Ind. Eng. Chem. Res. 1999, 38, 4068-4080
4068
Dry Absorption of HCL and SO2 with Hydrated Lime from
Humidified Flue Gas
Paul N. Chisholm and Gary T. Rochelle*
Department of Chemical Engineering, The University of Texas at Austin, Austin, Texas 78712-1062
The simultaneous absorption of HCl and SO2 by hydrated lime in a fixed-bed reactor has been
studied at conditions simulating humidified flue gas dry scrubbing. At 120 °C, an increase in
relative humidity from 0% to 19% increased HCl removal and sorbent utilization. At 19% relative
humidity (RH), the final loading of hydrated lime was 1.64 mol of HCl/mol of Ca2+. From 250 to
1000 ppm HCl, HCl removal was first order in HCl concentration. When SO2 was added to the
feed gas, the total utilization of the sorbent by HCl and SO2 was not a function of gas
concentration with 250-1000 ppm HCl and 0-2000 ppm SO2. However, the fraction of the
hydrated lime converted by SO2 increased as the SO2/HCl feed ratio increased. With oxygen
present in the feed, more SO2 was absorbed. Adding 150 ppm NO2 to the gas stream increased
the final SO2 loading from 0.06 to 0.17 mol of SO2/mol of Ca2+. As the reactivity of SO2 and NO2
increased, the reactivity of HCl decreased slightly as a result of the competition for alkalinity
with the other acid gases. The data were modeled using semiempirical flux equations based on
a modified shrinking core model. Results from the parameter estimation were used to predict
the absorption of HCl and SO2 on the surface of a bag filter. The predictions indicated that,
with a humidified flue gas and 50% sorbent utilization, less than 20% HCl penetration is possible.
However, even at the conditions with the greatest SO2 absorption (250 ppm HCl, 150 ppm NO2,
and 2.5% O2), 90% SO2 penetration is predicted at 50% sorbent utilization.
1. Introduction
Because of the passage of the Clean Air Act Amendments in 1990, many industrial processes must reduce
acid gases emissions, including HCl and SO2. These two
gases are emitted together, primarily in municipal
waste combustors (MWC), hazardous waste incinerators
(HWI), and coal-fired power plants. Published regulations (U.S. EPA, 1995) require new and large MWCs
undergoing major modifications to remove 95% of the
HCl and 75% of the SO2 emissions. Medical waste
incinerators, a common HWI, must remove 99% of their
HCl emissions and reduce the SO2 concentration to 55
ppm (U.S. EPA, 1997). New coal-fired power plants and
those undergoing major modifications, besides controlling SO2 emissions, must also report HCl releases if they
exceed 10 t/year.
If the combustor has a fabric filter in place for
particulate control, one strategy to meet these regulations is to inject a dry, alkaline sorbent into the flue
gas after it has been humidified. The solids are added
after the combustion chamber upstream of the fabric
filter. Some acid gas absorption will take place in the
duct. However, most will occur on the surface of the
fabric filter, provided care is taken to ensure that
injected solids do not fall out of the gas stream and are
deposited on the filter bags. The competing technologies
of dry sorbent injection are spray dryer absorption, wet
scrubbers, and furnace sorbent injection. The advantages of dry sorbent injection include the absence of
wastewater treatment and corrosion or steam plume
problems (wet scrubbing), the absence of nozzle plugging
(spray dryer absorption), greater sorbent utilization
(furnace sorbent injection), its relative simplicity, and
a small footprint allowing for easy retrofitting (Holzman
and Atkins, 1988).
Other researchers (Weinell et al., 1992; Pakrasi, 1992;
Jozewicz et al., 1990) have investigated the absorption
of HCl by hydrated lime in humid gas with bench-scale
apparatuses. They found that the absorption of HCl, like
SO2, is enhanced at low temperature, primarily by an
increase in the relative humidity (RH) of the gas.
Pakrasi found that at 79 °C and 45% RH, complete
utilization of hydrated lime by HCl is achieved. He
concluded that incomplete utilization was due to the
increasing barrier to diffusion presented by the calcium
chloride product layer.
Only limited studies of the simultaneous absorption
of HCl and SO2 by hydrated lime have been performed
(Matsukata et al., 1996; Jozewicz et al., 1990; Chang et
al., 1989). Those researchers have found that increased
HCl concentration increased SO2 absorption rates. Each
of these studies have also shown that the presence of
SO2 has little effect on HCl absorption.
The objective of this study was to gain a comprehensive understanding of the absorption of HCl alone and
of HCl and SO2 simultaneously by hydrated lime at lowtemperature flue gas conditions. Experiments were
performed simulating the reaction conditions of acid
gases with hydrated lime on a fabric filter for a MWC.
The acid gases contacted hydrated lime in a fixed-bed
reactor. The reactor outlet concentration of the acid
gases was measured continuously with an FT-IR spectrophotometer. One advantage of this investigation over
previous HCl-SO2-hydrated lime studies is that the
rate of acid gas absorption was measured continuously
through gas-phase analysis. Baseline studies were
performed with the single-gas systemssHCl and SO2s
and then with the combined gas system. The effects of
oxygen, NO, and NO2 in the gas phase were also
investigated. The bench-scale data were modeled using
a modification to the shrinking core theory. Results from
10.1021/ie9806601 CCC: $18.00 © 1999 American Chemical Society
Published on Web 09/14/1999
Ind. Eng. Chem. Res., Vol. 38, No. 10, 1999 4069
the modeling were used to predict the absorption of HCl
and SO2 by hydrated lime on the surface of a bag filter.
2. Reaction Chemistry
At the temperature and most of the relative humidity
ranges investigated, the most likely solid phases are
calcium chloride dihydrate (Collins and Menzies, 1936)
and calcium sulfite hemihydrate (Jones et al., 1976). The
absorption of HCl, SO2, and NO2 can be achieved by
irreversible acid-base reactions:
Ca(OH)2 + 2HCl f CaCl2‚2H2O (reaction 1)
Ca(OH)2 + SO2 f CaSO3‚1/2H2O + 1/2H2O
(reaction 2)
Ca(OH)2 + 3NO2 f Ca(NO3)2 + NO + H2O
(reaction 3)
In this study, it was discovered that when CaSO3‚
is exposed to HCl vapor, SO2 is evolved and
CaCl2‚2H2O is formed. Also, when CaCl2‚2H2O is exposed to SO2 in the absence of HCl, HCl is released by
the solids and CaSO3‚1/2H2O is formed. This reversible
reaction is given by Reaction 4:
1/ H O
2 2
CaSO3‚1/2H2O+ 2HCl + 1/2H2O T CaCl2‚2H2O +
SO2 (reaction 4)
When oxygen was present in the gas phase, S(IV)
could be oxidized to S(VI). According to Jones et al.
(1976), the dihydrate of calcium sulfate (gypsum) converts to a hemihydrate between 120 and 140 °C.
