Effect of processing temperatures on decarbonisation of calcium carbonate in sodium hydroxide solutions at ambient conditions

doi:10.21203/rs.3.rs-1597303/v1

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Effect of processing temperatures on decarbonisation of calcium carbonate in sodium hydroxide solutions at ambient conditions

https://doi.org/10.21203/rs.3.rs-1597303/v1

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The CO₂ released upon calcination of limestone (process CO₂ and combustion CO₂) accounts for the largest portion of the emissions from the cement manufacturing industry. Our previous works highlighted the possibility for a no-combustion decarbonisation of CaCO₃ through reaction with NaOH solutions to produce Ca(OH)₂ at ambient conditions, while sequestrating the process CO₂ in a stable mineral Na₂CO₃∙H₂O/Na₂CO₃. In this study, the effect of temperature was assessed within the range of 45–80°C, highlighting a higher impact at short residence times. The proportioning of the precipitated phases Na₂CO₃∙H₂O /Na₂CO₃ was also assessed at varying conditions, and the data was supported with related simulation. Experiments performed on targeted CaCO₃:NaOH:H₂O compositions also allowed for the detection of the Arrhenius parameters characterising each ternary system. An increasing activation energy (Ea) was found to be correlated to lower NaOH concentrations initially reacting with CaCO₃. The present work offers a further understanding of the effect of temperature on the process with the potential to minimise the GHG emissions from cement and lime industries, and correctly assess eventual industrial applicability.

The thermal decarbonisation of limestone to give lime (CaCO₃ → CaO + CO₂) is a crucial step for the production of Portland Cement (PC), whose worldwide market exceeded 4 billion metric tonnes in 2018 [1]. Given the heavy carbon footprint of the manufacturing process (approximately 0.8 t_CO2/t_PC [2]), the cement industry is currently responsible for about the 5–8% of the total CO₂ emissions worldwide [1, 3]. Also, considering the 2.9% increase in annual production expected by 2050 [4], the environmental burden of the cement industry is destined to intensify. The very energy-intensive [5] calcination of limestone itself occurs within the preheater tower of the cement plant pyro-processing system [6], and the emissions arising from this step are linked to both the process and combustion CO₂. While the former results from the stoichiometry of CaCO₃ decomposition, the latter refers to the combustion of the fuels (primarily fossil fuels) necessary to achieve the calcination temperature ( above 900°C) [7]. In order to meet the environmental goals set by the UN in 2015 (at least 80% of CO₂ emission cut from industry sector by 2050 [8]), a significant contribution is expected from the cement industry. The use of alternative binders with a lower calcium content [9], the efficiency enhancement of the existing carbon capture and storage (CCS) technologies [10], and the use of alternative fuels [11] are currently seen as important in mitigating these emissions. Despite that, all these solutions start from the assumption that the thermal calcination of limestone is unavoidable, as is the process CO₂. In contrast to this, recent works [12, 13] highlighted the possibility for alternative decarbonisation routes which apply no-combustion approaches for the synthesis of the essential chemical CaO from CaCO₃.

One of these alternatives has been recently discussed in our previous work [13], and it consists in a chemical decarbonisation of CaCO₃ to give Ca(OH)₂, occurring at ambient conditions by reaction with concentrated NaOH solutions (Eq. 1).

$$CaC{O}_{3}+2NaOH+x{H}_{2}O\to Ca(OH{)}_{2}+N{a}_{2}C{O}_{3}\bullet {xH}_{2}O (1)$$

Despite the application of the proposed alternative to the cement industry would prevent the process CO₂, the synthesis of the clinker phases still requires temperatures above 1450°C; that implies that the fuel combustion CO₂ would not be avoided at all.

The title study aims to gain a deeper insight into the chemistry of the heterogeneous system taken into account here, while providing essential information for an eventual scale-up. The full understanding of the system kinetic is a fundamental step towards the industrialisation, since it will be crucial to determine the feasibility of the scaled-up process. The reaction velocity may be defined by Eq. 2 [14], using the reaction time (t), temperature (T), and extent of reaction (0 < α < 1) [14]. The function f(α) represents the kinetic reaction model depending on the mechanism of the process [15], and f(T) is applied in the Arrhenius equation (Eq. 3) that varies with temperature, activation energy (E_a) and the pre-exponential factor (A). E_a is the minimal energy that needs to be overcome by the reactants in order for the reaction to occur [16], while the pre-exponential factor provides an estimation of the probability of the collisions between the reactants [17].

