Ind. Eng. Chem. Res. 1995,34, 607-612
607
Reaction Mechanisms of Dry Ca-Based Sorbents with Gaseous HCI Wojciech Jozewicz Acurex Environmental Corporation, 4915 Prospectus Drive, P.O. Box 13109, Research Triangle Park, North Carolina 27709
Brian K. Gullett" Air and Energy Engineering Research Laboratory, U.S. Environmental Protection Agency, Research Triangle Park, North Carolina 27711
The mechanisms of HC1 reaction with dry Ca(OH)2 or CaO sorbents in flue gas cleaning applications were investigated by differential scanning calorimetry (DSC), thermogravimetry, and X-ray diffraction. A short-time differential reactor (STDR) was used to contact 1000 ppm HC1 with dry sorbents at temperatures ranging from 100 to 600 "C. The product of HC1 reaction with Ca(OH)2 over the 100-600 "C range in the STDR was identified as a solid solution of CaC12-2H20 and CaClOH. Subsequently, the CVCa molar ratio in the product was determined to be less than 2. The product of the HC1 reaction with CaO in the STDR over the 100-600 "C range was identified as CaClOH. When scanned by DSC, samples of both CaO and Ca(OH)2 that were previously exposed to HC1 in the STDR revealed a n endotherm between 530 and 590 "C. The amount of enthalpy correlated well with the chemically determined sample conversion.
Introduction Hydrogen chloride is one of the acid gas emissions resulting from municipal and hazardous waste combustion and can be neutralized by reaction with calciumbased sorbents downstream of the furnace. The neutralization scheme frequently used is an injection of a dry Ca-based sorbent, most often calcium hydroxide or calcium carbonate. Depending on the injection temperature, the species reacting with HC1 may be Ca(OH)2 or a product of its decomposition-calcium o i d e (CaO)-if injection takes place at a temperature higher than 400 "C. Similarly, when calcium carbonate is injected, the reacting species may be CaC03 or CaO if injection takes place a t a temperature higher than 650 "C. Various types of dry Ca-based systems for HC1 removal are presented in the literature [e.g., Ettehadieh and Lee (1989), Hsieh et al. (1989), and Felsvang and Helvind (199111. The literature on dry Ca-based systems for HC1 removal discusses the overall effectiveness of HC1 removal across respective systems. There is also a body of reports describing fundamental, bench-scale work addressing the HC1 reaction with dry Ca-based sorbents (Karlsson et al. 1981; Gullett et al., 1992; Pakrasi et al., 1992; Weinell et al., 1992). More fundamental studies elucidate the kinetics and mechanism of dry Ca-based sorbent in the removal of HC1, assuming the overall reaction of Ca(OH),
+ 2HC1
--L
CaC12*2H20
(1)
or CaO
+ 2HC1-
CaCl,
+ H20
(2)
A knowledge of reaction products is necessary to accurately determine sorbent utilization following contact with HC1. Assuming the overall reaction proceeds according to reaction 1or 2 then the result is a product molar ratio of chlorine t o calcium (CVCa) equal to 2 in a completely utilized sorbent. Consistent with this assumption, any product CYCa molar ratio of less than
* Author t o whom correspondence should be addressed. 0888-588519512634-0607$09.00/0
2 would represent incomplete sorbent utilization or the formation of a reaction product that had an atomic ratio of CVCa less than 2. If the latter is the case, then complete sorbent utilization could be achieved even at product CVCa molar ratios of less than 2, and reaction 1 or 2 would not describe the reaction mechanism. Therefore, a knowledge of the reaction product of HC1 with dry Ca-based sorbents is necessary for an accurate determination of the sorbent's utilization. Sorbent utilization is often an important factor in determining the operating costs of a flue gas cleaning process. On the basis of the results of scanning electron microscopy (SEM),X-ray dispersive spectroscopy (XDS), and X-ray diffraction (XRD),Gullett et al. (1992) postulated the possibility of intermediate reactions resulting in the formation of calcium chloride hydroxide and calcium hypochlorite in addition to calcium chloride dihydrate. Weinell et al. (1992) hypothesized the existence of a