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. 2023 Oct 7;57(41):15580–15587. doi: 10.1021/acs.est.3c05777

Spontaneous Iodide Activation at the Air–Water Interface of Aqueous Droplets

Yunlong Guo †,‡, Kangwei Li ‡,§,*, Sebastien Perrier ‡, Taicheng An †,*, D James Donaldson ∥,*, Christian George ‡,*
PMCID: PMC10586319  PMID: 37804225

Abstract

graphic file with name es3c05777_0005.jpg

We present experimental evidence that atomic and molecular iodine, I and I2, are produced spontaneously in the dark at the air–water interface of iodide-containing droplets without any added catalysts, oxidants, or irradiation. Specifically, we observe I3– formation within droplets, and I2 emission into the gas phase from NaI-containing droplets over a range of droplet sizes. The formation of both products is enhanced in the presence of electron scavengers, either in the gas phase or in solution, and it clearly follows a Langmuir–Hinshelwood mechanism, suggesting an interfacial process. These observations are consistent with iodide oxidation at the interface, possibly initiated by the strong intrinsic electric field present there, followed by well-known solution-phase reactions of the iodine atom. This interfacial chemistry could be important in many contexts, including atmospheric aerosols.

Keywords: spontaneous iodide activation, air–water interface, iodine, aqueous droplets, atmospheric chemistry

Short abstract

Spontaneous interfacial production of I and I2 is observed in iodide-containing aerosol droplets. This suggests a previously unknown source of reactive iodine species and could be important in the atmospheric marine environment.

Introduction

Aerosols and cloud droplets are key and ubiquitous components in the atmospheric system, presenting a large air–water interface in the atmosphere.1,2 Many species display a minimum free energy at the air–water interface of such droplets, thus creating a favored and unique environment for chemical reactions.3,4 Very recently, it has been found that spontaneous H2O2 and OH radical generation and reduction of organic compounds take place at the air–water interface of microdroplets,5−9 while those reactions do not occur spontaneously in bulk solution. These intriguing phenomena observed from microdroplets do not rely on additional catalysts, external voltages, reducing agents, oxidants, or irradiation, which have received great attention in recent years and were intensively debated.10−12 In a recent study by Li et al.,9 an attempt was made to finally recombine all previous observations by underlying the fact that the appearance of such oxidants is not resulting from an oxidation process, but rather a charge separation process involving the ion–neutral pair of OH–. As initially suggested by Kloss,13 such interfacial reactivity may be driven by the intrinsic electric field (strength reported at the order of ∼109 V m–1) existing at the air–water interface of microdroplets, which is suspected to be strong enough to induce charge separation in ion pairs or facilitate certain reactions by reducing the activation energy.14−16

Molecular iodine (I2) plays a key role in various important atmospheric cycles,17,18 especially in marine environments. Iodine photochemistry has been shown to drive new particle formation in the marine boundary layer19,20 and to reduce global ozone levels by ca. 15%.21 Recently, it has been shown that I2 injections into the stratosphere may even represent a small but nonnegligible process inducing O3 depletion.22 Gas phase I2 in the marine environment originates from various sources such as photodissociation of organo-iodine in tidal and open ocean areas, as well as the direct I2 release from biota near the coastal areas.23−25 Currently, the reaction between ozone and iodide anions is believed to dominate abiotic atmospheric iodine production.18,26,27

Here, we report a novel mechanism for I2 production at the air–water interface of submicrometer- and micrometer-sized NaI-containing droplets. Both gas phase I2 and solution phase I3– are experimentally observed in the dark in the absence of ozone. A very recent publication has also reported on solution-phase I3– production (though not gas phase products) within droplets;28 the present study provides strong evidence that both I3(aq)– and I2(g) are products of the same water surface-mediated dark oxidation of iodide ions. The spontaneous formation of gas phase I2, although perhaps not environmentally significant on Earth due to the low iodide amounts in actual seawater (I– concentration of 10–200 nM),24,29 nevertheless represents another example of redox chemistry that is mediated by the high potentials present at the air–water interface of aqueous droplets.

