Introduction

Metal oxides have gained significant research attention owing to their chemical stability, catalytic, electrical and optical properties, and low cost [1, 2]. Because of their high density of surface sites and small size, metal oxide nanoparticles (NPs) have unique physical and chemical properties, compared to bulk materials [3]. Among the metal oxides NPs, cupric oxide (CuO), a p-type semiconductor material, exhibits remarkable electrical, optical, catalytic, adsorption, and biological properties [4, 5]. Since the morphologies of nanostructures have a strong influence on their efficiency, various CuO nanostructures have been synthesized using various methods namely, sol–gel, precipitation–pyrolysis, hydrothermal, solvothermal, microwave irradiation, and sonochemical methods [6, 7]. Another interesting metal oxide nanostructure is hematite (α-Fe2O3), which is an intriguing n-type semiconductor material, characterized by its chemical stability, natural abundance, nontoxicity, and low cost. Nanocomposites based on metal oxides have become increasingly popular in physics, chemistry, materials science, and engineering [8,9,10]. Various metal oxide nanocomposites such as CuO–ZnO, CeO2–MnOx, ZnO–MgO, ZnO–NiO, Co3O4–ZnO, and TiO2–WO3, have been fabricated and studied using various techniques. According to these studies, such nanocomposites have shown a higher photo-carrier separation performance than single oxides [8, 11]. They are mixtures of two or more different materials that combine the best qualities of the constituents to open new application opportunities, in which at least one material has nanoscale dimensions [10]. Numerous water treatment processes have used nanomaterials with special capabilities to remove diverse contaminants, including membrane separation, adsorption, coagulation, and others [12,13,14,15]. However, the primary objective of these techniques is to turn pollutants from an aqueous phase into a solid phase [16]. Advanced oxidation processes, or AOPs, are required for the complete degradation of these contaminants and to prevent secondary contamination. Due to its effectiveness in destroying organic molecules, heterogeneous photocatalysis has emerged as one of the most intriguing AOPs [17,18,19]. There have been a lot of laboratory-scale studies on photocatalysis, but before it can be used in an industrial setting, there are some issues that need to be solved, including the lack of effective and affordable catalysts that can be used with a wider range of solar spectra instead of UV lamps and the design of reactors that are suitable for potential applications on a large scale [20]. In addition to photocatalysis, microwave-assisted catalysis has also been used to reduce pollutants, but it still has several disadvantages, such as high costs, energy consumption, and difficulties with large-scale use. The difficulty in scaling up microwave-induced processes is primarily caused by greater heat loss, changes in absorption, and a shallower radiation penetration depth into the reaction media [21].

One of the emerging applications of metal oxide nanocomposites is the degradation of organic dyes, which are considered one of the most hazardous substances detected in water supplies and wastewater. As mixed metal oxides, iron and copper oxides have recently gained significant attention for a variety of applications [22,23,24]. The combination of Fe2O3 and CuO into an Fe2O3–CuO nanocomposite achieves superior physical properties compared to Fe2O3 and CuO separately. This study thus presents a simple method for degrading various organic dyes in the absence of light, addressing the issue of insufficient light, using an α-Fe2O3–CuO nanocomposite synthesized via the low-temperature decomposition of readily available copper hexacyanoferrate (Cu2[Fe(CN)6]; CHCF). In the presence of hydrogen peroxide (H2O2), α-Fe2O3–CuO is used as a catalyst to activate the thermally induced degradation of direct violet 4 (DV4), methylene blue (MB) and rhodamine b (RhB) dyes. In the context of AOP technologies, this technique saves more energy than photocatalysis. Various experimental parameters are studied to assess their influence on the degradation efficiency of DV4. Meanwhile, the reactive species and mechanism of H2O2 activation by α-Fe2O3–CuO are investigated considering the experimental results.

Experimental

Materials

Potassium ferrocyanide, copper chloride dihydrate, H2O2, sodium hydroxide (NaOH), hydrochloric acid (HCl), sodium nitrate (NaNO3), sodium chloride (NaCl) and sodium sulphate (Na2SO4) were obtained from Fluka. Humic acid was provided by Qualikems. Direct violet 4 (DV4), rhodamine B (RhB), and methylene blue (MB) were purchased from Acros and used without further purification (Scheme 1). Dilution and standard solutions were made with distilled water.

