Two-step mineral carbonation using seawater-based industrial wastewater: an eco-friendly carbon capture, utilization, and storage process

Abstract

Carbon dioxide is contained in the flue gas of many industrial plants, including those producing cement and power, and makes a significant contribution to climate change. Thus, methods to reduce carbon dioxide emissions are under intensive investigation. Carbon capture, utilization, and storage (CCUS) is a technology for converting captured carbon into useful chemical compounds while reducing CO₂ emissions. Thus, CCUS technology is a promising method for combating global warming. In this study, we focused on a mineral carbonation CCUS process using seawater-based industrial wastewater (SBIW). A two-step CCUS process using NaOH as an absorbent for CO₂ capture was used to produce two types of metal carbonate. Inductively coupled plasma optical emission spectroscopy, X-ray diffractometry, thermogravimetric analysis, and field-emission scanning electron microscopy techniques were used for analysis. Using our new technique, we obtained high-purity NaHCO₃ (92.2%) and MgCO₃ ·3H₂O (93%). We also found that it is possible to improve the utilization of the cations present in seawater.

Introduction

According to a study by the Carbon Dioxide Information Analysis Center, the global average land–sea temperature has increased since the 1880s, and carbon dioxide is widely recognized as the main reason for climate change [1]. Of the six main greenhouse gases (CH₄, HFC, H₂O, SF₆, perfluorinated carbons (PFC), and CO₂), CO₂ has the lowest global warming potential (GWP). However, CO₂ is emitted in vast quantities compared to the other greenhouse gases, resulting in significant changes to the atmospheric composition. The increased CO₂ concentration breaks the energy balance between the energy received from the sun and that lost from the earth’s surface. The captured infrared energy heats the surface of the earth, resulting in climate change [2]. Thus, the reduction of CO₂ gas emissions is regarded as a key method to prevent climate change. Carbon capture and storage (CCS) is one such technology for reducing atmospheric CO₂ concentrations. CCS is based on the capture of CO₂ gas using various absorbents or adsorbents, followed by the storage of the captured gas, which is often underground in geologically stable regions. Captured CO₂ can also be used in enhanced oil recovery (EOR) technology [3]. However, this application is only applicable to oil-producing countries. Furthermore, some regions are geologically unstable or have limited territory, making the underground storage of CO₂ impractical, especially because of the significant risks posed by CO₂ leakage. In addition, the acidic property of CO₂ gas can elute geological structures consisting of calcareous minerals [4]. Furthermore, the large amount of energy required for the desorption of CO₂ from the absorbent, which makes CCS an energy-intensive process, is another problem facing CCS technologies. For example, the desorption step can consume 70–80% of the total energy required for CCS operation [5, 6]. For example, monoethanolamine (MEA), the most popular absorbent in CCS technologies, requires 4.2 GJ/t-CO₂ for desorption [7]. Consequently, CCS technologies suffer from low cost-effectiveness, hindering their wide application in industry. An alternative process, carbon capture and utilization (CCU), has, thus come to attention. In CCU technology, CO₂ gas is converted into other usable chemical compounds, and these can be used as raw materials in industrial processes. Because this technology converts captured CO₂ into chemical compounds that can be used as feedstocks, no storage is required, thus removing this flaw of CCS technology [8]. There are two types of CCU technology, organic CCU and inorganic CCU. This categorization is based on the final product of CCU technology. In organic CCU technology, CO₂ is converted into organic compounds, such as formic acid or polymers, which are applicable in many industries. Despite its potentially broad applications, the use of CCU in industry is still far off at the present time, primarily because of the high energy consumption and unreasonable costs of organic CCU technology. Because CO₂ is a thermodynamically stable compound, the conversion of CO₂ into organic compounds requires a vast amount of energy. In short, currently, only laboratory-scale production is possible, and more studies are needed for the large-scale industrial application of this technology [8, 9]. In contrast, inorganic CCU technology converts CO₂ gas into inorganic compounds, usually metal carbonates. Inorganic CCU technology has many advantages: large quantities of CO₂ can be processed in a short time and the CO₂ is fixed permanently. Furthermore, the products of inorganic CCU can be used as raw materials in several industries, such as construction and medicine, making this process economically promising. Further, by using aqueous absorbents, such as MEA, CO₂ can be easily converted at ambient temperature and pressure, resulted in good economics [8]. Thus, inorganic CCU technology using aqueous absorbents is a very promising technology. Usually, inorganic CCU technology is based on mineral carbonation (also known as mineralization). In mineral carbonation, CO₂ gas is converted into metal carbonates by combination with cations such Ca²⁺ and Mg²⁺ [10]. Some of the cations for mineral carbonation come from natural sources such as rock material from mines. However, mining these resources can cause other kinds of pollution. Thus, the use of cations in industrial waste, such as slag and fly ash, has been studied as an alternative source of cations [11, 12]. Kang used seawater-based industrial wastewater (SBIW) originating from the sodium chloride manufacturing process to produce metal carbonates. SBIW contains various cations, including Ca²⁺, Mg²⁺, and Na⁺, in high concentrations, so it is suitable for mineral carbonation [13]. The cation components and concentrations of SBIW are listed in Table 1. According to Kang’s study, CO₂ is converted into metal carbonates, especially CaCO₃, which is the main product, and the byproduct liquid of the study contains high concentrations of Mg²⁺ and Na⁺ ions [13]. Considering that Mg²⁺ and Na⁺ can also be used for mineral carbonation, the byproduct of this first-pass CCU process can also be used for the mineralization of CO₂ [14, 15]. Thus, we have exploited the byproduct of the former study in order to utilize the remaining Mg²⁺ and Na⁺ ions. Our new process, which is a carbon capture, utilization, and storage (CCUS) process, produces MgCO₃·3H₂O and NaHCO₃, as shown in Fig. 1. Details of the process are given in the following sections.