Because this study was done at the lower border of the
temperature range, it was assumed that S(VI) existed
as gypsum:
CaSO3‚1/2H2O + 1/2O2 + 3/2H2O f
CaSO4‚2H2O (reaction 5)
NO2 is expected to act as a catalyst promoting S(IV)
oxidation by reaction 5 (Shen, 1997). NO2 may also
participate directly in S(IV) oxidation:
Ca(OH)2 + CaSO3‚1/2H2O + 2 NO2 + 1/2H2O f
CaSO4‚2H2O + Ca(NO2)2 (6)
3. Experimental Section
3.1. Apparatus. The reactivity of hydrated lime with
HCl and SO2 was studied in a fixed-bed reactor (see
Figure 1). The solids were placed on an extra coarse,
0.2-cm thick glass frit at the bottom of the Pyrex reactor
(1.9 cm in diameter and 13.4 cm in height). The pressure
drop through the bed did not exceed 7 kPa. Cylinders
of HCl, SO2, NO, and NO2, each 0.5% in nitrogen, were
used to make the simulated flue gas. HCl was fed
through a rotameter whose wetted surfaces were made
exclusively of glass and Teflon. The SO2, NO, and NO2,
along with filtered house air and nitrogen, were metered
into the system using mass flow controllers. Liquid
water, fed by a syringe pump, was vaporized in the
furnace and blended with the rest of the simulated flue
gas.
The simulated flue gas, usually flowing at 1.5 SLPM
(standard conditions: 1 atm, 21.1 °C), exited the furnace
Figure 1. Schematic of the fixed-bed reactor apparatus.
and was sent to a 0.8-L glass equilibration vessel that
was submerged in the sand bath. The purpose of the
vessel was to ensure the gas was at the bath temperature and to dampen out fluctuations caused by the
evaporation of water in the furnace. After leaving the
glass vessel, the gas was sent to the reactor where it
contacted the hydrated lime (HL). HL was dispersed in
quartz sand to minimize solids agglomeration and gas
channeling. Following Karlsson et al. (1981), a sandto-sorbent mass ratio of at least 50 was used for all
experiments. The reactor was immersed in a fluidized
alumina sand bath to maintain a constant ((1 °C)
temperature. After contacting the solids in the bed, the
reactor outlet gas was mixed with 40-60 SLPM of
filtered house air. Then, the diluted reactor effluent
entered the gas cell operating with an FT-IR spectrophotometer. Dilution was necessary to prevent condensation in the lines downstream of the reactor and to
reduce corrosion of the gas cell by the acid gases,
particularly HCl. The gas cell (Infrared Analysis, Model
G-5-22-V-AU) was 13 cm in diameter, was 56 cm long,
and had an equivalent path length of 24 m. The FT-IR
(Perkin-Elmer, System 2000) recorded infrared energy
absorbance over time during an experiment.
The tubing and fittings that contacted a high concentration (i.e., g20 ppm) of HCl were made of a perfluoroalkoxy polymer (PFA). The reactor was connected into
the system with polysulfone quick-connect fittings. All
the tubing and fittings from downstream of the furnace
to just after the dilution air feed point were submerged
in the sand bath.
The sorbent, supplied by the Mississippi Lime Co.
(MLC), was similar to the hydrated lime used by Nelli
and Rochelle (1998). The specific surface area, as
measured by nitrogen porosimetry (Micromeritics, Model
ASAP 2000), was 21.1 ( 0.4 m2/g. Assuming the pores
were cylindrical, the volume of the void space for pores
between 24 and 1830 Å in diameter was 0.110 cm3/g.
Approximately half of the pore volume consisted of pores
less than 300 Å in diameter. According to MLC, the
hydrated lime was composed of 97% calcium hydroxide
with the balance being mostly calcium carbonate. A
powder X-ray analysis exhibited calcium hydroxide
peaks only. As determined by sieving, all the particles
were smaller than 149 µm; 92% (by mass) were smaller
than 45 µm.
3.2. Methods. For a typical experiment, the reactor
was charged with HL and submerged in the sand bath.
Nitrogen, with a known relative humidity, was passed
through the reactor to humidify the solids. Flow was
then directed around the reactor through a bypass line
4070
Ind. Eng. Chem. Res., Vol. 38, No. 10, 1999
the two analyses confirmed the final loading of the
hydrated lime by HCl generally within 5%.
4. Experimental Results and Discussion
Figure 2. Typical data that provided reactor outlet concentrations
of HCl. At zero minutes, the flow was directed to the reactor from
the bypass line. Experiment performed at 19% RH, 1000 ppm HCl,
120 °C, 18 mg of hydrated lime, and 1.5 SLPM. Experiment No.
5.
so that a simulated flue gas could be synthesized and
allowed to reach steady state. The gas was then directed
to the reactor where the acid gases contacted the
sorbent.
At a resolution of 8 cm-1, the FT-IR scans from 6500
to 500 cm-1 every 3 s. To reduce the noise from these
scans, 5 consecutive spectra were averaged, giving an
absorbance versus wavelength plot every 15 s. HCl,
SO2, and NO2 each have a set of characteristic peaks
that are unique. To determine a concentration versus
time profile of an acid gas during an experiment, the
area under a peak of the gas on the absorbance plot was
determined for each time interval. The peak used for
HCl was located from 2828 to 2814 cm-1; for SO2, from
1360 to 1340 cm-1; for NO2, 1655 to 1615 cm-1. The area
next to each of these peaks was also calculated and
subtracted from the area of the acid gas peaks to
account for any baseline drift over the course of the
experiment. Typical data for the HCl reaction with HL
are shown in Figure 2.
The wavenumber range used to analyze HCl was in
the same vicinity as several minor water peaks. Consequently, the noise of the HCl signal was greater at
higher relative humidity experiments where fluctuations in the water level were more pronounced. Another
problem that was more evident during high relative
humidity experiments was due to the adsorption of
water on the walls of the gas cell. HCl absorbed into
this water layer led to a damping of sharp changes in
the HCl concentration. When flow was directed to the
reactor from the bypass line, the HCl concentration at
the outlet of the reactor dropped abruptly to almost zero
as HCl reacted with the solids. With no HCl entering
the gas cell, the HCl previously absorbed in the water
layer was stripped into the gas phase. While this
phenomenon does not affect final loading calculations,
it does limit the accuracy of HCl removal data early in
an experiment, particularly with those experiments that
have high relative humidity and low HCl concentration.
Analyses of the solids were periodically performed to
confirm the gas-phase mass balance of HCl. The reacted
solids were placed in a known volume of deionized
water. The chloride concentration was determined using
a chloride ion selective electrode (Fisher, Model 13-620519). An atomic absorption spectrophotometer (Varian,
Model AA-1475) was employed to determine calcium
concentration. The mole ratio of chloride to calcium from
Four sets of gas mixtures in nitrogen were studied:
(1) HCl; (2) HCl and SO2; (3) HCl, SO2, and O2; (4) HCl,
SO2, and O2 with NO or NO2. The experimental conditions, final loading of hydrated lime by HCl, SO2, and
NO2, and total sorbent utilization are given in Table 1.
4.1. HCl Only. As expected, the absorption of HCl
by hydrated lime was found to be sensitive to relative
humidity. Figure 3 illustrates that HCl removal and
loading increased dramatically with an increase in
relative humidity. These data agree well with Pakrasi
(1992). For example, he found that HCl reacted with
HL to a maximum loading of 0.54 mol/mol of Ca2+ at
107 °C and 5.4% RH. Loading of an individual gas is
defined as the total moles absorbed of that gas divided
by the moles of calcium in the sorbent. For each 15-s
time step, the moles of HCl absorbed in the bed was
calculated by comparing the outlet HCl molar flow to
the inlet molar flow. The difference was the amount of
HCl that had reacted in the bed during that time step.
The sum of the moles of HCl reacted through time τ
divided by initial moles of HL is the average HCl loading
of the sorbent throughout the bed at time τ.
Increased reactivity of acid gases with increasing
relative humidity has been seen throughout the literature (Pakrasi, 1992; Stouffer et al., 1989; Arthur,1998).
