(670e) Formic Acid Oxidation By the Fenton and Fenton-like Reaction: Experimental Study and Kinetic Modelling

Formic acid oxidation by the Fenton and Fenton-like
reaction: experimental study and kinetic modelling

The Fenton reaction is
considered to be an efficient purification technique for the treatment of
wastewaters from a wide range of industries [1].  The Fenton reaction,
known as one of the advanced oxidation processes, is the mixing of ferrous iron
with hydrogen peroxide to form hydroxyl radicals:

These hydroxyl radicals 
are non-selective and powerful in the oxidation of various organics. The
Fe(III) formed can react in a reaction know as the Fenton-like reaction:

The Fenton reaction is known
to be faster than the Fenton-like reaction (rate constants: 53 and 2·10^-3
M^-1 s^-1 for the Fenton and Fenton-like reaction
respectively). The Fenton and Fenton-like reactions start a sequence of 
radical combination and propagation reactions, as well as the reactions with
the organic compounds of interest.

 The large number of reactions that have to be
taken into account in the Fenton reaction makes the kinetic modelling
challenging. In addition, the radical reactions that have to be included in the
model are usually extremely fast and it is therefore difficult to determine the
reaction rate constants experimentally. Furthermore, the organic compounds can
form various intermediate compounds, which add to the reactions describing the
oxidation in the Fenton reaction. For example, Duesterberg et al. described the
oxidation of formic acid with a set of 20 reactions [2].
However, more reactions have been suggested in literature. For instance, it is known that Fe(III) can form
complexes with water and H₂O₂, such as [Fe^III]³⁺, [Fe^IIIOH]²⁺,
[Fe^III(OH)₂]⁺ and
[Fe₂^III(OH)₂]⁴⁺. In addition, it
has been shown that ferric-formate complexes could be
formed in the Fenton and Fenton-like reaction [3] [4].

In this study it was chosen to
study the oxidation of formic acid to limit the amount of intermediate
compounds that can be formed in the oxidation process of the organic. In
contrast to previously reported studies on the formic acid oxidation by the
Fenton reaction [2][5], our study mainly focuses on relatively high
concentrations of formic acid (0.01M) and low ratios of iron and H₂O₂
to formic acid. The aim of our work was to extend the current knowledge on the
oxiation of formic acid by the Fenton reaction and to find a kinetic model that
describes the formic acid oxidation in the both the Fenton and Fenton-like
reaction.

In the experimental study,
batch experiments were perfomed at constant temperature. The results of these
experiments show that the oxidation rate of formic acid in both the Fenton and
Fenton-like reaction increases with increasing H₂O₂
concentrations (Figure 1). The most important differences between the Fenton
and Fenton-like reactions are observed in the first minutes of the reaction. In
the Fenton reaction a steep decrease in formic acid is observed in the first
minute of the reactions, which is independent of the H₂O₂
concentration. Afterwards the reaction rate is constant. In contrast, the
reaction rate is relatively constant for the Fenton-like reaction for the
entire course of the reaction.

It was decided
to study the reactions spectrophotometrically to obtain more information on the
different iron species in the reaction solutions. Figure 2 shows the absorbance
spectra during the Fenton and Fenton-like reaction. Previous studies have shown
that he absorbance at  290 nm could be attributed to Fe^III]³⁺,
[Fe^IIIOH]²⁺, [Fe^III(OH)₂]⁺ and [Fe₂^III(OH)₂]⁴⁺
[6] [49] . [Fe₂^III(OH)₂]⁴⁺
and [Fe^III(OH)₂]⁺
are unlikely to form below pH 3 [6] and therefore not taken into account.
Although a fraction of the absorbed iron species can be regarded [Fe^III]³⁺, its decadic
molar absorptivity (315 M^-1 cm¹) is relatively low
compared to [Fe^IIIOH]²⁺ (2005 M^-1
cm^-1) and the absorbance at 290 nm is mainly attributed to [Fe^IIIOH]²⁺ [28].

Kinetic models
previously presented in the literature are unable to describe our experimental
results. For example, the model introduced by Duesterberg
et al. shows a fast oxidation rate of formic acid in the first minutes of the
reaction, which can be ascribed to the excess of Fe(II)
and H₂O₂ to formic acid in their study. The use of this
model for our experimental results gives a predicted initial decrease in formic
acid in the Fenton reaction that is proportional to the ratio of Fe(II) to formic acid, which is shown in Figure 3 (A).
However, the initial decrease observed is 30% for a ratio of formic to Fe(II) of 10:1. The fit to the Fenton-like data is
reasonably well although significant differences are observed for the highest
concentration of H2O2 (Figure 3B). To improve the fit of the model to the
experimental data, the reactions describing the formation of [Fe^IIIOH]²⁺
and the ferric-formate complex were added to
the model suggested by Duesterberg et al.

Figure 3 (C
and D) shows that the addition of these reactions describing the formation and
decomposition ferric complexes improved the fit of the kinetic model to the
experimental data. This could suggest that the formed complexes influence the
availability of free iron in the solution.

The limitation
regarding the kinetic modelling of the Fenton reaction is the number of
reactions and kinetic constants that have to be included to obtain an accurate
fit to the experimental data. Nonetheless, the suggested reactions and kinetic
constants could give a reasonable fit to our experimental data without using
fitting routines to obtain new kinetic constants.

Our work has
led us to the conclusion that the initial oxidation of formic acid in the first
minutes of the reaction is significantly different for the Fenton and
Fenton-like reaction. These experimental results could not be described by
kinetic models that are proposed in literature. The addition of the formation
of ferric-complexes (both inorganic and organic) improved the fit of the model
to the experimental data, especially for the initial decrease in formic acid in
the first minutes of the Fenton reaction. However, a perfect match with the
experimental data was not achieved.

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2693–2706, 1999.

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