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JOURNAL OF GEOPHYSICAL RESEARCH, VOL. 112, D18301, doi:10.1029/2006JD008207, 2007
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Heterogeneous oxidation of sulfur dioxide by ozone
on the surface of sodium chloride and its mixtures
with other components
L. Li,1 Z. M. Chen,1 Y. H. Zhang,1 T. Zhu,1 S. Li,1 H. J. Li,1 L. H. Zhu,1
and B. Y. Xu1
Received 31 October 2006; revised 4 May 2007; accepted 30 May 2007; published 18 September 2007.
[1] The heterogeneous oxidation of SO2 by O3 on NaCl particles has been studied using
diffuse reflectance infrared Fourier transform spectroscopy. The formation of sulfite and
sulfate on the surface was identified, and the roles of O3 and water in the oxidation processes
were determined. The results showed that in the presence of O3, SO2 could be oxidized to
sulfate on the surface of NaCl particles. The reaction is first order in O3 and zero order in
SO2. The initial reactive uptake coefficient for SO2 [(0.6–9.8) 1014 molecule cm3]
oxidation by O3 [(1.2–12) 1014 molecule cm3] was determined to be (4.8–0.7) 108
using the Brunauer-Emmett-Teller area as the reactive area and (9.8–1.4) 105 using
the geometric area at 40% relative humidity. A three-stage mechanism that involves
the adsorption of O3 results in an alkalescent surface, the adsorption of SO2
followed by O3 oxidation is proposed, and the adsorption of O3 on the NaCl surface is
the rate-determining step. The proposed mechanism can well explain the experiment
results. Furthermore, the surface oxidation on mixtures of NaCl with other components
such as CaCO3, Al2O3, TiO2, MgCl2 6H2O, MgO, elemental carbon, and soot
were studied. The reactivity of mixtures can be predicted from the reactivity of the single
component with each component weighted by its abundance in the mixture. The catalytic
and basic additives could enhance the production of sulfate on the NaCl surface.
Citation: Li, L., Z. M. Chen, Y. H. Zhang, T. Zhu, S. Li, H. J. Li, L. H. Zhu, and B. Y. Xu (2007), Heterogeneous oxidation of sulfur
dioxide by ozone on the surface of sodium chloride and its mixtures with other components, J. Geophys. Res., 112, D18301,
doi:10.1029/2006JD008207.
1. Introduction
[2] The sea-salt aerosol represents one of the largest
natural mass fraction of the global aerosol with an estimated
global atmospheric burden of 30 to 100 Tg [Graedel and
Keene, 1995; Intergovernmental Panel on Climate Change,
2001; Jacob, 2000]. The impact of sea-salt particles on the
Earth’s atmosphere is manifold. They are believed to have
direct and indirect effects on the radiation budget of the
atmosphere and therefore are expected to have impact on
climate [Cziczo et al., 2004; Tegen and Lacis, 1996]. In
addition, the surfaces of sea-salt particles provide sites for
heterogeneous reactions, and the potential role of the
heterogeneous reactions has been explored in several modeling studies [Bauer et al., 2004; Dentener et al., 1996;
Phadnis and Carmichael, 2000; Zhang and Carmichael,
1999].
[3] The sea-salt particles are formed by bursting bubbles
and generally have the composition of seawater [Harvey,
1
State Key Joint Laboratory of Environment Simulation and Pollution
Control, College of Environmental Sciences, Peking University, Beijing,
China.
Copyright 2007 by the American Geophysical Union.
0148-0227/07/2006JD008207$09.00
1928; Blanchard, 1985; Shaw, 1991]; however, they are
always enriched in sulfates from the sea surface. The sulfate
accounts for 0.01– 0.46% of the total soils in the samples of
arid area [Nishikawa et al., 1991]. However, in the particle
samples collected in Qingdao, a coastal city in China, the
sulfate accounted for about 0.7% of the TSP [Hu et al.,
2002]. Parungo et al. [1995] found that 50% – 80% of
coarse particles were coated with sulfates. Zhang et al.
[2000] and Zhang and Iwasaka [2001] also found the
sulfate-contained compounds, e.g., CaSO4, enriched in
particles collected in Qingdao. These studies implied that
the heterogeneous oxidation of SO2 probably is the main
source of sulfate on coarse particles [Ravishankara, 1997;
Ravishankara and Longfellow, 1999].
[4] It is well known that sulfate particles play a key role
in the global climate by participating as cloud condensation
nuclei and scattering solar radiation, thereby having a
cooling effect on the atmosphere [Dentener et al., 1996;
Charlson et al., 1990; Seinfeld and Pandis, 1998].
