1 Introduction

As a kind of semiconductor material with efficient photocatalytic activity, nanosized titanium dioxide (nano-TiO2) possesses nontoxicity, high chemical stability, and excellent photocatalytic activity (Zhou et al. 2017; Goudarzi and Shahabi-Ghahfarrokhi 2018). Due to its attractive properties, nano-TiO2 could be widely applied in a variety of fields such as sterilization, air purification, and water treatment (Wang et al. 2015; Zhang et al. 2017). However, nano-TiO2 also has some disadvantages. First of all, the disadvantage of easy agglomeration of nano-TiO2 greatly limits its performance. Due to the large specific surface area and high surface energy of pure nano-TiO2 particles, which belong to thermodynamic unstable systems, the surface particles cluster together with each other due to van der Waals force and Coulomb force, so that the nano-TiO2 particles are agglomerated and lose the physical properties and functions when applied. Therefore, it is imperative to prepare highly dispersed nano-TiO2 to maintain its advantages in photocatalytic performance. Secondly, nano-TiO2 particles in the water system are easy to lose, difficult to recycle, and high in recycling cost (Liu et al. 2018). Therefore, it is necessary to find a suitable carrier to immobilize nano-TiO2 for recycling. In order to solve the above problems, many scholars have extensively studied supported nano-TiO2 photocatalysts.

At present, the carrier materials commonly loading nano-TiO2 mainly include silicas (Sosnov et al. 2017), polymers (Vukoje et al. 2016), glasses (Yusoff et al. 2016), and carbon materials (Das et al. 2018; Xu et al. 2016a, b; Liang et al. 2018). All these materials have the characteristics of good stability, high strength, large specific surface area, and firm load. Because of large surface area, interwoven network structure of yarns, inherent strength, and flexibility, the cotton fabrics (CFs) can provide a good bearing platform for nano-TiO2, which is conducive to the formation of highly dispersed TiO2, thus maximizing its dye degradation performance. At the same time, TiO2 also provides a choice to impart multifunctional properties of the CFs (Mishra and Butola 2017; Giesz et al. 2016).

The annual output of waste CFs is huge, and they cannot be used reasonably, which will not only pollute the environment but also waste resources greatly (Pensupa et al. 2017). Therefore, using CFs as the carrier of TiO2 not only has a wide source and low cost but also protects the environment and makes effective utilization of resources (Ma et al. 2018). As reported by Giesz et al., using the methods of sol-gel process and microwave treatment successfully prepared TiO2-anatase film on cotton and viscose fabric surfaces (Giesz et al. 2016).

In this work, nano-TiO2 synthesized in situ and loaded on CFs was successfully prepared by using a hydrothermal method and low-temperature calcination, with cheap and readily available CFs as carrier and tetrabutyl titanate (TBT) as raw material. The loaded nano-TiO2 had anatase crystal form, high dispersion, and superior catalytic activity.

2 Experiments

2.1 Raw Materials and Reagents

CFs were used as substrate for loading nano-TiO2. Nano-TiO2 was prepared from TBT as raw material. The dyes, methylene blue (MB), methyl orange (MO), and rhodamine B (RhB) are all of analytical reagent (AR) grade and are commercially available.

2.2 Preparation of Loaded-TiO2 Cotton Fabrics (TiO2/CFs)

According to the analysis of this experiment, there are many factors affecting photocatalytic degradation rate, such as hydrolysis temperature, hydrolysis time, hydrothermal temperature, hydrothermal time, calcination temperature, calcination time, dye type, dye concentration, solution pH value, illumination time, illumination intensity, and hydrogen peroxide (H2O2) dosage. After comprehensive consideration, the hydrolysis temperature, hydrothermal temperature, hydrothermal time, and calcination time were determined as the experimental factors for the preparation of the experiment and were marked as A, B, C, and D, respectively. Three-level and four-factor orthogonal experiments were carried out. The factors and levels are shown in Table 1.

Table 1 Factors and levels of orthogonal design
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The specific preparation process of the experiment was as follows:1.6 g CFs were immersed in 8 mL TBT for 24 h at room temperature, followed by taking out the CFs and squeezing away as much liquid as possible from the surface of the CFs. Then, loaded TBT CFs were quickly immersed into 20 mL deionized water to hydrolyze for 30 min. Subsequently, it was transferred to a polytetrafluoroethylene autoclave for hydrothermal treatment. After a certain period of time, the products were poured into a beaker, washed three times with deionized water (to remove unloaded or not firmly loaded TiO2 particles from the CFs surface), and dried at 60 °C. Finally, the products were calcined at 200 °C and cooled to obtain TiO2/CFs. A simulation of the preparation process is shown in Fig. 1.