The addition of deliquescent salts such as calcium
chloride to calcium-based sorbents has also been shown
to improve gas absorption (Ruiz-Alsop and Rochelle,
1985). For solutions in contact with CaCl2‚2H2O and
saturated to calcium chloride at 120 °C, Collins and
Menzies (1936) found that the vapor pressure of water
over the solution was 35.3 kPa. Table 2 shows the
number of monolayers that adsorbed as a function of
the bulk-phase relative humidity, assuming the BET
equation for the adsorption of water.
Because calcium chloride reduced the water vapor
pressure over the solids, a thicker layer of water
adsorbed on the surface of the sorbent. In fact, an inlet
gas relative humidity of 19% led to the formation of bulk
water on the surface of the sorbent. As the product layer
became thicker, the reactivity of HCl with the hydrated
lime increased. Though Garea et al. (1997) have proposed an explanation of the role of relative humidity,
their theory has not been conclusively proven.
Close inspection of Figure 3 shows that, at 3.5% RH,
HCl removal reached a maximum of 45% at the beginning of the experiment, decreased to about 10% removal
at 0.1 mol of HCl/mol of Ca2+ loading, and then
increased to about 40% removal before finally tapering
off to 0% removal at a loading of 0.48 mol of HCl/mol of
Ca2+. The same trend, though less pronounced, was
evident in the experiment at 0.8% RH. The experiment
at 3.5% RH was reproduced five times; the experiment
at 0.8% RH was duplicated. Each of the seven experiments showed the same drop in reactivity at about 0.1
mol of HCl/mol of Ca2+ loading and then a resumption
of reactivity. While it is not clear in Figure 3, there was
a greater time lapse before the reaction restarted with
the experiment at 0.8% RH than with the experiment
at 3.5% RH.
The deliquescent nature of calcium chloride may
explain this unexpected phenomenon. The solids react
to a loading of about 0.1 mol of HCl/mol of Ca2+,
Ind. Eng. Chem. Res., Vol. 38, No. 10, 1999 4071
Table 1. Experimental Conditions and Results (All Experiments Performed at 120 °C and 1.5 SLPM, unless Indicated
Otherwise)
exp
no.
rel. hum.
(%)
mass
sorbent (mg)
time
(min)
HCl
(ppm)
SO2
(ppm)
O2
(%)
NOx
(ppm)
HCl
load.a
SO2
load.a
NO2
load.a
total util.b
(%)§
1
2
3
4
5
6
7
8
9c
10
11
12
13
14
15
16
17
18
19
20
21
22
23
24
25
26
27
0
1
3.5
9
19
19
19
19
19
19
19
19
19
19
19
19
19
19
19
19
19
19
19
19
19
9
9
18
18
18
18
18
18
18
18
18
35
35
35
35
50
50
50
100
35
35
35
35
35
35
35
35
36
18
24
45
45
45
45
55
45
45
45
110
65
45
45
45
45
45
40
45
110
110
60
100
45
100
100
45
45
1000
1000
1000
1000
1000
250
500
2000
3500
250
500
1000
2000
1000
1000
1000
0
0
250
250
1000
250
0
250
250
1000
500
0
0
0
0
0
0
0
0
0
1000
1000
1000
1000
250
500
2000
1000
1000
1000
1000
1000
1000
1000
1000
1000
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
5.5
1.2
5.5
5.5
0
0
0
2.5
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
NO:150
NO2:150
NO2:150
NO2:150
0
0
0.101
0.173
0.488
1.07
1.61
1.66
1.66
1.43
1.27
1.58
1.47
1.52
1.48
1.60
1.54
1.37
0
0
1.48
1.52
1.50
1.52
0
1.44
1.33
1.122
1.020
0
0
0
0
0
0
0
0
0
0.0606
0.0748
0.0469
0
1.79e-3
5.02e-3
0.0704
0.108
0.204
0.0903
0.152
0.106
0.143
0.126
0.174
0.310
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0
0.0534
0.123
0.100
0
0
5.05
8.65
24.4
53.5
80.5
83.0
83.0
71.5
63.5
85.6
81.0
80.7
74.0
80.2
77.5
75.5
10.8
20.4
83.0
91.2
85.6
90.3
15.3
95.6
102
56.1
51.0
a Loading of acid gas A is given by mol of A/mol of Ca2+. b Total utilization takes into account stoichiometry:
1
total utilization )
c Total flow rate:
/2nHCl + nSO2 + 1/2nNO2
0.69 SLPM.
Figure 3. Effect of relative humidity and loading on HCl removal.
Experiments performed at 1000 ppm HCl, 120 °C, 18 mg of
hydrated lime, and 1.5 SLPM. Experiments 1-5.
Table 2. Thickness of Water Layer on Hydrated Lime at
120 °C
bulk gas
rel. hum. (%)
water partial
pressure (kPa)
no. of water
monolayers
w/o CaCl2a
0
0.8
3.5
9
19
1.6
3.7
7.2
18
39
0
0.4
0.8
1.0
1.2
no. of water
monolayers
w/ CaCl2b
0
0.8
1.2
2.1
∞ (bulk liquid)
a Assumes that no CaCl is present in the liquid layer. b As2
sumes that the liquid layer is saturated with CaCl2.
producing a layer of CaCl2‚2H2O on the HL particles
per reaction 1. The calcium chloride then absorbed
water from the vapor phase. The reaction started up
again once there was a sufficient degree of surface water
nCa
× 100
Figure 4. Effect of HCl concentration and loading on HCl
removal. Experiments performed at 19% RH, 120 °C, 18 mg of
hydrated lime, and 1.5 SLPM (unless indicated otherwise). Experiments 5-9.
on the particles. The same phenomenon may have
occurred with the higher relative humidity experiments,
but because the water vapor rates were much greater,
the rehydration of the particles took place too quickly
to observe. While this behavior is interesting scientifically, further studies were not conducted to conclusively
explain its cause.
Over the typical HCl concentration range in flue gas
for MWCs and coal-fired boilers, the HCl removal rate
was first-order in the HCl concentration. Also, the final
HCl loading was not a function of concentration. However, at an HCl concentration of 2000 and 3500 ppm,
the increased HCl concentration decreased the final
loading. This trend is evident in Figure 4. From 250 to
1000 ppm HCl, removal was equal within experimental
error up to the final loading of 1.64 mol of HCl/mol of
4072
Ind. Eng. Chem. Res., Vol. 38, No. 10, 1999
Ca2+ loading. At a concentration of 2000 and 3500 ppm,
HCl removal remained high until a loading of approximately 0.6 mol of HCl/mol of Ca2+. However, the
absorption shut off more quickly to final loadings of 1.44
mol of HCl/mol of Ca2+ at 2000 ppm HCl and 1.28 mol
of HCl/mol of Ca2+ at 3500 ppm HCl. The experiments
at 2000 and 3500 ppm were each reproduced three
times. Experimental constraints prohibited experiments
with a concentration greater than 3500 ppm.
Previously, only Pakrasi (1992) had investigated the
effect of increased HCl concentration on the final
loading. The results of his experiments, which were
conducted at conditions similar to those in Figure 4, did
not reveal a decrease in HL utilization with the HCl
concentration from 1000 to 5000 ppm. However, his
analysis to determine HCl loading was done with a
maximum sorbent exposure time of only 5 min. It is
likely that, under highly reactive conditions, HL would
have continued to react after the first 5 min of the
experiment. Therefore, the loading at 5 min does not
represent an accurate final loading because lower molar
flows of HCl will take longer to completely react with
HL. In addition, the average residence time of reactive
solids is usually greater than the 5 min.