Scheme 1
scheme 1

Structure of organic dyes; methylene blue (MB), direct violet 4 (DV4), and rhodamine b (RhB)

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Synthesis of α-Fe2O3-CuO Nanocomposite

The α-Fe2O3–CuO nanocomposite was prepared via coprecipitation and low-temperature decomposition of the resulting complex (Scheme 2). Briefly, two aqueous solutions of potassium ferrocyanide (0.1 M) and copper chloride dihydrate (0.2 M) were prepared separately in 100 mL distilled water and then slowly mixed while stirring. Upon vigorous stirring for another 30 min, a dark brown precipitate was produced. The solid was collected, rinsed with distilled water several times, dried for 1 h at 90 °C, and further heated at 250 °C for 1 h to induce its decomposition. For comparative purposes, iron oxide was generated by combining the aqueous solutions of ferric chloride and potassium ferrocyanide, as previously reported [25]. The resulting complex (Prussian blue) was isolated, washed, and dehydrated at 90 °C for 1 h before heating at 250 °C for 1 h (Scheme 2). Meanwhile, cupric oxide (CuO) NPs are prepared through a precipitation process using NaOH as a precipitant. In a beaker of 50 mL distilled water, 0.2 M CuCl2·2H2O was dissolved. Next, 0.5 M NaOH was introduced gradually while being agitated constantly until pH 10 was achieved. The resulting precipitate was washed, dried, and heated for 1 h at 250 °C.

Scheme 2
scheme 2

Preparation of the α-Fe2O3-CuO nanocomposite

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Catalytic Activity of the α-Fe2O3-CuO Nanocomposite

In the presence of H2O2 (0.03 M), the α-Fe2O3–CuO nanocomposite (0.03 g) is introduced separately into 200 mL of (1.4–10) × 10−5 M solutions of DV4, MB, and RhB dyes to assess its catalytic activity. The flasks containing the solutions are put in a water shaker thermostat (50 °C) for 30 min in the dark to reach equilibrium adsorption. After the catalyst is added to the reaction mixture, the kinetics measurements are initiated immediately.

A decrease in the absorbance of unreacted dye is immediately recorded using an ultraviolet–visible (UV–vis) spectrophotometer, and the measurements are performed until no further decline in absorption was observed. The efficiency of catalytic degradation was estimated using Eq. (1) as follows:

$${\mathbf{Degradation}} \left( \% \right) = \frac{{{\mathbf{A}}_{{\mathbf{o}}} - {\mathbf{A}}_{{\mathbf{t}}} }}{{{\mathbf{A}}_{{\mathbf{o}}} }} \times 100$$
(1)

where A0 and At denote the absorbance of the dye before and after a specific period of the catalytic reaction, respectively.

Characterization

The XRD is used to characterize synthesized materials using a Rigaku MiniFlex 2 X-ray diffractometer. The FTIR spectra are obtained in a 4000–400 cm−1 range using a 670 (Nexus) Nicolet FTIR spectrophotometer in transmittance mode. A Cary Bio100 spectrophotometer is used to measure UV–vis spectra. SEM with EDX is used to analyze the nanocomposite’s surface properties and composition (SEM–EDX; JEOL, JSM-IT100LA). Using TEM (JEOL, JSM-6360A), the particle size is identified. XPS was performed on K-ALPHA (Themo Fisher Scientific, USA) employing monochromatic X-ray Al K-α radiation “with a spot size of 400 µm, at pressure 10–9 mbar with full spectrum pass energy 200 eV and at narrow spectrum 50 eV”.

The pHPZC, or point of zero charge, of the nanocomposite, is identified as described previously [30]. In brief, several Erlenmeyer flasks were filled with 0.01 M NaCl solutions (50 mL). By adding HCl or NaOH solutions (0.1 M), the original pH was changed to 2, 4, 6, 8, 10, and 12. After that, each flask received 0.1 g of the nanocomposite, which was agitated for 24 h at room temperature. The pHPZC is calculated by graphing the starting pH against the final pH of the solutions.

Results and Discussion

Characterization of the α-Fe2O3–CuO Nanocomposite

The surface morphology of the α-Fe2O3–CuO nanocomposite is investigated via SEM.

