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BY 4.0 license Open Access Published by De Gruyter Open Access October 31, 2022

Ferrate synthesis using NaOCl and its application for dye removal

  • Gunawan Gunawan EMAIL logo , Nor Basid Adiwibawa Prasetya , Abdul Haris and Eka Pratista
From the journal Open Chemistry


Ferrate salt is a powerful oxidant for dye degradation. This work demonstrates a new method for degrading dyes containing Fe(vi) by synthesizing NaOCl from the electrolysis of table salt. NaOCl is then reacted with Fe(OH)3 in an alkaline condition to form ferrate. Electrolysis of table salt was successfully carried out using platinum as an anode and zinc as a cathode. The obtained ferrate was characterized by using Fourier transform infrared, UV-Vis, and X-ray diffraction spectroscopy. The ferrate solution has a maximum wavelength of 505 nm with a characteristic purple color. Furthermore, the ferrate produced was utilized to remove methylene blue (MB), remazol black blue (RBB), and methyl orange (MO) dyes with varying contact times. The degraded dyes were then analyzed using LC/MS. The results showed that ferrate was effective to remove dyes with an optimum contact time of 60 min that follows an order one reaction. In this study, MB showed a percent degradation close to 100% with the fastest decolorization rate compared with MO and RBB. This research provides new insights into the benefits of table salt as a base material for NaOCl through electrolysis for synthesizing ferrate, used in dye removal applications.

1 Introduction

The availability of clean water is a crucial need at this time. This is due to reduced water sources because many lands are used for settlement. In addition, the increasing population makes clean water supply techniques an unavoidable need. One strategy to get clean water is to treat polluted water through physical, chemical, and microbiological methods. Some studies also use membranes to purify water [1,2,3,4]. Oxidizing agents are one of the most often used chemical treatments for water purification.

Ferrate ( FeO 4 2 ) has been known as an oxidizing agent of organic pollutants with high efficiency [5], low operating costs, and being environmentally friendly [6]. Ferrate has a wide range of applications, including oxidation, deodorization, coagulation, disinfection, adsorption, and sterilization [7,8,9,10]. As a coagulant, ferrate is particularly successful in dye removal [5,11,12] and gray water [13]. The main type of iron that acts as a coagulant is nanocrystalline Fe(iii) oxide/hydroxide resulting from the Fe(vi) reduction. Karelius and Nopriawan [11] used ferrate to remove methylene blue (MB) dye with a degradation percentage around 90%. Dwiasi and Kurniasih [14] also investigated the removal of methyl orange (MO) dye with ferrate.

Under highly alkaline circumstances, iron(iii) oxide (such as Fe2O3) or its salts (Fe(NO3)3, FeCl3) are reacted with sodium hypochlorite to produce wet ferrate [11,15]. People often utilize commercial sodium hypochlorite. However, in this study, NaOCl was obtained from electrolysis of the table salt solution, and then it was reacted with Fe(OH)3 in the alkaline condition to form ferrate solution. Then, liquid ferrate was used to remove dyes, which are a substantial source of nonaesthetic pollution. Remazol black blue (RBB), MB, and MO dye wastes are often used as the selected dyes and also used as model compounds in this study. Subsequently, the ability of potassium ferrate to remove dye was investigated. Through this research, researchers will be more aware of promoting the use of more environmentally friendly materials and green chemistry to support efforts to prevent sustainable environmental problems. In addition, the analysis presented in this study provides information that can be used in future research to explore various materials for dye removal.

2 Materials and methods

2.1 Materials

HCl (37%, Merck), NaOH (≥99%, Merck), KOH (≥85%, Merck), H2SO4 (95–97%, Merck), FeCl3 (99.99%), Remazol Black B (Sigma Aldrich), MB (Merck), MO (Merck), Na2S2O3 (≥97%, Merck), Starch (Merck), KI (≥99.5%, Merck), and table salt was obtained from the market.

