Photoenhanced sulfate formation by the heterogeneous uptake of SO₂ on non-photoactive mineral dust

Heterogeneous uptake of SO₂ on mineral dust is a predominant formation pathway of sulfates, whereas the contribution of photo-induced SO₂ oxidation to sulfates on the dust interfaces still remains unclear. Here, we investigated heterogeneous photochemical reactions of SO₂ on five mineral oxides (SiO₂, kaolinite, Al₂O₃, MgO, and CaO) without photocatalytic activity. Light enhanced the uptake of SO₂, and its enhancement effects negatively depended on the basicity of mineral oxides. The initial uptake coefficient (γ_0,BET) and the steady-state uptake coefficient (γ_s,BET) of SO₂ positively relied on light intensity, relative humidity (RH), and O₂ content, while they exhibited a negative relationship with the initial SO₂ concentration. Rapid sulfate formation during photo-induced heterogeneous reactions of SO₂ with all mineral oxides was confirmed to be ubiquitous, and H₂O and O₂ played key roles in the conversion of SO₂ to sulfates. In particular, triplet states of SO₂ (³SO₂) were suggested to be the trigger for photochemical sulfate formation. Atmospheric implications supported a potential contribution of interfacial SO₂ photochemistry on non-photoactive mineral dust to atmospheric sulfate sources.

1 Introduction

As an important trace gas in the atmosphere, SO₂ is mainly emitted by volcanic eruption and fuel combustion. There is an uneven distribution of atmospheric SO₂ concentrations that show a distinctive seasonal and regional differentiation. Typical mixing ratios of SO₂ in the troposphere are below 0.5 ppb for clean-weather days, rising to several hundred parts per billion during polluted days in urban regions (Ma et al., 2020). About half of SO₂ is oxidized to sulfates (He et al., 2012), which are key components of fine particulates in the atmosphere. High sulfate loading in PM_2.5 was observed (Shao et al., 2019), especially in polluted regions where high-sulfur fuels are usually used (Olson et al., 2021). They significantly alter physicochemical properties of aerosols in terms of hygroscopicity, acidity, and light absorption (Chan and Chan, 2003; Cao et al., 2013; Lim et al., 2018). Sulfates may lead to negative health outcomes, such as respiratory illnesses and cardiovascular diseases (Shiraiwa et al., 2017). In addition, the deposition of sulfates leads to adverse effects on ecosystems via the acidification of soils and lakes (Golobokova et al., 2020). Therefore, the oxidation of SO₂ to form sulfates has attracted widespread attention in the past decades.

The conversion of SO₂ to sulfates in the atmosphere usually occurs in different phases: gas-phase oxidation of SO₂ by hydroxyl radicals (•OH) or Criegee intermediate radicals (Mauldin et al., 2012; Davis et al., 1979); aqueous-phase reactions of SO₂ with O₃, peroxides, or transition metal ions dissolved in cloud and fog droplets (Alexander et al., 2009; Herrmann et al., 2000; Harris et al., 2013; Liu et al., 2020; Li et al., 2020); and heterogeneous SO₂ uptake on aerosols, including authentic mineral dust, soot, inorganic ion, and organic compounds (Adams et al., 2005; G. He et al., 2018; Ye et al., 2018; S. Wang et al., 2019; Yao et al., 2019; R. Zhang et al., 2020; Liu et al., 2020, 2021). However, the oxidation of SO₂ in gas and aqueous phases fails to explain high sulfate concentrations in polluted areas. Model simulation suggests that the rapid sulfate formation can be attributed to heterogeneous SO₂ uptake (Li et al., 2017). A positive relationship between the fraction of sulfates and mineral dust in haze days has been reported, implying that mineral dust may account for the formation of sulfates (J. Wang et al., 2020). Moreover, large numbers of sulfates were observed on the surface of mineral dust after long-distance transport (Prospero, 1999). Thus, investigating the heterogeneous oxidation of SO₂ on mineral dust can provide basic data for the model calculation to evaluate atmospheric sulfates.

Mineral dust, as the dominant component of particulate matters in the atmosphere, accounts for about 30 %–60 % mass fractions of global aerosols (Dentener et al., 1996; Peng et al., 2012). It primarily contains SiO₂ (40 %–80 %), followed by Al₂O₃ (10 %–15 %), Fe₂O₃ (6 %–13 %), CaO (3 %–10 %), MgO (1 %–7 %), and TiO₂ (0.1 %–5 %) (Urupina et al., 2019, 2021; Usher et al., 2003). Mineral dust can provide active sites for the adsorption and reaction of gases. Up to now, the heterogeneous SO₂ uptake on authentic mineral aerosols and model mineral oxides has been widely reported (Ma et al., 2019; Goodman et al., 2001; T. Wang et al., 2018, 2020), with the uptake coefficient (γ) of SO₂ varying from 10⁻⁹ to 10⁻⁴ (Urupina et al., 2019; Usher et al., 2002).

It was recognized that light could significantly enhance the heterogeneous conversion of SO₂ to sulfates on the surface of photocatalytic mineral dust (Chen et al., 2021; Li et al., 2019; T. Wang et al., 2020). Electron–hole pairs are produced via photo-induced electrons from the valence band to the conduction band of photocatalytic metal oxides and then react with H₂O and O₂ to generate reactive oxygen species (ROS) such as •OH and $\mathrm{•}{\mathrm{O}}_{\mathrm{2}}^{-}$ (Chu et al., 2019). Sulfates are produced by the heterogeneous reactions of SO₂ with ROS (Park and Jang, 2016; Park et al., 2017; Langhammer et al., 2020; Bounechada et al., 2017). In particular, due to the large abundance of non-photoactive mineral dust (more than 85 % mass of total mineral dust in the atmosphere) (Usher et al., 2003; Liu et al., 2022), revealing the photooxidation processes of SO₂ on these mineral dust is of great importance to better reevaluate the sulfate formation on aerosols at a global scale.