Reduced final sorbent loading at higher HCl concentrations may have been due to an increase in the
plugging of the pores of the sorbent. Simons’ (1992)
analysis of the reaction of SO2 with calcined limestone
showed that, under certain conditions, utilization is
incomplete because pores became plugged near the
external particle surface, leaving the interior of the
sorbent unavailable for reaction. When the HCl concentration in the gas was high, the rate of reaction at
the surface of the particle was greater. Therefore,
reaction products accumulated at the opening of the
pore, eventually leading to plugging. When the HCl
concentration was low, the lower reaction rate at the
surface of the solid allowed for increased diffusion of
the gas into the interior of the particle. The results
shown in Figure 4 imply that, at a concentration of 1000
ppm and below, the reaction rate at the surface of the
particle was low enough to allow the HCl to diffuse into
the pores and utilize more alkalinity.
4.2. HCl and SO2. As with HCl, the reactivity of SO2
with HL in the absence of HCl increased with increasing
relative humidity. Unlike HCl, decreasing SO2 concentration increased the fraction of SO2 removed. These
effects of relative humidity and SO2 concentration on
SO2 removal (Chisholm, 1999) followed the same trends
as those found by other researchers (Jorgensen et al.,
1986; Klingspor et al., 1984; Ruiz-Alsop and Rochelle,
1985).
Most of the experiments performed with HCl and SO2
were conducted at high relative humidity (19%). This
was done because, at lower relative humidity, the
reactivity of SO2 was too low to observe trends within
acceptable precision. At early experimental times, the
absorption of SO2 was limited by having to compete with
HCl for alkalinity. Increased SO2 absorbance with a
lower concentration of HCl is evident in Figure 5. Note
that the total loading ((0.5 mol of HCl + mol of SO2)/
mol of Ca2+) is defined somewhat differently than the
HCl loading (mol of HCl/mol of Ca2+). Later in the
experiment, as the alkalinity of the sorbent became
consumed, HCl reacted with CaSO3‚1/2H2O, releasing
SO2 as indicated by reaction 4.
Figure 5. Effect of HCl concentration and total loading on SO2
removal. Experiments performed at 19% RH, 1000 ppm SO2, 35
mg of hydrated lime, 120 °C, and 1.5 SLPM. Experiments 10-13.
Figure 6. Effect of SO2 concentration and total loading on HCl
removal. Experiments performed at 19% RH, 1000 ppm HCl, 35
mg of hydrated lime, 120 °C, and 1.5 SLPM. Experiments 12 and
14-16.
It was found that the concentration of SO2 in the gas
phase had little effect on the reactivity of HCl. This is
evident in Figure 6. The minor differences in the
breakthrough curves can be attributed to differences in
the quality of the distribution of HL in the reactor. If
there was more sorbent on one side of the reactor, the
curve would have broken early because some of the gas
was bypassing the reactive solids on the side with less
sorbent. The fact that all the experiments ended up at
approximately the same final total loading suggested
that there was, at most, only a minimal effect of SO2
concentration on HCl removal.
In all of the HCl-SO2 experiments, the sorbent
utilization achieved was the same as the utilization
achieved with only HCl in the gas stream. At 19% RH
and an HCl concentration of 1000 ppm and below, this
final total loading was 0.82 (0.5 mol of HCl + mol of
SO2)/mol of Ca2+. This fact implies that, with HCl and
SO2, there seemed to be a maximum available alkalinity
primarily defined by the relative humidity, regardless
of the feed gas composition. When the ratio of SO2/HCl
in the feed gas increased, slightly less CaCl2‚2H2O was
formed in favor of a greater amount of CaSO3‚1/2H2O.
Though with considerable scatter, this trend is evident
in Figure 7.
The fact that the final solids product composition is
a function of the feed gas composition suggests that a
pseudoequilibrium exists, probably in accordance with
reaction 4. Experiments (Chisholm, 1999) where solids
were exposed to just HCl or SO2 in sequence showed
that not only would HCl react to drive off SO2 but also
Ind. Eng. Chem. Res., Vol. 38, No. 10, 1999 4073
Figure 7. Effect of the SO2/HCl feed ratio on fraction of total
loading by SO2. Experiments performed at 19% RH, 35 mg of
hydrated lime, 120 °C, and 1.5 SLPM. Two data points had the
same fraction of total loading by SO2 of 9.3% at an inlet SO2/HCl
ratio of 2. Experiments 14-19.
Figure 8. Effect of O2 and HCl on the loading of SO2. Experiments
performed at 19% RH, 1000 ppm SO2, 35 mg of hydrated lime,
120 °C, and 1.5 SLPM. Experiments 10, 17-21.
the reverse was true. The concentration of SO2 emitted
when HCl was exposed to CaSO3‚1/2H2O increased as
the HCl concentration increased. The same trend was
observed when CaCl2‚2H2O was exposed to varying SO2
concentrations. Reagent calcium chloride also absorbed
low levels of SO2 while simultaneously emitting HCl.
One explanation for the pseudoequilibrium may be
reaction kinetics. Another explanation may be that the
chemistry dictates that HCl would react completely with
CaCl2‚2H2O. However, CaSO3‚1/2H2O became trapped
in the pores and was inaccessible to HCl.
4.3. HCl-SO2-O2 and Hydrated Lime System.
Adding air to a dilute HCl gas stream did not affect HCl
reactivity (experimental data not shown). With a similar
sorbent, Chu and Rochelle (1989) have found that the
presence of oxygen in the feed gas led to greater SO2
removal and final loading. With air present in the HClSO2 gas mixture, the oxidation of S(IV) to S(VI) by
reaction 5 could take place. Figure 8 shows the effect of
air when added to the HCl-SO2 mixture.
Several conclusions can be drawn from the experiments shown in Figure 8. First, the presence of O2
increased the reactivity of SO2 with HL. This is clear,
particularly in the experiments without HCl. When the
experiments containing 5.5% O2 with and without 250
ppm HCl are compared, it is apparent that there was a
slight increase in SO2 reactivity early in the experiment
due to the presence of HCl. However, once the HCl
moving front pushed to the end of the reactor and
alkalinity started to become scarce, HCl reacted with
CaSO3‚1/2H2O, causing SO2 emissions. As the concentration of oxygen increased, the amount of sulfur that
remained bound in the solids also increased. This
suggests that, with a greater oxygen concentration,
there was a greater amount of S(VI) formed through
reaction 5. When there was only 250 ppm HCl in the
gas stream, the SO2 did not have to aggressively
compete with HCl; the SO2 absorption probably took
place at the back of the reactor where HCl had not
broken through. But when the concentration of HCl was
increased to 1000 ppm, the rapid absorption of HCl and
consumption of alkalinity led to reduced SO2 reactivity.
This effect is seen clearly in the comparison of the
experiments with 250 and 1000 ppm HCl with a
constant O2 concentration of 5.5%.
In summary, if SO2 did not have to compete with HCl
for alkalinity, it reacted almost as if the HCl were not
in the feed gas. Once SO2 reacted, the presence of
oxygen in the gas stream allowed the irreversible
oxidation of S(IV). But the oxidation did not take place
quickly enough to force all the bound sulfur to S(VI).
Consequently, SO2 was emitted as HCl began to break
through to the back of the bed. At a high HCl concentration, SO2 had to compete with HCl for alkalinity. This
competition caused a reduction in SO2 reactivity, even
in the presence of oxygen.