2.2 Instrumentations

An electrochemical cell for ferrate electrolysis was prepared with platinum wire (≥99%, Osilla) as anode and zinc plate with an area of 3.4 cm × 3 cm (≥99%) as cathode. The electrodes were then attached to a DC power (Aditeg APS 3005) to carry out the electrolysis process. The instruments used were magnetic stirrer (Thermo Scientific), Ohaus analytical balance (Pioneer), thermometer, and AVOmeter. UV-Vis Spectrophotometer (Genesys 10 S UV-Vis) was used for dye analysis before and after the treatment, while solid Fe(OH)3 and ferrate were examined using instruments of Fourier transform infrared (FTIR) (Shimadzu IRPRESTIGE 21), SEM-EDX (JEOL JSM-6510LA), XRF (PANalytical XRF), and X-ray diffraction (XRD) (Shimadzu XRD-7000).

2.3 Synthesis of NaOCl by electrolysis of NaCl solution

Before use, platinum and zinc plates were washed with soap, aquadest, and alcohol to eliminate any contaminants adhered to the plate’s surface. Then the electrodes were attached to a DC power with platinum and zinc at anode and cathode, respectively, and they were immersed in 250 mL NaCl (table salt content 97% NaCl) solution with concentrations ranging from 0.25 to 5 M and electrolyzed at 3.5 V for 60 min. Previously, variations in the voltage and time of electrolysis were also carried out to see the ratio of the resulting sodium hypochlorite.

2.4 NaOCl determination using iodometric titration

The hypochlorite solution obtained through electrolysis was determined using the iodometric method, which included 25 mL of hypochlorite sample solution added with 15 mL distilled water, 20 mL KI 10%, 20 mL HCl 2 M, and 2 mL of starch solution as indicators and then titrated with 0.26 M Na2S2O3 until the solution was clear. The changing of the starch indicator from blue to clear (colorless) indicates the endpoint of the titration (colorless).

2.5 Synthesis of iron hydroxide (Fe(OH)3)

Sodium hydroxide 12 g and FeCl3 16 g were dissolved with 25 mL distilled water in separate two beakers. The NaOH solution was then added to the FeCl3 solution and reddish-brown precipitate was obtained. After 10 min the precipitate was filtered using filter paper, and dried in oven at 108oC and stored in desiccator. Part of the precipitate was characterized using FTIR and XRD and another part was used for ferrate synthesis.

2.6 Alkaline ferrate synthesis using Fe(iii) and NaOCl

A wet oxidation process was used to synthesize ferrate by reacting NaOH 28 g with 120 mL NaOCl solution (as a result of the electrolysis of NaCl 3 M for 6 h). Then the liquid was stirred until the NaOH was completely dissolved. After addition of Fe(OH)3 2 g to the solution, the solution was stirred for 30 min until the solution became dark purple. The solution was sealed and left for one day, then filtered through glass wool. The maximum wavelength of ferrate was measured at wavelength from 400 to 800 nm using a UV-Vis spectrophotometer.

2.7 Characterization

The presence of functional groups and bonds in the samples was determined using FTIR characterization. The crystalline phase of the sample was determined by XRD using Cu Kα radiation at an angle of 10–90°. SEM-EDX analysis and XRF were used to identify the elemental chemical composition. The absorption spectrum of the electrolysis solution was studied with UV-Vis to determine the ferrate produced and dyes. Absorbance measurements were done between 200 and 800 nm. After reaction with ferrate the concentration of dyes were measured with a UV-Vis spectrophotometer.

2.8 Ferrate application for dye removal

The dyes used in this study were MB, MO, and RBB. Each dye was prepared in a volume of 10 mL with concentration of 10 mg/L, then it was added with 3 drops of 170 mg/L ferrate solution (equal to 170 mg/L × 0.15 mL = 2.55 μg). The solution was stirred for varied times from 5 to 60 min. UV-Vis Spectrophotometer was used to determine the concentration of dye (before and after treated with ferrate).

3 Results and discussion

3.1 Synthesis of NaOCl through electrolysis of NaCl solution

Sodium hypochlorite was synthesized by electrolyzing 0.5, 1, 2, 3 and 5 M NaCl (from table salt) solutions using a platinum electrode as an anode and a zinc electrode as a cathode. In the electrolytic process, the voltage was scanned from 0 to 8 V. Currents were then recorded and displayed as in Figure 1(a).

