**1. Introduction**

The prevailing view today is that anthropogenic and geogenic greenhouse gases (GHGs) emitted into the atmosphere are the main cause of global warming. Energy production and use may be responsible for almost two-thirds of global greenhouse gas emissions (e.g., [1]). CO2 is considered to be the greenhouse gas with the greatest contribution to global warming, and global CO2 emissions are used as a clear indicator of global fossil energy consumption (e.g., [2]). Given the above assumptions, it can be concluded that capturing CO2 from industrial processes may be one of the main ways to control global temperature increases.

Mineral carbonation (MC) is considered to be one of the safest technologies for reducing CO2 emissions into the atmosphere and is used to capture and store CO2 in situ—in geological formations—or ex situ—as a potential solution for CO2 sequestration from smaller emitters where geological sequestration is not a viable option [3].

The advantage of mineral carbonation is the permanent storage of CO2 in the form of thermodynamically stable and environmentally friendly carbonates (e.g., [4–6]). This process is exothermic, and the raw materials for its operation are widely available (which is advantageous from an economic point of view) [7]. In the MC process, appropriately selected mineral substrates react with CO2 and form thermodynamically stable carbonates. This prevents emissions and ensures permanent CO2 sequestration [8,9]. One of proposed ex situ methods is the calcium looping technology—CaL [10].

Calcium looping (CaL) systems have been proposed as a less expensive method of CO2 capture for conventional power plants. In this process, a key role is played by calcium sorbent, which is used in alternating calcination and carbonation processes. The efficiency of the process varies depending on the properties of the sorbents used, which are expressed, inter alia, in the effects observed for the decreasing efficiency of gas capture with increasing number of CaL cycles. It is believed that this phenomenon is related to the reduction of the

active surface of the sorbent due to sintering and, possibly, the decrease in the chemical activity resulting from the reaction with sulfur oxides competing with carbonation. The reaction in which sulfur compounds are involved is largely similar to carbonation; however, it is irreversible under CaL conditions. It takes place in pores of small dimensions, and its products are deposited on the sorbent surface, which, in turn, makes carbonation difficult. The environments of carbonation are meso- and micropores, and especially in the latter ones, rapid filling with reaction products can take place [11].

The aim of this work was to assess the sequestration capacity of selected rocks using a simultaneous TGA-DSC analysis, thus simulating the calcium looping process. Such a method is suitable for small samples (e.g., drill cuttings, rock fragments), which are easy to obtain even at a very early stage of the raw material deposit recognition. Moreover, such tests do not require extensive reactors, and, in a relatively quick and simple way, they allow for the characterization of the material or screening of samples in terms of their suitability for CaL.

The calcium looping process—shown in Figure 1—uses a reversible chemical reaction,

$$\text{CaO} + \text{CO}\_2 = \text{CaCO}\_3,\tag{1}$$

between lime (CaO) and CO2, to capture CO2 from waste gas streams [10]. CO2 in the gas stream reacts with CaO in an exothermic carbonation reaction, forming calcium carbonate (CaCO3) at temperatures in the range of 600–700 ◦C. The CaCO3 from the carbonizer is then sent to a separate device called a calciner, where the calcination reaction takes place at a high temperature (around 900 ◦C). In these conditions, high-purity CO2 is released, which is suitable for transport to the sequestration site. The CaO produced is then sent back to the carbonator, closing the loop. Many researchers have proposed the oxy-combustion of coal in a calciner as a heat source for the calcination reaction [10,12]. The heat can be recovered from the exothermic carbonation reaction as well as from high-temperature gas and solid waste streams to generate electricity. As a result, CaL CO2 capture technology can be less energy intensive and more economical than the amine-based chemical absorption process.

