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# Water film-mediated photocatalytic oxidation of oxalate on TiO2

N Tan Luong, Khalil Hanna, Jean-François Boily

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N Tan Luong, Khalil Hanna, Jean-François Boily. Water film-mediated photocatalytic oxidation of oxalate on TiO2. Journal of Catalysis, 2024, Journal of Catalysis, 432, pp.115425. 10.1016/j.jcat.2024.115425. hal-04506368

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# Water film-mediated photocatalytic oxidation of oxalate on TiO₂

N. Tan Luong<sup>a,1,*</sup>, Khalil Hanna<sup>b</sup>, Jean-François Boily<sup>a</sup>

<figure><img src="image_2.png" /></figure>

Check for
updates

## ARTICLE INFO

**Keywords:**

Photocatalysis

TiO₂

Water films

Oxalate

Reactive oxygen species

Hydroxyl radicals

## ABSTRACT

Water films on minerals under humid environment can be photocatalytic hotspots when exposed to sunlight or artificial sources of ultraviolet (and visible) radiation. In this study, we resolved the water film-mediated photocatalysis of oxalate adsorbed on TiO₂ using *in situ* infrared spectroscopy. We found that 0.5 to 4 monolayer- (ML) thick water films enhanced the photodecomposition rates of oxalate under 21 kPa O₂. We explained this through the combined actions of direct hole transfer, ligand-to-metal-charge transfer, as well as the production of hydroxyl radicals and reactive oxygen species. Rates were, however, substantially slower in the absence of O₂ because charge recombination, together with water film-mediated charge localization, disrupted hole transfer and hydroxyl radical production. Our work adds insight into the impact of humidity on controlling important photocatalytic processes in nature (drying soils, atmospheric aerosols), and technology (water and air treatment).

## 1. Introduction

Air moisture captured by mineral particles can produce nanometrically-thick water films [1–3] that mediate compositional and structural transformation reactions. [4–8] They can also be photocatalytic hotspots when exposed to sunlight or artificial sources of ultraviolet radiation. For instance, water films formed on atmospheric mineral dust can experience periodic day-night cycles that are responsible for controlling the budget of important pollutants (e.g. O₃, SOₓ, NOₓ, organics) in the troposphere and terrestrial soils through photocatalytic reactions. [8] Understanding of mineral photo-reactivity in the presence of these films is important for fields including atmospheric chemistry, [5,8] geochemistry, [9] water and air purification technology, [10–13] CO₂ utilization, [14] energy production, [15,16] and even space exploration [17].

TiO₂ is an ideal semiconducting model mineral that is beneficial for advancing knowledge on water film-driven photocatalytic reactions. It has a reputably high photocatalytic efficiency on a wide range of substances, [8] is chemically stable, and is appreciably abundant in Earth's continental crust (0.7 % wt.) [18] and atmosphere (0.4–5 %). [8] Anatase and rutile are the most studied and applied TiO₂ phases in photocatalysis, each with different bulk and surface properties (i.e., crystal structure, band structure, defect sites, surface charges, surface OH population, etc.) that control charge carrier generation and

recombination, and consequently photocatalytic activity. [19,20] TiO₂-driven photocatalytic transformation studies range from fundamental aspects [8,9,19,21–25] to applications [8,9,16,24–27]. While there is an increasing body of work on TiO₂-based catalyst performance (e.g., expanding its working range to visible light), [16,28] novel insights on photocatalytic mechanisms are still urgently needed to broaden working conditions, such as environments that experience large variations in atmospheric humidity [8].

Photocatalytic reactions on TiO₂ generally begin with the photo-generation of electron-hole (e⁻/h⁺) pairs by light (Fig. 1 a). Unless they recombine, these charge carriers generate (i) hydroxyl radicals (•OH) via direct hole oxidation of surface hydroxo groups (≡OH + h⁺ → ≡□ + •OH) or water (H₂O + h⁺ → •OH + H⁺), or (ii) reactive oxygen species (ROS) beginning with the one-electron reduction of oxygen (O₂ + e⁻ → O₂⁻) (Fig. 1 b). These radicals can then degrade mineral-bound organic molecules, ultimately down to CO₂ and H₂O. Earlier studies revealed that the photo-reactivity of TiO₂ towards organics (e.g. phenol, [29] toluene, [30,31] acetic acid, [32] acetaldehyde [32,33] and ethanol [34]) was influenced by water films formed by the condensation of air moisture. Moreover, water films thickness can alter reaction mechanisms and rates [30,33,35,36] by controlling concentrations and distributions of charge carriers and radicals, [33,37–39]. Film thickness can also affect chemical speciation and adsorption modes on TiO₂, including metal-bonded (MB), hydrogen-bonded (HB), or outer-sphere

<footnote><sup>a</sup> Department of Chemistry, Umeå University, SE 901 87 Umeå, Sweden</footnote>

<footnote><sup>b</sup> Université de Rennes, École Nationale Supérieure de Chimie de Rennes, CNRS, ISCR-UMR 6226, F-35000 Rennes, France</footnote>

<footnote>* Corresponding author.</footnote>

<footnote>E-mail address: tan.luong@chalmers.se (N.T. Luong).</footnote>

<footnote n="1">¹ Current address: Department of Physics, Chalmers University of Technology, SE 412 96 Gothenburg, Sweden.</footnote>
# Water film-mediated photocatalytic oxidation of oxalate on TiO₂

N. Tan Luong<sup>a,1,*</sup>, Khalil Hanna<sup>b</sup>, Jean-François Boily<sup>a</sup>

<figure><img src="image_3.png" /></figure>

Check for
updates

## ARTICLE INFO

### Keywords:

Photocatalysis

TiO₂

Water films

Oxalate

Reactive oxygen species

Hydroxyl radicals

## ABSTRACT

Water films on minerals under humid environment can be photocatalytic hotspots when exposed to sunlight or artificial sources of ultraviolet (and visible) radiation. In this study, we resolved the water film-mediated photocatalysis of oxalate adsorbed on TiO₂ using *in situ* infrared spectroscopy. We found that 0.5 to 4 monolayer- (ML) thick water films enhanced the photodecomposition rates of oxalate under 21 kPa O₂. We explained this through the combined actions of direct hole transfer, ligand-to-metal-charge transfer, as well as the production of hydroxyl radicals and reactive oxygen species. Rates were, however, substantially slower in the absence of O₂ because charge recombination, together with water film-mediated charge localization, disrupted hole transfer and hydroxyl radical production. Our work adds insight into the impact of humidity on controlling important photocatalytic processes in nature (drying soils, atmospheric aerosols), and technology (water and air treatment).

## 1. Introduction

Air moisture captured by mineral particles can produce nanometrically-thick water films [1–3] that mediate compositional and structural transformation reactions. [4–8] They can also be photocatalytic hotspots when exposed to sunlight or artificial sources of ultraviolet radiation. For instance, water films formed on atmospheric mineral dust can experience periodic day-night cycles that are responsible for controlling the budget of important pollutants (e.g. O₃, SOₓ, NOₓ, organics) in the troposphere and terrestrial soils through photocatalytic reactions. [8] Understanding of mineral photo-reactivity in the presence of these films is important for fields including atmospheric chemistry, [5,8] geochemistry, [9] water and air purification technology, [10–13] CO₂ utilization, [14] energy production, [15,16] and even space exploration [17].

TiO₂ is an ideal semiconducting model mineral that is beneficial for advancing knowledge on water film-driven photocatalytic reactions. It has a reputably high photocatalytic efficiency on a wide range of substances, [8] is chemically stable, and is appreciably abundant in Earth's continental crust (0.7 % wt.) [18] and atmosphere (0.4–5 %). [8] Anatase and rutile are the most studied and applied TiO₂ phases in photocatalysis, each with different bulk and surface properties (i.e., crystal structure, band structure, defect sites, surface charges, surface OH population, etc.) that control charge carrier generation and

recombination, and consequently photocatalytic activity. [19,20] TiO₂-driven photocatalytic transformation studies range from fundamental aspects [8,9,19,21–25] to applications [8,9,16,24–27]. While there is an increasing body of work on TiO₂-based catalyst performance (e.g., expanding its working range to visible light), [16,28] novel insights on photocatalytic mechanisms are still urgently needed to broaden working conditions, such as environments that experience large variations in atmospheric humidity [8].

Photocatalytic reactions on TiO₂ generally begin with the photo-generation of electron-hole (e⁻/h⁺) pairs by light (Fig. 1 a). Unless they recombine, these charge carriers generate (i) hydroxyl radicals (•OH) via direct hole oxidation of surface hydroxo groups (≡OH + h⁺ → ≡□ + •OH) or water (H₂O + h⁺ → •OH + H⁺), or (ii) reactive oxygen species (ROS) beginning with the one-electron reduction of oxygen (O₂ + e⁻ → O₂⁻) (Fig. 1 b). These radicals can then degrade mineral-bound organic molecules, ultimately down to CO₂ and H₂O. Earlier studies revealed that the photo-reactivity of TiO₂ towards organics (e.g. phenol, [29] toluene, [30,31] acetic acid, [32] acetaldehyde [32,33] and ethanol [34]) was influenced by water films formed by the condensation of air moisture. Moreover, water films thickness can alter reaction mechanisms and rates [30,33,35,36] by controlling concentrations and distributions of charge carriers and radicals, [33,37–39]. Film thickness can also affect chemical speciation and adsorption modes on TiO₂, including metal-bonded (MB), hydrogen-bonded (HB), or outer-sphere

<footnote><sup>a</sup> Department of Chemistry, Umeå University, SE 901 87 Umeå, Sweden</footnote>

<footnote><sup>b</sup> Université de Rennes, École Nationale Supérieure de Chimie de Rennes, CNRS, ISCR-UMR 6226, F-35000 Rennes, France</footnote>

<footnote>* Corresponding author.</footnote>

<footnote>E-mail address: tan.luong@chalmers.se (N.T. Luong).</footnote>

<footnote n="1">¹ Current address: Department of Physics, Chalmers University of Technology, SE 412 96 Gothenburg, Sweden.</footnote>
<figure><img src="image_3.png" /><figcaption>**Fig. 1.** Schematic representation of TiO₂ photoexcitation producing •OH radicals in thin water films, and subsequent scavenging by oxalate. (a) Light activates charge separation in the TiO₂ bulk. While these charged carriers migrate from the bulk to the surface, they can be trapped and recombined. (b) Charges that are not trapped and/or recombined reach the surface and react with adsorbed (metal-bonded (MB), hydrogen-bonded (HB), and outer-sphere (OS)) species. Holes oxidize water film molecules and TiO₂ surface hydroxyl groups, producing hydroxyl radicals (•OH). Electrons combine with molecular oxygen through a series of reduction, protonation, and disproportionation reaction to generate reactive oxidative species (ROS, e.g., O₂⁻/HO₂, H₂O₂, •OH). These radicals react with adsorbed oxalate (electron donor) of different binding modes (MB and OS/HB), to produce CO₂ as a final product.</figcaption></figure>

(OS, with intervening water molecules) complexes (Fig. 1 b).[40,41] Uncertainties however remain regarding to the decomposition mechanisms, including whether water films can promote or inhibit TiO₂ photo-reactivity.[29–36].

To advance our knowledge of water film-driven photocatalysis, we resolved the water film-mediated photodecomposition of a dicarboxylate ligand, oxalate (OX), bound to TiO₂ nanoparticles. Oxalate is an abundant low-molecular weight organic compounds in atmospheric aerosols [42] and soil systems.[43] It is also a contaminant from industrial activities, such as in alumina production [44] and metallurgy. [45] Focusing on the photo-reactivity of oxalate on TiO₂ is consequently relevant for environmental problems, as well as making new fundamental chemistry accessible for study.

Previous solution-based work showed that oxalate and oxalic acid are easily decomposed by photo-activated TiO₂ to CO₂ or (bi)carbonates.[46] This can be understood by the low reduction potential of CO₂(g)/H₂C₂O₄(aq) (~0.49 V vs standard hydrogen electrode). Photo-degradation rates of oxalate were also shown to be larger at pH values below the isoelectric point of TiO₂ (pH 5.6–6).[47] This can be understood by vibrational spectroscopy work showing a dominance of MB oxalate surface complexes under acidic conditions, [48–50] namely where the surface is positively charged.[20] Under these conditions, oxalic acid decomposition to CO₂(g) by TiO₂ follows a Langmuir-Hindshelwood mechanism, one that involves two adsorbed species. [51] Other studies [50,52,53] also suggested that readsorption during irradiation changed the configuration of oxalate on TiO₂, a mechanism that, in turn, modified photodecomposition mechanisms.[54] Oxalate was, additionally, studied as an effective hole scavenger in the photocatalytic degradation of pharmaceutical pollutants by TiO₂ under visible light.[55]

In this work, we offer new insight into the photocatalytic decomposition of TiO₂-bound oxalate driven by nanometric water films, formed under atmospherically relevant conditions. This work fills the gaps between efforts in high-vacuum and aqueous systems, and gas-phase based TiO₂ photocatalysis. Using vibrational spectroscopy, we demonstrate that variations in water film thickness drove contrasting mechanisms on oxalate photodecomposition rates and yields under O₂-bearing and O₂-free environments. Our findings have direct implications for understanding water film-driven photocatalytic reactions that are important to terrestrial environments (e.g., vadose zones of soils), atmospheric chemistry (e.g., aerosol composition) and technology (e.g., indoor air treatment).

