Research Article | Open Access
Specific wetting enthalpies for investigating lime surface properties due to the high temperature reactivity of limestone impurities
A. Lagazzo*, R. Botter and DT. Beruto
*Corresponding author: Alberto Lagazzo
Department of Civil, Chemical and Environmental
Engineering (DICCA), University of Genoa, P.le J.F. Kennedy, pad. D, 16129 Genova, Italy; Tel. +390103536037; E-mail:firstname.lastname@example.org
Received: September 12th, 2017; Accepted: September 19th, 2017; Published: September 25th, 2017
Eng Press. 2017; 1(1): 4-11. doi: 10.28964/EngPress-1-102
Ⓒ 2017 Copyright by Lagazzo A, et al. Creative Commons Attribution 4.0 International License (CC BY 4.0).
In this paper, we address the role that limestone impurities play in the morphogenesis of different CaO-based particle surfaces. During the thermal decomposition of 00., such impurities can be the source of a complex pattern of reactions that form heterogeneous CaO-based surfaces. Our study reveals that this phenomenon is experimentally measurable once the level of impurities in the starting limestone is approximately 1.7% and the decomposition temperature is 1300°C. In limestones with impurity levels of 1% or less, the reactivity between them and the host CaO/CaCO3 phases is reduced to a negligible level. To explore their average surface properties, we dispersed a set of heterogeneous lime powders in liquid paraffin and measured the correspondent exothermic wetting heat. Combining these results with Brunauer, Emmett and Teller (BET) measurements of the specific surface of the limes allows one to introduce a new parameter: the specific wetting enthalpy ξ (J/m2). Limes with nearly identical specific surface areas of approximately 1 m2/g can differ greatly in their ξ values, which range from -8±0.2 J/m2 to -1±0.2 J/m2. The specific wetting heats can be used as an intensive average thermodynamic parameter to characterise the chemical and physical properties of lime surfaces with specific surface energies that are experimentally undeterminable via calorimetric measurements.
KEYWORDS:CaO; Limestone; CaCO3.
CaO-based limes produced by decomposition of limestone can have different morphologies due to the presence of little amount of impurities in the CaCO3. For some limes, powders are composed of a uniform set of grain aggregates similar to those generally observed.1-4 After the decomposition of calcium carbonate powders in air. These morphologies can be associated with the solid-solid sintering process catalysed by the escaping CO2.5,6 Sometimes, the edges of particles aggregates, where the particles have irregular shapes, tend to disappear and the external particle surface can be observed to have interesting patterns that resemble to Moire fringes.7 This is connected to the faceting mechanism and/or dissolution/recrystallization process of a low melting eutectic phase.8-11 Thus, those findings indicate that limestone impurities strongly influence the physical and most likely the chemical properties of the obtained lime surfaces.12 For many old and new technologies,13-22 lime is used for its surface properties, and there is a need to develop scientific knowledge in this complex and interesting field. Initially, evaluating the surface free energies of limes and the effect of lime impurities on these values seems to be a fairly reasonable way to proceed. Indeed, Brunuaer, Kantro, Weise23 and more recently, other researchers24 have measured the heat of dissolution for oxides in the liquid phase (HNO3, HCl and H2O). Coupling these values with Brunauer, Emmett and Teller (BET) surface area measurements, they obtained acceptable values for the specific surface enthalpies of these oxides. This value is similar to the specific surface free energy of the oxide neglecting entropy at room temperature. However, when the oxide is impure, as happens with limes, the surface enthalpy can be obscured by other bulk terms due to reactions between the impurities and the liquid phase. Therefore, it is impossible for this technique to determine the specific surface free energy of lime.