Table 1 Concentration of cation components in seawater, SBIW, and pretreated SBIW

Fig. 1

Concept of the two-step mineral carbonation CCUS process using SBIW

In our CCUS process, we used NaOH as the absorbent for CO₂ capture. NaOH can be obtained by electrolysis of NaCl solutions, such as seawater, and the chemical equation of this electrolysis step is given in Eqs. (1–3) [16]:

$$ 2 {\text{H}}_{ 2} {\text{O }} + {\text{ 2Na}}^{ + } + {\text{ 2e}}^{ - } \to {\text{H}}_{ 2} + {\text{ 2NaOH}} $$

(1)

$$ {\text{H}}_{ 2} {\text{O }} + {\text{ 2Cl}}^{ - } \to 2 {\text{HCl }} + { 1}/ 2 {\text{O}}_{ 2} + {\text{ 2e}}^{ - } $$

(2)

$$ {\text{NaCl }} + {\text{ H}}_{ 2} {\text{O}} \to {\text{NaOH }} + {\text{ HCl}} $$

(3)

By using electrolysis, we can lower the operating cost of the whole process. For example, 1 kg of MEA costs US$2–4, and this “cost problem” can be overcome by electrolysis using renewable energy. On the other hands, 1 kg of NaOH cost US$ 0.2–0.3 [17]. Moreover, HCl is also a widely used industrial material; thus, its production provides an additional economic benefit. Using the electrolysis process explained above, our CCUS process can be operated cheaply. Furthermore, in the CO₂ capture process, Na⁺ cations are precipitated in the form of NaHCO₃, which has industrial uses. Because of its solubility, some of the NaHCO₃ remains in aqueous solution. The remaining HCO₃⁻ is converted to MgCO₃·3H₂O [18]. It is well known that magnesium carbonate does not form easily under ambient conditions and requires significant energy. For that reason, the carbonation of magnesium is typically achieved by modifying the temperature and pressure. According to Swanson [19], only 29.7 wt% of magnesium carbonate can be produced at a CO₂ partial pressure of 15 atm and 150 °C. Further, the formation of magnesium carbonate in the SBIW system might be harder because of differences in the thermodynamic favorabilities of magnesium carbonate and calcium carbonate [13]. For that reason, many scientists produce magnesium carbonate by modifying reaction condition. But modifying reaction condition for precipitate Mg needs huge amount of energy [19]. Thus, we pretreated the SBIW to remove calcium. According to Kang et al. [20], the Ca²⁺ in SBIW can be removed by treatment with sodium sulfate (Na₂SO₄), leaving Mg²⁺ and Na⁺ in the SBIW, and MgCO₃ can be produced more easily in the pretreated SBIW than in the untreated SBIW system. Consequently, the pretreatment of SBIW is key to the production of magnesium carbonate, which can be used in various industrial fields.

As mentioned above, the final products of the CCUS process are MgCO₃·3H₂O and NaHCO₃. This two-step mineral carbonation process is an innovation for CCUS process because most CCUS processes produce only one type of product. Further, the byproduct of this two-step mineral carbonation CCUS process can be reused. Because the cations contained in the SBIW are removed by carbonation, the main components of the byproduct solution should be NaCl. As explained above, NaCl can be electrolyzed to yield NaOH, which can be reused as an absorbent for CO₂ capture. Thus, by electrolyzing the byproduct NaCl solution, the process becomes cyclic, as shown in Fig. 2. We have named this process two-step mineral carbonation CCUS using seawater-based industrial wastewater.