Figure 1 
                  The current–voltage curve for electrolysis of (a) 0.5; 1; 2, 3 and 5 M Nacl and (b) 0.25; 0.50; 0.75 and 1.0 M NaCl concentration used for the NaOCl synthesis.
Figure 1

The current–voltage curve for electrolysis of (a) 0.5; 1; 2, 3 and 5 M Nacl and (b) 0.25; 0.50; 0.75 and 1.0 M NaCl concentration used for the NaOCl synthesis.

Figure 1(a) shows that the greater the concentration of NaCl, the lower the potential for the generation of a positive current (E d). As the NaCl concentration decreases, the E d value will be even more significant, allowing a competitive reaction to form chlorine gas with oxygen gas. To avoid the formation of oxygen gas, the potential used is not too high, which is around 3–4 V. Figure 1(b) shows the electrolysis of NaOCl synthesis using a smaller NaCl concentration of 0.25, 0.5, 0.75, and 1 M. From the figure, it can be seen that the smaller the concentration of NaCl, the greater the potential for decomposition that occurs. This allows the formation of oxygen gas more significant in addition to the primary purpose of producing NaOCl.

Electrolysis during the NaOCl synthesis process forms chlorine at the anode and NaOH at the cathode. NaOCl formation due to chemical interaction between chlorine gas and NaOH solution [16].

At cathode:

H 2 O H + + OH ;

(1) 2H + + 2 e H 2 ; Na + + OH NaOH .

At anode:

(2) NaCl Na + + Cl ; 2 Cl 2 e Cl 2 .

In the beaker glass:

(3) NaOH + Cl 2 NaOCl + NaCl + H 2 O .

3.2 NaOCl measurement study using iodometric titration

The iodometric titration method was used, namely the titration of NaOCl with sodium thiosulfate after addition of KI solution. The changing of the starch indicator from blue to colorless indicates the endpoint of the titration (colorless) [17]. Figure 2 depicts the concentration of NaOCl produced by electrolysis using NaCl with varied times measured using iodometric method.

Figure 2 
                  Concentration of NaOCl electrolyzed using 3 and 5 M NaCl at varied times.
Figure 2

Concentration of NaOCl electrolyzed using 3 and 5 M NaCl at varied times.

Figure 2 shows that the greater the concentration of NaCl used, the greater the NaOCl produced. It can be seen that NaCl 5 M gives a NaOCl yield of about 130 mM by electrolysis for 240 min. Meanwhile, for NaCl 3 M with the same electrolysis time, the result is about 65 mM NaOCl. For electrolysis with 5 M NaCl, the maximum yield of NaOCl is about 130 mM, and then NaOCl begins to decrease. Meanwhile, for electrolysis with 3 M NaCl, the NaOCl concentration continued to increase until 300 min.

Figure 3 shows that at the same electrolysis time, the higher the NaCl concentration, the higher the NaOCl produced [18]. At the electrolysis time of 60 min with a NaCl concentration of 5 M was able to produce NaOCl close to 90 mM. But the result is still small compared to the pro analysis and technical NaOCl concentrations, namely 738.3, and 280.8 mM, respectively. However, the NaOCl obtained can be used for ferrate synthesis. For the synthesis of ferrate, NaOCl was used as a result of electrolysis with 3 M NaCl solution for 300 min. The ferrate is produced from the reaction of iron hydroxide precipitate with NaOCl under alkaline conditions.

Figure 3 
                  The influence of NaCl concentration on the resulting NaOCl concentration. Electrolysis time is 60 min.
Figure 3

The influence of NaCl concentration on the resulting NaOCl concentration. Electrolysis time is 60 min.

3.3 Synthesis of iron hydroxide and Characterization

After mixing the ferric chloride solution with the NaOH solution, Fe(OH)3 was formed. When OH of NaOH are substituted with chloride ions (Cl) in the solution, the following reaction occurs, resulting in the creation of iron(iii) hydroxide (Fe(OH)3) with no additional secondary phase.

