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The Effects of Oxyanion Adsorption on Reactive Oxygen Species Generation by Titanium Dioxide

Published online by Cambridge University Press:  01 January 2024

Mary R. Arenberg  and
Yuji Arai
Mary R. Arenberg
Affiliation:
Department of Natural Resources and Environmental Sciences, University of Illinois at Urbana-Champaign, 61801, Urbana, IL, USA
Yuji Arai*
Affiliation:
Department of Natural Resources and Environmental Sciences, University of Illinois at Urbana-Champaign, 61801, Urbana, IL, USA
*
*E-mail address of corresponding author: yarai@illinois.edu
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Abstract

The growing use of nano titanium dioxide (TiO2) in consumer and agricultural products has accelerated its introduction into terrestrial environments, where its impact has not been documented extensively. TiO2 toxicity arises primarily from its ability to photochemically generate reactive oxygen species (ROS), including hydrogen peroxide (H2O2). While common ligands in soil porewaters can either hinder or enhance the degradation of organic contaminants by TiO2, their effects on ROS production by TiO2 have not been understood clearly. The objective of this study was to assess the effect of phosphate (P) and nitrate on UV-irradiated anatase, nano-TiO2. Accordingly, H2O2-generation kinetics experiments were conducted in UV-irradiated TiO2 under environmentally relevant concentrations of the ligands (0, 50, 100, and 250 μM) and pH values (4.00 ± 0.02 and 8.00 ± 0.02) from 0–100 min. Under all conditions, H2O2 grew logarithmically and reached between 5.38 and 22.98 μM after 100 min. At pH 4.00 ± 0.02, H2O2 production was suppressed by P but not by nitrate. Conversely, at pH 8.00 ± 0.02, nitrate did not affect H2O2 concentration while P increased it. Non-specific, minimal adsorption of nitrate prevented interference with the photoreactivity of TiO2. Due to the pH-dependent behavior of suspended TiO2 and H2O2 degradation rates, specific adsorption of P on TiO2 impeded its ability to produce H2O2 photochemically at pH 4.00 ± 0.02 but amplified it at pH 8.00 ± 0.02.

Keywords

AdsorptionAnataseNitratePhosphatePhotocatalysisROSTiO2

Information

Type
Original Paper
Information
Clays and Clay Minerals , Volume 67 , Issue 5 , October 2019 , pp. 410 - 418
Copyright
Copyright © Clay Minerals Society 2019

Introduction

Engineered nanomaterials (ENM) are utilized in various consumer products, including cosmetics, plastics, fuel additives, textiles, paints, coatings, dietary supplements, medical applications, etc. (Keller & Lazareva Reference Keller and Lazareva2013; Manke et al. Reference Manke, Wang and Rojanasakul2013; Sajid et al. Reference Sajid, Ilyas, Basheer, Tariq, Daud, Baig and Shehzad2015). During their production, use, and disposal each year, ~51,600 metric tons of ENMs are released into the soil globally (Keller & Lazareva Reference Keller and Lazareva2013). Within terrestrial environments, ENMs have both beneficial and detrimental impacts (Gardea-Torresdey et al. Reference Gardea-Torresdey, Rico and White2014; Sajid et al. Reference Sajid, Ilyas, Basheer, Tariq, Daud, Baig and Shehzad2015). The positive effects of ENMs primarily involve increased yields of various crops and enhanced degradation of toxic chemicals (Ravichandran et al. Reference Ravichandran, Selvam and Swaminathan2010; Gardea-Torresdey et al. Reference Gardea-Torresdey, Rico and White2014). Conversely, the risks of ENMs include cyto- and genotoxicity to humans, animals, and plants as well as decreased crop biomass (Ghosh et al. Reference Ghosh, Bandyopadhyay and Mukherjee2010; Manke et al. Reference Manke, Wang and Rojanasakul2013; Fu et al. Reference Fu, Xia, Hwang, Ray and Yu2014; Gardea-Torresdey et al. Reference Gardea-Torresdey, Rico and White2014; Sajid et al. Reference Sajid, Ilyas, Basheer, Tariq, Daud, Baig and Shehzad2015). These impacts vary widely depending on the type and concentration of the ENM and the type of organism exposed (Manke et al. Reference Manke, Wang and Rojanasakul2013).

