SO2 Capture by Calcined Limestone from Hot

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Simultaneous HCl/SO2 capture by calcined limestone from hot gases Geng-Min Lin, and Chien-Song Chyang Energy Fuels, Just Accepted Manuscript • DOI: 10.1021/acs.energyfuels.6b02565 • Publication Date (Web): 22 Nov 2016 Downloaded from http://pubs.acs.org on November 27, 2016

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Simultaneous HCl/SO2 capture by calcined limestone from hot gases Geng-Min Lin, Chien-Song Chyang*

Department of Chemical Engineering, Chung Yuan Christian University, Chung Li District, Taoyuan City, Taiwan 32023, R.O.C.

Corresponding author: Tel.: +886 3 2654119. Fax: +886 3 4636242. E-mail: [email protected].

ABSTRACT Some confusion still exists in the literature regarding the simultaneous absorption of SO2 and HCl by calcined limestone in hot flue gases. Therefore, further efforts to understand and clarify the absorption behaviors of SO2 and HCl are presented in this study. Experiments were carried out under conditions that simulated combustion in a fixed-bed reactor coupled with an on-line Fourier Transform Infrared (FTIR) spectrometer. The fixed-bed reactor used was specially designed to be capable of handling high-temperature operations involving gas-solid reactions. This paper described and explained the phenomena of variations in the concentration profiles of SO2 and HCl during the simultaneous absorption process. This work suggested possible heterogeneous and homogeneous reactions that may occur under the conditions being studied. Particular emphasis is placed on elucidating why concurrent sulfation enhancement and

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chlorination suppression occurs; experimental evidence confirmed that these processes were caused because of the behavior of the sulfation process of chloride. In the present study, to continuously monitor SO2 and HCl discharged during reaction processes plays a crucial role in observing particular phenomena, providing important information regarding absorption characteristics. From the experimental results, it was observed that upon increasing the temperature from 650 to 700 ˚C, chlorination was considerably suppressed because chloride started to vigorously react with SO2, leading to the indirect enhancement of sulfation and the re-release of HCl into the gas phase. It was also surprising to find that at or above 700 ˚C, the overall uptake efficiency remained nearly unchanged despite the difference in reactivity of SO2 and HCl. Besides, the importance and impact of the gas atmosphere on the uptake efficiency was also analyzed and compared. Keywords: Sulfation; Chlorination; Calcined limestone; Combustion; FTIR

1. INTRODUCTION

Hydrogen chloride and sulfur dioxide are principal acidic pollutants that are commonly present in flue gases following waste incineration or fuel combustion processes. Other thermal treatment methods, such as biomass pyrolysis and gasification, can also generate these gaseous pollutants.1, 2 Coal-fired power plants are the leading source of SO2 emission, and HCl released


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during thermal disposal processes is derived from organic as well as inorganic chlorides in fuel. Incineration of wastes containing organic chloride is dominant source of HCl emission and has a greater potential to generate HCl.1-4 In certain cases, such as the incineration of municipal solid waste (MSW) or co-combustion of fossil fuel and refuse, both SO2 and HCl usually co-exist in exhaust gases.5 The emission levels can be in the range of 10 to 103 ppm, mainly depending on the sulfur/chlorine contents of the fuel. In addition to its well-known adverse impact on human health and the environment, HCl emission may also cause severe problems, such as fouling and hot corrosion, due to the formation of deposits in the equipment, resulting in increased capital costs. Another issue of concern pertaining to the chlorine content is the possible emission of toxic chlorinated organic compounds, especially polychlorinated dibenzo-p-dioxins and dibenzofurans (PCDD/Fs). This issue is of interest due to the highly toxic nature of PCDD/Fs to both human health and the environment, and thus, many countries have imposed stringent regulations for dioxin emissions. It is generally accepted that the presence of chlorine gas (Cl2) probably becomes a precursor for the formation of chlorinated aromatic compounds. HCl plays an important role in forming dioxin compounds because the oxidation of HCl through fly ash or metal catalysts can yield chlorine gas via the Deacon reaction, which is the initial step of dioxin formation mechanisms.6, 7 In other words, the presence of HCl probably indirectly contributes to the formation of PCDD/Fs. By 3 ACS Paragon Plus Environment

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contrast, sulfur has been thought to be beneficial for dioxin reduction as it acts as an inhibitor during the combustion process.8-10 For the above reasons, direct removal of SO2 and HCl from the furnace is an important requirement for reducing emissions. A feasible method to reduce the risk of PCDD/F formation and avoid the problem of hot corrosion is to decrease both SO2 and HCl emissions in the furnace to acceptable levels by using effective additives before the exhaust cools. For current combustion technologies, the fluidized bed combustor (FBC), which works by feeding sorbents into a hot furnace, is particularly suited for in-situ desulfurization and dechlorination due to its characteristics of continuous operation and beneficial gas-sorbent contact. Calcium-based sorbents have been considered to be one of the most promising classes of sorbents mainly due to the advantages of being abundantly available, inexpensive, and above all, suitable for high-temperature operation. Also, the addition of calcium-based sorbents into the FBC has a beneficial effect on the reduction of dioxin compounds, as has been demonstrated in previous investigations11-13. It is a general experience that the temperature is a major factor affecting the sorbent utilization for high-temperature acid gas control. On the other hand, it is also necessary to consider how a change in the flue gas composition might affect the sorbent reactivities because the gas atmosphere plays an important role during the absorption process.14-17 In the typical temperature-range for FBC operation, lime and limestone can also provide efficient 4 ACS Paragon Plus Environment

