Impact of Operational Parameters on the CO₂ Absorption Rate and Uptake in MgO Aqueous Carbonation—A Comparison with Ca(OH)₂

Abstract

The CO₂ absorption rate and total uptake by MgO aqueous suspensions were investigated in batch experiments by systematically varying MgO concentrations (0.5–5 wt.%), CO₂ flow rates (0.5–2 L/min), temperatures (278–363 K), NaCl salinities (0–7 wt.%), Na₂SO₄ and K₂SO₄ concentrations (0–10.5 wt.%), and gas–liquid mixing systems (pipe outlet and porous stone sparger). Results show that temperature strongly controls the carbonation process: increasing temperature above 303 K consistently reduced both the CO₂ absorption rate $\eta \left(t\right)$ and the total CO₂ uptake $V_{CO2}$ due to the destabilization of metastable Mg(HCO₃)₂ solutions and accelerated precipitation of less soluble hydrated magnesium carbonates. Under optimal low-temperature conditions (278–283 K, 1–1.5 wt.% MgO, sparger mixing, pure system), the average capture efficiency reached ≈ 35%, with maximum peaks over 70% and total CO₂ uptakes of ≈ 12–17 L. Adding NaCl at typical seawater levels (3.5–7 wt.%) slightly increased CO₂ uptake at temperatures above 323 K. Sulfate ions (Na₂SO₄ and K₂SO₄) were found to enhance the absorption rate at low concentrations (<2 wt.%) but reduce it at higher levels, with no significant impact on the total CO₂ uptake observed in this study. Using a CO₂ sparger significantly improved gas–liquid contact, achieving average CO₂ capture efficiencies ${\eta }_{max}\left(t\right)$ above 70% at low temperatures, compared to <20% with simple pipe bubbling. A direct comparison with Ca(OH)₂ aqueous carbonation confirmed that, despite its lower solubility and slower kinetics, MgO can outperform Ca-based systems under specific conditions. These results provide practical experimental benchmarks and process guidance for designing Mg-based aqueous carbonation systems, including applications that use brines, industrial wastewater or seawater.

1. Introduction

Engineered CO₂ mineralization is a promising carbon dioxide storage strategy that mimics natural weathering processes to achieve long-term CO₂ sequestration through the formation of stable carbonate minerals, primarily from calcium (Ca)-, magnesium (Mg)-, and iron (Fe)-bearing materials [1,2]. Ca- and Mg-rich materials have drawn particular attention owing to their higher carbonation reactivity compared to Fe-bearing phases, making them the most studied feedstocks for engineered carbonation [3,4].

Mg compounds are known to exhibit slower CO₂ uptake compared to Ca-based counterparts during aqueous carbonation [5]. Mg-based carbonation systems are inherently more challenging due to the high stability of Mg aqueous complexes, which hinders dehydration, and to the wider range of possible carbonate phases [6]. In water, Mg²⁺ predominantly exists as the strongly hydrated hexaaqua complex [Mg(H₂O)₆]²⁺, characterized by high hydration enthalpy and slow water-exchange kinetics compared with Ca²⁺ [7]. Carbonate binding and nucleation therefore require partial dehydration of Mg²⁺, introducing a kinetic barrier to precipitation. As a result, Mg systems often evolve through metastable aqueous or hydrated carbonate intermediates, whereas Ca systems carbonate more readily owing to the weaker hydration of Ca²⁺ and faster Ca(OH)₂ dissolution. In essence, the challenge lies in the aqueous-to-solid transition kinetics of Mg systems. Thus, while Ca minerals yield CaCO₃ polymorphs upon carbonation in typical conditions, Mg minerals yield chemically distinct carbonates. The most common phases obtained in laboratories are landsfordite (Lfd, MgCO₃·5H₂O), nesquehonite (Nes, MgCO₃·3H₂O), dypingite (Dyp, Mg₅(CO₃)₄(OH)₂·5H₂O) and hydromagnesite (Hmgs, Mg₅(CO₃)₄(OH)₂·4H₂O) [8]. Magnesite (Mgs, MgCO₃) is desirable for its stability but is kinetically hindered at low temperatures by dehydration barriers, and typically does not form below 363 K on laboratory timescales, unless high CO₂ pressures, seeding, or additives are employed [7,9,10,11]. In general, the thermodynamic stability of Mg carbonates increases from more to less hydrated phases, in the following order: Lfd < Nes < Dyp < Hmgs < Mgs. At near-ambient temperatures, Nes is the most commonly obtained phase, but it transforms into Dyp and/or Hmgs through a temperature-dependent kinetic process [12]. Extensive studies on Mg-carbonate physical chemistry reveal a wide metastability range, with small variations leading to different phases [13].

For carbon capture and storage (CCS) applications, Mgs formation is often preferred over Hmgs and Dyp because of its higher CO₂ uptake capacity, stemming from its superior Mg:C ratio. MgO-based systems have been widely proposed as an integral part of emerging CCS strategies, including the carbonation of olivine [14,15,16,17,18,19,20,21], ultramafic rocks such as peridotites [22,23], mine tailings [24,25,26,27], steel slags [28,29,30,31], MgO cements and mortars [32,33,34,35], water treatment [36,37], contaminated soil carbonation [38,39], and the magnesium looping approach (MgL) [40,41,42,43]. While promising at the global scale, the design of these processes depends critically on achieving favorable kinetics, predictable product speciation, and high CO₂ absorption efficiency under practical operating conditions.

Despite extensive research on MgO carbonation, key aspects of its aqueous system remain poorly quantified. In particular, the formation and stability of metastable magnesium bicarbonate solutions under varying conditions remain poorly constrained, despite their key role in MgO solubility and carbonate precipitation [44,45,46]. Furthermore, while the effect of temperature on Mg-carbonate phase transformations is well established, the influence of salinity, particularly chlorides and sulfates, on CO₂ absorption kinetics and uptake has received little systematic attention. This gap is critical, since industrial carbonation processes will likely use seawater or brines naturally containing these ions.

This study addresses these gaps by systematically investigating the aqueous carbonation of reactive MgO under controlled conditions, with particular focus on the effects of temperature, NaCl salinity, and sulfate ions on CO₂ absorption rate and total uptake. Building on our previous work on Ca(OH)₂ carbonation under comparable conditions [47], this study directly compares CO₂ absorption rate and total uptake between pure CaO and MgO aqueous systems, highlighting the key kinetic and operational implications.

A detailed analysis of operational parameters enables the design of efficient strategies to address a key question: how can a carbonation scrubber be optimized to absorb CO₂ as rapidly as possible? The investigated operational parameters include the following:

• Periclase (MgO) initial concentration.
• CO₂ flow rate.
• Temperature.
• NaCl, Na₂SO₄ and K₂SO₄ concentrations (salinity).
• Static gas−liquid mixing system, with a comparison of the pipe and the sparger.

An experimental dataset, based on batch experiments, was collected using high-precision gas flow systems to measure the percentage of flowing CO₂ absorbed by the reagents. The present study provides a novel contribution to the existing literature by implementing a new and accurate quantification method to assess the effects of chemical factors (reagent concentrations and chloride and sulfate ions) and physical variables (flow rate, temperature, and mixing system design) on the CO₂ absorption rate and uptake in MgO aqueous carbonation.

2. Materials and Methods

2.1. Reagents

MgO was obtained by calcination of Mg carbonate (Sigma-Aldrich, 99% purity, Darmstadt, Germany) at 1173 K for three hours. The calcined MgO was stored in a sealed container to prevent premature hydration. The BET SSA of the MgO was 31.886 m²/g ± 0.116 (R² = 0.999). Analytical-grade NaCl, Na₂SO₄ and K₂SO₄ were used to adjust salinity and sulfate concentrations in selected experiments (Sigma-Aldrich). All solutions were prepared using deionized water (18.2 MΩ·cm) to ensure consistent baseline chemistry. Phenolphthalein and methyl red indicators, along with hydrochloric acid (HCl), were also provided by Sigma-Aldrich. Each experiment was conducted in a double-wall Pyrex reactor with a capacity of 1.5 L, a height of 16.8 cm, and a diameter of 10 cm, equipped with a thermostatic bath, as reported by our previous Ca(OH)₂ study [47].

2.2. Laboratory Equipment

Each experiment was conducted in the setup shown in Figure 1, using a 1.5 L double-wall Pyrex reactor. The figure illustrates the modular design employed in this study, juxtaposing the standard bubbling configuration with the sparger setup. The sparger was constituted of a porous stone cylinder with a diameter of 25 mm × 40 mm height. The experimental setup was designed to maintain strict control over gas–liquid interactions during CO₂ absorption.

CO₂ flows were controlled and measured by a Bronkhorst modular system, composed of: Mass flow controller (“valve” hereafter) (MFC D-6321), with a maximum flow range of 2 L/min CO₂ and accuracy ± 1.0% relative deviation RD hereafter) plus ± 0.5% full scale (FS hereafter); a low ∆P (“flow meter” hereafter) F101E (max flow range of 3 L/min CO₂ and accuracy ± 1.0% FS). A digital pc-board provides self-diagnostics, alarm and counter functions, digital communication (RS232), remotely adjustable control settings, and an on-board interface based on the FLOW-BUS protocol, making it possible to communicate via a multi-bus system. CO₂ absorption acquisition data was performed every 0.25–5 s according to the specific accuracy requirements of the experimental run.

pH and EC (electrical conductivity, µS/cm) of the samples were measured by a Hanna HI H-ORP meter (Hanna Instruments, Woonsocket, RI, USA) and a Mettler Toledo Five Easy EC-meter (Mettler-Toledo, Greifensee, Switzerland), respectively. The samples’ masses were measured with a precision of ±0.1%.