Titanium dioxide is one of the most common ENMs and is frequently released into the environment through cosmetic, agricultural, medicinal, and food-additive applications (Weir et al. Reference Weir, Westerhoff, Fabricius, Hristovski and Von Goetz2012; Fu et al. Reference Fu, Xia, Hwang, Ray and Yu2014; Gondikas et al. Reference Gondikas, Von Der Kammer, Reed, Wagner, Ranville and Hofmann2014; Sajid et al. Reference Sajid, Ilyas, Basheer, Tariq, Daud, Baig and Shehzad2015). Unlike naturally occurring TiO2 particles, TiO2 environmental nanoparticles (ENPs) are manufactured to be homogeneous in size, shape, and structure. They also have smaller diameters and greater reactive surface area than natural TiO2 (Bernhardt et al. Reference Bernhardt, Colman, Hochella, Cardinale, Nisbet, Richardson and Yin2010). The homogeneity, purity, and large reactive surface area make engineered TiO2 NPs more reactive. Smaller particles become more photoreactive because of an increase in band-gap energies (Kavan et al. Reference Kavan, Stoto, Grätzel, Fitzmaurice and Shklover1993). Their release into soil and water environments, in particular, is extensive because they are utilized in various agricultural applications, including soil remediation nano-technologies and nano-enabled pesticides, fertilizers, soil additives, and growth regulators (Gardea-Torresdey et al. Reference Gardea-Torresdey, Rico and White2014). Additionally, TiO2 is added to the soil through the application of biosolids, which tend to have high concentrations of TiO2 (10–70 mg kg–1) (Keller & Lazareva Reference Keller and Lazareva2013; Gardea-Torresdey et al. Reference Gardea-Torresdey, Rico and White2014). Consequently, soil TiO2 concentrations can be significant. For instance, an exposure model predicted that nano-TiO2 concentrations in Swiss soils fall between 0.4 and 4.8 μg kg–1, depending on the extent of the emissions (Mueller et al. Reference Mueller, Som and Nowack2009). In the USA, soil nano-TiO2 increases by 0.43–2.13 μg kg–1 y–1 across all soil types, and by 34.5–170 μg kg–1 y–1 in sludge-treated soils (Gottschalk et al. Reference Gottschalk, Sonderer, Scholz and Nowack2009).

Elevated TiO2 concentrations in soils and porewater can lead to adverse effects. Past studies have reported that plant exposure to TiO2 altered the chlorophyll content and antioxidative enzyme activities and led to Ti accumulation in fruit, roots, stems, and leaves (Servin et al. Reference Servin, Morales, Castillo-Michel, Hernandez-Viezcas, Munoz, Zhao, Nunez, Peralta-Videa and Gardea-Torresdey2013; Gardea-Torresdey et al. Reference Gardea-Torresdey, Rico and White2014). Additionally, TiO2 can negatively affect wheat growth and soil enzyme activity and lead to phytotoxicity and DNA fragmentation (Sajid et al. Reference Sajid, Ilyas, Basheer, Tariq, Daud, Baig and Shehzad2015).

Furthermore, TiO2 toxicity may also detrimentally affect aquatic and terrestrial organisms, including arthropods, nematodes, and earthworms. The toxicity of anatase TiO2 arises from its ability to generate reactive oxygen species when exposed to UV light (Ma et al. Reference Ma, Brennan and Diamond2012). Reactive oxygen species cause oxidative stress to various soil organisms. While TiO2 exposure can increase the feeding rates and catalase activities of the terrestrial arthropod Porcellio scaber (Drobne et al. Reference Drobne, Jemec and Pipan Tkalec2009), it can also destabilize the cell membranes in these arthropods’ isolated digestive glands (Valant et al. Reference Valant, Drobne, Sepčić, Jemec, Kogej and Kostanjšek2009). Moreover, the fertility and survival of nematode Caenorhabditis elegans was compromised by exposure to nano-TiO2 (Wang et al. Reference Wang, Wick and Xing2009). In fact, the degree of those effects were inversely correlated with the size of the TiO2 particle (Roh et al. Reference Roh, Park, Park and Choi2010). Nano-TiO2 with a diameter <30 nm can enter plant cell-wall pores and be taken up into the roots. In roots, TiO2 can interrupt water flow and hydraulic potential. It can also be translocated from the roots into the leaves, where it influences pigments (i.e. chlorophyll, anthocyanin, and carotenoid) (Tan et al. Reference Tan, Peralta-Videa and Gardea-Torresdey2018). Anatase TiO2 accelerates the nitrogen metabolism of spinach by increasing the activities of nitrate reductase, glutamate dehydrogenase, glutamine synthase, and glutamic–pyruvic transaminase and by promoting the absorption of inorganic nitrogen and its subsequent conversion to organic forms (Yang et al. Reference Yang, Hong, You, Liu, Gao, Wu and Yang2006). Finally, TiO2 can bio-accumulate in earthworms and lead to oxidative stress, DNA and mitochondrial damage, and inhibition of cellulase (Hu et al. Reference Hu, Li, Cui, Li, Chen and Yang2010).

Many of these risks arise because TiO2 generates reactive oxygen species (ROS) under exposure to UV-light (Skocaj et al. Reference Skocaj, Filipic, Petkovic and Novak2011; Fu et al. Reference Fu, Xia, Hwang, Ray and Yu2014). Specifically, UV-irradiated TiO2 produces H2O2 in aqueous solutions (Pappas & Fischer Reference Pappas and Fischer1975; Konaka et al. Reference Konaka, Kasahara, Dunlap, Yamamoto, Chien and Inoue1999). When exposed to UV-radiation, TiO2 photochemically generates e–, which contributes to the reduction of O2 into H2O2, as seen below, in Eq. 1 (Burek et al. Reference Burek, Bahnemann and Bloh2019):

(1) O 2 → + e − · O 2 − → + H + · HO 2 → + e − · HO 2 − → + H + H 2 O 2

Subsequently, H2O2 can either be re-oxidized to O2 or be further reduced to water (Burek et al. Reference Burek, Bahnemann and Bloh2019).