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retention of SO2 and HCl despite having different temperature regimes. It has been reported by several earlier works that a temperature of ~ 850 ˚C is most suitable for the sulfation reaction of calcium-based sorbents. The optimal temperature for the capture of HCl has been reported to be ~ 600-650 ˚C, indicating that there are two distinct temperature domains for individual sulfation and chlorination processes.7, 14, 18-23 This initiated an interest in finding the proper conditions to effectively control both SO2 and HCl emissions at high temperatures, preferably within the range of the FBC operation temperatures. So far, a large volume of published research has been devoted to the fundamental investigation of either sulfation or chlorination; however, very few reports have discussed the simultaneous control of SO2 and HCl at high temperatures, suggesting that this area has still not been completely explored. Even so, some reports conclude that the SO2 retention efficiency is markedly promoted in the presence of HCl, while the concurrent chlorination reaction is suppressed.24-27 The reason for this phenomenon is not yet clear, indicating the need for further experimental evidence and analysis. To analyze how they affect each other is important and necessary to understand the competitive mechanism between sulfation and chlorination of calcined limestone. Thus, the present work was undertaken to elaborate on the simultaneous chlorination and sulfation behaviors of calcined limestone and their interactions under conditions similar to those typically encountered in fluidized bed combustors. The concept of gas-phase material balance 5 ACS Paragon Plus Environment

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will be used in this work to determine the levels of SO2 and HCl absorption, differing from those commonly used in previous studies. In this case, using the FTIR to simultaneously and continuously monitor SO2 and HCl discharged during reaction processes is particularly suited mainly because two or more reacting gases are involved in the gas-solid reaction. During the absorption process, the uptake efficiency of sorbent can then be calculated individually by means of analyzing the time-dependent SO2 and HCl concentration variations. The primary objective of this work was to demonstrate how the sulfation/chlorination behavior is affected by the coexistence of SO2 and HCl. Attempts were made to further analyze the reactivities of sulfur and chlorine in the context of their uptake efficiencies. Finally, this work also describes concurrent sulfation enhancement and chlorination suppression.

2. EXPERIMENTAL SECTION

Experiments on the simultaneous absorption of SO2 and HCl by calcined limestone were carried out in a self-made reactor system (as illustrated in Figure 1), mainly aimed at investigating the behavior of the gas-solid reaction. The reactor system was composed of several units, including a specially designed fixed-bed reactor, simulated flue gas supply system, electrical furnace, gas analyzer, and exhaust gas cleansing system.

2.1 Sorbent Materials 6 ACS Paragon Plus Environment

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In this work, domestic limestone was used as a calcium oxide (CaO) precursor. The solid sample was ground and screened to a particle size of 500 to 590 µm (35-30 mesh). In each experiment, approximately 0.1 g (±0.001 g) of solid sample, weighed by a high-precision analytical balance (Sartorius, BAS124S-CW), was pre-mixed with quartz sand (TOCHANCE technology Co., Ltd.), and the solid mixture was then loaded onto a sample holder. The calcined limestone needed for the experiment was obtained prior to the experiment by complete calcination at 850 ˚C in a pure-N2 atmosphere.

2.2 Apparatus

This fixed-bed reactor already met the following necessary criteria: (1) excellent leakproofness, (2) rapid temperature rise for the inlet reacting gas, (3) highly efficient cooling for the exhaust gas, and (4) low space requirement. As schematically depicted in Figure 2, the pre-mixed reacting gases were fed into the reactor via a side inlet and subsequently entered an annulus region for heating. For a particular gas flow rate, the reacting gas could be heated to the desired temperature by allowing a sufficient residence time. On the other hand, the exhaust flowing downwards through the inner tube also exchanged heat with the inlet gas mixtures. The exhaust gas was cooled in an air cooler with five tubes to achieve a temperature of approximately 32 to 36 ˚C, mainly depending on the operating temperature (650-850 ˚C), before being emitted


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from the reactor. The fixed-bed reactor was coupled with an on-line Fourier Transform Infrared spectrometer (Asea Brown Boveri, MB 3000 series) to monitor both the SO2 and HCl concentrations in the exhaust as a function of time. Finally, the exhaust gases were scrubbed through impingers containing 0.5 M NaOH solution to neutralize any remaining acid gases before discharging to the atmosphere. Prior to and after each run, the entire system was held under a N2-gas sweep until no SO2 and HCl could be detected. Two k-type thermocouples were inserted at different positions inside the reactor. One was used to control the reaction temperature, while the other was used to monitor the exhaust temperature.

2.3 Gas Supply

Cylinders containing 2000 ppm of SO2 in N2, 2000 ppm of HCl in N2, 99.9999 % O2, 99.9999 % CO2, and 99.99 % N2 were used to supply the reacting gas mixture. Each gas flow emitted from the corresponding gas cylinder was regulated and controlled by a thermal mass flow controller (Bronkhorst, Netherlands) before entering a blending chamber to achieve the desired gaseous concentrations. Throughout the experiment, the total gas flow rate was maintained at 1000 sccm (i.e., 1000 ml/min at 0 ˚C and atmospheric pressure). Moreover, the composition of CO2/O2 in the gas mixture was kept constant at 21% (e.g., 6% O2 and 15% CO2 in the reacting gas) for all tests. A fixed-bed reactor composed of quartz was designed to be


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capable of handling high-temperature operations involving gas-solid reactions, also accounting for the fact that the reactor system was slightly pressurized by the gases supplied from cylinders. Therefore, this customized method has been successfully applied to minimize the possibility of gas leakage.

