ORIGINAL ARTICLES

 

Catalytic wet peroxide oxidation of formic acid in wastewater with naturally-occurring iron ore

 

 

E Erasmus^{I, *}; JO Claassen^II; WA van der Westhuizen^II

^IDepartment of Chemistry, University of the Free State, Bloemfontein, 9300, Republic of South Africa
^IIDepartment of Geology, University of the Free State, Bloemfontein, 9300, Republic of South Africa

 

 

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ABSTRACT

The catalytic wet oxidation of formic acid, using hydrogen peroxide as the oxidizing agent over naturally-occurring iron ore, was explored. Firstly, the decomposition of hydrogen peroxide to its hydroxyl radicals (HO^- and HOO^-) over naturally-occurring iron ore was investigated. The reaction was monitored by ATR FTIR by following the disappearance of the O-H peak of H₂O₂ at 2 860 cm^-1. Decomposition occurred according to the Fenton mechanism and resulted in observed first-order rate constants one order of magnitude faster than that without the catalyst. Turnover frequencies (TOF) of 1.97-8.85 x 10^-9 s^-1 were obtained for the decomposition of H₂O₂. The wet oxidation of formic acid using hydrogen peroxide as the oxidizing agent over naturally-occurring iron ore reaction was also monitored by ATR FTIR, following the disappearance of the carbonyl stretching frequency of formic acid at 1 727 cm^-1. Experiments were performed at different hydrogen peroxide (2, 4, 6 and 8M) and formic acid (1.26, 2.52, 6.3 and 12.6 M) concentrations as well as with varying amounts of naturally-occurring iron ore catalyst, at pH = 2. Elevated hydrogen peroxide and formic acid concentrations led to increased observed first-order kinetics, as high as k_obs= 21.75 x 10⁴ min¹ with a TOF = 1.73 x 10^-8 - 1.12 x 10^-6 s^-1.

Keywords: iron ore, catalytic wet oxidation, formic acid, Fenton, hydrogen peroxide

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INTRODUCTION

Apart from the large-scale exploitation of iron containing ore as raw material feed for iron extraction, it is one of the few naturally-occurring minerals that have been exploited for alternative applications. As a result, numerous studies report on the use of iron-containing minerals as adsorbents (Nowack et al., 2006; Rovira et al., 2008; Zeng et al., 2004; Spiteri et al., 2008; Guo et al., 2007; Zhang et al., 2004; Pirillo et al., 2009) and catalysts (Costa et al., 2008; Huang et al., 2000; Centi et al., 2000; Andreozzi et al., 2003; Du et al., 2008) for the removal of pollutants from wastewaters.

Formic acid is a pollutant in wastewater from the effluents of some industries (Claudel et al., 1984). It is also a stable intermediate from the decomposition of many organic pollutants (Chou et al., 1998; Davis et al., 1999; Hwang et al., 1999; Dinsdale et al., 2000), which is difficult to oxidize further (Ogata et al., 1981; Harmsen et al., 1997). For this reason, formic acid is often used as a model organic pollutant in wastewater treatment studies (Kawaguchi et al., 1993; Kim et al., 1996; Butterfield et al., 1997; Candal et al., 1997; Miller et al., 1999).

Wet air oxidation is a procedure commonly used to reduce the total organic carbon (TOC) in industrial wastewater (Matatov-Meytal et al., 1998; Mishra et al., 1995; Levec et al., 1995; Luck et al., 1996; Gallezot et al., 1997; Klinghoffer et al., 1998; Romalho et al., 1983.). The reaction however requires high operating temperatures. The use of oxidation agents allows the process to precede at lower reaction temperatures. Ozone and hydrogen peroxide (H₂O₂) are the preferred oxidants. These compounds are active even at room temperature and decompose to environmentally friendly O₂ and H₂O.

Water treatment with ozone is well known (Masten et al., 1994; Luck et al., 1997; Allemane et al., 1993; Andreozzi et al., 1996). Its effectiveness is however limited by lower solubility at atmospheric pressure and the short lifetime of ozone renders the process more expensive. H₂O₂ is therefore the preferred oxidant.

The use of hydrogen peroxide to oxidize the pollutants in wastewater requires it to be in its radical form. The production of hydroxyl radicals from H₂O₂ and minimization of its decomposition to water is achieved either photochemically or by catalysis. The photochemical activation of hydrogen peroxide requires UV-radiation. However, due to its low extinction coefficient, the operation process is costly (Ojima et al., 1994). The formation of the hydroxyl radicals using iron ion catalysis is well known (Edwards et al., 1992; Walling et al., 1975). The process proceeds according to the Fenton mechanism as illustrated below (Centi et al., 2000):

Numerous examples are available for the oxidation of different organic pollutants in wastewater, using a combination of iron ion and hydrogen peroxide. There are however very few publications available on the use of naturally-occurring iron ore (such as hematite, ferrihydrite, goethite and lepidocrocite) as a cheap alternative for the oxidation of pollutants such as formic acid in wastewater.

