1 Introduction
There is scientific consensus that avoiding the most extreme effects of climate change will not only require decarbonization, but also the removal of carbon dioxide (CO2) from the atmosphere. Most scenarios that prevent more than 2 °C of warming require around 10-20 GtCO2/yr. of carbon dioxide removal (CDR) by 2,100 (IPCC, 2022; Rogelj et al., 2018). The scale of the problem means that an adoption of multiple methods will likely be required. Ocean alkalinity enhancement (OAE) refers to a promising set of CDR methods that leverage the natural role of the oceans to remove CO2 from the air (Renforth and Henderson, 2017; Committee on A Research Strategy for Ocean-based Carbon Dioxide Removal and Sequestration, Ocean Studies Board, Division on Earth and Life Studies, & National Academies of Sciences, Engineering, and Medicine, 2022). Alkalinity is a measure of the capacity of seawater to neutralize acids, in particular carbonic acid, which forms when CO2 reacts with water (Wolf-Gladrow et al., 2007). OAE methods increase the alkalinity of seawater, pulling additional CO2 from the atmosphere into the ocean, where it is stably stored as bicarbonate ions (HCO3−). As the ocean already contains a large inventory of bicarbonate (37,200 GtC, Keppler et al., 2020), the additional bicarbonate is not expected to appreciably change the ecological environment and is estimated to have a durability of 10,000 years (Renforth and Henderson, 2017) to 100,000 years (Cai et al., 2008). Multiple alkalinity sources have been proposed, including CaO (Kheshgi, 1995), Mg(OH)2 (Davies et al., 2018; Yang et al., 2023), crushed mafic and ultramafic rocks (Hangx and Spiers, 2009; Köhler et al., 2010; Meysman and Montserrat, 2017; Montserrat et al., 2017; Rigopoulos et al., 2018; and Schuiling and de Boer, 2011), limestone-derived Ca(OH)2 via electrochemistry (Rau, 2009), and aqueous NaOH generated electrochemically from aqueous brine streams (Eisaman et al., 2023; House et al., 2007; Schiffer et al. 2024).
Electrochemically generated aqueous NaOH is a promising approach because it avoids uncertainties in the dissolution rates of mined minerals (Hangx and Spiers, 2009; Montserrat et al., 2017) and also avoids environmental concerns regarding trace impurities released from geologically sourced alkalinity (Bach et al., 2019). However, aqueous alkalinity is gravimetrically and volumetrically less dense (0.5-1 mol(alk)/kg) than solid sources of alkalinity (20-25 mol(alk)/kg), so the costs and emissions of alkalinity transport tend to restrict NaOH alkalinity dispersal to coastal locations rather than the offshore ship-based dispersal that is possible with solid forms of alkalinity. The rate of alkalinity dispersal from any single location will be limited by permitted bounds of key metrics such as maximum allowable pH (He and Tyka, 2023). As an example, the addition of alkalinity in the surface ocean locally and temporarily increases seawater pH, with the pH increase being maximal at the point of introduction within the mixing zone and then decreasing as the added alkalinity mixes with untreated ocean water. OAE practitioners also aim to keep the carbonate chemistry in a range that avoids runaway precipitation of calcium carbonate (CaCO3; Moras et al., 2022; Fuhr et al., 2021; Suitner et al., 2024), which would tend to reverse the goal of OAE by removing net alkalinity (Zeebe and Wolf-Gladrow, 2001). The maximum rate of alkalinity addition at a given location will then be determined by the dissolution kinetics of mineral sources and the rate at which alkalinity in the mixing zone is diluted by the surrounding ocean. With this limitation, the maximum rate of OAE is expected to scale with the area over which the alkalinity can be added, and therefore also scale with the transportability of the alkalinity source.
Here an approach called “alkalinity exchange” is proposed that will allow OAE employing aqueous alkalinity to continue to grow in scale even once it starts to run up against practical limitations imposed by the need to avoid extreme chemical changes. The general concept of alkalinity exchange is to take advantage of the presence of divalent cations, like Mg2+ and Ca2+ in seawater and waste brines such as reverse osmosis concentrate (ROC), to precipitate alkalinity-containing solids, such as Mg(OH)2, upon addition of aqueous NaOH. This effectively exchanges one form of alkalinity (dilute aqueous NaOH) for another form (concentrated solid Mg(OH)2) that is more gravimetrically and volumetrically dense: from a molecular point of view, the hydroxide bound in the solid Mg(OH)2 is the same hydroxide that was added as NaOH. The Mg(OH)2 can then be economically dispersed further out at sea via ship, helping to overcome limitations on the rate at which alkalinity can be dispersed near the coast (He and Tyka, 2023). Mg(OH)2 also has a lower pKA than NaOH, lowering the risk of extreme perturbations to seawater pH compared to strategies that release NaOH directly. Alkalinity exchange is especially attractive because in some approaches to the electrochemical generation of alkalinity, the incoming brine must be pretreated to reduce the concentration of divalent cations such as Mg2+ and Ca2+ to prevent unwanted scaling in the electrochemical system (De Lannoy et al., 2018; Eisaman et al., 2012; Eisaman et al., 2018; Eisaman et al., 2023; Schiffer et al. 2024). In this case, alkalinity exchange can be used to partially pretreat the incoming brine by reducing the Mg2+ ion concentration to levels that are acceptable for input into the electrochemical system. While the Ca2+ ion concentration will still need to be reduced, the removal of Mg2+ during alkalinity exchange significantly reduces the cost of pretreatment. This process is conceptually illustrated in Figure 1A.