## 2. Experimental Section

### 2.1. Material characterization

Commercial TiO₂ (Aldrich, No. 718467) with properties similar to the well-studied commercial P25 Degussa (Evonik) [56,57] was used to study oxalate degradation in thin water films. These particles consist of a physico-chemical mixture of 90 % anatase and 10 % rutile, [56] as determined by powder X-ray diffraction (XRD, PANalytical X'Pert³, Supporting Information, Fig. S1). Scanning and transmission electron microscopy (SEM/TEM, Fig. S1) revealed a homogeneous distribution of 21 ± 6 nm sphere-like or cuboid particles exposing dominant (101), (100), (001) and (010) surfaces.[57,58] A 91-point N₂(g) adsorption/desorption isotherm (Micromeritics TriStar 3000) gave a B.E.T. specific surface area [59] of 53 m²/g. X-ray photoelectron spectroscopy (XPS, Kratos Axis Ultra) showed a surface composition of Ti:O:OH = 1:2.1:0.17 (Supporting Information, Table S1), indicating the presence of Ti vacancies and non-stoichiometric surface OH groups. Furthermore, XPS confirmed that all surface Ti sites were Ti(IV). Details on all characterization methods are accessible in Supporting Information. All experiments were performed on this TiO₂ sample without any further treatment.

### 2.2. Adsorption experiments and chromatography analyses of solutions

Oxalate was adsorbed in aqueous suspensions of TiO₂ nanoparticles under N₂(g). Experiments were performed in 15-mL Falcon® tubes covered by aluminum foil to avoid exposure to light. MilliQ water was used for preparing all solutions. A 0.5 mL aliquot of an aqueous suspension of 20 g/L TiO₂ was mixed with 9.5 mL of 0.5 mM oxalic acid solution (>99.5 %, Scharlau) under N₂(g), achieving a total oxalate concentration ([OX]tot) of 47.5 mM per g TiO₂. This amounts to a total concentration of ~ 5.4 oxalate molecules per nm² TiO₂ (OX/nm²). Small aliquots of 1 M HCl (Merck) and 1 M NaOH (Merck) were used to adjust the suspension pH to ~ 3.2, resulting in ionic strength values of ~ 1 mM. The resulting suspensions were then equilibrated using an end-to-end rotator for 24 h before they were centrifuged at 2900 g for 15 min. The supernatants were thereafter collected by filtration using a 0.2-µm filter, and their final oxalate concentration analyzed by ion chromatography (IC).

The IC system consisted of a modified liquid chromatography instrument (LC-10AD, Shimadzu), an anion separation column (Metrosep A Supp 5–150/4.0, Metrohm), a suppressor system equipped with a suppressor membrane (Xenoic® XAMS, Diduco AB) and a generation pump (Xenoic® ASUREX-A100, Diduco AB), and a conductivity detector (prototype, no brand) to detect ions. The mobile phase was a carbonate
<figure><img src="image_4.png" /><figcaption>**Fig. 1.** Schematic representation of TiO₂ photoexcitation producing •OH radicals in thin water films, and subsequent scavenging by oxalate. (a) Light activates charge separation in the TiO₂ bulk. While these charged carriers migrate from the bulk to the surface, they can be trapped and recombined. (b) Charges that are not trapped and/or recombined reach the surface and react with adsorbed (metal-bonded (MB), hydrogen-bonded (HB), and outer-sphere (OS)) species. Holes oxidize water film molecules and TiO₂ surface hydroxyl groups, producing hydroxyl radicals (•OH). Electrons combine with molecular oxygen through a series of reduction, protonation, and disproportionation reaction to generate reactive oxidative species (ROS, e.g., O₂⁻/HO₂, H₂O₂, •OH). These radicals react with adsorbed oxalate (electron donor) of different binding modes (MB and OS/HB), to produce CO₂ as a final product.</figcaption></figure>

(OS, with intervening water molecules) complexes (Fig. 1 b).[40,41] Uncertainties however remain regarding to the decomposition mechanisms, including whether water films can promote or inhibit TiO₂ photo-reactivity.[29–36].

To advance our knowledge of water film-driven photocatalysis, we resolved the water film-mediated photodecomposition of a dicarboxylate ligand, oxalate (OX), bound to TiO₂ nanoparticles. Oxalate is an abundant low-molecular weight organic compounds in atmospheric aerosols [42] and soil systems.[43] It is also a contaminant from industrial activities, such as in alumina production [44] and metallurgy. [45] Focusing on the photo-reactivity of oxalate on TiO₂ is consequently relevant for environmental problems, as well as making new fundamental chemistry accessible for study.

Previous solution-based work showed that oxalate and oxalic acid are easily decomposed by photo-activated TiO₂ to CO₂ or (bi)carbonates.[46] This can be understood by the low reduction potential of CO₂(g)/H₂C₂O₄(aq) (~0.49 V vs standard hydrogen electrode). Photo-degradation rates of oxalate were also shown to be larger at pH values below the isoelectric point of TiO₂ (pH 5.6–6).[47] This can be understood by vibrational spectroscopy work showing a dominance of MB oxalate surface complexes under acidic conditions, [48–50] namely where the surface is positively charged.[20] Under these conditions, oxalic acid decomposition to CO₂(g) by TiO₂ follows a Langmuir-Hindshelwood mechanism, one that involves two adsorbed species. [51] Other studies [50,52,53] also suggested that readsorption during irradiation changed the configuration of oxalate on TiO₂, a mechanism that, in turn, modified photodecomposition mechanisms.[54] Oxalate was, additionally, studied as an effective hole scavenger in the photocatalytic degradation of pharmaceutical pollutants by TiO₂ under visible light.[55]

In this work, we offer new insight into the photocatalytic decomposition of TiO₂-bound oxalate driven by nanometric water films, formed under atmospherically relevant conditions. This work fills the gaps between efforts in high-vacuum and aqueous systems, and gas-phase based TiO₂ photocatalysis. Using vibrational spectroscopy, we demonstrate that variations in water film thickness drove contrasting mechanisms on oxalate photodecomposition rates and yields under O₂-bearing and O₂-free environments. Our findings have direct implications for understanding water film-driven photocatalytic reactions that are important to terrestrial environments (e.g., vadose zones of soils), atmospheric chemistry (e.g., aerosol composition) and technology (e.g., indoor air treatment).

## 2. Experimental Section

### 2.1. Material characterization

Commercial TiO₂ (Aldrich, No. 718467) with properties similar to the well-studied commercial P25 Degussa (Evonik) [56,57] was used to study oxalate degradation in thin water films. These particles consist of a physico-chemical mixture of 90 % anatase and 10 % rutile, [56] as determined by powder X-ray diffraction (XRD, PANalytical X'Pert³, Supporting Information, Fig. S1). Scanning and transmission electron microscopy (SEM/TEM, Fig. S1) revealed a homogeneous distribution of 21 ± 6 nm sphere-like or cuboid particles exposing dominant (101), (100), (001) and (010) surfaces.[57,58] A 91-point N₂(g) adsorption/desorption isotherm (Micromeritics TriStar 3000) gave a B.E.T. specific surface area [59] of 53 m²/g. X-ray photoelectron spectroscopy (XPS, Kratos Axis Ultra) showed a surface composition of Ti:O:OH = 1:2.1:0.17 (Supporting Information, Table S1), indicating the presence of Ti vacancies and non-stoichiometric surface OH groups. Furthermore, XPS confirmed that all surface Ti sites were Ti(IV). Details on all characterization methods are accessible in Supporting Information. All experiments were performed on this TiO₂ sample without any further treatment.

### 2.2. Adsorption experiments and chromatography analyses of solutions

Oxalate was adsorbed in aqueous suspensions of TiO₂ nanoparticles under N₂(g). Experiments were performed in 15-mL Falcon® tubes covered by aluminum foil to avoid exposure to light. MilliQ water was used for preparing all solutions. A 0.5 mL aliquot of an aqueous suspension of 20 g/L TiO₂ was mixed with 9.5 mL of 0.5 mM oxalic acid solution (>99.5 %, Scharlau) under N₂(g), achieving a total oxalate concentration ([OX]tot) of 47.5 mM per g TiO₂. This amounts to a total concentration of ~ 5.4 oxalate molecules per nm² TiO₂ (OX/nm²). Small aliquots of 1 M HCl (Merck) and 1 M NaOH (Merck) were used to adjust the suspension pH to ~ 3.2, resulting in ionic strength values of ~ 1 mM. The resulting suspensions were then equilibrated using an end-to-end rotator for 24 h before they were centrifuged at 2900 g for 15 min. The supernatants were thereafter collected by filtration using a 0.2-µm filter, and their final oxalate concentration analyzed by ion chromatography (IC).

The IC system consisted of a modified liquid chromatography instrument (LC-10AD, Shimadzu), an anion separation column (Metrosep A Supp 5–150/4.0, Metrohm), a suppressor system equipped with a suppressor membrane (Xenoic® XAMS, Diduco AB) and a generation pump (Xenoic® ASUREX-A100, Diduco AB), and a conductivity detector (prototype, no brand) to detect ions. The mobile phase was a carbonate
<figure><img src="image_4.png" /><figcaption>**Fig. 2.** Oxalate binding on TiO₂ nanoparticles. (a) ATR-FTIR spectra of wet (blue) and N₂-dry (black) films of TiO₂ with oxalate loadings of 1.4 OX/nm², and ATR-FTIR of references (green): dry TiO₂ film, HC₂O₄ ions and C₂O₄²⁻ ions. The spectrum of the wet TiO₂-oxalate film was obtained on the centrifuged wet paste deposited on ATR surfaces, and the spectrum of the film was taken after this wet paste was dried under N₂(g) (c.f. Section 2.4). The spectra of HC₂O₄ ions and C₂O₄²⁻ ions were obtained from 100 mM oxalate solutions at pH 3.3 and at pH 6, respectively. Contributions from the bending region of liquid water were subtracted from the spectrum of the wet TiO₂ film and of the two oxalate solutions. To obtain the pure spectrum of HC₂O₄ ions, contributions of the C₂O₄²⁻ ion in the oxalic acid solution were subtracted. Both spectra (dry, wet) were scaled for the absorbance of an oxalate-free dry TiO₂ solid (reference TiO₂, below) at 600 cm⁻¹. The spectra of HC₂O₄ and C₂O₄²⁻ ions were not scaled. (b) Possible surface metal-bonded (MB) complexes of oxalate on TiO₂ anatase surfaces according to previous studies.[48,50] These complexes can be either protonated or deprotonated. (c) ATR-FTIR spectra of N₂-dry films of TiO₂ with oxalate loadings varied from 0 to 1.4 OX/nm², revealing an increase in loading of oxalate surface complexes. All spectra were scaled with respect to IR absorbance at 600 cm⁻¹. (For interpretation of the references to colour in this figure legend, the reader is referred to the web version of this article.)</figcaption></figure>

buffer solution (8 mM Na₂CO₃/2.5 mM NaHCO₃, Merck) to eluate oxalate ions in the separation column. Elution was performed at flow rate of 0.7 mL/min under a total pressure of 78 bars.

## 2.3. Water loadings by microgravimetry

Water loadings acquired on initially dried pure (unreacted) TiO₂ and TiO₂-oxalate samples were determined by dynamic vapor sorption (DVS), using a DVS Advantage ET 2 instrument (Surface Measurement Systems). These were acquired through a 21-point adsorption/desorption isotherm cycle between 0 and 90 % Relative Humidity (RH) (0.0095–2.84 kPa H₂O) at 25 °C. Measurements were made using ~30–60 mg samples initially dried at ~0 % RH (25 °C) for 10 h to remove adsorbed water from ambient atmosphere. The equilibrium criteria for each preselected RH step were 60 min for steps between 10 and 60 % RH, but 180 min for steps between 70 and 90 % RH. Gravimetric measurements were continuously taken every 1 min, and a complete adsorption–desorption isotherm cycle took up to 48 h. Using the B.E.T. specific surface area of TiO₂, the gravimetrically-determined water loadings were expressed in terms of H₂O/nm². These values were, in turn, expressed in terms of monolayer (ML), assuming that 1 ML is equivalent to 12 H₂O/nm². This value is more typical of water films adsorbed to (hydr)oxo-terminated metal oxide surfaces.[60–62]

## 2.4. Vibrational spectroscopy

Vibrational spectra were collected on a Fourier Transform Infrared (FTIR) spectrometer (Bruker Vertex 70/V) using an Attenuated Total Reflectance (ATR, diamond-bound Golden Gate® GS10563, Specac) accessory in a room kept at 25 ± 1 °C. The instrument was equipped with a deuterated L-alanine doped triglycine sulfate (DLaTGS) detector. Measurements were carried out in the 600–4500 cm⁻¹ range at a resolution of 4 cm⁻¹ and at forward/reverse scanning rate of 10 kHz, and the signals were treated with a Blackman-Harris 3-term apodization function with 16 cm⁻¹ phase resolution and the Mertz phase correction algorithm. Each spectrum was obtained by coadding 100 spectra for collected over a ~89 s period.

Prior to all experiments, an aliquot (5–10 µL) of the pre-equilibrated

TiO₂-oxalate wet paste was deposited on the diamond window of the ATR. It was then dried for 1 h under a flow of dry N₂(g) (~0% RH, 0.0095 kPa H₂O) in a closed-loop flow-through cell. FTIR spectra were continually recorded during this process to ensure that the samples were completely dried. This was confirmed by the loss of liquid water bending (~1640 cm⁻¹) and stretching (~3400 cm⁻¹) modes. This dry solid state TiO₂-oxalate film was then used for the following experiments.

### 2.4.1. Water vapor adsorption in the dark

A water adsorption/desorption isotherm cycle was carried out on the initially dry TiO₂-oxalate solid state film exposed under 500 mL flow of humidified N₂(g) (0 to 90 % RH, corresponding to 0.0095 to 2.84 kPa H₂O at 25 °C) in the dark. Relative humidity was controlled using a humidity generator module (MHG32, proUmid). FTIR spectra were continuously collected during the experiments.