In two of our papers,25,26 we studied the dispersions of silica and kaolin powders in liquid paraffin containing aromatic and naphthenic compounds as a function of their concentration in the dispersed phase at 30 °C. The paraffin interacted with the solid surfaces to form thin solid-liquid interfaces and produced an exothermic wetting heat that accounted for the stability of the colloidal dispersion.27 Theoretically, this mechanism involves three interactions: the van der Waals forces between solid particles, acid-base (Drago-Wayland)28 attractive interactions between the silica and both the naphtenic paraffin and aromatic carbons, and repulsive steric interactions29 between the layers adsorbed onto the silica surfaces. What matters to us is that non-bulk reactions occur between the liquid dispersant and the dispersed solid particles. Thus, for any given and selected liquid phase, the only information that can be derived from the measurable wetting heat relates to the nature of the dispersed solid particle surface. Whether the impurities are accessible to the lime surfaces or at the lime grain boundaries, their effect will change the corresponding exothermic wetting heat. This simple idea is the basis of the experimental method we use herein to obtain information on the surface impurities of lime and to gain insight into, as discussed below, how these impurities affect the morphogenesis of the CaO-based particles during the thermal decomposition of limestone. To the best of our knowledge, this is a completely new approach to investigating lime surface properties and might be of interest to those who design lime particles with optimal properties for saving energy, adsorbing pollutants and other technological applications.30-34
Limestones with a high calcium carbonate content ranging between 99.4 wt.% and 100 wt.%, are used to obtain the corresponding limes with CaO mol% ranging between 98.3 and 99.3. The chemical compositions of the limestones are reported in (Table 1). For all limestones, MgO is the primary impurity. LMST1 has a silica and alumina content greater than the other limestones.
Demineralised water and high purity liquid paraffin without the aromatic and naphtenic groups (0.88 g/cm3, boiling temperature of 68 °C, ignition point of 135 °C and viscosity of 20 mPas) were used as the liquid phase for the wetting experiments. The choose of this type of liquid paraffin is due to its high purity and to a degree of viscosity that allows a good dispersion of the lime powder inside the calorimetric cell.
Powder preparation: Limestone irregular blocks of 5-15 mm were decomposed in air for 2 h and 30 min at a constant temperature between 980 and 1300 °C according the samples producing limes with an average size between 1 and 8 µm. For some of these samples, a chemical analysis of the impurities was performed.
Reagent grade (r.g.) CaCO3 powders were decomposed both at 900 °C for 30 min and 1350 °C for 4 hours to obtain small-surface CaO samples for comparison. All the CaO-based oxides were stored in sealed containers in a desiccator. Table 2 lists the lime families based on the limestone decomposition temperature and rock.
Calorimetric measurements: The enthalpy of hydration23 and the heat of dispersion in liquid paraffin25-26 of the CaO-based oxides were evaluated through calorimetric analysis with a Setaram C80 calorimeter25 equipped with stainless steel wetting cells. Two under vacuum closed glass ampuoles of 5 mm in diameter, with a narrowing hook at one end to facilitate breaking, one filled with about 20 mg of lime powder sample and the other empty, were introduced into two identical cells (measure and reference respectively). The cells consisting in a cylinder of 2 cm in diameter and 10 cm in height, with a mobile piston inside, were then filled with 5 ml of liquid paraffin and then placed in the calorimetric at a temperature of 30±0.01 °C. After the reaching of the thermal equilibrium (about 12 h), the piston was manually lowered to brake curved tip of the ampoule and the powders were wetted by the liquid phase. The heat flow was recorded until the flow curve reached the initial reference value, after about 12 h, suggesting that the reaction was completed. The use of an empty ampoule in the reference cell is necessary to remove the mechanical effects of the glass breaking and of the liquid entering in the glass container. All of the tests were repeated three times. The experimental error is contained under the 0.5%.
Nitrogen adsorption measurements at 78K: BET measurements were obtained with a symmetrical Sartorius microbalance described in detail previously.33 The experimental error of the specific surface area was estimated to be approximately 3%.
Mercury porosimeter: Porosity and pore size distribution of the samples were analysed with a Thermo Scientific Pascal 240 mercury porosimeter equipped with a standard dilatometer for powders.
Scanning Electron Microscope (SEM): Morphological analyses of a carbon-coated sample were conducted using a scanning electron microscope (SEM) Hitachi-2500. The operating conditions were 20 kV.