(4) FeCl 3 ( aq ) + 3 NaOH ( aq ) Fe ( OH ) 3 ( s ) + 3 NaCl ( aq ) .

This iron hydroxide precipitate is gelatin which is very insoluble in water. By adding excess NaOH, there will be a common ion effect, and the chemical equilibrium shifts towards the formation of more perfect and insoluble precipitates. The yield of Fe(OH)3 obtained was 86.14%. The resulting iron hydroxide precipitate is brownish.

Characterization was carried out for Fe(OH)3 and compared with FeCl3 as the starting material for preparing Fe(OH)3. Figure 4 depicts the pattern of the FTIR and XRD spectra.

Figure 4 
                  XRD (a) and FTIR spectra (b) of FeCl3 and Fe(OH)3.
Figure 4

XRD (a) and FTIR spectra (b) of FeCl3 and Fe(OH)3.

The Fe–O bond was connected with the absorption at 578 cm−1, confirming the production of iron hydroxide. Absorption peak of –OH appeared at 3,400 cm−1 [19,20]. In sample FeCl3, there was a peak observed at 586 cm−1, which proved the existence of iron. The graphs show the diffraction peaks at 2θ = 15.47, 20.55, and 50.82°, which is in agreement with the standard FeCl3 diffraction peaks [CAS No. 10025-77-1] [21]. The peaks at 2θ = 35.00 matches well with standard of Fe(OH)3 (JCPDS number 22-0346) [19]. The resulted product, according to XRD and FTIR measurements, is iron hydroxide Fe(OH)3.

Table 1 shows the EDX of Fe(OH)3. The presence of FeO components in EDX indicates the presence of the target substance, Fe(OH)3. Carbon peaks in the EDX spectrum are caused by the carbon bands employed during the EDX testing.

Table 1

EDX of Fe(OH)3

No Sample name Component Unit Analysis result value Method
1. Fe(OH)3 FeO % Weight 86.01 EDX
C 10.16
ZrO2 1.93
Na2O 1.49
Al2O3 0.41

3.4 Alkaline ferrate synthesis using Fe(iii) and NaOCl

Using wet chemical techniques, a purple ferrate solution was obtained by mixing NaOCl as an oxidizing agent with Fe(OH)3 as a source of Fe(iii) in a strongly alkaline state (NaOH) [22,23,24]. The reaction for ferrate production is shown below:

(5) Oxidation: Fe ( OH ) 3 + 5 OH FeO 4 2 + 4 H 2 O + 3 e ,

(6) Reduction: ClO + H 2 O + 2 e Cl + 2 OH ,

(7) Result: 2 Fe ( OH ) 3 ( brown ) + 4 OH + 3 ClO 2 FeO 4 2 ( purple ) + 3 Cl + 5 H 2 O .

The wavelength and concentration of the resulting ferrate solution (sodium ferrate) were measured with UV-Vis. The maximum detection wavelength was 505 nm with an absorbance of 0.300 (with a dilution factor of 4×). Other investigations [7,15,25,26,27] have validated the absorption spectra of ferrate at 505 nm. By using the formula c = 166000A/1170 and by entering the absorbance at 505 nm and multiplying by the dilution factor, the ferrate concentration obtained was 170 mg/L (Figure 5).

Figure 5 
                  Visible spectrum of ferrate solution.
Figure 5

Visible spectrum of ferrate solution.

3.5 Ferrate application for dye removal

Sodium ferrate is the ferrate formed (Na2FeO4). Fe(iii) is an effective coagulant capable of removing suspended particles from wastewater [28]. Sodium ferrate is gradually converted into oxygen and (Fe(OH)3) according to equation (8). In wastewater treatment, oxygen is released, causing the concentration of dissolved oxygen (DO) in the wastewater to rise. A higher DO concentration in the wastewater can avoid anaerobic conditions and prevent odor formation in other parts of the wastewater treatment plant [29].

(8) 4 Na 2 FeO 4 + 10 H 2 O 8 NaH + 4 Fe ( OH ) 3 + 3 O 2 .