### 2.4.2. Photo-oxidation experiments

Photoreactions were performed under a 500 mL flow of oxygen-rich air (N₂:O₂ = 80:20 % p/p and CO₂ of ~300 ppm) and oxygen-free N₂(g). Initially the N₂(g) dried TiO₂-oxalate dried film was equilibrated with either one of these two gas compositions at a preselected humidity (0, 10, 30, 50, 70 and 90 % RH, controlled by MHG32) at 25 °C for 1 h in the dark. This procedure was used to ensure for water film stability. The samples were then irradiated using 150 W Xe arc lamp (broad UV-visible spectrum covering 200–1000 nm, LSH102, Quantum Design GmbH) through an optical fiber (fused silica, LSZ152, Quantum Design GmbH) connected at a 45° angle into the closed-loop flow-through cell. Although this irradiation angle was not optimized for maximal irradiance flux, it was still sufficient to provide 25 mW/cm² photon energy on the 2 × 2 mm² area of the sample. Photo-oxidation experiments were carried out for minimum of 20 h, and FTIR spectra were collected continuously during all experiments. We note that the temperature of the ATR cell did not increase by more than 1 °C over a 8-h irradiation period. It is thereby unlikely that hikes in light-induced temperatures affected oxalate decomposition.

### 2.4.3. Gaussian deconvolution and kinetic modeling

The 1500–1850 cm⁻¹ region was modeled using a linear
<figure><img src="image_5.png" /><figcaption>**Fig. 2.** Oxalate binding on TiO₂ nanoparticles. (a) ATR-FTIR spectra of wet (blue) and N₂-dry (black) films of TiO₂ with oxalate loadings of 1.4 OX/nm², and ATR-FTIR of references (green): dry TiO₂ film, HC₂O₄ ions and C₂O₄²⁻ ions. The spectrum of the wet TiO₂-oxalate film was obtained on the centrifuged wet paste deposited on ATR surfaces, and the spectrum of the film was taken after this wet paste was dried under N₂(g) (c.f. Section 2.4). The spectra of HC₂O₄ ions and C₂O₄²⁻ ions were obtained from 100 mM oxalate solutions at pH 3.3 and at pH 6, respectively. Contributions from the bending region of liquid water were subtracted from the spectrum of the wet TiO₂ film and of the two oxalate solutions. To obtain the pure spectrum of HC₂O₄ ions, contributions of the C₂O₄²⁻ ion in the oxalic acid solution were subtracted. Both spectra (dry, wet) were scaled for the absorbance of an oxalate-free dry TiO₂ solid (reference TiO₂, below) at 600 cm⁻¹. The spectra of HC₂O₄ and C₂O₄²⁻ ions were not scaled. (b) Possible surface metal-bonded (MB) complexes of oxalate on TiO₂ anatase surfaces according to previous studies.[48,50] These complexes can be either protonated or deprotonated. (c) ATR-FTIR spectra of N₂-dry films of TiO₂ with oxalate loadings varied from 0 to 1.4 OX/nm², revealing an increase in loading of oxalate surface complexes. All spectra were scaled with respect to IR absorbance at 600 cm⁻¹. (For interpretation of the references to colour in this figure legend, the reader is referred to the web version of this article.)</figcaption></figure>

buffer solution (8 mM Na₂CO₃/2.5 mM NaHCO₃, Merck) to eluate oxalate ions in the separation column. Elution was performed at flow rate of 0.7 mL/min under a total pressure of 78 bars.

## 2.3. Water loadings by microgravimetry

Water loadings acquired on initially dried pure (unreacted) TiO₂ and TiO₂-oxalate samples were determined by dynamic vapor sorption (DVS), using a DVS Advantage ET 2 instrument (Surface Measurement Systems). These were acquired through a 21-point adsorption/desorption isotherm cycle between 0 and 90 % Relative Humidity (RH) (0.0095–2.84 kPa H₂O) at 25 °C. Measurements were made using ~30–60 mg samples initially dried at ~0 % RH (25 °C) for 10 h to remove adsorbed water from ambient atmosphere. The equilibrium criteria for each preselected RH step were 60 min for steps between 10 and 60 % RH, but 180 min for steps between 70 and 90 % RH. Gravimetric measurements were continuously taken every 1 min, and a complete adsorption–desorption isotherm cycle took up to 48 h. Using the B.E.T. specific surface area of TiO₂, the gravimetrically-determined water loadings were expressed in terms of H₂O/nm². These values were, in turn, expressed in terms of monolayer (ML), assuming that 1 ML is equivalent to 12 H₂O/nm². This value is more typical of water films adsorbed to (hydr)oxo-terminated metal oxide surfaces.[60–62]

## 2.4. Vibrational spectroscopy

Vibrational spectra were collected on a Fourier Transform Infrared (FTIR) spectrometer (Bruker Vertex 70/V) using an Attenuated Total Reflectance (ATR, diamond-bound Golden Gate® GS10563, Specac) accessory in a room kept at 25 ± 1 °C. The instrument was equipped with a deuterated L-alanine doped triglycine sulfate (DLaTGS) detector. Measurements were carried out in the 600–4500 cm⁻¹ range at a resolution of 4 cm⁻¹ and at forward/reverse scanning rate of 10 kHz, and the signals were treated with a Blackman-Harris 3-term apodization function with 16 cm⁻¹ phase resolution and the Mertz phase correction algorithm. Each spectrum was obtained by coadding 100 spectra for collected over a ~89 s period.

Prior to all experiments, an aliquot (5–10 µL) of the pre-equilibrated

TiO₂-oxalate wet paste was deposited on the diamond window of the ATR. It was then dried for 1 h under a flow of dry N₂(g) (~0% RH, 0.0095 kPa H₂O) in a closed-loop flow-through cell. FTIR spectra were continually recorded during this process to ensure that the samples were completely dried. This was confirmed by the loss of liquid water bending (~1640 cm⁻¹) and stretching (~3400 cm⁻¹) modes. This dry solid state TiO₂-oxalate film was then used for the following experiments.

### 2.4.1. Water vapor adsorption in the dark

A water adsorption/desorption isotherm cycle was carried out on the initially dry TiO₂-oxalate solid state film exposed under 500 mL flow of humidified N₂(g) (0 to 90 % RH, corresponding to 0.0095 to 2.84 kPa H₂O at 25 °C) in the dark. Relative humidity was controlled using a humidity generator module (MHG32, proUmid). FTIR spectra were continuously collected during the experiments.

### 2.4.2. Photo-oxidation experiments

Photoreactions were performed under a 500 mL flow of oxygen-rich air (N₂:O₂ = 80:20 % p/p and CO₂ of ~300 ppm) and oxygen-free N₂(g). Initially the N₂(g) dried TiO₂-oxalate dried film was equilibrated with either one of these two gas compositions at a preselected humidity (0, 10, 30, 50, 70 and 90 % RH, controlled by MHG32) at 25 °C for 1 h in the dark. This procedure was used to ensure for water film stability. The samples were then irradiated using 150 W Xe arc lamp (broad UV-visible spectrum covering 200–1000 nm, LSH102, Quantum Design GmbH) through an optical fiber (fused silica, LSZ152, Quantum Design GmbH) connected at a 45° angle into the closed-loop flow-through cell. Although this irradiation angle was not optimized for maximal irradiance flux, it was still sufficient to provide 25 mW/cm² photon energy on the 2 × 2 mm² area of the sample. Photo-oxidation experiments were carried out for minimum of 20 h, and FTIR spectra were collected continuously during all experiments. We note that the temperature of the ATR cell did not increase by more than 1 °C over a 8-h irradiation period. It is thereby unlikely that hikes in light-induced temperatures affected oxalate decomposition.

### 2.4.3. Gaussian deconvolution and kinetic modeling

The 1500–1850 cm⁻¹ region was modeled using a linear
<figure><img src="image_5.png" /><figcaption>Fig. 3. Response of adsorbed oxalate to water films in the absence of light. (a) FTIR spectra of 1.4 OX/nm² on TiO₂ exposed to humidity from 0 to 90% RH. Spectra show that water bending (~ 1640 cm⁻¹) and oxalate bands concurrently increased in intensity. Moreover, the oxalate C-O/C-C band pair progressively shifted in position (1412 to 1375 cm⁻¹, and 1261 to 1219 cm⁻¹) upon exposure to water vapor, as a result of hydration. (b) Microgravimetry revealed that 1 monolayer (ML) water film formed on the oxalate-free (“pure TiO₂”) and oxalate-bearing (“1.4 OX/nm²”) TiO₂ at 70% RH, and that films grew up to ~ 4 MLs at 90% RH. (c-d) Schematic presentation of (c) partially hydrated TiO₂ and bound oxalate in sub-monolayer (< 1 ML) water coverages, and (d) fully hydrated TiO₂ and oxalate species in thicker water films (≥ 1 ML). Hydration shifted and enhanced the vibrational absorbances of oxalate bands in (a).</figcaption></figure>

combination of Gaussian components:

$$A(v) = \sum_{1}^{n} A_{n,max} e^{-(v-v_n)^2 / (2\sigma_n^2)} \quad (1)$$

with the maximal absorbance ($A_{n,max}$) of the n-th component centered at wavenumber $v_n$, and with the width of the distribution $\sigma_n$ method for our study as attempts using multivariate methods [63] were unsuccessful in clearly discerning the evolution of different bands over reaction time. Spectral deconvolution was performed in a sequence of time-resolved spectra with absorbance values offset to 0 at 1850 cm⁻¹, followed by normalization of all absorbance values to that of the Ti-O bonds at 600 cm⁻¹ to account for differences in sample mass in different experiments, and baseline correction using a simple linear function. All Gaussian and kinetic parameter optimization (non-linear least square) calculations were performed with MATLAB (version R2021b, The Mathworks, Inc.).

## 3. Results and discussion

Resolving the film-driven photocatalytic decomposition of oxalate on TiO₂ using vibrational spectroscopy first requires knowledge of the spectral signatures of bound oxalate species and their response to variations in hydration. To this end, we begin this study by presenting the chief vibrational markers of bound oxalate in the dark. These markers were then used to resolve how variations in water film thickness and dissolved oxygen impacted oxalate photodecomposition rates, efficiency and products.

### 3.1. Oxalate binding on TiO₂ in the dark

Equilibrating aqueous suspensions of TiO₂ with [OX]ₜot = 47.5 mM/g TiO₂ at pH 3 in the dark for 24 h adsorbed ~ 1.4 OX/nm². This loading was lower than the total O site density (e.g., ~15 O/nm² on the (101) plane of anatase), [23] and the population of non-stochiometric surface OH groups (~5 OH/nm²) [64]. From an additional adsorption isotherm (Fig. S2), we also find that this loading was close to the maximal value (1.44 OX/nm²) that can be achieved on the TiO₂ used for this study. This aligns with other previous efforts [65] on a similar type of TiO₂.

#### 3.1.1. Vibrational signatures of bound oxalate

Attenuated Total Reflectance-Fourier Transform Infrared (ATR-FTIR) spectra (Fig. 2 a) of wet centrifuged TiO₂ pastes revealed bands characteristic of metal-bonded (MB) species, in accordance with previous work. [48,49] This can be appreciated by the oxalate C = O

stretching (~ 1718 and ~ 1690 cm⁻¹), C-O or C-C stretching and even the O-C = O bending (1412 and 1261 cm⁻¹) frequencies, which were shifted in relation to the dominant solution HC₂O₄ species [48,49,66] (Fig. 2 a). Previous suggestions for MB species [50,54] include those of Fig. 2 b. Because oxalate forms these very strong MB species on TiO₂ under acidic conditions (pH < pHᵢep ~ 6.5, [47] Fig. S3), outer-sphere complexes (OS) or hydrogen-bonded (HB) species [67] were of comparatively low loadings. Furthermore, previous studies [49,66,68] reported that these OS and HB species have similar spectral profiles to the one of the free C₂O₄²⁻ ion (Fig. 2 a). In short, they can display bands at ~ 1560–1600 cm⁻¹, which possibly overlapped with surface Ti-O-H bending modes (1624 cm⁻¹; Fig. 2 c), as well as bands at ~ 1300 cm⁻¹ which were below the detection limit.

Removing water from the wet TiO₂ paste, using a stream of N₂(g), decreased the intensities of all bands, and especially the intensity ratio between the C = O and C-O/C-C band pair. It shifted the C-O/C-C band pair to lower wavenumbers (Δv = ~ 40–45 cm⁻¹, 1412 to 1375 cm⁻¹ and 1261 to 1219 cm⁻¹), yet it did not greatly affect C = O band positions (1718 to 1724 cm⁻¹). We interpret this to the loss of hydration water molecules, which shifted the C-O bond strength in the water films.