The numerical integration of Eq. 2 would lead to the expression of g(α) (Eq. 4), allowing for the experimental evaluation of E_a and A for single-step reactions only. In fact, the above equations would apply to each step involved in any reaction, therefore significantly affecting the final value.

Although the temperatures here considered (45–80°C) are not very high, keeping the system temperature constant requires an additional energy expenditure, which would increase the operational costs of the process. Therefore, the reaction efficiency was also tested without actively maintaining the temperature after the initial heating of the NaOH solutions to the targeted temperatures. Furthermore, some of the ternary starting compositions tested in our previous study [13] were subjected to the same processing conditions and increasing temperatures here, to obtain the parameters E_a and A.

2.1. Experimental procedure

The materials used in this work are reagent grade chemicals: Sigma-Aldrich CaCO₃ (≥ 99%), Honeywell Fluka NaOH (≥ 97%), and distilled water. The starting mixture of CaCO₃ (26.9 wt.%), NaOH (32.7 wt.%) and distilled water (40.4 wt.%) was first investigated, with the solution having an initial 20 m molality. Such a composition was chosen given the 86% conversion efficiency previously observed [13]. Also, the water to solids ratio of 1.5 g/g was ensuring a homogeneous mixing of the solid powders even at milder stirring conditions, as described below, in keeping with the previous investigation [13].

The experimental set up is illustrated in Fig. 1. The 20 m NaOH solutions were prepared in a polypropylene PP plastic vessel, and placed in the pre-heated water bath for 30 minutes at the target temperatures of 45°C, 60°C, or 80°C.

The direct contact between the hot plate and the bottom of the reaction vessel was avoided through the placement of an insulating glass disk. The CaCO₃ powders were also pre-heated at 80°C for 2 hours to remove any adsorbed water. The CaCO₃ was then quickly added to the NaOH solutions and manually mixed, at first, to prevent the agglomeration of the solids. Following, the vessel was covered with a lid to minimise the loss of water and heat, and a magnetic stirring of 300 rpm was set up. A constant temperature was ensured throughout the residence times investigated: 1, 2, 3, 4 and 5 minutes. Longer residence times were not considered, since the achievement of thermodynamic equilibrium was not the goal of the present study.

The reaction was also performed for five minutes without actively maintaining the temperature of the system after the initial heating of the NaOH solutions at the targeted temperature to assess the advantages of a controlled system temperature. For all systems tested, once the reaction was complete, 30 mL of methanol was added to the system to mix for a further five minutes, ensuring the removal of the unreacted NaOH. The solubility of NaOH in methanol is 238 g/L [18], while the negligible solubility of CaCO₃ [19], Ca(OH)₂ [20], Na₂CO₃ [21] and Na₂CO₃∙H₂O [21] would ensure the same solid phases assemblage remained upon methanol washing. The solid products and the liquid streams could then be collected separately using Buchner filtration. The solids were dried in an 80°C oven for two hours, before being manually ground and sieved below 63 µm for subsequent characterisation though TG and XRD analysis.

An additional set of experiments was conducted to estimate the activation energy E_a of the decarbonisation reaction at different starting compositions (Table 1), selected from our previous study [13]. The reaction was conducted at a constant temperature adopting the same procedure described above, and all the samples were reacted for one minute at 30⁰C, 45⁰C, 60⁰C and 70⁰C. The datasets used and/or analysed during the current study are available from the corresponding author on reasonable request.

Table 1

Starting compositions (w%) and the molar ratio NaOH/CaCO₃ of the samples.
ID	NaOH (w_%)	CaCO₃ (w_%)	H₂O (w_%)	NaOH/CaCO₃ (mol/mol)	H₂O/CaCO₃ (g/g)	H₂O/NaOH (mol/mol)	α [13]
NaOH_10M	14.3	14.3	71.4	2.5	3.0	5.5	0.07
NaOH_12M	24.5	24.5	51.0	2.5	2.1	4.6	0.46
NaOH_15M	30.1	20.0	49.9	3.8	2.5	3.7	0.69
NaOH_17M	37.2	8.1	54.7	11.5	6.7	3.3	0.96