3. RESULTS AND DISCUSSION

Figure 3 presents the results obtained from the fixed-bed operation, showing the SO2 and HCl concentration profiles during the simultaneous absorption process (also referred to as the breakthrough curve). The term "breakthrough curve", which is a plot of the effluent concentration as a function of time, was applied to describe the behavior of the simultaneous absorption of SO2 and HCl in terms of sorbent reactivity. These tests were performed with a gas mixture of 500 ppm SO2, 500 ppm HCl, 6% O2 and 15% CO2 in balanced N2 at three representative temperatures. In this case, the sulfation and chlorination processes of calcined limestone are usually expressed by the following overall reaction schemes:

CaO( s ) + SO2 ( g ) + 1 / 2O2 ( g ) → CaSO4 ( s )

(1)

CaO( s ) + 2HCl( g ) → CaCl2 ( s ) + H 2O( g )

(2)

or


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CaO( s ) + 2 HCl( g ) → CaCl2 ⋅ H 2O( s )

(3)

Under the above conditions, reactions (1)-(3) can be treated as an irreversible reaction because the reaction equilibrium is strongly favored to the right. It is also important to note that the end product (CaCl2) is not the only solid-phase product of chlorination because the formation of intermediates, including the most common intermediate, calcium hydroxychloride (CaClOH), has been confirmed to occur in the reaction of calcium-based sorbents with HCl.17, 28-35 Therefore, the chlorination reaction probably proceeds not only with CaCl2 but also with the intermediates. The reaction mechanism for the formation of intermediates and end products, as proposed in the literature, is listed below:30

CaO( s ) + HCl ( g ) → CaClOH ( s )

(4)

CaClOH( s) + HCl( g ) → CaCl2 ⋅ H 2O( s)

(5)

CaCl2 ⋅ H 2O(s ) → CaCl2( s) + H 2O( g )

(6)

CaCl2 ⋅ H 2O( s ) + O2 ( g ) → Ca(ClO) 2 ( s ) + H 2O( g )

(7)

It also seems that the reaction mechanism of the chlorination reaction is more complex than theoretically expected, and it is suggested that CaClOH may ultimately convert to either CaCl2·H2O or CaCl2.

3.1 Behaviors of Chlorination and Sulfation


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Figure 3, which plots individual concentration profiles, provides information regarding the SO2 and HCl uptake characteristics. It is clearly seen that the concentrations of both SO2 and HCl quickly reach a certain value over the first few minutes and then begin to change. As observed in Figure 3a, there is an initial rapid increase in the HCl concentration as the experiment starts, displaying a poor HCl retention efficiency under the conditions used in the present experiments. Similar results can also be noted for SO2 concentration profiles. From this "breakthrough" point, a sigmoidal increase in the HCl concentration as a function of time is shown in Figure 3a, which presents a real breakthrough curve for the reaction between HCl and a sorbent. Apart from the temperature dependence, a much higher space velocity, defined as the volumetric rate of disposal of HCl or SO2 containing gas per unit of bulk sorbent, is thought to be responsible for this result based on empirical evidence. As can be seen in Figure 4, an increase in the amount of sorbent up to 0.5 g (±0.001 g) obviously decreases the emitted levels of HCl during the initial stage of simultaneous absorption. Figure 3 also shows several unexpected results regarding the variation of the SO2 and HCl concentrations with the progress of the reaction, particularly at 750 ˚C and 850 ˚C. For HCl concentration profiles in Figure 3, it is interesting to note that the effluent HCl concentrations reach more than 500 ppm, which is the HCl concentration in the reacting gas mixture. This finding is especially surprising from the perspective of elementary material balance. More 11 ACS Paragon Plus Environment

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specifically, it does not satisfy the condition of chlorine material balance because the chlorine output is greater than the input when the HCl concentration reaches more than 500 ppm. Such a phenomenon, however, may not be impossible considering the overall chlorine material balance during the absorption process. On the other hand, as for the concentration profile of the effluent SO2, it is also surprising to find a decrease in the SO2 concentration in the early stages of the reaction, especially at 750 ˚C and 850 ˚C. However, the extent of the decrease is relatively small at 650 ˚C. The observed decrease in the SO2 concentration indicates that SO2 is being absorbed, presumably through some type of reaction of the sorbent particle other than the sulfation of calcined limestone. Assuming that the reaction products for the chlorination process are CaCl2·H2O or CaClOH, it appears that the chlorides that are initially formed most likely consumed according to the following reaction:

CaCl2 ⋅ H 2O( s ) + SO2( g ) + 1 / 2O2( g ) → CaSO4( s ) + 2HCl( g )

(8)

or

CaClOH ( s ) + SO2 ( g ) + 1 / 2O2 ( g ) → CaSO4 ( s ) + HCl( g )

(9)

Similar to the Hargreaves process (NaCl+SO2+1/2O2=Na2SO4+2HCl), which is widely used for the production of HCl and/or Glauber's salts, HCl can be released from alkali chlorides at appropriate temperatures and in the presence of SO2.36, 37 This HCl-producing reaction is called sulfation of chlorides. Therefore, it is conceivable that reactions (3) or (4) and (8) or (9) take 12 ACS Paragon Plus Environment

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place sequentially. In the present experiments, this reaction process provides the most promising and reasonable explanation for not only the decrease in the SO2 concentration level but also the previously described excessive HCl concentration level. The results in Figure 3 agree with this explanation, i.e., the excessive HCl concentration level is always accompanied by a decrease in the SO2 concentration as expected. In Figure 3a, which initially shows a slight, insignificant decrease in the SO2 concentration, a lower temperature induces lower favorability for the sulfation process through which chlorides are formed, leading to a less sigmoidal shape of the HCl concentration profile than that is typically seen in most gas-solid reactions. Another point to note is that in Figure 3b and Figure 3c, a further decrease in the effluent HCl concentration is observed after reaching the maximum level. Subsequently, the observed HCl concentrations tend to level off to a value lower than the input HCl concentration and remain constant thereafter. A possible explanation for the disappearance of a small amount of gaseous HCl is that it is oxidized to yield chlorine gas (Cl2). At temperatures above 600 ˚C, the oxidation process of HCl can take place vigorously without any metal catalyst, especially at HCl concentrations less than 1000 ppm.7 An elevated temperature is also helpful in promoting the oxidation of HCl to Cl2. This oxidation process, also well known as the Deacon reaction, is presented as follows:

4HCl( g ) + O2( g ) → 2Cl2( g ) + 2H2O( g )

(10) 13

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Similar to HCl, the effluent SO2 concentrations were also found to level out at a value that is less than that at the input level, implying that a fraction of SO2 in all likelihood yields SO3 via either or both of the following two reactions occurred under the conditions being studied.7, 38-40