We also investigated the effects of oxalate loadings (Fig. 2 c) on possible loading-dependent spectral responses during photocatalytic decomposition experiments. These experiments showed that band intensities scaled with oxalate loadings, and that bands were not shifted. The 1690 cm⁻¹ band became, however, more intense than the 1724 cm⁻¹ band at oxalate loadings below 0.5 OX/nm². In accordance with previous work [48–50,54], we interpret this as evidence that a distinct MB species [69] (MB1) was preferentially expressed at 1724 cm⁻¹ and that a second complex (MB2) was dominant at 1690 cm⁻¹. These strongly adsorbed species may possibly include those that preferentially bind on the more reactive (001) plane of anatase, yet the spectral profiles were not sufficiently informative to assign these bands to distinct coordination environments (Fig. 2b). [23] Still, we can foresee that variations in loadings could impact MB1/MB2 speciation and, in turn, photocatalytic efficiency. [54]

Exposing these dry oxalate-bearing TiO₂ films to water vapor (10–90% RH) formed water films, which can be seen by FTIR through the growth of O-H stretching (3400–3200 cm⁻¹, Fig. S4) and water bending (~ 1640 cm⁻¹, Fig. 3 a) bands. Spectra of the O-H stretching region (Fig. S4) also revealed free O-H groups (~ 3700 cm⁻¹) from water molecules which were presumably either directly bound to Ti⁴⁺ sites (e. g., on the (100) face) as similar to previous study [60] on water films forming on iron oxides, or from the topmost surface region (i.e. air/water interface) of the water films. [70] From microgravimetry (Fig. 3b),
<figure><img src="image_6.png" /><figcaption>Fig. 3. Response of adsorbed oxalate to water films in the absence of light. (a) FTIR spectra of 1.4 OX/nm² on TiO₂ exposed to humidity from 0 to 90% RH. Spectra show that water bending (~ 1640 cm⁻¹) and oxalate bands concurrently increased in intensity. Moreover, the oxalate C–O/C–C band pair progressively shifted in position (1412 to 1375 cm⁻¹, and 1261 to 1219 cm⁻¹) upon exposure to water vapor, as a result of hydration. (b) Microgravimetry revealed that 1 monolayer (ML) water film formed on the oxalate-free (“pure TiO₂”) and oxalate-bearing (“1.4 OX/nm²”) TiO₂ at 70% RH, and that films grew up to ~ 4 MLs at 90% RH. (c-d) Schematic presentation of (c) partially hydrated TiO₂ and bound oxalate in sub-monolayer (< 1 ML) water coverages, and (d) fully hydrated TiO₂ and oxalate species in thicker water films (≥ 1 ML). Hydration shifted and enhanced the vibrational absorbances of oxalate bands in (a).</figcaption></figure>

combination of Gaussian components:

$$A(v) = \sum_{1}^{n} A_{n,max} e^{-(v-v_n)^2 / (2\sigma_n^2)} \quad (1)$$

with the maximal absorbance ($A_{n,max}$) of the n-th component centered at wavenumber $v_n$, and with the width of the distribution $\sigma_n$ method for our study as attempts using multivariate methods [63] were unsuccessful in clearly discerning the evolution of different bands over reaction time. Spectral deconvolution was performed in a sequence of time-resolved spectra with absorbance values offset to 0 at 1850 cm⁻¹, followed by normalization of all absorbance values to that of the Ti–O bonds at 600 cm⁻¹ to account for differences in sample mass in different experiments, and baseline correction using a simple linear function. All Gaussian and kinetic parameter optimization (non-linear least square) calculations were performed with MATLAB (version R2021b, The Mathworks, Inc.).

## 3. Results and discussion

Resolving the film-driven photocatalytic decomposition of oxalate on TiO₂ using vibrational spectroscopy first requires knowledge of the spectral signatures of bound oxalate species and their response to variations in hydration. To this end, we begin this study by presenting the chief vibrational markers of bound oxalate in the dark. These markers were then used to resolve how variations in water film thickness and dissolved oxygen impacted oxalate photodecomposition rates, efficiency and products.

### 3.1. Oxalate binding on TiO₂ in the dark

Equilibrating aqueous suspensions of TiO₂ with [OX]ₜot = 47.5 mM/g TiO₂ at pH 3 in the dark for 24 h adsorbed ~ 1.4 OX/nm². This loading was lower than the total O site density (e.g., ~15 O/nm² on the (101) plane of anatase), [23] and the population of non-stochiometric surface OH groups (~5 OH/nm²) [64]. From an additional adsorption isotherm (Fig. S2), we also find that this loading was close to the maximal value (1.44 OX/nm²) that can be achieved on the TiO₂ used for this study. This aligns with other previous efforts [65] on a similar type of TiO₂.

#### 3.1.1. Vibrational signatures of bound oxalate

Attenuated Total Reflectance-Fourier Transform Infrared (ATR-FTIR) spectra (Fig. 2 a) of wet centrifuged TiO₂ pastes revealed bands characteristic of metal-bonded (MB) species, in accordance with previous work. [48,49] This can be appreciated by the oxalate C = O

stretching (~ 1718 and ~ 1690 cm⁻¹), C–O or C–C stretching and even the O–C = O bending (1412 and 1261 cm⁻¹) frequencies, which were shifted in relation to the dominant solution HC₂O₄ species [48,49,66] (Fig. 2 a). Previous suggestions for MB species [50,54] include those of Fig. 2 b. Because oxalate forms these very strong MB species on TiO₂ under acidic conditions (pH < pHᵢep ~ 6.5, [47] Fig. S3), outer-sphere complexes (OS) or hydrogen-bonded (HB) species [67] were of comparatively low loadings. Furthermore, previous studies [49,66,68] reported that these OS and HB species have similar spectral profiles to the one of the free C₂O₄²⁻ ion (Fig. 2 a). In short, they can display bands at ~ 1560–1600 cm⁻¹, which possibly overlapped with surface Ti–O–H bending modes (1624 cm⁻¹; Fig. 2 c), as well as bands at ~ 1300 cm⁻¹ which were below the detection limit.

Removing water from the wet TiO₂ paste, using a stream of N₂(g), decreased the intensities of all bands, and especially the intensity ratio between the C = O and C–O/C–C band pair. It shifted the C–O/C–C band pair to lower wavenumbers (Δv = ~ 40–45 cm⁻¹, 1412 to 1375 cm⁻¹ and 1261 to 1219 cm⁻¹), yet it did not greatly affect C = O band positions (1718 to 1724 cm⁻¹). We interpret this to the loss of hydration water molecules, which shifted the C–O bond strength in the water films.

We also investigated the effects of oxalate loadings (Fig. 2 c) on possible loading-dependent spectral responses during photocatalytic decomposition experiments. These experiments showed that band intensities scaled with oxalate loadings, and that bands were not shifted. The 1690 cm⁻¹ band became, however, more intense than the 1724 cm⁻¹ band at oxalate loadings below 0.5 OX/nm². In accordance with previous work [48–50,54], we interpret this as evidence that a distinct MB species [69] (MB1) was preferentially expressed at 1724 cm⁻¹ and that a second complex (MB2) was dominant at 1690 cm⁻¹. These strongly adsorbed species may possibly include those that preferentially bind on the more reactive (001) plane of anatase, yet the spectral profiles were not sufficiently informative to assign these bands to distinct coordination environments (Fig. 2b). [23] Still, we can foresee that variations in loadings could impact MB1/MB2 speciation and, in turn, photocatalytic efficiency. [54]

Exposing these dry oxalate-bearing TiO₂ films to water vapor (10–90% RH) formed water films, which can be seen by FTIR through the growth of O–H stretching (3400–3200 cm⁻¹, Fig. S4) and water bending (~ 1640 cm⁻¹, Fig. 3 a) bands. Spectra of the O–H stretching region (Fig. S4) also revealed free O–H groups (~ 3700 cm⁻¹) from water molecules which were presumably either directly bound to Ti⁴⁺ sites (e. g., on the (100) face) as similar to previous study [60] on water films forming on iron oxides, or from the topmost surface region (i.e. air/water interface) of the water films. [70] From microgravimetry (Fig. 3b),
<figure><img src="image_6.png" /><figcaption>**Fig. 4.** Gaussian deconvolution of the 1510–1780 cm⁻¹ in TiO₂ equilibrated with 1.4 OX/nm². (a) Example of deconvolution of sample exposed to 0 and 90 % RH in the absence of light for 20 h, here revealing four components. (b) Gaussian component intensities, showing congruent growth of the C = O stretching of MB1 (1724 cm⁻¹) and MB2 (1690 cm⁻¹) complexes and water bending (1644 cm⁻¹) bands with humidity. The C = O component of OS complexes (1675 cm⁻¹) remained unchanged regardless of water loadings.</figcaption></figure>

we find that films of less than monolayer (1 ML, ~ 12 H₂O/nm²) formed below 70 % RH, and of up to 4 MLs (~ 48 H₂O/nm²) at 90 % RH (Fig. 3 b). Also, microgravimetry retrieved similar masses of the dry sample before (57.470 ± 0.001 mg) and after (57.541 ± 0.006 mg) adsorption, a result indicating that no oxalate was evaporated or lost due to water adsorption. Hydration progressively and systematically enhanced the intensities and shifted position of the oxalate C–O/C–C band pair back to the positions (1412 cm⁻¹ and 1261 cm⁻¹) of the wet pastes. Additionally, because the positions of the C = O stretching band were not greatly changed, while their intensities enhanced by water loadings, we can infer that sub-monolayer coverages (< 70 % RH, Fig. 3 b) preferentially hydrated bound oxalate surface complexes (Fig. 3 c). Thicker water films (≥ 1 ML), on the other hand, developed full hydration shells on all bound oxalate (Fig. 3 d), and thus generated similar spectroscopic responses to those of the wet pastes (Fig. 2a). These spectroscopic responses were therefore not related to competitive adsorption of water to oxalate. By comparing the IR spectra of dry TiO₂-oxalate film before and

after water vapor adsorption (Fig. S4), and combining with microgravimetry (as described earlier), we conclude that there was no evidence showing that oxalate was desorbed from the water films. This conclusion is also supported by noting that (1) the available reactive crystalline oxygen site density (12.5 O/nm²) [71] is ~ 1 order of magnitude larger than oxalate loadings (1.4 OX/nm²), and (2) the volume of water in 4 ML-thick water films was substantially smaller than in the solutions in which the adsorption experiments were performed (15 mL, cf. Section 2.2).

### 3.1.2. Spectral components of bound oxalate species

To gain further insight into the complexation modes of oxalate and their response to hydration, we resolved the 1510–1780 cm⁻¹ region into four Gaussian components (Fig. 4 a). This revealed (i) the dominant 1724 cm⁻¹ ‘MB1’ component (C = O stretch), (ii) A less dominant ‘MB2’ component at 1690 cm⁻¹ (C = O stretch) [69], (iii) an even less dominant oxalate component at 1675 cm⁻¹ (OS), and finally (iv) a component

<figure><img src="image_7.png" /><figcaption>**Fig. 5.** FTIR spectral evidence for humidity- and oxygen-dependent photocatalytic (broad UV–vis 200–1000 nm, 25 mW/cm²) decomposition of bound oxalate (1.4 OX/nm²) on TiO₂. Reactions were conducted in (a,d) oxygen-rich (21 kPa O₂), and (a,e) oxygen-free atmospheres for 20 h. (a, top) Time-resolved spectra for reactions under 21 kPa O₂ and 90 % RH revealed a decrease in oxalate loadings through the loss of all oxalate bands. (a, bottom) Spectra of experiments under oxygen-free N₂(g) and 90 % RH revealed background changes caused by irradiation, alongside with the loss of oxalate bands. (b) The background increase in (a, bottom) results from the broad 885 cm⁻¹ band of shallow trap electrons (STIRA) within first ~ 60 min (c), which then slowly decayed for the rest of the irradiation period (inset in c and cf. Fig. 6 & S9). This band was revealed by taking the difference spectra from the raw data in (a, bottom). (d,e) Baseline-subtracted spectra showing comparable humidity-dependent losses of oxalate bands under (d) O₂-rich and (e) O₂-free conditions. Difference spectra showing the 885 cm⁻¹ in all O₂-free experiments are accessible in Supporting Information, Fig. S9.</figcaption></figure>
<figure><img src="image_7.png" /><figcaption>**Fig. 4.** Gaussian deconvolution of the 1510–1780 cm⁻¹ in TiO₂ equilibrated with 1.4 OX/nm². (a) Example of deconvolution of sample exposed to 0 and 90 % RH in the absence of light for 20 h, here revealing four components. (b) Gaussian component intensities, showing congruent growth of the C = O stretching of MB1 (1724 cm⁻¹) and MB2 (1690 cm⁻¹) complexes and water bending (1644 cm⁻¹) bands with humidity. The C = O component of OS complexes (1675 cm⁻¹) remained unchanged regardless of water loadings.</figcaption></figure>

we find that films of less than monolayer (1 ML, ~ 12 H₂O/nm²) formed below 70 % RH, and of up to 4 MLs (~ 48 H₂O/nm²) at 90 % RH (Fig. 3 b). Also, microgravimetry retrieved similar masses of the dry sample before (57.470 ± 0.001 mg) and after (57.541 ± 0.006 mg) adsorption, a result indicating that no oxalate was evaporated or lost due to water adsorption. Hydration progressively and systematically enhanced the intensities and shifted position of the oxalate C–O/C–C band pair back to the positions (1412 cm⁻¹ and 1261 cm⁻¹) of the wet pastes. Additionally, because the positions of the C = O stretching band were not greatly changed, while their intensities enhanced by water loadings, we can infer that sub-monolayer coverages (< 70 % RH, Fig. 3 b) preferentially hydrated bound oxalate surface complexes (Fig. 3 c). Thicker water films (≥ 1 ML), on the other hand, developed full hydration shells on all bound oxalate (Fig. 3 d), and thus generated similar spectroscopic responses to those of the wet pastes (Fig. 2a). These spectroscopic responses were therefore not related to competitive adsorption of water to oxalate. By comparing the IR spectra of dry TiO₂-oxalate film before and

after water vapor adsorption (Fig. S4), and combining with microgravimetry (as described earlier), we conclude that there was no evidence showing that oxalate was desorbed from the water films. This conclusion is also supported by noting that (1) the available reactive crystalline oxygen site density (12.5 O/nm²) [71] is ~ 1 order of magnitude larger than oxalate loadings (1.4 OX/nm²), and (2) the volume of water in 4 ML-thick water films was substantially smaller than in the solutions in which the adsorption experiments were performed (15 mL, cf. Section 2.2).