[5] Although a number of studies have been done on the
aqueous oxidation of SO2 [Chameides and Stelson, 1992;
Sievering et al., 1991; Knipping et al., 1995; Song and
Carmichael, 2001], there is little knowledge about the
heterogeneous oxidation of SO2 on sea-salt particles. Gebel
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et al. [2000] reported the uptake process of SO2 on the
surface of synthetic sea salt (SSS) and some components of
SSS using a Knudsen cell. They drew a conclusion that the
uptake was not apparent on the dry surface of NaCl, SSS
and MgCl2 6H2O, but was apparent on wet SSS and
MgCl2 6H2O particles. Laskin et al. [2003] studied the
physical and chemical changes on NaCl surface by means
of scanning electron microscopy with energy-dispersed
analysis of X rays (SEM-EDX) and time-of-flight secondary
ion mass spectrometry (TOF-SIMS). They found that the
increase in alkalinity would lead to an increase in the uptake
and oxidation of SO2 to sulfate on sea-salt particles. Up to
now, the mechanisms of sulfate formation on sea-salt particles are not completely elucidated [Dentener et al., 1996;
Ding and Zhu, 2003; Finlayson-Pitts and Hemminger, 2000;
Alexander et al., 2005].
[6] Although previous studies suggest an impact of the
heterogeneous reaction of SO2 with sea-salt particles on the
photo-oxidant budget of the atmosphere, the laboratory
studies are not yet available to quantify this process and
clarify its detailed mechanism [Rossi, 2003]. In the present
study, the uptake of SO2 on NaCl and its mixtures with
other components has been investigated at 293K and 1 atm
synthetic air using a diffuse reflectance infrared Fourier
transform spectroscopy (DRIFTS) reactor. From the investigation, uptake coefficients for SO2 and O3 as well as
condensed-phase products could be determined.
2. Experimental Setup
[7] The DRIFTS reactor used in this study has been
described in detail before [Li et al., 2006]. Briefly, A FTIR
Spectrometer (Nicolet Nexus) equipped with a mercury
cadmium telluride (MCT) detector and DRIFTS optics
(Model DRA-2CO, Harrick Scientific corp.) was used to
record in situ infrared spectra, ranging from 4000 to
600 cm1. The spectra were recorded at a resolution of
4 cm1, and 128 scans were usually averaged for each
spectrum corresponding to a time resolution of 2 min. The
reactant flow was forced to pass through the NaCl powder
in the reactor. The sample powders were put into the reactor,
and were flushed with nitrogen at room temperature before
an experiment began. A background spectrum of the initial
sample in the reactor was recorded before the gaseous
reactants were introduced. All the spectra for the reaction
were collected on this background. SO2 and O3 with the
specific concentration in the synthetic air were then introduced as gaseous reactants.
[8] A typical experiment lasted 180 min. The spectra
were collected at a 2-min interval during the initial 20 min
of the reaction; subsequently, the spectra were collected
every 10 or 20 min, according to the concentrations of the
reactants in the experiment.
[9] The experimental conditions were controlled at room
temperature and ordinary pressure in synthetic air with
change in relative humidity.
[10] The sodium chloride powders with high purity
(99.999%) were purchased from the commercial source
(Alfa Aesar). NaCl powders for the experiments were
prepared by grinding the commercial sample of NaCl,
and the powders were kept in a desiccator. The prepared
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powders were enough for all the experiments. The diameter distribution of the ground NaCl powder was determined by a Laser Sizer (MasterSizer 2000, Malvern
Instruments, U.K.) [Lobo et al., 2002; Yang et al., 2002]
as around 30 mm. The Brunauer-Emmett-Teller (BET)
surface area is usually used to represent the surface area
of particles [Brunauer et al., 1938; Knowles and Hudson,
1995]. The BET surface area of the NaCl powders was
determined by a measurement instrument of specific
surface area and pore size distribution (ASAP2010, Micromeritics, USA) as 0.81 m2 g1.
[11] In order to obtain the mixtures of NaCl with additives, other particles such as CaCO3, Al2O3, TiO2, MgCl2 6H2O, MgO, and elemental carbon were used as purchased.
The mixtures of NaCl with other components were made by
mechanically mixing and grinding. In the NaCl mixtures,
the mass percent concentration of the additive was 0.25 ±
0.02% by weight for elemental carbon and soot respectively,
but 5 ± 0.2% by weight for other additives respectively.
Soot samples were freshly collected on a SiO2 disc from the
flame by burning n-hexane.
[12] The gases of SO2 (National Research Center for
Certified Reference Material of China, 2000 ppm), N2
(Air products, 99.999%) and O2 (Air products, 99.999%)
were used in the experiments.
3. Results and Discussion
[13] In order to study the mechanisms of SO2 surface
oxidation by O3, the following aspects need to be investigated. First, the sequence of reaction steps should be
identified. Second, the kinetics parameters, for example,
uptake coefficients, should be determined. Then, the mechanism could be deduced finally.
3.1. Reactions on Pure NaCl Surface
[14] In order to explore the heterogeneous reactivity of
sea-salt particles, pure NaCl particles were used to study the
kinetics and mechanisms of the reaction with SO2 and O3 on
their surface at first.
3.1.1. Sequential Exposure of SO2 and O3
[ 15 ] Several sequential exposure experiments were
designed to probe the details of SO2 reaction with O3 on
the surface of NaCl.
