H₂O₂ evolution during the photocatalytic degradation of organic molecules on fluorinated TiO₂

Abstract

The effect of TiO₂ surface fluorination on the hydrogen peroxide evolution occurring in photocatalytic runs was investigated employing the azo dye Acid Red 1 (AR1) and two model organic molecules with acidic properties, i.e. formic acid (FA) and benzoic acid (BA), as substrates of oxidative degradation. While AR1 and BA photocatalytic degradation on fluorinated titanium dioxide (F–TiO₂) was markedly faster than on unmodified TiO₂, because of enhanced hydroxyl radical formation, H₂O₂ concentration during the photodegradation of both substrates on F–TiO₂ was lower, possibly because of the reduced rate of interfacial electron transfer. By contrast, FA underwent slower photocatalytic degradation on F–TiO₂, but, at the same time, hydrogen peroxide concentration was relatively high, while no H₂O₂ could be detected during FA photodegradation on unmodified TiO₂. Photocatalytic runs in the presence of the nitrate anion, able to react with the CO₂˙⁻ species produced from FA oxidation, but not with conduction band electrons, demonstrated that CO₂˙⁻ plays a relevant role in H₂O₂ formation during FA degradation on F–TiO₂. In fact, surface fluoride, having a shielding effect at the semiconductor–water interface, not only inhibits the photocatalytic decomposition of H₂O₂, but also favours CO₂˙⁻ desorption and reaction with dissolved O₂, generating H₂O₂. By contrast, CO₂˙⁻ mainly gives electron transfer to the conduction band of naked TiO₂ and surface reduction of the photocatalytically produced H₂O₂.

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Introduction

The electron transfer paths occurring at a water–semiconductor interface during photocatalytic processes^1,2 undergo strong modification upon titanium dioxide surface fluorination.^3–5 Indeed, the formation of ≡Ti–F species,⁶ which dominate at acidic pH, decreases the amount of surface hydroxyl groups, up to their almost complete displacement at pH 3.5–4.0.⁷ Significant variations of the semiconductor oxide photocatalytic behaviour have consequently been observed.^8–15

An increase in the photocatalytic degradation of phenol on fluorinated titanium dioxide (F–TiO₂) led to the hypothesis that an ˙OH radical mediated mechanism involving the homogeneous aqueous phase is almost exclusively responsible for the oxidation of this substrate in F–TiO₂ systems.^8,9 By contrast, hole transfer mediated oxidations are expected to be largely inhibited because of the hindered adsorption of substrates on F–TiO₂. At the same time, surface –F groups, because of the strong electronegativity of the fluorine atom, seem to act as electron-trapping sites, reducing the interfacial electron transfer rates of conduction band electrons.¹³ Spin-trapping EPR measurements provided direct experimental evidence that a ca. ten-fold higher concentration of hydroxyl radicals is generated on fluorinated titanium dioxide under irradiation, demonstrating that on F–TiO₂ organic molecules mainly undergo photodegradation via ˙OH radical attack.¹⁵

To have a deeper insight into the effects of surface fluorination on the electron transfer processes at the semiconductor–water interface under irradiation, we focused our attention on the main reductive counter reaction of most photocatalytic reactions proceeding oxidatively, i.e. hydrogen peroxide evolution. A preliminary note reporting a sustained production of H₂O₂ on irradiated TiO₂–fluoride systems appeared very recently.¹⁶ In the present study H₂O₂ evolution was monitored during the photocatalytic degradation of the azo dye Acid Red 1 (AR1, see Scheme 1), in the presence of naked or fluorinated titanium dioxide. This dye bears two sulfonic groups, capable of strong electrostatic interactions with the unmodified oxide surface.¹⁷ One of the first steps in its photocatalytic degradation is the cleavage of the azo bond, responsible for its bleaching in the visible region.¹⁸ The photocatalytic behaviour of this rather complex aromatic molecule on F–TiO₂ and the simultaneous H₂O₂ evolution were compared to those of benzoic acid (BA), chosen as a moderately acidic model aromatic molecule, and of formic acid (FA), chosen as a model aliphatic acidic compound undergoing direct mineralisation to CO₂ and H₂O.^19,20

Scheme 1 Molecular structure of the azo dye Acid Red 1 (AR1).

The origin of the peculiarly outstanding amount of H₂O₂ evolved during FA photocatalytic degradation on F–TiO₂ was clarified by investigating the effects of the addition of the nitrate anion, which evidenced the central role of the CO₂˙⁻ species, formed by photoinduced FA oxidation on the photocatalyst surface.