Fig. 2

Cyclic two-step mineral carbonation CCUS process using SBIW

To study this process, we carried out three experimental parts. First, we mixed Na₂SO₄ with SBIW to remove Ca²⁺ in the form of CaSO₄. Secondly, 15 vol% of CO₂ gas and 10 wt% NaOH solution as an aqueous absorbent were mixed. Finally, we mixed CO₂-saturated aqueous absorbent with pretreated SBIW and filtered out the precipitates from the supernatant. The precipitates were then analyzed using inductively coupled plasma optical emission spectroscopy (ICP-OES) to determine the cation components and concentrations and X-ray diffraction (XRD) to identify the phases in the precipitate. In addition, field-emission scanning electron microscopy (FE-SEM) was used to observe the precipitate morphologies, and thermogravimetric analysis (TGA) and derivative TGA (DTG) were used to estimate the purity of the precipitates. The experimental design and the terminology are presented in “Experimental methodology” section.

Experimental methodology

In this study, we carried out CO₂ absorption experiments using NaOH, followed by conversion to metal carbonates using pretreated SBIW. The experiment was carried out in three steps: CO₂ absorption, SBIW pretreatment, and CO₂ conversion. The experimental apparatus is shown in Figs. 3 and 4.

Fig. 3

SBIW pretreatment step for cation separation

Fig. 4

Cation concentration of seawater-based resources and cation removal ratio (mol/mol, %)

Experimental material

In this study, we investigated the use of the CCUS process using an adsorbent that could be prepared by the electrolysis of seawater. However, because the NaOH solution generated from seawater contains various chemical compounds that may affect the absorption experiments, we used pure chemical reagents as the aqueous absorbent. The aqueous absorbent for CO₂ capture was 400 mL of 5 or 10 wt% NaOH (purity > 99%, Sigma–Aldrich). SBIW, which contains concentrated cations including Na⁺, Ca²⁺, and Mg²⁺, was used for the mineral carbonation process. This SBIW was the byproduct of the sodium chloride manufacturing process. We obtained SBIW samples from the Hanju-Salt Corporation (Ulsan City, ROK). As mentioned above, this SBIW requires pretreatment for cation separation using Na₂SO₄ (purity > 99%, Sigma–Aldrich) to precipitate the Ca²⁺ in the form of Ca₂SO₄.

Experimental procedure

In this study, we carried out two-step cation conversion experiments. The first step was the SBIW pretreatment step for Ca²⁺ precipitation using Na₂SO₄. The second step was CO₂ absorption by the aqueous NaOH solution. The final step was carbon conversion using the pretreated SBIW. The details of these two steps are given in the following sections.

SBIW pretreatment

Seawater contains 10,550 ppm Na⁺, 400 ppm Ca²⁺, and 1270 ppm Mg²⁺. After the sodium chloride manufacturing process, the seawater is concentrated, resulting in an increase in each cation concentration and yielding SBIW. Thus, SBIW contains high concentrations of cations such as Na²⁺, Ca²⁺, and Mg²⁺. Each cation concentration is listed in Table 1 Because Ca²⁺ is present in high concentration in SBIW, the final products of many earlier studies dealing with carbon conversion with raw SBIW was CaCO₃ because of the thermodynamics of the reaction of Ca²⁺ ions with CO₂, which makes the precipitation of Mg²⁺ ions challenging. Because the purpose of this study is to exploit the Mg²⁺ cations as well, the Ca²⁺ cations were removed.

To remove the calcium from the SBIW, Na₂SO₄ was mixed with 400 mL SBIW in consideration of the stoichiometry of Ca²⁺ and SO₄²⁻ ions for the precipitation of Ca²⁺ as CaSO₄. The mixture was mixed at ambient temperature and pressure for 24 h using a magnetic stirrer bar. After 24-h stirring, the mixture was filtered to separate the precipitates and supernatant. A vacuum pump and glass fiber filter (Whatman, pore size: 1.2 µm) were used for filtering, as shown in Fig. 3. Next, the precipitate was washed with 50 wt% aqueous ethanol solution to remove the polar-solvent-soluble salts. According to Kang et al., the precipitate contains 99.88 wt% calcium sulfate [20]. This result means that the Ca²⁺ component of SBIW is totally removed, whereas cations such as Na²⁺ and Mg²⁺ remain in the supernatant. The cation components and concentrations in the supernatants were determined using ICP-OES, and the results are listed in Table 1 and shown in Fig. 5. The pretreated SBIW has a low Ca²⁺ concentration, but there is no difference between the Mg²⁺ and Na⁺ concentrations before and after pretreatment. Thus, the pretreatment step achieves 97.2% (mol/mol) Ca²⁺ removal; the cation removal ratios are shown in Fig. 4.

Fig. 5

Schematic of CO₂ absorption experiment

CO₂ absorption using the aqueous NaOH solution: first carbonation step

For the CO₂ absorption step, we prepared the experimental apparatus as shown in Fig. 5. Flue gas from coal-fired power plants usually contains 15 vol% CO₂. Thus, the flue gas was simulated using a mixture of 15 vol% CO₂ and 85 vol% N₂; both gases were of high purity (N₂ and CO₂ gas, Sam-Heung Gas, purity > 99.9%). A mass flow controller (MFC) was used to control the flow rate of the gas fed into the equipment. The MFC was connected to a mass flow indicator, which controlled the flux of CO₂ to 300 mL/min and that of N₂ gas to 1800 mL/min. A ceramic gas diffuser with micrometer-sized pores was attached to end of the air-flow tube to mix the gas with the absorbent solution equally. These experiments were carried out at 1 atm pressure and 298.15 K; these conditions were maintained until the end of the experiment.

