Development and design of the in-situ regeneration section of Vitrisol®, a novel, highly
selective desulphurization process
Wermink, Wouter Nicolaas
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Publication date: 2019
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Wermink, W. N. (2019). Development and design of the in-situ regeneration section of Vitrisol®, a novel, highly selective desulphurization process. Rijksuniversiteit Groningen.
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9
Chapter 2: The oxidation of Fe(II) in acidic sulphate solutions with air at
elevated pressures. Part 1. Kinetics above 1 M H
2SO
4Reproduced with permission from Ind. Eng. Chem. Res. 2017, 56, 14, 3775-3788, link. Copyright 2017 American Chemical Society.
Abstract
The oxidation of ferrous ions in acidic sulphate solutions at elevated air pressures was investigated. The effect of the Fe2+ concentration, initial H2SO4 concentration and partial oxygen pressure on the
reaction rate were determined at three different temperatures, i.e. T = 90 °C, 70 °C and 50 °C. The effect on the reaction rate of the components that H2SO4 dissociates into, i.e. HSO4-, H3O+ and SO42-,
was established as well. A second order of reaction in Fe2+ and a first order of reaction in O2 were
determined. No clear order in either H2SO4, or the components H2SO4 dissociates into, could be
established. For the experiments with initial concentrations of H2SO4 of 1 M and higher the oxidation
rate was not affected, i.e. a zero order of reaction in H2SO4 for these concentrations. Therefore the
kinetic rate expression for the oxidation of Fe2+ at concentrations of H2SO4 of 1 M and higher can be
calculated with:
[ ]
[ ] Where the activation energy was determined to be EA = 60.3 kJ/mol.
2.1 Introduction
Hydrogen sulphide (H2S) is a highly toxic and corrosive gas. Removal of H2S from acidic gas streams,
such as natural gas, industrial gas or biogas, is important for safety, health, environmental and economic reasons. Several regenerative and non-regenerative H2S removal processes are readily
available, which are economically viable only for specific gas compositions and gas flow rates. Apart from non-regenerative H2S removal by the use of e.g. adsorbents, all the regenerative aqueous liquid
redox desulphurization processes (e.g. THIOPAQ, LO-CAT, SulFerox) capture CO2 to varying extents
besides H2S.
The conventional method of removing H2S from natural gas is using an amine process. Subsequently,
the H2S is converted to elemental sulphur by a consecutive Claus process. For natural gas fields,
usually containing more CO2 than H2S, this will result in an inlet acid gas stream for the Claus process
that is low in H2S and high in CO2 content. The inlet gas stream should contain at least 20 mol.% of
H2S to be able to produce a stable flame in a Claus furnace. Modification of the Claus process is
needed between 20 and 50 mol.% H2S in the inlet acid gas stream. Above 50 mol.% H2S content no
modification of the Claus process is required.1,2 Moreover, owing to the coabsorption of CO2 the
regeneration costs of the amine process is substantially increased.
The novel Vitrisol® desulphurization process is based on the removal of H2S by precipitation with
copper sulphate (CuSO4) in an aqueous, acidic solution.3 Copper sulphide (CuS) and sulphuric acid are
formed in the gas treating process:4,5
( ) ( ) (2.1)
Ter Maat et al.4 developed an equilibrium model based on the precipitation of Cu2+ with sulphide and carbonate. A pH-region as function of H2S and CO2 partial pressures, in which CuS precipitation could
and CuCO3 and Cu(OH)2 could not occur, respectively, was derived from the equilibrium model for a
10
MPa is required at a pH of 4 for a 1.0 M CuSO4 solution to form CuCO3 precipitate. Lowering the pH
results in higher CO2 partial pressures required to form CuCO3 precipitate, whereas CuS precipitation
is possible. Therefore, it can be concluded that hardly no CO2 is co-absorbed, neither physically nor
chemically, due to the highly acidic environment of the Vitrisol® absorption liquid (e.g. typical H2SO4
concentrations are between 0.1 M and 1.0 M).
The current status of the Vitrisol® process is a scavenger-like application. Cu2+, the active compound in the absorption liquid, becomes depleted during H2S removal. It must be noted, however, that
nowadays copper is an expensive commodity; increasing amounts of H2S lead to increasing
operational costs. To reduce the operational costs for large amounts of H2S and/or large scale
applications, a regeneration step needs to be developed to replenish Cu2+.
CuS can be leached with ferric sulphate (Fe2(SO4)3), an operation encountered in copper ore
processing (a/o CuS).6,7 Copper sulphate, elemental sulphur (So) and ferrous sulphate (FeSO4) are
produced in this process:
( ) ( ) ( ) (2.2)
Ferrous sulphate can be re-oxidized to ferric sulphate with O2 according to:
( ) (2.3)
Resulting in the overall net reaction for the removal of H2S:
( ) (2.4)
In this study the mechanism and reaction kinetics of Reaction 2.3, the oxidation of ferrous sulphate by oxygen in the presence of sulphuric acid, are studied.
2.2 Literature review
The oxidation reaction of ferrous ions in acidic sulphate solutions has previously been investigated by several authors.
McBain8 studied the oxidation of ferrous ions at temperatures of 30 °C and 14.5 °C. One set of experiments was performed by adding a known amount of water, saturated with oxygen, to the solution to be oxidized. Oxygen was removed from water, used to prepare solutions to be oxidized, prior to the experiments. Initial concentrations of oxygen, FeSO4 and H2SO4 were 9.17 x 10-4 M to
1.85 x 10-3 M, 0.0184 M to 0.463 M and 0.01 M to 0.5 M, respectively. One additional experiment was performed to investigate the influence of sulphate ions by adding Na2SO4. Initial concentrations
of oxygen, FeSO4, H2SO4 and Na2SO4 of the respective experiment were 9.30 x 10-4 M, 0.0935 M, 0.01
M and 0.5, M respectively. These experiments were performed at a temperature of 30 °C.
Another set of experiments was performed with air and an O2/N2 mixture with a ratio of 90/10, both
at 1 atm. First a bottle was filled for 25 % with the solution to be oxidized, next the bottle was filled with either air or O2/N2 mixture and closed. The prepared bottles were placed on a shaking table.
Initial concentrations of H2SO4 and FeSO4 were 0.34 M and 0.68 M, respectively. These experiments
were performed at a temperature of 14.5 °C.
McBain concluded from the experiments performed at a temperature of 30 °C a first order of reaction in oxygen and a second order of reaction in ferrous sulphate, respectively. At low H2SO4
concentrations of 0.01 M and 0.05 M, respectively, no effect of H2SO4 on the oxidation rate was
11 acid concentration above 0.01 M H2SO4. An increase in oxidation rate was observed with the addition
of Na2SO4.
A first order of reaction in oxygen was concluded from the experiments performed at a temperature of 14.5 °C.
Lamb and Elder9 investigated the oxidation of ferrous ions by means of electromotive force measurements. Experiments were performed at a temperature of 30 °C. During the experiments either air or pure oxygen was passed through the solution. Initial concentrations of FeSO4 and H2SO4
were 0.146 M to 1.36 M and 0.008 M to 3.0 M, respectively.
Lamb and Elder reported a first order of reaction in oxygen and a second order of reaction in ferrous sulphate. According to their work the rate of oxidation of ferrous sulphate was independent of the concentration of sulphuric acid at concentrations starting at 0.23 M and higher, but increased progressively with lower concentrations. They ascribed this phenomenon to the formation of ferrous hydroxide at low acid concentrations.
Pound10 studied the oxidation of ferrous sulphate solutions exposed to air and at temperatures ranging from 7 °C to 28 °C. Initial concentrations of H2SO4 and FeSO4 were 5.0 x 10-5 M to 6.0 M and
0.1 M for pure sulphate solutions, respectively.
Pound observed a higher oxidation rate of neutral solutions of ferrous sulphate compared to the moderately acid solutions. At 25 °C, the rate of oxidation decreases, changing from a neutral solution to a sulphuric acid concentration of approximately 1 M. The formation of precipitates was observed at sulphuric acid concentrations below 0.005 M.