Theory of dissolution: It is known that the enthalpy of dissolution of a solid depends on its enthalpy of formation and on its surface area, whose contribute reduces the total amount of the heat developed during the dissolution.35 In the case of the CaO dissolution, when the surface enthalpy is accounted for, any pure CaO sample with a given BET specific surface area will dissolve in liquid water with a global enthalpy change equal to:
–ΔHexp = –ΔHCaO,bulk + σCaOSBET < 0
Where σCaO is the specific surface energy of the CaO and is always positive,36 SBET is the specific surface area of the CaO powders as evaluated using BET theory37 and ΔHexp and ΔHCaO,bulk respectively the global experimental enthalpy of the CaO dissolution and the theoretical enthalpy of CaO dissolution. Thus, the dissolution of a CaO powder must yield an enthalpy change greater than that derived from the thermodynamic tables.
RESULTS AND DISCUSSION
CaO powders containing negligible impurities react with liquid water according to the general equation below:
CaO(s) + H2O(l) = Ca(OH)2 (s)
The enthalpy change associated with reaction 1) can be derived from a selected thermodynamic table,37 and has a value of -1160 J/g at 30 °C.
Brunaer, Kantro and Weiss23 conducted experiments with fine and coarse CaO crystallites of 0.5 m2/g and of 7.8 m2/g, respectively, and found a consistent value for the CaO specific surface energy equal to 1310±200 mJ/m2. No significant variations were observed if a 2N solution of HNO3 was used instead of water. However, reproducing those experiments and obtaining coherent results when the specific surface area of the CaO reagents does not display significant differences is very arduous. We measured the surface energy of commercial grade CaO powders of 6 and 11 m2/g, and the results obtained of 1200±350 mJ/m2 were affected by an experimental error too high lead to a CaO specific surface energy, which agrees well with Brunuaer, Kantro and Weiss’s data.
For this reasons the difference that can be attributed to the thermal effect due to the presence of the impurities is about equal to the experimental error of the calorimetric measure, that can be estimated to around 4%. Therefore, these difference in the global heats developed during the limes dissolution in water cannot be accounted to evaluate the quality and the micro-structural nature of the samples. In any case, also overcoming these experimental difficulties, the evaluation of the impurities through the effect that they produce on the enthalpy of CaO dissolution in water is cannot be calculated.
Considering a specific surface energy, σ Lime, equal for all of the samples, then the plot standard heat of the dissolution reaction vs. surface area should form a straight line.24 If this not occurs, reasonably can be conclude that the σ Lime is different from lime to lime, even though they were obtained from the thermal decomposition of the same limestone parent.
In our previous paper,39 we demonstrated that the lime grain surfaces can be quite heterogeneous due to the presence of limestone impurities. The derivation of σ lime,i values from a plot of ΔHlime,i,exp vs. SBET, lime,i is a difficult task because of the following:
σ lime,i = (ΔHlime,i, bulk – ΔHlime,i, exp ) / SBET, lime,i
If the impurities enter the bulk phase of the lime and react, then the state of the standard reference lime also changes. Thus, ΔHlime,i, bulk is an unknown parameter. As a partial conclusion, any investigation into the surface properties of lime via the σ lime,i variable remains highly speculative and cannot be quantised from the experimental heat of reaction for lime powders dissolving in a liquid phase combined with the specific surface area measurements.
A kinetic approach, consisting in the measuring of parameter t60, the time necessary in order that a CaO–water dispersion reaches the temperature of 60 °C, is well known and used in the industry.
The data of water reactivity tests (t60) for limes obtained from different source of limestones (limestone A and limestone B), which have different amount of impurities is reported in (Figure 1). It is reasonable to assume that also the corresponding limes will have different amount of impurities, thus the comparison of their t60 values will give an idea of how the chemical impurities might influence the reactivity of the system lime-water. It is evident that limes with equal surface area, derived from limestones with greater amount of impurities, have a t60 parameter lower than the one for limes with less impurities. These data clearly prove that the degree of impurities in the limes does matter with its water reactivity.
It is interesting to discuss the evolution of the lime surface area as a function of the decomposition temperature of the parent limestone.