This study utilizes hard high-grade hematite ore as such an alternative. The ore is massive Gamagaran type ore from the Sishen Iron Formation in the Northern Cape Province of the Republic of South Africa. It comprises sand to silt size hematite grains in a matrix which appears as ferruginized, hematitic sandstone and shale. Small amounts of specular hematite are also present (Wilson et al., 1998).

 

EXPERIMENTAL

Chemicals and instruments

The hematite ore used in this study was sourced from Sishen Iron Ore Mine (see Fig. 1). The hematite was crushed and screened with only the particles in the size range 75-212 μm used. All reagents (Merck, Sigma-Aldrich) were used without further purification and water was double distilled. Attenuated total reflectance Fourier transform infrared (ATR FTIR) spectra were recorded on a Nicolet IS 50 infrared spectrophotometer equipped with a diamond ATR. X-ray photoelectron spectroscopy (XPS) data were recorded on a PHI 5000 Versaprobe system, with monochromatic Al Kcx X-ray source. Spectra were obtained, using the aluminium anode (Al Kcx = 1486.6 eV), operating at 50 μm, 12.5 W and 15 kV energy. The survey scans were recorded at constant pass energy of 187.85 eV. The detailed regional scans were recorded at constant pass energy of 29.35 eV and analyser resolution of < 0.5 eV. The background pressure was 2 x 10^-8 Pa. The XPS data were analysed with Multipak version 8.2c computer software (Moulder et al., 1995), applying Gaussian-Lorentz fits (the Gaussian/Lorentz ratios were always > 95%). The iron ore were characterized by X-ray diffraction (XRD) using a Panalytical Empyrean diffractometer for elemental analysis.

 

 

Hydrogen peroxide decomposition test without the catalyst

2M hydrogen peroxide solutions in water were prepared according to Table 2, while stirring at room temperature. The pH was adjusted with the addition of either H₂SO₄ or NaOH similar to published procedures (Li, 1999). The final test volume was 10 mL. Samples were taken at regular intervals and analysed by ATR FTIR.

 

 

 

Hydrogen peroxide decomposition test with the naturally-occurring iron ore added as catalyst

Hydrogen peroxide reaction mixtures of different concentrations (2, 4, 6 and 8 M) were prepared in water according to Table 2, while stirring at room temperature. After the pH was adjusted with the addition of H₂SO₄ to obtain a pH of 2, 20 mg crushed Fe₂O₃ ore was added to the reaction mixture. Samples were taken at regular intervals and analysed by ATR FTIR.

Hydrogen peroxide decomposition and formic acid oxidation test with Fe₂O₃ catalyst

Reaction mixtures (10 mL) were prepared according to Table 3, while stirring at 35°C. After the pH was adjusted to 2 with the addition H₂SO₄, the required amounts of formic acid and Fe₂O₃ catalyst (iron ore) were added. Samples were taken at regular intervals and analysed by ATR FTIR.

Turnover frequencies (TOF) were determined by the following equation:

In addition, pseudo first-order rate constants, k_oobs, were calculated by fitting kinetic data to the first-order equation

(Espenson, 1995) [A]t = [A]₀ e⁽k^obsi⁾ with [A]t and [A]₀ the absorbance of the indicated species at time t and at t = 0.

Preliminary stability testing of the natural iron ore

A 50 mg sample of the naturally-occurring iron ore was stirred in a solution of 4 M hydrogen peroxide at pH 2 for 10 days. This test showed that no iron leaching occurred during the time period, which was the same as was used for the catalytic testing.

 

RESULTS AND DISCUSSION

Chemical analysis of the iron ore

XRD analysis of the iron ore revealed that the sample contains 98.13% Fe₂O₃, 0.45% Al₂O₃ and 1.12% SiO₂.

X-ray photoelectron spectroscopy

The XPS spectra of the iron ore (Fig. 2) revealed that the sample contained Fe, O, C (possibly from CaCO₃) and Si. Table 1 summarizes the atomic percentages and binding energies (eV) of these elements. XPS samples most often exhibit adventitious carbon contamination which is detected by XPS analysis (Castaneda et al., 2010; Kyotani et al., 1997). The binding energy of these carbonaceous deposits is set at 284.9 eV and the other element's binding energies are referenced against it. The detected silicon is most probably from mineral contaminants such as quartz, as indicated by the XRD analysis.