Huffman and Davidson11 investigated the oxidation of ferrous ions in sulphuric acid solutions at temperatures ranging from 138 °C to 180 °C as well as 30.5 °C. Partial oxygen pressures were either 0.0213 MPa or 0.101 MPa oxygen. Ferrous ammonium sulphate was used as a source for ferrous ions, therefore ammonium ions were present in the experimental solutions.
The effect of Fe2+ on the reaction was studied at temperatures ranging from 138 °C to 180 °C with initial concentrations of Fe2+ and H2SO4 of 0.001 M to 0.025 M and 1 M, respectively. The effect of
Fe2+ was studied at a temperature of 30.5 °C as well with initial concentrations of Fe2+ of 0.001 M to
0.02 M, and either 1 M H2SO4, or 0.226 M H2SO4 and 0.354 M Na2SO4. The effect of Fe3+ on the
reaction was studied at a temperature of 160 °C with initial concentrations of Fe3+, Fe2+ and H2SO4 of
0.001 M to 0.02 M, 0.001 M to 0.005 M and 1 M, respectively. The effect of oxygen was studied at a temperature of 160 °C in 1 M H2SO4 solutions. Fe2+ concentrations were varied, but not mentioned.
The effect of SO42- was studied at a temperature of 30.5 °C with initial concentrations of Na2SO4, Fe2+
and H2SO4 of 0 M to 0.354 M, either 0.02 M or 0.04 M and 0.068 M to 0.226 M, respectively. The
effect of H3O+ was studied at a temperature of 30.5 °C with initial concentrations of Na2SO4, Fe2+ and
H2SO4 of 0 M to 0.0664 M, either 0.02 M or 0.04 M and 0.0203 M to 1 M, respectively. The solutions
used to study the effects of SO42- and H3O+ were kept at an ionic strength of 1.0 to 1.3 with NaClO4.
Speciation was calculated by assuming a value of 0.075 for the second ionization constant of H2SO4 at
an ionic strength of 1.0.
Huffman and Davidson reported that the oxidation of Fe2+ proceeded via two simultaneous paths; one with a first-order dependency in Fe2+, the other with a second-order dependency in Fe2+. It was determined that the oxidation reaction was not inhibited by ferric ions. They did not determine a clear first-order dependency in oxygen from first principal, though reported this in their kinetic equation. An increase in reaction rate, at a decreasing rate, was observed when the SO4
2-concentration was increased. The effect of H3O+ on the reaction could not adequately be
determined. However, the trend was observed that a decrease in H3O+ concentration resulted in a
slight increase in reaction rate. Activation energies were determined to be 56.1 kJ/mol for the first order term in Fe2+, and 68.2 kJ/mol for the second order term in Fe2+ in the temperature range of 140 °C to 180 °C. Activation energies varied when the experimental results at a temperature of 30.5 °C, as well as the experimental results obtained at high temperature, were used to derive kinetics.
12
Huffman and Davidson suggested that a first order dependence in Fe2+ is observed with anions forming strong complexes with ferric ions, e.g. H2P2O72-, F- and H2PO4-. A second order dependence in
Fe2+ is observed with anions forming weak complexes with ferric ions, e.g. ClO4-. The SO42- ion was
considered an anion that forms moderate complexes with ferric ions, because they determined a first, as well as a second-order dependency in Fe2+.
McKay12 and McKay and Halpern13 studied the oxidation of ferrous sulphate and pyrite separately. Ferrous sulphate oxidation experiments were performed at temperatures ranging from 100 °C to 130 °C and partial pressures of oxygen between 0.101 MPa and 0.405 MPa. Initial concentrations of H2SO4 and FeSO4 were 0.0125 M to 0.315 M and 0.048 M to 0.28 M, respectively. They mentioned
that experiments performed above a sulphuric acid concentration of 0.15 M could have resulted in corrosion of the reactor, a 316 stainless steel autoclave. Varying the stirrer speed between 450 and 830 rpm did not result in a change in reaction rate.
McKay and Halpern reported a first order of reaction in oxygen, a second order of reaction in ferrous sulphate and a fractional negative order in sulphuric acid. The addition of ferric ions (Fe3+) with a concentration of 0.054 M did not result in a change in reaction rate. Their rate law does not account for the inhibiting effect of sulphuric acid on the reaction rate, because the reaction rate constant k1
was determined from experiments carried out at an initial H2SO4 concentration of 0.08 M. The
activation energy was determined to be 69.0 kJ/mol.
Cornelius and Woodcock14 investigated the oxidation of ferrous sulphate in a stainless steel autoclave. Experiments were performed in the temperature range of 110 °C to 165 °C and at partial pressures of oxygen between 0.62 MPa and 2.4 MPa. Initial concentrations of H2SO4 and Fe2+ were
0.054 M and 0.018 M to 0.054 M, respectively. The paper does not clearly mention whether a stirrer was used for agitation.
Cornelius and Woodcock reported a first order of reaction in oxygen and a second order of reaction in ferrous ions. The activation energy was determined to be 61.9 kJ/mol.
Hotta and Suzuki15 studied the oxidation of ferrous ions in a glass reaction tube in a stainless steel reactor. Experiments were performed in the temperature range of 100 °C to 150 °C and at a partial pressure of oxygen of 3.04 MPa. Initial concentrations of H2SO4 and FeSO4 were 0.05 M and 0.01 M,
respectively.
Hotta and Suzuki reported a first order of reaction in ferrous ions. They claimed the activation energy to be near 56.1 kJ/mol.
Mathews and Robins16 investigated the oxidation of ferrous ions in an air lift percolator. Experiments were performed between 20 °C and 80 °C and at oxygen partial pressures (it must be noted that the unit of pressure is not clearly mentioned in the paper) of either 0.21 or 1.0. Studying the experimental setup suggests an open contact to the surroundings. Therefore the unit of pressure is assumed to be atm. H3O+ concentrations were estimated from pH measurements at ambient
temperature; H2SO4 concentrations were not reported. The initial pH and Fe2+ concentration varied
from a pH of 0 to 2 and 0.01 to 1.0 M, respectively. A pH above 2 was not investigated, because precipitation of ferric hydroxide could occur.
According to the authors the actual concentration of O2 could be determined from the saturated O2
concentration and the rates of dissolution and consumption of O2. The rate constant of dissolution of
O2 was estimated from experiments up to 50 °C in an air lift percolator. By changing the circulation
gas from nitrogen to air, the system response could be determined from dissolved oxygen concentrations, measured with an electrode. The liquid phase used was distilled water, adjusted to a desired level of pH with sulphuric acid in some instances. Rate constants of dissolution for higher temperatures were extrapolated.
Mathews and Robins reported a first order of reaction in oxygen and a second order of reaction in ferrous ions. According to Mathews and Robins the rate of oxidation of ferrous sulphate decreased
13 with increasing H3O+ concentration, according to a negative order of reaction of -0.25 in H3O+. The
activation energy was determined to be 73.6 kJ/mol.
Iwai et al.17 studied the oxidation of ferrous sulphate in a glass autoclave. Experiments were performed at temperatures between 343 K and 363 K and partial pressures of oxygen ranging from 0.3 MPa to 0.7 MPa. Initial concentrations of H2SO4 and ferrous ions were 0.1 M to 3.0 M and 0.2 M,
respectively. Experiments were conducted at a stirrer speed of 1100 rpm.
Iwai et al. reported a second order in ferrous ions, a first order in oxygen and a first order in sulphate ions, respectively. They assumed that the oxidation of ferrous ions proceeded through two parallel reaction paths. One reaction path is independent on the SO42- concentration; the other reaction path
is dependent on the SO42- concentration. Speciation calculations were performed by taking the
dissociation constant of HSO4- and the formation constants of ferrous sulfate complexes into
account. A power law was suggested for the situation where the H3O+ ion did not affect the oxidation
rate. Activation energies were determined to be 51.6 kJ/mol for the SO42- independent reaction path,
and 144.6 kJ/mol for the SO42- dependent reaction path.