Fig. 1
figure 1

The preparation of TiO2/CFs

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2.3 Photocatalytic Degradation of MB, MO, and RhB under UV Irradiation

MB, MO, and RhB were selected as model chemicals to evaluate the photocatalytic performance of TiO2/CFs. About 0.30 g of TiO2/CFs was placed separately in 60 mL of MB, MO, and RhB solution (10 mg/L). Before irradiation, these solutions were stirred (500 r/min) in the dark for 30 min to achieve the adsorption equilibrium (Wu and Long 2011). Then, they were exposed to UV radiation with the wavelength of 365 nm (UV-365) maintaining vigorous stirring (500 r/min). The MB, MO, and RhB solutions were taken out at regular intervals, and the absorbances were measured by a 721 visible spectrophotometer, at the maximum absorption wavelength of 664 nm (Phaltane et al. 2017), 464 nm (John et al. 2017), and 552 nm (Cai et al. 2018), respectively. Before measuring, the suspensions were filtered through an organic micropore filter (pore size 0.45 μm). Taking MB solution as an example, since the MB solution concentration C (mg/L) was proportional to the absorbance within a certain range as shown in Fig. 2, it could be approximately expressed as (Eq. (1)):

$$ A=0.12057\mathrm{C},{R}^2=0.99883. $$
(1)
Fig. 2
figure 2

Standard working curve of MB solution

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Therefore, the degradation rate D(%) could be calculated by replacing the concentration C with the absorbance A. Eventually, the degradation rate was calculated by the measured absorbance as follows (Eq. (2)) (Zhang and Yang 2012):

$$ D=\left[\left({A}_0-{A}_t\right)/{A}_0\right]\times 100\% $$
(2)

where A0 is the absorbance of the MB solution after dark reaction for 30 min; At is the absorbance of the MB solution irradiated for t.

3 Results and Discussions

3.1 Orthogonal Test

This experiment was a three-level and four-factor test, where only the effect of the four factors on the degradation rate was considered, ignorance of the interactions among the factors. Therefore, L9 (34) orthogonal table should be selected. The specific test conditions and results are shown in Table 2. The degradation rates of MB evaluated and recorded as DMB, as the final evaluation index.

Table 2 Specific test conditions and results
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As shown in Table 2, the orders for the average values (− K) of each factor were as follows: − KA2 > − KA1 > − KA3, − KB2 > − KB1 > − KB3, − KC3 > − KC1 > − KC2, and − KD1 > − KD3 > − KD2. Therefore, A2, B2, C3, and D1 could be determined as the optimal levels of factors A, B, C, and D, respectively. The combination of A2B2C3D1 was the optimal level combination in this experiment. This meant that the optimal preparation conditions were hydrolysis temperature 80 °C, hydrothermal temperature 150 °C, hydrothermal time 10 h, and calcination time 2 h. The range (R) calculation demonstrated that RB > RD > RA > RC. Hence, for MB, the greatest influence was hydrothermal temperature, followed by calcination time and hydrolysis temperature, while hydrothermal time was the least influence.

3.2 Characterization of Morphology and Structure

3.2.1 Scanning Electron Microscopy and Energy-Dispersive Spectrometer

Scanning electron microscopy (SEM) was used to characterize the surface morphology of the samples and the distribution of TiO2 particles on CFs as shown in Fig. 3. The chemical composition of TiO2/CFs was characterized by energy-dispersive spectrometer (EDS) as shown in Fig. 4.

Fig. 3
figure 3

SEM images of the two samples. a Pristine CFs. b TiO2/CFs

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Fig. 4
figure 4

EDS energy spectrum of TiO2/CFs

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Figure 3a is a SEM image of the original CFs, from which it could be observed that the surface had a small amount of natural impurities and warp (Yang et al. 2019), so it had a certain roughness and was beneficial to the loading of TiO2. Figure 3b is a SEM image of TiO2/CFs, from which it could be observed that CFs still maintained its original shape, and TiO2 was uniformly and firmly distributed on the surface of the CFs. This was because there were many free hydroxyl groups on the surface of CFs, which could react with hydroxyl radicals (·OH) on the surface of TiO2, so the combination between them was strong enough. Through EDS energy spectrum analysis, it could be found that the prepared TiO2/CFs contained C, Ti, and O elements, which further confirmed the in situ synthesis of TiO2 on CFs.