Materials and Methods

Quantification of I3–

UV–vis absorption spectra of I–, I2(aq), and I3– standard solutions were measured over a relevant range of concentrations. The I3– standard solution was prepared by adding a large excess of NaI to a solution with a known amount of I2(aq); due to the large equilibrium constant for I3– formation, the initial I2(aq) was assumed to fully convert to I3– (see Texts S1 and S2). As shown in Figure 1A, the absorption peaks for these three species can be identified and clearly distinguished; the I3– showed clear absorption peaks at ∼288 and ∼352 nm, respectively, consistent with previous literature.30,31 Using the absorption intensity at 352 nm with a molar absorptivity of 2.38 ± 0.056 × 104 L cm–1 mol–1 (Figure S1 and Text S2), the I3– concentration could be quantified, with the quantification limit at ∼0.20 μM.

Figure 1.

Figure 1

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Spontaneous I3– production from NaI-containing droplets based on UV–vis measurement. (A) Absorption spectra of NaI, I2(aq), and I3– standard solutions at selected concentrations. (B) Time-dependent absorption spectra for bulk solution (100 mM NaI, pH = 6.2) in mist chamber experiment under different spray gases (see inset for enlarged spectra), including pure N2, 20% O2 + 80% N2, and 0.72% C2H4ClOH(g) in N2. (C) Absorption spectra for droplets produced by atomizing the same 100 mM NaI bulk solution under different carrier gases, which were collected as a condensing liquid in two tandem bottles as shown in Figure S2B.

Mist Chamber and Atomizer Experiments

The homemade mist chamber has been reported in our recent study9 and is shown in Figure S2A. It has a volume of ∼110 mL. When the gas flow goes through the mist chamber, the reservoir solution at the bottom will be lifted up due to pressure difference and, therefore, microdroplets are continuously sprayed out. Such microdroplets impacted a hydrophobic PTFE membrane filter (0.2 μm pore size, ref: FGLP04700, Merck Millipore Ltd.) mounted at the top of the mist chamber and then dropped back into the bulk solution for respraying. As shown in Figure S2C, we measured the microdroplet number size distribution (APS, TSI 3321) at the outlet of the mist chamber during the spraying process, and no particles were observed when a membrane filter was added. For a typical mist chamber experiment, 20 mL of bulk NaI solution was added and the total gas flow was adjusted by mass flow controllers at 2.5 L min–1. C2H4ClOH(g), as an electron scavenger with a pKa value of 14.31,32 was generated by a gentle N2 flow passing through the headspace of a 120 mL bottle filled with 30 mL pure C2H4ClOH solution, and the C2H4ClOH(g) concentration was calculated based on its vapor pressure. It is expected that the solution in the mist chamber evaporated continuously during the spraying process, and the remaining volume of the solution decreased gradually over time. We calculated the liquid loss rate for each individual experiment, with the range of 1.7–2.6 mL h–1 (Tables S1 and S2). For these electron-scavenging experiments using C2H4ClOH, the solution pH decreased to 1–2 due to HCl formation through hydrolysis of the dissolved C2H4ClOH(g). Previous studies have shown the important role of solution acidity for I2 production, and iodide activation is favored at lower solution pH.33,34 Additionally, we performed macroscopic aqueous interfacial experiments, as shown in Figure S6, and saw whether spontaneous iodide activation can occur on flat macroscopic surfaces in the dark.

We also produced iodide-containing droplets by atomizing bulk NaI solution using a commercial constant output atomizer (TSI 3076) (Figure S2B), which is a different spraying procedure compared with the mist chamber. Since both spraying procedures used the same bulk NaI solution, we assumed that the composition of the droplets generated from the atomizer is the same as the microdroplets generated from the mist chamber. The droplets produced in the atomizer passed through two tandem 5 L glass bottles, where some large droplets were condensed and collected for UV–vis measurement. Note that we continuously atomized the bulk solution for about 20 h, until enough solution (i.e., 2–3 mL) accumulated in these bottles to be transferred into a quartz cuvette for UV–vis analysis. Although such atomizers are mostly used for generating small droplets in the submicron range (Figure S2E), micron-sized droplets are still produced, but with much lower number concentrations. As shown in Figure S2D, we measured the size distributions of microdroplets produced by such an atomizer, which typically peaked at ∼2 μm diameter with a broad size range.