These adverse effects have led lawmakers throughout the world to take different regulatory actions (Amenta et al. Reference Amenta, Aschberger, Arena, Bouwmeester, Botelho Moniz, Brandhoff, Gottardo, Marvin, Mech, Quiros Pesudo, Rauscher, Schoonjans, Vettori, Weigel and Peters2015; Gupta & Xie Reference Gupta and Xie2018). The World Health Organization has classified the carcinogenicity of TiO2 into group 2B, deeming it “possibly carcinogenic to humans” (World Health Organization 2019). Furthermore, the European Chemicals Agency is seriously considering classifying TiO2 as a class 1B carcinogen, deeming it “presumed to have carcinogenic potential for humans” (ANSES 2016; European Chemicals Agency Reference French, Jacobson, Kim, Isley, Penn and Baveye2017). The US Food and Drug Administration mandates that TiO2 makes 1.0% or less of food by weight (U.S. Food and Drug Administration 2019) while the US National Institute for Occupational Safety and Health has recommended that TiO2 be classified as a potential occupational carcinogen and that exposure be limited to 0.2 mg m–3 (National Institute for Occupational Safety and Health 2011).

Moreover, photocatalysis of TiO2 may be influenced by common, nearly ubiquitous environmental oxyanions in soil porewaters, including nitrate and phosphate). While the TiO2 photocatalytic reactions have been studied extensively in basic research, how environmentally relevant conditions (e.g. ligand interaction) affect its photoreactivity is not well understood. Researchers previously reported that at pH 7.9, large concentrations (5–25 mmol L–1) of a common ligand, phosphate (P), decreased hydroxyl radical (∙OH) generation by TiO2 in freshwaters while nitrate decreased∙OH generation to a lesser extent ( Budarz et al. Reference Budarz, Turolla, Piasecki, Bottero, Antonelli and Wiesner2017). Others have concluded that P adsorption promoted the photocatalytic degradation of organic compounds bound weakly to TiO2 while also inhibiting degradation of strongly sorbed organic compounds (Zhao et al. Reference Zhang, Lima, Lee, Zaera, Chi and Yin2008). Phosphorus adsorption accelerated organic contaminant degradation at pH 6.2 in another study (Long et al. Reference Long, Brame, Qin, Bao, Li and Alvarez2017). Furthermore, the presence of large concentrations of P (≥200 mM) under above neutral pH conditions boosted H2O2 production by TiO2 (Xiong et al. Reference Xiong, Zhang, Liu, Zhao and Xu2018; Burek et al. Reference Burek, Bahnemann and Bloh2019).

The purpose of the present study was to investigate the effect of common ligands (e.g. nitrate and phosphate) on the generation of H2O2 by UV-irradiated TiO2. The results should have implications for better assessing the toxicity of TiO2 in the aquatic and terrestrial environment. The consideration of lower, environmentally relevant nitrate and phosphate concentrations (i.e. 50 μM, 100 μM, 250 μM) and various pH conditions (4.00 ± 0.02 and 8.00 ± 0.02) differentiate the present study from previous ones. The hypothesis was that specific adsorption of P will suppress H2O2 production by UV-irradiated TiO2 while the non-specific adsorption of nitrate will not interfere with the photoreactivity of TiO2.

Materials and Methods

Materials

All reagents were American Chemical Society (ACS) grade and all solutions were prepared using ultrapure water.

Nanopowder TiO2 (Nanostructured and Amorphous Materials, Inc., Katy, Texas, USA) and anatase 99+% with a Brunauer, Emmett, and Teller (BET) surface area of 60 m2 g–1 (Nanostructured & Amorphous Materials, Inc., Texas, USA) were used in this study. Powder X-ray diffraction (XRD) was performed on TiO2 using a Siemens-Bruker D5000 XRD System (Bruker Corporation, Billerica, Massachusetts, USA). The voltage was adjusted to 40.0 kV and the current to 30.0 mA. The XRD angle ranged from 10 to 90°2θ, in accordance with previous studies (Sakthivel et al. Reference Sakthivel, Neppolian, Shankar, Arabindoo, Palanichamy and Murugesan2003; Djerad et al. Reference Djerad, Tifouti, Crocoll and Weisweiler2004; Anandan et al. Reference Anandan, Sathish Kumar, Pugazhenthiran, Madhavan and Maruthamuthu2008; Kathiravan & Renganathan Reference Kathiravan and Renganathan2009; Ravichandran et al. Reference Ravichandran, Selvam and Swaminathan2010). Furthermore, the scan speed was set at 1.2°2θ min–1 (Djerad et al. Reference Djerad, Tifouti, Crocoll and Weisweiler2004; Ravichandran et al. Reference Ravichandran, Selvam and Swaminathan2010). The 2θ position and intensity ratio of each peak on the collected spectra were compared to an anatase TiO2 reference pattern from the International Center for Diffraction Data (ICDD 1971).

Scanning electron microscopy (SEM) analysis was conducted on TiO2, using a Hitachi S-4700 High Resolution SEM. Because of its low conductivity, the sample was coated with gold-palladium (Au-Pd) before analysis. The conditions of the analysis included an accelerating voltage of 15.0 kV, an emission current of 10 μA, and magnification of 90,032. After the image was captured, the diameters of ~20 individual TiO2 particles were measured and averaged.