Scanning Electron Microscopy (SEM) imaging was conducted utilizing a TESCAN VEGA (Brno-Kohoutovice, Czech Republic), equipped with secondary electron (SE) and backscattered electron (BSE) detectors. Typical experimental settings included a tungsten (W) filament, an accelerating voltage of 10 kV, a beam current of 30 pA, and a working distance of 6 mm.

X-Ray Powder Diffraction (XRPD) measurements for the qualitative identification of the crystalline phases were carried out using a Rigaku MiniFlex 600 benchtop X-ray diffractometer (Rigaku, Tokyo, Japan, Bragg–Brentano geometry, CuKα radiation, X-ray source operating at 600 W (40 kV, 15 mA); D/teX Ultra2 silicon strip detector; 3° < 2θ < 70°, step width 0.02°, scan speed 1 °/min).

Specific surface area (SSA) analysis was performed using a Micromeritics ASAP 2020 instrument (Micromeritics, Norcross, GA, USA) employing the Brunauer–Emmett–Teller (BET) method with N₂ adsorption at −196 °C. Prior to measurement, samples underwent degassing at 150 °C for 240 min.

2.3. Experimental Procedures for MgO Carbonation

Table 1. Overview of the experimental parameters and their corresponding range.

Variables	Units	Ranges/Types
MgO initial concentration	wt.%	0.5–5
CO2 volumetric flow rate	L/min	0.5
Temperature	K	278–363
NaCl concentration	wt.%	0–7
Na2SO4	wt.%	0–10.5
K2SO4	wt.%	0–10.5
mixing system	-	pipe, sparger
Constants		Values/types
MgO initial SSA	m2/g	31.9
water mass	kg	1.2
CO2 partial pressure	bar	1.2
CO2 concentration	Vol.%	100
stirrer speed	rpm	300
cylinder height × diameter	cm	16.8 × 10

summarizes the experimental parameters investigated with their operational ranges. Detailed experimental conditions with their corresponding identifiers are provided in

Table A1. Detailed experimental parameters (variables) for the study with their corresponding experiment identifiers.

Experiment Identifier		[MgO]0	CO2 Flow Rate	Temperature	[Salt]	Mixing System
		wt.%	L/min	K	wt.%	-
Group 1	A1	1.5	0.5	278	0	sparger + stirrer
	A2	1.5	0.5	278	3.5	sparger + stirrer
Group 2	B1	0.5	0.5	278	0	sparger + stirrer
	B2	1	0.5	278	0	sparger + stirrer
	B3	1.5	0.5	278	0	sparger + stirrer
	B4	2	0.5	278	0	sparger + stirrer
	B5	2.5	0.5	278	0	sparger + stirrer
	B6	3	0.5	278	0	sparger + stirrer
	B7	3.5	0.5	278	0	sparger + stirrer
	B8	4	0.5	278	0	sparger + stirrer
	B9	4.5	0.5	278	0	sparger + stirrer
	B10	5	0.5	278	0	sparger + stirrer
Group 3	C1	1	0.5	303	0	sparger + stirrer
	C2	1	0.75	303	0	sparger + stirrer
	C3	1	1	303	0	sparger + stirrer
	C4	1	1.5	303	0	sparger + stirrer
	C5	1	2	303	0	sparger + stirrer
Group 4	D1	1	0.5	283	0	sparger + stirrer
	D2	1	0.5	303	0	sparger + stirrer
	D3	1	0.5	323	0	sparger + stirrer
	D4	1	0.5	343	0	sparger + stirrer
	D5	1	0.5	363	0	sparger + stirrer
	D6	1	0.5	283	3.5 [NaCl]	sparger + stirrer
	D7	1	0.5	303	3.5 [NaCl]	sparger + stirrer
	D8	1	0.5	323	3.5 [NaCl]	sparger + stirrer
	D9	1	0.5	343	3.5 [NaCl]	sparger + stirrer
	D10	1	0.5	363	3.5 [NaCl]	sparger + stirrer
	D11	1	0.5	283	7 [NaCl]	sparger + stirrer
	D12	1	0.5	303	7 [NaCl]	sparger + stirrer
	D13	1	0.5	323	7 [NaCl]	sparger + stirrer
	D14	1	0.5	343	7 [NaCl]	sparger + stirrer
	D15	1	0.5	363	7 [NaCl]	sparger + stirrer
Group 5	E1	1	0.5	303	1 [Na2SO4]	sparger + stirrer
	E2	1	0.5	303	3.5 [Na2SO4]	sparger + stirrer
	E3	1	0.5	303	7 [Na2SO4]	sparger + stirrer
	E4	1	0.5	303	10.5 [Na2SO4]	sparger + stirrer
Group 6	F1	1	0.5	303	1 [K2SO4]	sparger + stirrer
	F2	1	0.5	303	3.5 [K2SO4]	sparger + stirrer
	F3	1	0.5	303	7 [K2SO4]	sparger + stirrer
	F4	1	0.5	303	10.5 [K2SO4]	sparger + stirrer
Group 7	G1	1	0.5	303	0	pipe + stirrer

(Appendix A). Experimental runs were performed by changing the operational parameters one by one to isolate their impact on the CO₂ absorption rate. Specific conditions were carefully selected to optimize the representativeness of each parameter investigated.

Forty batch experiments were conducted, labeled A1–2, B1–10, C1–5, D1–15, E1–4, F1–4, and G1, divided in seven groups. Group 1 experiments, A1 and A2, focused on pH/EC and total alkalinity (TA) measurements at near-ambient temperature (303 K). Group 2 experiments, B1–10, examined the effect of MgO concentration, spanning ten concentrations from 0.5 to 5 wt.%. Group 3 experiments, C1–5, explored CO₂ flow rates from 0.5 to 2 L/min. Group 4, the D1–15 series assessed three NaCl concentrations (0%, 3.5%, and 7.0 wt.%) across five temperatures (283 K, 303 K, 323 K, 343 K, and 363 K) using a quadratic design. Groups 5 (E1–4) and 6 (F1–4) investigated the effects of Na₂SO₄ and K₂SO₄, respectively, at four concentrations (0%, 3.5%, 7.0 wt.% and 10.5 wt.%). Finally, G1 examined mixing systems, comparing regular pipe CO₂ bubbling with the sparger baseline. The selection of these variables and their respective ranges was guided by both practical constraints of the experimental setup and representative conditions of environmental and industrial carbonation systems. Temperature, salinity, CO₂ flow rate, and MgO concentration were varied within the operational limits of the apparatus while spanning values typically encountered in natural and engineered carbonation contexts. For example, NaCl and sulfate concentrations were chosen to reflect seawater salinity (3.5 wt.%) and ionic strengths relevant to industrial effluents and alkaline residues, whereas the 0.5–5 wt.% MgO range corresponds to typical solid contents of such materials. The CO₂ flow-to-liquid ratio (up to 2 L/min per 1 L reactor) was selected as a realistic gas–liquid contact condition for process-scale operation. For the cross-system context, the MgO runs used the same CO₂ partial pressure (1.2 bar), gas composition (100% CO₂), and stirrer speed (300 rpm) as the Ca(OH)₂ campaign. The liquid/solid (L/S) ratio window was broader here (20–200 for MgO vs. 10–100 in Ca(OH)₂) to ensure stable suspension of MgO with magnetic stirring at the lowest L/S values, while optimizing $\eta \left(t\right)$ by maximizing water column height using 1.2 kg of water. Regarding gas–liquid contactors, the Ca(OH)₂ dataset used a regular pipe bubbling system as the baseline, with several additional experiments employing porous-stone spargers. In contrast, the MgO dataset almost systematically used a porous-stone sparger (39/40 experiments), since the CO₂ absorption rate obtained with a simple pipe was too low to clearly highlight the influence of the other experimental variables, as demonstrated by experiment G1.

Prior to carbonation, MgO powder was added to deionized water at the desired liquid-to-solid (L/S) ratio directly in the reactor to avoid particle losses on the flask walls. The concentration of NaCl, Na₂SO₄, or K₂SO₄ was then adjusted as required. Suspensions were equilibrated to the target temperature using a thermostatic bath connected to the reactor’s double wall, and temperature was monitored directly within the solution via a PT100 thermocouple. In the instance of experiment D1–15, an extended preparation time was implemented to attain the targeted temperature, reaching from 283 to 363 K.

Air-free experiments were conducted by bubbling N₂ through the MgO aqueous suspensions during preparation, up to the start of the carbonation experiment, when the reactor was sealed and the air was flushed out with CO₂. Once thermal equilibrium was reached, to prevent immediate reaction between CO₂ and suspension during the initial flushing step, CO₂ was injected at 2 L/min through a dedicated top-mounted pipe, minimizing gas–liquid contact. The flushing duration was <1 min, resulting in less than 30 mL of CO₂ absorption, which represents 0.1% to 1% of the total CO₂ uptake during the experiments.

At t = 0, CO₂ was bubbled into the reactor at the desired flow rate. All experiments were executed under a constant CO₂ flow, ranging from 0.3 L/min to 2 L/min. The experiments reached completion when the flow meter indicated no further CO₂ absorption by the solution (flow_valve = flow_{flow meter}).

Preliminary tests showed that two MgO carbonation runs were sufficient to reliably represent the performance under each set of experimental conditions. Each experiment was therefore conducted in at least duplicate. Several series of “blank” experiments were performed with ultra-pure water and an empty reactor, using the procedures described in the ensuing sections, to correctly calibrate the setup. A linear proportionality was observed to hold between flow measurements at the valve and flow meter, so that FLOW_BLANK_valve = k_cal·FLOW_BLANK_{flow meter}, where FLOW_BLANK_valve (L/min) is the CO₂ inlet flow rate imposed by the mass-flow controller at the valve, FLOW_BLANK_{flow meter} (L/min) is the corresponding outlet flow rate measured by the flow meter during blank runs, and k_cal is the calibration coefficient obtained from the slope of their linear correlation. This coefficient accounts for the small offset between both gas-flow sensors.