Figure 1a–c displays the corresponding micrographs, which clearly reveal the presence of two distinct structures (Fig. 1a). The presence of monoclinic particles indicates the formation of CuO NPs (Fig. 1b). Meanwhile, several nanorods corresponding to α-Fe2O3 can be also observed (Fig. 1c). The TEM image depicted in Fig. 1d confirms the nanosize of the synthesized α-Fe2O3–CuO nanocomposite, which exhibits a monoclinic crystalline structure of CuO NPs with highly dispersed α-Fe2O3 nanorods on the surface. Figure 2 illustrates the XRD patterns of the synthesized materials. As previously mentioned, the heating of Prussian blue at 250 °C resulted in an amorphous phase of Fe2O3 (Fig. 2 inset) [25]. The peaks at 2θ = 17.64°, 25.04°, 35.79°, 40.1°, 44.0°, 51.4°, 55.0°, and 57.8° in the complex pattern (Fig. 2a) are indicative of the cubic crystalline structure of the CHCF complex [26]. Figure 2b exhibits the distinctive peaks of the monoclinic phase of CuO NPs at 2θ values of 35.60°, 39.14°, 49.22°, and 58.35° which correspond well with the standard JCPDS no. 45–0937. The XRD pattern obtained for the α-Fe2O3–CuO composite (Fig. 2c) displays the characteristic peaks of a high density CuO phase, albeit slightly shifted, and small diffractions of α-Fe2O3 at 2θ values of 24.20°, 33.07°, 35.98°, 40.75°, 49.40°, 53.90°, and 57.40°, indicating the successful synthesis of the composite α-Fe2O3–CuO structure. According to the Debye–Scherrer formula [27], the average crystalline size of the CuO/α-Fe2O3 nanocomposite is determined to be 14 nm. The FTIR spectra of the CHCF complex, the α-Fe2O3–CuO nanocomposite, amorphous Fe2O3, and CuO NPs are shown in Fig. 3. The spectrum of the CHCF complex exhibits a broad band at 3434 cm−1 and another band at 1630 cm−1 due to the O–H groups stretching and bending vibrations indicating the presence of water. A high intensity band at 2107 cm−1 is associated with the stretching vibration of cyanide group of the FeII–CN–CuII bonds, confirming the presence of Fe(II) in the compound structure [26, 28]. Additional bands at 595 and 500 cm−1 might be attributable to the Fe–C and Cu–N stretching vibrations, respectively [28]. In the spectrum of Fe2O3, besides the characteristic bands of OH groups, two bands appear at around 575 and 631 cm−1 corresponding to Fe–O stretching. The spectrum of CuO NPs shows bands in the range of 498–580 cm−1 attributable to Cu–O vibrations, confirming the formation of CuO NPs [29]. Meanwhile, the spectrum of the α-Fe2O3–CuO nanocomposite features a band at 544 cm−1 corresponding to the overlap of Cu–O and Fe–O bands, which is in accord with the formation of a composite between α-Fe2O3 and CuO. EDX spectroscopy is used to identify the elemental composition and the percentage of each element within the substance. Figure 4 shows the EDX distribution of the synthesized nanocomposite, which reveals the presence of Cu, Fe, and O as elementary components in α-Fe2O3–CuO. The weight percentages of Fe, Cu, and O in the synthesized sample were 21%, 49%, and 30%, respectively, suggesting that the Fe to Cu ratio is 1:2. The magnetic property of the synthesized sample was examined using a vibrating sample magnetometer, VSM. The saturation magnetization (Ms) of α-Fe2O3–CuO is 4.90 emu g−1, as outlined in Fig. 5a. This reflects the nearly superparamagnetic nature of the nanocomposite. The α-Fe2O3–CuO was swiftly extracted from the reaction solution using an external magnet. The optical band gap (Eg) is related to the absorption coefficient (α) and photon energy (hυ) by the Tauc equation αhυ = A (hυ–Eg)n, where A is a constant, and n is an index that varies depending on the mechanism of interband transitions, with n = 2 or 1/2 referring to indirect or direct transitions, respectively [30]. The Eg value of α-Fe2O3-CuO is obtained from the intersection of the extrapolation of the linear part of the curve (Fig. 5b). The best linear match was obtained by plotting (αhυ)2 against hυ, and the direct Eg value of nanocomposite is determined to be 2.92 eV. This value is higher than that of bulk CuO; in addition, the band gap energy increases as the particle size decreases due to quantum confinement effects [31].

Fig. 1
figure 1

SEM (a–c) and TEM (d) images of the α-Fe2O3–CuO nanocomposite

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Fig. 2
figure 2

XRD patterns of the synthesized materials; CHCF complex (a), CuO NPs (b), and α-Fe2O3-CuO nanocomposite (c), and the inset refers to amorphous Fe2O3

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Fig. 3
figure 3

FTIR spectra of the synthesized materials; (a) CHCF complex, (b) amorphous Fe2O3, (c) CuO NPs, (d) and (e) α-Fe2O3–CuO nanocomposite before and after the catalytic reaction

