares the rates of MgCO₃ formation as a function of time for MgO-(Na₂CO₃/RbNO₃), MgO-(Na₂CO₃/NaNO₃) and MgO-NaNO₃. The maximal observed rate of MgCO₃ formation for MgO-(Na₂CO₃/NaNO₃) is 14.6 mmolMgCO₃ molMgO⁻¹ s⁻¹, i.e., >10 times faster than the maximal rate that is observed for MgO-(Na₂CO₃/RbNO₃), viz. 1.4 mmolMgCO₃ molMgO⁻¹ s⁻¹ rate. Interestingly, the maximal MgCO₃ formation rates for MgO-(Na₂CO₃/RbNO₃) and MgO-NaNO₃ are similar, yet in MgO-(Na₂CO₃/RbNO₃) the maximal rate is reached earlier, viz. after 2.5 min as compared to 7 min for MgO-NaNO₃. This suggests that MgCO₃ nucleation is more favorable in MgO-(Na₂CO₃/RbNO₃) than in MgO-NaNO₃.

It is important to note that the rapid and high CO₂ uptake of MgO-(Na₂CO₃/RbNO₃) cannot be explained solely by the presence of Na₂CO₃ or Na₂Mg(CO₃)₂ as the sorbent MgO-Na₂CO₃ shows only a small CO₂ uptake (Figure S19).
<figure><img src="image_13.png" /><figcaption>**Figure 4.** Schematic representation of the proposed MgCO₃ formation mechanisms for (a) NaNO₃-promoted MgO and (b) (Na₂CO₃/NaNO₃)-promoted MgO.</figcaption></figure>

Furthermore, the structural evolution of MgO-Na₂CO₃ during the first 12 min of carbonation, as obtained from synchrotron-based *in situ* XRD data (Figure 3d), reveals that the observed CO₂ uptake is due to the formation of Na₂Mg(CO₃)₂ only, as no peaks due to MgCO₃ are observed. Even after 1 h of carbonation (Figure S22), no crystalline MgCO₃ phase can be detected. Hence, it can be concluded that the presence of a molten salt (RbNO₃ or NaNO₃) is a necessary requirement for MgCO₃ formation. It is also noted that the formation of Na₂Mg(CO₃)₂ is significantly slower in the absence of a molten salt.

The dissolution of [Mg²⁺...CO₃²⁻] ion pairs in molten alkali metal nitrates has been established as a key step in the formation of crystalline MgCO₃ for NaNO₃-promoted MgO-based CO₂ sorbents.²⁰ Thus, an explanation for the observation that RbNO₃ does not promote the CO₂ uptake of MgO on its own, while it does in the presence of Na₂Mg(CO₃)₂, could be related to a low solubility of [Mg²⁺...CO₃²⁻] ion pairs in RbNO₃ as compared to NaNO₃. To verify this hypothesis, we performed AIMD calculations. Previous AIMD calculations reported by Landuyt et al. have determined that the dissolution energy of [Mg²⁺...CO₃²⁻] ion pairs in NaNO₃ (E_MgCO₃/NaNO₃) is 1.1 eV.²⁰ Here, using an identical approach, we determined a dissolution energy of [Mg²⁺...CO₃²⁻] ion pairs in RbNO₃ (E_MgCO₃/RbNO₃) of 2.61 eV (Figure S23), which is approximately 2.5 times higher than E_MgCO₃/NaNO₃. The higher energy barrier for [Mg²⁺...CO₃²⁻] ion pair dissolution in MgO-(Na₂CO₃/RbNO₃) results in a lower [Mg²⁺...CO₃²⁻] ion pair concentration, preventing any XRD-detectable nucleation of MgCO₃.