SO2 ( g ) + 1 / 2O2 ( g ) → SO3 ( g )

(11)

SO2 ( g ) + Cl2 ( g ) + H 2O( g ) → 2 HCl( g ) + SO3 ( g )

(12)

Reaction (11) is a typical oxidation process of SO2, and reaction (12) presents another pathway that yields SO3. An insignificant conversion of SO2 to SO3 was observed by Liu et al.7 under conditions similar to those used in this work. Results from co-combustion of coal and MSW showed the emission of SO2 in flue gases was much more predominant than that of SO3.41 Furthermore, Scala et al.40 also found in their experiment that, when a gas mixture containing 1800 ppm SO2 and 8.5% O2 in balanced N2 was fed into the reactor kept at 850 ˚C, the SO2 concentration at the exhaust was slightly less than that at the input level. They attributed the loss of SO2 to the oxidation of SO2 to SO3, and the molar ratio of SO3 to SO2 was around 0.09. These results all suggested the presence of a minor SO3 concentration. As discussed above, the Deacon reaction provides the source of Cl2 required for reaction (12) to proceed. This would lead to the subsequent release of HCl via this secondary gas-phase reaction. On the other hand, reaction (12) has also been considered to be a possible mechanism for reducing the emission of chlorinated organic compounds. This hypothesis, that SO2 behaves as an inhibitor for PCDD/F formation, 14 ACS Paragon Plus Environment

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was originally proposed by Griffin42 and was later confirmed by other investigators.8-10, 43-45 It is suggested that the depletion of the Cl2-level by sulfur decreases the possible chlorination of aromatic compounds, i.e., through aromatic substitution reactions. As for the SO3 generated from these two reactions, it has been mentioned that the yields of SO3 from SO2 could be promoted in the presence of HCl.7 In addition, SO3 has more potential to yield calcium sulfate (CaSO4) than SO2.26 This may perhaps be one of the reasons for the enhancement of CaO to CaSO4 in the presence of HCl; however, its impact is still limited, mainly because only a small amount of SO2 is able to be converted to SO3.

CaO ( s ) + SO3 ( g ) → CaSO 4 ( s )

(13)

3.2 Sulfation of Chlorides

To confirm the possibility of reaction (8) or (9) occurring under the present conditions, tests on the sulfation of chloride were also performed in the fixed-bed reactor system. The spent calcium sorbent that had completely experienced the chlorination process was used as the source of chloride; it was prepared through an experiment in which CaO reacted with a gas mixture containing 500 ppm HCl, 6% O2 and 15% CO2 in balanced N2 at 650 ˚C. After the preparation of spent calcium sorbent was completed, the gas flow was switched to pure N2 to purge the reactor system. After all of the components in the reacting gas had been completely purged, the


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temperature inside the reactor was changed to the test temperature. As soon as the desired reaction temperature was reached, the reacting gas containing 500 ppm SO2, 6% O2 and 15% CO2 in balanced N2 was immediately introduced into the reactor. Unfortunately, tests at temperatures above 800 ˚C could not be executed due to the extensive reformation of HCl during the heating process. This basically implies that these chlorides (not necessarily CaCl2) are not thermodynamically stable above 800 ˚C, as explained by the reverse reactions of (3) and (4). It is also known that the chlorination reaction at elevated temperatures is not favored due to chemical equilibrium and thermodynamic considerations. Experiments were successfully carried out in the temperature range 650-750 ˚C, and these results are shown in Figure 5. The results demonstrate that the sulfation of spent sorbent containing chlorides occurs at these temperatures, resulting in the release of HCl as well as the simultaneous consumption of gaseous SO2 (not shown in this figure). In addition, it is clearly observed in this figure that the rate of production of HCl increases with the increasing reaction temperature. The release of HCl is more pronounced at 750 ˚C. These results confirm those observed in Figure 3, and further consolidate the validity of the proposed explanation concerning reactions (8) or (9). It is readily understood that simultaneous sulfation and chlorination of calcined limestone yields sulfates and chlorides, respectively; whereas, the sulfation of chlorides formed occurs sequentially, consuming SO2 as well as releasing chlorine in the form of gaseous 16 ACS Paragon Plus Environment

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HCl, thus giving rise to the excessive HCl concentration level and the decrease in the SO2 concentration level, as observed in Figure 3b and 3c. Although tests at temperatures above 800 ˚C could not be executed, the similarity of the breakthrough curve patterns in Figure 3b and 3c confirms indirectly that the sulfation of chlorides also takes place at or above 800 ˚C. Therefore, it is certain that, at temperatures of 700 ˚C or higher, a portion of the chlorides formed due to chlorination can be converted to CaSO4 via reaction (8) or (9) during the simultaneous absorption of SO2 and HCl; this is the reason for the concurrent enhancement of sulfation and suppression of chlorination.

3.3 Uptake Ability

For practical applications, the uptake efficiency of sorbent has always been the issue of greatest concern. In Figure 3, the breakthrough curves not only present the history of the reaction process but also include the concept of gas-phase material balance. Accordingly, the theoretical sulfur/chlorine content captured with solid lime during simultaneous absorption was evaluated on the basis of the breakthrough curves. The total moles of gas component, i, entering the reactor during the reaction process, Ni,total (mol), can be expressed as follows: N i ,total = Fi ,0 × teq

(14)

Fi,0 is the inlet molar flow rate of gas component, i (mol/min); teq is the reaction time required for


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reaching the equilibrium state (min). The inlet and outlet molar flow rate of gas component, i, Fi,0 and Fi,e, can be expressed respectively as follows:

Fi ,0 = Ci ,0 × υ m

(15)

Fi , e = Ci , e × υ m

(16)

where Ci,0 and Ci,e represent the inlet and outlet concentration of the gas component, i, in the gas mixtures (mol/Nm3), respectively; νm stand for the volumetric flow rate of the gas mixtures (Nm3/min). Note that in the experiments the measurement of Ci,e and Ci,e are expressed in terms of ppmv (parts per million by volume). The area under the breakthrough curve represents the unreacted moles of the gas component discharged from the reactor, Ni,lost (mol). Therefore, the Ni,lost can be obtained using the following equation:

N i , lost =

∫

t eq

0

Fi , e dt

(17)

Substitution of Fi,e from Eq. (16) into (17) yields: t =teq

Ni ,lost = υm ∫

t =0

Ci ,e dt

(18)

At any time, t (less than teq), the moles of the gas component reacted with the sorbent, Nrx,t (mol), can be written as: t =t

N rx , t = υ m × (Ci , 0 × t − ∫ Ci , e dt )

(19)

t =0

Figure 6 depicts the captured amount of HCl as a function of time during the simultaneous absorption processes of SO2 and HCl and also presents how the extent of chlorine uptake is 18 ACS Paragon Plus Environment

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affected when reaction (8) or (9) occurs. In this figure, the curve obtained at 650 ˚C shows a typical, asymptotic behavior. Nevertheless, the other curves are observed to rise at first and then decline. Such results are not surprising because some of the generated chlorides react with SO2 to produce HCl in the gas phase, causing a decline in the amount of captured HCl. This figure also reveals the relationship between the HCl uptake efficiency and temperature. For comparison, previous investigations7,

18, 23

have reported HCl absorption efficiency of calcined limestone

could achieve a maximum at 650 ˚C in the absence of SO2 and decrease rapidly as the temperature continues to rise. In this work, the optimal temperature for HCl absorption in the presence of SO2 is still 650 ˚C. Also, as expected from the results, the HCl uptake efficiency drops with rising reaction temperature. Apart from the negative impact of an elevated temperature, the presence of SO2 significantly limits the attainable extent of the chlorination reaction at or above 700 ˚C, whereas at 650 ˚C, the reaction extent is hardly affected. Obviously, the HCl uptake efficiency becomes particularly worse as the sulfation of chlorides starts to vigorously occur, restricting further absorption process. Because the formation of solid CaSO4 may cause pore blockage and closure, it is possible to hinder further chlorination/sulfation reactions and to restrict sorbent utilization. In summary, from the results presented thus far, one can conclude that during the simultaneous absorption of SO2 and HCl at or above 700 ˚C, chlorination will be considerably 19 ACS Paragon Plus Environment

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suppressed because the generated chlorides partly react with SO2 to form sulfates, thereby resulting in an overall sulfation enhancement and the re-release of HCl to the gas phase. Once the complex product layer consisting of a mixture of chlorides and sulfates is formed, the chlorides of which also react to some extent with SO2, thus reducing CaCl2 and forming more CaSO4 in the product shell. Basinas et al.46 widely discussed the role of the formation of CaSO4 layer in the progress of gas-solid reactions. One speculative explanation may be that an impenetrable CaSO4 shell forms on the periphery47, making it difficult to diffuse through the complex product layer into the interior of the sorbent, restricting, either directly or indirectly, further absorption process. In general terms, the process of the sulfation of chlorides plays a key role in the simultaneous removal of SO2 and HCl. The chlorides formed act as a sorbent to react with SO2, indicating an indirect sulfation process. Nevertheless, this is usually accompanied by the formation of gaseous HCl from the sulfation of chlorides. The role of chlorides was also known in the literature to have enhancing effects on the sulfation reaction as additives.25, 48-50 They attributed the enhancing sulfation by chlorides to the formation of eutectics25, 48, 50 or pore enlargement49. Hu et al.51 mentioned that the formation of eutectics kept the pores open and thus reduced the intraparticle diffusion resistance. Figure 7 shows the comparison of the SO2 and HCl uptake efficiencies at various temperatures and at various gas atmospheres. As mentioned above, increasing the reaction 20 ACS Paragon Plus Environment

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Page 22 of 42

temperature from 650 to 700 ˚C considerably suppresses HCl uptake due to the sulfation of chlorides; SO2 uptake is therefore slightly enhanced. This results in a sharp decline in the overall uptake efficiency, corresponding to approximately half of the observed uptake efficiency at 650 ˚C. It is also noted that increasing the temperature further to 850 ˚C has no pronounced impact on the overall uptake efficiency, i.e., it remains relatively stable at temperatures exceeding 700 ˚C. Even so, different uptake efficiencies of SO2 and HCl can still be distinguished from this figure. This suggests that there are two distinct temperature regimes for better SO2 and HCl absorption. For example, the calcined limestone maintains a stronger reactivity for HCl at temperatures of 700 to 750 ˚C, despite the fact that the presence of SO2 limits its uptake efficiency; whereas at temperatures at or above 800 ˚C, the SO2 uptake efficiency is found to be predominant. Although conditions with high temperatures and the presence of SO2 are unfavorable for HCl uptake, calcined limestone still has the ability to capture HCl. This observation supports the experience that alkali chlorides were found in the bottom ash when limestone was directly added to the commercial fluidized bed combustor.52 In addition, the influence of the gas atmosphere on the uptake efficiency is also compared in Figure 7. The overall uptake efficiency at each temperature is found to be slightly enhanced under conditions simulating poor combustion. Further examination of these results reveals that the change in gas atmosphere only has a small impact on the HCl uptake efficiency23, but enhances the SO2 uptake efficiency. The reason for this 21 ACS Paragon Plus Environment

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Energy & Fuels

enhanced sulfation could very well be that increasing the O2 concentration favors reaction (1), forming more sulfates according to the reaction equilibrium.