Experimental

Materials

Acid Red 1 (AR1), purchased from Aldrich, was purified by repeated crystallisation from methanol. Its purity from organic contaminants was verified by NMR analysis. Degussa P25 titanium dioxide (mainly anatase) was employed as photocatalyst. Formic acid (FA, purity 95–97%), benzoic acid (BA, purity >99.5%), 2-propanol (purity 99.5%), KNO₃ (purity 99.5%) and NaF (purity 99.99%) were purchased from Aldrich and employed as received. Water purified by a Milli-Q water system (Millipore) was used throughout.

Apparatus

All degradation runs were carried out at (35 ± 1) °C under atmospheric conditions in a magnetically stirred 400 mL cylindrical Pyrex reactor, employing an experimental set up similar to that already described.²¹ Illumination was performed through the reactor Pyrex walls by means of a 250 W iron alogenide lamp (Jelosil, model HG 200), emitting in the 315–400 nm wavelength range, with a mean emission intensity on the reactor of 3.5 × 10⁻⁴ einstein L⁻¹ s⁻¹, as periodically checked by ferrioxalate actinometry.²²

Procedures

The irradiated aqueous suspensions contained 0.1 g L⁻¹ of TiO₂; the initial concentration (C₀) of the substrates was 2.5 × 10⁻⁵ M for AR1, 1.0 × 10⁻⁴ M for BA and 5.0 × 10⁻⁴ M for FA. Titanium dioxide fluorination was achieved by adding 0.01 M NaF, corresponding to 0.1 mol of F⁻ per gram of TiO₂. Fluoride ions, able to very quickly displace the –OH groups on the surface of titanium dioxide, were added to the suspensions immediately before starting irradiation, to minimise the coagulation of the photocatalyst.^8,10 When investigating the effect of nitrate addition, the KNO₃ concentration was 0.05 M, high enough to guarantee a quantitative scavenging of the carbon dioxide radical anion.²³

The pH was monitored during the runs by means of an Amel Instruments 334-B pH-meter. A decrease in pH was observed during AR1 photocatalytic degradation under natural pH conditions, from an initial value of 5.8 to a final value of ca. 4.4, as a consequence of the production of stable acids, due to the fast removal of sulfonic groups and the oxidation of the azo double bond.²⁴ Thus, to guarantee an efficient adsorption of fluoride anions on the surface of the oxide, during AR1 photodegradation the pH was lowered to 3.7 by adding small amounts of HClO₄, which is known to have negligible influence on the photocatalytic activity, because of the low affinity of ClO₄⁻ anions for TiO₂ and their low reactivity towards hydroxyl radicals.²⁵ No pH change was noticed during the runs under such conditions. By contrast, the pH value of suspensions containing formic acid and benzoic acid increased during the runs, from 3.5 to 5.8 and from 4.2 to ca. 6.0, respectively, as a direct consequence of the mineralisation of the acids to CO₂ and H₂O.²⁶ With these acidic substrates, the runs were performed under so-called natural pH conditions, corresponding to the range of maximum fluoride adsorption on TiO₂, i.e. no buffer was added to the suspensions, to avoid possible interference of other species (mainly anions) on the photoredox processes at the TiO₂–water interface.

2-mL samples were periodically withdrawn from the reactor and analysed, after removal of TiO₂ particles by centrifugation at 3000 rpm for 30 min, employing an ALC 4225 centrifuge.²¹ The cleavage of the azo bond of AR1, leading to its bleaching (also mentioned as AR1 degradation), was monitored by spectrophotometric analysis at 531 nm (maximum AR1 absorption, ε = [3.13 ± 0.02] × 10⁴ M⁻¹ cm⁻¹) by means of a Perkin Elmer Lambda 16 spectrophotometer.²¹ BA concentration during the runs was detected by HPLC analysis, employing an Agilent 1100 Series apparatus, equipped with a μBondapack-C18 column and a UV-Vis detector. An acetonitrile∶water 60∶40 mobile phase was used for BA analysis,²⁷ flowing at 1.0 mL min⁻¹. FA concentration changes were detected using a total organic carbon (TOC) analyser in the not purgeable organic carbon (NPOC) mode (Shimadzu Instruments, TOC-5000A). All runs were repeated at least twice to check their reproducibility.