less stable precursor t o a more stable final product during the reaction of HC1 with Ca(OH)2. Except for the reports of Gullett et al. (1992) and Weinell et al. (1992),no other information was available on the mechanism of reaction 1or 2 in acid gas control equipment. In work investigating the formation of bleaching powder, Bunn et al. (1934) concluded that the reaction of chlorine and Ca(OH)2 never goes to completion [usually 15-25% of Ca(OH)2 remains unreactedl and that reaction products might contain a mixture of double salts-Ca*OH*OCl, CaCl-OCl, CwOC1-OC1, and CwCb C1-so that the CVCa atomic ratio of these product salts may vary from 1/1 t o 2/1. Rabovskii et al. (1972) investigated the structure of the reaction products of Cl2 and dry Ca(OH)2and concluded that approximately 50% of the chlorination product was not found in the free state; rather, it was bound to Ca(OC1)z as calcium hypochlorite bishydroxide salt [Ca(OC1)2*2Ca(OH)21. Markova (1973a,b) crystallized two calcium chloride hydroxide salts [CaC12-3Ca(OH)z-l2H20and CaCkCa(0H)zl out of CaC12-Ca(OH)2-H20 systems. Kosnyrev et al. (1990)investigated changes in the Ca(OH)2-CaCl2*2H20systems during thermal dehydration
0 1995 American Chemical Society
608 Ind. Eng. Chem. Res., Vol. 34, No. 2, 1995
and concluded that CaOHCl was a major phase incorporated into solid solutions. The amount of CaOHCl increased up to 500 "C over the 150-640 "C experimental range. As the aforementioned literature review demonstrates, a number of different salts may be produced in the Ca(OH)2-C12 (or HC1) system, depending on the reaction temperature and extent, as well as on the presence of water (moisture) in the system. This work presents results of research conducted to obtain information on the product(s) of HC1 reaction with Ca-based sorbents. The effects of sorbent type and reaction temperature on the nature of the products were investigated.
Experimental Section Samples of Ca-based sorbents were reacted with HC1 in the short-time differential reactor (STDR), an isothermal, fured-bed reactor described elsewhere (Gullett et al., 1990). A sample of Ca-based sorbent (100 mg) was evenly dispersed on quartz wool, placed in a pneumatically activated sample holder, and then rapidly moved under the flow of the process gas. The process gas was 1000 ppm HC1 with 5% oxygen in nitrogen at a flow rate of 23 L(STP)/min. The sample was reacted with HC1 for 60 s at temperatures ranging from 100 to 600 "C. This temperature range was intended to encompass the flue gas temperature range encountered during dry sorbent injection. The STDR was only used as a tool to produce chlorinated samples for product analysis during this work. No efforts were taken to minimize the effects of external flow diffusion on reactivity or to attain differentiality with respect to HC1. The Ca-based sorbents used throughout this work were Linwood Ca(OH)2 (from Linwood Lime and Cement Co. in Davenport, IA) and CaO derived from Ca(OH12 by heat treatment. The sorbent contained 89.1% Ca(OH)2by thermogravimetric analysis (TGA)and had a specific surface area of 14.5 m2/g, a porosity of 0.131, and a mass median diameter of 1.96 pm. CaO was prepared in a muffle furnace by exposing a thin layer of Ca(OH)2 to N2 flow at 870 "C for 30 min. The resulting CaO samples had a specific surface area of approximately 5 m2/g,a pore volume of 0.012 cm3/g,and a mass median diameter of 4 pm. Samples of CaO were produced daily to minimize their handling time. Chlorination products of Ca(OH)2 or CaO reaction with HC1 in the STDR were evaluated in a Perkin Elmer DSC7 differential scanning calorimeter (DSC)operated at a 20 "C/min scanning rate over the temperature range from ambient to 650 "C. For each DSC run, the weight of a sample (typically 12-20 mg) was entered into the DSC's memory. This allowed for the endothermic heat flow measured by the instrument to be expressed as enthalpy (joules per gram). Unless otherwise noted, aluminum pans were used for the DSC runs. Thermal decomposition of Linwood Ca(OH)2"as received" was measured first t o establish the baseline for the DSC and for comparison with available literature data for Ca(OH12 decomposition. Enthalpy of Ca(OH12 was determined for the 385-535 "C temperature range characteristic for thermal decomposition of Ca(OH12. The average of three measurements for Linwood Ca(OH)2was 1.168 kJ/g. This value was next normalized to account for impurities present in Ca(OH)2. A TGA measurement was used to determine Linwood Ca(OH12 purity. Based on the TGA-measured weight loss between 400 and 650 "C, which is associated with the loss