The zeta potential and particle size of the TiO2 was determined using a Malvern Zetasizer (Malvern Instruments, Malvern, Worcestershire, UK), using laser Doppler velocimetry and dynamic light scattering (DLS) techniques, respectively. TiO2 samples were prepared with suspension densities of 0.1 g L–1 and background solution of 0.001 M NaCl and in the presence or absence of 10 μM or 25 μM NO3 – or P, using sodium nitrate (NaNO3) and sodium phosphate (Na2HPO4), respectively. The samples were shaken on an end-over shaker at 50 rpm for 48 h and adjusted to various pH conditions, using 0.01 M HCl and NaOH. The refractive index of 2.488 was used (Zhang et al. Reference Zhang, Lima, Lee, Zaera, Chi and Yin2011). Prior to the measurements, the zetasizer was calibrated with a polystyrene latex standard.

Methods

Adsorption of Phosphate and Nitrate by TiO2

TiO2 suspensions (suspension density: 1 g L–1) were prepared in 0.001 M NaCl (Fisher Scientific International, Inc., Fair Lawn, New Jersey, USA) and immediately sonicated for 30 s. The suspensions were kept in the dark at ambient temperature on a stir plate at 300 rpm for 24 h. The pH values of the solutions were adjusted to 4.00 ± 0.02 or 8.00 ± 0.02, using 0.01 M HCl (Fisher Scientific International, Inc., Fair Lawn, New Jersey, USA) and NaOH (Fisher Scientific International, Inc., Fair Lawn, New Jersey, USA). After the pH had stabilized, Na2HPO4 (Fisher Scientific International, Inc., Fair Lawn, New Jersey, USA) also prepared in 0.001 M NaCl and adjusted to pH 4.00 ± 0.02 or 8.00 ± 0.02, was added to the treatment solutions to bring the final P concentrations of the eight samples to 0, 25, 50, 75, 100, 150, 200, and 250 μM, respectively. The solutions were mixed on a reciprocal shaker at 80 rpm, and aliquots of 10 mL were removed at 48 h and passed through a 0.2 μm polyvinylidene fluoride (PVDF) syringe filter. Preliminary centrifugation experiments of filtrate showed that quantifiable nanoparticles did not pass through the filter. Solution P concentrations were measured colorimetrically using the molybdenum blue method (Murphy & Riley Reference Murphy and Riley1962). The same procedure was carried out to determine nitrate adsorption, except NaNO3 (Fisher Scientific International, Inc., Fair Lawn, New Jersey, USA) was added to bring the concentration of the samples to 0, 50, 100, 150, 200, and 250 μM, respectively. Nitrate concentrations were determined using UV spectrometry at 220 nm (Patey et al. Reference Patey, Rijkenberg, Statham, Stinchcombe, Achterberg and Mowlem2008).

UV-irradiation of TiO2 and H2O2 Determination

TiO2 solutions (suspension density: 1 g L–1) were prepared in 0.001 M NaCl and 10% 2-propanol (VWR, Radnor, Pennsylvania, USA) and adjusted to pH 4.00 ± 0.02 or 8.00 ± 0.02, as described above. The addition of 2-propanol optimized the analytical determination of H2O2 by providing additional e– for the reduction of O2 to H2O2 and suppressing H2O2 degradation (Burek et al. Reference Burek, Bahnemann and Bloh2019). The reaction vessels, 250-mL beakers wrapped in aluminum foil, were filled with 100 mL of TiO2 solution and placed on a magnetic stir plate set to 350 rpm. A 60 W UV-A light bulb (Feit Electric Company, Pico Rivera, California, USA) (λ: 320–400 nm) was positioned 5.5 cm above the solution surface, which had an area of 36.32 cm2. Aliquots of 3 mL were removed from the samples at 0, 1, 5, 10, 15, 20, 30, 40, 60, 80, and 100 min and filtered through the 0.2 μm PVDF syringes. Immediately following filtration, the process of H2O2 determination was carried out on the aliquots, as described below.

Hydrogen peroxide was determined fluorometrically using homovanillic acid (HVA) (Acros Organics, New Jersey, USA) as a substrate (Guilbault et al. Reference Guilbault, Kramer and Hackley1967, Reference Guilbault, Brignac and Zimmer1968; Staniek & Nohl Reference Staniek and Nohl1999; Khosravi et al. Reference Khosravi, Vossoughi, Shahrokhian and Alemzadeh2013; Paital Reference Paital2014). Solutions, prepared in ultrapure water, included 0.025 M tris hydroxymethyl aminomethane (tris) buffer (Macron Fine Chemicals, Center Valley, Pennsylvania, USA) , adjusted to pH 7.5; 10 U (Universal Enzyme unit in μmol/min is defined as the amount of the enzyme that catalyzes the conversion of 1 μM of substrate per min under the specified conditions) horseradish peroxidase (HRP) (Alfa Aesar, Ward Hill, Massachusetts, USA), and 125 μM HVA. To determine H2O2, solutions were combined in dark 15-mL vessels in the following order: 6.6. mL tris buffer, 0.9 mL HRP, 0.6 mL HVA, and 0.9 mL irradiation aliquot (Khosravi et al. Reference Khosravi, Vossoughi, Shahrokhian and Alemzadeh2013). Solutions were gently shaken and left to incubate for 2 h (Paital Reference Paital2014). Then, each sample was transferred to a quartz fluorescence cuvette. The sample was illuminated with an excitation wavelength of 315 nm (Ci & Wang Reference Ci and Wang1991; Paital Reference Paital2014). Fluorescence emissions at 425 nm were measured using an Ocean Optics UV-Vis spectrophotometer (Ocean Optics, Inc., Largo, Florida, USA).