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Fig. 4
figure 4

EDX of the α-Fe2O3–CuO nanocomposite

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Fig. 5
figure 5

Magnetization curve (a), and Tauc plot (b) of the α-Fe2O3–CuO nanocomposite

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The α-Fe2O3–CuO nanocomposite’s composition and oxidation state are determined using XPS analysis. The XPS survey of α-Fe2O3–CuO reflects the existence of O 1 s (57.97%), Cu 2p3 (25.72%), and Fe 2p (14.15%), as seen in Fig. 6a. The Fe 2p3/2 and Fe 2p1/2 of Fe3+ species are indicated by two peaks in Fig. 6b, which are situated at 710.5 and 723.9 eV, respectively [32, 33]. The satellite peak at 717.7 eV is typical of Fe3+ in α-Fe2O3 [34]. The high-resolution spectrum for Cu was depicted in Fig. 6c, and the peaks at 953.5 and 933.27 eV, is attributed to Cu 2p1/2 and Cu 2p3/2 of Cu2+, respectively [30, 35]. This confirms the presence of CuO in the nanocomposite. CuO is also supported by the existence of the shake-up peaks (satellites) at 961.7, 943.4, and 941.0 eV, which are situated at higher binding energies to the main peaks and suggest the existence of an unfilled shell (Cu 3d9) of Cu2+ [30]. The main peak for the O 1 s XPS spectrum in Fig. 6d denotes the lattice oxygen of Fe–O and Cu–O with a binding energy of 529.9 eV. The other peak is located around 531.5 eV and represents the surface OH species [34, 36].

Fig. 6
figure 6

XPS survey (a) and high-resolution spectra of the Fe 2p (b), Cu 2p (c), and O 1 s (d) of the nanocomposite

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Degradation Kinetic Study

Organic dyes are common compounds that are widely used in many fields. As a result, massive amounts of dyes are released into streaming waste. AOPs have gained increasing attention in textile and dye wastewater treatment. In this study, to assess the catalytic deterioration efficiency of α-Fe2O3–CuO, DV4, MB, and RhB are selected as model dye contaminants in aqueous media because of their toxicity and resistance to degradation. The degradation tests were performed with H2O2 as an ecofriendly oxidant. The activation of H2O2 to create an efficient oxidizing species still constitutes a challenge. When a catalyst is added, oxygen reactive species, primarily OH radicals, are generated as the primary oxidizing species. In the absence of a catalyst, the mixture (dye/H2O2) remained stable for several hours with no change in absorbance, indicating that no reaction occurred between the dyes and H2O2. The decrease in absorbance caused by the catalytic degradation of DV4, RhB and MB in the presence of H2O2 is depicted in Fig. 7. The catalytic reaction kinetics followed a pseudo-first-order model (Eq. 2), where the initial H2O2 concentration was at least 150 times higher than the dye concentration:

$${\mathbf{ln}} \, ({\mathbf{A}}_{{\mathbf{0}}} /{\mathbf{A}}_{{\mathbf{t}}} ) \, = \, {\mathbf{k}}_{{{\mathbf{obs}}}} {\mathbf{t}}$$
(2)

where Kobs (min−1) denotes the observed rate constant and t (min) denotes the dye removal time.

Fig. 7
figure 7

Absorption spectra of direct violet 4 (DV4; 7 × 10 − 5) (a), methylene blue (MB; 3 × 10 − 5) (b), and rhodamine b (RhB; 2 × 10 − 5) (c) recorded at different time intervals in the presence of 0.03 g of α-Fe2O3–CuO nanocomposite and H2O2 (0.03 M) at 50 °C

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Effect of Operating Factors on Azo Dye DV4 Degradation

Figure 8a shows a comparison between the nanocomposite’s catalytic activity and that of amorphous Fe2O3 and CuO NPs, as well as the mechanical mixing of CuO and Fe2O3 with a 2 to 1 ratio, as in the nanocomposite. The synthesized α-Fe2O3–CuO nanocomposite displayed the highest DV4 degradation efficiency. This result reveals that the synergistic interaction of Fe2O3 and CuO significantly enriched the catalytic active sites and improved the DV4 degradation efficiency. Numerous variables, including catalyst dose, dye concentration, temperature, contact time, and reaction pH may have an impact on the catalytic activity of α-Fe2O3–CuO. To optimize the catalytic process, the variables having the highest impact must be identified.