Photochemical SO₂ uptake and sulfate formation on non-photoactive mineral oxides were investigated using a flow reactor and in situ diffuse reflectance infrared Fourier transform spectroscopy (DRIFTS). The conversion of SO₂ to sulfates was examined under various conditions, and the roles of light intensity, SO₂ concentration, H₂O, O₂, and basicity of mineral oxides were determined. Reaction mechanisms and atmospheric implications were proposed, highlighting an important pathway accounting for the photochemical uptake of SO₂ to form sulfates on the non-photoactive surfaces.

2 Experimental methods

2.1 Materials

Analytical-grade SiO₂ (Sinopharm Chemical Reagent Co., Ltd.), kaolinite (Sinopharm Chemical Reagent Co., Ltd.), Al₂O₃ (Alfa Aesar), MgO (Sigma-Aldrich), and CaO (Sigma-Aldrich) were used in the experiments. Through the nitrogen Brunauer–Emmett–Teller (BET) physisorption analysis, their specific surface areas were detected to be 0.419, 6.407, 8.137, 10.948, and 6.944 m² g⁻¹, respectively. With BaSO₄ used as the reference, the ultraviolet–visible (UV-vis) light absorption spectra of samples (Fig. S1 in the Supplement) in the wavelength range of 300–800 nm were obtained by the Shimadzu UV-2550 spectrophotometer, which was equipped with a diffuse-reflection integrating sphere attachment. The solid powder (0.2–5 g) was uniformly dispersed into 10.0 mL ethanol solution. The mixed liquid was poured into a rectangular quartz sample dish (14.0 cm × 7.0 cm) and dried to form a solid coating in an oven at 353 K for 10 h. SO₂ standard gas (50 ppm in N₂, Shenyang Air Liquide Co., Ltd.) and high-purity N₂ and O₂ (99.999 vol.%, Shenyang Air Liquide Co., Ltd.) were used as received. The solid sample powder (0.2 g) was immersed in 10 mL deionized water (20 mg mL⁻¹), and then the suspension was vigorously stirred for 10 min. The pH of SiO₂, kaolinite, Al₂O₃, MgO, and CaO suspension was measured to be 6.27, 6.58, 9.33, 10.61, and 12.72, respectively, using a pH meter.

2.2 Rectangular flow reactor

The uptake experiments of SO₂ on mineral dust were performed in a horizontal rectangular flow reactor (26.0 cm length × 7.5 cm width × 2.0 cm height), which is depicted in Fig. S2. In a previous study, a similar rectangular flow reactor was designed, and the feasibility of the reactor has been verified (Knopf et al., 2007). The reactor was made of quartz to allow the transmission of light. The temperature was maintained at 298 K by circulating temperature-controlled water through the outer jacket of the reactor. Synthetic air with a ${\mathrm{N}}_{\mathrm{2}}/{\mathrm{O}}_{\mathrm{2}}$ volume ratio of 4:1 was introduced into the flow reactor, and its total flow rate was 1000 mL min⁻¹. The Reynolds number (Re) was calculated to be 28.2 (Re < 200), as described in the Supplement, indicating a laminar flow state. SO₂, together with high-purity N₂ (100 mL min⁻¹), was introduced into the reactor through a movable T-shaped injector equipped with six exit holes (0.2 mm diameter) so that the gas could be uniformly distributed over the width of the reactor. The SO₂ concentration was 40–200 ppb and measured with a SO₂ analyzer (Thermo 43i). Wet N₂ generated with a bubbler containing deionized water was introduced by two parallel inlets on the side of a T-shaped injector. Relative humidity (RH, 10 %–75 %) was controlled by regulating the ratio of dry N₂ to wet N₂ and measured via a hygrometer (Center 314). The equivalent layer numbers of water on the surface were 0.9–4.0, according to the Brunauer–Emmett–Teller (BET) model (Sumner et al., 2004), and the thickness of the film of adsorbed water varied between 2.7–12 nm at RH = 10 %–75 %. There were three equally spaced exhaust ports to mitigate the outlet turbulence. A xenon lamp (CEL-LAX500, China Education Au-light Co., Ltd.) was used to simulate sunlight and vertically located above the reactor. A filter was placed on the reactor to remove the light with wavelengths shorter than 300 nm. The spectrum irradiance of the xenon lamp is displayed in Fig. S3 and was measured using a calibrated spectroradiometer (ULS2048CL-EVO, Avantes). The spectral irradiance was measured inside the reactor after passing the water cooling and in the absence of a sample. The total irradiance (0–7.93 × 10¹⁶ $\mathrm{photons}{\mathrm{cm}}^{-\mathrm{2}}{\mathrm{s}}^{-\mathrm{1}}$) on the coating can be adjusted by varying the distance of the xenon lamp to the reactor.