(decomposition)of chemically bound water, the purity of Linwood Ca(0H)z was determined to be 89.1% (average of four measurements). Normalization of the previously measured enthalpy of 1.168 kJ/g for "as received" Linwood Ca(OH)2 by the TGA-measured purity of Ca(OH12 yielded an enthalpy for the decomposition of "pure" Linwood Ca(OH)z equal to 1.311 kJ/g. Three available literature data sources give the enthalpy of decomposition of Ca(OH)2 as 1.477 (Hartman and Martinovsky, 19921, 1.341 (Dean, 1985), and 1.516 kJ/g (Criado and Morales, 1976). The DSC scan also revealed endotherm between 305 and 385 "C (48.7 J/g enthalpy) that was not accompanied by a corresponding weight loss. This may indicate the presence of an intermediate phase during the decomposition of Ca(OH)2, when CaO and HzO are at least partially dissociated, but H20 has not yet evaporated. Such a phase could be expressed as CaOaH20. The amount of unreacted Ca(OH)2 was determined in a Perkin Elmer TGA7 thermogravimetric analyzer operated at a 20 "C/min scanning rate over the temperature range from ambient to 1000 "C (platinum pans were used) and with 0.083 Umin N2 sweep. The N2 stream was passed through a desiccant bed before entering the TGA. A typical sample size was 12-20 mg and 10-20 mg for DSC and TGA analyses, respectively. The scanning rate of 20 "C/min for DSC and TGA runs was selected based on the review of the available literature data. Scanning rates reported in the literature varied widely. For example, l "C/min was used in a quantitative study of calcium sulfate (Schlichenmaier, 1975) while a scanning rate of 10 "C/min was selected to investigate the decomposition of calcite (Earnest, 1988). Rates as high as 60-200 "C/min were used to analyze flue gas scrubber materials (Dorsey and Buecker, 1988). It was felt that a 20 "C/min scanning rate was sufficiently slow to minimize temperature inhomogeneities within the sample. The sample size (typically 12-20 mg) for DSC and TGA scans was intended to be of sufficient amount to be representative of a reaction product resulting from the exposure of Ca(OH)2or CaO to HC1. Particle size distributions were determined in a Micromeritics Sedigraph using a gravity sedimentation method (Sedisperse A-11 dispersant). Porosity and surface area were measured in a Micromeritics ASAP 2600 using N2 adsorptioddesorption with a BrunauerEmmett-Teller (BET) method. Several samples were analyzed for CVCa molar ratios by ion chromatography (IC) and atomic absorption (AA), respectively. The procedure has been described in greater detail elsewhere (Gullett et al., 1992). XRD analyses were conducted in a Siemens diffractometer. Spectra were identified by computer comparison against the Joint Committee for Powder Diffraction Spectra (JCPDS) files (set 42).
Results and Discussion Thermal Decomposition of Ca(0H)Z. While DSC and TGA analyses of the direct reaction of CaO with HCl(2) are straightforward, analyses of the reaction of Ca(OH)2with HCl(1) must account for the simultaneous weight loss and decomposition energetics of the dehydration of Ca(0H)z. For this reason, direct decomposition of Ca(OH)2 in a N2 atmosphere was studied first in order to quantitatively separate these effects from the overall reaction with HC1.
Ind. Eng. Chem. Res., Vol. 34, No. 2, 1995 609
305-385'C enthalpy 385.535% enthalpy ( ) peaklempe,ralure,'C
A
(5021
(Normallzed)
Ca(OH),,~Oa'C,1 min, 1000 ppm HCI
I mln, no HCI
Penk temp 347.99627 54'C
Omel 406 08% Peak temp: 527.54-635.69-C Onset: 545.86% Heat flow: 265.00 Jlg
c 4 2 -
2
02 0
100
0 200
. 300
0
400
500
0
100
0
600
200
300
400
600
500
Figure 1. DSC results for samples of Linwood Ca(OH)2 exposed to hot Nz (STDR, 1 min).