Chmielewski and Charewicz18 investigated the oxidation of ferrous sulphate in an autoclave made of acid-resistant steel with a Teflon coated interior. Experiments were performed at temperatures between 313 K and 408 K, partial pressures of oxygen ranging from 0.132 MPa to 1 MPa and in the presence of Cu2+. Initial concentrations of H2SO4, copper ions and ferrous ions were 0.41 M to 0.82
M, 0.16 M to 1.42 M and 0.54 M, respectively. The effect of stirrer speed was studied. An increase in stirrer speed above 700 rpm did not result in a change in reaction rate, indicating that the reaction was not affected by mass transfer limitations and pure kinetics were measured.
The dependency of the species was determined from Fe2+ oxidation experiments performed with 0.16 M Cu2+. Chmielewski and Charewicz reported a second order of reaction in ferrous ions for concentrations exceeding 0.054 M to 0.14 M and a first order of reaction at lower concentrations. A first order of reaction in oxygen was determined. An enhancing effect of Cu2+ on the Fe2+ oxidation rate was observed. A decrease in Fe2+ oxidation rate was observed with increasing sulphuric acid concentrations. However, this observation was based on experiments performed with a stirrer speed of 120 rpm, and it was mentioned that some experiments were affected by mass transfer limitations. The temperature dependence of the reaction was determined from Fe2+ oxidation experiments performed with 0.79 M Cu2+. The activation energy was determined to be 56.9 kJ/mol.
Verbaan and Crundwell19 studied the oxidation of ferrous sulphate in a stainless steel autoclave. It is not clearly mentioned whether the autoclave was able to withstand the sulphuric acid concentrations at the temperatures investigated. The experimental conditions of three experiments are presented in the paper. These experiments were all performed with 0.25 M H2SO4, a total Fe
concentration of 0.39 M and a stirrer speed of 450 rpm. They did not examine whether the reaction was affected by mass transfer limitations at this stirrer speed. H3O+ concentrations were derived
from redox potential equations. Temperatures investigated were 333 K, 343 K and 353 K. Partial oxygen pressures investigated were 0.1 MPa, 0.3 MPa and 0.45 MPa. It is unclear if more experiments were used to obtain a power law for the oxidation of ferrous sulphate.
Verbaan and Crundwell reported a second order of reaction in ferrous ions, a first order of reaction in oxygen, and a -0.36 order of reaction in H3O+. The paper does not mention how the oxygen
concentration was determined. The activation energy was determined to be 68.6 kJ/mol. Their proposed rate law is only valid in the ferric ion concentration range of 0 M to 0.20 M, because in this concentration range a constant order of reaction in ferrous ions could be determined.
Dreisinger and Peters20 investigated the oxidation of ferrous sulphate for zinc sulphide leaching conditions. Experiments were performed in a Parr titanium autoclave. The effect of the addition of various sulphate salts was studied. Experiments were performed between 120 °C and 160 °C and at oxygen partial pressures of either 0.138 MPa or 0.207 MPa. The effect of H2SO4 on the reaction with
14
constant ZnSO4 concentration was studied with initial concentrations of ZnSO4, H2SO4 and FeSO4 of
2.0 M, 0.25 M to 0.7 M and 0.2 M, respectively. The effect of H2SO4 on the reaction with constant
total sulphate concentration was studied with initial concentrations of H2SO4 and FeSO4 of 0.2 M to
1.0 M and 0.2 M, respectively. The ZnSO4 concentration was varied to obtain a total sulphate
concentration of 2.7 M. The effect of ZnSO4 on the reaction with constant H2SO4 concentration was
studied with initial concentrations of ZnSO4, H2SO4 and FeSO4 of 0 M to 2.0 M, 0.5 M and 0.2 M,
respectively. Experiments were conducted at a stirrer speed of 650 rpm. It was not examined whether the reaction was affected by mass transfer limitations at this stirrer speed.
Dreisinger and Peters assumed that the ion pair FeSO4 is formed in solution, and assumed that the
FeSO4 ion pair in solution is more reactive than the unpaired Fe2+-ion in solution. They stated that
varying the total sulphate content resulted in a variation in FeSO4 ion pair concentration and
consequently in the oxidation rate. Speciation calculations were performed for Zn2+ - Fe2+ - H3O+ -
SO42- solutions. Dreisinger and Peters stated that equilibria involving Cu2+ and Fe3+ were disregarded
due to their limited effect on ferrous speciation. Furthermore, some of the thermodynamic parameters used in speciation calculations were estimated. E.g., the equilibrium constants for dissolved ferrous sulphate and zinc sulphate were not evaluated experimentally at their reaction temperature. Another study extrapolated equilibrium constants to elevated temperatures, and based on these results Dreisinger and Peters assigned values for these equilibrium constants. Activity coefficients for metal sulphate ion pairs were assumed to be 1.
Dreisinger and Peters determined a second order in Fe2+, as well as the temperature dependence of the reaction, for Fe2+ concentrations above 0.02 M. A deviation from second order in Fe2+ was observed for Fe2+ concentrations below 0.02 M. A first order of reaction in oxygen was reported, though it was not clearly mentioned whether this dependency was derived from first principles. Next, it was assumed that the FeSO4 ion pair is formed, and a kinetic equation was reported in which
the dependencies of Fe2+, as well as the FeSO4 ion pair, were assumed to be first and second order.
Their kinetic equation was fitted on experiments performed with a/o ZnSO4 and CuSO4, whereas the
activation energy of 80.3 kJ/mol was derived from experiments performed without CuSO4.
Vračar and Cerović21 studied the oxidation of ferrous ions in a stainless steel autoclave at
temperatures ranging from 50 °C to 200 °C and partial pressures of oxygen ranging from 0.203 MPa to 1.01 MPa. Initial concentrations of H2SO4 and ferrous ions were 0.0 M to 0.51 M and 0.036 M to
0.90 M, respectively. The authors do not clearly mention whether the stainless steel was able to withstand sulphuric acid in solution. Experiments were performed at a stirrer speed of 400 rpm. Vračar and Cerović mentioned that the reaction was affected by mass transfer limitations, and therefore no kinetics were measured.
Vračar and Cerović reported a second order of reaction in Fe2+. A retardation in oxidation rate was observed when experiments performed with 0.51 M H2SO4 were compared to experiments with no
H2SO4 present. The authors do not mention at which oxygen pressure and sulphuric acid
concentration the proposed rate equation is valid. The activation energy was determined to be 51.0 kJ/mol.
Rönnholm et al.22 investigated the oxidation of ferrous sulphate in a Parr autoclave with Teflon protection. Experiments were performed at temperatures between 60 °C and 130 °C and partial pressures of oxygen ranging from 0.4 MPa to 1 MPa. Initial concentrations of H2SO4 and FeSO4 were
1.3 M and 2 to 2.5 M, respectively. Experiments were conducted at a stirrer speed of 1000 rpm. Rönnholm et al. showed that the ferrous sulphate oxidation reaction was slow enough to neglect the influence of the reaction in the liquid film. The influence of solutes on oxygen solubility was accounted for by using the model proposed by Weisenberger and Schumpe.23 They reported a varying order in ferrous ions and a first order of reaction in oxygen. The activation energy was determined to be 34.7 kJ/mol.
15 Table 2.1: Experimental conditions reported in previous studies.
Reference T (K) PO2 (MPa) Composition (M)
FeSO4 H2SO4 McBaina 287.65 0.0101 - 0.101 0.68 0.34 303.15 9.17 x 10-4 M - 1.85 x 10-3 M 0.018 - 0.46 0.01 - 0.5 Lamb and Elder 303.15 0.0213 - 0.101 0.146 - 1.36 0.008 - 3.0
Pound 280.15 - 301.15 0.0213 0.1 5.0 x 10-5 - 6.0
Huffman and Davidsonb,c
413.15 - 453.15 0.0213 and 0.101 0.001 - 0.025 1 303.65 0.0213 and 0.101 0.001 - 0.02 1 McKay and Halpern 373.15 - 403.15 0.101 - 0.405 0.048 - 0.28 0.013 - 0.32 Cornelius and Woodcock 383.15 - 438.15 0.620 - 2.41 0.018 - 0.054 0.054
Hotta and Suzuki 373.15 - 423.15 3.04 0.01 0.05
Mathews and Robinsd 293.15 - 353.15 0.0213 - 0.101 0.01 - 1
Iwai et al. 343 - 363 0.3 - 0.7 0.2 0.1 - 3.0
Chmielewski and Charewicz 313 - 408 0.132 - 1 0.54 0.41 - 0.82
Verbaan and Crundwell 333 - 353 0.1 - 0.45 0.39 0.25
Dreisinger and Peterse 393.15 - 433.15 0.138 - 0.207 0.2 0.2 - 1.0 Vračar and Cerović 323.15 - 473.15 0.203 - 1.01 0.036 - 0.90 0 - 0.51
Rönnholm et al. 333.15 - 400.15 0.4 - 1 2 - 2.5 1.3
a. McBain8 reported oxygen partial pressures as well as oxygen concentrations b. Huffman and Davidson11 used Fe(NH4)2(SO4)2 as a source for Fe2+, not FeSO4
c. Huffman and Davidson11 performed some experiments in the presence of Na2SO4
d. Mathews and Robins16 performed experiments in the pH range of 0 to 2 e. Dreisinger and Peters20 performed experiments in the presence of ZnSO4
Orders of reaction reported in previous studies are summarized in Table 2.2: Table 2.2: Orders of reaction reported in previous studies.