2.2. Characterisation techniques

2.2.1. X-Ray Diffaction (XRD)

The reaction products were qualitatively identified by X-ray diffraction (XRD) and following analysis on the Highscore-Plus software with the PDF^− 4 2019 database. The measurements were performed using a Bruker D2 PHASER desktop X-ray diffractometer in the Bragg-Brentano geometry armed with a Cu-Kα radiation source running at 30 kV and 10 mA. The instrument is equipped with a one-dimensional LYNXEYE detector and a 1 mm divergence slit. All samples were in a powder form, and they were loaded onto the sample holder with 2.5 cm diameter and 1 mm deep. All the analyses were conducted between 5° and 80° (2θ) with a step size of 0.02° at 0.5 s per step, with the stage rotating at 15 rpm in order to improve counting statistics.

2.2.2. Thermogravimetry (TG/DTG)

A PerkinElmer TGA 4000 was used to provide Thermogravimetric analysis (TGA) for the reaction products; approximately 40 mg of sample were subjected to a temperature ramp from 30°C to 800°C at the heating rate of 10°C/min, with a 40 mL/min of N₂ flow. The heating program was set up to held the sample at 800°C for 1 hour to ensure complete loss of CO₂ from CaCO₃; the analysis temperature was not exceeding 800°C to avoid the melting or weight loss from Na₂CO₃. To identify evolving gases, a Hiden mass spectrometer (HPR-20 GIC EGA) was used to record the signals for CO₂ and H₂O. An example of TG analysis is presented in Fig. 2, where the mass loss events are linked to Ca(OH)₂, Na₂CO₃∙H₂O and CaCO₃. Together with the weight loss, the DTG trend is also reported to ensure the distinction of eventual overlapping signals. The extent of reaction was calculated based on the amount of Ca(OH)₂ [22] and unreacted CaCO₃ [23], based on the respective thermal decompositions. The content of Na₂CO₃∙H₂O could similarly be estimated in the temperature range 50–130°C [24].

The measurement was repeated 6 times for each sample to estimate the measurement error (± 0.16 wt.%, ± 0.10 wt.% and ± 0.16 wt.% for Na₂CO₃∙H₂O, Ca(OH)₂ and CaCO₃, respectively).

$$\alpha = \frac{w{\%}_{\left[Ca(OH{)}_{2}\right]}}{M{W}_{Ca(OH{)}_{2}}}/[\frac{w{\%}_{\left[Ca(OH{)}_{2}\right]}}{M{W}_{Ca(OH{)}_{2}}}+\frac{w{\%}_{\left[CaC{O}_{3}\right]}}{M{W}_{CaC{O}_{3}}}\left] \right(5)$$

3.1. Effect of temperature on reaction efficiency

The XRD analysis of the 45⁰C_n, 60⁰C_n, and 80⁰C_n samples are shown in Figs. 3A, B, and C, respectively. Upon reaction, only Ca(OH)₂, Na₂CO₃, Na₂CO₃∙H₂O and unreacted CaCO₃ could be detected, with respective main reflection angles at 2θ of 29.5⁰ [25], 34.1⁰ [26], 16.9⁰ [27] and 30.1⁰ [28]. No additional phases were detected, suggesting the absence of secondary and competing reactions to occur.

The TG data for the 45⁰C_n, 60⁰C_n and 80⁰C_n series is shown in Fig. 4A, B and C, respectively.

The phase quantification could be performed through the weight losses observed in the TG data for CaCO₃ (at 560–800°C), Ca(OH)₂ (at 310–470°C) and Na₂CO₃∙H₂O (at 50–130°C). The estimated quantities for all the samples discussed are summarised in Table 2, where the processing conditions used are stated in the sample IDs.