**Figure 1.** Diagram of the calcium looping (CaL) process for CO2 capture, according to [13].

The efficiency of CO2 capture (carbonation) by the sorbent and its regeneration (calcination) depends on the reaction kinetics, the sorbent grain size, its specific surface area, and the pore space characteristics. The cyclicity of the CaL process is accompanied by a decrease in the active surface of the sorbent particles due to the tight packing of CaO

(regular system) in comparison to CaCO3 (trigonal system). Capture of CO2 by the CaO phase occurs in two stages. The first stage, in which the surface of the sorbent is covered with calcium carbonate, is characterized by a fast reaction rate and is strongly dependent on the partial pressure of CO2 [14,15]. The second stage is slow—the contact of CO2 with the sorbent depends on diffusion through the CaCO3 coating. Therefore, CaL installations should be based on the use of the first stage [16]. The capture process can be described by, among others, the shrinking core model (SCM), in which both stages are included, as shown in Figure 2. According to this model, the reaction proceeds at a narrow front that moves into the solid particle. The reactant is completely converted as the front passes by [17].

**Figure 2.** Schematic diagram of the shrinking core model (SCM), according to [17].

The conversion of sorbent can be expressed on the basis of mass change as follows:

$$X = \left(\frac{m - m\_0}{m\_0}\right) \cdot \frac{M\_{\text{CaO}}}{M\_{\text{CO}\_2}}\tag{2}$$

where *m* is the sorbent mass at the time t, *m*0 is the initial mass of the sorbent, and *MCaO* and *MCO*2 are the molar masses of *CaO* and *CO*2.

According to [17], given the spherical shape of sorbent particles, the relationship of sorbent conversion with particle radius is given by the formula:

$$(1 - X = \left(\frac{volume\ of\ unreacted\ core}{total\ volume\ of\ particle}\right) = \frac{\frac{4}{3}\pi r\_c^3}{\frac{4}{3}\pi R^3} = \left(\frac{r\_c}{R}\right)^3\tag{3}$$

Let the time for complete conversion of a particle be τ (s), and t is the time of the carbonation reaction (s). Then, in terms of the fractional time for the carbonation reaction, the conversion of the sorbent is given by:

$$X = 1 - \left(\frac{r\_c}{R}\right)^3 = \frac{t}{\tau}.\tag{4}$$

Thus, we obtain the general relationship of time with the radius and with the conversion, and the progression of the chemical reaction in terms of the fractional conversion becomes:

$$1 - (1 - X)^{\frac{1}{f}} = \frac{t}{\tau'} \tag{5}$$

while the progression of diffusion is:

$$\mathbf{1} - \mathbf{3}(\mathbf{1} - \mathbf{X})^{\frac{\mathbf{i}}{\mathbf{f}}} + \mathbf{2}(\mathbf{1} - \mathbf{X}) = \frac{\mathbf{t}}{\tau}.\tag{6}$$

#### **2. Materials and Methods**

In this work, rock samples representing the listed lithological types were studied as potential sorbents in the CaL process (Table 1).


**Table 1.** Samples studied as potential sorbents in the CaL process.

> 1 % CaMg(CO3)2, 2 % MgCO3.

The rock samples were mechanically ground using a PM 100 CM ball mill (Retsch, Haan, Germany) to a fraction of less than 0.08 mm. CO2 sorption studies were carried out on the STA 449 F3 Jupiter (Netzsch, Selb, Germany) apparatus. Samples weighing about 15 mg were placed in an Al2O3 crucible. First, the mass change of the analyzed samples was measured (Table 1); this allowed the determination of the share of calcite (or dolomite) in carbonate rocks. Measurements were made in N2 atmosphere at a temperature of up to 1030 ◦C and a heating rate of 10 K/min. Next, the simulation of the CaL process was carried out using a temperature program (Figure 3), which consisted in heating (calcination) of the primary sample from a temperature of 40 ◦C to about 900 ◦C at a rate of 20 ◦C/min in a N2 atmosphere with a flow of 25 mL/min. Then, a 10 min isothermal section was introduced, and the temperature was lowered to about 650 ◦C. Then, the carbonation process was carried out, keeping the sample at this temperature for 10 min with the attached CO2 flow at the rate of 25 mL/min. The test was carried out in 10 cycles of alternating calcination and carbonation. During the measurements, changes in the mass of the sample over time (TGA) and heat flow (DSC) were recorded.

In this work, we the tested raw sorbents (unmodified) as well as selected limestone sorbents that were thermally modified by pre-heating the sample for one hour at 1000 ◦C in 100% N2 atmosphere.

**Figure 3.** Example of a temperature program sequence (purple), and gas flow (green dashed line—N2, blue solid line—CO2).