Reference Order of reaction
Fe2+ FeSO4 FeSO4 ion pair O2 H2SO4 SO42- H3O+
McBain 2 1
Lamb and Elder 2 1
Huffman and Davidson 1 and 2 1
McKay and Halpern 2 1 -0.33
Cornelius and Woodcock 2 1
Hotta and Suzuki 1 -
Mathews and Robins 2 1 -0.25
Iwai et al. 2 1 1
Chmielewski and Charewicz 2 1
Verbaan and Crundwell 2 1 -0.36
Dreisinger and Petersa 1 and 2 1 and 2 1
Vračar and Cerović 2 -
Rönnholm et al. 1 - 2 1
a. Dreisinger and Peters20 mentioned that the reported orders are valid for Fe2+ concentrations above 0.02 M
16
Rate laws and activation energies reported in previous studies are summarized in Table 2.3: Table 2.3: Rate laws and activation energies reported in previous studies.
Reference Rate law EA (kJ/mol)
EA,1 EA,2
Huffman and Davidson
[ ]
[ ] [ ] 56.1 68.2
McKay and Halpern [ ] [ ] 69.0 -
Cornelius and Woodcock [
]
[
]
61.9 -
Hotta and Suzuki [
]
[
] near
56,9 -
Mathews and Robins [
] [ ] [ ] [ ] 73.6 - Iwai et al. [ ] [ ] [ ] [ ] 51.6 144.6 Chmielewski and Charewicz [
]
[ ] 56.9 -
Verbaan and Crundwell [ ]
[
] [ ][ ]
68.6 -
Dreisinger and Peters
[ ]
( [ ] [ ][ ]
[ ] )( [ ] ) 80.3
Vračar and Cerović [
] [ ] 51 - Rönnholm et al. [ ] [ ] [ ] ( [ ]) 34.7 -
2.3 Experimental
2.3.1 Setup
The Fe2+ oxidation experiments were performed in a 0.5 l glass Büchi autoclave. The reactor was operated batchwise with regard to the liquid phase, and continuous with regard to the gas phase. Nitrogen could be fed to remove oxygen from the setup. During the experiment, air was supplied to the reactor via a mass flow controller. A water saturator was used to humidify the gas phase and maintain the water balance. The pressure in the reactor was controlled with a back-pressure controller. A pressure relief valve was added to the reactor as a safety precaution. The reactor temperature was regulated by a Julabo heater. A Büchi water bath heated the water saturator to the reactor temperature. A Teflon gas entrainment stirrer was used to obtain maximum gas dispersion in the liquid phase. The stirrer was operated with a magnetic drive. Glass baffles were introduced in the reactor to promote mixing. The steel interior of the reactor was protected by Teflon. Sampling was performed through Teflon tubing and Teflon valves; 3.5 ml UV cuvettes were used to collect samples. A Varian Cary 50 UV-Vis spectrophotometer (EL98123254) was used to analyze sample compositions. Samples were stored in a refrigerator. The temperature in the reactor was monitored with a Teflon-protected PT100 thermometer, the temperature in the water saturator was registered with a regular PT100 thermometer. The pressure inside the experimental setup was measured with a pressure transducer. The thermometers and pressure transducer were connected to a data acquisition computer. A schematic diagram of the experimental setup is shown in Figure 2.1.
17 Figure 2.1: Schematic diagram of the experimental setup.
2.3.2 Chemicals
ACS reagent grade FeSO4.7H2O [7782-63-0], 97 % and puriss p.a. grade Fe2(SO4)3.xH2O [15244-10-7]
and ACS reagent grade H2SO4 [7664-93-9] were used as supplied from Sigma-Aldrich. Demineralized
water was used.
2.3.3 Procedure
An amount of 400 ml of solution was prepared and added to the reactor. The reactor was flushed three times with N2 to remove oxygen from the setup. When the setup reached the required
temperature, the reactor was flushed two times with air. At the required air pressure the first sample was taken. The experiment was initiated by starting the stirrer. Sampling was performed periodically, samples were directly stored in a refrigerator after sampling and analyzed at the end of each experiment at room temperature. From each experiment the overall mass balance was determined.
2.3.4 Analytical techniques
Pure ferrous sulphate and ferric sulphate solutions, in the presence of sulphuric acid, absorb light at different wavelengths. A wavelength scan from 200 nm to 1100 nm was performed with the UV Vis spectrophotometer on a pure acidic ferrous sulphate solution and a pure acidic ferric sulphate solution. The scan indicated that at a wavelength of 410 to 430 nm, Fe2(SO4)3 had a significant
absorbance, whereas FeSO4 had almost no absorbance. A wavelength of 420 nm was chosen to
analyze the sample compositions. At this wavelength, the absorbance of Fe2(SO4)3 was 140 times
higher compared to the absorbance of FeSO4. As Fe2(SO4)3 is in low amounts present at the start of
an experiment, this is considered as an accurate method of determining compositions.
Calibration curves were made for each experiment, depending on the starting composition and resulting time dependent concentration profile of the experiment.
18
2.3.5 Consideration
To ensure a sufficiently large contact area between gas and liquid phase and high mass transfer coefficients, a gas entrainment stirrer was used. Experiments at two stirrer speed were performed to determine if the reaction was affected by mass transfer limitations (1200 rpm and 1700 rpm). No significant difference in conversion rates were observed, therefore it was concluded that the reaction was not affected by mass transfer limitations. All experiments were carried out at a stirrer speed of 1700 rpm. The supply of oxygen per time unit was at least five times higher than the amount consumed during the experiment.
Experiments were performed at temperatures of 50 °C, 70 °C and 90 °C and at air pressures ranging from 0.1 MPa to 0.5 MPa. Some experiments were performed using enriched air with composition 60 vol.% N2 and 40 vol.% O2 to be able to increase the oxygen partial pressure. Initial concentrations of
H2SO4 and FeSO4 were 0.25 M to 2.0 M and 0.125 M to 1.0 M, respectively.
The accuracy of the PT100 thermometers are temperature dependent. At temperatures of 50 °C, 70 °C and 90 °C the accuracies are ± 0.25 °C, ± 0.29 °C and ± 0.33 °C, respectively. The accuracy of the pressure transducer is ± 0.005 MPa. The photometric and wavelength accuracies of the UV-Vis spectrophotometer are ± 0.01 Abs and ± 0.5 nm, respectively.
The reproducibility of the experiments in this study is estimated to be within 5 %.
2.4 Results
2.4.1 Order of reaction
The order of reaction in a reactant, in a system with several reactants, can be determined through the so-called pseudo-first order approach. Generally, this implies that the concentration of one of the reactants is small, and the concentration(s) of the other reactant(s) in excess and hardly change(s) during the experiment. Therefore, the reactant(s) in excess are considered to remain constant during the reaction, and can be lumped together within the pseudo first order reaction rate constant. In the present study the pseudo-first order approach is used, but in a different way. By varying the initial concentration of one of the reactants in several experiments, the order in this reactant can be determined from the initial reaction rates of the experiments. The reactants with equal initial concentrations are lumped together with the reaction rate constant. Therefore it is not necessary to have reactants present in excess.