While the dynamics of SO2 absorption in the presence
of HCl and O2 were somewhat complex, HCl reactivity
was relatively uninteresting. For each of the experiments shown in Figure 8, the final loading of HCl was
1.5 mol of HCl/mol of Ca2+ within experimental error.
Therefore, in the experiments where the SO2 loading
was greater, the total loading was also greater. Without
oxygen present, sorbent utilization with HCl-SO2 is
about 82% at 19% RH. The experiment with 250 ppm
HCl, 1000 ppm SO2, and 5.5% O2 at 19% RH yielded a
final utilization of 91%s76% by HCl and 15% by SO2.
4.4. HCl-SO2-NOx-O2 and Hydrated Lime System. In general, NO accounts for approximately 90%
of the NOx emitted in flue gas; the remaining 10% is
NO2. Yuan and Fong (1992) studied the reactivity of the
HCl-SO2-NO system with a calcium hydroxide slurry
in a simulated spray-drying apparatus. Their results
show that fractional removal of NO by the slurry is an
order of magnitude less than those for HCl and SO2.
They concluded that increasing NO concentration had
a slight positive effect on SO2 absorption. In a study
done without HCl, Nelli and Rochelle (1998) found that
increasing NO2 concentration improved SO2 reactivity.
Greater SO2 concentration also enhanced NO2 absorption.
The experiments investigating the effect of NOx in the
feed gas were performed at 250 ppm HCl. This was done
so that the absorption of SO2 and NOx could be observed
without those gases having to compete aggressively with
HCl for alkalinity. A high concentration of SO2 and a
low concentration of HCl most closely represent coalfired boiler flue gas. However, it is expected that the
conclusions drawn below are qualitatively applicable to
flue gases with other SO2/HCl ratios. Experiments
performed in this study generally confirm the trends
found by previous researchers. Figure 9 shows the
effects of NO and NO2 on SO2 removal.
In 50 min, the presence of 150 ppm NO decreased the
SO2 loading by 0.05 mol of SO2/mol of Ca2+ when all
other variables were held constant. This effect was
probably due to the competition between SO2 and NO
4074
Ind. Eng. Chem. Res., Vol. 38, No. 10, 1999
Figure 9. Effect of HCl, NO, and NO2 on SO2 loading. Experiments performed at 19% RH, 250 ppm HCl, 1000 ppm SO2, 35
mg of hydrated lime, 120 °C, and 1.5 SLPM. Experiments 10, 20,
and 22-25.
for available alkalinity at the back of the sorbent bed.
In the experiment without NO, by 100 min the HCl
reactivity had forced the emission of SO2 by Reaction
4. Therefore, loading of SO2 in the experiments with and
without NO ended up within 1 mol of SO2/mol of Ca2+
of each other.
The loading of the solids by NO could not be accurately calculated because the NO peaks in the IR
spectrum were washed out by water peaks. Though
water was taken as part of the background, fluctuations
in the syringe pump and the water content of the
dilution air made the NO signal too noisy for a mass
balance. Qualitatively, it was apparent that NO reacted
at early experimental times but ceased absorbing (or
desorbing) after about 3 min.
One strategy for NO capture is to oxidize the relatively unreactive NO to NO2 by injecting a hydrocarbon
such as methanol into the flue gas (Lyon et al., 1990).
With this strategy in mind, experiments with NO2
added to the feed gas were conducted. The presence of
NO2 clearly increased the reactivity of SO2 with HL.
When 150 ppm NO2 was added to the gas stream in the
absence of oxygen, the final loading of SO2 was 0.11 mol
of SO2/mol of Ca2+ greater than that in the absence of
NO2. With 150 ppm NO2 and 2.5% O2 in the feed gas,
the final loading of SO2 reached 0.31 mol of SO2/mol of
Ca2+. Even as the solids were approaching complete
utilization, the absorption of SO2 continued. There are
several potential causes of the continued SO2 reactivity.
It is possible that NO2 was catalyzing the formation of
sulfuric acid (Duecker and West, 1959). However, the
analysis of reacted solids with a similar sorbent (Chisholm and Rochelle, 1999) did not support this conclusion. A free radical mechanism for the oxidation of S(IV)
by NO2 has also been reported (Shen, 1997). But this
mechanism was found to exist at high pH, a condition
probably not present on the surface of highly converted
solids. The improved SO2 absorption is probably due to
enhanced sorbent utilization from the formation of
sulfur-nitrogen compounds (Nelli and Rochelle, 1998).
The loading of NO2 was calculated by means of the
gas-phase mass balance. Figure 10 plots the loadings
of NO2 for two of the same experiments shown in Figure
9. The addition of 250 ppm HCl to the system increased
the NO2 loading. However, adding 2.5% O2 caused a
reduction in the loading of NO2 by 0.02 mol of NO2 mol
of Ca2+. This reduction in reactivity arose because with
the oxygen in the feed stream, S(IV) could be oxidized
Figure 10. Effect of HCl and O2 on NO2 loading. Experiments
performed at 19% RH, 250 ppm HCl, 1000 ppm SO2, 150 ppm NO2,
35 mg of hydrated lime, 120 °C, and 1.5 SLPM. Experiments 2325.
to S(VI). If S(IV) was oxidized by oxygen, it was not
available to react with NO2 by reaction 6.
As with the HCl-SO2-O2 system, the reactivity of
HCl when NOx is added did not change considerably.
Without NOx or oxygen in the gas, an experiment with
250 ppm HCl and 1000 ppm SO2 yielded a final loading
of HCl of 1.58 mol of HCl/mol of Ca2+. The final loading
of SO2 for the same experiment was 0.06 mol of SO2/
mole Ca2+. Under conditions with the most SO2 and NOx
reactivity (150 ppm NO2 and 2.5% O2), the final loading
of HCl dropped to 1.34 mol of HCl/mol of Ca2+ while
the final loading of SO2 and NO2 were 0.31 mol of SO2/
mol of Ca2+ and 0.10 mol of NO2/mol of Ca2+, respectively. So even when HCl must compete the most for
alkalinity, its reactivity still suffers only minimally.
5. Fixed-Bed Modeling and Bag Filter
Predictions
5.1. Model Description. The primary objective of
modeling the experimental results was to match the
data with a set of equations that accurately described
HCl and SO2 concentrations over the course of an
experiment. The primary objective was not to prove
conclusively any gas-solid reaction models. Results of
the fixed-bed modeling were used to predict the absorption of HCl and SO2 by hydrated lime on the surface of
a bag filter.
5.1.1. Fixed-Bed Mass Balance. The first step to
modeling the gas system was to set up the appropriate
mass-balance equations for the absorption of a gas by a
solid in a fixed-bed reactor. Non-steady-state experiments were performed at integral conditions where
there existed a concentration gradient through the bed
in the same direction as the gas flow. The gas concentration as a function of time and distance is described
by eq 1:
FA )
1 ∂CA
Q ∂CA
a ∂t
aAx ∂z
(1)
where FA is the flux of A to the sorbent in the fixed bed,
CA is the concentration of species A, t is the experimental time, z is the distance in the bed, a is the sorbent
surface area per unit reactor volume, Q is the gas flow
rate, and Ax is the cross-sectional area of the reactor.
In most cases, the time derivative is much less than the
spatial derivative (von Rosenberg et al., 1977). Therefore, eq 1 can be reduced to
Ind. Eng. Chem. Res., Vol. 38, No. 10, 1999 4075
FA ) -
Q dCA
aAx dz
(2)
The solid-phase reactivity is described by eq 3:
FA ) -
R dx
R dCS
)
a dt
Sm dt
(3)
where CS is the concentration of the sorbent, R is a
stoichiometric factor (R ) 2 for HCl; R ) 1 for SO2), Sm
is the molar surface area of the sorbent, and x is the
conversion.