Figure 3. DSC scan of a Linwood Ca(0H)z sample exposed to 1000 ppm HC1 for 1 min in STDR operated at 400 "C. ~
0.8
2-
$ 600 L,
f
0.6'
a
$ ?
0.5 0.4
A
0.3
0.1
"
1 0
A
100
200
300
400
300
2 U
A r
500
L
A
0
0.2
0
+
+
p 700
Lo
,g
m
CaC12*2H20 CaCl
'
0.7
200
L o 400
500
700
Temperature, "C
Temperature of Thermal Treatment, "C
600
700
Reaction Temperature, "C
Figure 2. Measured conversion of dry Ca-based sorbents (STDR, 1000 ppm HCl, 1 min, 100 mg sample size).
Thermal treatment consisted of the exposure of each sample of Linwood Ca(OH)2for 1min to hot Nz (in the STDR) at several temperatures, ranging from ambient to 600 "C. Immediately following the thermal treatment in the STDR, the sample was analyzed in the DSC and in the TGA. The DSC results are presented in Figure 1 giving enthalpies measured for thermally treated samples. The enthalpy between 385-535 "C and associated with the decomposing Ca(0H)z was generally decreasing with an increasing temperature of the preceding STDR thermal treatment. Enthalpies shown in Figure 1are expressed per gram of the sample and not per gram of Ca(OH)2. Therefore, decreasing of enthalpy with increasing temperature of treatment in the STDR is consistent with the increase in fraction of Ca(0H)z decomposing to CaO in the STDR with an increasing temperature of thermal treatment. The temperature of the peak (associated with the 385 t o 535 "C enthalpy) was decreasing with the increasing temperature of thermal treatment in the STDR (numbers in parentheses in Figure 1). The enthalphy associated with the intermediate phase during Ca(OH)2decomposition occurred between 305 and 385 "C for samples treated thermally at 100 and 200 "C. It was smaller for the sample exposed t o 200 "C than for the one exposed to 100 "C. Enthalpy a t the above temperature range was not measured for the as received sample of Ca(0H)z. No enthalpy was measured in the 305-385 "C temperature range for samples exposed to 300 "C or higher in the STDR. Lack of this phase change (lack of enthalpy from 305 to 385 "C) in samples that have been exposed to temperatures of 300 "C or higher is indicative of the completed transition
*
+ A + &
-
-
* H
20OJlg
A
CalOH), theat Cs(OH), +HCl
- + -
A
+@-A&
cm2 CaC12.2HzO A
@
* * - A -
e
+ 4
I
1
6
610 Ind. Eng. Chem. Res., Vol. 34, No. 2, 1995 Table 1. DSC Comparison of Intrinsic Ca(0H)z Decomposition with Na only vs loo0 ppm HCF Nz only 100 ppm HC1 STDR enthalpy peak enthalpy peak temp ("C) (kJ/g) location ("C) (kJ/g) location ("C) 100 1.083 502 1.031 497 200 1.011 504 0.854 493 300 0.695 494 0.615 479 400 0.408 484 0.189 463 500 0.522 447 0.375 437 600 0.238 422 0.212 420 Sample size, 12-20 mg; 20 "C/min. Samples of Ca(0H)Z pretreated in the STDR prior to the DSC scan.