2.3 Uptake coefficient of SO₂

The heterogeneous reaction kinetics of SO₂ with mineral dust can be described by a pseudo-first-order reaction. Figure S4 shows a linear relationship between the natural logarithm of the SO₂ concentration and the reaction time. The apparent rate constant ($k_{\text{obs},{\mathrm{SiO}}_{\mathrm{2}}}$) of SO₂ with SiO₂ can be calculated using Eq. (1),

$$\begin{array}{}\text{(1)}& {\displaystyle \frac{\ln (C_{\mathrm{0}}/C_{\mathrm{t}})}{t}}=k_{\text{obs},{\mathrm{SiO}}_{\mathrm{2}}},\end{array}$$

where C₀ and C_t (ppb) are the initial SO₂ concentration and the SO₂ concentration at the exposure time t, respectively; t was calculated by diving the length of the reactive surface by the average flow velocity. The loss of SO₂ on the internal wall of the reactor in blank experiments was carried out under various conditions (Fig. S5 as an example), and it has been deducted for the γ calculation. Assuming that the wall loss was constant in the experiments with and without samples, the geometric uptake coefficient (γ_geo) was determined by Eq. (2) (Knopf et al., 2007),

$$\begin{array}{}\text{(2)}& {\mathit{\gamma }}_{\text{geo}}={\displaystyle \frac{\mathrm{4}Vk}{S\mathit{\omega }}},\end{array}$$

where k (s⁻¹), V (4 × 10⁻⁴ m³), S (9.8 × 10⁻³ m²), and ω (314.05 m s⁻¹) are the reaction rate constant, the volume of the rectangular reactor, the surface area of the sample dish, and the mean molecular speed of SO₂, respectively.

The uptake process of SO₂ on SiO₂ depended on the reaction of SO₂ with SiO₂ and the mass transport of SO₂ to the surface. It can be expressed by Eq. (3),

$$\begin{array}{}\text{(3)}& k_{\mathrm{r},{\mathrm{SiO}}_{\mathrm{2}}}^{\prime }={\left[{\displaystyle \frac{\mathrm{1}}{k_{\text{obs},{\mathrm{SiO}}_{\mathrm{2}}}-k_{\text{obs},\text{wall}}}}-{\displaystyle \frac{a}{\mathit{Nu}D}}\right]}^{-\mathrm{1}},\end{array}$$

where $k_{\text{obs},{\mathrm{SiO}}_{\mathrm{2}}}$ and k_obs,wall (s⁻¹) are the apparent rate constants measured with and without SiO₂ samples, respectively. $k_{\mathrm{r},{\mathrm{SiO}}_{\mathrm{2}}}^{\prime }$ is the reaction rate constant of SO₂ accounting for the diffusion effect; D (0.1337 cm² s⁻¹) is the diffusion coefficient of SO₂ in air; a (1 cm) is one-half the height of the flow reactor; and Nu represents the Nusselt numbers obtained with a calculation method from Solbrig and Gidaspow (1967), which represents the mass transport. The corrected γ can be calculated by Eq. (2), where k was replaced by $k_{\mathrm{r},{\mathrm{SiO}}_{\mathrm{2}}}^{\prime }$. In our experiments, the correction for γ was estimated to be 10 %. Initial uptake coefficients (γ₀) and steady-state uptake coefficients (γ_s) were calculated by averaging the signals within the 1.0 and 40–60 min reaction time, respectively.

To understand the diffusion depth of SO₂ and determine the interaction of SO₂ with the underlying layers of SiO₂, the uptake of SO₂ as a function of the SiO₂ mass under irradiation is shown in Fig. S6. The γ exhibited a linear increase in the SiO₂ mass range of 0.05–2.0 g, while it remained unchanged at the SiO₂ mass > 3.0 g. Therefore, the uptake coefficient of SO₂ in the linear regions was normalized using the BET surface area of SiO₂ by Eq. (4) (Brunauer et al., 1938),

$$\begin{array}{}\text{(4)}& {\mathit{\gamma }}_{\text{BET}}={\displaystyle \frac{S_{\text{geo}}\times {\mathit{\gamma }}_{\text{geo}}}{S_{\text{BET}}\times m_{{\mathrm{SiO}}_{\mathrm{2}}}}},\end{array}$$

where γ_BET is the SO₂ uptake coefficient normalized to the BET surface area, S_geo (9.8 × 10⁻³ m²) is the geometric area of the sample dish, S_BET (0.419 m² g⁻¹) is the BET surface area of SiO₂, and $m_{{\mathrm{SiO}}_{\mathrm{2}}}$ (0.05–2.0 g) is the SiO₂ mass. The same method was also used to calculate the uptake coefficients of SO₂ on other mineral oxides.

2.4 In situ DRIFTS analysis

The changes in the chemical compositions on mineral oxides in the SO₂ uptake process were investigated by a Fourier transform infrared (FTIR) spectrometer (Thermo Nicolet iS50) equipped with an in situ diffuse reflectance accessory and a mercury cadmium telluride (MCT) detector. About 14 mg of mineral oxides was placed into a stainless-steel cup inside the reaction cell. To remove adsorbed impurities, SiO₂ was purged with a 150 mL min⁻¹ airflow (${\mathrm{N}}_{\mathrm{2}}/{\mathrm{O}}_{\mathrm{2}}$ volume ratio = 4:1) at RH = 40 % for 1 h. Then, a background spectrum of unreacted samples was collected. SO₂ (2 ppm) was introduced into the reaction cell, and the IR spectra were recorded as a function of time at a resolution of 4 cm⁻¹ by averaging 100 scans. The light from the xenon lamp (500 W) was transmitted into the DRIFTS reaction cell via a fiber. To verify the role of intermediates, Ru(bpy)₃(Cl)₂ and NaHCO₃, acting as ³SO₂ and •OH scavengers (Bulgakov and Safonova, 1996; Gen et al., 2019a), respectively, were mixed with SiO₂ powder in an agate mortar, and the mixture was put in the reaction cell of DRIFTS.