Hydrogen peroxide concentration was monitored during the photodegradation runs by fluorimetric analysis (λ_ex = 316.5 nm, λ_em = 408.5 nm) of the fluorescent dimer formed in the horseradish peroxidase-catalysed reaction of hydrogen peroxide with p-hydroxyphenylacetic acid,^28,29 using a 605-10S Perkin Elmer fluorescence spectrophotometer. H₂O₂ standard solutions employed in calibration were analysed iodometrically. The H₂O₂ concentration profiles obtained during BA photocatalytic degradation were corrected for the relatively small fluorescence signal originating from salicylic acid, one of the first degradation intermediates of BA. The signal was measured in blank photocatalytic runs under conditions identical to those employed in H₂O₂ analysis, but with no p-hydroxyphenylacetic acid and enzyme addition.

Adsorption studies were performed both in the presence and in the absence of 0.01 M NaF, on suspensions containing 1.0 g L⁻¹ of TiO₂ and 2.5 × 10⁻⁵ M AR1, 1.0 × 10⁻⁴ M BA or 5.0 × 10⁻⁴ M FA. After continuous stirring for 24 h in the dark at 35 °C, the photocatalyst was removed and the liquid phase was analysed for AR1, BA or FA residual content. The F⁻/TiO₂ ratio employed in the adsorption experiments, lower with respect to that of the photocatalytic degradation runs, guaranteed a similar displacement of surface –OH groups, the maximum value of adsorbed fluoride being 2.5 × 10⁻⁴ mol g⁻¹ of TiO₂.³⁰

The reduction potentials of AR1, BA and FA were determined by cyclic voltammetry measurement using a glassy carbon working electrode vs. Ag/AgCl in 0.1 M HClO₄ for AR1³¹ or 0.5 M NaClO₄ for BA and FA solutions, with a scan speed of 50 mV s⁻¹.

Results and discussion

Adsorption and photocatalytic degradation of the organic substrates

Preliminary adsorption studies evidenced that, while AR1 and BA are detectably adsorbed on the unmodified TiO₂ surface, their adsorption is almost completely inhibited on F-modified TiO₂. In fact, the adsorbed fraction on naked titanium dioxide at natural pH was 0.27 for AR1 and 0.30 for BA, but below our detection limit (adsorbed fraction ca. 0.001) on the fluorinated oxide. Similar results were obtained for the adsorption of the azo dye Acid Orange 7 on TiO₂ and on F–TiO₂.¹³ Thus, the displacement of –OH groups with the formation of stable ≡Ti–F species and the consequent negative surface charge¹³ induce significant alterations of the adsorption equilibria at the water–TiO₂ interface. On the other hand, under the adopted conditions FA adsorption could not be detected on both unmodified and F-modified TiO₂. FA adsorption on naked titanium dioxide has been recently measured in the presence of higher oxide amounts.³² Also the adsorption of molecular oxygen is expected to undergo deep modification upon TiO₂ fluorination, with great consequences on the rates of conduction band electron scavenging by O₂.^8,33

Upon TiO₂ fluorination, AR1 photocatalytic degradation at pH 3.7, i.e. under conditions of maximum F⁻ adsorption on the oxide surface, was found to proceed at a more than double rate,¹⁵ as shown by the first-order rate constants reported in Table 1.

Table 1 Rate constants for the photocatalytic degradation of the azo dye Acid Red 1 (C₀ = 2.5 × 10⁻⁵ M), of benzoic acid (C₀ = 1.0 × 10⁻⁴ M) and of formic acid (C₀ = 5.0 × 10⁻⁴ M) on TiO₂ or F–TiO₂

Substrate	TiO2	F–TiO2
AR1	(3.2 ± 0.4) × 10^−4 s^−1	(7.34 ± 0.15) × 10^−4 s^−1
BA	(4.65 ± 0.11) × 10^−4 s^−1	(7.84 ± 0.16) × 10^−4 s^−1
FA	(1.83 ± 0.05) × 10^−7 M s^−1	(1.58 ± 0.04) × 10^−7 M s^−1

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BA exhibits a behaviour similar to that of AR1, i.e. it underwent faster reaction on fluorinated TiO₂. The pseudo first-order rate constants for BA degradation on TiO₂ and F–TiO₂, reported in Table 1, also for this substrate demonstrate an almost double photodegradation rate upon fluoride addition under the conditions of the present study. Higher increases in rate have been very recently reported for higher initial BA concentrations.³⁴

FA was found to undergo photodegradation according to a zero order rate law on both unmodified and F-modified TiO₂. However, as evidenced by the rate constants reported in Table 1, formic acid exhibits an opposite behaviour compared to AR1 and BA, its photocatalytic degradation on F-modified TiO₂ being slower than on naked TiO₂.