### 3.1.2. Spectral components of bound oxalate species

To gain further insight into the complexation modes of oxalate and their response to hydration, we resolved the 1510–1780 cm⁻¹ region into four Gaussian components (Fig. 4 a). This revealed (i) the dominant 1724 cm⁻¹ ‘MB1’ component (C = O stretch), (ii) A less dominant ‘MB2’ component at 1690 cm⁻¹ (C = O stretch) [69], (iii) an even less dominant oxalate component at 1675 cm⁻¹ (OS), and finally (iv) a component

<figure><img src="image_8.png" /><figcaption>**Fig. 5.** FTIR spectral evidence for humidity- and oxygen-dependent photocatalytic (broad UV–vis 200–1000 nm, 25 mW/cm²) decomposition of bound oxalate (1.4 OX/nm²) on TiO₂. Reactions were conducted in (a,d) oxygen-rich (21 kPa O₂), and (a,e) oxygen-free atmospheres for 20 h. (a, top) Time-resolved spectra for reactions under 21 kPa O₂ and 90 % RH revealed a decrease in oxalate loadings through the loss of all oxalate bands. (a, bottom) Spectra of experiments under oxygen-free N₂(g) and 90 % RH revealed background changes caused by irradiation, alongside with the loss of oxalate bands. (b) The background increase in (a, bottom) results from the broad 885 cm⁻¹ band of shallow trap electrons (STIRA) within first ~ 60 min (c), which then slowly decayed for the rest of the irradiation period (inset in c and cf. Fig. 6 & S9). This band was revealed by taking the difference spectra from the raw data in (a, bottom). (d,e) Baseline-subtracted spectra showing comparable humidity-dependent losses of oxalate bands under (d) O₂-rich and (e) O₂-free conditions. Difference spectra showing the 885 cm⁻¹ in all O₂-free experiments are accessible in Supporting Information, Fig. S9.</figcaption></figure>
<figure><img src="image_8.png" /><figcaption>Fig. 6. Time-dependent absorbances revealing oxalate photodecomposition. (a,b) Normalized Gaussian component absorbances ($A_u/A_0$; Eqn. (13)) of the C = O band of MB1 (1724 cm⁻¹) during photoreactions over 60 min periods (cf. Fig. S10 for entire 20 h period) under (a) O₂-bearing (21 kPa O₂) and (b) O₂-free (0 kPa O₂) atmospheres over a range of humidity (0–90 % RH). (c) Time-dependent STIRAs (885 cm⁻¹ band) for the first 60 min of the reaction (cf. Fig. S9 for entire 20 h period). (d,e) Reaction half-lives of MB1 and (f,g) percent of photoconversion after a 20 h period, derived from the data in (a,b) for MB1 and in Fig. S10 for MB2 and OS.</figcaption></figure>

at 1644 ± 3 cm⁻¹ from overlapping contributions of Ti–O–H bending, water bending, and unresolved oxalate species. D₂O exchange (Fig. S5) displaced the water band and the Ti–O–H band but left a small oxalate component. Deconvolution also successfully separated contributions of the water and OH-related bending mode at ~ 1644 cm⁻¹ from the components at 1690 and 1675 cm⁻¹ at elevated humidity (Fig. 4 a,b). This operation will become useful for teasing out unwanted contributions generated from the photodecomposition at lower oxalate loadings (Fig. S6).

While the band intensities of the MB1 and MB2 species scaled with humidity, those of another oxalate component (1675 cm⁻¹) did not respond (Fig. 4 b & S6). One possibility is that this species retained its full hydration environment. For this reason, we assigned this component to a hydrated OS complex, an interpretation that could possibly explain previous work. [49] Additionally, from the total Gaussian absorbances, we estimate that total oxalate loadings of 1.4 OX/nm² were 63 % MB1, 16 % MB2, and 21 % OS species.

## 3.2. Oxygen- and humidity-dependent photoreactions

Using these spectroscopic markers for oxalate speciation, we resolved photodecomposition reactions as oxalate-bearing TiO₂ nano-particles were exposed to a broad spectrum (200–1000 nm) of UV–visible light over time. Photo-experiments in this study will be explored in O₂-bearing and O₂-free sub-monolayer to multilayered water films (Figs. 5 and 6) to uncover how these conditions impacted the fate of oxalate.

### 3.2.1. Oxygen-dependent photoreactions

The vibrational spectra revealed that exposing 4-ML water films (90 % RH) to 21 kPa O₂ and UV light triggered photodecomposition

reactions, detected through the loss of oxalate bands (Fig. 5a). No major band shifts occurred during the entire 20 h of the reaction, indicating that photodecomposition did not generate intermediate products (e.g., formic acid), as previously shown. [44,53] This situation thus contrasts with experiments in a liquid cell under flow-through conditions [53] where species re-adsorbed on TiO₂ contributed to oxalate bands growth during irradiation. Comparing these spectra with those of lower oxalate loadings under hydration (0–90 % RH, no light, Fig. 2c & S6) confirmed that the spectral response was still from oxalate.

These results consequently support the interpretation that the photodecomposition of oxalate MB complexes producing CO₂ in multilayer water films begins with the well-known [72] production of •OH radicals. These reactions begin with:

$$ \mathrm{TiO_2} + \text{light} \rightarrow e^- + h^+ \qquad (2) $$

then proceed through surface hydroxo groups:

$$ \equiv \mathrm{Ti}^{IV}-\mathrm{OH} + h^+ \rightarrow \equiv \mathrm{Ti}^{IV}-\cdot \mathrm{OH}^+ \qquad (3a) $$

or chemisorbed water:

$$ \equiv \mathrm{Ti}^{IV}-\mathrm{OH}_2^+ + h^+ \rightarrow \equiv \mathrm{Ti}^{IV}-\cdot \mathrm{OH}^+ + \mathrm{H}^+ \qquad (3b) $$

These radicals could also desorb into the multilayered water films:

$$ \equiv \mathrm{Ti}^{IV}-\cdot \mathrm{OH}^+ \rightarrow \equiv \mathrm{Ti}^{IV} + \cdot \mathrm{OH} (\mathrm{aq}) \qquad (4) $$

where they could, in turn, drive solvent-hosted reactions.

Both hole and the •OH radical can oxidize adsorbed oxalate (e.g., $\equiv \mathrm{Ti}^{IV}-\mathrm{C}_2\mathrm{O}_4^{2-}$) via [44]:

$$ \equiv \mathrm{Ti}^{IV}-\mathrm{C}_2\mathrm{O}_4^{2-} + h^+ \rightarrow \equiv \mathrm{Ti}^{IV}-\mathrm{C}_2\mathrm{O}_4^{-} \qquad (5) $$

$$ \equiv \mathrm{Ti}^{IV}-\mathrm{C}_2\mathrm{O}_4^{-} + \cdot \mathrm{OH} + \mathrm{H}^+ \rightarrow \equiv \mathrm{Ti}^{IV} + 2\mathrm{CO}_2 + \mathrm{H}_2\mathrm{O} \qquad (6) $$
<figure><img src="image_9.png" /><figcaption>Fig. 6. Time-dependent absorbances revealing oxalate photodecomposition. (a,b) Normalized Gaussian component absorbances ($A_u/A_0$; Eqn. (13)) of the C = O band of MB1 (1724 cm⁻¹) during photoreactions over 60 min periods (cf. Fig. S10 for entire 20 h period) under (a) O₂-bearing (21 kPa O₂) and (b) O₂-free (0 kPa O₂) atmospheres over a range of humidity (0–90 % RH). (c) Time-dependent STIRAs (885 cm⁻¹ band) for the first 60 min of the reaction (cf. Fig. S9 for entire 20 h period). (d,e) Reaction half-lives of MB1 and (f,g) percent of photoconversion after a 20 h period, derived from the data in (a,b) for MB1 and in Fig. S10 for MB2 and OS.</figcaption></figure>

at 1644 ± 3 cm⁻¹ from overlapping contributions of Ti–O–H bending, water bending, and unresolved oxalate species. D₂O exchange (Fig. S5) displaced the water band and the Ti–O–H band but left a small oxalate component. Deconvolution also successfully separated contributions of the water and OH-related bending mode at ~ 1644 cm⁻¹ from the components at 1690 and 1675 cm⁻¹ at elevated humidity (Fig. 4 a,b). This operation will become useful for teasing out unwanted contributions generated from the photodecomposition at lower oxalate loadings (Fig. S6).

While the band intensities of the MB1 and MB2 species scaled with humidity, those of another oxalate component (1675 cm⁻¹) did not respond (Fig. 4 b & S6). One possibility is that this species retained its full hydration environment. For this reason, we assigned this component to a hydrated OS complex, an interpretation that could possibly explain previous work. [49] Additionally, from the total Gaussian absorbances, we estimate that total oxalate loadings of 1.4 OX/nm² were 63 % MB1, 16 % MB2, and 21 % OS species.

## 3.2. Oxygen- and humidity-dependent photoreactions

Using these spectroscopic markers for oxalate speciation, we resolved photodecomposition reactions as oxalate-bearing TiO₂ nano-particles were exposed to a broad spectrum (200–1000 nm) of UV–visible light over time. Photo-experiments in this study will be explored in O₂-bearing and O₂-free sub-monolayer to multilayered water films (Figs. 5 and 6) to uncover how these conditions impacted the fate of oxalate.

### 3.2.1. Oxygen-dependent photoreactions

The vibrational spectra revealed that exposing 4-ML water films (90 % RH) to 21 kPa O₂ and UV light triggered photodecomposition

reactions, detected through the loss of oxalate bands (Fig. 5a). No major band shifts occurred during the entire 20 h of the reaction, indicating that photodecomposition did not generate intermediate products (e.g., formic acid), as previously shown. [44,53] This situation thus contrasts with experiments in a liquid cell under flow-through conditions [53] where species re-adsorbed on TiO₂ contributed to oxalate bands growth during irradiation. Comparing these spectra with those of lower oxalate loadings under hydration (0–90 % RH, no light, Fig. 2c & S6) confirmed that the spectral response was still from oxalate.

These results consequently support the interpretation that the photodecomposition of oxalate MB complexes producing CO₂ in multilayer water films begins with the well-known [72] production of •OH radicals. These reactions begin with:

$$ \mathrm{TiO_2} + \text{light} \rightarrow e^- + h^+ \qquad (2) $$

then proceed through surface hydroxo groups:

$$ \equiv \mathrm{Ti^{IV}-OH} + h^+ \rightarrow \equiv \mathrm{Ti^{IV}-\bullet OH}^+ \qquad (3a) $$

or chemisorbed water:

$$ \equiv \mathrm{Ti^{IV}-OH_2}^+ + h^+ \rightarrow \equiv \mathrm{Ti^{IV}-\bullet OH}^+ + \mathrm{H}^+ \qquad (3b) $$

These radicals could also desorb into the multilayered water films:

$$ \equiv \mathrm{Ti^{IV}-\bullet OH}^+ \rightarrow \equiv \mathrm{Ti^{IV}} + \bullet OH (aq) \qquad (4) $$

where they could, in turn, drive solvent-hosted reactions.

Both hole and the •OH radical can oxidize adsorbed oxalate (e.g., $\equiv \mathrm{Ti^{IV}-C_2O_4}^-$) via [44]:

$$ \equiv \mathrm{Ti^{IV}-C_2O_4}^+ + h^+ \rightarrow \equiv \mathrm{Ti^{IV}-C_2O_4}^- \qquad (5) $$

$$ \equiv \mathrm{Ti^{IV}-C_2O_4}^+ + \bullet OH + \mathrm{H}^+ \rightarrow \equiv \mathrm{Ti^{IV}} + 2CO_2 + \mathrm{H}_2O \qquad (6) $$
thus regenerating the reactive ≡TiIV sites after CO₂ generation.

Of note, the presence of oxygen plays an important role as an electron scavenger, and generates a variety of (adsorbed) ROS species through [72]:

$$O_2 + e^- \rightarrow O_2^\bullet^- \qquad (7)$$

$$O_2^\bullet^- + H^+ \rightarrow HO_2^\bullet \qquad (8)$$

$$O_2^\bullet^- + 2H^+ + e^- \rightarrow H_2O_2 \qquad (9)$$

$$H_2O_2 + e^- \rightarrow \bullet OH + OH^- \qquad (10)$$

These ROS can, in turn, oxidize adsorbed oxalate, for example, through [44]:

$$\equiv \mathrm{Ti}^{\mathrm{IV}}-\mathrm{C}_{2} \mathrm{O}_{4}^{-}+\mathrm{HO}_{2}^{\bullet}-\equiv \mathrm{Ti}^{\mathrm{IV}}-\mathrm{CO}_{3}^{-}+\mathrm{HCO}_{3}^{\bullet} \qquad (11)$$

as we foresee that HO₂ should be the dominant species in the acidic water films considered in this work (O₂̂/ HO₂, pKa = 4.8). We also note that ROS can decrease the rate of electron-hole recombination. [72]

Reactions in the absence of O₂ (Fig. 5a) generated, in contrast, a baseline increase peaking at 885 cm⁻¹ (Fig. 5b,c) over the course of the first ~ 30 min of irradiation. The baseline then decayed over the rest of the reaction period (Fig. 5b,c). We attribute this baseline to a Shallow Trap IR Absorbance (STIRA). [73] The origin of the STIRA is attributed to variations in Ti–O bond strengths in the surface and in the bulk, which are caused by the reduction of TiIV sites through:

$$\mathrm{Ti}^{\mathrm{IV}}-\mathrm{O}^{-}+\mathrm{e}^{-} \rightarrow \mathrm{Ti}^{\mathrm{III}}-\mathrm{O}^{-} \qquad (12)$$

as well as by charge-compensating protonation reactions of the type [73,74]:

$$\mathrm{Ti}^{\mathrm{III}}-\mathrm{O}^{-}+\mathrm{H}^{+} \rightarrow \mathrm{Ti}^{\mathrm{III}}-\mathrm{OH} \qquad (13)$$

The latter reaction is facilitated in proton-rich samples, such as those considered for this work.

We take the subsequent decay of the STIRA during irradiation as an indication that trapped electrons ultimately recombined with holes:

$$\mathrm{Ti}^{\mathrm{III}}-\mathrm{OH}+\mathrm{h}^{+} \rightarrow \mathrm{Ti}^{\mathrm{IV}}-\mathrm{O}^{-}+\mathrm{H}^{+} \qquad (14)$$

From the time-resolved STIRA (inset of Fig. 5c, S9), this recombination can take place over the course of at least 19 h.