## Carbonation Mechanisms

The results presented above provide key insights into the mechanisms of MgCO₃ formation in MgO-NaNO₃ and MgO-(Na₂CO₃/NaNO₃), suggesting fundamental differences between the mechanisms at play in these two sorbents. In MgO-NaNO₃, the rate of MgCO₃ formation is controlled by the nucleation of MgCO₃ as evidenced from the relatively long induction period in combination with the observation of a continuous conversion of individual MgO crystallites into MgCO₃. Initially, only a few MgCO₃ crystallites are formed, that grow into larger crystallites, while simultaneously promoting MgCO₃ nucleation in neighboring MgO crystallites. This leads to a continuous nucleation and growth process (or autocatalytic crystallization process), as illustrated in Figure 4.⁴¹,⁵⁶ Such an autocatalytic growth process could result from a

facilitated nucleation process on sites in proximity to the newly formed MgCO₃ crystallites.

In MgO-(Na₂CO₃/NaNO₃), the induction period of MgCO₃ formation is significantly shorter (19 s as compared to 3 min in MgO-NaNO₃) and the maximal rate of MgCO₃ formation is approximately 10 times higher than for MgO-NaNO₃. Furthermore, *in situ* XRD shows that the formation of Na₂Mg(CO₃)₂ crystallites precedes the formation of MgCO₃. Therefore, the Na₂Mg(CO₃)₂ crystallites act as nucleation seeds for MgCO₃ formation, greatly facilitating MgCO₃ nucleation (most likely by lowering the energy barrier for MgCO₃ nucleation), as illustrated in Figure 4b.

However, the presence of Na₂CO₃ or Na₂Mg(CO₃)₂ only is not sufficient to effectively promote MgCO₃ formation as shown by the poor CO₂ uptake performance of MgO-Na₂CO₃. Instead, a molten phase is required to dissolve the reactants, most likely in the form of [Mg²⁺...CO₃²⁻] ion pairs, that subsequently crystallize as MgCO₃. Notably, the addition of Na₂CO₃ which rapidly converts into Na₂Mg(CO₃)₂, allows for the formation of MgCO₃ when using a molten nitrate (RbNO₃) that on its own does not lead to an appreciable MgCO₃ formation due to a relatively high energy barrier for the dissolution of [Mg²⁺...CO₃²⁻] ion pairs in RbNO₃. We hypothesize that the presence of nucleation seeds allows for the nucleation of MgCO₃ at significantly lower concentrations of [Mg²⁺...CO₃²⁻] ion pairs.⁵⁷ The enhanced crystallization rate of MgCO₃ could simultaneously be a result of a higher initial concentration of carbonate ions in the molten nitrates (NaNO₃ or RbNO₃) due to the presence of dissolved Na₂CO₃. The slower rate of MgCO₃ formation in MgO-(Na₂CO₃/RbNO₃) compared to MgO-(Na₂CO₃/NaNO₃) indicates that the rate and concentration of ion dissolution in the melt, which is strongly affected by the nature of the salt (confirmed by AIMD) is a critical parameter for the rate of MgCO₃ formation. Overall, these studies suggest that the combined effect of Na₂Mg(CO₃)₂ seeding and a high dissolution of ions in the melt explain the promotional effect of adding Na₂CO₃. As reported previously, other carbonates (such as CaCO₃, BaCO₃ and K₂CO₃) also promote the formation of MgCO₃.³²,⁴²,⁵⁸ Future studies should assess whether also such (double) carbonates have a structural similarity with MgCO₃ and whether they can act as nucleation seeds.