As reported in (Table 2), the specific surface area decreased as the decomposition temperature increased, however, the change is quite different between limes. For the sake of comparison, CaO (r.g.) was characterised by a surface area of 11 m2/g when the CaCO3 was decomposed at 900 °C and 6 m2/g when kept for 4 hours at 1350 °C. These values are quite high relative to those obtained for the limes. Evidently, the impurities are an important variable in producing these effects. In particular limestone A decomposed at 980 °C produce lime with specific surface area of 2.8 m2/g while lime obtained at the same temperature from limestone B display a value of 4.8 m2/g. These experimental results support the hypothesis that limestone impurities are active at low temperatures where the catalytic action of the CO2 on the densification of the CaO grains is the rate-determining step,40 while others are more reactive at higher temperatures.
From the chemical analysis of the starting limestone, it can be observed that the molar fraction of the impurities covers only a narrow range. Thus, it is difficult to foresee which is the most effective for reducing the lime surface area.
To evaluate the surface reactivity of the limes samples, overcoming the difficulty connected with the calorimetric testes of CaO-based lime reacting with water, we describe the results obtained when the lime particles are dispersed in a liquid phase where the liquid molecules only adsorb onto the accessible lime surface. The dispersant liquid phase used herein is liquid paraffin, and the solid volume fraction, θ, of the dispersed lime particles was approximately 20 wt.%. This value can be considered greater than the percolate threshold.29 Because the liquid paraffin does not contain aromatic and naphthenic groups, only two forces should be present: the attractive van der Waals forces and the steric repulsive forces between solid CaO particles and the adsorbed alkane molecules.
We introduce the variable ξ [J/m2], defined as:
ξ = ΔHW / SBET
where ΔHW [J/g] is the wetting heat of a given lime and SBET is the associated specific surface area [m2/g]. This variable has been used in others paper25,41 and is strictly linked with the “quality” of the solid surface being investigated.
Figure 2 describes the parameter ξ vs. SBET for the limes from limestone A and B. It can be observed that the points A are best-fitted by an equation that obeys the following phenomenological law:
ξ = – (α SBET) / (S BET – β) [J/m2]
where α and β are the two best-fit constants for the lime/liquid paraffin system, which are equal to 0.43 and 0.675, respectively. It interesting to observe as the samples from limestone B do not follow the best-fit line.
ξ is an average intensive macroscopic parameter describing the capability of a lime surface to adsorb the liquid molecules and form the solid/liquid interface that make stable the dispersed solid powders. Thus, this variable depends on the accessibility of the internal surfaces (pores), external surfaces, and surface impurities to the liquid molecules, which can change the interactive forces between the solid surface and the liquid molecules.
The fact that sample A1 has a ξ value below that of lime A5 agrees with the morphological evidence (Figure 3), which supports the idea that limes with lower, but more accessible, surfaces can adhere more paraffin molecules per square metre than those with more surface area but smaller pores (Figure 4).
The grains of lime A1 (Figure 3a) can be describe as irregular polyhedral with an average dimension of 0.5-1 µm clustered in porous sintered aggregates similar to the one formed by the catalytic action of escaping CO2. The grains of lime A5 (Figure 3b) have a different shape, and some are composed of large, flat surfaces with an average dimension of 20 µm×0.2 µm.
The porosity and corresponding average pore size of sample A1 (0.55 cm3/g and 0.4 nm, respectively) are higher than those for sample A5 (0.15 cm3/g and 1 nm, respectively) and supports the inference from the SEM images in (Figure 3b).
However, morphological reasons alone are not enough to explain why the ξ values of limes from limestone A suddenly drop to -8 J/m2 when the BET surface area is approximately 1 m2/g, while for other limes from limestone B, the ξ values are fairly constant at approximately -1 J/m2 even if the specific surface area is approximately 1.2 m2/g.
In the author’s opinion, this interesting datum depends on the content, nature and high-temperature reactivity of the limestone impurities.