Fig. 8
figure 8

Degradation percentage of DV4 over time using 0.03 g of different catalysts and 0.03 M H2O2 (a), degradation percentage of DV4 over time (b), and the first-order plots (c) using 0.03 g of α-Fe2O3–CuO in the presence of diverse concentrations of H2O2 at 50 °C

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Effect of the H2O2 Concentration

A series of blank experiments revealed that H2O2 has no effect on the DV4 degradation in the absence of the catalyst and about 20% of the dye is degraded within 1 h in the presence of α-Fe2O3–CuO alone. When H2O2 was added to the reaction media, the percentage of dye degradation notably increased. The effect of the H2O2 concentration (0.01–0.07 M) on the degradation efficiency of DV4 (4 × 10−5 M) is investigated using a constant dose of catalyst (0.03 g) at 50 °C and pH 6. Figure 8b, c shows the degradation efficiency over time and first-order graphs for the oxidative reaction of DV4 at varied H2O2 concentrations. The percentage and rate constant values of the dye degradation (Table 1) increased with the H2O2 concentration. This could be due to the production of hydroxyl radicals (HO·), which attack DV4 molecules. The generation of highly active HO· has been previously connected with the activation of H2O2 by heterogeneous catalysts [37].

Table 1 Catalytic degradation of direct violet 4 (DV4) by nanocatalyst (α-Fe2O3-CuO) under different conditions after 1 h
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Effect of the Catalyst Dose

The effect of catalyst dose on the DV4 (4 × 10−5 M) degradation efficiency is investigated by changing the amount of α-Fe2O3–CuO from 0.01 to 0.06 g at 50 °C, while maintaining the H2O2 concentration at 0.03 M. The obtained results are shown in Fig. 9a and Table 1. The degradation efficiency and rate constant increased as the catalyst dose is increased to 0.04 g. This tendency could be explained by the existence of more active sites, which leads to the production of more reactive radicals (HO·). Upon further increasing the catalyst dose beyond the optimum level, the deterioration rate decreased, which could be due to the self-quenching of a large number of radicals produced by H2O2 instead of the reaction with dye molecules. The formation of aggregates between catalyst particles, which would decrease the number of available sites for the activation of H2O2, might be another reason behind the decrease in the degradation rate at a high catalyst dose.

Fig. 9
figure 9

The rate constant as a function of catalyst dose (a), degradation percentage of DV4 over time in the presence of 0.03 g of α-Fe2O3-CuO nanocomposite at various DV4 concentrations (b) and at various reaction temperatures (c)

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Effect of the DV4 Concentration

The effect of varying the initial DV4 concentration from 1.4 × 10−5 to 10 × 10−5 M was studied, while the H2O2 concentration and the catalyst dose were kept constant at 0.03 M and 0.03 g, respectively (Fig. 9b). At a low DV4 concentration (1.4 × 10−5 M), the degradation percentage reached almost 100% within 15 min; however, it declined upon further increasing the DV4 concentration. As can be seen in Table 1, raising the DV4 concentration resulted in a drop in the rate constant. This drop is attributable to excessive DV4 molecules covering the catalyst's active sites.

Effect of Temperature

In most oxidation processes, temperature is a critical parameter, especially in industrial and environmental applications. The temperature effect on the reaction rate is analyzed at 25, 37, 50, and 70 °C, while H2O2 (0.03 M), the concentration of DV4 (4 × 10−5 M), and the catalyst dose (0.03 g) were kept constant. The degradation efficiency increased dramatically upon increasing the reaction temperature. Thus, increasing the temperature from 25 to 70 °C resulted in a degradation efficiency of 87% within 2 min (Fig. 9c) and an increase in the reaction rate from 0.057 to 0.894 min−1 (Table 1). The considerable increase in the degradation efficiency at high temperature is attributable to the production of additional radicals, which contribute to the oxidation of DV4 molecules into reaction products. According to prior studies, heat speeds up the reaction and produces more active radicals [38, 39]. The activation energy (Ea) is calculated from the Arrhenius plot (not shown). Furthermore, thermodynamic parameters ΔG#, ΔS#, and ΔH# are determined from the Eyring and Gibbs equations and are listed in Table 2. The catalytic process was endothermic, as indicated by the positive enthalpy value. The Ea value was in the common range of chemical reactions.

Table 2 The activation parameters and observed rate constant values of the catalytic degradation of direct violet 4 (DV4) using nanocatalyst (0.03 g), and H2O2 (0.03 M), at different temperatures
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Effect of pH