[16] In the first experiment,when SO2 alone was introduced to the flow system in dry synthetic air, no apparent
bands were observed on surfaces of the surface of NaCl
(Figure 1 curve b). However, when SO2 was introduced
with 40% relative humidity (RH) in synthetic air, besides
the wide bands of adsorptive water at 3500 and 1640 cm1,
four bands at 1263, 1065, 903 and 801 cm1 were observed
(Figure 1 curve c). After the sample was evacuated, the
wide bands at 3500 and 1640 cm1 decreased, but other
bands kept stable (Figure 1 curve d). Dai et al. [1997] found
a uniform layer of water was formed above 35% RH, and
further more, Ghosal et al. [2005] found that there is a quasi
liquid layer on NaCl surface when RH is less than 40%, and
Br ions can affect the solvation and segregation properties
of NaCl. The purity of NaCl powder in this study is very
high, i.e., 99.999%. The purity of the commercial NaCl
sample used in this study was determined by ion chromatography (DIONEX, 2650, USA). The result shows that the
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Figure 1. In situ DRIFTS spectra when SO2 was exposed on NaCl surface. Product spectra before
addition of SO2 (curve a), 3 hours after addition of SO2 at dry condition (curve b), 3 hours after addition
of SO2 at 40% RH (curve c), and after evacuation of sample (curve d)). Spectra are shifted by +0.015,
+0.03 and +0.07 absorption units, for curves b, c, and d, respectively.
molar ratio of Br/Cl is at an upper limit of 0.0015, thus
the Br impurity content of the NaCl sample used in this
study is acceptable. Berg et al. [1996] found SO2 adsorbed
on NaCl(100) surfaces as monolayer. Thus, in this experiment, SO2 might react with the adsorptive water in the quasi
liquid layer to form species containing S– O groups, resulting in the four IR bands that do not disappear after the
sample was evacuated, as shown in Figure 1. However, on
the dry surface of NaCl, no bands were observed because
there was no adsorptive water.
[17] The second experiment included two steps, i.e., SO2
was exposed to NaCl particles followed by O3 exposure. At
first, SO2 was introduced into the synthetic air with 40%
RH, and then four bands at 1263, 1065, 903 and 801cm1
were observed (Figure 2 curve b). After 60 min, the SO2
flow was stopped and then O3 was introduced to the
reaction system. A band at 1124cm1 appeared (Figure 2
curve c), which corresponded to the stretching vibration (u3)
of sulfate [Nakamoto, 1997]. These results showed that SO2
could react with the wet surface of NaCl, and the products
could be oxidized into sulfate by O3.
[18] In the third experiment, the pure NaCl was replaced
by a mixture of NaCl externally mixed with NaOH to get a
basic surface (with a NaOH concentration of about 1% by
weight). The exposing sequence was the same as in the first
experiment mentioned above. When SO2 was introduced in
the synthetic air with 40% RH, a band at 1000 cm1 was
observed (Figure 3 curve a), which corresponded to the
stretching vibration (u3) of sulfite [Nakamoto, 1997]. After
60 min, the SO2 flow was stopped and O3 was introduced to
the reaction system. Then, a band of sulfate at 1124 cm1
appeared, whereas the band of sulfite at 1000 cm1
decreased gradually (Figure 3 curve b), and finally this
band disappeared (Figure 3 curve c). These results showed
that on the basic surface, SO2 adsorbed could react with OH
groups and subsequently be oxidized into sulfate.
[19] In the fourth experiment, the exposure sequence was
changed. SO2 was exposed to NaCl particles following the
Figure 2. In situ DRIFTS spectra recorded during the
sequential exposure experiment of SO2 and O3 on surface of
NaCl. Product spectra before addition of SO2 (curve a), at
saturation of adsorption after addition of SO2 at 40% RH
(curve b), and after cutting off SO2 supply and the addition
of O3 for 30 min (curve c) are shown. Spectra are shifted by
+0.005 and +0.01 absorption units, for curves b and c,
respectively.
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Figure 3. In situ DRIFTS spectra recorded during the
sequential exposure experiment of SO2 and O3 on surface of
NaCl mixture with NaOH. Product spectra at saturation of
adsorption after addition of SO2 at 40% RH (curve a), after
cutting off SO2 supply and the addition of O3 for 30 min
(curve b), and the subtractive spectrum of curves a and b
(curve c) are shown. Spectra are shifted by +0.02 and +0.05
absorption units, for curves b and c, respectively.
exposure of O3. When O3 was introduced in the synthetic
air with 40% RH, a band at 1100 cm1 was observed
(Figure 4 curve b), which corresponded to the vibration of
adsorptive O3. After 60 min, the O3 flow was stopped and
SO2 was introduced to the reaction system. Figure 4 curve d
is the subtractive spectrum of curve b to curve c, and
indicates that few changes could be found between these
two spectra. The adsorption of O3 could generate surface
OH [Keene et al., 1990], this would result in an alkalescent surface. It is found that SO2 could adsorb to the
alkalescent surface, but without further supply of O3 in this
experiment, little sulfate could be generated on the surface.