Table 1 contains the chemical composition of lime A5 and B5 derived respectively from limestone A and B. Is evident that the impurity level in sample B5 is clearly lower than that of sample A5 with a containing of SiO2 equal respectively to 0.051 wt.% and 0.322 wt.% and of Fe2O3 equal to 0.038 wt.% for B5 and 0.206 wt.% for A5. Thus, lime A5 has a ξ value of -8 J/m2 not only because of its low SBET surface area but also because the chemical active sites on its surface and the emerging grain boundaries are different from those for ξ values of approximately –1 J/m2. The high temperature reactivity of both the limestone impurities and the CaO and/or CaCO3 phases in these rocks are most likely the explanation for this experimental evidence. In agreement with previous studies on the Al-SiO2 system,36 it has been shown that the aluminium ions can diffuse into NaCl-type oxides at the tetrahedral sites that form the AlO4– groups. Other studies8-11 have proven the dramatic effects that adding impure powders at higher concentrations has on the shape of the particles after the thermal decomposition of pure CaCO3 powders.
When limes are obtained from limestone with impurities equal to or less than 1% that do not react during the decomposition of the parent limestone, the morphologies of the lime particles should be governed by the catalytic effect of escaping CO2, and the adsorption of the paraffin molecules would be primarily driven by the accessibility of the molecules to the lime surface.
In this paper, we analysed the role that limestone impurities have on the morphogenesis of different CaO-based particles. These impurities are a feature that geological history has left in the limestone. During the thermal decomposition of the parent limestone, these impurities can be the source of complex reaction patterns that lead to the morphogenesis of more heterogeneous aggregates of CaO-based grains. Our study reveals that decomposing limestones with 1.7% of impurities in air between 980 °C-1300 °C leads to lime grains with different shapes, while the limestones with impurities of 1% or less do not. In the latter case, the grain shapes look like those obtained from the decomposition of reagent grade CaCO3.
To explore the average surface energies of the heterogeneous lime particles, we have proposed dispersing the lime particles into a liquid vaseline and measuring their wetting heats. This parameter combines with the lime BET specific surface to create an important new parameter, ξ, which is the specific wetting heat of the dispersed powders. The plot of ξ vs. SBET turns out to be a unique phenomenological rule for limes obtained from all of the different sources. Based on the shape and value of ξ for each lime, it is possible to derive whether a scout molecule such as vaseline adsorbs under the action of weaker or stronger attractive forces. Stronger attractive fields were connected to the presence of surface active sites formed via high-temperature reactions between the limestone impurities and the CaO/CaCO3 phase. These reactions are linked to the morphogenesis of the different sets of CaO-based particles, which are often observed from data concerning the structural and chemical analysis of lime.
CONFLICTS OF INTEREST
The authors declare that they have no conflicts of interest.
1. Powell EK, Searcy AW. Surface-areas and morphologies of CaO produced by decomposition of large CaCO3 crystal in vacuum. J Am Cer Soc. 1982; 65(3): c42-c44. doi: 10.1111/j.1151-2916.1982.tb10395.x
2. Borgwardt RH. Calcination kinetics and surface area of dispersed limestone particles. AIChE J. 1985; 31(1): 103-110. doi: 10.1002/aic.690310112
3. Fuller EL, Yoos TR. Surface properties of limestones and their calcination products. Langmuir. 1987; 3: 753-760.
4. Beruto DT, Searcy AW, Kim MG. Microstructure, kinetic, thermodynamic analysis for calcite decomposition: Free surphace and powder bed experiments. Thermochimica Acta. 2004; 424(1-2): 99-109. doi: 10.1016/j.tca.2004.05.027
5. Erwing J, Beruto DT, Searcy AW. The nature of CaO produced by calcite powders decomposition in vacuum and in CO2. J Am Cer Soc. 1979; 62(11-12): 580-84. doi: 10.1111/j.1151-2916.1979.tb12736.x
6. Borgwardt RH. Calcium oxide sintering in atmosphere containing water and carbon dioxide. Ind Engineering Chem Res. 1989; 28(4): 493-500. doi: 10.1021/ie00088a019
7. Figueiredo M, Zerubia J, Jain AK. Energy Minimization Methods in Computer Vision and Pattern Recognition. Berlin, Germany: Springer; 2001.