420-500 "C range and at 350 "C (solid circles in Figure 4) were assumed to be representative of Ca(OH12 decomposition. However, it should be pointed out that these 420-500 "C enthalpies, representative of decomposing Ca(OH)2, were significantly lower for samples exposed to HC1 than for those exposed to N2 only during the thermal treatment. This comparison is shown in Table 1 and may be indicative of the enhanced decomposition of Ca(OH)2 in the presence of HC1. Similar trends occurred (except for the 500 and 600 "C samples) for the TGA scans. In general, the TGA-measured weight loss from the decomposing Ca(OH)2 was significantly lower for samples exposed to HC1 than for those exposed to N2 only. Because of the different TGA responses for samples exposed to HC1 a t STDR temperatures of 1500 "C, it was decided that the usual operational range of the DSC (ambient to 650 "C) should be expanded to 730 "C by using platinum pans. Additional endotherms were then seen on the DSC scans with peaks at 668 and 697 "C for the 500 and 600 "C STDR samples, respectively. No peaks above 600 "C were seen on the DSC scan for samples exposed to HC1 below 500 "C. Figure 4 shows that three endothermic reactions occur on exposure of Ca(OH)2to HC1 at STDR temperatures of 1500 "C. The number, magnitude, and temperature peak location of the endotherms are a function of the STDR temperature. "he occurrence of the additional endotherm on the DSC scan of the samples exposed to HC1 at 500 and 600 "C and simultaneous weight loss upon heating on the TGA scan are indicative of different HC1+ Ca(OH)2 reaction products at these temperatures than at temperatures within the 100-400 "C range. In an effort to identify the reaction products, known materials that were candidate products were scanned on the DSC and TGA. Reagent-grade calcium chloride dihydrate and anhydrous calcium chloride both from Fisher (C-70 and C-77, respectively), were scanned and the results are shown in the top portion of Figure 4 (solid squares). Enthalpies of 57 and 719 J/g were seen from 44 to 76 "C and from 124 to 260 "C on the DSC scan of CaC12.2H20 with peaks at 58 and 205 "C, respectively. The peak at 58 "C could not be detected following the overnight treatment of the sample in a 120 "C oven and is likely due to surface water. The DSC scan of CaC12 revealed enthalpies of 32 and 60 J/g from 134 to 165 "C and from 221 to 276 "C with peaks at 151 and 247 "C, respectively (solid diamonds in Figure 4). As the comparison of data in Figure 4 indicates, the temperature ranges for DSC heat flows and TGA weight loses from the known materials occurred a t temperatures that were different from those observed for reaction samples. Based on these results, it was concluded that pure CaCl2 and CaCk2H20 were not products of the HC1 + Ca(OH12 reaction in the STDR.
The possibility that Ca(OH)2and CaCk2H20 formed a solid solution was investigated by physically mixing the reagent-grade chemicals. Two mixtures were prepared with 9.8 and 20.2 wt % CaCl2.2H20 in Ca(OH)2 to simulate 0.104 and 0.226 CYCa atomic ratios, respectively. The 200 "C peak on the DSC scan previously measured for reagent-grade CaC12.2H20 was not observed. Instead, peaks at approximately 590 "C were observed in both mixtures that match with samples of Ca(OH12 exposed to HC1 (solid triangles in Figure 4). Therefore, it was concluded that the product found in the physical mixture of Ca(OH12 and CaClg2H20 was the same product as previously seen in Ca(OH)2samples exposed to HC1 in the STDR. Based on the known amount of reagent grade CaClp2HzO in the mixture and the difference between the added and remaining (following the contact of reagent grade substances) amounts of Ca(OH)2,the CYCa molar ratios in the product were calculated to be 0.78 and 1.0 in the 9.8 and 20.2 wt % CaC12-2H20 samples, respectively. This is consistent with the DSC results indicating the absence of peaks representative of CaCly2H20 (if CaCl2.2H20 were the product, then the calculated CYCa molar ratio would equal 2). Both Ca(OH12 and CaCk2H20 are known to attract moisture when exposed to ambient air, eventually forming surface moisture film. The presence of the moisture film may in turn promote the formation of products that would not be stable at temperatures representative of the STDR operation. In order to determine the stability of the product in each reagentgrade mixture, both mixtures were exposed to 300 and 500 "C N2 flow in the STDR for 1min. No changes in peak temperatures or in the enthalpy were observed on the DSC scans, suggesting that the product of reagentgrade physical mixing was stable between 300 and 500 "C and not simply a product of