The solution pH is a significant parameter that affects the oxidative degradation of organic pollutants. The efficiency of the catalyst in the degradation of DV4, RhB and MB was investigated at the pH levels of 4, 6, 8, and 10. As depicted in Fig. 10a, the highest degradation percentage (100%) for DV4 is obtained at a pH of 6. However, the degradation efficiency decreased upon further increasing the pH. This is because the surface properties of the catalyst are influenced by pH changes. The pHPZC of α-Fe2O3–CuO is estimated to be 7.8 following the reported method [10]. At pH < pHPZC (7.8), the catalyst’s surface becomes positively charged, whereas at pH > pHPZC, it becomes negatively charged. DV4 is an anionic dye, and at pH 6, the catalyst surface owns a positive charge that attracts the DV4 dye molecules more strongly, increasing the degradation rate. For RhB, the maximum degradation (95.3%) is achieved at pH 4. At this pH, the positively charged catalyst surface attracts RhB molecules, which is a zwitterionic structure in a polar solvent [40]. The formation of the zwitterion is favored as the pH increases, which contributes to the RhB aggregation and formation of dimers [40]. Therefore, the degradation efficiency of the catalyst diminishes at higher pH values. In contrast, the MB degradation increased efficiency upon increasing the pH value from 4 to 10. MB is a cationic dye, and at pH > pHPZC the catalyst surface becomes negatively charged which attracts the MB molecules more strongly, resulting in an enhancement of the degradation rate.

Fig. 10
figure 10

Degradation percentage of DV4, MB, and RhB at various pH values (a), and the effect of ultraviolet irradiation (b) on the degradation of DV4 (4 × 10 − 5) using 0.03 g of α-Fe2O3-CuO nanocomposite and H2O2 (0.03 M) at pH 6 and 50 °C

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Effect of Ultraviolet Irradiation

Generally, the combination of ultraviolet (UV) light irradiation and H2O2 results in a complete degradation of organic dyes. During the oxidation process, highly oxidative species such as HO· are formed. In the present study, an aqueous solution of DV4 (4 × 10−5 M) remained relatively stable upon UV irradiation in the absence of H2O2 or the catalyst (not shown). In contrast, when the reaction mixture (DV4, catalyst, and H2O2) is exposed to UV light (254 nm), the degradation rate constant increased dramatically from 0.163 to 0.54 min−1, as seen in Fig. 10b. This enhancement reveals that UV irradiation produces high concentrations of free radical species.

Effect of Inorganic Ions

In natural and wastewater, various inorganic anions including chloride (Cl−), nitrate (NO3−), and sulfate (SO42−) are frequently found [41]. So, it is essential to investigate how they affect the degradation rate of DV4 dye. The effects of added SO42−, Cl−, and NO3− ions on the degradation rate are investigated separately at constant doses of the α-Fe2O3–CuO catalyst and DV4 with various quantities of Na2SO4, NaCl, and NaNO3. The degradation rate of DV4 is decreased by the SO42− ion concentration, as seen in Fig. 11a. The thermally induced catalytic activity in the presence of SO42− ions appears to produce SO4·− radicals, according to eqs. 3 and 4.[42] The ·OH radicals or holes are captured by the SO42− anions, which prevents deterioration. The excess SO42− decelerated the degradation rate of DV4 because the generated SO4·− is less reactive than ·OH and h+ [42].

$${\mathbf{SO}}_{{\mathbf{4}}}^{{{\mathbf{2}} - }} + \, {\mathbf{h}}^{ + } \to \, {\mathbf{SO}}_{{\mathbf{4}}}^{ \bullet - }$$
(3)
$${\mathbf{SO}}_{{\mathbf{4}}}^{{{\mathbf{2}} - }} + \,^{ \bullet }{\mathbf{OH}} \, \to \, {\mathbf{SO}}_{{\mathbf{4}}}^{ \bullet - } + \, {\mathbf{OH}}^{ - }$$
(4)
Fig. 11
figure 11

Degradation rate constants of DV4 (4 × 10 − 5 M) using 0.03 g of α-Fe2O3-CuO and 0.03 M H2O2 in the presence of various concentrations of a Na2SO4, b NaCl, and c NaNO3 at 50 °C

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By contrast, the rate of DV4 degradation rises with increasing NaCl concentration (Fig. 11b). This is primarily caused by an increase in reactive chlorine species (RCS: Cl·, Cl2·−, ClOH·−, etc.), which hastened the degradation rate of DV4. Earlier studies reported a similar outcome [41, 43]. Like Cl−, rising NO3− ion concentrations accelerate the rate of deterioration (Fig. 11c). This considerable boost in performance is attributed to the production of more reactive species (·OH and NO2·) because of NO3 photolysis [41, 44].

Effect of Humic Acid

Natural organic matters, or NOMs, are a group of carbon-based substances that are found in a variety of groundwater resources and surface waters. Humic substances make up a sizable component of NOMs [45]. So, the impact of various humic acid, HA, concentrations on the rate of DV4 degradation is assessed (Fig. 12). At a low dose of HA (1–5 mg L−1), the DV4 degradation was ≥ 98%, as depicted in Fig. 12a. At 1 and 5 mg L−1 HA, the rate constants reduced marginally from 0.163 min−1 to 0.162 and 0.158 min-1, respectively (Fig. 12b). With increasing HA dose to 10, 20, and 35 mg L−1, the degradation rates decreased to 0.136, 0.115, and 0.088 min−1, respectively. This can be attributed to HA's capacity to operate as a radical scavenger of ·OH as well as the radical competition between HA and DV4 molecules [33, 46].