Equations 2 and 3 were transformed to algebraic
equations and Euler’s method was used to predict the
gas concentration through the bed. Because Euler’s
method was employed, care was taken to ensure that
sufficiently small time and distance steps were taken
during parameter estimation calculations.
To estimate the parameters in a flux equation (discussed below), calculations began at the top of the
reactor at time zero. A spatial mass balance in a finite
element generated the inlet to the next finite element
downstream. Another mass balance in the finite element
determined the conversion for the next time step in the
same element. Calculations were performed through the
bed to generate an outlet concentration, CA,L|τ from the
reactor at time τ.
For SO2, CSO2,L|τ was compared directly to the datum
taken at time τ, CSO2|τ. However, with the more soluble
HCl, water absorption on the surfaces of the gas cell
caused confounding effects. Therefore, the gas cell
surfaces were modeled as a well-mixed absorber/stripper
where CHCl,L|τ was the inlet to the gas cell. The equation
used to generate a modeled concentration, CHCl0|τ, to a
concentration datum obtained at the same t, CHCl|τ, is
given below:
CHCl0|τ ) CHCl,L|τ +
c
(C
| - CHCl0|τ-∆τ) (6)
∆τ HCl,L τ
where c is an empirical parameter estimated uniquely
for each experiment. The driving force in eq 6 accounted
for molecules of HCl stripping out of or absorbing into
the water layer on the surfaces of the gas cell.
Once the outlet concentrations were generated from
the model at time τ, they were compared to the data
taken for the same time. The same calculations through
the bed were performed for each of the time intervals
and compared to the experimental data. Once this was
completed for an entire experiment, the parameters
were adjusted to improve the fit and the calculations
started again from the top of the bed at time zero.
5.1.2. Flux Equation. Challenges with modeling
arose not with the mass balances discussed above but
with the identification of an accurate equation describing the flux of the acid gases, FA. Many researchers
(Gullett et al., 1992; Wang et al., 1996; Fonseca et al.,
1998) have employed shrinking core models, either with
particles or grains within particles. These models
described the rate of absorption of an acid gas as being
dictated by mass-transfer resistances in series. In most
systems with small particles, the gas film and pore
diffusion resistances do not limit the rate of gas absorption. At early times in an experiment before a product
layer has deposited on the surface of a reactant particle,
the rate of gas absorption was limited only by the
kinetics of the surface reaction. As the acid gas contin-
ued to react with the sorbent particle, a product layer
formed. This product layer then acted as a barrier
between the gas and the reactive solid core of the
particle, thereby limiting the rate at which the gas was
absorbed. Assuming first-order kinetics and diffusion
through the product layer were the limiting rates in gas
absorption, the following equation describes the flux of
gas A:
FA )
CA
1
δ
+
k s DA
(7)
where ks is the first-order surface rate constant, δ is
the product layer thickness, and DA is the diffusion
coefficient of the diffusing species through the product
layer.
In developing eq 7, planar rather than spherical
geometry was used. This was done primarily to simplify
the flux expression. When the sorbents absorbed the
acid gases studied, the volume of the particles expanded
because the absorption products have greater molar
volumes than hydrated lime. This effect was ignored
again in the interest of simplicity. As shown in eq 7,
the flux was a function of the product layer thickness.
This thickness was not measured. Because conversion,
x, was proportional to product layer thickness, δ was
replaced by δox.
Throughout the literature and in this study, it was
apparent that most sorbents generally were not totally
converted from the alkaline particle to the salt product.
At some point during an experiment, the absorption rate
declined to zero, despite the fact that there was still
unreacted sorbent in the reactor. However, the shrinking core model predicts that the rate will not drop to
zero until all the sorbent has been consumed. Incomplete utilization may be due to pore plugging (Simons
and Garman, 1986) or reaction product expansion (Duo
et al., 1994). Because of the discrepancy between the
shrinking core model and the experimental results, a
driving force based on conversion was added to eq 7:
FA )
CA(xT - x)
1 δ0x
+
k s DA
(8)
xT is defined as the conversion when reactivity of the
sorbent terminated. The three regressed parameters in
eq 8 are xT, ks, and DA/δ0. The value DA/δ0 will be
referred to as DA,eff. Principles leading to eq 8 are used
with both the single-gas system (HCl) and with the
multi-gas system (HCl and SO2). However, another
modification to the flux equation for SO2 was found to
be necessary. This modification will be discussed later.
This modification to the shrinking core equation and
the other assumptions discussed above led to a semiempirical model for the flux. However, the primary objective of the modeling of experiments was to describe the
data. The goal of modeling was not to determine
conclusively an absorption mechanism or accurate
physical constants.
The parameter estimation package used was Generalized REGression (GREG) (Caracotsios, 1986; Stewart
et al., 1992). GREG was employed in this work to
estimate parameters with multivariable inputs but only
one response vector. The modeled concentration of the
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Ind. Eng. Chem. Res., Vol. 38, No. 10, 1999
outlet of the reactor normalized by its inlet concentration was compared to the same ratio determined experimentally. Code was written to link GREG with the
fixed-bed integration routine. Parameters with confidence intervals were estimated and reported.
5.1.3. Bag Filter Performance Model. The objective of the modeling effort was to project how well
hydrated lime and calcium silicate would perform
absorbing HCl and SO2 on the surface of a bag filter.
The bag filter system is physically different from the
experimental system used in this study. In the fixedbed reactor, high concentration acid gas contacts mostly
converted the sorbent in the front of the reactor.
Conversely, in a bag filter the high-concentration gas
will pass through fresh sorbent first and will contact
the most highly converted solids just before it escapes
the filter. In addition, the bag filter system has sorbent
continually added to the filter cake, whereas the experimental system has a batch loading of sorbent.
Because of these differences, predicting how well a
sorbent would perform under given circumstances was
not straightforward using only the experimental data.
The bag filter performance model assumed that the flux
expressions that govern the absorption of HCl and SO2
in the experimental system also described the flux
expressions for the bag house. Also, it was assumed that
the sorbent contacts the surface of the bag to form a
uniformly thick layer. No reaction took place in the duct
between the introduction of the sorbent and when it
reached the bag house. In addition, it was assumed that
the presence of any other contaminates, gas or solid,
would not influence the acid gas absorption. Changes
in pressure drop with increased filter cake thickness
were not calculated.
One of the key operational conditions that impacted
the performance projections was the sorbent feed ratio.
This ratio was the number of moles of sorbent fed to
the system divided by the number of moles needed to
remove all the acid gas, assuming 100% utilization of
the solids by simple acid-base reactions. For example,
if 1 mol each of SO2 and HCl entered a flue duct, a
sorbent feed ratio of 2 meant 3 mol (2 x (1 mol + 0.5
mol)) of sorbent was fed into the duct.
The second important operational variable in a bag
filter system was the cycle timesthe period between
cleanings of the bag. Bag filters are generally cleaned
by a reversed air stream, by a pulse jet, or by simply
shaking it. After a bag was cleaned, it was assumed that
there was no cake on the surface of the filter. Consequently, removal from that bag was lowest immediately
after being cleaned and increased with increasing cycle
time.
The calculations for bag filter performance used the
surface of the filter as the origin. As the cycle time
continued, the filter cake became thicker. Over time,
t*, the average penetration of gas A was defined as the
time average of the instantaneous penetration of A from
0 to t*:
average penetration of A ) 1 - removal of A )
1 t* CA(t) dt
(9)
t* 0 CA,in
∫
where CA(t) is the instantaneous concentration of A
leaving the bag house and CA,in is the inlet concentration
of A to the bag house.