reaction in the moisture film. A physical mixture of Ca(OH)2 and CaCk2H20 was then prepared and analyzed by XRD on the same day. The phases detected were Ca(OH12 (JCPDS No. 4-733) and CaClOH (JCPDS No. 36-983). An identical sample was stored for a week and then analyzed by the XRD. No Ca(OH12 was detected in this sample, with the only phase detected being CaClOH. Based on the XRD work, the eventual formation of CaClOH is proposed to occur when Ca(OH)2 and CaCl2.2H20 are contacted. Following the XRD analysis of the reagent-grade mixture, a sample of Ca(OH)2 exposed to HC1 was prepared for XRD analysis. The sample of Ca(OH)2 exposed for 1 min to 1000 ppm HC1 in the STDR operated at 500 "C was divided in half. Half of it was analyzed by XRD, and the following phases were detected: CaO (JCPDS No. 37-14971, CaCk2H20 (JCPDS No. 1-9891, CaClOH, and Ca(OH12. The DSC scan on the other half of the sample revealed peaks at 438 and 576 "C associated with enthalpy in the range representative of the decomposing Ca(OH)2 and of the product, respectively. No enthalpy representative of reagent-grade CaC12.2H20 was identified on the DSC scan. XRD analysis of another sample exposed to 1000 ppm HC1 for 1min at 200 "C revealed Ca(OH12 and CaClOH. The DSC scan was significant with 486 and 587 "C peaks. No enthalpy representative of CaC12.2H20 was identified on the DSC scan. It would appear, therefore, that upon the contact of Ca(OH)2 and HC1, the following reactions take place:
Ind. Eng. Chem. Res., Vol. 34, No. 2, 1995 611
Ca(OH), + 2HC1- CaC12*2H20
Ca(OH), Ca(OH),
+ HC1-
CaClOH*H,O
+ CaC12*2H20- 2CaC10H*2H20
(3)
(4) (5)
The resulting product of reactions 3-5 is a solid solution of CaCl2.2H20, CaClOHxHZO, and unreacted Ca(OH)2. It is possible that CaClOH is the primary reaction product between Ca(OH)2 and HC1. CaClOH then further reacts with HC1 t o form CaCly2H20. CaO Reaction with HCl. Samples of CaO prepared in the muffle furnace from Linwood Ca(OH)2, in the manner described in the experimental section, were exposed to HC1 for 1min in the STDR operated at 100600 "C. Following exposure to HC1, samples were analyzed in the DSC. The results are presented in Figure 5, giving peak temperatures of endotherms observed on DSC scans. The length of horizontal bars in Figure 5 is proportional to the amount of enthalpy measured for a given temperature range associated with the peak. Only reaction product peaks are shown in Figure 5 ; small heat flow peaks associated with rehydration of CaO (2%) are not shown in Figure 5. Significant peaks were detected on DSC scans a t temperatures ranging from 526 to 589 "C in CaO samples exposed to HC1 in the STDR. Additionally, in samples exposed to HC1 at 300,400, and 500 "C, peaks were detected at 118, 127, and 133 "C, respectively. In an approach similar to that described for Ca(OH)2, reagent-grade CaCly2H20 was mixed with CaO to investigate the possibility of a solid solution's being formed. A sample of 10.2 wt % CaC12.2H20 in CaO was prepared. The DSC scan of this sample revealed endotherm in the temperature range from 515 t o 587 "C with a peak at 563 "C. No peaks at the 44-76 "C or the 124-260 "C ranges (characteristic of reagent-grade CaCl2.2H20) were detected. Another physical mixture of reagent-grade chemicals was prepared and contained 5.1 wt % CaClz in CaO. The use of anhydrous salt (CaC12) was intended as a verification of the role of chemically bound water in the formation of solid solutions. The DSC scan of this sample revealed endotherm in the 508-561 "C range with the peak at 536 "C. By comparing the results of two experiments (CaO mixed with CaCl2-2H20 or with CaC12), it was concluded that chemically bound water was not a necessary step for the formation of a new phase. It should be kept in mind, however, that approximately 2% rehydration of CaO took place upon handling of the sample and prior to its contact with the chlorine-containing salt. In the next step, a sample for XRD was prepared by exposing CaO to 1000 ppm HC1 for 1 min in the STDR operated at 500 "C. XRD analysis was performed on half of the sample. The detected phases were CaO and CaClOH. The second half of the sample was used for DSC analysis, which determined peaks at 564 (heat flow in the 515-587 "C range) and 127 "C (heat flow in the 86-147 "C range). The analysis of an XRD spectrum of another sample of CaO reacted with HC1 a t 600 "C indicated a moderate amount of CaClOH product. A major spectral line in the 600 "C sample could not be identified from the JCPDS file. It appeared that the unresolved spectra somewhat resembled those of the CaC12.2H20, which could mean that the unknown phase was in transition to another phase.