Table 2

Processing conditions and phases composition gained from thermogravimetric analysis, together with the Loss On Ignition (LOI) expressed in % and the Na₂CO₃.H₂O/Ca(OH)₂ molar ratio registered for all the samples discussed.
Sample ID	CaCO₃ (wt.%)	Ca(OH)₂ (wt.%)	Na₂CO₃∙H₂O (wt.%)	Na₂CO₃ (wt.%)	LOI (%)	Na₂CO₃.xH₂O/Ca(OH)₂ (mol_%/mol_%)
45⁰C_1min	13.7	34.5	6.6	45.2	15.4	1.1
45⁰C_2min	6.4	35.6	23.6	34.4	14.9	1.1
45⁰C_3min	4.2	34.3	55.9	5.6	18.3	1.1
45⁰C_4min	4.8	32.9	59.6	2.7	18.8	1.1
45⁰C_5min	3.2	33.5	61.6	1.7	18.5	1.1
45⁰C_RT_5min	4.6	34.7	57.0	3.7	18.7	1.1
60⁰C_1min	8.9	34.8	15.1	41.2	14.6	1.1
60⁰C_2min	4.5	36.8	18.0	40.7	13.5	1.1
60⁰C_3min	4.2	37.7	8.6	49.5	12.2	1.1
60⁰C_4min	4.3	36.7	8.9	50.1	15.0	1.1
60⁰C_5min	3.1	37.8	11.1	48.0	12.2	1.1
60⁰C_RT_5min	3.2	33.5	61.6	1.7	18.5	1.1
80⁰C_1min	8.5	34.7	11.1	45.7	13.8	1.2
80⁰C_2min	6.1	35.9	11.0	47.0	13.0	1.2
80⁰C_3min	5.2	36.1	6.6	52.1	12.0	1.2
80⁰C_4min	5.5	36.1	7.9	50.5	12.3	1.2
80⁰C_5min	5.2	35.8	11.2	47.8	12.6	1.2
80⁰C_RT_5min	6.0	35.2	15.0	43.8	13.4	1.2

It must be mentioned that the content of Na₂CO₃ was calculated by subtracting the sum of the other quantified phases Ca(OH)_2¸ CaCO₃ and Na₂CO₃∙H₂O from the total mass (100%). In fact, since it would start decomposing above 851⁰C [29], it could not be directly quantified by TG analysis (up to 800⁰C) through the detection of the relevant peak. A higher content in Na₂CO₃ could also be suggested by the lower LOI registered for samples with similar reaction efficiencies (Table 2). This aspect will be extensively discussed in Section 3.2. However, the quantification was considered reliable since the XRD analysis confirmed the absence of any additional phases in the solid products. Moreover, the ratio between the mol% of Na₂CO₃.H₂O/Na₂CO₃ and Ca(OH)₂ revealed a good accordance to the stoichiometry expressed in Eq. 1, since the values were close to the unity. In fact, the conversion of 1 mole of CaCO₃ should theoretically lead to the precipitation of 1 mole of both Na₂CO₃.xH₂O and Ca(OH)₂ (Eq. 1). Specifically, the ratios were suggesting a slight over-production of Na₂CO₃.H₂O and Na₂CO₃ with respect to Ca(OH)₂ for all the systems studied, and that could possibly be referring to the distribution of the ionic sites Ca²⁺ and CO₃²⁻ of the surface of the CaCO₃ crystals used, with an average particle size of 27.5 µm (D_x50) and specific surface area of 126.8 m² kg^− 1. Statistically, a 27% excess of negatively charged sites may be found on the CaCO₃ surface [30], justifying the greater affinity of the CaCO₃ to interact with the cationic species Na⁺ in the liquid bulk.

Based on the TG data, extent of reaction (α) was calculated for each system, and the outcomes are reported in Fig. 5.

The conversion of CaCO₃ was high in the tested reaction conditions, since extents of 0.7–0.8 were obtained. The general trends of the extent of reaction coincides with that of XRD data (Figs. 3A, B, and C) which shows progressive decrease in the intensity of the CaCO₃ main peak at 29.5⁰ 2θ. In the first 60 seconds, temperature had a significant impact on the extent of reaction, showing enhanced efficiency of CaCO₃ conversion at higher temperatures, while limited effects of temperature were observed at longer reaction times. This is because the reaction is nearly completed with the longer reaction time even at the lower temperatures. Likely, the higher conversion registered at short residence time and higher temperature could be linked to the lower viscosity of the NaOH solutions, favouring the ionic mobility and the further interaction between the dissolved species and the solid surface and bulk.

The efficiency of the system may also be expressed in terms of CO₂ capture, expressed as moles of CO₂ precipitated as Na₂CO₃.xH₂O per second of reaction progression. As reported in Fig. 6, the CO₂ capture rate was decreasing from ~ 4.5*10^− 4 mol_CO2·sec^− 1 in the first 60 seconds of reaction down to two order of magnitude below (~ 10^− 6 mol_CO2·sec^− 1) after 300 seconds of contact time. In other terms, around the 80% of the total process CO₂ initially introduced was effectively captured after 60 seconds of reaction.