The absorption reaction was carried out in a 500-mL double-jacket reactor made of Pyrex glass and connected to a water bath. The temperature of the water bath (LC-LT407, LK Lab Korea, ROK) was set to 298.15 K, and the temperature was maintained constant. On the upper side of the reactor, the condenser was connected to a chiller (LC-LT408, LK Lab Korea, ROK) set to 276.15 K. The condenser prevented the aqueous solvent inside the reactor from evaporating, thus avoiding concentration changes because of the heat from the exothermic reaction. A gas mixer was connected after the MFC (MFC-3660, KOFLOC) to obtain a uniform mixture of CO₂ and N₂ gases. Then, a gas analyzer was set to analyze the change in the gas concentration during the absorption experiments. The gas analyzer (Gas Master, Sensoronic, ROK) monitored the concentration of the gas at the reactor output every 30 s; the output was displayed on a computer connected to the analyzer.

The aqueous absorbent solutions contained 5 or 10 wt% sodium hydroxide (Sigma–Aldrich, purity > 99%). The different absorbent concentrations were used to compare the CO₂ absorbing capacity and the amount of final product.

Before starting the CO₂ absorption experiments, the reactor was purged with high-purity N₂ gas (Sam-Heung Gas, purity > 99.9%) to remove impurities. After 5 min purging, the 5 wt% NaOH aqueous solution was placed in the reactor, which was sealed with a silicon plug and vacuum grease. Then, 300 mL/min of CO₂ gas and 1800 mL/min of N₂ gas from the MFC were mixed in the gas mixer. Subsequently, the mixed gas was flown into the reactor through a valve. The gas diffuser, which contains many small pores, was used to create bubbles. This instrument aids the dissolution of CO₂ gas in the aqueous solvent. To aid the absorption step, a magnetic bar and a stirrer located on the underside of the reactor were used, and the stirring rate was set to 300 rpm during the absorption experiments. The mixed gas flowing out from the gas diffuser was allowed to pass through the aqueous solvent. Some portion of the mixed gas was absorbed by the aqueous absorbent, and the remaining component passed out through the condenser and was transported to the gas analyzer. During the absorption experiments, the gas analyzer measured the CO₂ gas concentration of the gas mixture flowing out of the reactor. The concentration of CO₂ gas was 0% at the initial stage because most of the CO₂ gas was rapidly absorbed by the aqueous absorbent; as time passed, the concentration increased because of the decreased CO₂ absorbing capacity of the aqueous absorbent. Once the concentration of CO₂ gas displayed on the analyzer screen reached 15%, it was presumed that the absorbent was fully saturated with CO₂; at this point, the flow of mixed gas was stopped, the valve was closed, and the CO₂ absorption experiment was completed. At the end of CO₂ absorption step, we found that a white solid had precipitated. Because the precipitate was produced, we denote this as the first carbonation step. We separated this precipitate from the supernatant, and the precipitate was analyzed by XRD and FE-SEM. The supernatant was removed from the reactor for the carbon conversion step.

Carbon conversion via SBIW: second mineral carbonation step

After the CO₂ absorption step, CO₂ conversion using mineral carbonation with pretreated SBIW, denoted the second mineral carbonation step, was carried out. For this step, 400 mL of supernatant from the CO₂-saturated NaOH (5 and 10 wt%) solution was taken out of the reactor and mixed with 400 mL of the pretreated SBIW, i.e., SBIW with Ca²⁺ ions removed. Then, 800 mL of each mixture was placed in beakers equipped with magnetic stirrer bars and stirred using an electric stirrer machine at 300 rpm for 24 h. During the stirring time, the temperature was controlled to be 298.5 K using a water bath (LC-LT407, LK Lab Korea). The mixtures produced a white precipitate within a few seconds to a few minutes. After stirring, the mixtures were filtered through a glass fiber filter with 1.2-μm pores. An eggplant-type flask connected to a vacuum pump was used to filter out the supernatant and the precipitate. The supernatants were sampled in Falcon tubes and analyzed by ICP-OES (OPTIMA 8300, Perkin-Elmer, USA). The precipitates were dried in a convection oven set to 313.15 K for 48 h and ground to a powder. Then, the samples were placed in vials and analyzed by SEM (7800F, JEOL, Japan), XRD (ULTIMA IV, Rigaku, Japan), ICP-OES (Perkin-Elmer, OPTIMA 8300), and TGA (TA Instruments, SDT Q600).

Results and discussion

In this section, we report and discuss the experimental results and observations. This section contains a detailed discussion of the experimental results.