Gas-Phase I2(g) Measurement with CI-Orbitrap

Bromide chemical ionization mass spectrometry has been frequently deployed to directly measure gas-phase iodine species, where dibromomethane (CH2Br2) is used to produce bromide as reagent ions due to its good affinity to various iodine species.25,35 An ultrahigh-resolution Orbitrap mass spectrometer coupled by a chemical ionization interface (CI-Orbitrap) has emerged as a novel tool, which combines advantages of ultrahigh mass resolving power (m/Δm ≈ 140,000) and minimal fragmentation from soft atmospheric pressure ionization.36,37 The experiment setup is shown in Figure S3, where NaI-containing droplets were produced by a commercial atomizer or our homemade mist chamber and then removed by a hydrophobic PTFE membrane filter to allow a particle-free measurement of the remaining gas phase. Since CH2Br2 was used to produce reagent ions, I2(g) was detected as a bromide adduct such as 79Br(I2)− and 81Br(I2)−, with their ratio consistent with the theoretical isotope prediction (Figure S4). We also semicalibrated the instrument by introducing a known quantity of I2 vapor into the CI-Orbitrap. The instrument showed a good linear response for I2(g) at the ppt level (Text S3 and Figure S5). The data were processed using the Xcalibur 2.2 software (Thermo Scientific), and the selected ions (i.e., m/z = 332.7278 for 79Br(I2)−, m/z = 334.7258 for 81Br(I2)−) were exported with a mass tolerance of 10 ppm. For consistency, all I2(g) intensity reported in this study was shown as 79Br(I2)− and detected at m/z = 332.7278. More detailed information about the CI-Orbitrap can be found elsewhere.36,37

Results and Discussion

Spontaneous I3– Production from I–-Containing Droplets

We carried out experiments in microdroplets (1–10 μm); a few experiments investigated iodide oxidation at macroscopic aqueous interfaces as well. The microdroplet experiments used two droplet sources. In the first one, NaI-containing microdroplets were produced through a homemade mist chamber by spraying a 100 mM NaI bulk solution using various bath gases. The mist chamber was operated in recirculation mode (Figure S2A), where the microdroplets continuously impact a membrane filter and drop back into the bulk solution before being resprayed. Under such recycling conditions, any possible reaction products in the reservoir bulk water of the mist chamber accumulate over time. The bulk reservoir solution was analyzed for the appearance of I3– by UV–vis spectroscopy, typically every hour. Figure 1A displays the absorption spectrum of I3– and its precursors. Figure 1B clearly indicates an increase in absorbance of I3– at ∼352 nm after spraying a 100 mM NaI solution for 7 h using synthetic air (20% O2 + 80% N2) as the bath gas.

Interestingly, when such experiments are performed under nitrogen instead of synthetic air, Figure 1B shows that the increase in absorbance at ∼352 nm is reduced, highlighting a possible role of oxygen in the observed chemistry. Normally, iodide is not oxidized by molecular oxygen in solution or gas phase. To test whether our observation relates to electron capture by oxygen at the air–water interface, 2-chloroethanol (C2H4ClOH)—a strong electron scavenger—was added to the pure nitrogen gas flow and gave rise to a notable increase of the I3– production, as illustrated in Figure 1B.

To test whether the observed iodide activation is dependent on the spraying procedure, we employed a commercial atomizer (TSI 3076), rather than the mist chamber, to nebulize the same NaI solution into microdroplets, which were subsequently trapped in two tandem glass bottles for UV–vis measurement. Figure S2B illustrates the experimental arrangement used for these experiments. As shown in Figure 1C, we also observed the peak at ∼352 nm in this case, indicating the formation of I3– in the microdroplets collected using four different spraying gases. Again, adding oxygen or 2-chloroethanol to the nitrogen gas yielded much greater amounts of I3– product than using pure nitrogen. Compared with the first vessel, the microdroplets trapped in the second vessel generally showed a higher absorbance at ∼352 nm due to a longer residence time.

Finally, we demonstrated that iodide anions are oxidized even on flat macroscopic surfaces in the dark. For this purpose, one glass bottle (30 mL size) was filled with a 3 mL volume of NaI-containing solution with a macroscopic air–water interface (surface area of ca. 8.3 cm2), and another identical glass bottle was fully filled with the same solution with no headspace (i.e., with no air–water interface) (Figure S6). Under such conditions, I3– production was observed after several days from the solution that had a macroscopic air–water interface, while no I3– formation was observed in the solution with no interface, even days after preparing the stock solution (Figure S6). These experiments confirm that iodide oxidation does not occur in bulk solution but requires an interface and that it is not driven by impurities present in our water supply or in the chemicals used.