Fig. 12
figure 12

Degradation percentage (a) and rate constants (b) of DV4 (4 × 10−5 M) using 0.03 g of α-Fe2O3-CuO and 0.03 M H2O2 in the presence of diverse doses of humic acid at 50 °C

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Dyes Mineralization

The purpose of the catalytic degradation process is not only the decolorization of the dye but also its mineralization, that is, its decomposition into CO2 and H2O. The dye decolorization was measured using a UV–vis spectrophotometer. The dye absorption bands for DV4, MB, and RhB disappeared entirely, indicating that the dyes were rapidly degraded, and the aromatic system was fully destroyed. The total organic carbon (TOC) analysis technique was used to measure the total quantity of carbon in organic compounds converted to CO2 during an oxidation reaction. The TOC removal percentage is about 91% following 2 h of catalytic degradation of DV4 (4 × 10−5 M) in the presence of H2O2 at pH 6 and 50 °C. The reduction in TOC is due to ring-opening mechanisms, which convert the aromatic molecules to aliphatic ones. For MB, the catalyst reached a 90% TOC removal, whereas a TOC removal of 81% is obtained in the case of RhB. These experiments are performed at the optimum pH values for the degradation of MB and RhB. Moreover, the chemical oxygen demand (COD) is measured by using a powerful oxidizing agent to completely oxidize any organics into CO2 under acidic conditions. After the catalytic degradation process, the COD removal percentage for DV4, MB, and RhB were 95%, 93%, and 85%, respectively. The observed reduction of COD demonstrates that the starting organic dyes were oxidized to mineral ions.

Degradation Mechanism

To investigate the mechanism of the α-Fe2O3–CuO nanocatalyst in the oxidative degradation of DV4 dye, studies were conducted to trap reactive species during the thermo-induced catalytic reaction. The selected scavengers are tert-butyl alcohol (TBA, ·OH radical scavenger), benzoquinone (BQ, O2· scavenger), and disodium ethylenediaminetetraacetic acid (EDTA, h+ scavenger). They were added individually to a 100 mL reaction solution at a concentration of 0.4 M. In the absence of a scavenger, the degradation rate constant of DV4 was 0.163 min−1. As shown in Fig. 13a, adding BQ to the DV4 solution decreased the rate of degradation slightly (0.147 min−1), whereas the addition of EDTA resulted in a significant reduction in the degradation rate (0.003 min−1). When TBA was used, the degradation rate of DV4 dropped from 0.163 to 0.070 min−1. These findings suggest that ·OH and h+ are responsible for DV4 dye's degradation.

Fig. 13
figure 13

The effect of different scavengers (a) on the degradation of DV4 (4 × 10−5) using 0.03 g of α-Fe2O3-CuO nanocomposite and H2O2 (0.03 M) at pH 6 and 50 °C, proposed illustration of the thermally induced catalytic degradation of DV4 dye by a p–n heterojunction (b) and the effect of nanocatalyst recycling on the DV4 degradation percentage (c)

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One of the most effective strategies that have been considered to improve photocatalytic performance is the formation of a heterojunction between numerous semiconductors [47,48,49,50]. In the current study, a possible degradation mechanism is suggested in the schematic design of Fig. 13b based on the active species trapping results. A p-n heterojunction will form at the interface of the synthesized α-Fe2O3-CuO nanocomposite because α-Fe2O3 (n-type) and CuO (p-type) have differing band gaps and electronegativity [50,51,52]. As a result, an internal electric field is formed at the interface [51, 52]. Under the circumstances of thermal activation, the excited electrons from the valence bands (VB) of α-Fe2O3 and CuO migrate to the corresponding conduction bands (CB) [38]. CuO’s high CB position and the interface’s internal electric field cause electrons to move from CuO to α-Fe2O3 and holes to move the other way [50].

Thus, a successful partition of electrons and holes happens at the interface. Furthermore, according to the double charge transfer mechanism the electrons will be accumulated in the conduction band of α-Fe2O3 and the holes will be accumulated in the valence band of CuO [50, 53]. When α-Fe2O3–CuO is present alone, the catalytic degradation efficacy is limited because of electron and hole recombination. With the addition of H2O2, the results demonstrate a rapid rate of degradation (Fig. 8b, c). Where H2O2 inhibits the recombination of free carriers by trapping electrons. H2O2 molecules are then reduced and converted to OH− and ·OH. Furthermore, ·OH can be produced by trapping holes in the valence band with surface-bound H2O or OH– (adsorbed and free) [54].