[20] To further prove the effect of O3 on the oxidation, an
external mixture of Na2SO3 and NaCl (with a sulfite
concentration of about 12% by weight) was exposed to
O3 [4.92 1014 molecule cm3] in the synthetic air with
40% RH. The unexposed sample was used to collect the
background spectrum. The sulfate positive bands at 1124 and
620 cm1 (u3 and u4 vibrations) and the sulfite negative
band at 1000 cm1 (u3) were observed in the subsequent
reaction (Figure 5). This phenomenon suggested that
sulfite could be easily oxidized into sulfate by O3.
3.1.2. Simultaneous Exposure of SO2 and O3
[21] When NaCl particles were exposed simultaneously
to SO2 [4.92 1014 molecule cm3] and O3 [4.92 1014 molecule cm3] in synthetic air, the main product
detected by DRIFTS analysis was sulfate, and little sulfite
was observed. Figure 6 presents a time series of adsorption
spectra recorded during this reaction. The band at 1124 cm1,
which is the u3 vibrations of sulfate, increases with reaction
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Figure 4. In situ DRIFTS spectra recorded during the
sequential exposure experiment of O3 and SO2 on surface of
NaCl. Product spectra before addition of O3 (curve a),
saturation of adsorption after addition of O3 at 40% RH
(curve b), after cutting off O3 supply and the addition of
SO2 for 2 hours (curve c), and the subtractive spectrum of
curves c and b (curve d). Spectra are shifted by +0.01,
+0.015, and +0.025 absorption units for curves b, c, and d,
respectively.
Figure 5. In situ DRIFTS spectra of O3 oxidation of
Na2SO3 mixed into NaCl powder. The unexposed sample
was used to collect a background spectrum, and the sulfate
positive band at 1200 and 620 cm1 and the sulfite negative
band at 1000 cm1 changed with the reaction time.
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Figure 6. In situ DRIFTS spectra during the reaction of SO2 with O3 on the surface of NaCl with the
change in reaction time ([SO2] = 4.92 1014 molecule cm3 and [O3] = 4.92 1014 molecule cm3).
time. Few changes were observed on the bands of surface
adsorbed water, and this demonstrates the surface hygroscopy kept unchanged. After the exposure stopped, the
sample in the reactor was evacuated to 10 Pa pressure and
a spectrum was recorded, showing little change compared to
that before evacuation. This indicated that under a vacuum
condition, the sulfate would not be desorbed from the
surface of NaCl.
[22] In order to separate overlapping sulfate bands in the
region from 1300 to 950 cm1, a curve-fitting procedure
was employed [Börensen et al., 2000]. On the basis of the
Lorenz and Gaussian functions, three bands could be
deconvolved, indicating several different chemical forms
of sulfate (Figure 7a). The temporal behavior of the three
bands was then analyzed to study their possibly different
kinetics behavior. An example was plotted as Figure 7b,
showing the fitted bands as a function of the reaction time.
As seen from Figure 7b, all of the three bands grew at a
similar rate and showed a similar change trend with the
reaction time. All of the three bands grew fast at the
beginning of the reaction, and then the growing rate
slowed down as the reaction proceeded. This suggested
that the total integrated absorbance of the overlapping
bands could be used to derive the formation rate of sulfate,
thus the individual bands need not be separated when we
analyze the kinetics of the heterogeneous process in this
study.
3.1.3. Reaction Kinetics
[23] The amount of sulfate on the particulate sample was
determined in order to quantify the sulfate formation rate
d{SO2
4 }/dt in terms of the reactive uptake coefficient. The
reactive uptake coefficient, g, is defined as the rate of
product formation (d{SO2
4 }/dt) divided by the rate of
surface collisions per unit time (Z).
g¼
=dt
d SO2
4
Z
ð1Þ
1
Z ¼ cAsurface ½SO2 4
ð2Þ
sffiffiffiffiffiffiffiffiffiffiffiffiffi
8RT
c ¼
pMSO2
ð3Þ
where Z is the rate of collisions between SO2 and particles,
c is the mean molecular velocity of SO2, Asurface is the
effective sample surface, R is the gas constant, T is the
temperature and MSO2 is the molecular weight of SO2
[Baltensperger et al., 1996; Finlayson-Pitts and Pitts, 1999;
Molina and Molina, 1996]. Concentrations marked with
braces indicate surface species, whereas brackets indicate
the concentration of gas-phase species. Two extreme cases
of effective sample surface were considered for calculating
the uptake coefficient. If the reaction probability is high, the
reactants would have no time to diffuse into the sample
before reacting and the effective surface area will be the
geometric surface area of the sample. If the reaction
probability is low, the reactants may have enough time for
diffusion into the entire sample and thus the BET surface
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Figure 7. Deconvolution of overlapping bands using Lorenz and Gaussian curves. (a) An absorption
spectrum recorded in the reaction of SO2 and O3 with NaCl. (b) Normalized integrated absorbance of all
three bands as a function of reaction time. All three bands grew at similar rates and showed similar trends
with reaction time.
area would more appropriately represent the effective area.