The samples reacted without actively maintaining the temperature after the initial heating at 45, 60 and 80°C (temperatures of 21.5, 42.1, and 53.6°C were registered at the end of the reaction, respectively), indicated extent of reactions similar to those reacted at a constant temperature after 300 seconds (Fig. 5). Since the reaction proceeds fast, majority of the reaction must have happened before the temperature decreases significantly.

This is beneficial because a 20m NaOH solution (c_p = 3,370 J kg^− 1 ⁰C^− 1 [31]) would only require an initial energy input of 67.4, 117.9 and 185.3 kJ∙kg^− 1 to reach the temperature of 45, 60 and 80°C, respectively, from ambient conditions (25°C, 1 atm). As a result, energy inputs of 206, 350 and 589 kJ are required to chemically decarbonise 1 kg of CaCO₃ at T_i of 45, 60 and 80⁰C, respectively. It follows that the decarbonisation route proposed here requires a much lower energy input with respect to the conventional calcination route, where 1819.4 kJ is required (thermodynamically) to decarbonise 1 kg of CaCO₃ [32]. Naturally, this comparison can only be done on a laboratory-scale, whereas additional energy expenditures should be taken into account when considering an industrial case, such as the handling of the materials, transport of the reactants and products, maintenance of the reactors.

3.2. Proportioning of Na₂CO₃∙H₂O and Na₂CO₃

The reaction at different conditions resulted in sequestration of CO₂ by forming Na₂CO₃∙H₂O and Na₂CO₃ in different proportion, with the x value (Eq. 1) equal to 1 and 0, respectively. To further study this phenomenon, the proportion between Na₂CO₃∙H₂O and Na₂CO₃ was examined. Based on the data obtained from TGA (Figs. 4A, B and C), the molar fraction ν of precipitated Na₂CO₃ was calculated by dividing the moles of Na₂CO₃ in the samples by the total moles of Na₂CO₃∙H₂O and Na₂CO₃. As reported in Fig. 7, at a constant temperature of 45⁰C, νNa₂CO₃ decreased significantly from 0.88 to 0.11 in the first 3 minutes, whereas the precipitation of Na₂CO₃ was dominant at 60 and 80°C for all the residence times investigated.

The samples reacted at ambient conditions showed enhanced precipitation of Na₂CO₃∙H₂O at initial temperatures T_i of 45 and 60⁰C. Such a different proportion gained at different values of T_i might be explained by taking into account the thermodynamics of the system. Referring to Eq. 1 and considering the x value equal to 0 or 1, the calculation of the reaction standard enthalpy of reaction ΔH_R led to values of -54.5 kJ mol^− 1 and − 69.2 kJ mol^− 1, respectively. The calculation was performed by application of the Hess law and considering standard enthalpies of formation of -1207.4, -426.7, -285.8, -986.1, -1429.7, and − 1129.2 kJ∙mol^− 1 for CaCO₃ [33], NaOH [34], H₂O [34], Ca(OH)₂ [34], Na₂CO₃∙H₂O [33], and Na₂CO₃ [33], respectively. That would suggest that the precipitation of Na₂CO₃ would be slightly favoured at higher temperatures with respect to Na₂CO₃∙xH₂O, i.e. the reaction with x equal to 1 shows a slightly more exothermic behaviour. In fact, lower Na₂CO₃ contents were registered for those samples reacted at ambient conditions from a T_i of 45 and 60°C, respect with 80⁰C, with respective final temperatures T_f of 21.5⁰C, 42.1⁰C and 53.5⁰C. This reflects the slightly less exothermic, i.e higher ΔH_R, precipitation of Na₂CO₃ respect with Na₂CO₃∙H₂O. The results of the TG analysis were supported by the XRD patterns reported in Figs. 3A, B and C. In fact, weak signals could be detected at 2θ of 16.9⁰, main peak for Na₂CO₃∙H₂O, for all the 80⁰C_n series, whereas increased intensities were observed for the 45⁰C_n series above 2 minutes of residence time and the sample reacted at ambient conditions with T_i of 60⁰C.

The equilibrium of Na₂CO₃∙H₂O and Na₂CO₃ was studied by performing targeted simulations of a simplified system using the PHREEQC software [35] with the PITZER₍₂₀₁₈₎ database, and an overview of the outcomes is reported in Figs. 8A and B.