8. Beruto DT, Searcy AW, Fulrath RM, Basu T. Effects of carbon dioxide and sodium chloride on the sintering of calcium oxide. LBL. 1977; 76: 101.
9. Beruto DT, Knudsen GF, Searcy AW. Effect of LiCl on the rate of calcite decomposition. J Am Cer Soc. 1982; 65(4): 219-22. doi: 10.1111/j.1151-2916.1982.tb10409.x
10. Beruto DT, Barco L, Belleri G, Longo V. Interactions of LiBr with calcite and calcium oxide powders. Cer Int. 1983; 9(2): 53-58. doi: 10.1016/0272-8842(83)90023-8
11. Beruto DT, Kim MG, Barco L. Effect of Li2CO3 on the reaction between CaO and CO2. J Am Cer Soc. 1984; 67(4): 274-278. doi: 10.1111/j.1151-2916.1984.tb18846.x
12. Jing JY, Li TY, Zhang XW, et al. Enhanced CO2 sorption performance of CaO/Ca3Al2O6 sorbents and its sintering-resistance mechanism. App Energy. 2017; 199: 225-233. doi: 10.1016/j.apenergy.2017.03.131
13. Boyton RS. Chemistry and Technology of Lime and Limestone. 2nd ed. New York, USA: Wiley-Interscience; 1980.
14. Konigsberger E, Konigsberger LG, Gamsjager H. Geochim Cosmochim Acta. 1999; 63(19-20): 3105-3119.
15. Salvador C, Lu D, Anthony EJ, Abanades JC. Enhancement of CaO for CO2 capture in an FBC environment. Chem Eng J. 2003; 96(1-3): 187-195. doi: 10.1016/j.cej.2003.08.011
16. Lanas J, Alvarez JI. Dolomitic limes: Evolution of the slaking process under different conditions. Thermochimica Acta. 2004; 423(1-2): 1-12. doi: 10.1016/j.tca.2004.04.016
17. Fennel PS, Davidson JF, Dennis JS, Hayhurst AN. Regeneration of sintered limestone sorbents for the sequestration of CO2 from combustion and other systems. J Energy Institute. 2007; 80(2): 116-119.
18. Daniele V, Taglieri G, Quaresima R. The nanolimes in cultural heritage conservation: characterization and analysis of the carbonation process. J Cult Heritage. 2008; 9(3): 294-301.
19. Ströhle J, Galloy A, Epple B. Feasibility study on the carbonate looping process for post-combustion CO2 capture from coal-fired power plants. Energy Procedia. 2009; 1(1): 1313-1320. doi: 10.1016/j.egypro.2009.01.172
20. Folger P. Carbon Capture: A tecnology assessement. Congressional Res Service 7-5700. 2013. Web site.http://digitalcommons.unl.edu/crsdocs/19/. Acccessed September 11, 2017.
21. Skoufa Z, Antzara A, Heracleous E, Lemonidoua AA. Evaluating the activity and stability of CaO-based sorbents for postcombustion CO2 capture in fixed-bed reactor experiments. Energy Procedia. 2016; 86: 171-180. doi: 10.1016/j.egypro.2016.01.018
22. Granados-Pichardo A, Granados-Correa F, Sanchez-Mendieta V, Hernandez-Mendoza H. New CaO-based adsorbents prepared by solution combustion and high-energy ball-milling processes for CO2 adsorption: Textural and structural influences. Arabian J Chem. 2017. Web site. http://www.sciencedirect.com/science/article/pii/S1878535217300631. Accessed March 21, 2017.