The quality of the bag filter performance projections
were only as good as the model fit of the data. The
experiments attempted to quantify the phenomena as
accurately and precisely as possible. Imperfections with
the data and the modeling, then, may have manifested
in the performance projections. Consequently, caution
should be used such that the precision of the performance predictions is not overestimated.
5.2. HCl Only. 5.2.1. Fixed-Bed Modeling. The
focus of parameter estimation for the HCl-hydrated lime
system was the experiments conducted at 9% and 19%
humidity. As discussed earlier, the experiments at 0.8%
and 3.5% RH exhibited unexpected, transient behavior.
Because of the potential difficulty in modeling the data
at these low relative humidities, the emphasis of modeling was placed on the high-humidity experiments. In
addition, the high-concentration data (2000 and 3500
ppm HCl) were not modeled because those experiments
exhibited final conversion different from those obtained
at low and intermediate HCl concentrations.
The data for the absorption of HCl by hydrated lime
was first modeled with the flux described by eq 8. Using
eq 8, the estimation routine systematically underpredicted the outlet concentration of HCl absorption in the
first 5 min of each experiment. For each of the data sets,
100% removal of HCl was predicted for nearly all the
experiments. For this reason, eq 8 was modified to a
pseudo-second-order dependence on HCl concentration
at early experimental times to better describe absorption. Beyond the time for initial reaction, HCl absorption
was still described as first-order in HCl. This first-order
dependence over the majority of the active absorption
regime reflected the same concentration dependence
seen in Figure 4. Equation 10 was used as the flux
equation for the absorption of HCl by hydrated lime:
FA )
CHCl(xT - x)
1
x
+
ksCHCl Deff,H
(10)
The results of parameter estimation for experiments run
at 9% and 19% at concentrations less than 2000 ppm
are given in Tables 3 and 4, respectively.
Note that, in Tables 3 and 4, the experimental data
were modeled both individually and simultaneously
with the other experiments performed at the same
relative humidity. Different values for the gas cell
correction constant, c, were estimated for each experiment and were all between 5.6 and 12.6 s. The values
of ks were determined mostly during the first few
minutes of the experiment before much of a product
layer of calcium chloride had developed on the surface
of the sorbent. Values of Deff,H were determined as HCl
was breaking through the bed. The 95% confidence
intervals of Deff,H for the two humidities almost overlapped. The values of xT were considerably different, as
expected because of the different chloride capacities
observed in the bench-scale studies.
Figure 11 compares the model fit to the data for
experiment 6, the experiment with the greatest error
of those regressed. The fit shown represents a worst case
indication of the accuracy of the model.
5.2.2. Bag Filter Predictions. The regressed parameters for the simultaneous regressions shown in
Tables 3 and 4 were used to predict the absorption of
HCl by hydrated lime on the surface of a bag filter. The
plot showing the trade-off between sorbent utilization
Ind. Eng. Chem. Res., Vol. 38, No. 10, 1999 4077
Table 3. Results of Parameter Estimation for the HCl-Hydrated Lime System at 9% RH
4
26
27
sorb.
load.
(mg)
HCl
conc.
(ppm)
exp.
no.
simult. model
2σ:
1000
2σ:
1000
2σ:
500
2σ:
18
36
18
Deff,H
× 104
(m/s)
ks
(m4/mol/s)
xT
3.20
0.36
3.03
0.55
3.90
0.97
4.33
0.93
0.613
0.185
0.352
0.230
0.172
0.063
0.872
0.568
0.497
0.006
0.465
0.011
0.508
0.008
0.465
0.010
mean
model
errora
simult.
mean
model errorb
4.22
3.09
4.13
3.48
4.20
4.29
4.76
a The error reported is the square root of the average model error squared multiplied by 100 when the experiment was modeled
individually. b The error reported is the square root of the average model error squared multiplied by 100 when the experiments were
modeled simultaneously.
Table 4. Results of Parameter Estimation for the
HCl-Hydrated Lime System at 19% RH
HCl sorb. Deff,H
ks
exp. conc. load. × 104
no. (ppm) (mg) (m/s) (m4/mol/s)
simult. model
2σ:
5
1000 18
2σ:
6
250 18
2σ:
7
500 18
2σ:
2.67
0.20
3.38
0.46
4.23
0.98
1.87
0.11
0.160
0.015
0.112
0.029
0.129
0.015
0.243
0.045
xT
0.795
0.008
0.748
0.012
0.802
0.012
0.794
0.010
mean
simult.
model mean model
errora
errorb
4.09
2.63
3.08
4.76
5.20
2.59
3.14
a The error reported is the square root of the average model
error squared multiplied by 100 when the experiment was modeled
individually. b The error reported is the square root of the average
model error squared multiplied by 100 when the experiments were
modeled simultaneously.
Figure 12. HCl penetration as a function of hydrated lime
utilization at varying cycle times. Performance projection with 19%
RH and 1000 ppm HCl at 120 °C. The curves use the global
parameters given in Table 3.
concentration and was similar in form to the flux
equation for HCl:
FSO2 )
CSO2(xT - x)
1
x
+
ksSCHCl Deff,S
(11)
where x at time τ is the sum of the moles of HCl, nHCl,
and SO2, nSO2, that have been absorbed since time zero,
adjusted for reaction stoichiometry and divided by the
initial moles of alkalinity:
Figure 11. Comparison of model fit to the data for experiment
6: 19% RH, 250 ppm HCl, 18 mg of hydrated lime, 120 °C, and
1.5 SLPM.
and HCl penetration is shown in Figure 12. At low
utilization (and a very high sorbent feed ratio), nearly
0% penetration is possible. As the sorbent feed ratio
decreases, the utilization increases to almost xT. But
when the sorbent is being utilized as much as possible,
almost 100% HCl penetration is predicted. Figure 12
projects that to reduce HCl penetration to 10%, slightly
more than half the sorbent will be utilized.
5.3. HCl and SO2. 5.3.1. Fixed-Bed Modeling. The
addition of SO2 to the inlet gas stream led to an increase
in the complexity of the modeling approach. Though
analysis of the gas system became more complex, the
flux equation for HCl was the same as the one used for
the HCl-only system given by eq 10. With respect to SO2
absorption, it was shown above that the presence of O2
and NO2 as oxidizing agents of S(IV) had a positive
influence on SO2 absorption. The flux equation for SO2
absorption did not explicitly account for O2 or NO2
1
xτ )
/2nHCl|τ + nSO2|τ
nCa
(12)
At early experimental time (x f 0), eq 11 described the
SO2 flux as being first-order in both SO2 and HCl
concentrations. The same effect of HCl concentration on
SO2 absorption was seen with calcium silicate as the
sorbent (Chisholm and Rochelle, 1999). It is believed
that the deliquescent nature of CaCl2 salts improved
SO2 absorption. This same effect was seen by other
researchers (Matsukata et al., 1996; Jozewicz et al.,
1990). Other experiments (Chisholm, 1999) demonstrated that an increase in relative humidity over
hydrated lime increased SO2 absorption. As the relative
humidity increased, the thickness of the layer of liquid
water adsorbed on the sorbent increased. If the solids
had surface solution that was saturated to CaCl2, the
vapor pressure of water over the solids would have been
reduced. This reduction may have led to a thicker
adsorbed water layer and increased SO2 absorption.