### 3.2.2. Humidity-dependent photoreactions

To test the water film thickness dependence on photoreaction yields, we performed another set of experiments at 10–90 % RH, again in the presence (Fig. 5d & S8) and in the absence of oxygen (Fig. 5e & S8). To compare the spectral response of oxalate in the presence and absence of O₂, we extracted all STIRA-generated baselines (Fig. S7) from the spectra.

These experiments revealed that water film thickness played (i) no strong effects in O₂-bearing but (ii) important effects in O₂-free systems. Also, from the intensities of the 1724 cm⁻¹ bands after 20 h of irradiation (Fig. 5d & S8), we find that O₂-bearing films decomposed ~ 80–90 % of the oxalate, irrespective of film thickness. On the other hand, while this level of photodecomposition was also achieved in the absence of O₂ on dry TiO₂ (Fig. 5e & S8) and in sub-ML water films (10–30 % RH), thicker films formed at and above 50 % RH decomposed only ~ 30–50 % of the oxalate.

### 3.3. Kinetics of photoreactions

The time-dependent intensities of the Gaussian components were used to gain insight into the kinetics of the photocatalytic decomposition reactions for oxalate (Fig. 6a,b), and to track the STIRAs (Fig. 6c). Preliminary analyses on the MB1 (1724 cm⁻¹) and the C–O/C–C band pair (Fig. S10) revealed highly comparable responses. Because we could not extract other contributions of the 1690/1675 cm⁻¹ pair from the C–O/

C–C bands, we focused our analyses on three dominant components of the C = O stretching region (1724, 1690 and 1675 cm⁻¹). Furthermore, because absorbances of the 1724 and 1690 cm⁻¹ scaled with humidity (Fig. 4 b), we normalized their absorbances to those at the onset of the reaction (A₀) at each humidity (Fig. 6 a,b & S10).

Kinetic modeling of these time-dependent values (Aₜ/A₀) showed that the reactions did not follow a simple first order kinetic (Fig. S11). As a result, we focused our discussion on the time-resolve absorbance profiles of these bands, as well as their half-lives (Fig. 6 d,e; Aₜ/A₀= ½ from Fig. 6a,b), as they can still provide insights into reaction kinetics. We also derive photoreaction yields (Fig. 6 f,g) from these absorbance profiles. We begin our discussion on the dominant MB1 (1724 cm⁻¹) as it represented ~ 63 % of total oxalate that undergone humidity-dependent photodecomposition in the presence and absence of O₂. We thereafter highlight the influence of complexation modes of oxalate on photocatalytic efficiency in water films.

#### 3.3.1. O₂-rich water films

To tease out the specific contributions of water on the photoreactions, we begin by highlighting the intrinsic properties of dry TiO₂ on the photodecomposition of oxalate. Photoreactions in the absence of water (0 ML, 0 % RH) in 21 kPa O₂ effectively decomposed ~ 80 % of the bound oxalate after 20 h, and with a half-life of ~ 3 h (Fig. 6d). We explain this high photodecomposition yield by direct hole transfer (HT) from TiO₂ to the bound oxalate:

$$\equiv \mathrm{Ti}^{\mathrm{IV}}-\mathrm{C}_{2} \mathrm{O}_{4}^{-}+\mathrm{h}^{+} \rightarrow \equiv \mathrm{Ti}^{\mathrm{IV}}-\mathrm{C}_{2} \mathrm{O}_{4}^{-} \qquad (15)$$

$$\equiv \mathrm{Ti}^{\mathrm{IV}}-\mathrm{C}_{2} \mathrm{O}_{4}^{-}+2 \mathrm{h}^{+} \rightarrow \equiv \mathrm{Ti}^{\mathrm{IV}}-\mathrm{C}_{2} \mathrm{O}_{4}^{\bullet+} \qquad (16)$$

$$\equiv \mathrm{Ti}^{\mathrm{IV}}-\mathrm{C}_{2} \mathrm{O}_{4}^{\bullet+} \rightarrow \equiv \mathrm{Ti}^{\mathrm{IV}}-\mathrm{CO}_{2}^{\bullet+}+\mathrm{CO}_{2} \qquad (17)$$

$$\equiv \mathrm{Ti}^{\mathrm{IV}}-\mathrm{CO}_{2}^{\bullet+}+\mathrm{O}_{2}^{\bullet-} \rightarrow \equiv \mathrm{Ti}^{\mathrm{IV}}+\mathrm{CO}_{2}+\mathrm{O}_{2} \qquad (18)$$

Note that superoxide (O₂̂) in Eq. 18 is a result of direct reduction of oxygen via Eq. (7). Additionally, irradiation of TiO₂ bound oxalate facilitated electron transfer from oxalate to the TiO₂ conduction band through ligand-to-metal charge transfer (LMCT) and generated CB electrons (e⁻) and the oxalate radical anion (C₂O₄⁻). The C₂O₄⁻ radical anion then spontaneously decomposed into CO₂ and a CO₂⁻ radical anion, which in turn reacted with oxygen:

$$\equiv \mathrm{Ti}^{\mathrm{IV}}-\mathrm{C}_{2} \mathrm{O}_{4}^{\bullet-} \rightarrow \equiv \mathrm{Ti}^{\mathrm{III}}-\mathrm{C}_{2} \mathrm{O}_{4}^{\bullet-} \qquad (19)$$

$$\equiv \mathrm{Ti}^{\mathrm{III}}-\mathrm{C}_{2} \mathrm{O}_{4}^{\bullet-} \rightarrow \equiv \mathrm{Ti}^{\mathrm{III}}-\mathrm{CO}_{2}^{\bullet-}+\mathrm{CO}_{2} \qquad (20)$$

$$\equiv \mathrm{Ti}^{\mathrm{III}}-\mathrm{CO}_{2}^{\bullet-}+2 \mathrm{O}_{2} \rightarrow \equiv \mathrm{Ti}^{\mathrm{IV}}+\mathrm{CO}_{2}+2 \mathrm{O}_{2}^{\bullet-} \qquad (21)$$

The generated O₂⁻ can then proceed through Eqs. (8)–(10).

Photoreaction yields (Fig. 6 f) directly scaled with humidity, increasing from ~ 80 % conversion at 0 % RH to full decomposition at 70 % RH (1 ML), and then reversing back to ~ 80 % conversion at 90 % RH (4 MLs). Reaction half-lives (Fig. 6 d) mirrored this trend by decreasing, in turn, from ~ 120 min down to ~ 28 min at 70 % RH and slightly increasing to ~ 36 min at 90 % RH. In addition to the HT/LCMT processes and superoxide/hydroperoxyl radical production, we interpret the water loading-dependent trends in yield and half-life to arise from the water film-mediated production of •OH radicals [72] in films of at least ~ 0.8 ML via Eqs. (3a) and (4).

It is thus this synergistic HT/LMCT, ROS production, and water-assisted •OH radical production that explains the faster reaction rates (shorter half-lives) with water film loadings. Additionally, it is possible that the decreased acidity of thicker water films produced less H₂O₂, and therefore lowered •OH radical concentrations (Eqs. (9)–(10)). This can consequently explain the minor reversal in trend at 90 % RH. Furthermore, because the diffusion of dissolved oxygen to TiO₂ surface is likely to be slower in thicker water films, it also can lower ROS production (Eqs. (7)–(10)) and consequently slow down oxalate photodecomposition.
thus regenerating the reactive ≡TiIV sites after CO₂ generation.

Of note, the presence of oxygen plays an important role as an electron scavenger, and generates a variety of (adsorbed) ROS species through [72]:

$$O_2 + e^- \rightarrow O_2^\bullet^- \qquad (7)$$

$$O_2^\bullet^- + H^+ \rightarrow HO_2^\bullet \qquad (8)$$

$$O_2^\bullet^- + 2H^+ + e^- \rightarrow H_2O_2 \qquad (9)$$

$$H_2O_2 + e^- \rightarrow \bullet OH + OH^- \qquad (10)$$

These ROS can, in turn, oxidize adsorbed oxalate, for example, through [44]:

$$\equiv \mathrm{Ti}^{\mathrm{IV}}-\mathrm{C}_{2} \mathrm{O}_{4}^{-}+\mathrm{HO}_{2}^{\bullet}-\equiv \mathrm{Ti}^{\mathrm{IV}}-\mathrm{CO}_{3}^{-}+\mathrm{HCO}_{3}^{\bullet} \qquad (11)$$

as we foresee that HO₂ should be the dominant species in the acidic water films considered in this work (O₂̂/ HO₂, pKa = 4.8). We also note that ROS can decrease the rate of electron-hole recombination. [72]

Reactions in the absence of O₂ (Fig. 5a) generated, in contrast, a baseline increase peaking at 885 cm⁻¹ (Fig. 5b,c) over the course of the first ~ 30 min of irradiation. The baseline then decayed over the rest of the reaction period (Fig. 5b,c). We attribute this baseline to a Shallow Trap IR Absorbance (STIRA). [73] The origin of the STIRA is attributed to variations in Ti–O bond strengths in the surface and in the bulk, which are caused by the reduction of TiIV sites through:

$$\mathrm{Ti}^{\mathrm{IV}}-\mathrm{O}^{-}+\mathrm{e}^{-} \rightarrow \mathrm{Ti}^{\mathrm{III}}-\mathrm{O}^{-} \qquad (12)$$

as well as by charge-compensating protonation reactions of the type [73,74]:

$$\mathrm{Ti}^{\mathrm{III}}-\mathrm{O}^{-}+\mathrm{H}^{+} \rightarrow \mathrm{Ti}^{\mathrm{III}}-\mathrm{OH} \qquad (13)$$

The latter reaction is facilitated in proton-rich samples, such as those considered for this work.

We take the subsequent decay of the STIRA during irradiation as an indication that trapped electrons ultimately recombined with holes:

$$\mathrm{Ti}^{\mathrm{III}}-\mathrm{OH}+\mathrm{h}^{+} \rightarrow \mathrm{Ti}^{\mathrm{IV}}-\mathrm{O}^{-}+\mathrm{H}^{+} \qquad (14)$$

From the time-resolved STIRA (inset of Fig. 5c, S9), this recombination can take place over the course of at least 19 h.

### 3.2.2. Humidity-dependent photoreactions

To test the water film thickness dependence on photoreaction yields, we performed another set of experiments at 10–90 % RH, again in the presence (Fig. 5d & S8) and in the absence of oxygen (Fig. 5e & S8). To compare the spectral response of oxalate in the presence and absence of O₂, we extracted all STIRA-generated baselines (Fig. S7) from the spectra.

These experiments revealed that water film thickness played (i) no strong effects in O₂-bearing but (ii) important effects in O₂-free systems. Also, from the intensities of the 1724 cm⁻¹ bands after 20 h of irradiation (Fig. 5d & S8), we find that O₂-bearing films decomposed ~ 80–90 % of the oxalate, irrespective of film thickness. On the other hand, while this level of photodecomposition was also achieved in the absence of O₂ on dry TiO₂ (Fig. 5e & S8) and in sub-ML water films (10–30 % RH), thicker films formed at and above 50 % RH decomposed only ~ 30–50 % of the oxalate.

### 3.3. Kinetics of photoreactions

The time-dependent intensities of the Gaussian components were used to gain insight into the kinetics of the photocatalytic decomposition reactions for oxalate (Fig. 6a,b), and to track the STIRAs (Fig. 6c). Preliminary analyses on the MB1 (1724 cm⁻¹) and the C–O/C–C band pair (Fig. S10) revealed highly comparable responses. Because we could not extract other contributions of the 1690/1675 cm⁻¹ pair from the C–O/

C–C bands, we focused our analyses on three dominant components of the C = O stretching region (1724, 1690 and 1675 cm⁻¹). Furthermore, because absorbances of the 1724 and 1690 cm⁻¹ scaled with humidity (Fig. 4 b), we normalized their absorbances to those at the onset of the reaction (A₀) at each humidity (Fig. 6 a,b & S10).

Kinetic modeling of these time-dependent values (Aₜ/A₀) showed that the reactions did not follow a simple first order kinetic (Fig. S11). As a result, we focused our discussion on the time-resolve absorbance profiles of these bands, as well as their half-lives (Fig. 6 d,e; Aₜ/A₀= ½ from Fig. 6a,b), as they can still provide insights into reaction kinetics. We also derive photoreaction yields (Fig. 6 f,g) from these absorbance profiles. We begin our discussion on the dominant MB1 (1724 cm⁻¹) as it represented ~ 63 % of total oxalate that undergone humidity-dependent photodecomposition in the presence and absence of O₂. We thereafter highlight the influence of complexation modes of oxalate on photocatalytic efficiency in water films.