Upon completion, a fraction of the solid residues was recovered by filtration, washed with deionized water, and dried at ambient temperature. Subsamples were analyzed for mineralogical characterization using XRPD. Measurement uncertainties for CO₂ uptake were estimated to be below 5%, based on repeated trials and instrument calibration.

2.4. Phreeqc Geochemical Models

The PHREEQC-3.7.3 program [48] was used for equilibrium and kinetic calculations. The minteq.v4.dat database was used for equilibrium simulations of the pure Mg(OH)₂–H₂O–CO₂ system, whereas the sit.dat and pitzer.dat databases were employed for simulations involving NaCl, as salt concentrations in these systems lie beyond the applicability range of the Debye–Hückel theory [48].

The aqueous carbonation of Mg(OH)₂ was simulated at 20 °C, under a fixed CO₂ partial pressure of 1.2 bar, to explore the speciation pathways and the saturation state of Mg carbonates during progressive CO₂ dissolution. The system consisted of 0.506 mol of brucite equilibrated with 0.5 kg of water, and a total of 1 mol of CO₂ (g) added incrementally in 50 steps. This configuration reproduces the quasi-equilibrium conditions of a CO₂ titration, where alkalinity from brucite dissolution gradually shifts to a bicarbonate-controlled regime as CO₂ loading increases. The SI of Hmgs and Nes, along with ion activities and pH, were calculated as functions of the absorbed CO₂.

The dissolution kinetic model, implemented in this work for brucite using the BASIC interpreter embedded in Phreeqc, has already been adopted by several other authors [49,50,51]. The accuracy of the model predictions was assessed by a comparison with the observed rates. Brucite dissolution rate, its saturation ratio and pH all depend on the amount of water, the crystals’ surface area $\mathrm{S}\mathrm{A}\left(\mathrm{t}\right)$, temperature, Arrhenius pre-exponential factor and activation energy for dissolution Ea.

The initial surface area $\mathrm{S}{\mathrm{A}}_0$, which was determined by BET, decreases as the reaction proceeds, since the particles become smaller in size until they are completely dissolved. This is simulated by calculating the ratio between the initial surface area of brucite and the amount of undissolved solid remaining after each step, according to the following formula:

$$\mathrm{S}\mathrm{A}\left(\mathrm{t}\right)=\mathrm{S}{\mathrm{A}}_0·n\left(t\right)·a_{exp}$$

(1)

where $\mathrm{S}\mathrm{A}\left(\mathrm{t}\right)$ is the surface area at time t (cm²); $\mathrm{S}{\mathrm{A}}_0$ is the initial surface area (cm²/mol); $n\left(t\right)$ the amount of brucite at time t (mol); $a_{exp}$ is an empirical parameter ($a_{exp}$ = 0.5) [52].

The dissolution rate of brucite is described by the following equation [52]:

$$\mathrm{R}\mathrm{a}\mathrm{t}\mathrm{e}=\left(k_1{+k}_2{+k}_3\right)·(1-\Omega )·SA\left(t\right)$$

(2)

where $\Omega$ is the saturation ratio between the ion activity products $IAP$ and the solubility product $K_{sp}$ of Mg(OH)₂, and $k_1{,k}_2{,k}_3$ are the rate constants, respectively, at acidic, neutral and alkaline pH:

$$k_1={\mathrm{A}}_1·\mathrm{e}\mathrm{x}\mathrm{p}(-{\displaystyle \frac{E_{A1}}{RT}})·a(H^{+}{)}^{{a1}_{exp}}$$

(3)

$$k_2={\mathrm{A}}_2·\mathrm{e}\mathrm{x}\mathrm{p}(-{\displaystyle \frac{E_{A2}}{RT}})$$

(4)

$$k_3={\mathrm{A}}_3·\mathrm{e}\mathrm{x}\mathrm{p}(-{\displaystyle \frac{E_{A3}}{RT}})·a({HO}^{-}{)}^{{a2}_{exp}}$$

(5)

${\mathrm{A}}_i$ = pre-exponential factor (${\mathrm{A}}_1=4.0·{10}^5; {\mathrm{A}}_2=1.3·{10}^{-1}; {\mathrm{A}}_3=4.8·{10}^{-6}$); $E_{Ai}$ = activation energy of dissolution ($E_{A1}=59000 \mathrm{J}·{\mathrm{m}\mathrm{o}\mathrm{l}}^{-1}$; $E_{A2}=42000 \mathrm{J}·{\mathrm{m}\mathrm{o}\mathrm{l}}^{-1}$; $E_{A3}=25000 \mathrm{J}·{\mathrm{m}\mathrm{o}\mathrm{l}}^{-1}$) [52]; $R$ = gas constant ($\mathrm{R}=8.31451 \mathrm{J}·{\mathrm{K}}^{-1}·{\mathrm{m}\mathrm{o}\mathrm{l}}^{-1}$); $a\left({HO}^{-}\right)$ and $a\left(H^{+}\right)$ = thermodynamic activities of ${HO}^{-}$ and $H^{+}$, respectively; ${ai}_{exp }=$ reaction orders with respect to H⁺ and OH⁻ activities (acid and base terms) (${a1}_{exp }=$ ${a2}_{exp }= 0.5$).

The model was implemented at temperatures and [NaCl] ranging between 283–343 K and 0–7 wt.%, respectively. Appendix B reports all the equilibrium and kinetic models employed to calculate the results presented in Section 3.1. Additionally, Appendix C (Figure A1) shows the solubility modeling of typical Ca- and Mg-based minerals.

2.5. Data Treatment

The CO₂ absorption rate as a function of time $\eta \left(t\right)$ (%) was determined by measuring the influent and effluent CO₂ volumetric flow with the valve and the flow meter, placed at the reactor inlet and outlet, respectively. In doing so, the instantaneous absorption rate as a function of time was measured, according to the equation below

$$\eta \left(t\right)=100\times \left(1-{\displaystyle \frac{k_{cal}\times \mu \left(t\right)}{B\left(t\right)}}\right)$$

(6)

where $B\left(t\right)$ (L/min) = in-flow (valve); μ(t) (L/min) = out-flow (flow meter); k_cal: calibration factor defined in Section 2.3. Note that η includes the fractions of CO₂ that are either mineralized or dissolved into the active solution.

We introduce the average CO₂ absorption rate over the step $i$ of the carbonation run (defined in Section 3.1), $\overline{{\eta }_i}$ (%), as followed

$$\overline{{\eta }_i}={\displaystyle \frac{\sum _{t_1}^{t_2}\eta \left(t\right)}{n}}$$

(7)

where $\left[t_{1,}{t}_2\right]$ = interval of time for which $\overline{{\eta }_i}$ is being calculated; $n$ = size of the sample studied with $\eta \left(t\right)$ values.

In addition to $\eta \left(t\right)$, we introduce $V_{CO2}$ (L), the volume of CO₂ absorbed during the interval of time $\left[t_{1,}{t}_2\right].$

$$V_{CO2}={\int }_{t_2}^{t_1}\eta \left(t^{\prime }\right)\times B\left(t^{\prime }\right) d{t}^{\prime }$$

(8)

The uncertainty on $\eta \left(t\right)$ and $V_{CO2}$ was estimated from the accuracy of the flow sensors. Based on manufacturer specifications (±1.0% of reading + ±0.5% of full scale) and standard propagation of errors through Equations (1) and (3), the instantaneous uncertainty on $\eta \left(t\right)$ was found to range between ±3% and ±4 percentage points under low-flow conditions, decreasing to ≤2 points at higher flows. After time averaging, the effective uncertainty typically remained within ±2–3 points. For the integrated CO₂ uptake ($V_{CO2}$), the overall propagated relative uncertainty is estimated to be ≤5%.

To measure total alkalinity ($TA$) over time, 80 mL samples were titrated using pH-sensitive indicators. Suspensions were settled for a few minutes and filtered through 0.22 µm syringe filters to release dissolved CO₂ and separate particulates. Samples were then stored in 100 mL sealed bottles. $TA$ (mg CaCO₃/L) was approximated as follows using the equation

$$TA=\left[{HCO}_3^{-}\right]+2\left[{CO}_3^{2-}\right]+\left[{B\left(OH\right)}_4^{-}\right]+\left[{OH}^{-}\right]-\left[H^{+}\right]$$

(9)

where $\left[{HCO}_3^{-}\right]$ = bicarbonate ions; $\left[{CO}_3^{2-}\right]$ = carbonate ions; $\left[{B\left(OH\right)}_4^{-}\right]$= borate ions; $\left[{OH}^{-}\right]$ = hydroxide ions.

Diluted HCl served as the titrant, providing a known source of H⁺ to neutralize the alkalinity species listed in Equation (9). The main acid–base reactions during titration are:

$${CO}_3^{2-}+H^{+}= {HCO}_3^{-}$$

(10)

$${ HCO}_3^{-}+H^{+}=H_{2}C{O}_3=C{O}_2+H_{2}O$$

(11)

$${B\left(OH\right)}_4^{-}+H^{+}=B(OH{)}_3+H_{2}O$$

(12)

These reactions correspond to the two equivalence points observed during titration with phenolphthalein and methyl red indicators. Each titration used 10–25 mL of the sample, with color changes from pale yellow to orange indicating completion.

Phenolphthalein alkalinity (P) and total alkalinity were calculated using the equations below

$$P(meq l^{-1})={\displaystyle \frac{d_{1}N\ast 1000}{V}}$$

(13)

$$T(meq l^{-1})={\displaystyle \frac{d_{2}N\ast 1000}{V}}$$

(14)

where $P$ = phenolphthalein alkalinity; $TA$ = total alkalinity; $d_1, d_2$ (mL) = volume used until the color change with phenolphthalein and methyl red, respectively ($d_2$ includes $d_1$), N = normality of HCl; $V$ (mL) = volume of sample used.