600 -
- I
500
-
400
-
1 - 1
-
-
;
I
t
+-+-
*
300 -
zoo -
+--+--
100 -
40 Jig
*
0
v
100
0
200
300
400
500
700
600
800
Product Peak Temperature, 'C
Figure 5. DSC results for samples of CaO exposed to 1000 ppm HCl (STDR,1min). Length of horizontal bars denotes the amount of heat flow.
v'7
A A
0.05tA
0
0 Ca(OH), t HCI
A
0
0
0.1
0.2
0.3
0.4
0.5
0.6
0.7
0.8
Conversion, mols CI 7mol Ca'+
Figure 6. Correlation of AA/IC-measured sample conversion and the amount of heat flow within 530-590 "C.
Based on the results of XRD and DSC analyses, the reaction between CaO and HC1 may be described as proceeding in the following way:
CaO
+ HC1-
CaClOH
(6)
On the basis of the XRD spectral intensity, background, and widths, it appeared that the crystallinity of the chlorinated CaO sample was higher than that of chlorinated Ca(OH)2(both samples were exposed to HC1 at 500 "C). Higher crystallinity means that the crystallites are larger and/or the internal strains are lower as a result of slower product formation. Slower product formation (slower reaction) with CaO than with Ca(0H)z is consistent with the results presented in Figure 2 where, for each reaction temperature investigated, a lower level of conversion was measured for CaO than for Ca(OH)2. Summary
On the basis of the above discussion, a mixture of CaCly2H20 and CaClOH was proposed as a product of Ca(OH)2reaction with HC1. CaClOH was proposed as a product of the CaO and HC1 reaction. DSC scans on chlorinated samples revealed peaks within the 537-592 "C range and within the 526-589 "C range for samples of chlorinated Ca(OH)2 and CaO, respectively. An attempt was made to correlate sample conversion (measured by AA/IC) with the enthalpy measured by the DSC (for peaks within the two temperature ranges given above). The results are given in Figure 6 for samples of CaO and Ca(OH12 exposed for 1min t o 1000 ppm HC1 in the STDR. The conversion is expressed as
612 Ind. Eng. Chem. Res., Vol. 34,No. 2, 1995
a ratio of number of C1 moles to the number of Ca moles. The amount of enthalpy measured by the DSC correlates well with the sample conversion measured by the M C , consistent with identification of this product (characteristic with the enthalpy within the 530- 590 "C range) as a main product of reaction between CaO and HC1 or Ca(OH)2 and HC1. The occurrence of CaClOH as a product of reaction implies a CYCa molar ratio in the product of less than 2. Therefore, true utilizations of Ca-based sorbents used in flue gas cleanup are higher than those based on the assumption of a CaC12*H20product.
Conclusions The mechanisms of HC1 reaction with Ca(OH)2 and CaO were investigated by DSC, TGA, and XRD to reveal the following: 1. The product of HC1 reaction with Ca(OH)2 over a 100-600 "C range in the STDR was identified as a solid solution of CaCly2H20 and CaClOH. Subsequently the CYCa atomic ratio in the product was determined to be less than 2. 2. The product of HC1 reaction with CaO in the STDR over a 100-600 "C range was identified as CaClOH. Again the CYCa atomic ratio in the product was determined to be less than 2. 3. When scanned in the DSC, samples of both CaO and Ca(OH)2 that were previously exposed t o HC1 in the STDR revealed endothermic heat flow between 530 and 590 "C whose magnitude correlates with the chemically determined sample conversion and therefore identifies the phase responsible as a main reaction product of Ca(OH)2(or CaO) with HC1. True utilizations of Cabased sorbents used in flue gas cleanup are higher than those based on the assumption of a CaC12.H20 product. Acknowledgment The authors thank Wojciech Kozlowski (Acurex Environmental Corp.) for reactor operatiodanalyses and DSCPTGA scans. Special thanks to Frank E. Briden and the late George R. Gillis (both of the US.Environmental Protection Agency, Air and Energy Engineering Research Laboratory) for XRD analyses and equipment maintenance support, respectively.
Literature Cited Bunn, C. W.; Clark, L. M.; Clifford, I. L. The Constitution and Formation of Bleaching Powder. Proc. R. SOC.1934,141-167. Criado, J. M.; Morales J. On the Thermal Decomposition Mechanism for Dehydroxylation of Alkaline-Earth Hydroxides. J . Therm. Anal. 1976,10,103-110. Dean, J. A., Ed. Lange's Handbook of Chemistry; McGraw-Hill: New York 1985; pp 9-112. Dorsey, D. L.; Buecker, B. J. Analysis of Flue Gas Scrubber Materials from a Coal-Fired Power Plant by Thermogravimetry.