The enhanced iodide activation due to the addition of electron scavengers suggests the involvement of free electrons in the mechanism. This, in turn, implies that I– may be oxidized to produce a free electron according to (E1), in a manner similar to that previously suggested for the dissociation of hydroxide anions (OH–) into OH radicals and electrons under the naturally formed electric field (∼109 V m–1) at the air–water interface of microdroplets.6,9,13−15,38

graphic file with name es3c05777_m001.jpg E1

Similar to OH– as initially proposed by Kloss,13 the forward step in E1 is suggested to be a charge separation process induced by an interfacial electric field, and hence it should not be regarded as an oxidation process that is costly in energy. For instance, iodide is serving as a typical example for investigating charge-transfer-to-solvent transitions and the corresponding solvation transitions have been investigated.39 However, the transition is typically triggered by the appropriate UV light, in contrast to what is suggested here. The electron affinity of the iodine atom is determined as 3.059 eV.40 It is known that solvation shells do differ between the bulk and the interfacial region and that this may facilitate or promote some redox reactions.41−43 Here, specifically, modified solvation processes at the interface will affect equilibrium E1. If the electron is scavenged, the equilibrium in (E1) will be shifted to the right and enhance iodide activation. Once produced, the atomic iodine may react through a sequence of different reactions (E2–E5) producing I2, I2– and finally I3–, of which many I2– may decay back to I– (E6),44 according to

graphic file with name es3c05777_m002.jpg E2
graphic file with name es3c05777_m003.jpg E3
graphic file with name es3c05777_m004.jpg E4
graphic file with name es3c05777_m005.jpg E5
graphic file with name es3c05777_m006.jpg E6

Under pure nitrogen, the solvated electron generated in E1 can recombine on an approximately hundred picosecond time scale with the nearby I radical due to a solvent cage effect,45 resulting in a null cycle. By contrast, enhanced I3– formation was observed when an electron scavenger such as C2H4ClOH(g), O2(g), or N2O(g) was added to the pure N2 gas spraying flow,46,47 either in the mist chamber (Figures 1B and S7A) or using a commercial atomizer (Figures 1C and S8). Note that C2H4ClOH(g) has a high solubility in water, which can lower the solution pH by forming HCl through hydrolysis of the dissolved C2H4ClOH(g), as illustrated in Figure S7D,E. The lower pH may also affect the electron lifetime in the solution. In addition, by dissolving over time in the aqueous solution, C2H4ClOH may also change the physical characteristics of the nebulized aerosol droplets due to changes in colligative properties and specific molecular interactions. However, our observation of enhanced I3– production in the presence of O2(g) and N2O(g), neither of which is strongly soluble, gives strong support to electron capture at the air–aqueous interface as being a key driver of I3– production (see Text S4 and Figure S8).

The presence of an electron scavenger will shift the equilibrium in E1, allowing I atoms to live sufficiently long to produce I2, I2– as well as I3–, via E2–E6 above. This reaction scheme gives a plausible explanation for the observation of I3– formation. As many of such intermediates decay back to iodide, our observations represent only a lower limit to the actual activation of the iodide anion. It is worth mentioning that a recent work by Xing et al.28 also reported similar I3– formation from spraying NaI bulk solution into microdroplets, and they proposed that the oxidation of iodide was explained by spontaneous production of OH radicals at the interface and subsequent OH-oxidation chemistry. Given the ubiquitous interfacial production of OH radicals,9 this pathway is certainly taking place as well under our conditions, as the experimental approach is overall similar. However, we suggest that this OH-initiated iodide oxidation should be regarded as minor. In fact, the addition of an organic compound, in large excess and acting as an OH radical scavenger, should prevent the formation of I2 or I3–. By contrast, the addition of 2-chloroethanol did not prevent the formation of these products but rather enhanced their appearance, showing indeed that the key element in this chemistry is the electron trapping process that displaces equilibrium E1 to the right. Additionally, our observation of gas phase I2 formation and Langmuir–Hinshelwood behavior (vide infra) suggests that spontaneous iodide activation is more likely driven by a surface-mediated process involving the loss of free electrons (E1).