23. Brunauer S, Kantro DL, Weise CH. The surface energies of calcium oxide and calcium hydroxide. Canadian J Chemistry. 1956; 34(6): 729-742. doi: 10.1139/v56-096
24. Beruto DT, Rossi PF, Searcy AW. MgO of very high energy content from decomposition of Mg(OH)2. J Phys Chem. 1985; 89(9): 1695-1699. doi: 10.1111/j.1151-2916.1982.tb10395.x
25. Beruto DT, Lagazzo A, Botter R. Silica–paraffin and kaolin–paraffin dispersions: Use of rheological and calorimetric methods to investigate the nature of their dispersed microstructure units. Colloids Surface A. 2012; 396: 153-160. doi: 10.1016/j.colsurfa.2011.12.062
26. Beruto DT, Lagazzo A, Botter R. Nanoscopic water layers adsorbed onto mesoporous silica aggregates and their effect on the stability and the static yield stress of their dispersion in liquid paraffin. Colloids Surface A. 2012; 407: 133-140. doi: 10.1016/j.colsurfa.2012.05.020
27. Dékàny I. Microcalorimetric control of liquid sorption on hydrophilic/hydrophobic surfaces in nonacqueous dispersions. In: Dekker M, ed. Thermal Behaviour of Dispersed Systems. New York, USA: CRC Press; 2001.
28. Drago RS, Wayland BB. A Double–Scale Equation for Corre-ating Enthalpies of Lewis Acid–Base Interactions. J Am Chem Soc. 1965; 87(16): 3571-3577. doi: 10.1021/ja01094a008
29. Napper DH. Polymeric Stabilization of Colloidal Dispersions. New York, USA: Academic Press; 1983.
30. Abanades JC, Anthony EJ, Lu DY, Salvador C, Alvarez D. Capture of CO2 from combustion gases in a fluidised bed of CaO. AIChE J. 2004; 50(7): 1614-1622. doi: 10.1002/aic.10132
31. Hughes RW, Lu D, Anthony EJ, Yu Y. Improved long-term conversion of limestone-derived sorbents for in situ capture of CO2 in a fluidized bed combustor. Ind Eng Chem Res. 2004; 43(18): 5529-5539. doi: 10.1021/ie034260b
32. Alvarez D, Abanades JC. Pore-size and shape effects on the recarbonation performance of calcium oxide submitted to repeated calcination/recarbonation cycles. Energy & Fuels. 2005; 19(1): 270-278. doi: 10.1021/ef049864m
33. Husmann M, Zuber C, Maitz V, Kienberger T, Hochenauer C. Comparison of dolomite and lime as sorbents for in-situ H2S removal with respect to gasification parameters in biomass gasification. Fuel. 2016; 181: 131-138. doi: 10.1016/j.fuel.2016.04.124
34. Beruto DT, Botter R, Searcy AW. H2O catalyzed sintering of 2-nm-crosssection particles of MgO. J Am Ceram Soc. 1987; 70(3): 155-159. doi: 10.1111/j.1151-2916.1987.tb04950.x
35. Costa GCC, Ushakov SV, Castro RHR, Navrotsky A, Muccillo R. Calorimetric measurement of surface and interface enthalpies of Yttria-stabilized Zirconia (YSZ). Chem Mater. 2010; 22(9): 2937-2945. doi: 10.1021/cm100255u
36. Defay R, Prigogine I. Surface tension and adsorption. London: Longmans; 1966.
37. Brunauer S, Emmett PH, Teller E. Adsorption of Gases in Multimolecular Layers. J Am Chem Soc. 1938; 60(2): 309-319. doi: 10.1021/ja01269a023
38. The NBS tables of chemical thermodynamic properties. J Phy Chem Ref Data. 1982; 11(2).
39. Beruto DT, Botter R, Cabella R, Lagazzo A. A consecutive decomposition-sintering dilatometer method to study the effect of limestone impurities on lime microstructure and its water reactivity. J European Cer Soc. 2010; 30(6): 1277-1286. doi: 10.1016/j.jeurceramsoc.2009.12.012
40. Busca G, Lorenzelli V. Infrared spectroscopic identification of species arising from reactive adsorption of carbon oxides on metal oxide surfaces. Mater Chem. 1982; 7(1): 89-126. doi: 10.1016/0390-6035(82)90059-1
41. Peng L, Qisui W, Xi L, Chaocan Z. The molar formation enthalpy of nano-SiO2 with different surface area. J Therm Anal Calorim. 2009; 95(2): 667-670. doi: 10.1007/s10973-008-9325-3