4. CONCLUSION

The results of the experiments on the simultaneous absorption of SO2 and HCl by calcined limestone under conditions simulating combustion were fairly complex because in addition to the expected heterogeneous reactions, homogeneous reactions were also found to occur in the gas phase. At the studied temperatures, the presence of O2 either directly or indirectly led to the simultaneous oxidation of a small amount of SO2 and HCl, producing SO3 and Cl2, respectively. Although SO3 had more potential than SO2 to be absorbed, it was considered to have an insignificant effect on the enhancement of sulfation, mainly because only a small amount SO3 was formed. The results also suggested that the chloride formed is not CaCl2 but is rather CaCl2·H2O or CaClOH. At a temperature of 650 ˚C, chlorination is much more predominant than sulfation. Nevertheless, at or above 700 ˚C, chlorination was found to be considerably suppressed, resulting in the enhancement of sulfation. This phenomenon was confirmed to be due to the sulfation of the chlorides that were initially formed. It was observed that the sulfation of chlorides also restricts further absorption process and re-release of HCl into the gas phase. The


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results revealed that at or above 700 ˚C, the overall uptake efficiency remained nearly unchanged despite the difference in reactivity of SO2 and HCl. Changes in the gas atmosphere (either CO2 or O2 concentrations) had a small impact on HCl uptake; nevertheless, increasing the O2 concentration slightly enhanced SO2 uptake at the studied temperatures.


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Energy & Fuels

NOMENCLATURE Ci,0

[mol/Nm3]

the inlet concentration of gas component i in the gas mixtures

Ci,e

[mol/Nm3]

the outlet concentration of gas component i in the gas mixtures

Fi,0

[mol/min]

the inlet molar flow rate of gas components i

Fi,e

[mol/min]

the outlet molar flow rate of gas components i

Ni, total

[mol]

the total moles of gas components i entering the reactor

Ni, lost

[mol]

the unreacted moles of gas component i discharging from the reactor

Nrx,t

[mol]

the moles of gas component i that have reacted

νm

[Nm3/min]

the volumetric flow rate of gas mixtures

teq

[s]

the reaction time required for reaching the equilibrium state


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Wang, Z.; Huang, H.; Li, H.; Wu, C.; Chen, Y.; Li, B. HCl Formation from RDF Pyrolysis

and Combustion in a Spouting-Moving Bed Reactor. Energy Fuels 2002, 16, 608-614. 2.

Zhu, H. M.; Jiang, X. G.; Yan, J. H.; Chi, Y.; Cen, K. F. TG-FTIR analysis of PVC thermal

degradation and HCl removal. J. Anal. Appl. Pyrolysis 2008, 82, 1-9. 3.

Wey, M.-Y.; Chen, J.-C.; Wu, H.-Y.; Yu, W.-J.; Tsai, T.-H. Formations and controls of HCl

and PAHs by different additives during waste incineration. Fuel 2006, 85, 755-763. 4.

Wey, M. Y.; Liu, K. Y.; Yu, W. J.; Lin, C. L.; Chang, F. Y. Influences of chlorine content on

emission of HCl and organic compounds in waste incineration using fluidized beds. Waste Manage. 2008, 28, 406-415. 5.

Wei, X.; Wang, Y.; Liu, D.; Sheng, H.; Tian, W.; Xiao, Y. Release of Sulfur and Chlorine

during Cofiring RDF and Coal in an Internally Circulating Fluidized Bed. Energy Fuels 2009, 23, 1390-1397. 6.

Takasuga, T.; Makino, T.; Tsubota, K.; Takeda, N. Formation of dioxins (PCDDs/PCDFs)

by dioxin-free fly ash as a catalyst and relation with several chlorine-sources. Chemosphere 2000, 40, 1003-1007. 7.

Liu, K.; Pan, W. P.; Riley, J. T. A study of chlorine behavior in a simulated fluidized bed

combustion system. Fuel 2000, 79, 1115-1124. 25 ACS Paragon Plus Environment

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8.

Thomas, V. M.; McCreight, C. M. Relation of chlorine, copper and sulphur to dioxin

emission factors. J. Hazard. Mater. 2008, 151, 164-170. 9.

Anthony, E. J.; Jia, L.; Granatstein, D. L. Dioxin and Furan Formation in FBC Boilers.

Environ. Sci. Technol. 2001, 35, 3002-3007. 10. Raghunathan, K.; Gullett, B. K. Role of Sulfur in Reducing PCDD and PCDF Formation. Environ. Sci. Technol. 1996, 30, 1827-1834. 11. Tagashira, K.; Torii, I.; Myouyou, K.; Takeda, K.; Mizuko, T.; Tokushita, Y. Combustion characteristics and dioxin behavior of waste fired CFB. Chem. Eng. Sci. 1999, 54, 5599-5607. 12. Takeshita, R.; Akimoto, Y. Control of PCDD and PCDF formation in fluidized bed incinerators. Chemosphere 1989, 19, 345-352. 13. Li, Y.; Wang, H.; Jiang, L.; Zhang, W.; Li, R.; Chi, Y. HCl and PCDD/Fs emission characteristics from incineration of source-classified combustible solid waste in fluidized bed. RSC Adv. 2015, 5, 67866-67873. 14. Partanen, J.; Backman, P.; Backman, R.; Hupa, M. Absorption of HCl by limestone in hot flue gases. Part I: the effects of temperature, gas atmosphere and absorbent quality. Fuel 2005, 84, 1664-1673.


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15. Manovic, V.; Anthony, E. J. Competition of Sulphation and Carbonation Reactions during Looping Cycles for CO2 Capture by CaO-Based Sorbents. J. Phys. Chem. A 2010, 114, 3997-4002. 16. Lawrence, A. D.; Bu, J. The reactions between Ca-based solids and gases representative of those found in a fluidized-bed incinerator. Chem. Eng. Sci. 2000, 55, 6129-6137. 17. Xie, X.; Li, Y.; Wang, W.; Shi, L. HCl removal using cycled carbide slag from calcium looping cycles. Appl. Energy 2014, 135, 391-401. 18. Chyang, C.-S.; Han, Y.-L.; Zhong, Z.-C. Study of HCl Absorption by CaO at High Temperature. Energy Fuels 2009, 23, 3948-3953. 19. Sun, Z.; Yu, F.-C.; Li, F.; Li, S.; Fan, L.-S. Experimental Study of HCl Capture Using CaO Sorbents: Activation, Deactivation, Reactivation, and Ionic Transfer Mechanism. Ind. Eng. Chem. Res. 2011, 50, 6034-6043. 20. Daoudi, M.; Walters, J. K. The reaction of HCl gas with calcined commercial limestone particles: The effect of particle size. The Chemical Engineering Journal 1991, 47, 11-16. 21. Wang, C.; Zhang, Y.; Jia, L.; Tan, Y. Effect of water vapor on the pore structure and sulfation of CaO. Fuel 2014, 130, 60-65. 22. Bie, R.; Li, S.; Yang, L. Reaction mechanism of CaO with HCl in incineration of wastewater in fluidized bed. Chem. Eng. Sci. 2005, 60, 609-616. 27 ACS Paragon Plus Environment