#### 3.3.1. O₂-rich water films

To tease out the specific contributions of water on the photoreactions, we begin by highlighting the intrinsic properties of dry TiO₂ on the photodecomposition of oxalate. Photoreactions in the absence of water (0 ML, 0 % RH) in 21 kPa O₂ effectively decomposed ~ 80 % of the bound oxalate after 20 h, and with a half-life of ~ 3 h (Fig. 6d). We explain this high photodecomposition yield by direct hole transfer (HT) from TiO₂ to the bound oxalate:

$$\equiv \mathrm{Ti}^{\mathrm{IV}}-\mathrm{C}_{2} \mathrm{O}_{4}^{-}+\mathrm{h}^{+} \rightarrow \equiv \mathrm{Ti}^{\mathrm{IV}}-\mathrm{C}_{2} \mathrm{O}_{4}^{-} \qquad (15)$$

$$\equiv \mathrm{Ti}^{\mathrm{IV}}-\mathrm{C}_{2} \mathrm{O}_{4}^{-}+2 \mathrm{h}^{+} \rightarrow \equiv \mathrm{Ti}^{\mathrm{IV}}-\mathrm{C}_{2} \mathrm{O}_{4}^{\bullet+} \qquad (16)$$

$$\equiv \mathrm{Ti}^{\mathrm{IV}}-\mathrm{C}_{2} \mathrm{O}_{4}^{\bullet+} \rightarrow \equiv \mathrm{Ti}^{\mathrm{IV}}-\mathrm{CO}_{2}^{\bullet+}+\mathrm{CO}_{2} \qquad (17)$$

$$\equiv \mathrm{Ti}^{\mathrm{IV}}-\mathrm{CO}_{2}^{\bullet+}+\mathrm{O}_{2}^{\bullet-} \rightarrow \equiv \mathrm{Ti}^{\mathrm{IV}}+\mathrm{CO}_{2}+\mathrm{O}_{2} \qquad (18)$$

Note that superoxide (O₂̂) in Eq. 18 is a result of direct reduction of oxygen via Eq. (7). Additionally, irradiation of TiO₂ bound oxalate facilitated electron transfer from oxalate to the TiO₂ conduction band through ligand-to-metal charge transfer (LMCT) and generated CB electrons (e⁻) and the oxalate radical anion (C₂O₄⁻). The C₂O₄⁻ radical anion then spontaneously decomposed into CO₂ and a CO₂⁻ radical anion, which in turn reacted with oxygen:

$$\equiv \mathrm{Ti}^{\mathrm{IV}}-\mathrm{C}_{2} \mathrm{O}_{4}^{\bullet-} \rightarrow \equiv \mathrm{Ti}^{\mathrm{III}}-\mathrm{C}_{2} \mathrm{O}_{4}^{\bullet-} \qquad (19)$$

$$\equiv \mathrm{Ti}^{\mathrm{III}}-\mathrm{C}_{2} \mathrm{O}_{4}^{\bullet-} \rightarrow \equiv \mathrm{Ti}^{\mathrm{III}}-\mathrm{CO}_{2}^{-}+\mathrm{CO}_{2} \qquad (20)$$

$$\equiv \mathrm{Ti}^{\mathrm{III}}-\mathrm{CO}_{2}^{\bullet-}+2 \mathrm{O}_{2} \rightarrow \equiv \mathrm{Ti}^{\mathrm{IV}}+\mathrm{CO}_{2}+2 \mathrm{O}_{2}^{\bullet-} \qquad (21)$$

The generated O₂⁻ can then proceed through Eqs. (8)–(10).

Photoreaction yields (Fig. 6 f) directly scaled with humidity, increasing from ~ 80 % conversion at 0 % RH to full decomposition at 70 % RH (1 ML), and then reversing back to ~ 80 % conversion at 90 % RH (4 MLs). Reaction half-lives (Fig. 6 d) mirrored this trend by decreasing, in turn, from ~ 120 min down to ~ 28 min at 70 % RH and slightly increasing to ~ 36 min at 90 % RH. In addition to the HT/LCMT processes and superoxide/hydroperoxyl radical production, we interpret the water loading-dependent trends in yield and half-life to arise from the water film-mediated production of •OH radicals [72] in films of at least ~ 0.8 ML via Eqs. (3a) and (4).

It is thus this synergistic HT/LMCT, ROS production, and water-assisted •OH radical production that explains the faster reaction rates (shorter half-lives) with water film loadings. Additionally, it is possible that the decreased acidity of thicker water films produced less H₂O₂, and therefore lowered •OH radical concentrations (Eqs. (9)–(10)). This can consequently explain the minor reversal in trend at 90 % RH. Furthermore, because the diffusion of dissolved oxygen to TiO₂ surface is likely to be slower in thicker water films, it also can lower ROS production (Eqs. (7)–(10)) and consequently slow down oxalate photodecomposition.
### 3.3.2. *O₂*-free water films

Reactions in the absence of O₂ had contrasting humidity-dependent trends in half-lives (Fig. 6e) and reaction yields (Fig. 6g). Firstly, yields were on par with those of O₂-bearing films with up to 0.5 ML (30 % RH), but thereafter dramatically dropped in thicker films. Half-lives were, in contrast, comparable on dry TiO₂ (~ 2 h), then rose to ~ 15 h in the 0.8 ML (50 % RH) film, and dropped to ~ 3 h in 4 ML-thick films.

From the relatively constant and comparable rates in dry TiO₂ in the presence and absence of O₂, we propose that LMCT and HT were the chief mechanisms for oxalate photodecomposition. The hike in half-lives in O₂-free water films (10–90 % RH) signaled, on the other hand, that films could not sustain an important production of *OH radicals compared to O₂-bearing films. We explain the drop in rates in O₂-free water films (10–90 % RH) using two primary processes: (i) the suppression of ROS, and (ii) faster charge recombination caused by the absence of electron-scavenging oxygen [19,72]. These effects nevertheless failed to account for the similar MB1 degradation rates in the presence (Fig. 6a) and absence (Fig. 6b) of O₂ in the first ~ 10 min of the reaction. These initially comparable rates can, however, be explained by the entrapment of electrons at Ti⁴⁺ sites (Eq. (12)), which gave rise to stable STIRAs (Fig. 6c and S9). Again, a STIRA signals the effective capture of electrons, thus enabling holes to migrate to the surface and oxidize oxalate through HT (Eqs. (15)–(18)) and/or generate *OH (Eqs. (3)–(4)). However, within this context, the rates were rapid and comparable for only the first 10 min, and followed by a subsequent deceleration. We cannot explain this slow period by electron/hole recombination, as the existence of stable STIRAs (Fig. S9) during this period indicated that electron entrapment persisted, yet the holes were unable to effectively facilitate oxalate oxidation. Consequently, these findings suggest that recombination was not the key factor to slow down the photocatalytic rates in anerobic and hydrated conditions.

Another influential factor affecting photocatalytic rates and efficiency under anerobic conditions was the spatial confinement of electrons and holes. Previous modeling studies showed that hydration in aqueous solutions trapped electrons at surface Ti⁴⁺ sites, [38], and we can substantiate this through the STIRAs (Fig. 6c and S9) which we observed in the thicker, multilayered, water films. Additionally, based

on recent work, [39] hydration could have rapidly localized holes at lattice/bulk O sites, a process that potentially decreased hole populations at the surface. Such a process could have hindered HT (Eqs. (15)–(17)) and even *OH production (Eqs. (3)–(4)). It could have consequently diminished oxalate photodecomposition rates in the thickest O₂-free water films considered in this study. We however note that previous work [75] showed that trace (sub-ML) amounts of water could, instead, enhance hole mobility, at least at the scale of minutes. This consequently implies that these effects of (i) hindered HT and *OH production in sub-ML and (ii) hydration-promoted electron trapping in larger ML films were manifested in the large range of water loadings (0.1–4 ML) considered in our study.

We consequently propose that (i) water film thickness, (ii) oxygen deficiency and (iii) accumulation of electrons in the shallow traps collectively contributed to hole localization. It also possible that photogenerated holes near sub-surfaces were more able to transfer to oxalate or water for photocatalytic reactions than photogenerated holes in the bulk. This would explain the fast regime in the first 10 min, during which time only photogenerated holes on the surfaces were consumed, while other holes in the deeper regions were less dynamics and more localized at lattice/bulk O sites.

We also note that based on previous work, [73] proton rich water films could have also stabilized shallow trapped electrons (Eq. (12)), and thereby hampered electron-hole recombination and HT to oxalate. Once the trapping sites became saturated from the high electron density, which arose from hole consumption by oxalate, bulk holes slowly began to recombine, while still decomposing oxalate. This can be seen from the ~ 1st up to the 20th h of the reaction (Figs. Inset of 5c and S9). The drop in half-lives (Fig. 6e) and increase in yields (Fig. 6g) in the thicker water films revealed, in contrast, that (i) bulk holes successfully migrated to the surface to react with oxalate, or (ii) fresh holes were generated near the surface. The latter can occur if electron-hole recombination (Eq. (14)) happened, and a new cycle of electron-hole pair generation (Eq. (2)) under photoirradiation occurred.

Finally, we must emphasize that photoreactions in water-free TiO₂ had shorter half-lives (~ 2 h) and higher yields (~ 80 %) (Fig. 6g) than in the presence of water films. From the small and unstable STIRAs

<figure><img src="image_9.png" /><figcaption>**Fig. 7.** Schematic representation of humidity- and oxygen-dependent photocatalytic decomposition of oxalate on TiO₂. (a) Nearly (0.0095 kPa H₂O) dry O₂-bearing conditions promote electron-scavenging ROS which enhances hole lifetimes. (b) Hydrated O₂-bearing conditions promote ROS and OH radical formation, as well as charge mobility. Oxalate oxidation proceeds via HT, LMCT and radical-driven reactions, which enhance reaction rates compared to the dry condition in (a). (c) Nearly (0.0095 kPa H₂O) dry O₂-free conditions promote oxalate oxidation by HT, and possibly LMCT because of the cycle of generation-recombination of electron-hole pairs. (d) Hydrated O₂-free conditions limit both charge mobility and ROS formation, thus decreasing oxalate oxidation rates. These conditions also generated STIRAs (Fig. 5 a-c).</figcaption></figure>
### 3.3.2. *O₂*-free water films

Reactions in the absence of $O_2$ had contrasting humidity-dependent trends in half-lives (Fig. 6e) and reaction yields (Fig. 6g). Firstly, yields were on par with those of $O_2$-bearing films with up to 0.5 ML (30 % RH), but thereafter dramatically dropped in thicker films. Half-lives were, in contrast, comparable on dry $TiO_2$ (~ 2 h), then rose to ~ 15 h in the 0.8 ML (50 % RH) film, and dropped to ~ 3 h in 4 ML-thick films.

From the relatively constant and comparable rates in dry $TiO_2$ in the presence and absence of $O_2$, we propose that LMCT and HT were the chief mechanisms for oxalate photodecomposition. The hike in half-lives in $O_2$-free water films (10–90 % RH) signaled, on the other hand, that films could not sustain an important production of •OH radicals compared to $O_2$-bearing films. We explain the drop in rates in $O_2$-free water films (10–90 % RH) using two primary processes: (i) the suppression of ROS, and (ii) faster charge recombination caused by the absence of electron-scavenging oxygen [19,72]. These effects nevertheless failed to account for the similar MB1 degradation rates in the presence (Fig. 6a) and absence (Fig. 6b) of $O_2$ in the first ~ 10 min of the reaction. These initially comparable rates can, however, be explained by the entrapment of electrons at $Ti^{IV}$ sites (Eq. (12)), which gave rise to stable STIRAs (Fig. 6c and S9). Again, a STIRA signals the effective capture of electrons, thus enabling holes to migrate to the surface and oxidize oxalate through HT (Eqs. (15)–(18)) and/or generate •OH (Eqs. (3)–(4)). However, within this context, the rates were rapid and comparable for only the first 10 min, and followed by a subsequent deceleration. We cannot explain this slow period by electron/hole recombination, as the existence of stable STIRAs (Fig. S9) during this period indicated that electron entrapment persisted, yet the holes were unable to effectively facilitate oxalate oxidation. Consequently, these findings suggest that recombination was not the key factor to slow down the photocatalytic rates in anerobic and hydrated conditions.

Another influential factor affecting photocatalytic rates and efficiency under anerobic conditions was the spatial confinement of electrons and holes. Previous modeling studies showed that hydration in aqueous solutions trapped electrons at surface $Ti^{IV}$ sites, [38], and we can substantiate this through the STIRAs (Fig. 6c and S9) which we observed in the thicker, multilayered, water films. Additionally, based

on recent work, [39] hydration could have rapidly localized holes at lattice/bulk O sites, a process that potentially decreased hole populations at the surface. Such a process could have hindered HT (Eqs. (15)–(17)) and even •OH production (Eqs. (3)–(4)). It could have consequently diminished oxalate photodecomposition rates in the thickest $O_2$-free water films considered in this study. We however note that previous work [75] showed that trace (sub-ML) amounts of water could, instead, enhance hole mobility, at least at the scale of minutes. This consequently implies that these effects of (i) hindered HT and •OH production in sub-ML and (ii) hydration-promoted electron trapping in larger ML films were manifested in the large range of water loadings (0.1–4 ML) considered in our study.

We consequently propose that (i) water film thickness, (ii) oxygen deficiency and (iii) accumulation of electrons in the shallow traps collectively contributed to hole localization. It also possible that photogenerated holes near sub-surfaces were more able to transfer to oxalate or water for photocatalytic reactions than photogenerated holes in the bulk. This would explain the fast regime in the first 10 min, during which time only photogenerated holes on the surfaces were consumed, while other holes in the deeper regions were less dynamics and more localized at lattice/bulk O sites.

We also note that based on previous work, [73] proton rich water films could have also stabilized shallow trapped electrons (Eq. (12)), and thereby hampered electron-hole recombination and HT to oxalate. Once the trapping sites became saturated from the high electron density, which arose from hole consumption by oxalate, bulk holes slowly began to recombine, while still decomposing oxalate. This can be seen from the ~ 1st up to the 20th h of the reaction (Figs. Inset of 5c and S9). The drop in half-lives (Fig. 6e) and increase in yields (Fig. 6g) in the thicker water films revealed, in contrast, that (i) bulk holes successfully migrated to the surface to react with oxalate, or (ii) fresh holes were generated near the surface. The latter can occur if electron-hole recombination (Eq. (14)) happened, and a new cycle of electron-hole pair generation (Eq. (2)) under photoirradiation occurred.