As an example: reactants x, y and z participate in a reaction according to the following power law rate equation:
[ ] [ ] [ ] (2.5)
Only the initial concentration of reactant x is varied in different experiments, therefore reactants y and z can be lumped together with reaction rate constant k, forming the pseudo-first order reactant rate constant kps:
[ ] (2.6)
Equation 2.6 can be linearized by applying the natural logarithm:
[ ] (2.7)
19 The initial reaction rates of experiments are determined by fitting the data points via a polynomial equation. From the slope at t = 0 (calculated from the derivative of the polynomial fit) the initial reaction rate is derived. An example of the fitting of the experimental data is shown in Figure 2.2.
Figure 2.2: [Fe2+] versus time with polynomial fit. Initial conditions: T = 70 °C, Pair = 0.5 MPa, [FeSO4] =
1.0 M, [H2SO4] = 1.0 M. ◊ experimental data points, — polynomial fit.
2.4.2 Order of reaction in Fe
2+The order of reaction in Fe2+ was determined at three different temperatures, i.e. T = 90 °C, 70 °C and 50 °C. The experimental data for the initial reaction rate are given in Figure 2.3. Based on the curve fittings in Figure 2.3 it can be concluded that slopes varied between 2.08 and 2.21. This indicates that within the experimental accuracy for conditions studied, a second-order dependency in Fe2+ is observed.
2.4.3 Order of reaction in O2
The order of reaction in the partial pressure of O2 was determined at three different temperatures; T
= 90 °C, 70 °C and 50 °C. The actual oxygen partial pressure was determined via correction of the total pressure for water vapor pressure. The saturated water vapor pressure was calculated according to the Antoine equation.24
( ( ))
(2.8) Where A = 8.07131, B = 1730.63 and C = 233.426 in the range of 274.15 K < T < 372.15 K.
The experimental data for the initial reaction rate are given in Figure 2.4. Based on the curve fittings in Figure 2.4 it can be concluded that slopes varied between 0.94 and 1.01. This indicates that within the experimental accuracy for conditions studied, a first-order dependency in O2 is observed.
0.84 0.86 0.88 0.90 0.92 0.94 0.96 0.98 1.00 0 500 1000 1500 2000 2500 3000 [Fe 2+] (M ) t (s)
20
Figure 2.3: lnRFe2+, t = 0 versus ln[Fe2+]. Initial conditions: T = 90, 70 & 50 °C, Pair = 0.5 MPa, [H2SO4] = 1.0
M. Initial Fe2+ concentrations at T = 90 °C: 0.125, 0.25, 0.5 and 1.0 M. Initial Fe2+ concentrations at T = 70 & 50 °C: 0.25, 0.5 and 1.0 M.
Figure 2.4: lnRFe2+, t = 0 versus lnPO2. Initial conditions: T = 90 °C: [Fe2+] = 0.5 M, [H2SO4] = 0.5 M, Pair =
0.124, 0.178, 0.285 and 0.5 MPa. T = 70 °C: [Fe2+] = 0.5 M, [H2SO4] = 0.5 M, Pair = 0.139, 0.246 and
0.461 MPa. T = 50 °C: [Fe2+] = 1.0 M, [H2SO4] = 1.0 M, Pair = 0.12, 0.228, 0.443 and 0.463 (40 vol.% O2)
MPa.
2.4.4 Order of reaction in H2SO4
The order of reaction in H2SO4 was determined at three different temperatures; T = 90, 70 and 50 °C.
The experimental data for the initial reaction rate are given in Figure 2.5. 2.17 2.21 2.08 -15 -14 -13 -12 -11 -10 -9 -8 -2.5 -2 -1.5 -1 -0.5 0 0.5
ln
R
Fe 2+ln[Fe
2+]
T = 90 °C T = 70 °C T = 50 °C 0.99 1.01 0.94 -13 -12.5 -12 -11.5 -11 -10.5 -10 -9.5 -9 9 9.5 10 10.5 11 11.5 12 12.5ln
R
Fe 2+lnP
O2 T = 90 °C T = 70 °C T = 50 °C21 Figure 2.5: lnRFe2+, t = 0 versus ln[H2SO4]. Initial conditions: T = 90, 70 & 50 °C, Pair = 0.5 MPa, [H2SO4] =
0.25, 0.5, 1.0 and 2.0 M. Initial Fe2+ concentrations at T = 90 & 70 °C, [Fe2+] = 0.25 M; at T = 50 °C, [Fe2+] = 0.5 M.
A uniform order of reaction in H2SO4 was not observed. The oxidation rate of Fe2+ decreases with
increasing H2SO4 concentration up to 1.0 M H2SO4. At concentrations above 1.0 M H2SO4 it seems
justified to conclude that there is no effect on the reaction rate and at concentrations below a negative order is observed of about -0.62 to -0.72 (average -2/3). Chmielewski and Charewicz18 concluded that Bielopolskij and Urosov observed a similar phenomenon. They too found that the Fe2+
oxidation rate decreased with increasing H2SO4 concentration up to 1.0 M H2SO4, and that this effect
was not observed at concentrations above 1.0 M H2SO4. Iwai et al.17 reported also a decrease in Fe2+
oxidation rate with increasing H2SO4 concentration up to 1.0 M H2SO4. Their Figure 1 shows
experiments performed with varying H2SO4 concentrations. Experiments performed with 1.0 M, 2.0
M and 3.0 M H2SO4 resulted in a constant reaction rate. This confirms the present observation that
above H2SO4 concentrations of 1.0 M no effect is observed on the reaction rate. Furthermore, Iwai et
al. reported a negative order of -0.6 in H2SO4 at concentrations below 1.0 H2SO4 (their Figure 4). This
agrees with the negative order observed in the present study of about -2/3 for H2SO4 concentrations
below 1.0 M.
2.4.5 Order of reaction in HSO4
-, H3O
+and SO4
2-The order of reaction in H2SO4 is not uniform and has a discrete change at a concentration around 1.0
M. Therefore, as H2SO4 dissociates in water into HSO4-, H3O+ and SO42-, it might be possible that any
of these species is affecting the reaction mechanism.
The dissociation constants of H2SO4 are given in Equations 2.9 and 2.10:
(2.9)
(2.10)
The first dissociation step of aqueous H2SO4 can be regarded to be complete for concentrations up to
approximately 14 M in the temperature range of 273 K to 323 K.25 Moreover, in computational -0.72 -0.63 -0.62 -13 -12.5 -12 -11.5 -11 -10.5 -10 -2 -1.5 -1 -0.5 0 0.5 1
ln
R
Fe 2+ln[H
2SO
4]
T = 90 °C T = 70 °C T = 50 °C22
studies of aqueous sulphuric acid, the first dissociation has been assumed to be complete.26,27,28 Consequently, in this study the first dissociation of sulphuric acid was assumed to be 100 %.
The dissociation of the bisulphate ion (HSO4-) in water depends strongly on the temperature.25,26,29
Knopf et al.29 described the second dissociation constant as a function of temperature: ( ) - [( - ) ( - ) - ( ) ( - )] (2.11)
Where K2o = 1.0576 x 10-2, ΔH2o/R = -2231.620793, cpo/R= -24.7273 and (dcp/dT)/R = -0.11967 in the
temperature range of 180 K to 473 K.
The relation presented by Knopf et al. for the second dissociation constant K2, together with the
assumption of complete first dissociation of sulphuric acid to bisulphate, was used to determine the speciation of the initial solution. The initial amount of water per volume of solution could be determined from calibration mixtures. This parameter was used to convert between molalities and molarities.
Figures 2.6, 2.7 and 2.8 show the experimental data to determine the orders in HSO4-, H3O+ and SO42-,
respectively.