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23. Lin, G.-M.; Chyang, C.-S. Effect of CO2 on high temperature chlorination behavior of calcined limestone in an innovated fixed-bed reactor. Journal of the Taiwan Institute of Chemical Engineers 2016, 62, 60-67. 24. Partanen, J.; Backman, P.; Backman, R.; Hupa, M. Absorption of HCl by limestone in hot flue gases. Part III: simultaneous absorption with SO2. Fuel 2005, 84, 1685-1694. 25. Matsukata, M.; Takeda, K.; Miyatani, T.; Ueyama, K. Simultaneous chlorination and sulphation of calcined limestone. Chem. Eng. Sci. 1996, 51, 2529-2534. 26. Xie, W.; Liu, K.; Pan, W. P.; Riley, J. T. Interaction between emissions of SO2 and HCl in fluidized bed combustors. Fuel 1999, 78, 1425-1436. 27. Kwak, Y.-H.; Lee, Y.-M.; Bae, W.-K.; Kim, W.-H.; Bae, S.-K. Absorption behavior of chlorine and sulfur by Ca(OH)2 under simulated conditions with a fluidized bed combustor. J. Mater. Cycles Waste Manage. 2011, 13, 314-320. 28. Partanen, J.; Backman, P.; Backman, R.; Hupa, M. Absorption of HCl by limestone in hot flue gases. Part II: importance of calcium hydroxychloride. Fuel 2005, 84, 1674-1684. 29. Xie, X.; Li, Y.-J.; Liu, C.-T.; Wang, W.-J. HCl absorption by CaO/Ca3Al2O6 sorbent from CO2 capture cycles using calcium looping. Fuel Process. Technol. 2015, 138, 500-508. 30. Gullett, B. K.; Jozewicz, W.; Stefanski, L. A. Reaction kinetics of calcium-based sorbents with hydrogen chloride. Ind. Eng. Chem. Res. 1992, 31, 2437-2446. 28 ACS Paragon Plus Environment

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31. Jozewicz, W.; Gullett, B. K. Reaction Mechanisms of Dry Ca-Based Sorbents with Gaseous HCl. Ind. Eng. Chem. Res. 1995, 34, 607-612. 32. García, J.; López, T.; Álvarez, M.; Aguilar, D. H.; Quintana, P. Spectroscopic, structural and textural properties of CaO and CaO–SiO2 materials synthesized by sol–gel with different acid catalysts. J. Non-Cryst. Solids 2008, 354, 729-732. 33. Chin, T.; Yan, R.; Liang, D. T. Study of the Reaction of Lime with HCl under Simulated Flue Gas Conditions Using X-ray Diffraction Characterization and Thermodynamic Prediction. Ind. Eng. Chem. Res. 2005, 44, 8730-8738. 34. Chin, T.; Yan, R.; Liang, D. T.; Tay, J. H. Hydrated Lime Reaction with HCl under Simulated Flue Gas Conditions. Ind. Eng. Chem. Res. 2005, 44, 3742-3748. 35. Bausach, M.; Krammer, G.; Cunill, F. Reaction of Ca(OH)2 with HCl in the presence of water vapour at low temperatures. Thermochim. Acta 2004, 421, 217-223. 36. Boonsongsup, L.; Iisa, K.; Frederick, W. J. Kinetics of the Sulfation of NaCl at Combustion Conditions. Ind. Eng. Chem. Res. 1997, 36, 4212-4216. 37. Uchida, S.; Kamo, H.; Kubota, H.; Kanaya, K. Reaction kinetics of formation of hydrochloric acid in municipal refuse incinerators. Industrial & Engineering Chemistry Process Design and Development 1983, 22, 144-149.


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38. Iisa, K.; Lu, Y.; Salmenoja, K. Sulfation of Potassium Chloride at Combustion Conditions. Energy Fuels 1999, 13, 1184-1190. 39. Dennis, J. S.; Hayhurst, A. N. The formation of SO3 in a fluidized bed. Combust. Flame 1988, 72, 241-258. 40. Scala, F.; Montagnaro, F.; Salatino, P. Enhancement of Sulfur Uptake by Hydration of Spent Limestone for Fluidized-Bed Combustion Application. Ind. Eng. Chem. Res. 2001, 40, 2495-2501. 41. Lindbauer, R. L.; Wurst, F.; Prey, T. Combustion dioxin supression in municipal solid waste incineration with sulphur additives. Chemosphere 1992, 25, 1409-1414. 42. Griffin, R. D. A new theory of dioxin formation in municipal solid waste combustion. Chemosphere 1986, 15, 1987-1990. 43. Wu, H.-L.; Lu, S.-Y.; Li, X.-D.; Jiang, X.-G.; Yan, J.-H.; Zhou, M.-S.; Wang, H. Inhibition of PCDD/F by adding sulphur compounds to the feed of a hazardous waste incinerator. Chemosphere 2012, 86, 361-367. 44. Gullett, B. K.; Dunn, J. E.; Raghunathan, K. Effect of Cofiring Coal on Formation of Polychlorinated Dibenzo-p-Dioxins and Dibenzofurans during Waste Combustion. Environ. Sci. Technol. 2000, 34, 282-290.