The simulation was conducted for a system composed of both 1 mol of Na₂CO₃∙H₂O and 1 mol of Na₂CO₃ dissolved in 0.1 kg of water at 45, 60 or 80°C, with increasing NaOH introduction up to 20 m. The simulated system does not contain Ca ions as in the experimentally tested systems, but it is useful to understand the general behaviour of Na₂CO₃∙H₂O and Na₂CO₃ in the aqueous system with the presence of NaOH at high concentrations. As seen in Fig. 8A, lower molality of NaOH would lead to the dominant precipitation of Na₂CO₃∙H₂O while Na₂CO₃ is the major precipitate at higher NaOH molality. In the NaOH concentration range of approximately 10 to 12.5 m, co-precipitation of Na₂CO₃∙H₂O and Na₂CO₃ occurs. This may explain the data presented in Fig. 7 for the samples reacted at 45°C. In fact, a low consumption of NaOH has taken place at the very start of Reaction 1, when the concentration of NaOH is still high (20 m), and Na₂CO₃ is the main product. When more NaOH is consumed, at enhanced reaction progression, the precipitation of Na₂CO₃∙H₂O is favoured. Although the reaction also produces Ca(OH)₂, its solubility is much lower than that of NaOH, and thus, the behaviour of the aqueous system is expected to be dominated by NaOH. Despite the NaOH concentration, the data for 60 and 80°C presented in Fig. 7 indicated that Na₂CO₃ was the main product throughout the tested period even after significant NaOH consumption. This may be related to the activity of water in the system. As shown in Fig. 8B, the activity of water (aH₂O) generally decreases at higher NaOH concentration, in accordance with the literature [36]. The activity of water is lower also when the temperature is higher. By comparing this with Fig. 8A, it can be deduced that Na₂CO₃ forms when the activity of water is lower (the NaOH concentration is higher), and Na₂CO₃∙H₂O can form when it is higher (the NaOH concentration is lower). Most likely, in these systems, the reaction temperatures (60 and 80 ^oC) were high enough to maintain the activity of water sufficiently low to form Na₂CO₃ even after consumption of NaOH. This also suggests that 45^oC is not high enough to form Na₂CO₃, as confirmed with the experiments conducted at ambient conditions. In fact, the system initially heated at 60^oC mostly formed Na₂CO₃∙H₂O at the end of the reaction, where the temperature became 42.1°C.

The temperature significantly affects the aH₂O values up to a NaOH molality around 10, and, above such a value, the system assumes a transitionary aH₂O value of 0.603. That corresponds to a chemical potential u of -241.3 kJ∙mol^− 1 (Eq. 6), where u₀ is the standard chemical potential of formation of pure water [37], R the gas constant, and T the temperature in K.

$$u={u}_{0}+RTln{a}_{w} \left(6\right)$$

At this point, the transition Na₂CO₃∙H₂O /Na₂CO₃ is swapped (Fig. 8A), promoting the precipitation of Na₂CO₃∙H₂O and Na₂CO₃ above and below an activity of water of 0.603, respectively. To explain the constant water activity value of 0.603, the gradual formation of Na₂CO₃ to the detriment of Na₂CO₃∙H₂O must be taken into account; as a result, water is released into the liquid bulk and it counterbalances the further addition of NaOH. In fact, when all the sodium into the system is converted to Na₂CO₃ and no more H₂O is released by the dissolution of Na₂CO₃∙H₂O, the activity of water rapidly drops down to just above 0.400 at NaOH 20 m. That corresponds to a higher water activity aH₂O within the liquid phase of the system (Fig. 8B), and therefore resulting in a higher proportioning of Na₂CO₃∙H₂O with respect to Na₂CO₃ (Fig. 8A).

3.3. Calculation of the Arrhenius parameters E_a and A

An additional set of experiments was conducted to estimate the activation energy E_a of the decarbonisation reaction at the different starting compositions reported in Table 1, selected from our previous study [13]. Based on the amount of CaCO₃ and Ca(OH)₂ in the reaction products estimated from their TGA data (Supplementary Electronic Information I - IV), the extent of reaction (α) was calculated for each reaction. As reported in Fig. 9, beneficial effects of raising temperatures were observed at different levels depending on the starting composition.