If sulfate formed is evenly distributed into the sample, the
effective surface area might approach the BET surface area.
[24] The amount of sulfate ions formed during the reaction was determined by DRIFTS and proved by ion chromatography. Both the absorbance [Vogt and Finlayson-Pitts,
1994; Ullerstam et al., 2002] and Kubelka-Munk (K-M)
function [Tsai and Kuo, 2006; Averett and Griffiths, 2006]
integrated over the (u3) region can be used to quantify the
surface products. The K-M method is known to vary with
baseline position error that give rise to unacceptable uncertainty levels in quantitative experiments [Samuels et al.,
2006]. Hence the surface products of the SO2 reactions on
NaCl particles in this study were quantified using the
integrated absorbance.
[25] There is an overlap on the bands of sulfate and
absorbed ozone in the region of 1300 – 950 cm1 (Figures 4
and 6). However, the absorbance of ozone is much less than
that of sulfate, as shown in Figure 8. When SO2 together with
O3 was introduced to the surface of NaCl particles, the
integrated absorption of sulfate formed was over 10 times
higher than that of O3 absorbed when O3 was introduced
alone. So in the kinetic studies, we thought the bands in the
region of 1300– 950 cm1 should be attributed to sulfate
when SO2 and O3 are simultaneously introduced to the NaCl
surface.
[26] Two stages of every experiment were considered for
the kinetic study. For the initial stage, the reactive gas was
introduced into the fresh particle layers, the IR bands of
sulfates increased quickly, and the initial uptake coeffi-
cient, g 0, was determined. After a certain time, when the
IR bands of sulfates increased at a stable speed, namely,
the steady state stage, the steady state uptake coefficient,
g ss, was determined.
[27] Three sets of experiments were performed: in the
first experiment, the SO2 concentration was varied and the
O3 and water vapor concentrations were kept constant; in
Figure 8. Integrated absorption of the band between 1300
and 950 cm1 of two typical exposures of O3 alone (circles)
and SO2 + O3 (squares) introduced to the surface of NaCl
particles at 40% RH.
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Table 1. Experimental Concentrations and Reactive Uptake Coefficients for SO2 for the Reaction of SO2 and O3
on NaCl at 293 K and 1 atm Air With 40% RH
g SO2
g O3
[SO2],
1014 molecule cm3
[O3],
1014 molecule cm3
Geometric Area,
105
BET Area,
108
0.6 – 9.8
4.9
9.8 – 1.4 (g 0)
7.3 – 1.0 (g ss)
4.8 – 0.7 (g 0)
3.6 – 0.5 (g ss)
4.9
1.2 – 12
the second, the O3 concentration was varied and the SO2
and water vapor concentrations were kept constant; and in
the third, O3 and SO2 concentrations were kept constant and
water vapor concentration was varied.
[28] The BET surface area has been determined to be
0.81 (m2 g1) and the uptake coefficients calculated using
these data are listed in Table 1. The uptake coefficients
obtained using the geometric and BET surface area as the
reactive surface area should be respectively considered as
higher and lower limits. The initial uptake coefficient for
SO2 is (9.8 – 1.4) 105 if the geometric surface area is
used and (4.8 –0.7) 108 if the BET surface area is
applied; the steady state uptake coefficients for SO2 are
(7.3 –1.0) 105 and (3.6 – 0.5) 108, respectively.
[29] Ravishankara and Longfellow [1999] distinguished
between two kinds of solids: dynamic solids, where the
flux of the solid’s constituent through the surface is much
larger than that of the atmospheric reactant; and the second
kind are rigid solids, where the surface is not refreshed via
deposition or evaporation. The g ss is similar to the g 0 for
the SO2 on NaCl surface, thus NaCl can be considered as
a dynamic solid. As a comparison, CaCO3 considered as a
rigid solid in our previous work [Li et al., 2006].
Geometric Area,
105
BET Area,
108
8.8 ± 2.6 (g 0)
5.0 ± 2.2 (g ss)
4.3 ± 1.3 (g 0)
2.5 ± 1.1 (g ss)
[30] Equation (1) shows that in the case of a first-order
reaction, the uptake coefficient is independent of the reactant concentration, while the reaction rate is proportional to
the reactant concentration. In case of a zero-order reaction,
the reaction rate is concentration-independent.
[31] A double-logarithmic plot of the initial and steady
state formation rates of sulfate on NaCl particles versus the
concentration of SO2 (Figure 9a) gives a slope of 0.11 ±
0.01 (s) and 0.07 ± 0.07 (s), respectively, and these values
represent the reaction orders in SO2. The reaction orders,
which are close to zero, indicate that the sulfate formation
rate is independent of the SO2 concentration at both the
initial stage and steady state stage. Hence the uptake
coefficient was inversely dependent on SO2 concentration.