The binding energy of the Fe 2p_3/2 (see Fig. 2 right) was detected at 710.8 eV, which is in agreement with published values of 710.8 eV for Fe₂O₃ (Brion, 1980; Nefedov et al., 1975). The detailed scan of the oxygen 1s area (see Fig. 2 right) showed that there are two distinct peaks present. The peak observed at 531.5 eV, which accounts for 85.7% of the oxygen present in the sample, is from adsorbed oxygen and other adventitious sources, normally present in all samples. The peak observed at 529.3 eV, which accounts for 14.3% of the oxygen present, is allocated to oxygen present in iron oxide Fe₂O₃. When calculated, the peak at 529.3 eV accounts for 4.3% of the total atomic percentage. This value is in agreement with the amount of oxygen expected if the iron content in Fe₂O₃ is 2.8%. The binding energy of the O 1s (529.3 eV) in Fe₂O₃, is in close agreement with published values of 529.5 and 529.6 eV for Fe₂O₃ (Haber et al., 1976; Nefedov et al., 1977).

Decomposition of hydrogen peroxide

Hydrogen peroxide is a thermodynamically unstable compound that decomposes to form oxygen and water. Various factors influence the rate of decomposition, including temperature and pH. The decomposition of hydrogen peroxide is enhanced when catalysts such as Pt and MnO₂ are added (Petrucci, 2007). However when certain metal ions, including Fe²⁺, are used as a catalyst the decomposition proceeds according to the Fenton mechanism to produce the free radicals HO' and HOO'.

Since this study is concerned with the decomposition of hydrogen peroxide using naturally-occurring iron ore, it is also important to investigate the decomposition of hydrogen peroxide in the absence of iron ore.

Figure 3 shows the comparative ATR FTIR graphs of hydrogen peroxide, water, a 1:1 mixture of hydrogen peroxide and water, and an 8 M hydrogen peroxide solution in water. The distinctive frequencies associated with water are the -OH bending frequency at 1 646 cm^-1 and the -OH stretching frequency in the area 2 800-3 800 cm^-1. Hydrogen peroxide shows 4 peaks: 1 370, 1 646, 2 860 cm^-1, and 2 800-3 800 cm^-1. These peaks are assigned to the asymmetric deformation band of H₂O₂ (v₁), the -OH bending (v₂), a combination of v₁, v₂ and v₃, and the -OH stretching frequency (v₃), respectively. These assignments are in agreement with published data from Giguere (1950). The peak observed at 2 860 cm^-1 was used to follow the decomposition of hydrogen peroxide.

 

 

The decomposition of 2M hydrogen peroxide solutions in water was followed by ATR FTIR at different pH levels. Figure 4 shows the ATR FTIR spectra of the decomposition of a 2 M solution of hydrogen peroxide in water at a pH of 9. It is known from literature that hydrogen peroxide is more stable at lower pH levels, as indicated in Table 2. Thus the rest of the experiments were conducted at a pH of 2 to keep self-decomposition to a minimum.

 

 

The decomposition of hydrogen peroxide solutions of different concentrations (2, 4, 6 and 8 M) were followed by ATR FTIR at pH = 2 in the presence of naturally-occurring iron ore (Fe₂O₃).

Figure 5 depicts the ATR FTIR spectra of the decomposition of an 8 M solution of hydrogen peroxide in water at a pH of 2 in the presence of a catalytic amount (20 mg) of naturally-occurring iron ore.

 

 

The observed rate constants obtained for the different concentrations of hydrogen peroxide (see Table 2) are, within experimental conditions, the same, at approx. 1.0 x 10^-5min^-1. This is approximately one order of magnitude larger than values observed for the decomposition of hydrogen peroxide in the absence of iron ore. The slower than expected rate of decomposition of the hydrogen peroxide in the presence of the iron ore can be attributed to, firstly, the small amount of active catalyst (the small amount of Fe₂O₃ at the surface of the powdered sample available for catalysis). Secondly, it is known that iron in its oxide form is a more ineffective catalyst for the decomposition of hydrogen peroxide than other iron ions such as FeCl₂, FeCl₃, Fe(NO₃)₃, etc. Lastly, the low solubility of the iron ore at the chosen conditions affects diffusion during the reaction.