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Figure 1 (a) The temporal variation in the SO₂ concentration on SiO₂ in the dark and under irradiation (7.93 × 10¹⁶ $\mathrm{photons}{\mathrm{cm}}^{-\mathrm{2}}{\mathrm{s}}^{-\mathrm{1}}$); the background changes in the SO₂ concentration in the blank reactor have been deducted. (b) The γ_0,BET and γ_s,BET of SO₂ on SiO₂ as a function of the light intensity. (c) The γ_0,BET and γ_s,BET of SO₂ on SiO₂ at different SO₂ concentrations under irradiation (7.93 × 10¹⁶ $\mathrm{photons}{\mathrm{cm}}^{-\mathrm{2}}{\mathrm{s}}^{-\mathrm{1}}$). The fitting lines for γ_0,BET and γ_s,BET were based on the Langmuir–Hinshelwood mechanism using Eq. (6). Reaction conditions: SiO₂ mass of 0.2 g, temperature of 298 K, RH of 40 %, and O₂ content of 20 %.

3 Results and discussion

3.1 Photoenhanced uptake of SO₂

Acting as the most abundant of the mineral oxides, SiO₂ was used to investigate the uptake behaviors of SO₂. As shown in Fig. 1a, when SO₂ (80 ppb) was exposed to SiO₂ in the dark, the SO₂ concentration decreased to 70 ppb, and then it quickly increased and reached a steady state after 20 min. Upon exposure to SiO₂ under irradiation, the SO₂ concentration exhibited a greater drop than that in the dark. The deactivation of SO₂ uptake on SiO₂ was very slight after 20 min under irradiation. These results suggest that light can promote the heterogeneous reaction of SO₂ on SiO₂. Few studies have observed the photochemical uptake of SO₂ on non-photoactive minerals (Xu et al., 2021; Zhang et al., 2022). When SO₂ did not come into contact with SiO₂, its concentration recovered rapidly. The desorption of SO₂ was observed when SO₂ was isolated from SiO₂ in the dark and under irradiation, indicating that the physical adsorption partially contributed to the SO₂ loss during the photochemical process. The proportion of the desorbed SO₂ during the uptake process can be quantified by dividing the integral of reversible desorption of SO₂ (t = 80–100 min) into the integral of the SO₂ uptake (t = 20–80 min), which was calculated to be 95 % and 12 % in the dark and under irradiation, respectively. This implies that SO₂ uptake in the dark was primarily ascribed to the physical adsorption of SO₂, while SO₂ uptake under irradiation was mainly attributed to chemical processes or irreversible adsorption.

The uptake coefficients of SO₂ on SiO₂ as a function of irradiation intensity are shown in Fig. 1b. The errors in all figures are the standard deviations of three repetitive experiments. Both γ_0,BET and γ_s,BET displayed a strong linear relationship with the irradiation intensity. The γ_0,BET and γ_s,BET under the irradiation of 7.93 × 10¹⁶ $\mathrm{photons}{\mathrm{cm}}^{-\mathrm{2}}{\mathrm{s}}^{-\mathrm{1}}$ were 1.75 and 2.25 times those in the dark, respectively. This further confirms the photochemical nature of the reactions of SO₂ on SiO₂. In particular, γ_0,BET and γ_s,BET on SiO₂ under simulated solar irradiation were comparable with those (10⁻⁷–10⁻⁶) on Gobi Desert dust (GDD) and Arizona test dust (ATD) under UV irradiation, which contained photocatalytic metal oxides (Park et al., 2017). As for the SO₂ uptake on TiO₂, γ_0,BET and γ_s,BET were measured to be 10⁻⁶ and 10⁻⁷, respectively, using the flow tube (Ma et al., 2019), which were similar to our results. It should be pointed out that the similar uptake coefficient did not mean the comparable ability of photoactive and non-photoactive mineral oxides to SO₂ uptake, since the uptake coefficient was highly dependent on environmental conditions (SO₂ concentration, relative humidity, mineral oxides mass, light source, and pressure) and reactor type (chamber and flow tube reactor). Table S1 in the Supplement shows that the percentage of SiO₂ in the sample was 99.02 %, accompanied by a small amount of Al₂O₃, K₂O, Fe₂O₃, and CaO. Photoactive substances (Fe₂O₃) were very few in the sample, and they should not be the main contributor to the photochemical uptake of SO₂.

Figure 1c shows the evolution of γ_0,BET and γ_s,BET at different SO₂ concentrations under irradiation. An inverse dependence of γ_0,BET and γ_s,BET on the SO₂ concentration was observed, meaning that the efficiency of SO₂ uptake decreased with increasing the SO₂ concentration. The uptake of gases on the solid surfaces usually follows the Langmuir–Hinshelwood (L–H) mechanism (Ammann et al., 2003; T. Zhang et al., 2020), suggesting that gaseous molecules are quickly absorbed on the surfaces, and then the reactions occur among the absorbed species. Assuming that the adsorption of SO₂ on SiO₂ is in accordance with the Langmuir isotherm, the dependence of γ on the SO₂ concentration can be described by Eq. (5) (T. Zhang et al., 2020):

$$\begin{array}{}\text{(5)}& \mathit{\gamma }={\displaystyle \frac{(\mathrm{4}V/S\mathit{\omega })k[{\mathrm{SiO}}_{\mathrm{2}}{]}_{\mathrm{T}}{K}_{{\mathrm{SO}}_{\mathrm{2}}}}{\mathrm{1}+K_{{\mathrm{SO}}_{\mathrm{2}}}[{\mathrm{SO}}_{\mathrm{2}}{]}_{\mathrm{g}}}},\end{array}$$