CO₂ absorption process using aqueous NaOH solution

In this study, we used aqueous NaOH solution for the absorption of CO₂ gas. NaOH is highly soluble in aqueous systems, producing solutions with a high pH. It is well known that the CO₂ speciation in aqueous solution is changed by the pH swing technique. CO₃²⁻ is dominant over HCO₃⁻ or CO₂ above pH 10, whereas, from pH 6 to 10, most of the CO₂ in solution exists as HCO₃⁻ [21]. The pH values of the 5 and 10 wt% NaOH solutions are 12.9 and 13.5, respectively. CO₂, an acidic gases, is, thus, easily captured by the very basic aqueous NaOH solution. The chemical reactions between aqueous NaOH solution and CO₂ gas are shown in Eqs. (4)–(11) [22]:

$$ {\text{CO}}_{ 2} \left( {\text{g}} \right) \to {\text{CO}}_{ 2} \left( {\text{aq}} \right) $$

(4)

$$ 2 {\text{NaOH}} \to 2 {\text{Na}}\left( {\text{aq}} \right) \, + {\text{ 2OH}}^{ - } \left( {\text{aq}} \right) $$

(5)

$$ {\text{CO}}_{ 2} \left( {\text{aq}} \right) \, + {\text{ OH}}^{ - } \left( {\text{aq}} \right) \leftrightarrow {\text{HCO}}_{ 3}^{ - } \left( {\text{aq}} \right) $$

(6)

$$ {\text{HCO}}_{ 3}^{ - } \left( {\text{aq}} \right) \, + {\text{ OH}}^{ - } \left( {\text{aq}} \right) \leftrightarrow {\text{H}}_{ 2} {\text{O}}\left( {\text{l}} \right) \, + {\text{CO}}_{ 3}^{ 2- } \left( {\text{aq}} \right) $$

(7)

$$ 2 {\text{Na}}^{ + } \left( {\text{aq}} \right) \, + {\text{ CO}}_{ 3}^{ 2- } \left( {\text{aq}} \right) \to {\text{Na}}_{ 2} {\text{CO}}_{ 3} \left( {\text{aq}} \right) $$

(8)

The total reaction can be expressed by Eq. (9):

$$ 2 {\text{NaOH}}\left( {\text{aq}} \right) + {\text{ CO}}_{ 2} \left( {\text{gas}} \right) \to {\text{Na}}_{ 2} {\text{CO}}_{ 3} \left( {\text{aq}} \right) \, + {\text{ H}}_{ 2} {\text{O}}\left( {\text{l}} \right) $$

(9)

Because CO₂ gas was fed into the Na₂CO₃ containing solution continuously, Na₂CO₃ becomes NaHCO₃, as shown in Eq. (8).

$$ {\text{Na}}_{ 2} {\text{CO}}_{ 3} \left( {\text{aq}} \right) + {\text{ H}}_{ 2} {\text{O}}\left( {\text{l}} \right) \, + {\text{ CO}}_{ 2} \left( {\text{aq}} \right) \to 2 {\text{ NaHCO}}_{ 3} \left( {\text{aq}} \right) $$

(10)

The final reaction of the absorption process is given by Eq. (11):

$$ {\text{NaOH}}\left( {\text{aq}} \right) \, + {\text{ CO}}_{ 2} \left( {\text{gas}} \right) \to {\text{NaHCO}}_{ 3} \left( {\text{aq}} \right) $$

(11)

We used two NaOH concentrations (5 and 10 wt%) to allow us to compare the absorption capacity of different absorbent concentrations. To compare the absorption capacities of the 5 and 10 wt% NaOH solutions, we calculated the loading value (mol-CO₂/L-solution) using Eq. (12) [13]:

$$ {\text{MOA}}_{\text{CO2}} = {\text{FR}}_{\text{CO2}} \times \left( {\frac{{{\text{C}}_{{{\text{R}},{\text{CO2}}}} - {\text{C}}_{\text{v,CO2}} }}{{{\text{C}}_{{{\text{v}},{\text{CO2}}}} }}} \right) \times \Delta \tau , $$

(12)

where MOA_CO2 is number of moles of absorbed CO₂ (mol); FR_CO2, CO₂ flow rate into the reactor (mol/s); C_V,CO2, concentration of CO₂ flowing out of vent (vol%); C_R,CO2, concentration of CO₂ flowing into from reactor (vol%), and ∆τ is the fixed time interval (s).