3.2.2 X-Ray Diffraction

X-ray diffraction (XRD) was carried out to determine the crystal structure (brookite, anatase, or rutile) of TiO2 particles on CFs at 2θ from 10° to 70° with Cu Kα radiation, as shown in Fig. 5. Figure 5a showed the XRD pattern of the pristine CFs as blank samples, from which only the characteristic peaks of CFs were observed. Figure 5b shows the XRD pattern of the prepared TiO2/CFs, from which not only the characteristic diffraction peaks 15.6°, 16.8°, and 23.1° of CFs but also the characteristic diffraction peaks 25.6°, 38.2°, 48.2°, 54.3°, and 62.2° of typical anatase TiO2 were observed, corresponding to the crystal faces of (101), (112), (200), (211), and (204) (Sun et al. 2000; Zhou et al. 2017). In addition, using the XRD diffraction pattern of Fig. 4b, the crystallite diameter (D) of anatase TiO2 crystal was calculated based on (101) crystal face reflection according to the Scherer formula (Eq. (3)) (Kinsinger et al. 2011; Sun et al. 2000):

$$ D= k\lambda /\left(\upbeta \cos \theta \right) $$
(3)

where k is the shape constant (typical around 0.9) (Tang et al. 2009), λ is the wavelength of the Cu Kα radiation (λ = 0.154 nm), β is the full width at half maximum (FWHM), and θ is the diffraction angle. The TiO2 particle size was about 9.4 nm. Therefore, TiO2 was a nanosized particle.

Fig. 5
figure 5

The XRD patterns of the samples. a Pristine CFs. b TiO2/CFs

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From above, it could be determined that in situ synthesis of anatase nano-TiO2 particles on CFs could be achieved by the hydrothermal method and low-temperature (200 °C) calcination. Therefore, in order to make use of the photocatalytic properties of nano-TiO2, the prepared TiO2/CFs would be used in the research of dye-simulated wastewater treatment in the subsequent work.

3.3 Photocatalytic Degradation Performance

3.3.1 Effects of UV Irradiation Time on Photocatalytic Activity

Figure 6 shows the effects of irradiation time on photocatalytic performance of TiO2/CFs under the conditions of UV-365 irradiation, MB solution concentration of 10 mg/L, and TiO2/CFs dosage of 4.8 g/L. As can be seen from Fig. 6, the degradation rate of pristine CFs to MB solution in 9 h under UV-365 irradiation was about 15%, which was basically negligible in comparison with the higher degradation rate of TiO2/CFs. For TiO2/CFs, the degradation efficiency of MB solution gradually decreased with the increase of irradiation time, but the degradation rate gradually increased. This was because the intermediate products produced in the process of MB degradation occupied the active parts of TiO2, and a competition for degradation may occur between the MB and the intermediate products (Rauf et al. 2011), resulting in a decrease in degradation efficiency. Especially after 5 h, the degradation efficiency of MB solution was only 1.6%. However, with the gradual reaction of intermediate products and the continuous adsorption and degradation of MB solution, the degradation rate showed an increasing trend. However, after 5 h, the degradation rate also tended to be basically balanced.

Fig. 6
figure 6

The effects of illumination time on degradation rate of MB solution under UV-365

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3.3.2 Effects of TiO2/CFs Dosage on Photocatalytic Activity

Figure 7 reveals the effects of TiO2/CFs dosage on photocatalytic performance of TiO2/CFs under the conditions of UV-365 irradiation for 5 h and MB solution concentration of 10 mg/L. When the dosage of TiO2/CFs was less than 4.8 g/L, the degradation rate of MB solution increased with the increase of TiO2/CFs. This was because the more TiO2/CFs could provide, the more active sites there are and the more dyes could be adsorbed and degraded. However, when the dosage of TiO2/CFs exceeded 4.8 g/L, the degradation rate of MB solution did not significantly increase with the addition of TiO2/CFs. This was because the presence of excess TiO2/CFs would lead to a decrease in light transmittance, thus limiting the irradiated areas of TiO2/CFs. Therefore, for this experiment, the optimum dosage of TiO2/CFs was 4.8 g/L.