## CONCLUSIONS

In this work, we elucidated the mechanism through which the addition of Na₂CO₃ enhances and accelerates the CO₂ uptake
of NaNO₃-promoted MgO. Utilizing in situ synchrotron-based XRD with a high time resolution, we observed that Na₂CO₃ converts rapidly into Na₂Mg(CO₃)₂, i.e., within seconds during carbonation and precedes the formation of MgCO₃. Na₂Mg-(CO₃)₂ crystallites act as nucleation seeds, accelerating the nucleation of MgCO₃ and significantly enhancing the CO₂ uptake kinetics. Specifically, the maximal CO₂ uptake rate increases from 1.4 to 14.6 mmol MgCO₃ (mol MgO)⁻¹ s⁻¹ when Na₂CO₃ is added to NaNO₃-promoted MgO. Additionally, the presence of a high concentration of carbonate ions in the melt, due to a partial dissolution of Na₂CO₃ in NaNO₃, further accelerates the carbonation kinetics in MgO-(Na₂CO₃/NaNO₃). We also observed that the addition of Na₂Mg(CO₃)₂ as nucleation seeds promotes MgCO₃ formation even in the presence of a molten salt (RbNO₃) that poorly dissolves [Mg²⁺...CO₃²⁻] ion pairs, likely by lowering the critical ion pair concentration required for MgCO₃ nucleation. These findings open new possibilities for the development of more effective MgO-based CO₂-sorbents by using promoter mixtures that contain nucleation seeds or lead to the rapid, in situ formation of nucleation seeds.

<figure><img src="image_14.png" /></figure>

# EXPERIMENTAL SECTION

## Materials

Magnesium nitride (Mg₃N₂, 99.5%), sodium nitrate (NaNO₃, 99.5%) and rubidium nitrate (RbNO₃, 99.7%) were purchased from Sigma-Aldrich. Sodium carbonate (Na₂CO₃, 99.5%) and magnesium nitrate hexahydrate (Mg(NO₃)₂·6H₂O, ≥99%) were purchased from Acros Organics.

## Preparation of the Sorbents

MgO was prepared by calcining high surface area Mg(OH)₂ (synthesized through a stoichiometric reaction of Mg₃N₂ with H₂O at room temperature) for 4 h at 500 °C in air. For the preparation of Na₂CO₃-promoted MgO, the following amounts of precursor were used: 0.2 g Na₂CO₃, 1.5 g MgO (Na₂CO₃:MgO molar ratio of 0.05) and 10 mL isopropanol. The precursors were mixed and ball-milled with a Pulverisette 7 planetary micro mill at 100 rpm for 20 h (60 cycles of 10 min, followed by a 10 min break) using 1 mm ZrO₂ balls in a Si₃N₄ cylinder. After ball milling, the sample was collected in a crystallization dish and dried at 50 °C and 1 mbar for 2 h. For unpromoted MgO, a similar milling procedure was used but without the addition of Na₂CO₃. To obtain the alkali metal nitrate-promoted sorbents, the milled materials were mixed with 10 mol % NaNO₃ or RbNO₃ by grinding for 20 min with a mortar and pestle.

## Evaluation of Sorbents

CO₂ capture experiments were carried out in a TGA DSC3 + instrument. In a typical experiment, 15 mg of sample was loaded into an alumina crucible. The sample was first pretreated at 450 °C for 30 min in N₂, followed by carbonation at 315 °C for 1 h in CO₂ and cooldown to room temperature in CO₂. The flow rate of CO₂ or N₂ was set to 80 mL min⁻¹ and the heating and cooling rates were set to 50 °C min⁻¹.

## Characterization

Ex situ XRD patterns were collected using a PANalytical Empyrean X-ray powder diffractometer equipped with a Bragg–Brentano HD mirror and operated at 45 kV and 40 mA using Cu Kα radiation. The scans were collected in the range of 20–70° using a step size of 0.033° and a time per step of 2.4 s. (S)TEM images and STEM-EDX maps were acquired on a Talos F200X microscope at 200 kV equipped with 4 Super-X EDS detectors. ICP-OES analysis was performed on an Agilent 5100 VDV instrument to determine the Na:Mg ratio in the as-prepared Na₂CO₃-promoted sorbent.