Using Eq. (12), we obtained the loading values for both NaOH absorbent solutions, which are shown in Fig. 6. The 5 and 10 wt% solutions absorbed 1.22 and 2.68 mol-CO₂/L-absorbent, respectively, as listed in Table 2. The difference in the loading values could arise from the amount of NaOH absorbent used. However, considering the amount of absorbent used and the CO₂ loading value of water, the difference in loading value is acceptable. Further, based on these data, the 10 wt% aqueous NaOH solution was chosen as an MEA replacement. A 30 wt% aqueous MEA solution is regarded as the standard CO₂ absorbing solvent, and this has a loading capacity of 2.5–2.6 mol-CO₂/L-absorbent. Comparing the 10 wt% aqueous NaOH solution and 30 wt% aqueous MEA solution, the NaOH solution has sufficient capacity for CO₂ capture to be used in a CCUS plant. However, the cost and energy usage required to reuse the absorbent must also be considered. One kilogram of MEA costs US$ 2–4 and requires 4.3 GJ/t-CO₂ of energy for desorption [7]. Unlike MEA, the CO₂-saturated aqueous NaOH solution, that is, the aqueous NaHCO₃ solution, can be converted into aqueous NaCl solution through mineral carbonation. These reactions are shown in Eqs. (13) and (14) [23]:

Fig. 6

CO₂ loading of the aqueous NaOH solutions

Table 2 Properties of the absorbents

$$ {\text{MCl}}_{ 2} \left( {\text{aq}} \right) \, + {\text{ NaHCO}}_{ 3} \left( {\text{aq}} \right) \to {\text{MCO}}_{ 3} \downarrow \left( {\text{s}} \right) \, + {\text{ NaCl }}\left( {\text{aq}} \right) $$

(13)

$$ {\text{MCl}}_{ 2} \left( {\text{aq}} \right) \, + {\text{ Na}}_{ 2} {\text{CO}}_{ 3} \left( {\text{aq}} \right) \to {\text{MCO}}_{ 3} \downarrow \left( {\text{s}} \right) \, + {\text{ NaCl }}\left( {\text{aq}} \right) $$

(14)

The mineral carbonation reaction is spontaneous, so the conversion of NaHCO₃ into NaCl does not require additional energy. Further, the aqueous NaCl solution can be electrolyzed to yield NaOH and HCl, as shown in Eqs. (1)–(3). The electricity for electrolysis could be obtained from a renewable source, such as solar energy. Thus, captured CO₂ is not emitted because a desorption process is not required, and the process can be carried out at low cost. Consequently, our proposed CCUS process is not only environmentally friendly, but also highly cost-effective.

Analysis of precipitates

In this study, a CCUS process involving two-step mineral carbonation with pretreated SBIW yielded two types of precipitate. Each precipitate was subjected to analysis to determine the key characteristics and components.

Precipitate of the first mineral carbonation step: NaHCO₃

The formation reactions of the precipitate in the first CO₂ mineralization step are given by Eqs. (4)–(11), and we expected that the precipitate produced in this step would be NaHCO₃. To check this assumption, we analyzed the precipitate using XRD, ICP-OES, and FE-SEM measurements. Figure 7 shows the results of XRD analysis of the precipitates formed from the 5 and 10 wt% aqueous NaOH solutions. The XRD patterns are the same; thus, the precipitates are identical. As shown in Fig. 7, the XRD pattern matches that of NaHCO₃. Furthermore, FE-SEM analysis of the precipitates (Fig. 8) reveals that the precipitates have a hexahedral morphology, which is typical of NaHCO₃, further indicating that the precipitate is NaHCO₃ [24]. We also weighed the produced precipitate, and the results are given in Table 3. Considering the number of moles of NaOH in the absorbent solutions (0.500 and 1.000 mol for 5 and 10 wt% solutions, respectively), the amount of precipitate produced during the first mineral carbonation step is consistent with the expected amount of NaHCO₃ within 10% error. These results further indicate that the precipitate was NaHCO₃. On the basis of three analytical results, we concluded that the precipitate was pure NaHCO₃. NaHCO₃ can be used in various fields such as medicine and healthcare, and, every year, 100,000 t of NaHCO₃ is produced for industrial use [25]. Concerning practical industrial use, the produced NaHCO₃ might contain some impurities, but the impurity is likely to be H₂O, which can be easily removed by heating [25]. If the manufacture of NaHCO₃ were replaced by our process, 52,377 t/year of CO₂ could be captured. Thus, our conceptual process is an eco-friendly advanced CCUS process.

Fig. 7

XRD pattern of the precipitate from the first step

Fig. 8

FE-SEM images of the first-step precipitates from both absorbent solutions. The scale bar is 10 µm

Table 3 Amount of precipitate produced in the first step

Precipitate of the second mineral carbonation step: MgCO₃·3H₂O

As mentioned in “Experimental methodology” section, we mixed SBIW with the CO₂-saturated absorbent, thus fixing the CO₂ as metal carbonates. This process is denoted by the second mineral carbonation step and resulted in a second type of precipitate.