where [SO₂]_g is the concentration of gaseous SO₂, [SiO₂]_T is the total number of active sites on SiO₂, k is the reaction rate constant of SO₂ absorbed on SiO₂, and $K_{{\mathrm{SO}}_{\mathrm{2}}}$ represents the Langmuir adsorption constant of SO₂. Because the SiO₂ mass remained constant during the reaction, Eq. (5) can be written as Eq. (6):

$$\begin{array}{}\text{(6)}& \mathit{\gamma }={\displaystyle \frac{a}{\mathrm{1}+K_{{\mathrm{SO}}_{\mathrm{2}}}[{\mathrm{SO}}_{\mathrm{2}}{]}_{\mathrm{g}}}},\end{array}$$

where $a=(\mathrm{4}V/S\mathit{\omega })k\left[{\mathrm{SiO}}_{\mathrm{2}}\right]K_{{\mathrm{SO}}_{\mathrm{2}}}$. As shown in Fig. 1c, Eq. (6) can well describe the correlation of the SO₂ uptake coefficient with the SO₂ concentration, suggesting that the L–H mechanism can explain the influence of the SO₂ concentration on γ_0,BET and γ_s,BET.

3.2 Photo-induced formation of sulfates by the SO₂ uptake

To investigate the products formed on SiO₂, in situ DRIFTS spectra were recorded, as shown in Fig. 2. The band at 1359 cm⁻¹ was assigned to physically adsorbed SO₂ on SiO₂ (Urupina et al., 2019). The bidentate sulfates and bisulfates contributed to the bands at 1260 and 1229/1074 cm⁻¹ (Urupina et al., 2019; Yang et al., 2020), respectively. The band at 1038 cm⁻¹ may be related to the monodentate sulfites (Yang et al., 2019; Z. Wang et al., 2019). It was noted that the intensity of physically absorbed SO₂ (1359 cm⁻¹) under irradiation was lower than that in the dark (Fig. S7), which may be ascribed to further conversion of SO₂ absorbed on SiO₂ under irradiation. The sulfate bands in particular (1260, 1229, and 1074 cm⁻¹) only appeared under irradiation, while the sulfites (1038 cm⁻¹) were only detected in the dark. This suggests that light changed the SO₂ conversion pathways on SiO₂. As shown in Fig. S7, the bands at 1157/1055 cm⁻¹ were assigned to the asymmetric stretching of Si–O (Hu et al., 2003). Sulfates generated on the surface may interact with SiO₂, leading to a decrease in the intensity of peaks (1157/1055 cm⁻¹).

Figure 2 In situ DRIFTS spectra of SiO₂ during the uptake process of SO₂ (2 ppm) in the dark (a) and under irradiation (b). Reaction conditions: RH of 40 %, temperature of 298 K, and O₂ content of 20 %.

Figure 3 (a) The dependence of γ_0,BET and γ_s,BET on RH. (b) Integrated area of sulfates in DRIFTS spectra (1289–1202 cm⁻¹) as a function of time. (c) The dependence of γ_0,BET and γ_s,BET on O₂. Reaction conditions: SiO₂ mass of 0.2 g, irradiation intensity of 7.93 × 10¹⁶ $\mathrm{photons}{\mathrm{cm}}^{-\mathrm{2}}{\mathrm{s}}^{-\mathrm{1}}$, temperature of 298 K, O₂ content of 20 % for (a), and RH of 40 % for (b).

3.3 Key roles of H₂O and O₂ in photochemical conversion of SO₂ to sulfates

Figure S8a shows temporal variations in the SO₂ concentration in the reaction with SiO₂ at RH = 10 % and 60 % under irradiation. The uptake of SO₂ was very weak at RH = 10 %, whereas it was obvious at RH = 60 %. Moreover, H₂O prolonged the time to reach the steady-state uptake of SO₂. This means that H₂O plays an enhancement role in the photochemical uptake of SO₂. As shown in Fig. 3a, γ_0,BET had a continuous increase from (1.20 ± 0.04) × 10⁻⁷ to (1.54 ± 0.07) × 10⁻⁶ with increasing the RH in the 10 %–60 % range, but it decreased to (1.05 ± 0.09) × 10⁻⁶ at RH = 75 %. The γ_s,BET linearly depended on the RH, and linear fitting to γ_s,BET versus RH yielded an equation γ_s,BET = 1.31 × 10⁻⁸ × RH − 1.02 × 10⁻⁷. Adsorbed H₂O promoted the hydration and dissociation of SO₂ (Huang et al., 2015), and it may generate reactive oxygen species (ROS) such as •OH or HO₂ radicals to oxidize SO₂ under irradiation (Li et al., 2020; Ma et al., 2019), which would lead to positive effects of RH on the SO₂ uptake. Adsorbed H₂O also occupied adsorptive and active sites on the surface, leading to a decrease in SO₂ adsorption. When this competitive role was dominant, the uptake of SO₂ was hindered.