I3– production rates (in nmol h–1; see Text S5) measured under different spraying gases were summarized in Figure 2A. The clear evidence for activation of iodide seen in the droplets was not observed in bulk solution and was enhanced in the presence of gas phase electron scavengers. These observations are consistent with a surface-mediated mechanism taking place at the air–water interface. This conjecture is supported by the dependence of I3– production on the bulk iodide concentration. Figure 2B shows that the I3– production rate does not depend linearly on bulk iodide concentration (as expected for a first-order bulk process), but exhibits limiting behavior, with a linear dependence at lower I– concentrations (0.1–20 mM), and reaching a plateau at higher I– concentrations (20–100 mM). The detailed results for each individual mist chamber experiment are summarized in Figure S9 and Table S1. Such a relationship between the I3– production rate and the NaI bulk concentration is well represented by a Langmuir–Hinshelwood model, which is specific for surface-driven processes.48−51 This idea is explored further below. Altogether, these indications are strongly supportive of an interfacial oxidation of iodide anion, giving rise to I3– production.

Figure 2.

Figure 2

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I3– production rate from mist chamber experiments under different conditions. (A) Fixed NaI bulk solution (100 mM) under different spraying gases. (B) Varied NaI bulk solution (0.1–100 mM) under the same spraying gas (0.72% C2H4ClOH(g) in N2). Error bars represented standard deviation (1σ) from 2 to 10 measurements from different spraying time (Tables S1–S2).

Release of Interfacial I2 into the Gas Phase

Molecular iodine is only moderately soluble in water, with a Henry’s law constant of ca. 2.8 M atm–1 at room temperature.52 As a consequence, if I2 production occurs at the air–water interface, then one would expect a significant fraction of that compound to desorb into the gas phase. To test whether gaseous iodine is produced at the droplet surface, we performed a direct measurement of I2(g) using an online chemical ionization orbitrap (CI-Orbitrap) instrument, with bromide as a reagent ion.

As shown in Figure 3A, we observed clear steady-state I2(g) signal (detected as bromide-adduct such as 79Br(I2)−) when atomizing a 100 mM NaI solution under N2 or pure air, while a 100 mM (NH4)2SO4 solution nebulized by the same procedure led to signal at the level of instrumental background. Compared with the results using N2 as a carrier gas, an enhanced I2(g) signal was observed using pure air (Figures 3A–B and S10A,B), consistent with the observations of enhanced I3– production in the presence of O2(g) discussed above. This comparison reinforces the fact that electron scavengers such as oxygen can play an important promotion role in interfacial iodide activation. Figures 3C and S10C show the I2(g) signal as a function of pH of a 10 mM NaI bulk solution. The production of I2 seems to be more efficient under acidic conditions, again similar to what we see for I3– production in Figure S7C. Such pH dependence agrees with previous findings that abundant protons (H+) at lower pH could enhance the production of I3(aq)– and I2(g).33,53 As shown in Figure 3D, the I2(g) signal also shows a nonlinear dependence on the bulk NaI concentration, similar to what is shown above in Figure 2B for I3– production. This is compelling evidence to support the proposal that both spontaneous I2(g) production and I3– production via iodide activation occur as an interfacial process. Such Langmuir–Hinshelwood behavior also rules out the potential iodide activation route from the interfacial production of H2O2. For instance, if I2 would arise from the titration of iodide by H2O2, then the amount of I2 produced should not depend on the I– concentration due to the great excess of I– to H2O2, which does not agree with the above observation in this study.

Figure 3.

Figure 3

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Direct I2(g) measurement from NaI-containing droplets under various experimental conditions. (A) Time-dependent raw I2(g) intensity measured at the output of the atomizer by atomizing 100 mM NaI under N2 and air, respectively. (B) I2(g) intensity measured at the output of atomizer or mist chamber with 100 mM or 10 mM NaI under air vs N2. (C) I2(g) intensity measured at the output of atomizer with 10 mM NaI under different solution pH. (D) I2(g) intensity measured at the output of the atomizer as a function of bulk NaI concentration.

The CI-Orbitrap instrument demonstrated a linear response to different I2(g) concentrations and was only calibrated for semiquantitative purposes (see Figure S5 and Text S3 for details). The corresponding semiquantitative I2(g) concentrations for all experimental conditions tested here are summarized in Table S3. Total surface area concentrations could be calculated based on the droplet size distribution measured at the output of the atomizer or mist chamber (Figure S2C–E), respectively. As shown in Table S4, the surface area normalized I2(g) concentrations were calculated to be in the range of 2–8 × 1015 mol m–2 for droplets produced from the atomizer and mist chamber respectively, which agrees in between within a factor of 4.