TA was expressed as mg of CaCO₃ equivalent. From the two values, it is also possible to determine which ion in Equation (9) is responsible for the measured alkalinity. Bicarbonates are present only when P < ½ T, and their presence automatically excludes hydroxides. This means that if only bicarbonates are present in solution, the measured pH will be below 8, P will be zero, and the solution will not turn pink with the first indicator, showing only one equivalence point. Conversely, if T ≅ 2P, only carbonates are present in solution, and the same volume of acid is required to reach both the first and second equivalence points. Ultimately, the concentration of bicarbonate in solution was determined as

$${HCO3}^{-}\left(meq l^{-1}\right)=T-2P$$

(15)

3. Results and Discussion

3.1. Geochemical Models

3.1.1. Mg(OH)₂ Aqueous Carbonation Equilibrium

The equilibrium results of the Mg(OH)₂–H₂O–CO₂ system are shown in Figure 2, illustrating the evolution of pH, ion activities, and saturation indices during the progressive equilibrium addition of CO₂, together with the solubility behavior of Mg(OH)₂ as a function of temperature and salinity.

At the start of the simulation (CO₂ < 0.1 mol), the system is buffered by brucite dissolution, maintaining a strongly alkaline pH of ≈ 12.2. Under these hydroxyl-rich conditions (act(OH⁻) ≈ 10⁻²), Mg speciation is dominated by MgOH⁺, and the solution is strongly supersaturated with Hmgs (SI ≈ +12 to +15), while Nes SI is already positive and rising from ≈+0.08 at low CO₂ to ≈+0.87 at 0.1 mol. This confirms that Hmgs is the preferred phase under hydroxyl-rich, low-DIC conditions.

As CO₂ uptake increases (≈0.1–0.5 mol), pH decreases modestly from ≈12 to ≈11.4, act(Mg²⁺) increases several-fold owing to acid-enhanced brucite dissolution, act(MgOH⁺) declines, and act(MgHCO₃⁺) rises to ≈10⁻³, marking a window of maximum Mg solubility. Over this interval, Hmgs SI further increases to ≈+17.6, while Nes SI rises to ≈+1.6, indicating a shift in relative phase stability as DIC accumulates and the system begins transitioning from carbonate- to bicarbonate-dominated chemistry.

At higher CO₂ loading (≈1 mol), the system stabilizes near pH 7.6, and both phases remain supersaturated (Hmgs SI ≈ +6; Nes SI ≈ +0.47), approaching equilibrium but without undersaturation. Overall, the model defines three distinct regimes: (1) an alkaline, Hmgs-dominated domain (pH > 10); (2) a mid-titration window near pH ≈ 11 where MgHCO₃⁺ and Nes formation intensify; and (3) a CO₂-rich, near-neutral regime where both carbonates remain moderately supersaturated, delineating the thermodynamic envelope of Mg(OH)₂ aqueous carbonation under ambient conditions.

The solubility of Mg(OH)₂ decreases slightly with temperature, from 9.4 × 10⁻³ g L⁻¹ at 273 K to 8.9 × 10⁻³ g L⁻¹ at 373 K, consistent with its retrograde solubility and the exothermic nature of brucite dissolution.

As a function of NaCl concentration, Mg(OH)₂ solubility exhibits a non-monotonic trend, increasing up to ≈0.019 g L⁻¹ at 50–60 g L⁻¹ NaCl before decreasing beyond this point, while CO₂ solubility declines continuously.

3.1.2. Mg(OH)₂ Dissolution Kinetic

The dissolution rate of Mg(OH)₂ increases strongly with temperature (Figure 3A), following a clear Arrhenius-type behavior. The initial rate rises from 0.0369 M/s at 283 K to 1.2047 M/s at 343 K, corresponding to an acceleration of about 33 times. The time required to reach 95% of equilibrium decreases accordingly, from 8.64 ms at 283 K to 0.24 ms at 343 K, indicating a 36-fold reduction in equilibration time. Meanwhile, the final equilibrium concentration of dissolved Mg(OH)₂ slightly decreases with temperature, from 6.44 × 10⁻³ g/kgw (283 K) to 6.27 × 10⁻³ g/kgw (343 K). This inverse dependence between solubility and temperature is consistent with the retrograde solubility of Mg(OH)₂ obtained under equilibrium conditions (Figure 2B). Altogether, these results show that increasing temperature accelerates surface-controlled dissolution kinetics while only marginally reducing the equilibrium solubility of brucite.

At 298 K, NaCl concentration exerts a marked influence on Mg(OH)₂ dissolution kinetics (Figure 3B). The initial rate decreases from 0.0724 M/s at 0 M to 0.053–0.054 M/s between 0.05 M and 2.5 M, indicating a 25–30% reduction in the early-time dissolution flux with increasing ionic strength. In contrast, the final dissolved Mg(OH)₂ concentration increases continuously over the studied salinity range, from 0.0054 g/kgw at 0 M to 0.0099 g/kgw at 0.5 M, 0.0078 g/kgw at 1 M, and 0.0040 g/kgw at 2.5 M, consistent with the progressive enhancement of solubility predicted by the equilibrium results (Figure 2C). The time to reach 95% of equilibrium mirrors this evolution, increasing from 4.32 ms (0 M) to 7.44 ms (0.5 M) before slightly decreasing to 3.72 ms (2.5 M). These results demonstrate that ionic strength modulates both the kinetics and apparent solubility of Mg(OH)₂ through thermodynamic activity effects captured by the SIT model, in agreement with the equilibrium solubility trends shown in Figure 2C.

3.2. Aqueous Carbonation Phenomena: Insights from Literature, SEM and XRPD

Table 2 presents the chemical reactions and key thermodynamic parameters involved in the aqueous carbonation of MgO, including the main intermediate and final carbonate phases relevant to this study. The most accepted mechanism leading to the release of Mg²⁺ from MgO in water occurs through a three-step process: (1) MgO reacts with water, forming surface MgOH⁺ groups and releasing OH⁻ into solution; (2) additional OH⁻ anions are adsorbed onto the positively charged MgOH⁺ surface; and (3) desorption of these OH⁻ groups leads to the dissolution of Mg²⁺ ions into the solution [53,54].

Upon CO₂ bubbling, dissolved CO₂ reacts with water to form carbonic acid (H₂CO₃), which partially dissociates into HCO₃⁻ and CO₃²⁻, thereby lowering the solution pH. The Mg²⁺ ions released from MgO hydration readily form transient Mg bicarbonate complexes or amorphous hydrated Mg carbonate (AMC) at low to moderate supersaturation [27]. These metastable precursors transform over time into more thermodynamically stable crystalline phases, most commonly Nes or Hmgs, depending on temperature, CO₂ partial pressure, and solution composition [55]. These phase transitions are governed by local supersaturation and kinetic barriers, and may be influenced by impurities, mixing regimes, or the presence of competing anions such as Cl⁻ and SO₄²⁻ [56,57,58].

While the mineralogy and morphology of Mg-carbonate products under varying operational parameters have been extensively studied in the literature [59] and are therefore beyond the scope of this work, Figure 4 highlights the two prevalent phases observed. Overall, the analyses confirmed that the carbonation process produced mixtures of hydrated Mg carbonates, either Nes, Hmgs or Dyp, whose speciation depended strongly on operational conditions, especially temperature and solution composition. SEM analysis (Figure 4) reveals that Nes formed as micro- and nano-sized crystals at temperatures of 323 K and below, whereas microsized rosette-like crystals of Hmgs predominated at higher temperatures, as confirmed by XRPD. These results are in agreement with established observations for Mg-carbonate precipitation [24,60,61]. This temperature boundary condition, occurring at 323 K, was previously reported by [5,57]. Consistent with the literature, spherical or plate-like aggregates were occasionally observed, indicating possible co-precipitation of Hmgs [6,53]. Notably, no Mgs (MgCO₃) peaks were observed, confirming that direct precipitation of the anhydrous carbonate is kinetically hindered below ≈363, likely due to the strong hydration shell around Mg²⁺ ions [9,62]. As highlighted by Oliver [46], the presence of metastable Mg(HCO₃)₂ species can delay immediate nucleation, promoting the initial formation of amorphous or poorly crystalline precursors. This was reflected in some runs, where broadened XRPD peaks suggested a fraction of poorly ordered carbonate phases. Overall, the insights from SEM and XRPD confirm that any engineered carbonation process using MgO must account for this variability to predict product stability and CO₂ uptake reliably.

Table 2. Chemical reactions and key thermodynamic variables involved in the aqueous carbonation of MgO, including the main intermediate and final carbonate phases. Note that the thermodynamic variables for the formation of common minerals in the MgO–CO₂–H₂O system at standard parameters have recently been reviewed by Santos [7], including also S°, $C_p^0$, and V_m.