## In Situ Synchrotron XRD Experiments

In situ XRD experiments were performed at the BM01 of the European Synchrotron Radiation Facility (ESRF, Grenoble, France).⁵⁹ The monochromatic X-ray beam with an energy of 17.8 keV (λ = 0.69668 Å) was focused to a size of 0.5 × 0.5 mm² and the data was acquired using a PILATUS 2 M 24-0111 detector. XRD patterns were collected at a sample-to-detector distance of ca. 34 cm (Qmax = 8 Å⁻¹, Q is the modulus of the scattering vector; Q = 4π/λ·sin α). The in situ experiments were carried out using a quartz capillary cell (1 mm diameter and a wall thickness of 0.01 mm) in which the material (ca. 2 mg) was fixed using quartz wool (see Figure S24 for an image of the setup). The temperature of the sample was controlled using a gas blower, which was calibrated using a thermocouple placed inside an empty quartz capillary. The sample-to-blower distance was fixed in the setup allowing for high reproducibility of the temperature and thus the determined cell parameters. Data acquisition was performed during pretreatment, carbonation and cooling down. The samples were first pretreated at 450 °C for 30 min under a flow of He (7.1 mL min⁻¹) to remove any adsorbates (e.g., CO₂ or hydroxyl groups) from the MgO surface. Subsequently, the samples were cooled down to the carbonation temperature of 315 °C. At 315 °C, the gas flow was switched to a flow of CO₂ (13.5 mL min⁻¹) for a duration of 60 min unless stated otherwise. During pretreatment and cooling down, data was acquired every 3 min (acquisition time: 1 s), during carbonation, data were acquired every 1 s. A LaB₆ NIST (0.3 mm) diffraction standard was used for calibration of the instrument. Data normalization and azimuthal integration was performed using the SNBL software Bubble.⁵⁹

### Rietveld Refinement of the In Situ XRD Data during Carbonation.
Parametrical Rietveld refinement of the in situ XRD data was performed using the software TOPAS 6.0 (Bruker AXS).⁶⁰ The weight percentages, cell parameters and average crystallite sizes of each phase were determined as a function of time. In general, every 10th scan was used for refinement to reduce the amount of data to be analyzed. For the material (Na₂CO₃/NaNO₃)-MgO, every scan was refined during the first 60 s of the carbonation reaction, to follow the rapid dynamics. Subsequently, every 10th scan was used. For the refinement, the crystal structures of MgCO₃ (ICSD-40119), MgO (ICSD-101007) and Na₂Mg(CO₃)₂ (ICSD-100482) from the Inorganic Crystal Structure Database (ICSD) were used. The determination of the MgO conversion was based on Rietveld refinement. Here, first the weight percentages of the phases, obtained through Rietveld refinement, were converted into molar percentages. Subsequently, the MgO conversion was calculated as (n<sub>MgO</sub>(t = 0) – n<sub>MgO</sub>(t = x)) × n<sub>MgO</sub>(t = 0)⁻¹ × 100%, where n<sub>MgO</sub>(t = 0) is the molar percentage of MgO at time = 0 and n<sub>MgO</sub>(t = x) is the molar percentage of MgO after x seconds of carbonation. It should be noted that this calculation is based on crystalline phases only. The instrumental peak function was obtained from a measurement of a microcrystalline MgO powder in a 1 mm capillary, i.e., identical to the in situ experiment (the LaB₆ NIST standard in a 1 mm capillary could not be used due to its high absorption at the used energy).

## AIMD Modeling of MgCO₃ Dissolution

AIMD simulations were performed using a plane-wave basis set, as implemented in VASP.⁶¹,⁶² The projector augmented-wave method⁶³ was used to describe the core electrons with the Perdew–Burke–Ernzerhof exchange–correlation (XC) functional.⁶⁴ The kinetic energy cutoff for the wave function and charge density was set to 450 eV, and single γ k-point was used. Two MD runs were carried out using the NVT ensemble and a Nose-Hoover thermostat. In the first run, the systems were annealed from 300 to 600 K over 5 ps. In the second run, the resultant final configurations from the first run were annealed at 600 K for another 5 ps. We computed the total energy by taking the average of the final 2 ps of the second run. The solubilities of [Mg²⁺...CO₃²⁻] ion pairs were investigated in NaNO₃ and RbNO₃ molten matrices. The afore-described MD procedure was used to obtain the total energies of the systems with and without the [Mg²⁺...CO₃²⁻] ion pairs dissolved in NaNO₃ or RbNO₃. The Packmol code
was used to build these configurations using a unit cell with a size of
10 × 10 × 10 Å.⁶⁵

■ ASSOCIATED CONTENT

SI Supporting Information

The Supporting Information is available free of charge at
https://pubs.acs.org/doi/10.1021/jacsau.4c00782.