3.5.1 The influence of contact time on dye degradation by ferrate

Contact time is a crucial element influencing material adsorption capability. Figure 6 depicts the influence of contact time on dye removal utilizing the ferrate oxidation procedure. Contact time is an essential factor affecting the adsorption capacity of materials. Figure 6 depicts the effect of contact time on removal dyes using the ferrate oxidation procedure.

Figure 6 
                     Degradation of dyes by ferrate with the effect of contact time.
Figure 6

Degradation of dyes by ferrate with the effect of contact time.

Figure 6 demonstrates that the destruction of MB produces the highest yield (almost 100 %), followed by MO and RBB, and the reaction time is classified as fast. After entering the reactor, ferrate needs times to diffuse in dye solution [30]. For MB, from the first 5 min, the removal percentage was already high, and with increasing contact time, there was an increase but not too significant in removal percentage. In such cases, the active sites are saturated and the dye molecules in the adsorbent (ferrate) have achieved equilibrium [31]. When it was considered by the amount of dye taken against the weight of ferrate as a oxidator, it can be seen that MB gives a large yield of about 1,400 mg/g of ferrate used. For MO and RBB, there was a significant increase in destruction with contact time. The more complicated the structure of the dyestuff, the more it will be destroyed by ferrate. The results revealed that the dye clearance % rose with contact duration, with 60 min producing the greatest outcomes. The contact duration of 60 min, according to Talaiekhozani et al., is the optimum for eliminating COD and hydrogen sulfide using ferrate(vi) from household wastewater [32].

3.5.2 Reaction kinetics of dyes removal

Figure 7 shows the reaction kinetics of dye degradation (MB, MO, and RBB). Degradation of dyes by ferrate was the first order reaction indicated by the value of R 2 which was close to 1. The corresponding first-order kinetic plots shown in Figure 8 indicate that the initial rate of decolorization was much faster for MB than for MO and RBB.

Figure 7 
                     Kinetics curves of the dyes (MB (a), MO (b), and RBB (c)) destruction by ferrate.
Figure 7

Kinetics curves of the dyes (MB (a), MO (b), and RBB (c)) destruction by ferrate.

Figure 8 
                     The maximum wavelength of (a) MB, (b) MO, and (c) RBB at various pH.
Figure 8

The maximum wavelength of (a) MB, (b) MO, and (c) RBB at various pH.

3.5.3 Study of dyes on pH

To evaluate the influence of pH on the shift of the maximum wavelength peak used for concentration measurements before and after ferrate treatment was done by measurement the dye UV-Vis spectrum with different pH. Figure 8, especially for MB, there is a drastic color change when the pH is high from blue to purple, and it occurs because the ferrate used is derived from extreme base synthesis. So that measurements at alkaline pH must be carried out at the maximum peak of the dye under alkaline conditions.

3.6 Characterization

3.6.1 Ferrate characterization

The Fe–O bond stretching vibration peaks in ferrate is detected at roughly 700, 766, and 877 cm−1, demonstrating the presence of Fe–O bonds in the crystal, namely sodium ferrate salt [27,33]. The H–O bond from water is responsible for the peaks at 1,651 cm−1 and between 2,492 and 3,500 cm−1 [34,35]. The strength of the ferrate characteristic vibrational peak is strong in Figure 9a, suggesting that the purity of the produced ferrate is greater. The graph shows the diffraction peaks at 2θ = 27.68, 30.16, 32.33, 34.53, 35.28, 40, 45.41, and 56.75° which is in agreement with the reference (Na2FeO4). Figure 9b shows a significant resemblance and proves the crystal structure of ferrates, as discovered by El Maghraoui et al. [27]. Therefore, the obtained product is sodium ferrate (Na2FeO4), according to the XRD and FTIR measurements.

Figure 9 
                     XRD (a) and FTIR spectra (b) of ferrate.
Figure 9

XRD (a) and FTIR spectra (b) of ferrate.

The elemental composition of ferrate is shown in Table 2. From the table, we know that the sample contains Na, Fe, and most of the other Fe was found in its oxide form with the value 93.774%. The existence of the components in XRF demonstrates that the target substance, Na2FeO4, was formed.