Step/Phase Formation	Chemical Reaction	Log K° (298 K)/∆fG° (kJ/mol)/∆fH°	Comments	Reference
MgO hydration	$\mathrm{M}\mathrm{g}\mathrm{O}\left(\mathrm{s}\right)+{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right)={\mathrm{M}\mathrm{g}\left(\mathrm{O}\mathrm{H}\right)}_2\left(\mathrm{s}\right)$	Log K° = −11.16; ∆fG° = −832.2; ∆fH° = −924.66	Rapid at ambient conditions; forms poorly soluble Mg(OH)2	[7,53,63]
Mg(OH)2 dissolution	${\mathrm{M}\mathrm{g}\left(\mathrm{O}\mathrm{H}\right)}_2\left(\mathrm{s}\right)=\mathrm{M}{\mathrm{g}}^{2+}+2\mathrm{O}{\mathrm{H}}^{-}$	Log K° = 17.01	Controls Mg2+ availability in solution and pH	[64]
CO2 dissolution	${\mathrm{C}\mathrm{O}}_2\left(\mathrm{g}\right)+{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right)={{\mathrm{H}}_2\mathrm{C}\mathrm{O}}_3={\mathrm{H}}^{+}+\mathrm{H}\mathrm{C}{\mathrm{O}}_3^{-}=2{\mathrm{H}}^{+}+\mathrm{C}{\mathrm{O}}_3^{2-}$	pKa1 = 6.35; pKa2 = 10.33	CO2 forms carbonic acid; equilibrium shifts based on pH	-
Nes formation	$\mathrm{M}{\mathrm{g}}^{2+}+\mathrm{C}{\mathrm{O}}_3^{2-}+3{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right)={\mathrm{M}\mathrm{g}\mathrm{C}\mathrm{O}}_3·{3\mathrm{H}}_2\mathrm{O}\left(\mathrm{s}\right)$	Log K° = −5.27; ∆fG° = −1724.4; ∆fH° = 1981.7	Forms at T < 35 °C, pH 7–9; metastable	[7,12,65]
Hmgs formation	$5\mathrm{M}{\mathrm{g}}^{2+}+4\mathrm{C}{\mathrm{O}}_3^{2-}+2\mathrm{O}{\mathrm{H}}^{-}+4{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right)={\mathrm{M}{\mathrm{g}}_5(\mathrm{C}\mathrm{O}}_3{)}_4(\mathrm{O}\mathrm{H}{)}_2·{4\mathrm{H}}_2\mathrm{O}\left(\mathrm{s}\right)$	Log K° = − 37.08; ∆fG° = –5866.6; ∆fH° = −6514.9	Stable at pH > 9.5, common in long-term carbonation	[7,66,67]
Dyp formation	$5\mathrm{M}{\mathrm{g}}^{2+}+4\mathrm{C}{\mathrm{O}}_3^{2-}+2\mathrm{O}{\mathrm{H}}^{-}+5{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right)={\mathrm{M}{\mathrm{g}}_5(\mathrm{C}\mathrm{O}}_3{)}_4(\mathrm{O}\mathrm{H}{)}_2·{5\mathrm{H}}_2\mathrm{O}\left(\mathrm{s}\right)$	Log K° = −34.94; ∆fG° = −6081.7; ∆fH° = −6792.8	More hydrated than Hmgs; intermediate stability	[7,12]
Lfd formation	$\mathrm{M}{\mathrm{g}}^{2+}+\mathrm{C}{\mathrm{O}}_3^{2-}+5{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right)={\mathrm{M}\mathrm{g}\mathrm{C}\mathrm{O}}_3·{5\mathrm{H}}_2\mathrm{O}\left(\mathrm{s}\right)$	Log K° = −5.24; ∆fG° = −2197.8; ∆fH° = − 2574.3	Forms at T < 10 °C; unstable at room temperature	[7]
Mgs formation	$\mathrm{M}{\mathrm{g}}^{2+}+\mathrm{C}{\mathrm{O}}_3^{2-}={\mathrm{M}\mathrm{g}\mathrm{C}\mathrm{O}}_3\left(\mathrm{s}\right)$	Log K° = −8.29; ∆fG° = −1029.3; ∆fH° = −1112.9	Requires elevated temperature and/or pressure to crystallize	[7,68]

Table 2. Chemical reactions and key thermodynamic variables involved in the aqueous carbonation of MgO, including the main intermediate and final carbonate phases. Note that the thermodynamic variables for the formation of common minerals in the MgO–CO₂–H₂O system at standard parameters have recently been reviewed by Santos [7], including also S°, $C_p^0$, and V_m.

Step/Phase Formation	Chemical Reaction	Log K° (298 K)/∆fG° (kJ/mol)/∆fH°	Comments	Reference
MgO hydration	$\mathrm{M}\mathrm{g}\mathrm{O}\left(\mathrm{s}\right)+{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right)={\mathrm{M}\mathrm{g}\left(\mathrm{O}\mathrm{H}\right)}_2\left(\mathrm{s}\right)$	Log K° = −11.16; ∆fG° = −832.2; ∆fH° = −924.66	Rapid at ambient conditions; forms poorly soluble Mg(OH)2	[7,53,63]
Mg(OH)2 dissolution	${\mathrm{M}\mathrm{g}\left(\mathrm{O}\mathrm{H}\right)}_2\left(\mathrm{s}\right)=\mathrm{M}{\mathrm{g}}^{2+}+2\mathrm{O}{\mathrm{H}}^{-}$	Log K° = 17.01	Controls Mg2+ availability in solution and pH	[64]
CO2 dissolution	${\mathrm{C}\mathrm{O}}_2\left(\mathrm{g}\right)+{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right)={{\mathrm{H}}_2\mathrm{C}\mathrm{O}}_3={\mathrm{H}}^{+}+\mathrm{H}\mathrm{C}{\mathrm{O}}_3^{-}=2{\mathrm{H}}^{+}+\mathrm{C}{\mathrm{O}}_3^{2-}$	pKa1 = 6.35; pKa2 = 10.33	CO2 forms carbonic acid; equilibrium shifts based on pH	-
Nes formation	$\mathrm{M}{\mathrm{g}}^{2+}+\mathrm{C}{\mathrm{O}}_3^{2-}+3{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right)={\mathrm{M}\mathrm{g}\mathrm{C}\mathrm{O}}_3·{3\mathrm{H}}_2\mathrm{O}\left(\mathrm{s}\right)$	Log K° = −5.27; ∆fG° = −1724.4; ∆fH° = 1981.7	Forms at T < 35 °C, pH 7–9; metastable	[7,12,65]
Hmgs formation	$5\mathrm{M}{\mathrm{g}}^{2+}+4\mathrm{C}{\mathrm{O}}_3^{2-}+2\mathrm{O}{\mathrm{H}}^{-}+4{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right)={\mathrm{M}{\mathrm{g}}_5(\mathrm{C}\mathrm{O}}_3{)}_4(\mathrm{O}\mathrm{H}{)}_2·{4\mathrm{H}}_2\mathrm{O}\left(\mathrm{s}\right)$	Log K° = − 37.08; ∆fG° = –5866.6; ∆fH° = −6514.9	Stable at pH > 9.5, common in long-term carbonation	[7,66,67]
Dyp formation	$5\mathrm{M}{\mathrm{g}}^{2+}+4\mathrm{C}{\mathrm{O}}_3^{2-}+2\mathrm{O}{\mathrm{H}}^{-}+5{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right)={\mathrm{M}{\mathrm{g}}_5(\mathrm{C}\mathrm{O}}_3{)}_4(\mathrm{O}\mathrm{H}{)}_2·{5\mathrm{H}}_2\mathrm{O}\left(\mathrm{s}\right)$	Log K° = −34.94; ∆fG° = −6081.7; ∆fH° = −6792.8	More hydrated than Hmgs; intermediate stability	[7,12]
Lfd formation	$\mathrm{M}{\mathrm{g}}^{2+}+\mathrm{C}{\mathrm{O}}_3^{2-}+5{\mathrm{H}}_2\mathrm{O}\left(\mathrm{l}\right)={\mathrm{M}\mathrm{g}\mathrm{C}\mathrm{O}}_3·{5\mathrm{H}}_2\mathrm{O}\left(\mathrm{s}\right)$	Log K° = −5.24; ∆fG° = −2197.8; ∆fH° = − 2574.3	Forms at T < 10 °C; unstable at room temperature	[7]
Mgs formation	$\mathrm{M}{\mathrm{g}}^{2+}+\mathrm{C}{\mathrm{O}}_3^{2-}={\mathrm{M}\mathrm{g}\mathrm{C}\mathrm{O}}_3\left(\mathrm{s}\right)$	Log K° = −8.29; ∆fG° = −1029.3; ∆fH° = −1112.9	Requires elevated temperature and/or pressure to crystallize	[7,68]

3.3. CO₂ Capture Rate vs. Time $\eta \left(t\right)$: General Insights

The temporal evolution of CO₂ capture efficiency, expressed as $\eta \left(t\right)$, revealed distinct absorption patterns that align with typical multi-stage gas–liquid reaction kinetics for alkaline slurries [69,70,71,72,73,74].

Figure 5 (Experiment A1) shows that the EC is very low at the beginning of the tests, with 0.489 mS/cm and a pH of 11.28 measured at t = 0, consistent with the low solubility of MgO in water. Figure 5A shows the CO₂ absorption rate vs. time pattern $\eta \left(t\right)$ in four distinct rate-controlling steps (hereafter denoted as step 1: initial rise; step 2: maximum absorption; step 3: downfall; step 4: final decline). Throughout these steps, the pH decreases almost linearly until stabilizing at 7.15 at the conclusion of Step 3. During step 1, $\eta \left(t\right)$ rapidly rises to nearly 67%, indicating that the chemical reaction is initially much faster than CO₂ dissolution in water. It then gradually decreases to about 53% during step 2, before falling to around 15% in step 3. Thus, across all runs, the $\eta \left(t\right)$ curves exhibited an initial rapid uptake phase, followed by a progressive slowdown as the system approached equilibrium. This behavior is consistent with a transition from reaction-controlled to diffusion-controlled carbonation, as reported in comparable MgO slurry studies [5,75].

Figure 5 (Experiment A2) illustrates that the TA attributed to carbonates reaches a peak value of 3740 mg CaCO₃/L before declining to 0, whereas the bicarbonate TA increases and stabilizes at approximately 23,000 mg CaCO₃/L. This transition occurs as the pH decreases into the bicarbonate stability range. Comparing the two graphs in Figure 5, the observed pH correlations indicate that the decrease in TA attributed to carbonates coincides with the onset of the step 2 $\eta \left(t\right)$ decline. This suggests that the decline in step 2 is primarily driven by the conversion of CO₃²⁻ to HCO₃⁻ in the solution. This conversion reduces the availability of CO₃²⁻ for further reactions, thereby inhibiting CO₂ dissolution and subsequently slowing its absorption rate in the suspension. The Step 4 $\eta \left(t\right)$ decline coincides with the stabilization of both EC and bicarbonate levels, while the pH reaches its steady plateau at 7.15.