Prior to measuring the H2O2 concentration of irradiated samples, H2O2 polynomial standard curves were prepared, using 0–15 μM H2O2 standards. All H2O2 standards were made in 0.001 M NaCl and 10% 2-propanpol. The effect of background oxyanions on H2O2 measurements was tested by making additional standard curves using 50, 100, and 250 μM Na2HPO4 and 50, 100, and 250 μM NaNO3 background solutions.

Kinetic and Statistical Analyses

To analyze the rate of H2O2 production under the different pH and oxyanion conditions, time vs H2O2 data for each treatment were substituted into the logarithmic regression curve (i.e. y = a ∙ ln(x) + b) and the statistical fitting parameters were determined.. At a given time x, H2O2 increased at a rate of a∙x –1 μM min–1. Coefficients of determination (R2 values) were calculated for each regression curve. The H2O2 data were analyzed using two sample t-tests to determine if the population means of H2O2 varied significantly at 100 min based on treatment effect (i.e. the presence or absence of nitrate or phosphate).

Results

TiO2 Characterization

Various analyses were utilized to characterize TiO2. The particle size of TiO2 (0.1 g L–1) in 0.001 M NaCl, measured by DLS, was 1040.67 ± 107.30 nm at pH 4.00 ± 0.05 and 808.57 ± 76.15 nm at pH 8.00 ± 0.05 (Fig. 1a). In 10 μM P, the particle size of TiO2 was 2071.67 ± 186.92 nm at pH 4.00 ± 0.05 and 764.73 ± 7.81 nm at pH 8.00 ± 0.05. In 25 μM P, the particle size of TiO2 was 2141.67 ± 189.87 nm at pH 4.00 ± 0.05 and 498.00 ± 28.52 nm at pH 8.00 ± 0.05.

Fig. 1.

Characterization of TiO2: ( a) particle size diameter (nm) of TiO2 (0.1 g L–1) in 0.001 M NaCl, as affected by pH and phosphate concentration; (b) X-ray diffraction pattern of anatase TiO2 powder, with d-spacing values (Å) displayed over major peaks; (c) SEM image of anatase TiO2 powder; and (d) zeta potential (mV) of TiO2 (0.1 g/L) in 0.001 M NaCl in the presence and absence of phosphate and nitrate, as affected by pH (n = 3).

XRD patterns of the TiO2 are shown in Fig. 1b. The XRD reflections (12.65, 18.47, 24.02, 26.94, 27.53, 31.34, 70.29, 35.14, 37.043, and 41.34°2θ) confirmed the presence of anatase. A SEM image of anatase TiO2 powder revealed that the average individual particle size was 35.89 ± 4.18 nm (n = 19) (Fig. 1c).

Zeta Potential of TiO2 in the Presence and Absence of Oxyanions

From approximately pH 3 to 9.5, the zeta potential of TiO2 (0.1 g L–1) in 0.001 M NaCl followed very similar trends in the presence of 10 μM and 25 μM nitrate and the absence of any oxyanion ligands (Fig. 1d). For instance, the zeta potential was 11.42 ± 1.34 mV at pH 4.58 in the control, 11.67 ± 0.42 mV at pH 4.55 in 10 μM nitrate, and 13.73 ± 0.91 mV at pH 4.69 in 25 μM nitrate. The isoelectric point (IEP) occurred at around pH 5.5 in all three samples.

In the presence of 10 μM and 25 μM P, however, the zeta potential of TiO2 deviated greatly and was much smaller compared to the other conditions from approximately pH 3 to 7. For instance, the zeta potential was –20.43 ± 1.07 mV at pH 4.76 in 10 μM P and –19.97 ± 0.58 mV at pH 4.40 in 25 μM P. In 10 μM P, the IEP of TiO2 shifted to ~pH 3.25 while, in 25 μM P, the isoelectric point occurred at a pH lower than that of the most acidic sample. Nevertheless, at ~pH 7, the zeta potential of TiO2 in both P concentrations became similar. At pH 9.15 ± 0.23, the zeta potential of TiO2 in all conditions averaged –37.34 ± 0.89 mV.

Adsorption of Phosphate and Nitrate at the TiO2–Water Interface

Adsorption of P by TiO2 was greater at pH 4.00 ± 0.02, than at pH 8.00 ± 0.02. This was anticipated, as the positively charged TiO2 attracts negatively charged P (Connor & McQuillan Reference Connor and McQuillan1999). At pH 8.00 ± 0.02, the surface of TiO2 is negatively charged, weakening this attraction. The trend line of C eq vs adsorbed q for P (Fig. 2A) followed y = 0.261 ln(x) + 2.322 (R2 = 0.9647) at pH 4.00 ± 0.02 and y = 0.110 ln(x – 0.015) + 1.331 (R2 = 0.9950) at pH 8.00 ± 0.02.

Fig. 2.

Adsorption of (a) phosphate and (b) nitrate by TiO2 (1 g L–1) in 0.001 M NaCl, as affected by pH (n = 3).