At neutral pH, F⁻ ions displace basic –OH groups (estimated around 0.14 mmol g⁻¹), with an equilibrium constant³⁵ equal to 8 × 10⁻⁷; at lower pH even acidic –OH groups are substituted by –F atoms.⁶ This induces a relevant change in the adsorption properties on the photocatalyst surface, especially for acidic molecules strongly interacting with basic sites,⁹ as is the case for our substrates. Moreover, the displacement of hydroxyl groups implies the absence of an effective trap for photogenerated valence band holes as ≡TiO˙ species, formed through reaction (1). In fact, –F species are stable and cannot be oxidised by valence band holes h_vb⁺ even in acidic media. Indeed, the potential of the F˙/F⁻ couple in homogeneous aqueous phase is 3.6 V,³⁶ while the valence band energy is E⁰_vb(V) = 3.00 − 0.059 pH vs. NHE.²³ Thus, valence band holes photoproduced in F–TiO₂ directly react with water molecules at the interface, producing ˙OH radicals according to reaction (2).

≡Ti–OH + h_vb⁺ → ≡Ti–O˙ + H⁺ (1)

≡Ti–F + h_vb⁺ + H₂O → ≡Ti–F + ˙OH + H⁺ (2)

The presence of fluoride anions on the photocatalyst surface, having a shielding effect on surface reactions leading to ˙OH radical decomposition, favours their desorption and their higher accumulation in the aqueous phase.¹⁵ By contrast, ˙OH radicals formed by water oxidation on unmodified TiO₂ are unable to leave the surface, because they rapidly react with surface –OH groups.³⁷

Thus, all ˙OH radical mediated paths are favoured in F–TiO₂ suspensions and substrates that easily undergo ˙OH radical attack are more rapidly degraded. This is the case for AR1 and BA, the second-order rate constants for hydroxyl radical attack having been reported to be k = 4.3 × 10⁹ M⁻¹ s⁻¹ for BA and k = 1.7 × 10¹⁰ M⁻¹ s⁻¹ for an azo dye similar to AR1.³⁸ On the other hand, substrates, such as FA, having a relatively lower second-order rate constant for hydroxyl attack (k = 1.3 × 10⁸ M⁻¹ s⁻¹),³⁸ whose degradation is preferably initiated by direct hole transfer, upon TiO₂ fluorination undergo slower photocatalytic degradation because of their hindered adsorption (or complexation) on the F–TiO₂ surface.¹³

Moreover, the presence of aromatic rings in AR1 and BA can also play a role in the observed increase of photocatalytic activity of F–TiO₂. In fact, the radical or cationic species formed upon reaction of aromatic rings with ˙OH radicals or with valence band holes are relatively long-lived, being highly stabilised by resonance. Thus, they have a high probability to back react with conduction band electrons, giving no net change and enhanced recombination.³⁹ The lower adsorption of AR1 and BA on F–TiO₂ can inhibit the detrimental back reaction with conduction band electrons, with a consequent increase of the overall degradation rate.

Hydrogen peroxide evolution during photocatalytic degradation

Hydrogen peroxide, a key intermediate in photocatalytic processes, was demonstrated to form on TiO₂ exclusively via reduction of molecular oxygen by conduction band electrons.⁴⁰Fig. 1 shows the concentration profile of H₂O₂ evolved during AR1 photodegradation on naked TiO₂ and on F–TiO₂ at pH 3.7. A lower H₂O₂ amount accumulated on fluorinated TiO₂ and its formation was retarded, with respect to naked TiO₂, so that it could hardly be detected on F–TiO₂ during the first 30 min of irradiation. Moreover, H₂O₂ concentration always remained below 4 × 10⁻⁵ M during AR1 photocatalytic degradation on F–TiO₂, i.e. lower than one third of the maximum concentration attained in the absence of fluoride.

Fig. 1 Hydrogen peroxide evolution during AR1 photodegradation (C₀ = 2.5 × 10⁻⁵ M) on naked (○) and on fluorinated TiO₂ at pH 3.7 (●).

Based on these results, AR1 photocatalytic degradation on fluorinated TiO₂ may appear in competition with molecular oxygen reduction by conduction band electrons. This would be compatible with our cyclic voltammetry measurements on AR1, yielding an AR1 reduction potential value of −0.33 V (NHE), and with a reported⁴¹ shift of the band edge potentials of TiO₂ towards more negative values upon addition of a fluoride salt in acetonitrile. Thus, electron transfer from the conduction band to AR1 is in principle allowed on unmodified TiO₂ [E⁰_cb(V) = −0.13 − 0.059 pH vs. NHE]⁴² and even thermodynamically more favourable on the F-modified oxide, under the hypothesis of the above mentioned negative band shift.⁴¹ Therefore, besides through hydroxyl radical attack or reaction with photogenerated holes, the azo group of AR1, which ensures extensive conjugation between its two aromatic moieties and hence its capability of absorbing light in the visible, might also undergo degradation through a reductive pathway,⁴³ directly or indirectly involving conduction band electrons.