Figure 2.6: lnRFe2+, t = 0 versus ln[HSO4-]. Initial conditions: T = 90, 70 & 50 °C, Pair = 0.5 MPa, [H2SO4] =
0.25, 0.5, 1.0 and 2.0 M. Initial Fe2+ concentrations at T = 90 & 70 °C, 0.25 M; at T = 50 °C, 0.5 M. As it can be seen from Figures 2.6, 2.7 and 2.8, the orders of reaction in HSO4-, H3O+ and SO42- vary
with temperature up to a concentration of 1.0 M H2SO4, with coefficients of determination of R2 <
0.99. Negative orders of reaction in HSO4- and H3O+, between -0.76 and -1.01 and between -0.21 and
-0.30, respectively, are obtained for H2SO4 concentrations below 1.0 M. Positive orders of reaction in
SO42-, between 0.30 and 0.42, are obtained for H2SO4 concentrations below 1.0 M.
-1.01 -0.82 -0.76 -13 -12.5 -12 -11.5 -11 -10.5 -10 -1 -0.5 0 0.5 1
ln
R
Fe 2+ln[HSO
4-]
T = 90 °C T = 70 °C T = 50 °C23 Figure 2.7: lnRFe2+, t = 0 versus ln[H3O+]. Initial conditions: T = 90, 70 & 50 °C, Pair = 0.5 MPa, [H2SO4] =
0.25, 0.5, 1.0 and 2.0 M. Initial Fe2+ concentrations at T = 90 & 70 °C, 0.25 M; at T = 50 °C, 0.5 M.
Figure 2.8: lnRFe2+, t = 0 versus ln[SO42-]. Initial conditions: T = 90, 70 & 50 °C, Pair = 0.5 MPa, [H2SO4] =
0.25, 0.5, 1.0 and 2.0 M. Initial Fe2+ concentrations at T = 90 & 70 °C, 0.25 M; at T = 50 °C, 0.5 M. The observation that the reaction rate does not change for H2SO4 concentrations of 1.0 M and above
could be explained by approximately no net change in concentration of one of the species. From Figures 2.6 to 2.8 it must be concluded that the concentration of SO42- does not change significantly
with an increase in H2SO4 concentration above 1.0 M. Therefore the oxidation of ferrous ions in
acidic sulphate solutions could be related to SO42-, though the orders in SO42- vary with temperature
and do not have a significant coefficient of determination. -0.30 -0.27 -0.21 -13 -12.5 -12 -11.5 -11 -10.5 -10 -5 -4 -3 -2 -1 0 1
ln
R
Fe 2+ln[H
3O
+]
T = 90 °C T = 70 °C T = 50 °C 0.42 0.39 0.30 -13 -12.5 -12 -11.5 -11 -10.5 -10 -7 -6 -5 -4 -3 -2 -1 0ln
R
Fe 2+ln[SO
42-]
T = 90 °C T = 70 °C T = 50 °C24
A possible route to reveal the influence of H2SO4 on the reaction rate and mechanism, respectively, is
to interpret the obtained results in terms of activities or chemical potentials; however, then a thorough and extended thermodynamic model is required.
2.5 Kinetic equation
A power law kinetic equation is determined from experiments performed with an initial sulphuric acid concentration of 1.0 M and higher, because at these conditions the H2SO4 concentration seems
to have no influence on the observed reaction rates. Moreover, for the process under development, the used solvent compositions usually contain H2SO4 at concentrations above 1.0 M.
2.5.1 Power law kinetic equation
The overall reaction rate equation for the oxidation of ferrous ions in acidic sulphate solutions with H2SO4 concentrations above 1.0 M is
[ ]
[
] (2.12)
Table 2.4 summarizes the orders in Fe2+ and O2.
Table 2.4: Orders in Fe2+ and O2 at varying temperatures.
Component
Temperature (°C)
50 70 90
Fe2+ 2.08 2.21 2.17
O2 0.94 1.01 0.99
The order of reaction in a component is averaged from the orders of reaction determined at varying temperatures. From Table 2.4 it can be concluded that the order of reaction in Fe2+ becomes 2 and the order of reaction in O2 becomes 1. As mentioned before, at H2SO4 concentrations above 1.0 M,
the orders of reaction in H2SO4, HSO4-, H3O+ and SO42-, respectively, can be regarded to be equal to
zero.
Reaction rate constants at the varying temperatures were calculated using Equation 2.13. For fitting the reaction rate constants, 9 experimental data points at T = 90 °C, 4 experimental data points at T = 70 °C and 8 experimental data points at T = 50 °C were used. Table 2.5 summarizes the reaction rate constants.
Table 2.5: reaction rate constants at varying temperatures. Temperature (°C) k (m3.kmol-1.Pa-1.s-1)
50 1.94 x 10-10
70 7.15 x 10-10
90 2.30 x 10-9
By using the Arrhenius equation, the temperature dependence of the reaction rate constant was determined:
{ } (2.13)
Resulting in a value of the activation energy of EA = 60.3 kJ/mol (see Figure 2.9).
25 Figure 2.9: ln(k) versus T-1.
2.5.2 Parity plot
To evaluate the derived kinetic expression according to Equations 2.12 and 2.14, the
experimentally obtained initial reaction rates are compared to the rates calculated. A parity
plot of the comparison is shown in Figure 2.10. Experimental reaction rates varied about three
decades (from -9 x 10-7 kmol.m-3.s-1 to -4 x 10-4 kmol.m-3.s-1).The parity plot shows good agreement between experimental reaction rates and reaction rates calculated according to the postulated kinetic equation. The maximum deviation between the experimental and calculated reaction rate is 12.8 %, the average deviation is 4.9 %.
2.6 Discussion
2.6.1 Previous studies
Huffman and Davidson11 stated that the partial pressure of oxygen was approximately constant during the reaction, due to a sufficiently large gas hold-up above the solution. From the dimensions of the Pyrex glass tube reactor, as reported by Huffman and Davidson, it can be concluded that the glass reactor had a total volume of 16.3 ml, of which 5 ml was liquid volume. Using the ideal gas law, the worst case scenario for 0.025 M Fe2+ and air can be determined. For complete conversion of Fe2+, approximately 32 % of the oxygen is required. Therefore it seems unlikely that the partial pressure of oxygen remained constant for this particular experiment. Consequently, this could be an explanation for their inability to report a clear first order of reaction in oxygen. Moreover, the use of these experiments in the determination of the kinetic equation could have resulted in an error in some of the kinetic parameters.
McKay and Halpern13 reported an order of -0.33 in H2SO4. However, upon study of the master thesis
of McKay12 it can be concluded that the order was determined from pH measurements, and consequently is an order in the component H3O+.
26
Figure 2.10: Parity plot of experimental and calculated reaction rates.
Hotta and Suzuki15 claimed that the activation energy of the Fe2+ oxidation reaction was near 56.1 kJ/mol. They gave the reaction rate constants at three temperatures, i.e. 1.24 x 10-4 s-1, 5.17 x 10-4 s-1
and 6.10 x 10-4 s-1 at 100 °C, 130 °C and 150 °C, respectively. A linear fit of the natural logarithm of
these rate constants against 1/T results in an activation energy of 43.7 kJ/mol with a coefficient of determination R2 of 0.92.
Chmielewski and Charewicz18 determined a kinetic equation based on experiments performed with initial Cu2+ concentrations of either 0.16 M or 0.79 M, respectively. As Chmielewski and Charewicz explained, the rate of oxidation is affected by the concentration of Cu2+. Therefore it was incorrect of them to report a kinetic equation that did not account for the concentration of Cu2+.
The proposed rate law by Verbaan and Crundwell19 is only valid in the ferric ion concentration range of 0 M to 0.20 M, because in this concentration range a constant order of reaction in ferrous ions could be determined. This concentration range applies for experiments performed at temperatures of 343 K and 353 K. However, the experiment performed at a temperature of 333 K showed a second order of reaction over the whole concentration range investigated (their Figure 5). This deviation from second order of reaction in Fe2+ at higher temperatures could occur due to the reaction being diffusion controlled. It is questionable whether pure kinetics were measured at their applied stirrer speed of 450 rpm.
The first-order dependency in oxygen, reported by Verbaan and Crundwell, does not seem to be determined from first principles.
Vračar and Cerović21 mentioned that the reaction was diffusion controlled at their applied stirrer speed of 400 rpm. Therefore their rate law can only be used for their setup and experimental conditions, and is not representative for pure kinetics of the oxidation of ferrous sulphate.