In Compositional Analysis by Thermogravimetry;Earnest, C. M., Ed.; ASTM STP 997; American Society for Testing and Materials: Philadelphia, PA, 1988; pp 254-258. Earnest, C. M. The Modern Thermogravimetric Approach to the Compositional Analysis of Materials. In Compositional Analysis by Thermogravimetry;ASTM STP 997; Earnest, C. M., Ed.; American Society for Testing and Materials: Philadelphia, PA, 1988; pp 1-18. Ettehadieh, B.; Lee, S. Y . Reducing Acid Gas Emissions from MassBurn Waste-To-Energy Plants by Direct Lime Injection. 1989 International Conference on Municipal Waste Combustion, Hollywood, FL, 1989; EPA-600/R-92-052~ (NTISPB92-174689). Felsvang, K. S.; Helvind, 0. Results of Full Scale Dry Injection Testa at MSW Incinerators Using A New Active Absorbent. 1991 International Conference on Municipal Waste Combustion, Tampa, FL, 1991; EPA-600/R-92-209~ (NTIS PB93-124196). Gullett, B. K.; Bruce, K. R.; Machilek, R. M. Apparatus for Short Time Measurements in a Fixed-Bed GadSolid Reactor. Rev. Sci. Znstrum. 1990,61,904-906. Gullett, B. K.; Jozewicz, W.; Stefanski, L. A. Reaction Kinetics of Ca-Based Sorbents with HC1. Znd. Eng. Chem. Res. 1992,31, 2437-2446. Hartman, M.; Martinovsky A. Thermal Stability of the Magnesian and Calcareous Compounds for Desulfurization Processes. Chem. Eng. Commun. 1992,111,149-160. Hsieh, J. Y.; Teller, A. J.; Zmuda, J. T.; Nakazato, K.; Wheless, E. Comparison of Performance of System Configurations of Pollution Control Systems for MSW Incineration. 1989 International Conference on Municipal Waste Combustion, Hollywood, FL, 1989; EPA-600/R-92-052d (NTIS PB92-174697). Karlsson, H. T.; Klingspor, J.; Bjerle, I. Adsorption of Hydrochloric Acid on Solid Slaked Lime for Flue Gas Clean Up. J . Air PolZut. Control Assoc. 1981,31 (111, 1177-1180. Kosnyrev, G. T.; Desyatnik, V. N.; Nosonova, E. N. Changes in the Ca(OH)z-CaCl~-2Hz0System During Thermal Dehydration. J. Appl. Chem. USSR 1990,63 (l),Part 2, 171-174. Markova, 0. A. Physical Chemistry of Calcium Hydroxide Chlorides. Russ. J . Phys. Chem. 1973a,4 , 608. Markova, 0. A. Synthesis of Pure Preparations of Calcium, Strontium, and Barium Chloride Preparation. Sb. Tr.-Mosk. 1nzh.-Stroit. Inst. 197313,No. 109,16-24. Pakrasi, A,; Davis, W. T.; Parham, M. C. Kinetic Studies on Removal of Hydrogen Chloride from MSWC Flue Gas by Dry Hydrated Lime. Presented at the 85th AWMA Annual Meeting & Exhibition, Kansas City, MO, 1992. Rabovskii, B. G.; Kukso, V. M.; Rogozhina, S. A.; Sokolov, V. A.; Khromenkov, L. G. Phase Composition of the Chlorination Products of Calcium Hydroxide. Russ. J. Inorganic Chem.,1972, 17 (3), 325-327. Schlichenmaier, V. Detection and Quantitative Determination of CaSO&HzO in CaSO~1/2Hz0Using Calorimetric DTA. Thermochim. Acta 1975,11,334-338. Weinell, C. E.; Jensen, P. I.; Dam-Johansen, K.; Livbjerg, H. Hydrogen Chloride Reaction with Lime and Limestone: Kinetics and Sorption Capacity. Znd. Eng. Chem. Res. 1992,31,164171.
Received for review March 21, 1994 Accepted October 11, 1994 @
IE940178S ~~
Abstract published in Advance ACS Abstracts, January 1, 1995. @