Finally, we must emphasize that photoreactions in water-free $TiO_2$ had shorter half-lives (~ 2 h) and higher yields (~ 80 %) (Fig. 6g) than in the presence of water films. From the small and unstable STIRAs

<figure><img src="image_10.png" /><figcaption>**Fig. 7.** Schematic representation of humidity- and oxygen-dependent photocatalytic decomposition of oxalate on $TiO_2$. (a) Nearly (0.0095 kPa $H_2O$) dry $O_2$-bearing conditions promote electron-scavenging ROS which enhances hole lifetimes. (b) Hydrated $O_2$-bearing conditions promote ROS and OH radical formation, as well as charge mobility. Oxalate oxidation proceeds via HT, LMCT and radical-driven reactions, which enhance reaction rates compared to the dry condition in (a). (c) Nearly (0.0095 kPa $H_2O$) dry $O_2$-free conditions promote oxalate oxidation by HT, and possibly LMCT because of the cycle of generation-recombination of electron-hole pairs. (d) Hydrated $O_2$-free conditions limit both charge mobility and ROS formation, thus decreasing oxalate oxidation rates. These conditions also generated STIRAs (Fig. 5 a-c).</figcaption></figure>
(Fig. 6c & S8) we suggest that charge mobility was not greatly affected. We explain this by adapting a suggestion by Litke et al. [75], in which surface OH groups [76], formed by the dissociative binding of water in sub-ML films, trapped holes at faster rates than chemisorbed (undisso- ciated) water. Because of these OH groups, holes can either recombine with electrons, producing unstable STIRAs, or participate in photore- actions with oxalate via bound *OH radicals (Eqs. (3)–(4)) and HT (Eqs. (15)–(18)). The photodecomposition rate did not slow down because fresh electron-hole pairs were continuously produced during irradiation after recombination. These holes were then able to preferentially oxidize the smaller surface loadings of oxalate (1.4 OX/nm²) over the more abundant OH sites (~ 5 OH/nm²) [64] – the sources of bound *OH radicals (Eq. (3b)). Still, we cannot discard LMCT as an additional pathway for decomposition.

### 3.3.3. Photodecomposition affected by oxalate complexation modes

Finally, we discuss how the oxalate binding modes controlled photodecomposition rate. In general, the rate loss of MB1 (1724 cm⁻¹) was more rapid than the MB2 (1690 cm⁻¹) and OS (1675 cm⁻¹) at all humidity ranges (Fig. 6f-g & S10), and this difference became important in O₂-free atmosphere. The photodecomposition trends of MB1 and MB2 were similar in both O₂-bearing and O₂-free atmospheres. OS complexes, on the other hand, shared similar photoreaction trends in the absence of O₂, yet their decomposition rates were unaffected by variations of hu- midity in the presence of O₂. While similar observations were previously made, [77] we can here explain these in terms of the different photo- chemical responses of MB and OS species.

These differences illustrate the case that oxalate decomposition did not solely proceed through a solvent-mediated *OH radical because this process would have unselectively oxidized all forms of oxalate. They show that HT and LMCT mechanisms in the absence of O₂ targeted the MB1 species, while decomposition rates of MB2 and OS were slower because *OH radical production was hampered by charge localization in thick water films. In this case, the hydration shell of the OS complex contributed to charge migration, explaining its reversal in reaction rates with humidity. Additionally, MB1 in both reaction atmospheres could undergo structural changes to MB2 before being finally decomposed, as previously suggested, [54] resulting in the slower disappearance of MB2. This mechanism could not be supported by our IR spectra because it likely happened at a faster rate than the spectral acquisition time (89 s).

As a result, we propose that MB2 and OS species were more resilient to decomposition via LMCT than MB1 species. The greater stability of MB2 and OS, again, can be justified by their first appearance upon adsorption experiment (Fig. 2c), even at low oxalate loadings. It was possible that MB2 adapted bridging (bidentate binuclear) or chelating (bidentate mononuclear) structures (Fig. 2b) for this high stability. MB1 could, on the other hand, have a monodentate structure, which is easier to decompose, and is a precursor species to MB2. [54]

## 4. Conclusions

Using vibrational spectroscopy, we resolved the water film-mediated photodecomposition of oxalate bound to TiO₂ nanoparticles. We find that thicker water films accelerated photocatalytic reaction rates for complete oxalate decomposition under atmospheric oxygen (21 kPa O₂) but decelerated the reactions in the absence of O₂. A schematic repre- sentation summarizing these effects is presented in Fig. 7. This con- trasting effect can be attributed to the interplay of five key processes: (i) direct charge transfer between oxalate and TiO₂, involving hole transfer (HT) and ligand-to-metal charge transfer (LMCT), (ii) ROS production, (iii) solvent-mediated hydroxyl (*OH) radical oxidation, (iv) charge localization, and (v) charge recombination.

In the presence of oxygen (Fig. 7 a,b), HT/LMCT and ROS played a dominant role in driving the photodecomposition of oxalate on dehy- drated TiO₂. The hydration environment of the films promoted water-

mediated *OH radical production, which enhanced the reaction rates. Conversely, in the absence of oxygen (Fig. 7 c,d), water films induced charge localization on TiO₂ that stabilized shallow trap electrons. This, in turn, restricted hole oxidation for *OH production. In addition, the absence of O₂ as an electron scavenger suppressed the formation of O₂/ HO₂ radicals. Consequently, the photodecomposition of oxalate pri- marily proceeded through HT and LMCT, competed by electron-hole recombination that ultimately decreased reaction rate.

It is thus this interplay between surface loading, coordination envi- ronment (MB1 vs. MB2 vs. OS), water film thickness and oxygen avail- ability that control the photostability of oxalate on TiO₂. By shedding new light on the water film-driven photodecomposition of oxalate by TiO₂, this work can help address the fate of organics in important natural (e.g. vadose zones of soils, atmospheric aerosol, etc.) and technological (e.g. wastewater treatment, air purification) settings.

## CRediT authorship contribution statement

**N. Tan Luong:** Writing – review & editing, Writing – original draft, Visualization, Methodology, Investigation, Formal analysis, Data cura- tion, Conceptualization. **Khalil Hanna:** Writing – review & editing. **Jean-François Boily:** Writing – review & editing, Supervision, Re- sources, Project administration, Methodology, Funding acquisition, Conceptualization.

## Declaration of competing interest

The authors declare that they have no known competing financial interests or personal relationships that could have appeared to influence the work reported in this paper.

## Data availability

Data will be made available on request.

## Acknowledgements

This work was supported by the Swedish Research Council (2016- 03808, 2020-04853) and FORMAS (2022-01246) to J.-F.B., and IRP CNRS CHEMICY (2023-2027) to J.-F.B. and K.H., and Styrelsen för Stiftelsen J C Kempes Minnes Stipendiefond to N.T.L. We thank Prof. Knut Irgum (Department of Chemistry, Umeå University), Chau Huynh (Northvolt AB), and Dr. Ngoc Phuoc Dinh (Klaria AB) for great support and assistance in the ion chromatography measurements.

## Appendix A. Supplementary data

Supplementary data to this article can be found online at https://doi. org/10.1016/j.jcat.2024.115425.

## References

[1] G.E. Ewing, Ambient thin film water on insulator surfaces, Chem. Rev. 106 (2006) 1511–1526.

[2] M. Yeşilbaş, J.-F. Boily, Particle size controls on water adsorption and condensation regimes at mineral surfaces, Sci. Rep. 6 (2016) 32136.

[3] M. Tang, D.J. Cziczo, V.H. Grassian, Interactions of water with mineral dust aerosol: water adsorption, hygroscopicity, cloud condensation, and ice nucleation, Chem. Rev. 116 (2016) 4205–4259.

[4] N.T. Luong, E.S. Ilton, A. Shchukarev, J.-F. Boily, Water film-driven mn (oxy) (hydr)oxide nanocoating growth on rhodochrosite, Geochim. Cosmochim. Acta 329 (2022) 87–105.

[5] C.E. Nanayakkara, W.A. Larish, V.H. Grassian, Titanium dioxide nanoparticle surface reactivity with atmospheric gases, CO₂, SO₂, and NO₂: roles of surface hydroxyl groups and adsorbed water in the formation and stability of adsorbed products, J. Phys. Chem. C 118 (2014) 23011–23021.

[6] G.R. Bourret, O. Diwald, Thin water films covering oxide nanomaterials: stability issues and influences on materials processing, J. Mat. Res. 34 (2019) 428–441.

[7] G. Rubasinghe, V.H. Grassian, Role(s) of adsorbed water in the surface chemistry of environmental interfaces, Chem. Commun. 49 (2013) 3071–3094.
(Fig. 6c & S8) we suggest that charge mobility was not greatly affected. We explain this by adapting a suggestion by Litke et al. [75], in which surface OH groups [76], formed by the dissociative binding of water in sub-ML films, trapped holes at faster rates than chemisorbed (undisso- ciated) water. Because of these OH groups, holes can either recombine with electrons, producing unstable STIRAs, or participate in photore- actions with oxalate via bound *OH radicals (Eqs. (3)–(4)) and HT (Eqs. (15)–(18)). The photodecomposition rate did not slow down because fresh electron-hole pairs were continuously produced during irradiation after recombination. These holes were then able to preferentially oxidize the smaller surface loadings of oxalate (1.4 OX/nm²) over the more abundant OH sites (~ 5 OH/nm²) [64] – the sources of bound *OH radicals (Eq. (3b)). Still, we cannot discard LMCT as an additional pathway for decomposition.

### 3.3.3. Photodecomposition affected by oxalate complexation modes

Finally, we discuss how the oxalate binding modes controlled photodecomposition rate. In general, the rate loss of MB1 (1724 cm⁻¹) was more rapid than the MB2 (1690 cm⁻¹) and OS (1675 cm⁻¹) at all humidity ranges (Fig. 6f-g & S10), and this difference became important in O₂-free atmosphere. The photodecomposition trends of MB1 and MB2 were similar in both O₂-bearing and O₂-free atmospheres. OS complexes, on the other hand, shared similar photoreaction trends in the absence of O₂, yet their decomposition rates were unaffected by variations of hu- midity in the presence of O₂. While similar observations were previously made, [77] we can here explain these in terms of the different photo- chemical responses of MB and OS species.

These differences illustrate the case that oxalate decomposition did not solely proceed through a solvent-mediated *OH radical because this process would have unselectively oxidized all forms of oxalate. They show that HT and LMCT mechanisms in the absence of O₂ targeted the MB1 species, while decomposition rates of MB2 and OS were slower because *OH radical production was hampered by charge localization in thick water films. In this case, the hydration shell of the OS complex contributed to charge migration, explaining its reversal in reaction rates with humidity. Additionally, MB1 in both reaction atmospheres could undergo structural changes to MB2 before being finally decomposed, as previously suggested, [54] resulting in the slower disappearance of MB2. This mechanism could not be supported by our IR spectra because it likely happened at a faster rate than the spectral acquisition time (89 s).

As a result, we propose that MB2 and OS species were more resilient to decomposition via LMCT than MB1 species. The greater stability of MB2 and OS, again, can be justified by their first appearance upon adsorption experiment (Fig. 2c), even at low oxalate loadings. It was possible that MB2 adapted bridging (bidentate binuclear) or chelating (bidentate mononuclear) structures (Fig. 2b) for this high stability. MB1 could, on the other hand, have a monodentate structure, which is easier to decompose, and is a precursor species to MB2. [54]

## 4. Conclusions

Using vibrational spectroscopy, we resolved the water film-mediated photodecomposition of oxalate bound to TiO₂ nanoparticles. We find that thicker water films accelerated photocatalytic reaction rates for complete oxalate decomposition under atmospheric oxygen (21 kPa O₂) but decelerated the reactions in the absence of O₂. A schematic repre- sentation summarizing these effects is presented in Fig. 7. This con- trasting effect can be attributed to the interplay of five key processes: (i) direct charge transfer between oxalate and TiO₂, involving hole transfer (HT) and ligand-to-metal charge transfer (LMCT), (ii) ROS production, (iii) solvent-mediated hydroxyl (*OH) radical oxidation, (iv) charge localization, and (v) charge recombination.

In the presence of oxygen (Fig. 7 a,b), HT/LMCT and ROS played a dominant role in driving the photodecomposition of oxalate on dehy- drated TiO₂. The hydration environment of the films promoted water-

mediated *OH radical production, which enhanced the reaction rates. Conversely, in the absence of oxygen (Fig. 7 c,d), water films induced charge localization on TiO₂ that stabilized shallow trap electrons. This, in turn, restricted hole oxidation for *OH production. In addition, the absence of O₂ as an electron scavenger suppressed the formation of O₂/ HO₂ radicals. Consequently, the photodecomposition of oxalate pri- marily proceeded through HT and LMCT, competed by electron-hole recombination that ultimately decreased reaction rate.

It is thus this interplay between surface loading, coordination envi- ronment (MB1 vs. MB2 vs. OS), water film thickness and oxygen avail- ability that control the photostability of oxalate on TiO₂. By shedding new light on the water film-driven photodecomposition of oxalate by TiO₂, this work can help address the fate of organics in important natural (e.g. vadose zones of soils, atmospheric aerosol, etc.) and technological (e.g. wastewater treatment, air purification) settings.

## CRediT authorship contribution statement

**N. Tan Luong:** Writing – review & editing, Writing – original draft, Visualization, Methodology, Investigation, Formal analysis, Data cura- tion, Conceptualization. **Khalil Hanna:** Writing – review & editing. **Jean-François Boily:** Writing – review & editing, Supervision, Re- sources, Project administration, Methodology, Funding acquisition, Conceptualization.

## Declaration of competing interest

The authors declare that they have no known competing financial interests or personal relationships that could have appeared to influence the work reported in this paper.

## Data availability

Data will be made available on request.

## Acknowledgements

This work was supported by the Swedish Research Council (2016- 03808, 2020-04853) and FORMAS (2022-01246) to J.-F.B., and IRP CNRS CHEMICY (2023-2027) to J.-F.B. and K.H., and Styrelsen för Stiftelsen J C Kempes Minnes Stipendiefond to N.T.L. We thank Prof. Knut Irgum (Department of Chemistry, Umeå University), Chau Huynh (Northvolt AB), and Dr. Ngoc Phuoc Dinh (Klaria AB) for great support and assistance in the ion chromatography measurements.

## Appendix A. Supplementary data

Supplementary data to this article can be found online at https://doi. org/10.1016/j.jcat.2024.115425.

## References

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