Moreover, the higher H₂O₂ amount observed during AR1 photocatalytic degradation on naked TiO₂ (Fig. 1) might simply be a consequence of a relatively higher H₂O₂ production occurring through a dye sensitised path.⁴⁴ Indeed, AR1 is able to absorb a fraction of the incident radiation and electron transfer from the electronically excited state of AR1 to the semiconductor conduction band is expected to be favoured by the higher adsorption of AR1 on TiO₂ with respect to F–TiO₂.

However, ≡Ti–F groups on the oxide surface have been recently shown to act as electron-trapping sites, decreasing interfacial electron transfer rates by tightly holding trapped electrons, due to the strong electronegativity of fluorine.¹³ Thus, the slower formation of hydrogen peroxide in fluorinated systems might well be just the direct consequence of this effect.

Hydrogen peroxide evolution during BA photodegradation clarifies this point; in fact, for BA a reductive path is thermodynamically impossible, because its reduction potential, measured by cyclic voltammetry, is −1.22 V (NHE). The H₂O₂ concentration profiles obtained during BA photodegradation in the presence and in the absence of fluoride (Fig. 2) appear very similar to those obtained during the degradation of the azo dye (Fig. 1). This points to a decreased rate of interfacial electron transfer in the presence of surface fluoride as the main reason for the lower H₂O₂ evolution and also excludes any relevant role of AR1 photocatalytic reduction and of a dye sensitised path for H₂O₂ production, which are not possible for BA.

Fig. 2 Hydrogen peroxide evolution during BA photodegradation (C₀ = 1.0 × 10⁻⁴ M) on naked (○) and on fluorinated TiO₂ (●).

The reactivity of the superoxide radical anion towards aromatic radical species may be invoked to account for the shape of the H₂O₂ concentration profiles obtained during AR1 and BA photocatalytic degradation on F–TiO₂ (Fig. 1 and 2). In fact, this species and its protonated form, the perhydroxyl radical HO₂˙ , are precursors of H₂O₂ and may easily attack aromatic radical species.⁴⁵ On fluorinated titania the active species, i.e. the radicals generated by both reduction and oxidation surface reactions, do not reside on the surface. This makes their reactions in the homogeneous phase more competitive, compared to further direct electron transfer at the water–semiconductor interface. As a result, the reactions between aromatic molecules, such as AR1 and BA, and the superoxide radical anion are expected to be enhanced on F–TiO₂, with the consequent faster degradation of both substrates and the lower production of hydrogen peroxide, as long as the aromatic moieties are present.

The evolution profiles of H₂O₂ during FA photodegradation, both in the absence and in the presence of fluoride, are shown in Fig. 3. No hydrogen peroxide could be detected on naked TiO₂; in fact, H₂O₂ undergoes extremely fast decomposition on the “clean” surface of titania during formic acid photodegradation, having a high affinity for the oxide surface and being able to complex Ti(IV) ions at the interface.^16,46 By contrast, H₂O₂ could be easily detected on F–TiO₂. Its concentration increased from the beginning of the run, up to values even higher than those attained during AR1 and BA photodegradation, and started to decline after almost complete FA degradation, i.e. when all species able to scavenge the photogenerated valence band holes had been consumed. After this time, the electron–hole fast recombination would make less conduction band electrons available for O₂ reduction.

Fig. 3 Hydrogen peroxide evolution during FA photodegradation (C₀ = 5.0 × 10⁻⁴ M) on naked (○) and on fluorinated TiO₂ (●).

Fluoride anions adsorbed on the TiO₂ surface clearly have a shielding effect on the quite fast photodegradation of hydrogen peroxide occurring at the water–semiconductor interface. In fact, adsorbed F⁻ ions inhibit the formation of surface peroxides (≡Ti^IV–OOH),⁸ a preliminary key step for their subsequent degradation by photogenerated species according to eqn (3) and (4):⁴⁰

H₂O₂ + 2e_cb⁻ + 2H⁺ → 2H₂O (3)

H₂O₂ + 2h_vb⁺ → O₂ + 2H⁺ (4)

The shielding effect of adsorbed fluoride anions was confirmed by monitoring the photocatalytic degradation of H₂O₂ on naked and F-modified TiO₂, starting from an initial concentration around 2 × 10⁻⁴ M. Both experiments were performed at pH 3.7, i.e. under conditions of F⁻ maximum adsorption on TiO₂. The H₂O₂ concentration profiles, shown in Fig. 4, clearly demonstrate the slower photodegradation of hydrogen peroxide on the fluorinated surface.