Conversely, adsorption of nitrate by TiO2 was low at pH 4.00 ± 0.02 and did not occur at pH 8.00 ± 0.02. At pH 4.00 ± 0.02, the trend line of C eq vs adsorbed q for nitrate followed y = 9.097 x 2 – 0.00601 x + 0.00165 (R2 = 0.9043) (Fig. 2b).

H2O2 Generation by UV-Irradiated TiO2

UV-irradiated TiO2 generated H2O2 to different degrees, depending on pH and the type and concentration of oxyanion ligand (Fig. 3; Table 1). In the control at pH 4.00 ± 0.02, H2O2 increased at a rate of 4.1494 x –1 μM min–1 and reached 18.61 μM at 100 min. The addition of P suppressed this production. At pH 4.00 ± 0.02, the rates of H2O2 generation were 1.76 x –1 μM min–1 in the presence of 50 μM P, 0.84 x –1 μM min–1 in the presence of 100 μM P, and 1.99 x –1 μM min–1 in the presence of 250 μM P. H2O2 reached 8.70 μM, 5.38 μM, and 9.85 μM at 100 min in 50 μM P, 100 μM P, and 250 μM P, respectively. After 100 min of irradiation, H2O2 was significantly (p < 0.01) less in the presence of P compared to its absence.

Fig. 3.

H2O2 generation by UV-irradiated nano TiO2 over the course of 100 min, as affected by various concentrations of (a) phosphate at pH 4, (b) nitrate at pH 4, (c) phosphate at pH 8, and (d) nitrate at pH 8 (n = 2)

Table 1.

H2O2 generation rates, where x = time, in min, by UV-irradiated nano TiO2, as affected by pH and oxyanion type and concentration. Solutions were prepared in 0.001 M NaCl and 10% 2-propanol, in addition to the oxyanion solutions below, and then adjusted to the specified pH

Compared to P, the addition of nitrate did not cause the same inference in photocatalysis. At pH 4.00 ± 0.02, the rates of H2O2 generation were 4.15 x –1 μM min–1 in the presence of 50 μM nitrate, 4.19 x –1 μM min–1 in the presence of 100 μM nitrate, and 4.76 x –1 μM min–1 in the presence of 250 μM nitrate. At 100 min, H2O2 reached 22.98, 19.91, and 22.44 μM in 50 μM nitrate, 100 μM nitrate, and 250 μM nitrate, respectively. After 100 min of irradiation, H2O2 was not significantly different in the presence of nitrate compared to its absence.

At pH 8.00 ± 0.02, H2O2 concentration increased at markedly lower rates. In the control, H2O2 yield followed 0.99 x –1 μM min–1 and hit 5.79 μM at 100 min. In the presence of nitrate, H2O2 similarly reached 6.29, 5.58, and 5.53 μM, at 100 min in 50 μM nitrate, 100 μM nitrate, and 250 μM nitrate, respectively. The production follows the trends of 1.1430 x –1 μM min–1 in 50 μM nitrate, 1.1081 x –1 μM min–1 in 100 μM nitrate, and 0.97 x –1 μM min–1 in 250 μM nitrate. After 100 min of irradiation, H2O2 was not significantly different in the presence of nitrate compared to its absence.

In the presence of P at pH 8.00 ± 0.02, H2O2 grew at rates of 2.02 x –1 μM min–1, 2.08 x –1 μM min–1, and 1.51 x –1 μM min–1 and reached 10.04, 10.48, and 8.11 μM at 100 min in 50 μM P, 100 μM P, and 250 μM P, respectively. After 100 min of irradiation, H2O2 was significantly (p < 0.01) greater in the presence of P compared to its absence.

Discussion

Physicochemical Properties of TiO2

Physicochemical properties of TiO2 were studied using XRD, SEM, and zeta potential analyses. No differences were observed when the measured XRD pattern was compared to the anatase TiO2 reference from the ICDD database (ICDD 1971), confirming that the TiO2 used in the present study was in fact anatase (ICDD 1971). SEM analysis revealed that the diameter of the particles of the TiO2 powder was larger at 35.89 ± 4.18 nm (n = 19) than the size advertised by the vender (i.e. 10 nm). Furthermore, the SEM image illustrated the aggregation of the dry powder into large clusters, which is consistent with previous characterizations (French et al. Reference French, Jacobson, Kim, Isley, Penn and Baveye2009; Yin et al. Reference Yin, Liu, Ehrenshaft, Roberts, Fu, Mason and Zhao2012).

The zeta potential values of the control aligned with those reported previously for TiO2 under similar conditions (Nelson et al. Reference Nelson, Candal, Corn and Anderson2000; Dutta et al. Reference Dutta, Ray, Sharma and Millero2004; Kataoka et al. Reference Kataoka, Gurau, Albertorio, Holden, Lim, Yang and Cremer2004; Kim & Choi Reference Kim and Choi2011; Wan et al. Reference Wan, Yan, Liu, Tan, He and Feng2016). Below pH 5.5, TiO2 became increasingly positively charged as the surface became more and more protonated (i.e. TiOH2 +) (Kataoka et al. Reference Kataoka, Gurau, Albertorio, Holden, Lim, Yang and Cremer2004). The reverse was true above pH 5.5, as the surface became deprotonated (i.e. TiO–) (Kataoka et al. Reference Kataoka, Gurau, Albertorio, Holden, Lim, Yang and Cremer2004). Moreover, in the presence of P, the zeta potential and IEP of TiO2 shifted significantly downward due to the specific adsorption of P (Nelson et al. Reference Nelson, Candal, Corn and Anderson2000; Wan et al. Reference Wan, Yan, Liu, Tan, He and Feng2016).