In Figure 1B, a simulated balanced mass flow of a bipolar membrane electrodialysis (BPMED) alkalinity exchange reactor is shown (Eisaman et al., 2012). The initial brine was assumed to have twice the ionic concentration of seawater. Additional Cl− ions were introduced to adjust for charge differences with Mg2+. The desired Mg(OH)2 slurry output was 3 mol AT/kg based on water content. The BPMED system uses energy to continuously exchange and separate Cl− and OH− creating basic and acidic streams which enter the NaOH and HCl Tanks. The generated NaOH is then pumped into the “Settling Tank” where it meets fresh RO reject brine (which is rich in Mg2+). Precipitation of Mg(OH)2 occurs, which removes essentially all Mg2+ in solution together with a stochiometric amount of OH−. A pH below 14 is maintained to prevent Ca(OH)2 precipitation which would reduce the yield of Mg(OH)2. CaCO3 precipitation is limited by the availability of CO32− in the solution. The values displayed in the “Settling Tank” represent the ion concentrations after precipitation. The Mg(OH)2 or “Magnesium Slurry” is continuously removed from the tank and forms the desired solid alkalinity carrier. Excess liquid also leaves the settling tank in order to balance the overall volume flows. On the other side of the BPMED stack, the acidic effluent is circulated into the “HCl Tank.” The HCl byproduct is then removed for beneficial use in other processes such as enhanced mineral weathering (Li and Hitch, 2015; Zhang et al., 2024), algal cultivation (Hibbeln et al., 2024), and upgrading of concrete aggregates (Jin et al., 2025). To maintain a constant pH in the HCl tank, “clarified” brine from the “Settling Tank” is pumped to mix with the acidic solution to generate the acidic waste stream. Instead of using the clarified brine, as shown here, it is also possible to use a different source of water for this purpose.
Understanding the Mg(OH)2 precipitation process and the kinetics of its subsequent dissolution in seawater when used as an alkalinity source for OAE is of critical importance for an OAE approach based on alkalinity exchange. Precipitation of Mg(OH)2 from brines using NaOH has been explored as a means of mineral extraction (Vassallo et al., 2021) and pretreatment (Ayoub et al., 2014; El-Manharawy and Hafez, 2003), but not in the context of its subsequent dissolution in seawater. Re-dissolution kinetics are critical for OAE approaches because sinking particulates that dissolve slowly release the alkalinity deeper in the water column, or potentially not at all, thereby delaying or decreasing the impact of the mineral addition on surface-ocean air-sea exchange of CO2. On the other extreme, highly soluble chemicals like CaO can dissolve rapidly and lead to large unwanted perturbations in surface water chemistry. The dissolution speed of magnesium hydroxide particles in water depends on the particle size distribution, the hydration state, and the surface dissolution rate. Dissolution is highly pH and temperature dependent and dissolution surface rates at seawater pH of 8.1 have been reported by a number of studies: Pokrovsky and Schott (2004) reports a surface dissolution rate of ground Brucite, 100-200 μm diameter particles, using a flow reactor as R = 1 × 10−8 mol/m2/s at pH 8.1 at 25 °C. Bharadwaj et al. (2013) reported dissolution of lab grade Mg(OH)2, 6 μm diameter particles, in a stirred, pH-controlled reactor, once converted to surface rate, as R= 𝑘𝑟r0p/M𝑟=4.07 ×10−6 mol/m2/s (r0 is initial particle radius, p is density, M𝑟 is molar mass) at pH 8.6 at 22 °C. Vermilyea (1969) finds rates from 1.5 × 10−5 to 3 × 10−7 mol/m2/s with 20 mm diameter particles. Kudoh et al. (2006) found the rate to be R = 5.0 × 10−9 mol/m2/s at pH 8 and 25 °C. These rates (Figure 2) span nearly 3 orders of magnitude suggesting that factors other than size, temperature, and pH may be at play, such as the degree of crystallinity and/or the presence of crystal defects. In this paper we examine the re-dissolution of precipitated Mg(OH)2, and the dependence of its material properties.
In this study, the role of morphological changes to Mg(OH)2 is investigated by altering the time that the material is left in its saturated supernatant after precipitation. By storing it in this type of environment, changes in chemical composition and crystallinity could occur that impact the kinetics of dissolution. This tests the sensitivity of Mg(OH)2 precipitation to this delay time as well as the ability to control the subsequent dissolution characteristics of the product to meet the needs of OAE. We also separately test the impacts of stirring during Mg(OH)2 precipitation and its subsequent dissolution. This research aims to inform potential future efforts to effectively manufacture and store Mg(OH)2 as a transportable alkalinity source for use in OAE.