The DRIFTS spectra of SiO₂ during the SO₂ uptake at different RHs are shown in Fig. S9a. The band intensities of sulfates (1260 and 1229 cm⁻¹) at RH = 60 % were stronger than those at RH = 10 %, suggesting that H₂O promotes the sulfate formation. To further investigate the influence of H₂O on the sulfate formation, the integrated area of sulfates in the DRIFTS spectra (1289–1202 cm⁻¹) as a function of the time at different RHs is illustrated in Fig. 3b. Sulfates exhibited a fast formation in the initial 30 min, and then they were continuously generated at a relatively slow rate. SO₂ absorption on the surface was blocked because of the accumulation of H₂O and products (sulfites and sulfates), resulting in the gradual deactivation of the surface. It was noted that sulfate formation was more prominent at higher RHs, revealing that H₂O can act as an important participator in the production of sulfates by the photochemical uptake of SO₂ on SiO₂.

Figure S8b displays effects of O₂ on the photochemical uptake of SO₂ on SiO₂. Negligible SO₂ uptake occurred in N₂, while there was a significant decrease in the SO₂ concentration in air. The γ_0,BET greatly increased from (1.37 ± 0.45) × 10⁻⁷ under anaerobic conditions to (1.19 ± 0.13) × 10⁻⁶ under 20 % O₂ content conditions (Fig. 3c), confirming that O₂ was involved in the reaction of SO₂ on SiO₂. The γ_s,BET increased from (7.10 ± 2.85) × 10⁻⁸ under anaerobic conditions to (4.37 ± 0.58) × 10⁻⁷ under 15 % O₂ content conditions, whereas it remained unchanged with O₂ content further increasing.

Figure 4 (a) The dependence of γ_s,BET under irradiation on the basicity (pH) of mineral oxides. (b) The ratios of steady-state uptake coefficients under irradiation to those in the dark (${\mathit{\gamma }}_{\mathrm{s},\text{Light}}/{\mathit{\gamma }}_{\mathrm{s},\text{Dark}}$). Reaction conditions: mineral oxides mass of 0.2 g, irradiation intensity of 7.93 × 10¹⁶ $\mathrm{photons}{\mathrm{cm}}^{-\mathrm{2}}{\mathrm{s}}^{-\mathrm{1}}$, temperature of 298 K, RH of 40 %, and O₂ content of 20 %.

DRIFTS spectra of SiO₂ during the SO₂ uptake in N₂ and air were compared in Fig. S9b. In both air and N₂, the bands of absorbed SO₂ were 1359 cm⁻¹ and those of sulfates were 1260, 1229, and 1074 cm⁻¹. Nevertheless, their intensities in N₂ were weaker than those in air. According to the integrated area of sulfates in the DRIFTS spectra (1289–1202 cm⁻¹), the formation trends of sulfates were similar in N₂ and air (Fig. 3b), while the sulfate formation rate in N₂ was obviously lower than that in air, meaning that O₂ enhanced the sulfate production. It was reported that the production rate of sulfates from the SO₂ uptake on TiO₂ and by the photolysis of nitrates under UV irradiation in N₂ was also smaller than that in air (Ma et al., 2019; Gen et al., 2019b). In addition, it was noted that sulfates can be generated in N₂, meaning that O₂ was not necessary and some pathways contributed to sulfates without O₂.

3.4 Ubiquitously photoenhanced conversion of SO₂ to sulfates

To better assess the potential for photochemical conversion of SO₂ to sulfates, the SO₂ uptake experiments were further performed on typical mineral oxides without photocatalytic activity. As displayed in Fig. S10, more obvious SO₂ uptake on kaolinite, Al₂O₃, MgO, and CaO was observed under irradiation compared to that in the dark. Figure 4a shows that there was the largest γ_s,BET for CaO among five minerals, and γ_s,BET positively depended on the basicity (pH) of mineral oxides. Basic oxides generally contain more surface hydroxyls, which enhance the heterogeneous uptake of SO₂ (Zhang et al., 2006). The ratios of steady-state uptake coefficients under irradiation to those in the dark (${\mathit{\gamma }}_{\mathrm{s},\text{Light}}/{\mathit{\gamma }}_{\mathrm{s},\text{Dark}}$) were larger than 1.0 for all mineral oxides (Fig. 4b). The experiments for the pH dependence on SiO₂ have also been performed (Fig. S11). The pH of SiO₂ suspension was adjusted to pH = 9, and γ_s,BET and ${\mathit{\gamma }}_{\mathrm{s},\text{Light}}/{\mathit{\gamma }}_{\mathrm{s},\text{Dark}}$ were determined to be (8.79 ± 0.85) × 10⁻⁶ and 1.31, respectively (Fig. 4a and b). These results suggest that light can generally enhance the SO₂ uptake on minerals at a wide pH range. However, the ${\mathit{\gamma }}_{\mathrm{s},\text{Light}}/{\mathit{\gamma }}_{\mathrm{s},\text{Dark}}$ had smaller values with an increase in the basicity, suggesting that the promotion effect of the light was less remarkable for basic oxides.

Figure 5 In situ DRIFTS spectra of kaolinite (a, b), Al₂O₃ (c, d), MgO (e, f), and CaO (g, h) during the uptake process of SO₂ (2 ppm) for 600 min in the dark (black lines) and under irradiation (colorful lines). Reaction conditions: RH of 40 %, temperature of 298 K, and O₂ content of 20 %.