Fig. 4 Photocatalytic degradation of hydrogen peroxide (C₀ ∼ 2 × 10⁻⁴ M) at 35 °C at pH 3.7 on naked (○) and on fluorinated TiO₂ (●).

The addition of a rather high concentration of FA (2 × 10⁻³ M) during AR1 photodegradation on both naked and fluorinated titanium dioxide also confirmed this effect. As shown by the H₂O₂ concentration profiles reported in Fig. 5, FA addition first of all caused an increase of H₂O₂ concentration: molecular oxygen reduction, yielding hydrogen peroxide, is promoted by a higher availability of conduction band electrons, as a consequence of a reduced recombination rate with photogenerated holes. This occurs when valence band holes can directly or indirectly react with compounds, such as FA, acting as scavengers of photoproduced oxidant species. Therefore, the results in Fig. 5 indicate that H₂O₂ also forms on F–TiO₂ (as it forms on TiO₂) via O₂ reduction by conduction band electrons. Moreover, in the presence of FA, H₂O₂ concentration continuously increased on F–TiO₂, mainly thanks to the shielding effect of surface fluorides. By contrast, on naked TiO₂ rather high amounts of H₂O₂ were present only as long as AR1 acted as a shield on the TiO₂ surface, hindering the formation of photoactive peroxo complexes.

Fig. 5 Hydrogen peroxide evolution during AR1 photodegradation at pH 3.7 on TiO₂ (open symbols) and on F–TiO₂ (full symbols) in the absence of scavengers (circles) and in the presence of 2 × 10⁻³ M formic acid (triangles) or 2-propanol (squares).

Similar results were obtained when 2-propanol (2 × 10⁻³ M), a well known oxidisable scavenger with relatively low affinity for the TiO₂ surface, was added during AR1 photodegradation (Fig. 5). The behaviour of the two scavenger systems is qualitatively similar, though a lower concentration of H₂O₂ was observed when 2-propanol, instead of FA, was added to the suspensions containing either naked or fluorinated TiO₂.

The main peculiarity of FA consists in the fact that a strongly reductant species, i.e. CO₂˙⁻ or HCO₂˙ depending on pH, with a pK_a of 1.4, forms in its photocatalytic oxidation on TiO₂ particles,⁴⁷ originating the so-called current doubling effect.⁴⁸ In fact, its redox potential, i.e. E⁰_{(CO2/˙CO2⁻)} = −1.8 V,²³ makes the carbon dioxide radical anion able to inject an electron into the conduction band of titanium dioxide, see reaction (5), and also to mediate the reduction of a wide variety of molecules, in particular dissolved O₂, according to reaction (6):

CO₂˙⁻ → CO₂ + e_cb⁻ (5)

CO₂˙⁻ + O₂ → CO₂ + O₂˙⁻ (6)

The role of these reactions was discerned by verifying the effects of nitrate ion addition on hydrogen peroxide formation. In fact, nitrate anions are able to react with the carbon dioxide radical anion according to reaction (7), but not with conduction band electrons,²³ their standard redox potential being E⁰_{(NO3⁻/NO3²⁻)} = −1.0 V vs. NHE, i.e. more negative than the TiO₂ conduction band potential (vide supra).

NO₃⁻ + CO₂˙⁻ → NO₃²⁻ + CO₂ (7)

The effects of the addition of 0.05 M nitrate anions on both the FA photocatalytic degradation rate and hydrogen peroxide evolution on F–TiO₂ are shown in Fig. 6. Of course, this type of experiment could only be performed on F–TiO₂, H₂O₂ evolution remaining always below the detection limit on naked TiO₂. The rate of FA photocatalytic degradation was not modified in the presence of nitrate anions, the measured rate constants being k = (1.58 ± 0.04) × 10⁻⁷ M s⁻¹ in the absence of nitrate and k = (1.57 ± 0.12) × 10⁻⁷ M s⁻¹ in its presence. By contrast, hydrogen peroxide formation significantly decreased during FA photodegradation on F–TiO₂, indicating that the carbon dioxide radical anion, which is scavenged by nitrate, plays a major role in H₂O₂ production under these conditions, yielding e_cb⁻ and O₂˙⁻ through reactions (5) and (6).