3.4. Impact of Operational Parameters on CO₂ Absorption During MgO Aqueous Carbonation

3.4.1. Impact of MgO Concentration and CO₂ Flow Rate

Figure 6A,B illustrate the effects of MgO initial concentration (B1–10 experiments) and CO₂ flow rate (C1–5 experiments), respectively, on the CO₂ absorption vs. time patterns $\eta \left(t\right)$.

Table 3. Average CO₂ absorption rate ($\overline{\eta }\left(t\right)$), maximum absorption rate (${\eta }_{max}\left(t\right)$), duration and volume of absorbed CO₂ ($V_{CO2}$) in the B1–10 experiments.

Exp. ID.	[MgO]0	$\overline{\mathit{\eta }}\left(\mathit{t}\right)$	${\mathit{\eta }}_{\mathit{m}\mathit{a}\mathit{x}}\left(\mathit{t}\right)$	σ	t	${\mathit{V}}_{\mathit{C}\mathit{O}2}$
	wt.%	%	%	-	s	L
B1	0.5	34.3	59.7	22.5	2450	7.0
B2	1.0	35.3	66.1	23.5	4253	12.5
B3	1.5	33.1	73.9	27.1	6155	17.0
B4	2.0	32.1	65.5	24.2	7432	19.9
B5	2.5	30.5	66.3	24.2	8798	22.4
B6	3.0	28.9	68.6	23.9	10,285	24.8
B7	3.5	31.3	68.6	23.2	10,405	27.2
B8	4.0	33.0	67.8	25.3	10,826	29.8
B9	4.5	34.2	65.9	24.4	10,766	30.7
B10	5.0	35.0	64.9	21.5	11,502	33.5

reports the calculations of the average CO₂ absorption rate ($\overline{\eta }\left(t\right)$), maximum absorption rate (${\eta }_{max}\left(t\right)$), duration and volume of absorbed CO₂ ($V_{CO2}$) in the B1–10 experiments. Figure 6A shows that, despite altering the appearance of $\eta \left(t\right)$, increasing the initial MgO concentration from 0.5 to 5 wt.% does not lead to a significant change in the CO₂ absorption rate intensity despite the higher reactive surface area. A superior ${\eta }_{max}\left(t\right)$ value, approaching 74%, was achieved only at an initial MgO concentration of 1.5 wt.%. However, calculations of total volumetric CO₂ uptake ($V_{CO2}$) reveal diminishing returns when MgO concentrations exceed 1.5–2 wt.%. At lower concentrations, most Mg forms Mg-bicarbonates with a Mg:CO₂ ratio of 1:2, leading to higher CO₂ uptake. For instance, the 1 wt.% MgO experiment absorbed 12.5 L of CO₂, whereas the 5 wt.% experiment absorbed 33.5 L, demonstrating that increasing MgO concentrations beyond 2 wt.% results in a proportionally lower $V_{CO2}$. At higher solid concentrations, the benefit of increased reactive sites is partly offset by enhanced slurry viscosity, which can limit gas–liquid contact and mass transfer efficiency. This behavior is consistent with the kinetic constraints described by [53] who reported that dense MgO suspensions may exhibit local CO₂ undersaturation due to hindered bubble dispersion and reduced mixing. Thus, an optimal range of MgO loading must be defined to balance solid availability and slurry flowability.

On the other hand, the CO₂ flow rate also played a critical role in controlling the absorption dynamics. As expected, Figure 6B shows that increasing the CO₂ flow rate decreases the CO₂ absorption rate and reduces the carbonation time, in agreement with previous observations in both MgO and Ca(OH)₂ aqueous carbonation systems [76]. The ${\eta }_{max}\left(t\right)$ value was 59.5 at a CO₂ flow rate of 0.5 L/min and decreased to 32.2 at 2 L/min. Nevertheless, when the flow rate exceeded the absorption capacity of the suspension, the excess gas passed through the reactor unreacted, resulting in diminished overall capture efficiency. Similar behavior has been noted in Mg looping scenarios, where the residence time of CO₂ bubbles must be carefully matched to the reaction kinetics [40].

Overall, these results highlight the need to co-optimize MgO concentration and CO₂ flow rate in scaled aqueous carbonation systems. Excessive solid content or over-supply of CO₂ can reduce system efficiency due to kinetic limitations and phase separation constraints. Identifying these operational trade-offs is crucial for designing industrial carbonation scrubbers or mineralization reactors using Mg-rich feedstocks.

3.4.2. Impact of Temperature and NaCl Concentration

Figure 7 shows the $\eta \left(t\right)$ for experiments D1–5, D6–10, and D11–15, as well as pH and EC measurements for the pure MgO-H₂O-CO₂ system at 303 K.

Table 4. Average CO₂ absorption rate ($\overline{\eta }\left(t\right)$), maximum absorption rate (${\eta }_{max}\left(t\right)$), duration and volume of absorbed CO₂ ($V_{CO2}$) in the D1–15 experiments.

Exp. ID.	[NaCl]	Temperature	$\overline{\mathit{\eta }}\left(\mathit{t}\right)$	${\mathit{\eta }}_{\mathit{m}\mathit{a}\mathit{x}}\left(\mathit{t}\right)$	σ	t	${\mathit{V}}_{\mathit{C}\mathit{O}2}$
	wt.%	K	%	%	-	s	L
D1	0	283	47.63	74.83	27.25	3118	12.38
D2	0	303	39.13	81.52	31.61	3155	10.29
D3	0	323	36.93	76.52	26.01	2500	7.69
D4	0	343	22.58	41.09	13.1	3000	5.65
D5	0	363	15.7	25.63	6.65	2953	3.87
D6	3.5	283	45.40	81.39	26.35	3545	13.41
D7	3.5	303	33.76	73.98	27.95	3880	10.92
D8	3.5	323	30.78	64.04	24.64	3419	8.77
D9	3.5	343	31.88	65.46	22.09	2644	7.02
D10	3.5	363	19.87	48.36	13.00	2468	4.09
D11	7	283	37.45	64.57	13.73	4400	13.73
D12	7	303	34.67	65.60	22.75	4000	11.56
D13	7	323	33.66	60.87	21.01	3500	9.82
D14	7	343	28.97	64.39	22.58	2850	6.88
D15	7	363	21.67	41.44	15.10	2200	3.97

reports the average CO₂ absorption rate ($\overline{\eta }\left(t\right)$), maximum absorption rate (${\eta }_{max}\left(t\right)$), duration and volume of absorbed CO₂ ($V_{CO2}$) in the D1–15 experiments. Temperature significantly affected MgO carbonation, influencing speciation, crystal morphology, CO₂ absorption rate, and uptake. The lower the temperatures, the higher the $\eta \left(t\right)$ and $V_{CO2}$ values. Experiments at 283 K absorbed more than three times the $V_{CO2}$ compared to those at 363 K. At 283 K, the measured $V_{CO2}$ values were 12.4 L, 13.4 L, and 13.7 L depending on NaCl concentration, while at 363 K they dropped to 3.9 L, 4.1 L, and 4.0 L, respectively. A 10 K reduction in temperature led to approximately 1 L less CO₂ uptake. The CO₂ absorption was two to three times faster in experiments at 283 K compared to those at 363 K. $\overline{\eta }\left(t\right)$ values were, on average, 43.5% at 283 K against 19.1% at 363 K. The ${\eta }_{max}\left(t\right)$ values were, on average, 73.6% at 283 K against 38.5% at 363 K. At 283 K (D1, D6, and D11), a limpid solution indicated complete reactant dissolution and Mg(HCO₃)₂ formation, whereas higher temperatures produced a white slurry of Mg carbonates. At temperatures exceeding 303 K, following carbonate precipitation, the suspension exhibited a flocculant texture, attributed to the aggregation of fine MgO particles and precipitates.

In this study, NaCl concentration did not appear to affect precipitated carbonates speciation, though a slight increase in CO₂ uptake was measured. Although NaCl addition generally lowers the physical solubility of CO₂ in water, the observed increase in CO₂ uptake likely reflects ionic-strength effects that enhance MgO dissolution and stabilize bicarbonate complexes (HCO₃⁻, MgHCO₃⁺). This promotes greater chemical fixation of CO₂ despite its slightly reduced solubility. However, a significant increase in $\eta \left(t\right)$ values, particularly at higher temperatures (from 343 K), was observed. Indeed, $\overline{\eta }\left(t\right)$ and ${\eta }_{max}\left(t\right)$ values were 22.6% and 41.1% at 343 K in ultrapure water. In the presence of 3.5% NaCl, these values increased to 31.9% and 65.5%, respectively. With 7 wt.% NaCl, $\overline{\eta }\left(t\right)$ was 29%, and ${\eta }_{max}\left(t\right)$ reached 64.4%. A similar trend was detected at 363 K.

Practically, these results indicate that using seawater or industrial brines with significant Cl⁻ content could shift the optimum operating temperature for MgO carbonation. At moderate salinity and temperatures exceeding ≈323 K, the CO₂ absorption rate surpasses that in pure water.

3.4.3. Impact of Na₂SO₄ and K₂SO₄ Concentration

It is known that Cl⁻ and SO₄²⁻ anions affect both the kinetics and thermodynamics of MgO hydration. According to Amaral [53], these ions modify ionic strength and common-ion equilibria, which can either retard or accelerate MgO dissolution and brucite formation depending on their concentration. Such effects may therefore influence the CO₂ absorption rate measured in this study.