Table 2

XRF Analysis of Ferrate

Component Unit Analysis result value
Na % Weight 5.1802
Si 0.2461
Cl 0.8394
K 0.0204
Ca 0.0099
Fe 0.0237
Zn 0.0519
Ag 0.0545
Balance 93.574

3.6.2 Characterization of dye destruction by ferrate

Tests with LC/MS gave the spectral results as shown in Figure 10 for MB, MO, and RBB.

Figure 10 
                     LC/MS spectra of MB (a) & (b), MO (c) & (d), and RBB (e) & (f) before and after treated with ferrate.
Figure 10

LC/MS spectra of MB (a) & (b), MO (c) & (d), and RBB (e) & (f) before and after treated with ferrate.

MB degradation begins with the breakdown of the C–S═C bond, followed by the dissociation of the two benzene rings. The aromatic ring is subsequently hydroxylated, yielding phenolic metabolites. The last aromatic component discovered before the ring opened was hydroxyhydroquinone. The amino group can either produce an ammonium ion, which is slowly reduced to nitrate, or it can be immediately converted to hydroxylamine, which is then oxidized to nitrate. The existence of the MB molecule was established by the signal at m/z 284. Scheme 1 depicts the sequential hydroxylation processes of the MB benzene ring. The addition of hydroxyl promotes ring-opening and the formation of mineralized products. The intermediate will decompose and mineralize into CO2, H2O, SO 4 2 and NO 3 [36,37,38].

Scheme 1 
                     Degradation of MB by ferrate.
Scheme 1

Degradation of MB by ferrate.

It can be observed that the MO spectra show a peak at m/z 306, which indicates MO characteristics as shown in Scheme 2. Meanwhile, the MO chromatograms degraded by ferrate showed a total loss of MO peaks and new m/z peaks at 165, 153, and 97, corresponding to the intermediate product of MO degradation. The product formed at m/z 165 is dimethyl-(4-nitro-phenyl)-amine, followed by further degradation of the intermediate product. MO breakdown products can be further degraded to inorganic ions (ammonium, nitrate, and sulfate), H2O, and CO2 [39,40,41].

Scheme 2 
                     Degradation of MO by ferrate.
Scheme 2

Degradation of MO by ferrate.

Based on the intermediates identified by LC-MS, possible degradation pathways of RBB dyes have been suggested and are shown in Scheme 3. The schematic shows that RBB dyes are degraded through the breakdown of azo bonds (N═N) into intermediate products. The breakdown of dyestuffs via the breaking of azo linkages results in the creation of aromatic amines. Fragmentation continues until CO2 and H2O are formed as the end result of the ferrate degradation of remazol black B [42,43].

Scheme 3 
                     Degradation of remazol black B by ferrate.
Scheme 3

Degradation of remazol black B by ferrate.

4 Conclusion

The synthesis of NaOCl by electrolysis of table salt has been successfully obtained. NaOCl reacted with iron(iii) ions in alkaline condition to produce purple ferrate. The resulted product is sodium ferrate (Na2FeO4) according to the XRD and FTIR measurements. The ferrates were used to remove MB, MO, and RBB dyes with contact times of 5–60 min. The ferrate dyes can be removed through destruction process, resulting in almost 100% removal of the MB sample without any harmful by-products.


The authors are thankful to the DIKTI (PDUPT scheme) for funding this work.

  1. Funding information: The Ministry of Research and Technology/National Research and Innovation Agency financed this work through the Directorate of Research and Community Service, Deputy for Strengthening Research and Development (PDUPT research grant with contract number 225-91/UN7.6.1/PP/2021).

  2. Author contributions: The published version of the work has been reviewed and approved by all authors.

  3. Conflict of interest: The authors state that they have no conflict of interest.

  4. Ethical approval: The research was not undertaken for human or animal use.

  5. Data availability statement: All the data is available within the manuscript.


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Received: 2022-08-08
Revised: 2022-09-28
Accepted: 2022-10-04
Published Online: 2022-10-31

© 2022 Gunawan Gunawan et al., published by De Gruyter

This work is licensed under the Creative Commons Attribution 4.0 International License.

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