In these terms, the effects of temperature (T) is commonly expressed in the reaction rate constant (k) using Arrhenius equation (Eq. 7) [38], depending on the pre-exponential factor (A), activation energy (Ea) and gas constant (R).

$$k=A{e}^{\left(-{E}_{a}/RT\right)} \left(7\right)$$

The rate constant k is a change in the extent of reaction (α) per unit time (t). Since the reactions in this sub-set of experiment were all conducted for 1 minute, the extent of reaction obtained (Fig. 9) was directly used in Eq. 7 as representative of the reaction rate, and thus:

$$\alpha =A{e}^{\left(-{E}_{a}/RT\right)} \left(8\right)$$

Taking the natural logarithm, Eq. 8 can also be expressed as:

$$ln\alpha =lnA-\frac{{E}_{a}}{RT} \left(9\right)$$

By plotting ln α against T^− 1, the activation energy E_a and pre-exponential factor A can be calculated from the slope (-E_a/RT) and the intercept at T^− 1 = 0, respectively.

Figure 10. Linearized Arrhenius plot showing the correlation between lnk and − 1/T (K^− 1) for the samples series NaOH_10M_n, NaOH_12M_n, NaOH_15M_n and NaOH_17M_n, reflected by the equations, 9, 10, 11, and 12, respectively.

$$y=3865.1x+10.254 {R}^{2}=1.00 \left(10\right)$$

$$y=2054.3x+5.638 {R}^{2}=0.91 \left(11\right)$$

$$y=1242.1x+3.507 {R}^{2}=0.94 \left(12\right)$$

$$y=933.3x+2.591 {R}^{2}=0.93 \left(13\right)$$

Based on these, the apparent activation energies for NaOH_10M, NaOH_12M, NaOH_15M and NaOH_17M are estimated as 32.1, 17.1, 10.3 and 7.8 kJ mol^− 1, respectively, and therefore reflects the limited kinetics at lower NaOH molality.

Despite that, faster kinetics could be observed for all the conditions tested here, with respect to the conventional calcination route of CaCO₃ under inert atmospheres (164–225 kJ∙mol^− 1 [39]) and with varying CO₂ partial pressures (213.3–2142.2 kJ∙mol^− 1 [40]). Since the pre-exponential factor A represents the probability of collisions between the reacting components [41], we expected larger values of the pre-exponential factor for the higher NaOH concentration. However, this trend was not clearly observed in the present analysis, as shown in Fig. 11.

Given the lower efficiency reported, it is supposed that such collisions would not fulfil the energy barrier to be overcome for the reaction to occur. According to this data, a first-order kinetic description, only depending on a single parameter (i.e. NaOH concentration), would be supposed; on the other hand, the results reported by Hanein et al. [13] were also suggesting lower conversions at NaOH concentrations higher than 17.0 M. The other factors, such as the concentration of other reactants, and intrinsic properties of both the NaOH solutions [42] used and CaCO₃ [43] may also play a determining role for collision of the reacting components. In these terms, previous studies [44–46] already discussed the interaction between solid calcite and dissolved solutes, concluding that the incorporation would occur at the level of the solid surface by adsorption, followed by nucleation/precipitation of the products within the core of solid CaCO₃. This may justify the outcomes highlighted here. In fact, lower NaOH molalities could only allow for the saturation of the surface binding sites of calcite, whereas a higher gradient would be required to ensure the following nucleation/precipitation step. Finally, events of ionic competition could occur at too high NaOH concentrations [44, 47], as explained above, leading to hindered conversions; that would also justify the outcomes outlined in [13].

The effects of temperature on the efficiency of a low temperature alternative decarbonisation process for CaCO₃ was assessed. The reaction appeared effective for the sequestration of the process CO₂ embodied into CaCO₃, with over the 80% of it captured within the first 60 seconds of reaction time. These outcomes might suggest the potential for application to the cement industry, upon significant modifications of the current manufacturing process. In fact, the reaction discussed here would provide the decarbonised calcareous source necessary for the following clinkerisation step at above 1450°C. In other terms, a two-step process would avoid the totality of the process CO₂ usually involved in the cement manufacturing, whereas the emissions from the fuel combustion would not be limited at all. Although, an extensive literature highlights the suitability of different types of alternative fuels for the cement manufacturing process, leading to lower emissions from the combustion step.