Figure 8A,B show the impacts of Na₂SO₄ (E1–4 experiments) and K₂SO₄ (F1–4 experiments) concentrations on $\eta \left(t\right)$.

Table 5. Average CO₂ absorption rate ($\overline{\eta }\left(t\right)$), maximum absorption rate (${\eta }_{max}\left(t\right)$), duration and volume of absorbed CO₂ ($V_{CO2}$) in the E1-E4 and F1-F4 experiments.

Exp. ID.	[Na/-K2SO4]	$\overline{\mathit{\eta }}\left(\mathit{t}\right)$	${\mathit{\eta }}_{\mathit{m}\mathit{a}\mathit{x}}\left(\mathit{t}\right)$	σ	t	${\mathit{V}}_{\mathit{C}\mathit{O}2}$
	wt.%	%	%	-	s	L
E1	1	27.4	63.6	24.0	3490	8.0
E2	3.5	26.3	57.2	21.5	4100	9.0
E3	7	26.2	53.7	20.5	4280	9.3
E4	10.5	23.5	42.7	12.6	4320	8.5
F1	1	29.4	64.0	24.9	2970	7.3
F2	3.5	26.6	59.8	22.6	3415	7.6
F3	7	25.5	54.3	19.1	3650	7.7
F4	10.5	25.5	53.4	19.6	4100	8.7

reports the average CO₂ absorption rate ($\overline{\eta }\left(t\right)$), maximum absorption rate (${\eta }_{max}\left(t\right)$), duration and volume of absorbed CO₂ ($V_{CO2}$) in the E1-E4 and F1-F4 experiments. The dotted $\eta \left(t\right)$ represents the baseline: the pure H₂O–MgO–CO₂ system. Both salts yield a slight increase in $\eta \left(t\right)$ at a low concentration, 1 wt.%, while higher concentrations result in a steadily decreasing $\eta \left(t\right)$. Thus, the ${\eta }_{max}\left(t\right)$ values approached 64% at 1 wt.% salt concentration, while they were only 42.7% and 53.4% at 10.5 wt.% for Na₂SO₄ and K₂SO₄, respectively. Interestingly, the results show that the volumetric CO₂ uptake $V_{CO2}$ increases with higher Na₂SO₄ and K₂SO₄ concentrations. The $V_{CO2}$ value for the baseline, without any salt, was 7.2 L, increasing to 9.3 L at 7 wt.% Na₂SO₄ and 8.7 L at 10.5 wt.% K₂SO₄. This suggests that sulfates significantly enhance the solubility of Mg-bicarbonates, promoting the formation of metastable solutions with higher concentrations.

3.4.4. Impact of Mixing System

It is known that the uptake of CO₂ by metal (hydr)oxide suspension increased with higher turbulence due to better gas/water exchange [5]. This principle has been confirmed for both Ca(OH)₂ and MgO aqueous carbonation systems, as well as for various alkaline waste suspensions [26,30].

Figure 9 depicts the impact of using a regular pipe bubbling gas–liquid mixing system (G1 experiment) instead of a gas sparger on $\eta \left(t\right)$. The G1 experiment results in a very low $\overline{\eta }\left(t\right)$ value of 8.6%, a $V_{CO2}$ of 8.6 L, and an ${\eta }_{max}\left(t\right)$ of 16.4%. These results demonstrate that employing a sparger, which produces finer bubbles and improves gas dispersion, significantly enhances the kinetics of MgO aqueous carbonation.

This behavior is consistent with the findings of [24], who showed that improved mixing regimes sustain higher local CO₂ concentrations and delay local equilibrium saturation, thus extending the period of metastable bicarbonate stabilization. In practical terms, the increased interfacial area achieved with spargers or porous diffusers accelerates CO₂ dissolution and its reaction with Mg²⁺. This helps overcome the inherently slower hydration–dehydration dynamics of Mg compared with calcium systems [5].

In summary, these findings confirm that effective mixing system design is essential for maximizing CO₂ uptake in MgO aqueous carbonation systems, supporting both higher capture rates and improved utilization of metastable solution pathways.

3.5. Comparing Pure CaO- and MgO-Based Mineral Carbonation Systems: Impact of Operational Parameters During Aqueous Carbonation

The comparative analysis of pure free CaO- and MgO-based aqueous carbonation systems highlights distinct operational constraints and opportunities for process design. Reference [5] previously compared CaO and MgO aqueous carbonation systems, reporting a higher pH (>11.7) in Ca-based systems compared to Mg (<10.3), which can be attributed to the lower solubility of MgO relative to CaO. Table 6 compares the effects of operational parameters on carbonation performance in aqueous free CaO and MgO systems, providing practical guidelines based on literature, our previous work, and the present study. It is important to note that the higher BET surface area of MgO (≈31.9 m²/g) compared with Ca(OH)₂ (≈16.0 m²/g) significantly favors higher CO₂ absorption rates in MgO suspensions. A larger specific surface area enhances gas–liquid–solid contact, accelerating MgO hydration and dissolution as well as early CO₂ uptake, though its effect on total uptake is less significant once equilibrium is reached.

In this section, a comparison with our previous work [47] highlights the main differences between Ca- and Mg-(hydr)oxide carbonation systems:

3.5.1. Impact of Mixing System on $\eta \left(t\right)$

As expected, MgO carbonation is generally slower than Ca(OH)₂, as shown by comparing the $\eta \left(t\right)$ patterns, particularly when the mixing system is not optimized, i.e., when CO₂ is bubbled through a regular pipe. However, surprisingly, our results suggest that using a CO₂ sparger at low temperatures (≤303 K) in a pure system may yield similar or slightly higher $\eta \left(t\right)$ values in MgO carbonation than in Ca(OH)₂ carbonation. Specifically, at very close operational conditions (similar pressure, PCO₂, CO₂ flow rate, CO₂ vol.%, mixing speed), experiment SPR1 [47] using Ca(OH)₂ yielded $\eta \left(t\right)$ values reaching about 65%, while experiments D2 and B3 using MgO yielded values exceeding 70–80% in this study.

Table A2. Comparison of Ca(OH)₂ and MgO aqueous carbonation performances under optimized conditions of low temperature, sparger-enhanced gas–liquid mixing, and pure Ca(OH)₂/MgO-H₂O-CO₂ system.

Experiment Identifier	SSA	[Reagent]	Temperature	Constants	$\overline{\mathit{\eta }}\left(\mathit{t}\right)$	${\mathit{\eta }}_{\mathit{m}\mathit{a}\mathit{x}}\left(\mathit{t}\right)$	${\mathit{V}}_{\mathit{C}\mathit{O}2}$
	m2/g	wt.%	K	-	%	%	L
SPR1	16	1.5 [Ca(OH)2]	303	1.2 kgw, CO2 0.5 L/min, PCO2 1.2 bar, CO2 100 vol.%, sparger, stirrer 300 rpm, pure system (no NaCl, no salts)	37.9	64.8	3.73
B3	31.9	1.5 [MgO]	278		33.1	73.9	17
D2	31.9	1 [MgO]	303		39.13	81.52	10.29

in Appendix D reports the exact operational conditions and the corresponding carbonation performances related to this observation.

The higher CO₂ absorption rates $\eta \left(t\right)$ observed for MgO under sparger-assisted, low-temperature conditions can be explained by the strong enhancement of gas–liquid transfer kinetics and the thermodynamic favorability of CO₂ dissolution at low temperature. The sparger-generated microbubbles greatly accelerate the cascade of dissolution–aqueous complexation reactions, promoting rapid conversion of CO₂ into hydrated and bicarbonate species. Simultaneously, the continuous hydration of MgO maintains alkalinity and sustains the chemical driving force for CO₂ absorption. In contrast, Ca(OH)₂ systems consume dissolved CO₂ more rapidly and locally equilibrate, benefiting less from enhanced gas–liquid mixing. This combination of efficient microbubble-driven dissolution dynamics, higher aqueous CO₂ availability, and sustained alkalinity explains why MgO systems can transiently exhibit higher $\eta \left(t\right)$ values than Ca(OH)₂ under identical operating conditions.

3.5.2. Impact of NaCl, Na₂SO₄ and K₂SO₄ on $\eta \left(t\right)$

While seawater-level NaCl (3.5 wt.%) markedly enhances $\eta \left(t\right)$ values in Ca(OH)₂ carbonation, its enhancing influence in MgO carbonation was only significant at higher temperatures (>323 K), particularly between 323 and 363 K, with slight decreases observed at lower temperatures. The opposite trend observed here compared with Ca(OH)₂ systems indicates that the presence of common anions (Cl⁻, SO₄²⁻) does not always promote CO₂ hydration and absorption. According to the framework proposed by Dennard [77], oxyanions can influence CO₂ hydration through direct interactions with the CO₂ molecule or by interacting with neighboring water molecules that coordinate CO₂, thereby modifying its hydration kinetics. In Mg-based suspensions, however, such molecular-scale effects appear to be slightly outweighed by the ionic-strength-induced decrease in CO₂ solubility and the reduced Mg²⁺ activity in solution, resulting overall in a lower CO₂ absorption rate $\eta \left(t\right)$ in the presence of high concentrations of salts.

Table 6. Comparison of operational parameter effects on carbonation performance in aqueous free CaO and MgO systems, with practical guidelines.