Fig. 6 Photocatalytic degradation of formic acid (open symbols, left axis) and hydrogen peroxide evolution (full symbols, right axis) on F–TiO₂ in the absence of nitrate (circles) and in the presence of 0.05 M nitrate (triangles).

This conclusion was checked using BA as a photodegradation substrate on TiO₂. The H₂O₂ concentration profiles in the presence and in the absence of nitrate were very similar, attaining a 7 × 10⁻⁵ M maximum concentration in both cases. Thus, when the radical species produced by substrate oxidation does not have reductive properties, the nitrate anion has no influence on hydrogen peroxide production, which proceeds in this case only through direct dioxygen reduction by conduction band electrons.

The crucial role of the CO₂˙⁻ radical in hydrogen peroxide formation during FA photodegradation can be appreciated only on F–TiO₂, because on naked TiO₂ this highly reactive species is adsorbed at the semiconductor–water interface and rapidly further oxidised to CO₂ according to reaction (5). Moreover, this radical may also react with the photocatalytically formed hydrogen peroxide, also adsorbed on the TiO₂ surface:

H₂O₂ + CO₂˙⁻ → ˙OH + OH⁻ + CO₂ (8)

Reaction (8) provides another decomposition path for hydrogen peroxide, which can concur in lowering its concentration during FA photodegradation on naked TiO₂.

Finally, the difference in H₂O₂ concentration observed on both TiO₂ and F–TiO₂ in the presence of FA compared to 2-propanol (Fig. 5) indicates that a radical with reductive properties is generated also by the one electron oxidation of 2-propanol. Such a species is able to produce H₂O₂ through reactions analogous to (5) and (6), though its lower redox potential²³ makes it less powerful in reducing molecular oxygen.

Conclusions

Surface TiO₂ fluorination inhibits H₂O₂ evolution during the photocatalytic degradation of the aromatic substrates AR1 and BA, mainly because of the reduced rate of interfacial electron transfer through the fluorinated surface. By contrast, H₂O₂ production is sustained during FA photocatalytic degradation on F–TiO₂: surface fluoride, having a shielding effect at the interface, inhibits the photocatalytic decomposition of H₂O₂ and favours the desorption of the CO₂˙⁻ radical anion produced from FA, thus favouring its reaction with dissolved O₂, to yield H₂O₂. No H₂O₂ evolution can be observed during FA photocatalytic degradation on naked TiO₂, because CO₂˙⁻ preferentially gives electron transfer to the conduction band, originating the current doubling effect, and also surface reduction of the photocatalytically produced H₂O₂.