Therefore, the thermally generated positive hole (h+) in (VB) and an electron (e–) in (CB) are considered powerful oxidizing and reductive agents respectively and the oxidation/reduction reactions can be represented by the following Eqs. (5–9):

$${{\varvec{\upalpha}}} - {\mathbf{Fe}}_{{\mathbf{2}}} {\mathbf{O}}_{{\mathbf{3}}} - {\mathbf{CuO}} \, + \, {\mathbf{thermal}} \, {\mathbf{excitation}} \, \to \, {\mathbf{e}}^{{{-} }} \left( {{\mathbf{CB}}} \right) \, + \, {\mathbf{h}}^{ + } \left( {{\mathbf{VB}}} \right)$$
(5)
$${\mathbf{H}}_{{\mathbf{2}}} {\mathbf{O}}_{{\mathbf{2}}} + \, {\mathbf{e}}^{-} \to \,^{\bullet }{\mathbf{OH}} \, + \, {\mathbf{OH}}^{ - }$$
(6)
$${\mathbf{H}}_{{\mathbf{2}}} {\mathbf{O}} \, + \, {\mathbf{h}}^{ + } \to \,^{ \bullet }{\mathbf{OH}} + {\mathbf{H}}^{ + }$$
(7)
$${\mathbf{h}}^{ + } + \, {\mathbf{OH}}^{ - } \to \,^{ \bullet }{\mathbf{OH}}$$
(8)
$${\mathbf{Dye}} \, + \,^{ \bullet }{\mathbf{OH}} \, \to \, {\mathbf{CO}}_{{\mathbf{2}}} + \, {\mathbf{H}}_{{\mathbf{2}}} {\mathbf{O}}$$
(9)

Catalyst Reuse

The catalyst’s reusability is a crucial factor in determining how cost-effective the process is. Therefore, the recovery and recyclability of the α-Fe2O3–CuO catalyst were examined in the oxidation of DV4 with H2O2. The catalyst can be regenerated in a straightforward manner. Using an external magnet, the catalyst was magnetically recovered from the media after two hours of disintegration. It was thoroughly washed with distilled water, dried (at 60 °C for 12 h), and then used once more in the reaction’s subsequent cycle. Figure 13c depicts the degradation efficiency of DV4 as a function of cycle number. The results reveal that the DV4 degradation percentage remained constant at approximately 100% throughout five cycles. The nanocomposite’s FTIR spectrum is assessed following three regeneration cycles to investigate its stability. The spectrum revealed the distinctive bands of the α-Fe2O3–CuO nanocomposite, as illustrated in Fig. 3. Moreover, ICP/OES (inductively coupled plasma-optical emission spectrometry) analysis showed that almost no Cu or Fe leached from the catalyst surface into the solution proving the stability of the catalyst. The nanocomposite (α-Fe2O3–CuO) is therefore an appropriate option for the remediation of dye-contaminated water.

Conclusion

This study provides an ecofriendly and cost-effective technique for the synthesis of an α-Fe2O3–CuO nanocomposite. The α-Fe2O3–CuO nanocomposite is successfully prepared via oxidative conversion of CHCF in air at a low decomposition temperature. Techniques like XRD, SEM, EDX, FTIR, XPS, VSM, and TEM described the α-Fe2O3–CuO nanocomposite. It showed excellent activity for the degradation of DV4, MB, and RhB in the presence of H2O2 with the aid of temperature (≥ 50 °C). During thermal activation, the electrons on the VB are transferred to the CB because a higher temperature can promote the charge transfer of electrons and holes at semiconductors. Meanwhile, the rest of the electrons on the CB are trapped and oxidized by H2O2 to produce ·OH. The released radicals and the existing free holes oxidize the organic dyes. The degradation rate of the dye is influenced by the H2O2 concentration as well as the catalyst dose. Inorganic anions (Cl−, NO3−, and SO42−) and humic acid, especially at high doses, have a notable impact on the degradation as well. According to pH impact, RhB and DV4 are more favored in acid media, whereas MB degradation proceeded better in alkaline media. TOC and COD measurements revealed that a significant portion of DV4, MB, and RhB dyes are degraded to CO2, H2O, and some inorganic acids. Moreover, the catalyst showed excellent recyclability after five runs. The α-Fe2O3–CuO nanocomposite’s capacity to degrade the three dyes and its reusability make it an effective catalyst for treating wastewater.