4078
Ind. Eng. Chem. Res., Vol. 38, No. 10, 1999
Table 5. Results of Parameter Estimation for the HCl-SO2 System
exp.
no.
HCl
conc.
(ppm)
SO2
conc.
(ppm)
NO2
conc.
(ppm)
O2
conc.
(%)
Deff,H
× 104
(m4/mol/s)
ks,H
(m/s)
25
250
150
2.5
24
250
150
0
20
250
0
5.5
21
1000
0
5.5
16
1000
0
0
12
1000
1000
2σ:
1000
2σ:
1000
2σ:
1000
2σ:
2000
2σ:
1000
2σ:
0
0
2.33
0.23
3.87
1.05
2.08
0.18
0.443
0.025
1.36
0.06
0.987
0.054
0.185
0.029
0.0660
0.0075
0.0728
0.0084
0.0735
0.0062
0.146
0.023
0.108
0.016
xT
Deff,S
× 107
(m/s)
ks,S
× 102
(m4/m/s)
mean
model
errora
0.951
0.028
0.939
0.031
0.974
0.029
0.802
0.018
0.794
0.011
0.861
0.015
9.80
2.10
5.35
1.63
2.92
0.86
1.52
∞b
0.405
0.235
2.97
1.14
5.87
2.90
4.47
7.22
15.4
∞b
1.43
0.42
2.40
0.11
1.93
0.29
4.08
4.91
5.03
3.38
2.53
2.93
a The error reported is the square root of the average model error squared multiplied by 100. b ∞ means the parameter was determined
to be very large or had a very large confidence interval.
While the flux equation employed for SO2 absorption
clearly did not evolve from first principles, it is important to recall that the primary objective was to describe
the experimental data. With HCl and SO2 in the inlet
gas stream, eqs 11 (HCl) and 12 (SO2) indeed described
the data effectively. The results of the parameter
estimations for six experiments from the HCl-SO2hydrated lime system are given in Table 5. For each of
the experiments modeled, a unique set of regressed
parameters were determined.
The HCl and SO2 flux equations do not explicitly
account for the effects of O2 or NO2. Instead, the values
for the regressed parameters were allowed to vary at
different O2 and NO2 concentrations. Consequently, the
parameters themselves account for different degrees of
S(IV) oxidation.
There were large confidence intervals for some SO2
flux parameters due to the fact that only a small amount
of SO2 absorption took place over a rather narrow range
of experimental time. Consequently, there were not
enough data in some cases to accurately estimate
parameters. Values for Deff,H and ks,H were approximately equal to the same values estimated in the
absence of SO2. This was expected because the addition
of SO2 to the gas did not have a strong influence on HCl
absorption. The values for Deff,S and ks,S were determined to be at least 2 orders of magnitude less than
Deff,H and ks,H. Because HCl reacted considerably more
than SO2, it was expected that the SO2 parameters
would be less than the HCl parameters. Deff,S, like SO2
loading, generally decreased in the absence of NO2 or
O2 or in the presence of excess HCl. Figure 13 shows a
typical fit of the HCl and SO2 data.
5.3.2. Bag Filter Predictions. The bag filter performance calculations predicted what was qualitatively
seen in the fixed-bed experiments: HCl penetration was
low while SO2 penetration was high. Projections of HCl
and SO2 penetration under two sets of conditions that
allowed for improved SO2 absorption are illustrated in
Figure 14. On one hand, high levels of HCl removal were
achieved, regardless of NO2 and O2 concentration. On
the other hand, almost 70% of the inlet SO2 penetrated
through the bag filter at 60-min cycle time, even in the
presence of 150 ppm NO2 and 2.5% O2.
HCl and SO2 penetration as a function of sorbent
utilization is given in Figure 15. To reduce HCl penetration to 10%, hydrated lime utilization was approximately 40%. At this utilization, 85% of the SO2 fed
penetrated through the bag filter. The conditions for the
Figure 13. Comparison of model fit to the data for experiment
24: 19% RH, 250 ppm HCl, 1000 ppm SO2, 150 ppm NO2, 35 mg
of hydrated lime, 120 °C, and 1.5 SLPM.
Figure 14. Penetration of HCl and SO2 with varying NO2 and
O2 concentrations in a bag filter system. Performance predictions
with a sorbent feed ratio of 1.5 at 19% RH and 120 °C with 250
ppm HCl and 1000 ppm SO2. Parameters used are taken from
Table 5 for experiments 20 and 24.
prediction depicted in Figure 15 delivered the greatest
SO2 loading in the fixed-bed experiments. So the projection given in the figure below represents a best case
scenario for SO2 absorption with hydrated lime. The bag
filter model assumed that the sorbent undergoes only
one pass. It is possible that an analysis that includes
sorbent recycle may make the use of hydrated lime more
attractive for both HCl and SO2. In the absence of any
sorbent recycle, the small measure of SO2 reactivity does
not make hydrated lime an attractive option for acid
gas emissions control from municipal waste combustors
or coal-fired boilers.
Ind. Eng. Chem. Res., Vol. 38, No. 10, 1999 4079
Literature Cited
Figure 15. HCl and SO2 penetration as a function of cycle time
in a bag filter system. Performance prediction at 19% RH, 120
°C, 250 ppm HCl, 1000 ppm SO2, 150 ppm NO2, and 2.5% O2.
Parameters used are taken from Table 5 for experiment 24.
Conclusions
In the single-gas system, the reactivity of HCl with
hydrated lime (HL) increased considerably with increasing relative humidity. At 1000 ppm HCl and below, the
rate of absorption was first-order with respect to HCl
concentration and the final loading was constant. Greater
than 1000 ppm HCl, increasing concentration decreased
the final loading. The decrease in final loading may have
been due to pore plugging near the external particle
surface.
When HCl and SO2 were simultaneously exposed to
HL, the final total loading of the solids did not depend
on the SO2/HCl feed ratio. However, as the SO2/HCl
ratio increased, the fraction of the converted solids that
consisted of CaSO3‚1/2H2O increased while the fraction
of CaCl2‚ 2H2O decreased slightly. This phenomenon
suggested that there was a pseudoequilibrium between
the reaction products and the acid gases. The addition
of oxygen to the HCl-SO2 system enhanced SO2 absorption but did not affect HCl reactivity.
On one hand, the presence of NO in the HCl-SO2O2 system did not affect the final loading by HCl or SO2.
On the other hand, the presence of NO2 enhanced SO2
absorption. While the presence of O2 improved SO2
reactivity, the presence of O2 caused a decrease in NO2
loading. Increasing SO2 and NO2 reactivity had a minor
negative effect on the absorption of HCl. The conditions
that maximize sorbent utilization are high relative
humidity, SO2 concentration, NO2 concentration, and
oxygen concentration and low HCl concentration.
A model was developed to describe the experimental
data. Semiempirical flux equations for HCl and SO2
were developed that fit the data well. Parameters
defined by the flux equations were estimated. An
expression to correct for transient effects due to absorption of HCl in the gas cell was included. Regressed
parameters were used to predict the penetration of HCl
and SO2 in a bag filter system.
Injection of hydrated lime was projected to be an
effective sorbent for HCl absorption only at high relative
humidity. Almost 10% HCl penetration was predicted
for a sorbent utilization of 50% at 19% RH. Low levels
of HCl and high levels of SO2 penetration
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lime. At conditions that optimize SO2 reactivity but do
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ppm NO2), predictions for penetration at 50% sorbent
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Received for review October 16, 1998
Revised manuscript received July 7, 1999
Accepted July 21, 1999
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