The Effect of Nitrate on H2O2 Generation by UV-Irradiated TiO2

H2O2 generation by UV-irradiated TiO2 was similar between the control and nitrate treatments at both pH 4.00 ± 0.02 and 8.00 ± 0.02, while adsorption of nitrate by TiO2 was minimal to zero. Nitrate is considered an indifferent ligand because it does not specifically interact with TiO2 (Budarz et al. Reference Budarz, Turolla, Piasecki, Bottero, Antonelli and Wiesner2017). This lack of specific interaction between nitrate and TiO2 was supported by zeta potential measurements, in which differences were not discernible between the control TiO2 and the TiO2 in 10 and 25 μM nitrate from pH 3 to 9.5. Consequently, the presence of nitrate did not effectively hinder or interfere with the ability of TiO2 to produce H2O2 photochemically.

The Effect of Phosphate on H2O2 Generation by UV-Irradiated TiO2

Phosphate strongly adsorbs to TiO2, including the anatase form used in the present study, through inner-sphere monodentate and bidentate complexes (Hadjiivanov et al. Reference Hadjiivanov, Klissurski and Davydov1989; Connor & McQuillan Reference Connor and McQuillan1999; Zhao et al. Reference Zhang, Lima, Lee, Zaera, Chi and Yin2008; Alshameri et al. Reference Alshameri, Yan and Lei2014). Past infrared spectroscopic studies revealed that multinuclear complexation of P at the TiO2 surfaces dominates (Connor & McQuillan Reference Connor and McQuillan1999; Gong Reference Gong2001; Ronson & McQuillan Reference Ronson and McQuillan2002).

Throughout the irradiation, adsorption of P by TiO2 remained relatively constant under both acidic and alkaline conditions. Within the eleven aliquots taken from 0 to 100 min, aqueous P averaged 0.54 ± 0.46 μM at pH 4 in the 50 μM treatment, 22.23 ± 0.91 μM at pH 4 in the 100 μM treatment, 217.63 ± 2.15 μM at pH 4 in the 250 μM treatment, 23.74 ± 1.97 μM at pH 8 in the 50 μM treatment, 70.52 ± 0.73 μM at pH 8 in the 100 μM treatment, and 224.13 ± 5.03 μM at pH 8 in the 250 μM treatment. Across the treatments, aqueous P varied, on average, 1.88 μM P during the 100 min irradiations.

The effects of P adsorption by TiO2 photoreactivity were more complex than nitrate and varied with pH. On one hand, much adsorption of P by TiO2 at pH 4.00 ± 0.02 corresponded to suppressed H2O2 generation. After 100 min of UV-irradiation, H2O2 was more than three times less in the presence of 100 μM P than in the absence of P. On the other hand, comparatively less adsorption of P by TiO2 at pH 8.00 ± 0.02 corresponded to enhanced H2O2 production. After 100 min of UV-irradiation, H2O2 was nearly two times greater in the presence of 100 μM P than in the absence of P. These results can be explained by the contrasting behavior of suspended TiO2 and H2O2 degradation under the different pH conditions.

Adsorbed P caused a shift in the behavior of suspended TiO2. At pH 4.00 ± 0.05, the particle size of TiO2 was approximately two times larger in the presence of 10 μM and 25 μM P than in the control. This is consistent with past research, which demonstrated that the adsorption of P (>50 μM) by TiO2 led to its aggregation at pH ~3, which is lower than the IEP of the pure solid, 6.3 (Wan et al. Reference Wan, Yan, Liu, Tan, He and Feng2016). Conversely, at pH 8.00 ± 0.05, the particle size of TiO2 in the presence of 10 μM and 25 μM P was smaller than in the control. Past researchers have similarly observed that the adsorption of P (>50 μM) by TiO2 under neutral to alkaline conditions (i.e. pH 7.0 and 7.9) facilitates the dispersion (Wan et al. Reference Wan, Yan, Liu, Tan, He and Feng2016; Budarz et al. Reference Budarz, Turolla, Piasecki, Bottero, Antonelli and Wiesner2017).

In the present work, specific adsorption of P at pH 4.00 ± 0.02 caused the TiO2 to aggregate and, thus, reduced its ability to generate H2O2 photochemically under all three concentrations (Connor & McQuillan Reference Connor and McQuillan1999; Fujishima et al. Reference Fujishima, Zhang and Tryk2008; Budarz et al. Reference Budarz, Turolla, Piasecki, Bottero, Antonelli and Wiesner2017). Inner-core particles within TiO2 aggregates will not be as accessible as outer-core particles. The degree to which H2O2 production was suppressed correlated positively with the magnitude of P adsorption (100 μM > 50 μM > 250 μM) during irradiation. Previous studies confirmed that the aggregation of nano-TiO2 reduces its photoreactivity, due to the reduction of reactive surface area and photon penetration (Maira et al. Reference Maira, Yeung, Lee, Yue and Chan2000; Hotze et al. Reference Hotze, Phenrat and Lowry2010; Jassby et al. Reference Jassby, Farner Budarz and Wiesner2012). Thus, the presence of P reduced the production of H2O2 by UV-irradiated TiO2 at pH 4.00 ± 0.02.