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45. Ogawa, H.; Orita, N.; Horaguchi, M.; Suzuki, T.; Okada, M.; Yasuda, S. Dioxin reduction by sulfur component addition. Chemosphere 1996, 32, 151-157. 46. Basinas, P.; Wu, Y.; Grammelis, P.; Anthony, E. J.; Grace, J. R.; Jim Lim, C. Effect of pressure and gas concentration on CO2 and SO2 capture performance of limestones. Fuel 2014, 122, 236-246. 47. Anthony, E. J.; Granatstein, D. L. Sulfation phenomena in fluidized bed combustion systems. Prog. Energy. Combust. Sci. 2001, 27, 215-236. 48. Zhao, Y.; Lin, W.-C. Multi-functional sorbents for the simultaneous removal of sulfur and lead compounds from hot flue gases. J. Hazard. Mater. 2003, 103, 43-63. 49. Davini, P.; DeMichele, G.; Ghetti, P. An investigation of the influence of sodium chloride on the desulphurization properties of limestone. Fuel 1992, 71, 831-834. 50. Hu, G.; Dam-Johansen, K.; Wedel, S.; Hansen, J. P. Enhancement of the Direct Sulfation of Limestone by Alkali Metal Salts, Calcium Chloride, and Hydrogen Chloride. Ind. Eng. Chem. Res. 2007, 46, 5295-5303. 51. Hu, G.; Dam-Johansen, K.; Wedel, S.; Peter Hansen, J. Review of the direct sulfation reaction of limestone. Prog. Energy. Combust. Sci. 2006, 32, 386-407.


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52. Wan, H.-P.; Chang, Y.-H.; Chien, W.-C.; Lee, H.-T.; Huang, C. C. Emissions during co-firing of RDF-5 with bituminous coal, paper sludge and waste tires in a commercial circulating fluidized bed co-generation boiler. Fuel 2008, 87, 761-767.


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Figure captions

Schematic of the bench-scale fixed-bed reactor system: (1) mass flow controller; (2) Figure 1

air purging pump; (3) cooler; (4) drying column; (5) blending chamber; (6) electrical heater; (7) fixed-bed reactor; (8) FTIR spectrometer; (9) NaOH solution.

Figure 2

Schematic diagram of the specially designed fixed-bed reactor. SO2 and HCl emissions during the simultaneous sulfation and chlorination of calcined

Figure 3 limestone at (a) 650 ˚C; (b) 750 ˚C; (c) 850 ˚C. Effect of sorbent mass on HCl emissions during the simultaneous sulfation and Figure 4 chlorination of calcined limestone at 650 ˚C. Re-release of HCl from the sulfation process of calcined limestone containing Figure 5

chlorides. chlorination process: 500 ppm HCl, 6% O2 and 15% CO2 in balanced N2; sulfation process: 500 ppm SO2, 6% O2 and 15% CO2 in balanced N2 Effect of temperature on the chlorination behavior of calcined limestone in the

Figure 6 presence of SO2. Comparison of the SO2 and HCl uptake efficiency of calcined limestone under various Figure 7 conditions; gas composition: 500 ppm SO2, 500 ppm HCl, α % O2, β % CO2 in N2.


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Energy & Fuels

Figure 1 Schematic of the bench-scale fixed-bed reactor system: (1) mass flow controller; (2) air purging pump; (3) cooler; (4) drying column; (5) blending chamber; (6) electrical heater; (7) fixed-bed reactor; (8) FTIR spectrometer; (9) NaOH solution.


Energy & Fuels

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Figure 2 Schematic diagram of the specially designed fixed-bed reactor.


Page 36 of 42

Page 37 of 42

600

emitted concentration (ppm)

500 492 ppm 474 ppm 462 ppm

400

300

200

100

HCl SO2

0 0

25

50

75

100

125

150

175

200

time (min)

(a) 600 526 ppm

485 ppm

500 emitted concentration (ppm)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Energy & Fuels

481 ppm 400 404 ppm

300

350 ppm

200

100

HCl SO2

0 0

25

50

75

100

125

150

time (min)

(b)


175

200

Energy & Fuels

600 515 ppm

478 ppm

500 emitted concentration (ppm)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

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400

471 ppm 314 ppm

300

200 233 ppm 100

HCl SO2

0 0

25

50

75

100

125

150

175

200

time (min)

(c) Figure 3 SO2 and HCl emissions during the simultaneous sulfation and chlorination of calcined limestone at (a) 650 ˚C; (b) 750 ˚C; (c) 850 ˚C.


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600

emitted concentration (ppm)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Energy & Fuels

HCl (0.5 g calcined limestone) HCl (0.1 g calcined limestone)

500 400 300 200 100 0

0

25

50

75

100

125

150

175

200

time (min)

Figure 4 Effect of sorbent mass on HCl emissions during the simultaneous sulfation and chlorination of calcined limestone at 650 ˚C.


Energy & Fuels

700

emitted HCl concentration (ppm)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

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665 ppm

o

650 C o 700 C o 750 C

600 500 400 300 210 ppm 200

88 ppm

100 0

0

50

100

150

200

250

300

350

400

time (min)

Figure 5 Re-release of HCl from the sulfation process of calcined limestone containing chlorides. chlorination process: 500 ppm HCl, 6% O2 and 15% CO2 in balanced N2; sulfation process: 500 ppm SO2, 6% O2 and 15% CO2 in balanced N2


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0.0015 accumulated captured amount (mole)

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Energy & Fuels

650 oC 700 oC 750 oC 800 oC 850 oC

0.0010

0.0005

0.0000

0

25

50

75

100

125

150

175

200

Time (min)

Figure 6 Effect of temperature on the chlorination behavior of calcined limestone in the presence of SO2.


Energy & Fuels

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Figure 7 Comparison of the SO2 and HCl uptake efficiency of calcined limestone under various conditions; gas composition: 500 ppm SO2, 500 ppm HCl, α % O2, β % CO2 in N2.


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SO2 Capture by Calcined Limestone from Hot

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