An equivalent plot of the initial and steady state formation
rate of sulfate versus the O3 concentration (Figure 9b) gives
a slope of 0.85 ± 0.16 (s) and 0.83 ± 0.19 (s), respectively.
The reaction order is close to 1 in the initial time stage and
steady state stage. This suggests that the uptake coefficients
are independent of O3 concentration in the concentration
range used in this study. The corresponding reactive uptake
coefficients are shown in Table 1. The results indicate that
the reaction order is 1 in O3 and zero in SO2 for the
Figure 9. Double-logarithmic plot of the rate of sulfate formation as a function of (a) [SO2] and (b) [O3]
at initial time and steady state. The reaction order in SO2 was determined from a linear regression
yielding n0 = 0.11 ± 0.01 (s) and nss = 0.07 ± 0.07 (s). For O3, corresponding values are n0 = 0.85 ± 0.16
(s) and nss = 0.83 ± 0.19 (s).
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Figure 10. Production rates of sulfate as a function of relative humidity in the reaction of SO2 and O3
on NaCl ([SO2] = 4.92 1014 molecule cm3 and [O3] = 4.92 1014 molecule cm3).
heterogeneous reaction of SO2 and O3 on NaCl particles,
suggesting that the supply of O3 determines the reaction
rate. The results also prove that in the reaction process, O3
adsorbed would consumed immediately and the bands of O3
should be very weak.
[32] A plot of the initial and steady state formation rates
of sulfate versus the relative humidity RH shows a different
regime (Figure 10). In the initial stage, the uptake coefficients increase with the increase in RH. In the steady state
regime, the uptake coefficient in 20% RH is twice as much
as that in 0% RH. However, when the RH rise above 20%
RH, the uptake coefficient did not increase any more.
3.1.4. Mechanism
[33] Because the gas-phase oxidation of SO2 by O3 can
be neglected in the experiment, the main path of SO2
oxidation is a heterogeneous process on the particulate
surface in this study. According to the results of the
sequential exposure experiments mentioned above, the
following conclusions could be drawn: (1) SO2 can be
absorbed on a basic surface and form sulfite with adsorbed
water. (2) O3 can be absorbed on the surface of NaCl to
generate an alkaline surface. (3) Sulfite can be easily
oxidized into sulfate by O3. Therefore we propose a reaction
mechanism as follows: SO2 oxidation by O3 on the surface
of NaCl particles proceeds mainly via three major stages. At
first, an adsorption of O3 to the NaCl surface occurs
k1
O3 þ H2 O þ 2Cl ! Cl2 þ 2OH þ O2
[35] The second stage is the adsorption of SO2 on the
surface of NaCl, then the adsorptive SO2 convert to sulfite
at the basic surface. Considerations of nucleophilic reactivity indicate that SO2
3 should react more rapidly with ozone
than HSO
3 , and HSO3 should react more rapidly in turn
than SO2 H2O [Hoffmann and Calvert, 1985]. The results
of sequential exposure of SO2 and O3 also indicates that
SO2 can easily absorbed on the alkaline surface, so the
reaction between the adsorbed SO2 and alkaline surface can
be expressed as following
SO2 ð gÞ
! SO2 ðadsÞ
k3
SO2 ðadsÞ þ 2OH ! SO2
3 ðadsÞ þ H2 O
ð5Þ
ð6Þ
[36] The third stage is the irreversible reaction, in which
the sulfite is oxidized to sulfate by O3. This reaction occurs
very quickly. The surface reaction can be expressed as
k4
2
SO2
3 ðadsÞ þ O3 ! SO4 ðadsÞ þ O2 ð g Þ
ð7Þ
[37] Thus the reaction rate of sulfate formation can be
described by a general equation:
ð4Þ
[34] Keene et al. [1990] proposed the above reaction of
ozone on NaCl surface to explain a measured loss of
chloride from sea-salt particles that was larger than the
increase in nitrate and non-sea-salt sulfate. The generated
OH raises the alkalinity of NaCl surface.
k2
k2
d SO2
4
r¼
½O3 ¼ k4 SO2
3
dt
ð8Þ
[38] Because the generated OH can react with SO2
rapidly, the concentration of OH can be considered constant in the reaction. According to the steady state approximation, the net rate of change in the intermediates may be
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LI ET AL.: HETEROGENEOUS OXIDATION OF SO2 ON NaCL
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Figure 11. Schematic diagram of the mechanism of SO2
oxidation by O3 on the NaCl surface. The curve above the
surface represents the quasi liquid layer on the NaCl
surface. The O3 reaction with the NaCl surface results in
the alkalescent surface. Gas-phase SO2 can be adsorbed on
the alkalescent surface. Adsorbed SO2 is transformed into
sulfite, and then the sulfite is rapidly oxidized into sulfate
by O3.