References

1. D. W. Bahnemann, J. Cunningham, M. A. Fox, E. Pelizzetti, P. Pichat and N. Serpone, in Aquatic and Surface Photochemistry, ed. G. R. Helz, R. G. Zepp and D. G. Crosby, Lewis Publisher, Boca Raton, FL, 1994, p. 261.
2. M. R. Hoffmann, S. T. Martin, W. Choi and D. W. Bahnemann, Chem. Rev., 1995, 95, 69.
3. C. Chen, X. Li, W. Ma, J. Zhao, H. Hidaka and N. Serpone, J. Phys. Chem. B, 2002, 106, 318.
4. W. Choi, A. Termin and M. R. Hoffmann, J. Phys. Chem., 1994, 98, 13669.
5. S. Ikeda, N. Sugiyama, B. Pal, G. Marcí, L. Palmisano, H. Noguchi, K. Uosaki and B. Ohtani, Phys. Chem. Chem. Phys., 2001, 3, 267.
6. M. Herrmann, U. Kaluza and H. P. Boehm, Z. Anorg. Allg. Chem., 1970, 372, 308.
7. J. Sanchez and J. Augustynski, J. Electroanal. Chem., 1979, 103, 423.
8. C. Minero, G. Mariella, V. Maurino and E. Pelizzetti, Langmuir, 2000, 16, 2632.
9. C. Minero, G. Mariella, V. Maurino, D. Vione and E. Pelizzetti, Langmuir, 2000, 16, 8964.
10. M. S. Vohra, S. Kim and W. Choi, J. Photochem. Photobiol., A, 2003, 160, 55.
11. M. Lewandowski and D. F. Ollis, J. Catal., 2003, 217, 38.
12. K. Chiang, R. Amal and T. Tran, J. Mol. Catal. A: Chem., 2003, 193, 285.
13. H. Park and W. Choi, J. Phys. Chem. B, 2004, 108, 4086.
14. Y. C. Oh and W. S. Jenks, J. Photochem. Photobiol., A, 2004, 162, 323.
15. M. Mrowetz and E. Selli, Phys. Chem. Chem. Phys., 2005, 7, 1100.
16. V. Maurino, C. Minero, G. Mariella and E. Pelizzetti, Chem. Commun., 2005, 2627.
17. J. Bandara, J. A. Mielczarski and J. Kiwi, Langmuir, 1999, 15, 7670.
18. K. Vinodgopal, D. E. Wynkoop and P. V. Kamat, Environ. Sci. Technol., 1996, 30, 1660.
19. L. Davydov and P. G. Smirniotis, J. Catal., 2000, 191, 105.
20. C. J. G. Cornu, A. J. Colussi and M. R. Hoffmann, J. Phys. Chem. B, 2001, 105, 1351.
21. M. Mrowetz and E. Selli, J. Photochem. Photobiol., A, 2004, 162, 89.
22. C. G. Hatchard and C. A. Parker, Proc. R. Soc. London, Ser. A, 1956, 235, 518.
23. L. L. Perissinotti, M. A. Brusa and M. A. Grela, Langmuir, 2001, 17, 8422.
24. J. M. Joseph, H. Destaillats, H. M. Hung and M. R. Hoffmann, J. Phys. Chem. A, 2000, 104, 301.
25. M. Abdullah, J. K. C. Low and R. W. Matthews, J. Phys. Chem., 1990, 94, 6820.
26. R. J. Candal, W. A. Zeltner and M. A. Anderson, Environ. Sci. Technol., 2000, 34, 3443.
27. A. A. Ajmera, S. B. Sawant, V. G. Pangarkar and A. A. C. M. Beenackers, Chem. Eng. Technol., 2002, 25, 173.
28. G. G. Guilbault, P. J. Brignac Jr. and M. Juneau, Anal. Chem., 1968, 40, 1256.
29. C. Kormann, D. W. Bahnemann and M. R. Hoffmann, Environ. Sci. Technol., 1988, 22, 798.
30. A. Torrents and A. T. Stone, Environ. Sci. Technol., 1993, 27, 1060.
31. J. Bandara and J. Kiwi, New J. Chem., 1999, 23, 717.
32. N. Serpone, J. Martin, S. Horikoshi and H. Hidaka, J. Photochem. Photobiol., A, 2005, 169, 235.
33. A. H. Boonstra and C. A. H. A. Mutsaers, J. Phys. Chem., 1975, 79, 1694.
34. D. Vione, C. Minero, V. Maurino, M. E. Carlotti, T. Picatonotto and E. Pelizzetti, Appl. Catal., B, 2005, 58, 79.
35. J. A. R. van Veen, F. T. G. Veltmaat and G. Jonkers, J. Chem. Soc., Chem. Commun., 1985, 1656.
36. D. M. Stanbury, Adv. Inorg. Chem., 1989, 33, 69.
37. D. Lawless, D. Meisel and N. Serpone, J. Phys. Chem., 1991, 95, 5166.
38. G. V. Buxton, C. L. Greenstock, W. P. Helman and A. B. Ross, J. Phys. Chem. Ref. Data, 1988, 17, 513.
39. W. Choi and M. R. Hoffmann, Environ. Sci. Technol., 1995, 29, 1646.
40. A. J. Hoffman, E. R. Carraway and M. R. Hoffmann, Environ. Sci. Technol., 1994, 28, 776.
41. C. M. Wang and T. E. Mallouk, J. Phys. Chem., 1990, 94, 4276.
42. F. Forouzan, T. C. Richards and A. J. Bard, J. Phys. Chem., 1996, 100, 18123.
43. G. T. Brown and J. R. Darwent, J. Chem. Soc., Faraday Trans. 1, 1984, 80, 1631.
44. T. Wu, G. Liu, J. Zhao, H. Hidaka and N. Serpone, J. Phys. Chem. B, 1999, 103, 4862.
45. L. Cermenati, P. Pichat, G. Guillard and A. Albini, J. Phys. Chem. B, 1997, 101, 2650.
46. F. Shiraishi and C. Kawanishi, J. Phys. Chem. B, 2004, 108, 10491.
47. G. V. Buxton and R. M. Sellers, J. Chem. Soc., Faraday Trans. 1, 1973, 69, 555.
48. E. R. Carraway, A. J. Hoffman and M. R. Hoffmann, Environ. Sci. Technol., 1994, 28, 786.

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