Fig. 7
figure 7

The effects of TiO2/CFs dosage on degradation rate of MB solution

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3.3.3 Effects of Dye Concentration on Photocatalytic Activity

Figure 8 displays the effects of MB concentration on photocatalytic performance of TiO2/CFs under the conditions of UV-365 irradiation for 5 h and TiO2/CFs dosage of 4.8 g/L. As can be seen from Fig. 8, the degradation rate of MB solution decreased with the increase of initial concentration of MB solution. The reason was that when the concentration of MB solution was relatively low, the number of dyes in the solution was relatively small, so there were relatively few dyes migrating and enriching to the TiO2/CFs surface. However, there were relatively many active sites on the surface of TiO2/CFs, and fewer dyes were not enough to enable all active sites of TiO2/CFs to exert the catalytic ability. When the concentration of MB solution was 5 mg/L, the degradation rate of TiO2/CFs to MB solution was the highest, reaching 98%. With the increase of MB solution concentration, the amount of MB adsorbed on the TiO2/CFs surface gradually reached saturation, and the catalytic capacity also reached maximum. In addition, intermediate products that were difficult to decompose would also accumulate around TiO2/CFs due to lack of time for reaction, hindering the progress of photocatalytic reaction. Therefore, the prepared TiO2/CFs were suitable for the degradation of low concentration dyes (5 mg/L, 10 mg/L, and 15 mg/L).

Fig. 8
figure 8

The effects of dye concentration on degradation rate of MB solution

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3.3.4 Effects of pH Value on Photocatalytic Activity

Figure 9 describes the effects of initial pH value on photocatalytic performance of TiO2/CFs under the conditions of UV-365 irradiation for 5 h, TiO2/CFs dosage of 4.8 g/L, and MB concentration of 10 mg/L. It could be seen that the degradation of MB by TiO2/CFs first increased and then decreased with the increase of pH value in the pH range of 4.0–11.0. The isoelectric point of TiO2 in water was about pH = 6.0 (Devipriya and Yesodharan 2010), which meant that the surface of TiO2 was neutral. When the solution was pH < 6.0, the TiO2 surface had a net positive charge, as shown in Eq. (4) (Kaur et al. 2018). Because MB was a cationic dye, the existence of TiOH2+ was not conducive to the adsorption of MB by TiO2/CFs. So, the degradation of MB by TiO2/CFs was obviously inhibited under acidic conditions. When the solution was pH > 6.0, the TiO2 surface carried a net negative charge, as shown in Eq. (5) (Kaur et al. 2018). This was also conducive to the migration of holes to the surface of TiO2 and the formation of ·OH with strong oxidation by reacting with OH−, H2O, and other electron donors on the surface of TiO2. In addition, TiO2 was negatively charged in neutral and weakly alkaline solutions and electrostatically adsorbed MB cations (Azeez et al. 2018), thus greatly increasing the degradation of MB. Particularly, when the solution was pH = 7.0, the degradation rate of MB was the highest, up to 80.4%.

$$ \mathrm{TiOH}+{\mathrm{H}}^{+}\to {{\mathrm{TiOH}}_2}^{+}\ \left(\mathrm{pH}<6.0\right) $$
(4)
$$ \mathrm{TiOH}+{\mathrm{OH}}^{-}\to {\mathrm{TiO}}^{-}+{\mathrm{H}}_2\mathrm{O}\ \left(\mathrm{pH}>6.0\right) $$
(5)
Fig. 9
figure 9

The effects of the initial pH value on degradation rate of MB solution

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However, when the pH the solution was pH > 9.0, the degradation rate of MB solution decreased significantly. To sum up, TiO2/CFs was more beneficial to the degradation of MB under neutral or weakly alkaline conditions. Since the pH value of MB solution itself was about 7, there was no need to adjust the pH value later.

3.3.5 Effects of H2O2 Dosage on Photocatalytic Activity

As a kind of strong oxidative reagent and electron scavenger, the dosage of H2O2 is considered to be an important parameter affecting the photocatalytic degradation process because it is directly related to the amount of ·OH generated in the photocatalytic reaction.

First of all, in order to prove that H2O2 could promote TiO2 to degrade dyes, three experiments were designed as shown in Fig. 10. The curve (1) showed the degradation rate of MB solution only in the presence of TiO2/CFs. The curve (2) showed the degradation rate of MB solution only in the presence of H2O2. The dashed line (1) + (2) was the sum of the degradation rates of MB solution under conditions (1) and (2) alone. The curve (3) showed the degradation of MB solution in the presence of TiO2/CFs and H2O2. As can be seen from Fig. 11, the degradation rate was obviously (3) > (1) + (2). Therefore, it came to a conclusion that H2O2 could promote TiO2 to degrade dyes.

