Modern climate change is considered one of the greatest threats to humankind (Hoegh-Guldberg et al., 2019; IPCC, 2021; The Royal Society and Royal Academy of Engineering, 2018). Global mean temperature has increased by 1.0 ∘C since pre-industrial times and could reach +1.2-1.9 ∘C in the next 20 years and +2.1-5.7 ∘C by the end of this century (IPCC, 2021). Furthermore, about 26 % of all anthropogenic carbon dioxide (CO2) emissions were taken up by the ocean through air-sea gas exchange between 1750 and 2020 (Friedlingstein et al., 2022). This has led to a decrease in the average open-ocean pH by 0.1 units in a process termed ocean acidification – OA (Bates et al., 2012; Canadell et al., 2007; Carter et al., 2019; Cyronak et al., 2014; Doney et al., 2009; Hoegh-Guldberg et al., 2007).
The aim of the 2015 Paris Agreement is to minimise the negative impacts of global warming and OA by limiting global warming to less than +2.0 ∘C, ideally below +1.5 ∘C, by the end of this century (Goodwin et al., 2018). However, the current and pledged reductions will likely not be enough, and additional CO2 mitigation strategies are needed, such as ocean alkalinity enhancement – OAE (Gattuso et al., 2015; GESAMP, 2019; Lenton and Vaughan, 2009; The Royal Society and Royal Academy of Engineering, 2018). OAE could be an efficient approach for CO2 removal (current emissions of 40 Gt yr−1), with models suggesting a potential of 165 to 790 Gt (1 Gt = 1015 g) of atmospheric CO2 removed by the year 2100 on a global scale if OAE were implemented today (Burt et al., 2021; Feng et al., 2017; IPCC, 2021; Keller et al., 2014; Köhler et al., 2013; Lenton et al., 2018). However, empirical data on OAE efficacies are limited, and safe thresholds for mineral dissolution are particularly lacking (National Academies of Sciences and Medicine, 2022).
OAE typically relies on the dissolution of alkaline minerals in seawater, releasing alkalinity similarly to natural rock-weathering processes (Kheshgi, 1995). Suitable candidates are magnesium-rich minerals such as brucite, periclase or forsterite and calcium-rich minerals such as quick and hydrated lime (Renforth and Henderson, 2017). Quick lime and hydrated lime are of particular interest due to their high solubility in seawater and rapid dissolution. Quick lime, i.e. calcium oxide (CaO), is obtained by the calcination of limestone, composed primarily of calcium carbonate (CaCO3), which is present in large quantities within the earth’s crust. Once heated to temperatures of ∼ 1200 ∘C, each molecule of CaCO3 breaks down into one molecule of CaO and one molecule of CO2 (Ilyina et al., 2013; Kheshgi, 1995). Hence, for maximum OAE potential, carbon capture during calcination and subsequent storage would be necessary (Bach et al., 2019; Ilyina et al., 2013; Kheshgi, 1995; Renforth et al., 2013; Renforth and Kruger, 2013). CaO can be hydrated into calcium hydroxide (Ca(OH)2), also known as hydrated lime. The addition of either CaO or Ca(OH)2 to seawater leads to the dissociation of Ca(OH)2 into one calcium Ca2+ and two hydroxyl ions OH− (Feng et al., 2017; Harvey, 2008). Ignoring the non-linearities of the seawater carbonate system (i.e. changes in total alkalinity, TA, and dissolved inorganic carbon, DIC, are not 1:1), the chemical reaction of CO2 and Ca(OH)2 dissolution and the subsequent uptake of atmospheric CO2 can be written as
The dissolution of CaO and Ca(OH)2 and the subsequent addition of TA increase the pH of seawater, which changes the carbonate chemistry speciation (Zeebe and Wolf-Gladrow, 2001). DIC can be approximated as the sum of HCO3- and CO32- (ignoring the small contribution by CO2). Similarly, TA can be approximated as the sum of HCO3- and 2 CO32- (ignoring the smaller contributions by boric and silicic acids and other minor components). Combining both DIC and TA equations reveals that CO32- concentrations can be expressed as [CO32-] = TA − DIC. Hence, increasing TA at a constant DIC, e.g. by dissolving CaO or Ca(OH)2, increases [CO32-], shifting the carbonate chemistry speciation towards a higher pH (Fig. A1) (Dickson et al., 2007; Wolf-Gladrow et al., 2007; Zeebe and Wolf-Gladrow, 2001). The subsequent shift in DIC speciation leads to a decrease in dissolved CO2 concentrations, reducing the partial pressure of CO2 (pCO2) in seawater and increasing its atmospheric CO2 uptake potential.