As shown in Fig. 5a and b, the band at 1300 cm⁻¹ should be ascribed to the sulfates. The intensity of sulfates (1300 and 1220 cm⁻¹) under irradiation was larger than those in the dark. Compared to weaker peaks of sulfates (1200 and 1260 cm⁻¹) for Al₂O₃ in the dark (Fig. 5c), a stronger band of bisulfates appeared at 1220 cm⁻¹ under irradiation (Fig. 5d). In contrast to the generation of sulfates for kaolinite and Al₂O₃, both sulfite and sulfate formations were observed for MgO and CaO (Fig. 5e–h). Sulfites were dominant in the dark, as shown by the peaks at 966 and 1020 cm⁻¹ for MgO and 943 cm⁻¹ for CaO, whereas the sulfate formation was significantly enhanced under irradiation according to peak intensities at 1163 cm⁻¹ for MgO and 1137 cm⁻¹ for CaO. It should be noted that these mineral oxides were non-photoactive because of their poor light absorption properties (Fig. S1). Nevertheless, the light can promote the formation of sulfates via the SO₂ uptake process on mineral oxides without photocatalytic activity.

Figure 6 The proposed photochemical conversion mechanisms of SO₂ to sulfates on non-photoactive mineral dust.

3.5 Conversion mechanisms of SO₂ to sulfates

Heterogeneous photochemical reaction mechanisms of SO₂ on non-photoactive mineral dust were proposed in light of experimental observations (Fig. 6). Gaseous SO₂ was adsorbed on the surface (R1) and then reacted with H₂O to form sulfites (R2). Under irradiation, adsorbed SO₂ accepted photons to form its singlet states (¹SO₂) and ³SO₂ (R3–R5) (Sidebottom et al., 1972; Martins-Costa et al., 2018). The reaction between ³SO₂ and H₂O resulted in the formation of HOSO• and •OH (R6), which can combine with SO₂ to produce HOSO₂• (R7). HOSO• and HOSO₂• can be transformed into SO₃, which reacted with H₂O to drive the sulfate formation (R8 and R9). The interaction between ³SO₂ and O₂ may also generate SO₃ directly, which would be subsequently converted to sulfates (R10). Theoretical calculations suggested that the multistep reactions between ³SO₂ with H₂O and O₂ had small energy barriers or were barrier-free (Gong et al., 2022), which could enhance the generation of ROS and the transformation of S(IV) to S(VI). As displayed by R11 and R12, SO₂ and H₂SO₃ adsorbed on the surface may be oxidized to form sulfates via the reactions with ROS, including •O, •OH, or HO₂•, which were produced in R6 and R8–R10. In addition, gaseous SO₂ could be dissolved into adsorbed H₂O to generate bisulfites, which would be finally converted to sulfates by ROS (R13) (Urupina et al., 2019). As displayed in Fig. S12a, the IR peaks of sulfates were not observed when tris(2,2^′-bipyridine)ruthenium dihydrochloride (Ru(bpy)₃(Cl)₂) was employed as the quencher of ³SO₂. The peaks were assigned to the vibrations of excited Ru(bpy)₃(Cl)₂ (Mukuta et al., 2014). This definitely proves that ³SO₂ is the key trigger for the sulfate formation. Figure S12b shows that the peaks of sulfates were weaker in the presence of NaHCO₃, confirming the dominant contribution of •OH formed in R6 and R9 to the formation of sulfates.

Several photochemical mechanisms have been reported to explain the sulfate formation via the SO₂ uptake on various surfaces. Photoactive mineral oxides (such as TiO₂, F₂O₃, and ZnO) can accept photons to produce electron–hole pairs, which generated ROS for the conversion of SO₂ to sulfates (Ma et al., 2019; Li et al., 2019; T. Wang et al., 2020). For example, •OH and HO₂•, generated from the reaction of the hole with H₂O and the electron with O₂, respectively, can act as oxidizing agents for the reaction with SO₂ (Ma et al., 2019). Similarly, the reaction of SO₂ with photo-induced •OH obviously enhanced the formation of sulfates on diesel soot and actual PM_2.5 (Zhang et al., 2022; Y. Zhang et al., 2020). NO₂ and ${\mathrm{NO}}_{\mathrm{2}}^{-}$/HNO₂ can be formed in the nitrates photolysis and primarily contribute to the oxidation of SO₂ to sulfates on nitrates (Gen et al., 2019b, a). Theoretically, the mechanism proposed in this study should also occur on photo-excited substrates. Taking TiO₂ as an example, SO₂ competed with TiO₂ for photons, and the production efficiency of ³SO₂ and excited state of TiO₂ (${\mathrm{TiO}}_{\mathrm{2}}^{\ast }$) depended on their light absorption properties. Meanwhile, ³SO₂ has competition with electron–hole pairs generated from ${\mathrm{TiO}}_{\mathrm{2}}^{\ast }$ for O₂ and H₂O. Thus, the dominant mechanism for the SO₂ uptake on TiO₂ should be related to light absorption properties of precursors and the reactivity for ³SO₂ and ${\mathrm{TiO}}_{\mathrm{2}}^{\ast }$ to O₂ and H₂O. By contrast, all mineral oxides used here cannot be excited under irradiation according to their light absorption spectra (Fig. S1). Nevertheless, SO₂ adsorbed on mineral oxides can absorb ultraviolet radiation (290–400 nm) to form the excited states of SO₂ (${\mathrm{SO}}_{\mathrm{2}}^{\ast }$) (Kroll et al., 2018), which subsequently reacted with H₂O and O₂, finally converting SO₂ to sulfates. The SO₂ uptake experiment in the dark and the visible light (> 420 nm) was carried out (Fig. S13). An ignorable difference was observed for the SO₂ concentration with or without visible light, suggesting that visible light had a minor contribution to the photoenhanced SO₂ uptake.