Iron oxide may be a relatively slow catalyst for the decomposition of hydrogen peroxide. Since the aim of the study is to catalytically oxidize formic acid using the hydroxyl radical, fast decomposition of the hydrogen peroxide is not required.

The observed rate of hydrogen peroxide decomposition over naturally-occurring iron ore (measured in this study) compares well with studies conducted on hematite (0.9 x 10^-5 min^-1) and other iron-containing catalysts reported in literature (Huang et al., 2001).

Catalytic wet oxidation of formic acid

Figure 6 depicts the comparative ATR FTIR graphs of formic acid and a variety of different concentrations (2, 4, 6 and 8 M) of formic acid dissolved in water at pH = 2. The arrows (in the bottom spectra) indicate how the peaks increase or decrease as the concentration of the formic acid is increased. The distinctive frequencies observed for formic acid dissolved in water appear at 2 856, 1 727, 1 413 and 1 200 cm^-1, representing the C-H stretching frequency (v C-H), the C=O stretching frequency (v C=O) and the two CH₂ bending frequency (u>CH₂), respectively. These assignments were made in agreement with published data from Popova et al. (2000).

 

 

Since the C-H stretching frequency of formic acid overlaps with the distinctive combination peaks (v₁, v₂ and v₃) of H₂O₂ at 2 860 cm^-1, the C=O stretching frequency at 1 727 cm^-1 was used to follow the oxidation of formic acid in the following experiments.

The wet oxidation of formic acid by decomposed hydrogen peroxide radicals (HO' and HOO') over naturally-occurring iron ore solutions was followed by ATR FTIR using different concentrations of hydrogen peroxide (2, 4, 6 and 8 M) and formic acid (1.26, 2.52, 6.3 and 12.6 M) as well as varying amounts of Fe₂O₃ catalyst at pH = 2. Figure 7 shows the ATR FTIR spectra of the wet oxidation of an 8 M formic acid solution by the hydroxyl radicals (HO' and HOO') from decomposed 2 M H₂O₂ in water at a pH of 2 in the presence of a catalytic amount of naturally-occurring iron ore. The different experimental conditions and its observed first-order rate constants are presented in Table 3. These observed rate constants are two orders of magnitude larger than that observed for the decomposition of hydrogen peroxide. This could be attributed to an increased amount of catalyst used and the formation of two hydroxyl radicals that can react with the formic acid.

 

 

From Table 3 (Experiments A1-A6) and Fig. 8, it is evident that the decomposition rate increases with an increase in the hydrogen peroxide concentration. This is expected since higher concentrations of hydrogen peroxide will result in the formation of more hydroxyl radicals, which in turn would oxidize the formic acid faster.

 

 

Experiments A1, B1, C1 and D1, where the hydrogen peroxide concentration and amount of iron ore catalyst are kept constant while the concentration of the formic acid is being varied, indicated that the rate of decomposition also increases as the concentration of the formic acid increases (see Fig. 8). More formic acid molecules present in solution would lead to more interaction between the formic acid and the hydroxyl molecules which would in turn lead to a faster formic acid oxidation rate being observed.

In the next set of experiments, A1 and A7-9, the hydrogen peroxide and formic acid concentrations were kept constant while the amount of catalyst was varied (see Table 3). From Fig. 9, it can be seen that the observed rate constant for the catalytic wet oxidation of formic acid exponentially increases as the amount of catalyst increases.

 

 

CONCLUSION

It has been confirmed that formic acid can be effectively wet oxidized by the decomposed products of hydrogen peroxide over a naturally-occurring iron ore catalyst. The decomposition of hydrogen peroxide to its hydroxyl radicals (HO' and HOO') is influenced by the pH of the solution. At higher pH, hydrogen peroxide is more unstable and decomposition can occur spontaneously. At a constant pH of 2, an increase in the concentration of hydrogen peroxide in aqueous solutions does not influence the rate of catalytic decomposition over naturally-occurring iron ore with an approx. k_obs= 1 x 10⁵ min¹ observed under the chosen experimental conditions.

The catalytic wet oxidation of formic acid by decomposed hydrogen peroxide radicals (HO' and HOO') over naturally-occurring iron ore indicated that higher levels of hydrogen peroxide and formic acid concentration result in higher formic acid oxidation rates. This can be attributed to increased amounts of reacting molecules present in solution.

 

ACKNOWLEDGEMENTS

The authors acknowledge the financial support provided by the University of the Free State.

 

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Received 20 July 2015
Accepted in revised form 3 June 2016

 

 

* To whom all correspondence should be addressed.+27 51 401-9656 / +27 62 704 6350; e-mail: erasmuse@ufs.ac.za