set equal to zero. This means the generation rate of OH is
equal to their consumption rate:
d fOH g
¼ 2k1 ½O3 2k3 fSO2 gfOH g2 ¼ 0
dt
ð9Þ
So
fOH g2 ¼
k1 ½O3 k3 fSO2 g
ð10Þ
[39] When NaCl was exposed to SO2 and O3 simultaneously, no observable adsorbed SO2 and sulfite was
observed. This phenomenon implies that the SO2 adsorbed
on the surface is quickly transformed into sulfite, then the
sulfite is immediately oxidized into sulfate. Also according
to the steady state approximation, for the intermediates in
the SO2 surface oxidation, i.e., surface SO2 and SO2
3 on the
surface, the generation rates are then equal to their consumption rate:
d fSO2 g
¼ k2 ½SO2 k2 fSO2 g k3 fSO2 gfOH g2 ¼ 0
dt
ð11Þ
d SO2
3
½O3 ¼ 0
¼ k3 fSO2 gfOH g2 k4 SO2
3
dt
ð12Þ
Thus
d SO2
4
r¼
¼ k1 ½O3 dt
ð13Þ
[40] From equation (13), it can be seen that the reaction is
first order for O3 and zero order for SO2. Therefore the
reaction orders deduced from the proposed mechanism are
consistent with the experimental results. The mechanism that
we propose was illustrated in Figure 11. It should be pointed
out that the above mechanism is deduced in confined experimental conditions.
3.2. Reactions on Mixtures of NaCl With Other
Components
[41] After pure NaCl particles were used as model for seasalt particles to study the kinetics and mechanisms of
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reaction with SO2 and O3, some mixtures of NaCl with
other components, including CaCO3, Al2O3, TiO2, MgCl2 6H2O, MgO, elemental carbon, and soot were used to probe
their impact on the heterogeneous reactions. MgCl2 6H2O
is an important component of sea salt. CaCO3 and Al2O3 are
important components of mineral dust. TiO2, as a component of rutile in mineral dust, is considered for its high
catalysis. MgO, selected as the base model of atmospheric
particles, is chosen as a comparison of MgCl2 6H2O. Soot
and elemental carbon represent the carbon components of
atmospheric particles. These substances are selected to
represent the major components of atmospheric particles.
[42] In the atmosphere, the particles usually contain a lot
of components. Different components have different chemical properties, thus results in difficulty of the study of
atmospheric heterogeneous chemistry [Al-Abadleh and
Grassian, 2003; Bizjak et al., 1999; DeHaan et al., 1999;
Hemminger, 1999].
3.2.1. Reactions on the Mixture of NaCl With CaCO3
[43] A previous work demonstrated that the heterogeneous uptake coefficient of a China loess sample could be
predicted from the fractional amount and reactivity of the
single component oxides and carbonates [Usher et al.,
2002]. CaCO3 is a good candidate for studying such a
reaction because the uptake coefficient of CaCO3 has been
determined in our previous work [Li et al., 2006]. The
mixture of NaCl and CaCO3 is a typical representative of
externally mixed particles.
[44] NaCl particles mixed with CaCO3 (with a CaCO3
concentration of about 5% by weight) were exposed
simultaneously to SO2 [4.92 1014 molecule cm3] and
O3 [4.92 1014 molecule cm3] at 40% RH, the main
product detected by DRIFTS analysis was sulfate. The loss
of CaCO3 was also observed. Figure 12 presents a time
series of adsorption spectra recorded during the reaction of
SO2 and O3 on NaCl particles mixed with CaCO3. The band
at 1148 cm1 increases with reaction time which is the u3
vibrations of sulfates. The double peaks at 670 and 602 cm1
correspond to the u4 vibrations of CaSO4 [Nakamoto,
1997]. The negative bands at 1433, 878 and 713 cm1
correspond to the u3, u2 and u4 vibrations of CaCO3
[Nakamoto, 1997]. The results showed that CaSO4 was
generated and CaCO3 was consumed.
[45] According to the integrated absorbance of the sulfate
band as a function of time in the reaction of SO2 and O3 on
the surface of NaCl particles mixed with CaCO3, the
production rate of sulfate can be calculated as 1.62 1015 ion s1 g1. In this study, the production rate of sulfate
on the surface of pure NaCl was determined as (0.9 ± 0.3) 1015 ion s1 g1. The production rate of sulfate on the
surface of pure CaCO3 is (1.1 ± 0.2) 1016 ion s1 g1
[Li et al., 2006]. If the uptake was additive, then it would be
expected that the mixture would have a reactivity that can
be calculated from the linear combination of the reactivity
of pure components. So the calculated production rate of
sulfate on the mixture of NaCl and CaCO3 is (1.4 ± 0.4) 1015 ion s1 g1, which is consistent with the result from
our experiment. This indicates that the reactivity of a
mixture can be predicted from the reactivity if the single
component with each component weighted by its abundance
in the sample. This result agree with the work on loess
reported by Usher et al. [2002].
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Figure 12. In situ DRIFTS spectra during the reaction of SO2 with O3 on the surface of NaCl externally
mixed with CaCO3 with the change of reaction time at 40% RH. The spectra of CaSO4 2H2O and CaCO3
are listed for comparison ([SO2] = 4.92 1014 molecule cm3 and [O3] = 4.92 1014 molecule cm3).