Parameter/Effect	CaO System	MgO System	Reference
Typical pH	>11.7	<11.3	[5]; this study
Precipitated phases	Cal is the predominant phase, being thermodynamically stable, whereas Vat, the intermediate between amorphous CaCO3 and Cal, can persist at ≈30 °C.	Nes, Hmgs, and Dyp predominate under near-ambient P/T conditions, whereas Mgs dominates at higher P/T.	[78,79]
CO2 absorption rate $\eta \left(t\right)$	Faster than MgO under most equivalent thermo-hydro-mechanical-chemical conditions.	Generally slower, but can match or slightly exceed CaO at low temperatures with enhanced gas–solid–liquid mixing.	This study; [47]
Effect of CO2 partial pressure	Rapid carbonation at low pCO2 due to high solubility; limited further enhancement beyond saturation	Strongly dependent on pCO2; higher pCO2 accelerates Mg2+ release and favors stable carbonate precipitation such as Mgs	[70,80,81,82]
Effect of Temperature	Stabilizes Cal, increases carbonation kinetics up to ≈343 K and CO2 uptake	Precipitation of less hydrated and more stable carbonate phases, steep decrease in carbonation kinetics and CO2 uptake with temperature due to lower CO2 solubility and bicarbonate formation. However, Mg-silicate carbonation is dissolution-controlled and requires high temperatures.	This study; [47,83]
Chloride impurities	Strong, near-optimal enhancement at 3.5 wt.% (seawater level) at ambient temperature	Only significant at >323 K	This study; [47,84,85]
Sulfate impurities	Among the tested salts (NaCl, KCl, K2SO4, Na2SO4), Na2SO4 was the most efficient, enhancing the CO2 absorption rate by up to 75%. K2SO4 was also highly effective at low concentrations (<2 wt.%)	Slightly enhance the absorption rate at low concentrations (<2 wt.%) but reduce it at higher levels, with no significant impact on the CO2 uptake	[47,86]; this study; [87]
Effect of mixing system	A gas–liquid spargers with an impeller generate microbubbles and well-dispersed suspensions, enhancing CO2 absorption.	Optimizing the mixing system is even more critical than for CaO, due to slower dissolution–precipitation kinetics.	[76,88,89,90,91]
Metastable bicarbonate formation	Negligible due to the very low concentration and limited persistence of Ca(HCO3)2 in solution.	Mg(HCO3)2 forms at <303 K, increasing CO2 uptake up to 2× that of the CaO system.	[45,92]; This study; [47]
Operational/fouling behavior	Heavy deposit formation; frequent clogging of spargers and porous tools.	Fewer reactor deposits; slurries harden at flask bottom, requiring direct powder addition.	This study; [47]

Table 6. Comparison of operational parameter effects on carbonation performance in aqueous free CaO and MgO systems, with practical guidelines.

Parameter/Effect	CaO System	MgO System	Reference
Typical pH	>11.7	<11.3	[5]; this study
Precipitated phases	Cal is the predominant phase, being thermodynamically stable, whereas Vat, the intermediate between amorphous CaCO3 and Cal, can persist at ≈30 °C.	Nes, Hmgs, and Dyp predominate under near-ambient P/T conditions, whereas Mgs dominates at higher P/T.	[78,79]
CO2 absorption rate $\eta \left(t\right)$	Faster than MgO under most equivalent thermo-hydro-mechanical-chemical conditions.	Generally slower, but can match or slightly exceed CaO at low temperatures with enhanced gas–solid–liquid mixing.	This study; [47]
Effect of CO2 partial pressure	Rapid carbonation at low pCO2 due to high solubility; limited further enhancement beyond saturation	Strongly dependent on pCO2; higher pCO2 accelerates Mg2+ release and favors stable carbonate precipitation such as Mgs	[70,80,81,82]
Effect of Temperature	Stabilizes Cal, increases carbonation kinetics up to ≈343 K and CO2 uptake	Precipitation of less hydrated and more stable carbonate phases, steep decrease in carbonation kinetics and CO2 uptake with temperature due to lower CO2 solubility and bicarbonate formation. However, Mg-silicate carbonation is dissolution-controlled and requires high temperatures.	This study; [47,83]
Chloride impurities	Strong, near-optimal enhancement at 3.5 wt.% (seawater level) at ambient temperature	Only significant at >323 K	This study; [47,84,85]
Sulfate impurities	Among the tested salts (NaCl, KCl, K2SO4, Na2SO4), Na2SO4 was the most efficient, enhancing the CO2 absorption rate by up to 75%. K2SO4 was also highly effective at low concentrations (<2 wt.%)	Slightly enhance the absorption rate at low concentrations (<2 wt.%) but reduce it at higher levels, with no significant impact on the CO2 uptake	[47,86]; this study; [87]
Effect of mixing system	A gas–liquid spargers with an impeller generate microbubbles and well-dispersed suspensions, enhancing CO2 absorption.	Optimizing the mixing system is even more critical than for CaO, due to slower dissolution–precipitation kinetics.	[76,88,89,90,91]
Metastable bicarbonate formation	Negligible due to the very low concentration and limited persistence of Ca(HCO3)2 in solution.	Mg(HCO3)2 forms at <303 K, increasing CO2 uptake up to 2× that of the CaO system.	[45,92]; This study; [47]
Operational/fouling behavior	Heavy deposit formation; frequent clogging of spargers and porous tools.	Fewer reactor deposits; slurries harden at flask bottom, requiring direct powder addition.	This study; [47]

3.5.3. Impact of Temperature on $\eta \left(t\right)$

Increasing temperature led to a steep decrease in both $\eta \left(t\right)$ and CO₂ uptake ($V_{CO2}$) values in MgO carbonation, whereas our previous study shows a slight improvement in the Ca(OH)₂ system [47]. Finally, working at mild temperatures (≤323 K) results in far fewer deposits and reduced corrosivity on reactor surfaces and mixing system tools compared to Ca(OH)₂, which tends to clog porous mixing tools, spargers, and diffusors. This operational advantage could be significant for designing scrubbers or loop reactors where ease of maintenance is critical.

3.5.4. Formation of Metastable Bicarbonate

Working at ambient to low temperatures (<303 K) yielded a metastable Mg(HCO₃)₂ solution, resulting in ($V_{CO2}$) values more than double those in the Ca(OH)₂ system. This mechanism, rooted in the unique aqueous speciation behavior of Mg²⁺, has no direct analog in Ca-based systems where bicarbonate formation is minimal due to CaCO₃’s low solubility product [45].

4. Conclusions

In conclusion, this study provides a comprehensive assessment of the factors controlling CO₂ absorption $\eta \left(t\right)$ and total uptake ($V_{CO2}$), during the aqueous carbonation of MgO suspensions. By systematically varying temperature, salinity, CO₂ flow rate, and mixing system design, we demonstrate how each operational parameter shapes the carbonation performances.

Our results confirm that temperature is a critical factor, with increasing temperatures above ≈303 K leading to a steep decline in both $\eta \left(t\right)$ and CO₂ uptake ($V_{CO2}$). This trend reflects the destabilization of metastable Mg(HCO₃)₂ solutions and the faster precipitation of hydrated carbonates, which limit the time window for effective CO₂ absorption.

Furthermore, the introduction of NaCl at 3.5 and 7 wt.% provides a slight improvement at high temperatures (>323 K), particularly between 323 and 363 K. Sulfate ions also influenced the system by altering nucleation kinetics and metastable bicarbonate stability, highlighting the importance of solution chemistry in real brine or seawater scenarios.

Additionally, although Ca(OH)₂ carbonation is generally faster than that of MgO, the results show that using a CO₂ sparger at low temperatures (≤303 K) in a pure system can lead to higher $\eta \left(t\right)$ values for MgO carbonation than for Ca(OH)₂, despite nearly identical operational conditions. Under these specific conditions, $\eta \left(t\right)$ reached up to 70–80%, while Ca(OH)₂ achieved only about 65%. This confirms that enhancing MgO carbonation through a combination of static and dynamic mixing systems is highly effective, enabling $\eta \left(t\right)$ values significantly higher than typically expected.

Importantly, the use of an optimized gas sparger and static–dynamic mixing system combination proved highly effective for enhancing gas–liquid mass transfer. Under low-temperature conditions (<303 K), this configuration yielded $\eta \left(t\right)$ values for MgO carbonation that, in some cases, exceeded those for Ca(OH)₂ under equivalent settings. This finding highlights the unique potential of metastable bicarbonate formation in Mg-based systems, a mechanism playing only a minor role in Ca-based carbonation where precipitation occurs directly as CaCO₃.

Compared with our previous work on Ca(OH)₂ and the relevant literature, this study shows that MgO aqueous carbonation is inherently limited by the strong hydration shell around Mg²⁺ and its low mineral solubility. However, these kinetic barriers can be mitigated through careful control of operating conditions, allowing rapid CO₂ absorption and substantial uptake.

Furthermore, the experiments were performed under controlled laboratory conditions using pure MgO–H₂O–CO₂ systems, which simplifies the complexity of real industrial or brine environments. The interpretation of absorption rate and uptake is also constrained by the absence of direct speciation monitoring, such as time-resolved DIC partitioning or solid-phase quantification. Moreover, the batch configuration used here does not fully capture the hydrodynamic and mass transfer behavior expected in continuous or scaled-up systems. Addressing these limitations will require coupling kinetic experiments with in situ geochemical monitoring and pilot-scale validation under realistic water chemistries.

Overall, this work closes a key gap in understanding how metastable speciation, temperature, salinity, and mixing regimes control CO₂ absorption rate and uptake in MgO aqueous carbonation systems. It provides new quantitative benchmarks for optimizing these parameters under realistic operating conditions, establishing a practical basis for designing and scaling up Mg-based aqueous carbonation strategies that use natural, industrial, or brine-derived feedstocks. Particular attention should be given to heat and mass-transfer limitations, gas-dispersion efficiency, and the stabilization of metastable bicarbonate phases while preventing CO₂ degassing.

Future research should explore continuous flow or pilot-scale systems, integrate real brine or seawater streams, and investigate how biological or chemical additives might further stabilize metastable bicarbonate pathways, thereby enhancing MgO’s practical role in large-scale CO₂ mineralization and carbon removal technologies.