Langmuir–Hinshelwood Behavior

The dependence of product formation on the bulk iodide concentration is consistent between both aqueous phase I3– production (Figure 2B) and gas phase I2(g) production (Figure 3D). This implies that the production of both species has the same dependence on the presence of iodide anions at the air–water interface. To describe this process by a Langmuir–Hinshelwood kinetic model, we normalize all the production rates to their maxima and plot these as a function of bulk NaI concentration in Figure 4, fitting the data to a Langmuir–Hinshelwood equation (eq 1)

graphic file with name es3c05777_m007.jpg 1

Figure 4.

Figure 4

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Observed I3– and I2(g) production from spontaneous iodide activation as a function of bulk NaI concentration. The data from Figures 2B and 3D were normalized to their respective maxima and the combined data fit with a Langmuir–Hinshelwood model.

In this equation, [I–aq] is the bulk concentration of iodide; parameter A represents the maximum observed rate; parameter B represents the ratio of the desorption to adsorption rate constants for I– adsorption from the bulk solution to the interface, which is equivalent to the inverse of the equilibrium constant of I–aq for surface adsorption (Kads).49,50

By applying a Langmuir–Hinshelwood fit to the combined data in Figure 4, we obtain an adsorption constant Kads of 295 ± 100 M–1 for I– partitioning from bulk to the interface. Note that this does not represent a true thermodynamic equilibrium constant, as we have ignored the activity coefficient, and the surface standard state is not defined. However, comparing this value to other similar analyses shows reasonable agreement within a factor of 4–13, with the reported Kads values (i.e., 23 ± 3 M–1 and 70 ± 16 M–1), from analysis of the heterogeneous reaction between O3(g) and NaI solution.50,51

Atmospheric Implications

It is worth mentioning that some previous studies also reported I3(aq)– and I2(g) production as a result of the addition of an external trigger, i.e., ozone, irradiation, catalyst, etc.33,34,53,54 We observed spontaneous I3– and I2(g) production from NaI droplets through an iodide activation process that does not require catalysts, external voltages, oxidants, or irradiation. This represents a previously unknown source of reactive iodine species. Our results suggest that such spontaneous dark production of I3– and I2(g) is an interfacial process, which can be described well using a Langmuir–Hinshelwood formalism. The production of I3– and I2(g) is strongly affected by the presence of electron scavengers such as N2O, O2, or C2H4ClOH, consistent with oxidation of iodide at the interface.

We propose that this spontaneous iodide activation could be induced by the strong electric field (∼109 V m–1) which is thought to be formed naturally at the air–water interface of microdroplets.14,15 However, molecular dynamics simulations show large fluctuations across the interface due to partial solvation effect and the associated variation of the solvent electrostatic potential, which suggests the incomplete and large uncertainty in the current understanding of interfacial electric field.55,56 Similar to the recent observation of spontaneous H2O2 and OH radical production at the interface of microdroplets,9,38 we hypothesize that such a high electric field can dissociate iodide anions into iodine atoms and electrons (e–) (E1), following by the formation of reactive iodine species (I and I2), which can further accommodate into the aqueous phase or be released into the gas phase and then potentially trigger more iodine-related reactions with other compounds. For instance, the aqueous iodine species may react with dissolved organic matter in the sea-surface microlayer to produce organo-iodine species,21,23 while the interfacial released I2(g) could deplete tropospheric O3, initiate new particle formation and oxidize hydrocarbons.17,57,58 The Supporting Information (Text S6 and Figures S11–S12) gives details of some preliminary experiments that show new particle formation by reaction of I2 released from droplet surfaces with ozone under visible irradiation.

Acknowledgments

This study was supported by the European Research Council under the Horizon 2020 research and innovation programme/ERC Grant Agreement 101052601—SOFA. T.A. and Y.G. acknowledge the financial support from the National Natural Science Foundation of China (42020104001 and 42377221).

Supporting Information Available

The Supporting Information is available free of charge at https://pubs.acs.org/doi/10.1021/acs.est.3c05777.

  • The Supporting Information includes additional experimental details, materials and methods, and further results from mist chamber and atomizer experiments, I2(g) calibration with CI-Orbitrap and aerosol flow tube experiments (PDF)

Author Contributions

⊥ Y.G. and K.L. contributed equally to this work.

The authors declare no competing financial interest.

Supplementary Material

es3c05777_si_001.pdf (1.2MB, pdf)

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