Control experiments

Kang, Lee, and Jo studied mineral carbonation using waste resources such as SBIW, fly ash, or construction waste to fix CO₂ as metal carbonates. The waste material used to prepare MCO₃ contained high concentration of cations such as Ca²⁺, Mg²⁺, and Na⁺ or other kinds of heavy metal ions [11,12,13]. However, most of the metal carbonates formed were CaCO₃. The predominant formation of CaCO₃ is due to the thermodynamic characteristics of CaCO₃. The standard Gibbs free energy of formation for CaCO₃ is − 1129.1 kJ/mol [26]. This spontaneous reaction easily occurs in the presence of Ca²⁺ and CO₃²⁻ ions. Although there are other cations in the system, because the formation of CaCO₃ is thermodynamically favorable, its formation is favored over those of other metal carbonates, making the formation of these other metal carbonates, such as MgCO₃, challenging.

Pokrovsky studied mineral carbonation in Na⁺/Ca²⁺/Mg²⁺ systems and found that the metal carbonates produced depend on the cations present and their concentrations. It means that the trends of CaCO₃ formation, as like we mentioned above, would not be observed in our study because all the chemical reactions and phenomena observed in our study were in Na⁺/Ca²⁺/Mg²⁺ systems, as Pokrovsky’s study did [27]. Thus, we also carried out a control experiment where we used “raw SBIW” in the second mineral carbonation step. The precipitate was analyzed using ICP-OES and XRD, as shown in Fig. 9. In Fig. 9a, the XRD pattern of the precipitate is shown, and the XRD pattern matches that of calcite (CaCO₃). The ICP-OES analysis is shown in Fig. 9b and reveals the presence of a high concentration of Ca²⁺ ions. Thus, the main phase of the precipitate is likely calcite. In addition, as shown in Fig. 10, the FE-SEM images of the precipitate reveal the typical shape of calcite [13]. Park et al. [28] studied calcite precipitation using NaOH and CaO. In their study, Na⁺ cations hindered the formation of CaCO₃ because the Gibbs free energy of NaHCO₃ formation is relatively high compared to that of CaCO₃. However, in our study, CaCO₃ was formed more easily that in Park’s study. Possibly, this is due to the phase difference of the Ca²⁺ ions between the SBIW and CaO solutions. In the former study, Ca²⁺ ions originated from CaO in the solid form in an aqueous system and only a small proportion was present in the ionic state. In contrast, in SBIW, most of the Ca²⁺ ions are present in the ionic state, resulting in the easier formation of CaCO₃ [28]. In summary, the control experiments reveal that CaCO₃ is produced without the SBIW pretreating step, as reported in earlier studies. Thus, the SBIW pretreatment step is necessary to prevent CaCO₃ production.

Fig. 9

XRD and ICP-OES results of the control group precipitates. a XRD and b ICP-OES results of the control group precipitates from both absorbent solutions

Fig. 10

FE-SEM images of the control precipitates from both absorbent solutions. The scale bar is 100 nm

Second precipitate prepared using pretreated SBIW

As mentioned in the experimental section, we mixed the CO₂-saturated absorbent solution with pretreated SBIW. On mixing, precipitation occurred, and we analyzed the precipitate with XRD, ICP-OES, and FE-SEM. The results are shown in Figs. 11 and 12. In Fig. 11a, the XRD patterns of the precipitates are shown, and the patterns are consistent with that of MgCO₃·3H₂O (nesquehonite). In other words, MgCO₃·3H₂O was produced during the second mineral carbonation step with pretreated SBIW. The XRD pattern of the precipitate formed from the 10 wt% aqueous NaOH solution contains sharper peaks and has a greater similarity to the standard MgCO₃·3H₂O pattern than that of the precipitate obtained from the 5 wt% aqueous NaOH solution. This difference could be due to the different CO₃²⁻ concentrations of the absorbents. Because the 10 wt% aqueous NaOH solution has a higher CO₃²⁻ concentration than the 5 wt% aqueous NaOH solution, metal carbonates formed more readily in the former than the latter [29]. Ca was hardly detected in the precipitate.

Fig. 11

XRD and ICP-OES result of the second-step precipitates. a XRD and b ICP-OES results of the second-step precipitates from both absorbent solutions

Fig. 12

FE-SEM images of the second-step precipitates from both absorbent solutions. The scale bar is 10 µm

Figure 11b shows the ICP-OES results of the precipitates formed in the second stage of absorption. This figure shows us that all precipitates from each type of absorbent solution contain high concentrations of Mg²⁺. In addition, very little Ca²⁺ or Na⁺ was detected in the precipitate. Thus, the use of pretreated SBIW in the second mineral carbonation step enhances the precipitation of magnesium carbonates. Figure 12 shows the FE-SEM image of the second-step precipitate. The precipitate particles obtained from both absorbent solutions have an orthorhombic morphology, which is characteristic of MgCO₃·3H₂O [18, 30]. This result supports our conclusions obtained from the XRD and ICP-OES analyses.