Conversely, at pH 8.00 ± 0.02, greater H2O2 was produced in the presence of P than in its absence. Interestingly, the extent to which H2O2 production was enhanced correlated with the degree of P adsorption (50 μM ~ 100 μM > 250 μM) during irradiation. The degradation of H2O2 at pH 8.00 ± 0.02 was approximately three times greater than at pH 4.00 ± 0.02 (Burek et al. Reference Burek, Bahnemann and Bloh2019). However, under alkaline conditions, the adsorption of P by TiO2 suppressed the adsorption of H2O2 onto TiO2 and, thus, hindered its photocatalytic degradation, enabling greater aqueous H2O2 accumulation in the presence of P than in its absence. This phenomenon has been documented by several researchers (Moon et al. Reference Moon, Kim, Bokare, Sung and Choi2014; Xiong et al. Reference Xiong, Zhang, Liu, Zhao and Xu2018). Moreover, the dispersion of TiO2 caused by P adsorption under alkaline conditions optimized its photoreactivity. Consequently, the presence of P increased aqueous H2O2 at pH 8.00 ± 0.02.

Environmental Implications and Uncertainties

The presence of ≥50 μM P at pH 4.00 ± 0.02 facilitated the aggregation of TiO2, reducing its photoreactivity and ability to produce H2O2. Because P is nearly ubiquitous within porewaters and surface waters throughout the world, TiO2 toxicity in terrestrial environments, under acidic conditions, may be of less concern than originally thought. However, water-soluble P varies widely from <1 to 1000 μM (Fried & Broeshart Reference Fried and Broeshart1967; Sharpley & Smith Reference Sharpley and Smith1989), so the extent of this suppression by P in soil environments likely varies widely. At the same time, P enhanced H2O2 generation at pH 8.00 ± 0.02 because P adsorption reduced the degradation of H2O2.

Additionally, dissolved nitrate was up to 103 times more prevalent in soil pore waters than P (Fried & Broeshart Reference Fried and Broeshart1967) and did not effectively suppress or increase H2O2 production. Lastly, the effects of other compounds (e.g. other anions, organic compounds, colloids) commonly found in the soil matrix and their complex interactions on the photochemistry of TiO2 are not yet understood and should be evaluated further. Thus, many uncertainties remain regarding the extent of ROS generation by TiO2 in terrestrial and aquatic environments.

Conclusions

The ability of nano-TiO2 to generate reactive oxygen species under UV light exposure raises concerns about its potentially adverse effects in soil, where agricultural applications and the disposal of consumer products have enabled its accumulation. In the present study, the production of H2O2 by UV-irradiated TiO2 was monitored in the presence and absence of common environmental oxyanion ligands (i.e. nitrate and P) of different concentrations under acidic and alkaline conditions. At pH 4.00 ± 0.02, aggregation, caused by the specific adsorption of P onto TiO2, suppressed the generation of H2O2. Conversely, adsorption of P at pH 8.00 ± 0.02 reduced H2O2 degradation, enabling higher H2O2 accumulation in solution. Nitrate, which does not specifically interact with TiO2, did not affect the photoreactivity of TiO2 at pH 4.00 ± 0.02 or 8.00 ± 0.02. The present work provided evidence that P positively or negatively influences TiO2 photoreactivity at concentrations that are environmentally relevant in soil porewaters, depending on pH, while these same small concentrations of nitrate do not influence the photoreactivity of TiO2.

Acknowledgements

The authors gratefully acknowledge the United States Department of Agriculture (Hatch #1002831-ILLU-875-939) for supporting this project financially.

Compliance with Ethical Standards

Conflict of Interest

The authors declare that we have no conflict of interest.

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Figure 0

Fig. 1. Characterization of TiO2: (a) particle size diameter (nm) of TiO2 (0.1 g L–1) in 0.001 M NaCl, as affected by pH and phosphate concentration; (b) X-ray diffraction pattern of anatase TiO2 powder, with d-spacing values (Å) displayed over major peaks; (c) SEM image of anatase TiO2 powder; and (d) zeta potential (mV) of TiO2 (0.1 g/L) in 0.001 M NaCl in the presence and absence of phosphate and nitrate, as affected by pH (n = 3).

Figure 1

Fig. 2. Adsorption of (a) phosphate and (b) nitrate by TiO2 (1 g L–1) in 0.001 M NaCl, as affected by pH (n = 3).

Figure 2

Fig. 3. H2O2 generation by UV-irradiated nano TiO2 over the course of 100 min, as affected by various concentrations of (a) phosphate at pH 4, (b) nitrate at pH 4, (c) phosphate at pH 8, and (d) nitrate at pH 8 (n = 2)

Figure 3

Table 1. H2O2 generation rates, where x = time, in min, by UV-irradiated nano TiO2, as affected by pH and oxyanion type and concentration. Solutions were prepared in 0.001 M NaCl and 10% 2-propanol, in addition to the oxyanion solutions below, and then adjusted to the specified pH