The laboratory-scale experimental results are promising in terms of scale-up, since short residence times and mild heating conditions were found to lead to high conversion degrees of CaCO₃. The beneficial effect of a higher temperature was observed at short residence time, whereas the benefits appeared limited at longer ones. The results showed also that the temperature can be controlled only at the beginning of the reaction, and it may not be necessary to actively maintain it throughout the reaction. In fact, the reaction took place rapidly even under ambient conditions, regardless of the initial temperature considered. That would reflect a significant advantage from an industrial point of view, proving robustness of the process.

The precipitation of Na₂CO₃∙H₂O was found to be the main process CO₂ sequestration path at higher activities of water, milder temperatures and longer residence times. Despite the stoichiometry of the reaction would suggest a 1 : 1 molar ratio between the reaction products Ca(OH)₂ and Na₂CO₃.xH₂O, the experiments outlined a slightly excess of the latter species. That is likely related to the statistically higher density of negative (CO₃²⁻) sites on the surface of CaCO₃ crystals.

The Arrhenius parameters Ea and A were determined for the four starting mix compositions previously tested. Generally, the activation energy required was smaller for the system with the higher initial NaOH concentration. On the other hand, the obtained pre-exponential factors A did not indicate a clear correlation with the NaOH concentration, suggesting the importance of other factors such as the physical properties of the solutions for the collisions of the reactants, such as the density of the NaOH solutions used. However, the decarbonisation route discussed here showed much faster kinetics with respect to the conventional CaCO₃ calcination for all the conditions tested.

The alternative decarbonisation of CaCO₃ investigated in the present work may have the potential to drastically reduce the carbon footprint of the related industries, demonstrating the possibility to avoid the thermal calcination worldwide applied. The next steps will involve the testing and efficiency comparison of the reaction performed on the industrial grade raw materials, to assess the feasibility of the process from an industrial point of view.

Author’s Contribution

Theodore Hanein and Hajime Kinoshita discovered and conceptualized the technology. Theodore Hanein, Hajime Kinoshita, John L. Provis, and Marco Simoni developed the technology and designed the methodology and experiments. Marco Simoni carried out experiments and drafted the original manuscript. Theodore Hanein, Hajime Kinoshita, and John L. Provis acquired funding, and supervised Marco Simoni and Chun Long Woo about authorship. Theodore Hanein, Hajime Kinoshita, John Provis, Magnus Nyberg, and Mark Tyrer reviewed and edited the manuscript.

Author’s Contribution

Conflicts of interest

The authors declare that they have no competing interests as defined by Scientific Reports, or other interests that might be perceived to influence the interpretation of the article.

Acknowledgements

This work was funded by the Engineering and Physical Science Research Council (EPSRC) and CEMEX under grant ID EP/R025959/1. The group would also express gratitude to Alan Maries, Juan-Carlos Martinez, John Stennet, and Nestor I. Quintero-Mora for the support and knowledge received.

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No competing interests reported.

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Effect of processing temperatures on decarbonisation of calcium carbonate in sodium hydroxide solutions at ambient conditions

Status:

Version 1

Abstract

Figures

1. Introduction

2. Materials And Methods

2.1. Experimental procedure

2.2. Characterisation techniques

2.2.1. X-Ray Diffaction (XRD)

2.2.2. Thermogravimetry (TG/DTG)

3. Results And Discussion

3.1. Effect of temperature on reaction efficiency

3.2. Proportioning of Na₂CO₃∙H₂O and Na₂CO₃

3.3. Calculation of the Arrhenius parameters E_a and A

4. Conclusions

Declarations

References

Additional Declarations

Supplementary Files

Status:

Version 1

Effect of processing temperatures on decarbonisation of calcium carbonate in sodium hydroxide solutions at ambient conditions

Status:

Version 1

Abstract

Figures

1. Introduction

2. Materials And Methods

2.1. Experimental procedure

2.2. Characterisation techniques

2.2.1. X-Ray Diffaction (XRD)

2.2.2. Thermogravimetry (TG/DTG)

3. Results And Discussion

3.1. Effect of temperature on reaction efficiency

3.2. Proportioning of Na2CO3∙H2O and Na2CO3

3.3. Calculation of the Arrhenius parameters Ea and A

4. Conclusions

Declarations

References

Additional Declarations

Supplementary Files

Status:

Version 1

3.2. Proportioning of Na₂CO₃∙H₂O and Na₂CO₃

3.3. Calculation of the Arrhenius parameters E_a and A