Depending on the amount of TA added and the initial seawater pCO2, the TA-enriched seawater would either take up CO2 from the atmosphere or reduce outgassing of CO2. Factoring in the non-linearities of the carbonate system, about 1.6 mol of atmospheric CO2 could be taken up per mole of dissolved CaO or Ca(OH)2 (Köhler et al., 2010). Furthermore, dissolving CaO and Ca(OH)2 can also counteract ocean acidification. During the dissolution of alkaline minerals, both pH and the CaCO3 saturation state of seawater (Ω) increase through increasing Ca2+ and CO32- concentrations. This makes OAE a dual solution for removing atmospheric CO2 and mitigating OA (Feng et al., 2017; GESAMP, 2019; Harvey, 2008). However, there are important knowledge gaps in our understanding surrounding basic mineral dissolution in seawater (Feng et al., 2016; González and Ilyina, 2016; Mongin et al., 2021; Renforth and Henderson, 2017).
One knowledge gap is the critical Ω threshold beyond which CaCO3 starts to precipitate inorganically. Such secondary precipitation constitutes the opposite of alkaline mineral dissolution and would decrease pH and Ω while simultaneously increasing the CO2 concentration in seawater. This would decrease the ocean uptake’s capacity for atmospheric CO2, causing the opposite of the intended effect. Additionally, if all added alkalinity were precipitated, only 1 mol of atmospheric CO2 per mole of Ca2+ would be removed, instead of ∼ 1.6 mol in the absence of CaCO3 precipitation. If even more CaCO3 precipitated, the efficiency of OAE would be further reduced. Under typical seawater conditions, CaCO3 precipitation does not occur due to the absence of mineral-phase precipitation nuclei and the presence of precipitation inhibitors such as dissolved organic compounds, magnesium (Mg) or phosphate (Chave and Suess, 1970; De Choudens-Sanchez and Gonzalez, 2009; Pytkowicz, 1965; Rushdi et al., 1992; Simkiss, 1964). There are three types of CaCO3 precipitation, (1) homogeneous (in the absence of any precipitation nuclei), (2) heterogeneous (in the presence of mineral phases) and (3) pseudo-homogeneous (in the presence of colloids and organic materials) (Marion et al., 2009; Morse and He, 1993). For pseudo-homogeneous precipitation, the critical threshold at which calcite precipitates spontaneously is at a calcite saturation state (ΩC) of ∼ 18.8 (at a salinity of 35 and at a temperature of 21 ∘C) (Marion et al., 2009). Assuming typical open-ocean carbonate chemistry (e.g. TA ∼ 2350 µmol kg−1 and DIC ∼ 2100 µmol kg−1), this threshold would be reached through an increase in TA of ∼ 810 µmol kg−1. This corresponds to a critical threshold for Ω with respect to aragonite, i.e. ΩA, of ∼ 12.3. The two other types of precipitation (i.e. homogeneous and heterogeneous) are more poorly constrained (Marion et al., 2009). Importantly, at current dissolved Mg and Ca concentrations in seawater, the CaCO3 polymorph that is favoured during inorganic precipitation is aragonite rather than calcite (Morse et al., 1997; Pan et al., 2021). Therefore, aragonite saturation state ΩA may be a more important determinant of critical runaway precipitation thresholds. No matter what mineral phase is precipitating, a better understanding of CaCO3 precipitation under conditions relevant to OAE is needed.
To gain a better understanding of the consequences of CaO and Ca(OH)2 dissolution for OAE, we conducted several dissolution experiments with CaO and Ca(OH)2 to determine (1) how much alkaline material can be dissolved without inducing CaCO3 precipitation, (2) what causes secondary CaCO3 precipitation and (3) how secondary precipitation can be avoided.