Example Rietveld refinements and additional details of the *in situ* XRD setup and the Rietveld refinements. Supplementary results include *in situ* XRD data, (S)TEM and EDX images, TGA results and *ex situ* XRD patterns (PDF)

■ AUTHOR INFORMATION

Corresponding Authors

**Paula M. Abdala** — Laboratory of Energy Science and Engineering, Department of Mechanical and Process Engineering, Eidgenössische Technische Hochschule (ETH) Zürich, 8092 Zürich, Switzerland; ● orcid.org/0000-0002-2011-1707; Email: abdalap@ethz.ch

Christoph R. Müller — Laboratory of Energy Science and Engineering, Department of Mechanical and Process Engineering, Eidgenössische Technische Hochschule (ETH) Zürich, 8092 Zürich, Switzerland; ● orcid.org/0000-0003-2234-6902; Email: muelchri@ethz.ch

Authors

Annelies Landuyt — Laboratory of Energy Science and Engineering, Department of Mechanical and Process Engineering, Eidgenössische Technische Hochschule (ETH) Zürich, 8092 Zürich, Switzerland; ● orcid.org/0000-0003-4045-5360

Ilia Kochetygov — Paul Scherrer Institut, PSI Center for Energy and Environmental Sciences, Villigen CH-5232, Switzerland; ● orcid.org/0000-0001-8013-5885

Charles J. McMonagle — Swiss-Norwegian Beamlines (SNBL), European Synchrotron Radiation Facility (ESRF), Grenoble 38000, France

Priyank V. Kumar — School of Chemical Engineering, The University of New South Wales (UNSW Sydney), 2052 Sydney, New South Wales, Australia; ● orcid.org/0000-0002-8203-7223

Jodie A. Yuwono — School of Chemical Engineering, The University of Adelaide, Adelaide, SA 5005, Australia

Wendy L. Queen — Institute of Chemical Sciences and Engineering (ISIC), École Polytechnique Fédérale de Lausanne (EPFL), 1051 Sion, Switzerland; ● orcid.org/0000-0002-8375-2341

Complete contact information is available at:
https://pubs.acs.org/10.1021/jacsau.4c00782

Author Contributions

CRediT: **A.L.** conceptualization, data curation, formal analysis, investigation, methodology, visualization, writing-original draft; **I.K.** formal analysis, software, writing-review and editing; **C.J.M.** methodology, writing-review and editing; **J.A.Y.** formal analysis, software, writing-review and editing; **P.V.K.** formal analysis, software, writing-review and editing; **W.L.Q.** resources, writing-review and editing; **P.M.A.** conceptualization, supervision, validation, writing-review and editing; **C.R.M.** conceptualization, funding acquisition, project admin-

istration, resources, supervision, validation, writing-review and editing.

Notes

The authors declare no competing financial interest.

■ ACKNOWLEDGMENTS

The European Synchrotron Radiation Facility (ESRF) is acknowledged for providing access to the beamline through proposal CH-6089. Drs. Vadim Diadkin, Dmitry Chernyshov, and Dragos Stoian at SNBL are acknowledged for their excellent assistance and discussions during the beamtime. Dr. A. Kierzkowska is acknowledged for performing the ICP measurements. This project has received funding from the European Research Council (ERC) under the European Union's Horizon 2020 Research and Innovation Program Grant Agreement 819573. We are grateful to ScopeM (ETH Zürich) for the use of their electron microscopy facilities.

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