Considering the analysis of the second-step precipitate, we conclude that the SBIW pretreatment step for cation separation is required for MgCO₃ ·3H₂O production. It is known that MgCO₃ ·3H₂O is metastable, subsequently undergoing conversion into hydromagnesite (Mg₅(CO₃)₄(OH)₂∙5H₂O). Yoo et al. [30] showed that a high concentration of Na⁺ ions in aqueous systems can prevent the conversion of MgCO₃·3H₂O into (Mg₅(CO₃)₄(OH)₂∙5H₂O). Since we have used NaOH for absorbing CO₂ gas and because there is a high concentration of Na⁺ in the pretreated SBIW system, the Na⁺ from the absorbent and pretreated SBIW should prevent the conversion of MgCO₃ ·3H₂O to Mg₅(CO₃)₄(OH)₂∙5H₂O. Thus, our goal, i.e., the utilization of cations in seawater for carbon capture, was successfully achieved, yielding Mg²⁺-containing precipitates.

Amount of fixed carbon

As mentioned in “Precipitate of the first mineral carbonation step: NaHCO₃” section, the first-step precipitate was pure NaHCO₃. This conclusion was reached easily and reasonably because only Na⁺ cations were present in the absorbent solution. In addition, we expected that the amount of NaHCO₃ produced would be linearly related to the Na⁺ concentration of the system. In contrast, for the second-step precipitate, a more in-depth characterization method was needed because the second carbonation step involved a complex system containing both Ca²⁺ and Na⁺. We estimated the purity of MgCO₃·3H₂O in the second-step precipitate using XRD and ICP-OES analysis. However, this analysis might suffer from severe error. Consequently, we also used TGA to investigate the purity of MgCO₃·3H₂O in the second-step precipitate. In an earlier study, Yoo found that MgCO₃·3H₂O loses 2.5 molecules of H₂O at 250.15–300.15 °C and 0.5 molecules of H₂O at 300–450 °C. Furthermore, 1 molecule of CO₂ is lost at 350–550 °C [30,31,32]. Based on Yoo’s results, we estimated the purity of MgCO₃·3H₂O using TGA/DTG measurements. Figure 13 shows the thermal decomposition curves depending on temperature (TG/DTG curves) of the second-step precipitate. The purities of MgCO₃·3H₂O from the 5 and 10 wt% absorbent solutions were 92.2% and 93%, respectively, and the results are listed in Table 4.

Fig. 13

TG/DTG analysis of the second-step precipitates (left) from the 5 wt% and (right) 10 wt% aqueous NaOH solutions

Table 4 Purity and conversion yield of precipitates

Based on our estimates of the purity of the metal carbonates from both mineral carbonation steps, it is possible to estimate the CO₂ conversion to metal carbonates via Eq. (15):

$$ {\text{Conversion yield }}\left( {\text{\% }} \right) = \frac{{{\text{Mol of converted CO}}_{2} }}{{{\text{Mol of absorbed CO}}_{2} }}. $$

(15)

Based on Eq. (15) and the measured data, the CO₂ conversions for the 5 and 10 wt% absorbent solutions were 68.4% and 77.6%, respectively, in the first mineral carbonation step and 19.4% and 14.4%, respectively, in the second mineral carbonation step. In sum, 87.8% and 92.0%, respectively, of the absorbed CO₂ gases can be fixed permanently by using the two-step mineral carbonation CCUS process, as shown in Fig. 14.

Fig. 14

CO₂ conversion yield for mineral carbonation

Conclusions

In this study, we investigated a two-step mineral carbonation process using seawater-based industrial wastewater. The basic idea of this study was to test a two-step mineral carbonation process for the all-in-one capture and conversion of CO₂ in seawater. Before the process began, we electrolyzed seawater to produce NaOH for CO₂ capture. This highly basic solution was then used for the carbonation process. In the first step, the aqueous NaOH solution was fully saturated with CO₂ and NaHCO₃ was produced as a result. This material could be used as a raw material in various industries. In the next step, we pretreated SBIW with Na₂SO₄ to remove the Ca²⁺ ions to ensure a high purity of the metal carbonates in the second mineral carbonation step. The pretreatment resulted in Ca²⁺-free SBIW, which was used to produce MgCO₃·3H₂O in the second carbonation step. Thus, CO₂ gas was fixed in the form of MgCO₃·3H2O using only seawater. Thus, overall, we have reported an absorbent solution for CO₂ capture and the subsequent production of high-purity metal carbonates, which have industrial applications, particularly NaHCO₃ and MgCO₃. The byproduct of this process is a solution containing a high concentration of sodium chloride, which can be electrolyzed to yield NaOH and HCl. The electrolysis of the liquid byproduct of our process thus yields NaOH, the absorbent for CO₂ capture. Consequently, the process for two-step mineral carbonation using SBIW is cyclic, as shown in Fig. 1. We believe that, using our new environmentally friendly, cyclic process, greenhouse gas emissions could be reduced while producing high-purity raw materials that can be used in various industries. Because this study proposes a new concept for mineral carbonation, it is not only academic, but also practical.