Rönnholm et al.22 accounted for the influence of solutes on oxygen solubility by using the model proposed by Weisenberger and Schumpe.23 Because Weisenberger and Schumpe do not provide a model parameter for HSO4-, Rönnholm et al. estimated the change in oxygen solubility based solely
8.E-07 8.E-06 8.E-05 8.E-04
8.E-07 8.E-06 8.E-05 8.E-04
R
cal cR
exp T = 90 °C T = 70 °C T = 50 °C +10% -10%27 on H3O+, SO42-, Fe2+ and Fe3+. It is questionable whether reasonably correct oxygen concentrations
were obtained, as HSO4- is present in the experimental solution in adequate amounts.
2.6.2 Orders of reaction
The second-order dependency in Fe2+ determined in this work agrees well with the order in either FeSO48,9,13 or Fe2+ 14,16,17,18,19,21 reported in most of the previous studies (Table 2.2). Huffman and
Davidson11 reported a first, as well as a second order in dependency in Fe2+. Dreisinger and Peters20 reported a first and a second-order dependency in Fe2+, as well as a first and a second-order dependency in the FeSO4 ion pair complex. Rönnholm et al.22 reported a varying order in Fe2+
(between 1 and 2), as a result of their proposed reaction mechanism. One study reported a first-order dependency in Fe2+.15
The first-order dependency in O2 determined in this work agrees with the order reported in the
previous studies (Table 2.2).
The kinetic equation derived in this study does not incorporate H2SO4, or any of the species H2SO4
dissociates into, because it is valid for H2SO4 concentrations above 1.0 M. At these concentrations
H2SO4 does not influence the Fe2+ oxidation rate. However, an order in H2SO4 between 0.62 and
-0.72 is derived in this study (Figure 2.5) that agrees reasonably with the order of -0.6 reported by Iwai et al.17 for H2SO4 concentrations below 1.0 M.
The order in H3O+ (Figure 2.7) agrees reasonably with the order reported in previous studies for
H2SO4 concentrations below 1.0 M.12,13,16,19 The order in H3O+ reported in this study is between -0.21
and -0.30. The orders reported by Mckay and Halpern, Mathews and Robins and Verbaan and Crundwell were -0.33, -0.25 and -0.36, respectively. The orders derived by McKay and Halpern and Mathews and Robins are based on pH measurements performed at ambient conditions. The order derived by Verbaan and Crundwell is based on redox potential measurements.
2.6.3 Activation energy
The activation energy of 60.3 kJ/mol, derived in this study, agrees well with the activation energy of 61.9 kJ/mol reported by Cornelius and Woodcock14 (Table 2.3).
The activation energies reported by Huffman and Davidson,11 McKay and Halpern,13 Hotta and Suzuki,15 Chmieliewski and Charewicz18 and Verbaan and Crundwell,19 as well as the first activation energy reported by Iwai et al.,17 are within a range of 10 kJ/mol with the activation energy reported in this study.
The activation energies reported by Mathews and Robins,16 Dreisinger and Peters20 and Rönnholm et al.22 do not agree well with the activation energy reported in this study (see Table 2.3).
2.6.4 Reaction rate
The derived kinetic equation in this study is compared to kinetic equations in previous studies as summarized in Table 2.3 by calculating the Fe2+ oxidation rate for one unique experimental setting. The kinetic equations reported by Huffman and Davidson,11 Hotta and Suzuki,15 Iwai et al.,17 Chmielewski and Charewicz,18 Dreisinger and Peters20 and Vračar and Cerović21 will not be compared. Huffman and Davidson used ferrous ammonium sulphate as a source for ferrous ions, therefore ammonium ions were present in the experimental solutions. Hotta and Suzuki did not incorporate oxygen in their kinetic equation as they performed experiments at one partial pressure of oxygen of 3.04 MPa. Iwai et al. did not provide clearly the pre-exponential factors in their kinetic equation. Chmielewski and Charewicz derived a kinetic equation based on experiments performed with Cu2+, which acts as a catalyst. Dreisinger and Peters incorporated the FeSO4 ion pair in their kinetic
28
mentioned that their experiments were affected by mass transfer limitations. Consequently, they did not determine kinetics.
The remaining kinetic equations were derived from experiments with different experimental conditions compared to this study. Only Rönnholm et al.22 reported a kinetic equation based on experiments performed with H2SO4 concentrations of 1.0 M and/or above.
The oxygen concentration, reported in the kinetic equations by Mathews and Robins,16 Verbaan and Crundwell19 and Rönnholm et al.,22 was estimated with the actual oxygen partial pressure via Henry’s law for oxygen solubility in pure water using the Procede Process Simulator.30 The influence of solutes on the oxygen solubility was not accounted for.
(2.15)
{ ( ) } (2.16)
Where A = 155.92, B = -7775, C = -18.397 and D = -0.0094435 in the range of 273 K < T < 600 K
The H3O+ concentrations, reported in the kinetic equation by Mathews and Robins and Verbaan and
Crundwell, were determined from pH measurements and redox potential measurements, respectively. The H3O+ concentration was estimated from the dissociation constants of H2SO4 as
described in the present study.
Table 2.6 reports the calculated reaction rates for the experimental setting: T = 90 °C, PO2 = 0.1 MPa,
initial concentrations of FeSO4 and H2SO4 of 0.25 M and 1.0 M, respectively. The O2 concentration
was estimated to be 7.94 x 10-4 M.
Deviations between the kinetic relations in Table 2.6 may occur as it was not possible to derive for each study the required information on the values that were used for e.g. the oxygen concentrations or the pH. Therefore the relations were used as presented and compared to each other to have a first overview on the outcome of the various reaction rates.
Table 2.6: Reaction rates calculated with kinetic equations from previous studies and this study. Experimental setting: T = 90 °C, PO2 = 0.1 MPa, concentrations of FeSO4 and H2SO4 of 0.25 M and 1.0
M, respectively.
Reference R (kmol.m-3.s-1) McKay and Halpern -7.19 x 10-6 Cornelius and Woodcock -2.40 x 10-5 Mathews and Robins -1.39 x 10-5 Verbaan and Crundwell -1.52 x 10-5 Rönnholm et al. -1.55 x 10-10
This study -8.98 x 10-6
McKay and Halpern,13 Cornelius and Woodcock14 and Rönnholm et al.22 reported kinetic equations that are not a function of H2SO4, or any of the species H2SO4 dissociates into. Consequently, the
calculated reaction rate is not determined as a function of H2SO4.
Except for Rönnholm et al., all the reaction rates reported in Table 2.6 are based on kinetic equations derived from experiments below a concentration of 1.0 M H2SO4. The reaction rates derived from
Cornelius and Woodcock, Mathews and Robins and Verbaan and Crundwell are between a factor of 1.55 to 2.67 times higher compared to the reaction rate derived in this study. The reaction rate derived from McKay and Halpern is a factor of 1.25 times lower, respectively. Except for McKay and
29 Halpern, the reaction rates derived are in agreement with the trend that a decrease in H2SO4
concentration below 1.0 M increases the reaction rate.
The reaction rate derived from the kinetic equation reported by Rönnholm et al. is more than a factor 103 lower compared to the other reaction rates observed by the other authors in Table 2.6. Their kinetic equation was fitted on experiments with initial FeSO4 concentrations varying from 2 M
to 2.5 M. The initial FeSO4 concentration of 0.25 M used in deriving the reaction rates in Table 2.6 is
significantly lower; therefore the calculated reaction rate is an extreme extrapolation.
2.7 Conclusion
The oxidation of ferrous ions in acidic sulphate solutions at elevated air pressures was investigated. The effect of the Fe2+ concentration, initial H2SO4 concentration and partial oxygen pressure on the
reaction rate were determined at three different temperatures, i.e. T = 90 °C, 70 °C and 50 °C. From experiments performed with varying stirrer speed it was concluded that the reaction was not affected by mass transfer limitations.
A second order of reaction in Fe2+ and a first order of reaction in O
2 were determined. The oxidation
rate of Fe2+ decreased with increasing H
2SO4 concentration up to 1.0 M H2SO4. At concentrations
above 1.0 M H2SO4 it was concluded that H2SO4 did not affect the reaction rate, hence a zero order of
reaction in H2SO4.