Fig. 10
figure 10

The degradation rate of MB solution under different conditions

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Fig. 11
figure 11

The effects of H2O2 dosage on degradation rate of MB solution

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Subsequently, the effects of H2O2 dosage on photocatalytic performance of TiO2/CFs under the conditions of UV-365 irradiation for 2 h and 4 h, TiO2/CFs dosage of 4.8 g/L, MB concentration of 10 mg/L, and initial pH value of 7 were discussed, as shown in Fig. 11. According to the curve of illumination for 2 h (sealed), it could be concluded that in the presence of H2O2, the degradation rate of MB solution first increased and then tended to be stable with the increase of H2O2 dosage; the degradation efficiency of MB solution was the highest when 0.24 mol/L H2O2 was added to 60 mL MB solution. This was because when the dosage of added H2O2 was less than 0.24 mol/L, H2O2 could capture electrons on the surface of TiO2 (Eq. (6)), thus reducing the probability of electron-hole recombination and improving hole utilization (Eq. (7) and Eq. (8)) (Akpan and Hameed 2009). In addition, H2O2 may be decomposed directly to generate ·OH under UV irradiation (Eq. (9)) (Navarro et al. 2005), which was the strongest oxidant in aqueous phase systems and could degrade dyes into harmless products (Eq. (10) and Eq. (11)) such as CO2 and H2O. When the dosage of added H2O2 was above 0.24 mol/L, H2O2 made quenching reaction due to excess, i.e., H2O2 reacted with ·OH to form ·OOH (Eq. (12)) and the reaction (Eq. (13)) (Fayazi et al. 2016), thus reducing the degradation efficiency.

$$ {\mathrm{H}}_2{\mathrm{O}}_2+{\mathrm{e}}^{-}\to {\mathrm{O}\mathrm{H}}^{-}+\cdotp \mathrm{OH} $$
(6)
$$ {\mathrm{h}}^{+}+{\mathrm{OH}}^{-}\to \cdotp \mathrm{OH} $$
(7)
$$ {\mathrm{h}}^{+}+{\mathrm{H}}_2\mathrm{O}\to \cdotp \mathrm{OH}+{\mathrm{H}}^{+} $$
(8)
$$ {\mathrm{H}}_2{\mathrm{O}}_2+\mathrm{h}\upnu \to 2\cdotp \mathrm{OH} $$
(9)
$$ \mathrm{OH}+\mathrm{dye}\to \mathrm{dye}\ \mathrm{intermediate}\ \mathrm{products} $$
(10)
$$ \mathrm{OH}+\mathrm{dye}\ \mathrm{intermediate}\ \mathrm{products}\to \mathrm{harmless}\ \mathrm{products} $$
(11)
$$ {\mathrm{H}}_2{\mathrm{O}}_2+\cdotp \mathrm{OH}\to {\mathrm{H}}_2\mathrm{O}+\cdotp \mathrm{OOH} $$
(12)
$$ \mathrm{OOH}+\cdotp \mathrm{OH}\to {\mathrm{H}}_2\mathrm{O}+{\mathrm{O}}_2 $$
(13)

According to the curve of illumination for 4 h (unsealed), it could be seen that the degradation rate of MB solution kept increasing obviously with the increase of H2O2 dosage. This may be because the opening of the preservative film after 2 h of light would bring in air and promote the photocatalytic degradation.

Therefore, in the following work of this experiment, when studying the effect of UV-365 and H2O2 on other dyes, the optimal H2O2 dosage was 0.24 mol/L.

3.3.6 Degradation Effect of TiO2/CFs on Different Dyes

Figure 12 presents the effects of dye type on photocatalytic performance of TiO2/CFs under the conditions of UV-365 irradiation for 2 h and 4 h, TiO2/CFs dosage of 4.8 g/L, dye concentration of 10 mg/L, initial pH value of 7, and H2O2 dosage of 0.24 mol/L. The degradation rates of MB, MO, and RhB were 68.5%, 52.0%, and 23.5%, respectively, when the irradiation time was 2 h (sealed). The degradation rate increased by 20–30% to 90.4%, 81.4%, and 58.3% when the illumination time was further extended to 4 h (unsealed). Therefore, the prepared TiO2/CFs not only had excellent degradation effect on MB but also had certain degradation effect on MO and RhB.

Fig. 12
figure 12

The effects of dye type on degradation rate

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4 Conclusion

Anatase TiO2 was synthesized with uniform distribution and firm loaded in situ with CFs as carrier and TBT as raw material by the hydrothermal method and low-temperature calcination. Therefore, the excellent photocatalytic performance of nano-TiO2 permanently endowed CFs. The photocatalytic properties of the prepared TiO2/CFs were tested by visible spectrophotometry. The results showed that TiO2/CFs had good degradation effects on MB, MO, and RhB. The degradation rates of MB, MO, and RhB were 90.4%, 81.4%, and 58.3%, separately, at TiO2/CFs dosage of 4.8 g/L, initial dye concentration of 10 mg/L, pH of 7, and H2O2 dosage of 0.24 mol/L, after 4 h of UV-365 irradiation.