According to the experimental results, some surfaces providing absorptive sites for SO₂ can enhance the photooxidation of SO₂ to sulfates. However, the promotion effect would vary with different substances. For example, the current experiments on some basic minerals indicate that light plays a minor enhancement role in the SO₂ uptake (Fig. 4), but it could still enhance the sulfate formation (Fig. 5). The solubility and effective Henry's law constant of SO₂ were positively dependent on pH. Thus, SO₂ was more liable to be dissolved to form ${\mathrm{HSO}}_{\mathrm{3}}^{-}/{\mathrm{SO}}_{\mathrm{3}}^{\mathrm{2}-}$ on a more alkaline surface, leading to a strong SO₂ uptake in the dark (Fig. 4a and b) and abundant sulfites on surfaces (Fig. 5). Nevertheless, gaseous SO₂ tends to be adsorbed on kaolinite and Al₂O₃ due to less solubility of SO₂ on these surfaces, and then it is converted to sulfates under irradiation (Fig. 6). Accordingly, a promotion effect of light on SO₂ uptake was observed on neutral and weakly alkaline surfaces (Fig. 4b).

4 Atmospheric implications

The lifetime (τ) for photochemical loss of SO₂ on mineral dust was given using Eq. (7),

$$\begin{array}{}\text{(7)}& \mathit{\tau }={\displaystyle \frac{\mathrm{4}}{\mathit{\gamma }\mathit{\omega }A}},\end{array}$$

where γ and ω are the uptake coefficient and the mean molecular speed of SO₂, respectively. A is the surface area density of mineral dust, and it is estimated to be (1.4–4.8) × 10⁻⁵ cm² cm⁻³ (Zhang et al., 2019; P. He et al., 2018). In this work, γ_s,BET of SO₂ on several mineral oxides was measured to be 4.39 × 10⁻⁷–3.45 × 10⁻⁵ (reaction conditions: SO₂ concentration of 40 ppb, irradiation intensity of 7.93 × 10¹⁶ $\mathrm{photons}{\mathrm{cm}}^{-\mathrm{2}}{\mathrm{s}}^{-\mathrm{1}}$, and RH of 40 %). Thus, the τ of SO₂ with respect to the photooxidation on mineral dust was calculated to be 0.9–240 d, which was shorter than that (54 years) for the photochemical uptake of SO₂ on TiO₂ and the corresponding one (346 d) for the heterogeneous oxidation of SO₂ on ATD in the presence of nitrates (Ma et al., 2019; Zhang et al., 2019). The reaction conditions in this study and those in the literature are different in some respects, and the previously reported SO₂ uptake coefficient (10⁻⁷–10⁻⁶) had a lower value (Ma et al., 2019). The huge difference in the τ of SO₂ was also ascribed to the variation in the surface area density. The content of TiO₂ in mineral dust was only about 1 %. Thus, the surface area density of TiO₂ was about 10⁻⁷ cm² cm⁻³, leading to a longer τ (54 years) for SO₂ on TiO₂ (Ma et al., 2019). It was comparable to the lifetime (3.6–20 d) of SO₂ for the gas-phase reaction with •OH at a concentration of ∼ 10⁻⁶ $\mathrm{molec}.{\mathrm{cm}}^{-\mathrm{3}}$ (Huang et al., 2015; Zhang et al., 2019). Therefore, the photochemical process with the excited state SO₂ acting as a driver on mineral dust was an important pathway for the SO₂ sink in the atmosphere.

Sulfates have a significant influence on the atmosphere, contributing to the haze formation, affecting the activity of aerosols acting as cloud condensation nuclei (CCN) and ice nuclei (IN), and modifying optical property and acidity of aerosols. A sulfate formation rate (R) can be obtained using γ in Eq. (8) (Cheng et al., 2016):

$$\begin{array}{}\text{(8)}& R={\displaystyle \frac{\mathrm{d}\left[{\mathrm{SO}}_{\mathrm{4}}^{\mathrm{2}-}\right]}{\mathrm{d}t}}={\left[{\displaystyle \frac{R_{\mathrm{P}}}{D}}+{\displaystyle \frac{\mathrm{4}}{\mathit{\gamma }\mathit{\omega }}}\right]}^{-\mathrm{1}}A\left[{\mathrm{SO}}_{\mathrm{2}}\right],\end{array}$$

where R_P (m) is the radius of mineral dust, which can be estimated using Eq. (9) (Li et al., 2020), which states

$$\begin{array}{}\text{(9)}& R_{\mathrm{P}}=(\mathrm{0.254}\times \left[{\mathrm{PM}}_{\mathrm{2.5}}\right]/(\mathrm{\mu }\mathrm{g}{\mathrm{m}}^{-\mathrm{3}})+\mathrm{10.259})\times {\mathrm{10}}^{-\mathrm{9}},\end{array}$$

where [PM_2.5] is average PM_2.5 mass concentration and 300 µg m⁻³ is used for the polluted periods in typical China cities (Li et al., 2020; Guo et al., 2014). It was assumed that mineral dust accounted for 50 % of the mass of PM_2.5 (Tohidi et al., 2022), and the mass fraction of SiO₂, Al₂O₃, MgO, and CaO in mineral dust was 60 %, 12.5 %, 4 %, and 6.5 %, respectively (Urupina et al., 2019, 2021; Usher et al., 2003). Thus, R was determined to be 2.15 $\mathrm{\mu }\mathrm{g}{\mathrm{m}}^{-\mathrm{3}}{\mathrm{h}}^{-\mathrm{1}}$. This suggests that the SO₂ uptake on non-photoactive surfaces may be an important sulfate formation pathway under irradiation in some dust-rich conditions.