[46] It should be noticed that the above results were
obtained with the assumption that the samples have a
similar surface area. The BET areas of NaCl and CaCO3
particles used in this study were determined as 0.81 and
1.95 m2 g1, respectively. It can be said that they have a
similar specific surface area.
3.2.2. Comparison on Uptake Coefficients on Mixtures
of NaCl With Other Components
[47] When sea-salt particles are transported into an inland
area, they can lose water and take up other atmospheric
constituents, and they can also mixed with other particles
such as mineral dust [Tseng et al., 1992; Rossi, 2003]. The
additives could change the heterogeneous reactivity of seasalt particles [Hara et al., 2002, 2005]. In this study, the
uptake coefficients of NaCl mixtures with different addi-
tives were determined at 40% RH. Some interesting results
have been obtained, as shown in Figure 13.
[48] For the mixtures of NaCl with MgO and MgCl2
6H2O respectively, the former additive can greatly increase
the uptake coefficient, while the effect of the later is limited.
Although Mg2+ does not have a catalytic effect, it can
provide a acid-basic effect. MgO is a basic oxide, so the
additive of MgO can greatly increase the basic property of
the surface of NaCl. The basic surface can accelerate the
adsorption of SO2, resulting in the increase of uptake
coefficient. MgCl2 is an important component of sea salt,
but it does not have an acid-basic effect. So the uptake
coefficient of its mixture does not increase significantly.
Gebel et al. [2000] found in Knudsen cell experiments that
no measurable uptake was observed when MgCl2 6H2O
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Figure 13. Comparison of the change in uptake coefficients among different chemicals externally
mixed in NaCl particles at 40% RH. The Y axile is the logarithmic of the ratio of uptake coefficient of
mixtures with that of pure NaCl. The mass percent concentration of elemental carbon and soot was kept
at 0.25 ± 0.02%, and others were kept at 5 ± 0.2%.
was free of large amounts of water. This result is consistent
with that of our study.
[49] The initial and steady state uptake coefficients of
NaCl mixtures using Al2O3 and TiO2 as additives are much
larger than those of pure NaCl. Both Al2O3 and TiO2 are
usually considered as catalyst, and the results of uptake
coefficients determined in this study show that surface
catalysis reaction can greatly enhance the uptake of SO2.
[50] For MgO and CaCO3 as additives, the initial and
steady state uptake coefficients are larger than that from
pure NaCl. Both MgO and CaCO3 are basic chemicals,
which can enhance the uptake of SO2 on the surface. MgO
has a stronger basic property than CaCO3, thus the former
has a greater uptake ability.
[51] For the elemental carbon and soot, they showed
different reactivities. Both the initial and steady state uptake
coefficients for elemental carbon as additive are less than
those of pure NaCl. This is because of the competitive
adsorption of element carbon. For soot as additive, the
initial uptake coefficient is larger than that of pure NaCl.
This result agrees with other works [Akhter et al., 1985;
Ammann et al., 1998]. As a representative of atmospheric
soot, the reactivity of the soot generated by combustion is
higher than that of the elemental carbon. As the surface of
soot has a lot of active functional groups, it can take part in
the atmospheric heterogeneous reaction. However, soot is a
typical rigid solid, the reactivity would decrease largely
after the surface functional groups have been consumed.
This is the reason why its steady state uptake coefficient is
lower than that of pure NaCl. In summary, the addition of
catalytic and basic chemicals can enhance the production of
sulfate on NaCl particles.
4. Conclusions
[52] The kinetics and the mechanism of SO2 oxidation by
O3 on NaCl particles were studied using the DRIFTS
technique. The initial and steady state reactive uptake
coefficients were determined. The initial reactive uptake
coefficient was determined to be (9.8 –1.4) 105 when
the reactive surface area was assumed to the geometric
surface area, or (4.8 – 0.7) 108 when a BET surface area
was assumed. The steady state uptake coefficient is lower
by 1 order of magnitude than the initial coefficient.
[53] The reaction mechanism was studied by sequential
exposure experiments. A three-stage mechanism is proposed, involving the absorption of O3 on the NaCl surface
to make a basic surface, the absorption of SO2 on basic
surface to form SO2
3 in the presence of water, followed by
O3 oxidation to form sulfate. The adsorption of O3 is the
rate-determining step. The reaction was determined to be
zero order in SO2 and first order in O3 within the experimental concentration range. The reaction order deduced
from the proposed mechanism is in good agreement with
the experimental result.
[54] The reactivity of mixture can be predicted from the
reactivity of the single component with each component
weighted by its abundance in the mixture. The addition of
catalytic and basic chemicals can enhance the production of
sulfate on NaCl particles.
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[55] Acknowledgments. The authors wish to express thanks for financial support to the Project of Development Plan of the State Key Fundamental
Research of China (grants 2002CB410802 and 2002CB410801) and the
National Natural Science Foundation of China (grants 20107001 and
20077001). The authors thank Jie Chongyu for his determining bromine
content of NaCl sample.
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