The possibility existed that one of the species H2SO4 dissociates into, i.e. HSO4-, H3O+ and SO42-, was
affecting the reaction mechanism. Speciation calculations, based on the dissociation constants of H2SO4, were performed to determine the concentrations of HSO4-, H3O+ and SO42-. No clear order in
any of the components H2SO4 dissociates into could be established. However, the observation was
made that an increase in H2SO4 concentration up to 1 M resulted in a decrease in SO4
2-concentration, whereas an increase in H2SO4 concentration above 1 M resulted in a minimal change
in SO42- concentration. Therefore the oxidation of ferrous ions in acidic sulphate solutions could be
related to SO42-.
A power law kinetic equation for the oxidation of Fe2+ at concentrations of H2SO4 of 1.0 M and higher
was postulated:
[ ]
[ ] Where the activation energy was determined to be EA = 60.3 kJ/mol.
2.8 Nomenclature
A, B, C, D constants
cP heat capacity [J.mol-1.K-1]
EA activation energy [kJ/mol]
H enthalpy [J/mol]
k reaction rate constant [m3.kmol-1.Pa-1.s-1]
K dissociation constant [-]
KH Henry’s law [Pa]
m molality [mol/kg H2O]
P pressure [Pa]
R reaction rate [kmol.m-3.s-1]
R universal gas constant [J.mol-1.K-1]
30
T temperature [˚C, K]
x mole fraction [-]
Subscripts and superscripts
a, b, c orders x, y, z components ps pseudo-first order Fe2+ ferrous ion O2 oxygen H3O+ hydronium ion
HSO4- bisulphate ion
H2SO4 sulphuric acid
SO42- sulphate ion
2.9 References
[1] Mcintyre, G. and Lyddon, L. Claus sulphur recovery options. Pet. Technol. Q. Spring 1997, 57─61. [2] Perry, D., Fedich R.B. and Parks, L.E. Better acid gas enrichment. Sulphur 2010, 326, 38─42. [3] Versteeg, G.F. and Ter Maat, H. Method and system for selective removal of contamination from
gas flows. WO patent 1998055209 A1, assigned to Procede Twente B.V., priority date June 2, 1997.
[4] Ter Maat, H., Hogendoorn, J.A. and Versteeg, G.F. The removal of hydrogen sulfide from gas streams using an aqueous metal sulfate absorbent. Part I. The absorption of hydrogen sulfide in metal sulfate solutions. Sep. Purif. Technol. 2005, 43 (3), 183─197.
[5] Ter Maat, H., Al-Tarazi, M., Hogendoorn, J.A., Niederer, J.P.M. and Versteeg, G.F. Theoretical and experimental study of the absorption rate of H2S in CuSO4 solutions. The effect of enhancement
of mass transfer by a precipitation reaction. Chem. Eng. Res. Des. 2007, 85 (1), 100─108.
[6] Peacey, J., Guo, X.-J. and Robles, E. Copper hydrometallurgy – current status, preliminary economics, future direction and positioning versus smelting. Trans. Nonferrous Met. Soc. China
2004, 14 (3), 560─568.
[7] Dutrizac, J.E. and MacDonald, R.J.C. The kinetics of dissolution of covellite in acidified ferric sulphate solutions. Can. Metall. Q. 1974, 13 (3), 423─433.
[8] McBain, J.W. Oxidation of ferrous solutions by free oxygen. J. Phys. Chem. 1901, 5, 623─638. [9] Lamb, A.B. and Elder, L.W. The electromotive activation of oxygen. J. Am. Chem. Soc. 1931, 53,
137─166.
[10] Pound, J.R. The oxidation of solutions of ferrous salts. J. Phys. Chem. 1939, 43, 955─967.
[11] Huffman, R.E. and Davidson, N. Kinetics of the ferrous iron-oxygen reaction in sulfuric acid solution. J. Am. Chem. Soc. 1956, 78, 4836─4842.
[12] McKay, D.R. A kinetic study of the oxidation of pyrite in aqueous suspension. Thesis, University of British Columbia, Vancouver, 1957.
[13] McKay, D.R. and Halpern, J. A kinetic study of the oxidation of pyrite in aqueous suspension. Trans. Metall. Soc. AIME 1958, 212, 301─309.
[14] Cornelius, R.J. and Woodcock, J.T. Pressure leaching of a manganese ore. Part I: kinetic aspects. Proc. Aust. Inst. Min. Met. 1958, 185, 65─107.
[15] Hotta, H. and Suzuki, N. Thermal and radiation oxidations of benzene to phenol in aqueous solutions containing metal ions at elevated temperatures. I. The thermal behavior of iron sulfate and chloride in aqueous solutions without benzene. Bull. Chem. Soc. Jpn. 1963, 36, 717─721. [16] Mathews, C.T. and Robins, R.G. The oxidation of aqueous ferrous sulphate solutions by molecular
oxygen. Proc. Aust. Inst. Min. Met. 1972, 242, 47─57.
[17] Iwai, M., Majima, H. and Awakura, Y. Oxidation of Fe(II) in sulfuric acid solutions with dissolved molecular oxygen. Metall. Trans. B 1982, 13B, 311─318.
31 [18] Chmielewski, T. and Charewicz, W.A. The oxidation of Fe(II) in aqueous sulphuric acid under
oxygen pressure. Hydrometallurgy 1984, 12, 21─30.
[19] Verbaan, B. and Crundwell, F.K. An electrochemical model for the leaching of a sphalerite concentrate. Hydrometallurgy 1986, 16, 345─359.
[20] Dreisinger, D.B. and Peters, E. The oxidation of ferrous sulphate by molecular oxygen under zinc pressure-leach conditions. Hydrometallurgy 1989, 22, 101─119.
[21] Vračar, R.Ž. and Cerović, K.P. Kinetics of oxidation of Fe(II) ions by gaseous oxygen at high temperatures in an autoclave. Hydrometallurgy 1997, 44, 113─124.
[22] Rönnholm, M.R., Wärna, J., Salmi, T., Turunen, I. and Luoma, M. Kinetics of oxidation of ferrous sulfate with molecular oxygen. Chem. Eng. Sci. 1999, 54, 4223─4232.
[23] Weisenberger, S. and Schumpe, A. Estimation of gas solubilities in salt solutions at temperatures from 273 K to 363 K. AIChE J. 1996, 42 (1), 298─300.
[24] Dortmund Data Bank. http://www.ddbst.com/ddb.html (accessed 10-08-2016).
[25] Young, T.F., Maranville, L.F. and Smith, H.M. Raman spectral investigations of ionic equilibria in solutions of strong electrolytes. The structure of Electrolytic Solutions; W. J. Hamer, ed., Wiley: New York, 1959.
[26] Clegg, S.L., Rard, J.A. and Pitzer, K.S. Thermodynamic properties of 0-6 mol kg-1 aqueous sulfuric acid from 273.15 to 328.15 K. J. Chem. Soc., Faraday Trans. 1994, 90 (13), 1875─1894.
[27] Clegg, S.L. and Brimblecombe, P. Application of a multicomponent thermodynamic model to activities and thermal properties of 0-40 mol kg-1 aqueous sulfuric acid from <200 to 328 K. J. Chem. Eng. Data 1995, 40, 43─64.
[28] Cherif, M., Mgaidi, A., Ammar, M.N., Abderrabba, M. and Fürst, W. Representation of VLE and liquid phase composition with an electrolyte model: application to H3PO4-H2O and H2SO4-H2O.
Fluid Phase Equilib. 2002, 194–197, 729─738.
[29] Knopf, D.A., Luo, B.P., Krieger, U.K. and Koop, T. Thermodynamic dissociation constant of the bisulfate ion from Raman and ion interaction modeling studies of aqueous sulfuric acid at low temperatures. J. Phys. Chem. A 2003, 107, 4322─4332.
[30] Hamborg, E.S, Derks, P.W.J., Elk, E.P.V. and Versteeg, G.F. Carbon dioxide removal by alkanolamines in aqueous organic solvents. A method for enhancing the desorption process. Energy Procedia 2011, 4 (1), 187─194.
